Chemistry of the Upper and Lower Atmosphere: Theory, Experiments, and Applications

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Chemistry of the Upper and Lower Atmosphere

This Page Intentionally Left Blank

Chemistry of the Upper and Lower Atmosphere Theory, Experiments, and Applications

Barbara J. Finlayson-Pitts Department of Chemistry School of Physical Sciences University of California, Irvine Irvine, California

James N. Pitts, Jr. Department of Chemistry School of Physical Sciences University of California, Irvine Irvine, California

@ ACADEMIC PRESS A Harcourt Science and Technology Company

San Diego San Francisco New York Boston London Sydney Tokyo

Front cover photograph: Polar stratospheric clouds in the upper atmosphere seen from the NASA DC-8 aircraft at nearly 39,000 feet, north of Stavanger, Norway; the authors are grateful to NASA and O. B. Toon for providing this photograph. Photograph of breaking waves at "the Wedge," Newport Beach, California, and of multiple light beam reflections in aerosol chamber by J. N. Pitts, Jr. Back cover photograph: Image 9 1999 Photodisc, Inc. Dedication quote (facing page): All rights relating to the interest of Herman Hupfield in Canada and the reversionary territories are controlled by BIENSTOCK PUBLISHING COMPANY on behalf of REDWOOD MUSIC LTD. Used by permission. All rights reserved.

This book is printed on acid-free paper. @ Copyright 9 2000 by ACADEMIC PRESS All Rights Reserved. No part of this publication may be reproduced or transmitted in any form or by any means, electronic or mechanical, including photocopy, recording, or any information storage and retrieval system, without permission in writing from the publisher. Requests for permission to make copies of any part of the work should be mailed to: Permissions Department, Harcourt Inc., 6277 Sea Harbor Drive, Orlando, Florida, 32887-6777.

Academic Press A Harcourt Science and Technology Company 525 B Street, Suite 1900, San Diego, California 92101-4495, USA http://www.apnet.com

Academic Press 24-28 Oval Road, London NWl 7DX, UK http://www.hbuk, co.uk/ap / Library of Congress Catalog Card Number: 99-63218 International Standard Book Number: 0-12-257060-x PRINTED IN THE UNITED STATES OF AMERICA 99 00 01 02 03 04 MM 9 8 7 6

5

4

3

2

1

To our parents, Colin and Jean Finlayson and James and Esther Pitts, who, across gaps of a continent and time, imbued us with common interests in the environment, values reflected in this book, and with the philosophy, "The fundamental things apply

as time goes by." (From the song "As Time Goes By," Herman Hupfield, 1931.)


Contents

2

Preface xvii About the Authors xix Acknowledgments xxi

The Atmospheric System A. EMISSIONS

15

1. Oxides of Nitrogen

Overview of the Chemistry of Polluted and Remote Atmospheres

2. Volatile Organic Compounds (VOC)

A. REGIONS AND CHARACTERISTICS OF THE ATMOSPHERE 2

26

28 28

C. REMOVAL FROM THE ATMOSPHERE: WET AND DRY DEPOSITION 30 D. TYPICAL AMBIENT CONCENTRATIONS AND AIR QUALITY STANDARDS 33

a. Historical 8 b. Overview of Acidic Rain and Fogs

1. Units of Concentrations and Conversions

9

D. CHEMISTRY OF THE STRATOSPHERE E. GLOBAL CLIMATE CHANGE

10

2. Criteria and Noncriteria Pollutants and Air Quality Standards 35

11

13

G. DISCUSSION TOPIC AND OZIPR MODEL 1. Discussion Topic: "Background Ozone"

33

a. Parts per Million, Parts per Hundred Million, Parts per Billion, and Parts per Trillion 33 b. Number per Cubic Centimeter 34 c. Micrograms per Cubic Meter 34

C. CHEMISTRY OF THE NATURAL TROPOSPHERE: REMOTE ATMOSPHERES 9

13

26

3. Temperature Inversions

8

REFERENCES

25

2. Potential Temperature

a. Historical 4 b. Photochemical Air Pollution 5 c. Nighttime Chemistry of NO2 7

13

20

1. Lapse Rate: Temperature and Altitude

3. "Los Angeles" Smog: Ozone and Photochemical Oxidants 4

2. OZIPR Model

4. Sulfur Compounds

B. METEOROLOGY

2. "London" Smog: Sulfur Dioxide, Acidic Aerosols, and Soot 3

Some Points for Discussion

20

6. Lead

1. Historical Perspectives: Ancient and Medieval Times 3

F. INDOOR AIR POLLUTION

3. Carbon Monoxide

18

5. Total Suspended Particles (TSP), PM10, and PM2.5 21

B. AIR POLLUTION AND THE CHEMISTRY OF OUR TROPOSPHERE 3

4. Acid Deposition

17

E. EFFECTS ON VISIBILITY AND MATERIALS

13

13

F. ECONOMICS

13

37

38

G. ATMOSPHERIC CHEMISTRY: RISK ASSESSMENTS AND PUBLIC POLICIES FOR AIR POLLUTION CONTROL 38

vii

viii

CONTENTS

H. PROBLEMS

4

39

REFERENCES

Photochemistry of Important Atmospheric Species

39

3

A. MOLECULAR OXYGEN

Spectroscopy and Photochemistry: Fundamentals A. BASIC PRINCIPLES

1. Absorption Spectra 2. Photochemistry B. OZONE

43

1. Molecular Energy Levels and Absorption and Emission Spectroscopy 43 a. Diatomic Molecules 43 b. Polyatomic Molecules 49

86

89

90

1. Absorption Spectra 2. Photochemistry

90

91

C. NITROGEN DIOXIDE

2. Fates of Electronically Excited Molecules

1. Absorption Spectra

50

a. Photophysical Processes 50 b. Photochemical Processes 51 c. Quantum Yields 51

1. Basic Relationships

52

95

95

2. Photochemistry

96

D. NITRIC ACID

98

E. NITROUS ACID B. ABSORPTION OF LIGHT

86

99

F. PEROXYNITRIC ACID

100

52

2. The Beer-Lambert Law

G. NITRATE RADICAL

53

100

H. DINITROGEN PENTOXIDE C. ATMOSPHERIC PHOTOCHEMISTRY

55

I. NITROUS OXIDE

1. Solar Radiation and Its Transmission through the Atmosphere 55 a. The Sun and Its Relationship to the Earth: Some Important Definitions for Atmospheric Chemistry b. Solar Spectral Distribution and Intensity in the Troposphere 57

2. Calculating Photolysis Rates in the Atmosphere

101

J. ORGANIC NITRATES AND PEROXYACETYL NITRATE 102 55

1. Organic Nitrates

102

2. Peroxyacetyl Nitrate 61

a. Photolysis Rate Constant/kp ], Radiance [L(A)], Actinic Flux [F(A)], and Irradiance [E(A)] 61 b. Estimates of the Actinic Flux, F(A), at the Earth's Surface 64 c. Effects of Latitude, Season, and Time of Day on F(a) 65 d. Effect of Surface Elevation on F(a) 65 e. Effect of Height above Earth's Surface on F(a) 67 f. Sensitivity of Calculated Actinic Fluxes to Input Values for Surface Albedo and Ozone and Particle Concentrations 67 g. Effects of Clouds on F(a) 72 h. Comparison of Calculated Actinic Fluxes to Experimentally Measured Values 75 i. Actinic Fluxes in the Stratosphere 76

3. Procedure for Calculating Photolysis Rates

76

4. Example: Photolysis of Acetaldehyde at the Earth's Surface 81

103

K. SULFUR DIOXIDE AND SULFUR TRIOXIDE

REFERENCES

2. SO 3

105

L. HYDROGEN PEROXIDE AND ORGANIC HYDROPEROXIDES 107 M. ALDEHYDES AND KETONES

107

N. CHLORINE NITRATE (C1ONO 2) AND BROMINE NITRATE (BrONO 2) 111 O. HC1 AND HBr

113

P. THE HALOGENS Q. ClO, BrO, AND IO R. C1OOC1 S. OC10

114 114

114 115 115

83 84

103

1. SO 2 103

T. HOC1, HOBr, AND HOI

D. PROBLEMS

101

U. NITROSYL CHLORIDE (C1NO) AND NITRYL CHLORIDE (C1NO 2) 117

CONTENTS

V. HALOGENATED METHANES AND ETHANES 117 W. PROBLEMS

E. LABORATORY TECHNIQUES FOR STUDYING HETEROGENEOUS REACTIONS 156 1. Analysis of Systems with Gas- and Liquid-Phase Diffusion, Mass Accommodation, and Reactions in the Liquid Phase or at the Interface 158

117

REFERENCES

ix

126

2. Knudsen Cells

5

3. Flow Tube Studies

A. FUNDAMENTAL PRINCIPLES OF GAS-PHASE KINETICS 130

131

2. Termolecular Reactions and Pressure Dependence of Rate Constants 133 3. Temperature Dependence of Rate Constants

138

a. Arrhenius Expression 138 b. Predictions of Collision Theory and Transition State Theory 139 c. Example of Importance of Temperature Dependence: PAN Decomposition 141

B. LABORATORY TECHNIQUES FOR DETERMINING ABSOLUTE RATE CONSTANTS FOR GAS-PHASE REACTIONS 141

142 145

168

7. Liquid Jet Apparatus

169

171

10. Other Methods

171

172

F. COMPILATIONS OF KINETIC DATA FOR ATMOSPHERIC REACTIONS 172 G. PROBLEMS

174

REFERENCES

175

6 Rates and Mechanisms of Gas-Phase Reactions in Irradiated Organic- N O x - Air Mixtures

1. OH

146

5. Cavity Ring Down Method 6. Static Techniques

6. Aerosol Chambers

A. SOURCES OF OXIDANTS IN THE TROPOSPHERE: OH, 0 3, NO 3, HO z, AND C1 179

142

3. Flash Photolysis Systems 4. Pulse Radiolysis

168

167

9. Surface Science Techniques

a. Elementary vs Overall Reactions 130 b. Rate Laws, Reaction Order, and the Rate Constant c. Half-Lives and Lifetimes 132

2. Fast-Flow Systems

5. Bubble Apparatus

8. DRIFTS

130

1. Kinetic Analysis

167

4. Falling-Droplet Apparatus

Kinetics and Atmospheric Chemistry

1. Definitions

165

2. 0 3 147

3. NO 3

148

180 180

4. HO 2 180

C. LABORATORY TECHNIQUES FOR DETERMINING RELATIVE RATE CONSTANTS FOR GAS-PHASE REACTIONS 149 D. REACTIONS IN SOLUTION

179

5. C1

180

B. LIFETIMES OF TYPICAL ORGANICS IN THE TROPOSPHERE 181

151 C. REACTIONS OF ALKANES

1. Interactions of Gaseous Air Pollutants with Atmospheric Aqueous Solutions 151

1. Hydroxyl Radical (OH)

2. Diffusion-Controlled Reactions of Uncharged Nonpolar Species in Solution 152 3. Reactions of Charged Species in Solution

153

4. Experimental Techniques Used for Studying Solution Reactions 155

2. Nitrate Radical (NO 3)

182

182 184

3. Chlorine Atoms (C1) 184 D. REACTIONS OF ALKYL (R), ALKYLPEROXY (RO 2), AND ALKOXY (RO) RADICALS IN AIR 185

X

CONTENTS

1. Alkyl Radicals (R)

185

4. Arctic

2. Alkylperoxy Radicals (RO 2) a. b. c. d. e.

185

Reactions with N O 185 Reactions with HOe and ROe 186 Reactions with N O 3 187 Reactions with N O e 187 Fate of RO e under Typical Tropospheric Conditions

3. Alkoxy Radicals (RO)

K. ATMOSPHERIC CHEMISTRY AND BIOMASS BURNING 244 L. PROBLEMS 188

1. Hydroxyl Radical (OH)

191

196


7

5. Nitrogen Dioxide (NO 2)


3. Reactions of NO and NO 2 with Water and Alcohols 268

206

206

a. Uptake into and Reaction with Liquid Water 268 b. "Heterogeneous" Reaction of NO2 with Water Vapor c. Reaction with Alcohols 272

a. Organics 272 b. Sea Salt Particles and Mineral Oxides

212

3. Chlorine Atoms (C1) 212

272

C. ATMOSPHERIC CHEMISTRY OF HONO

H. REACTIONS OF OXYGEN-CONTAINING ORGANICS 213 1. Reactions of OH, NO 3, and C1 2. Hydroperoxyl Radical (HO 2)

213

1. Formation of HONO

273

273

2. Atmospheric Fates of HONO

274

D. REACTIONS OF NO 3 AND N 2 0 5 276

216

1. Reactions of NO 3 276

I. REACTIONS OF NITROGENOUS ORGANICS 217 1. Peroxyacetyl Nitrate and Its Homologs

217

220

a. b. c. d.

Reactions with N O e and N O 276 Reactions with Organics 276 Thermal Decomposition 276 Reaction with Water 277

2. Reactions o f N 2 0 5 279

a. Alkyl Nitrates 220 b. Alkyl Nitrites 221

3. Amines, Nitrosamines, and Hydrazines

221

a. Amines 221 b. Nitrosamines 223 c. Hydrazines 223

a. Hydrolysis 279 b. Other Reactions 280

E. ATMOSPHERIC CHEMISTRY OF HNO 3 281 1. Formation

J. CHEMISTRY OF REMOTE REGIONS 1. Emissions of Biogenic Organics

225

281

2. Tropospheric Fates F. "MISSING" NOy

225

231

281 286

G. AMMONIA (NH 3)

a. Biogenic Hydrocarbons 231 b. General Remote Boundary Layer Chemistry

239

269

4. Other Reactions of NO 2 272

207

2. Alkyl Nitrates and Nitrites

266

2. Nighttime Reactions to Form NO 3 and N 2 0 5 267 206

G. REACTIONS OF SIMPLE AROMATIC HYDROCARBONS 207 1. Hydroxyl Radical (OH)

A. OXIDATION OF NO TO NO 2 AND THE LEIGHTON RELATIONSHIP 265

1. Daytime Gas-Phase Reaction with OH

F. REACTIONS OF ALKYNES 1. Hydroxyl Radical (OH)

Chemistry of Inorganic Nitrogen Compounds

B. OXIDATION OF NO 2 266 201

4. Chlorine Atoms (C1) 205

3. Upper Troposphere

248

188

E. REACTIONS OF ALKENES (INCLUDING BIOGENICS) 191

2. Chemistry

247

REFERENCES

4. Summary of R, RO 2, and RO Radical Reactions in the Troposphere 191

2. Ozone ( 0 3)

241

234

H. PROBLEMS

287

REFERENCES

288

286

CONTENTS

8

F. PROBLEMS

Acid Deposition: Formation and Fates of Inorganic and Organic Acids in the Troposphere A. CONTRIBUTION OF H2SO4, HNO 3, HONO, AND ORGANIC ACIDS 294 B. SOLUBILITY OF GASES IN RAIN, FOGS, AND CLOUDS: HENRY'S LAW AND AQUEOUS-PHASE EQUILIBRIA 295 C. OXIDATION OF SO 2 296 1. Field Studies

a. b. c. d.

3. Oxidation in the Aqueous Phase

301

Particles in the Troposphere A. PHYSICAL PROPERTIES 1. Some Definitions

349

2. Size Distributions

351

a. Number, Mass, Surface, and Volume Distributions 351 b. Atmospheric Aerosols and Log-Normal Distributions 358

2. Reactions of Gases at Particle Surfaces 3. Reactions in the Aqueous Phase

1. Reactions of Dimethyl Sulfide (CH3SCH 3) Reaction with OH 329 Reaction with the Nitrate Radical (NO 3) 332 Oxidation by Chlorine Atoms 332 Oxidation by Halogen Oxides: IO, BrO, and ClO Oxidation by Ozone in the Aqueous Phase 334

334

375

379

380

4. Relative Importance of Various Aerosol Growth Mechanisms 380 C. CHEMICAL COMPOSITION OF TROPOSPHERIC AEROSOLS 38O

E. OXIDATION OF SULFUR COMPOUNDS OTHER THAN SO 2 328

334

362

1. Nucleation, Condensation, and Coagulation

326

2. Dimethyl Disulfide (CH3SSCH 3)

349

B. REACTIONS INVOLVED IN PARTICLE FORMATION AND GROWTH 375

324

3. Methyl Mercaptan (CH3SH)

9

a. Light Scattering and Absorption 365 b. Relationship of Light Scattering and Absorption to Visibility Reduction 368

301

5. Relative Importance of Various Oxidation Pathways for SO 2 325

a. b. c. d. e.

337

4. Light Scattering and Absorption and Their Relationship to Visibility Reduction 365

a. S(IV) Aqueous Equilibria 301 b. Physical and Chemical Steps in Aqueous-Phase Oxidation 306 c. Oxidation by 0 e (Catalyzed and Uncatalyzed) 308 d. Oxidation by 03 311 e. Oxidation by He 02 and Organic Peroxides 313 f . Oxidation by Oxides of Nitrogen 314 g. Free Radical Reactions in Clouds and Fogs 315 h. Effect of Droplet Size on S(IV) Oxidation 322 i. Fog- Smog- Fog Cycles 323

D. ORGANIC ACIDS

REFERENCES

a. Gravitational Settling 362 b. Brownian Diffusion 363

298

Hydroxyl Radical 298 Criegee Biradical 299 Other Gas-Phase Reactants 300 Computing Oxidation Rates in the Atmosphere

4. Oxidation on Surfaces

336

3. Particle Motion

296

2. Oxidation in the Gas Phase

xi

329

333

1. Inorganic Species a. b. c. d. e.

381

Size Distribution 381 Trace Elements as Tracers 386 Source Apportionment Models 386 Forms of Inorganics in Tropospheric Particles 388 Unique Sources of Particles 391

2. Organics

393

a. Biogenically Derived Organics in Aerosol Particles 393 b. Organics in Anthropogenically Influenced and Aged Aerosol Particles 396

D. G A S - PARTICLE DISTRIBUTION OF SEMIVOLATILE ORGANICS 412 1. Adsorption on Solid Particles 2. Absorption into Liquids

413

417

4. Hydrogen Sulfide (H2S)

335

3. Octanol- Air Partitioning Coefficients

5. Carbon Disulfide (CS2)

335

E. PROBLEMS

6. Carbonyl Sulfide (COS)

336

423

REFERENCES

423

420

xii

CONTENTS

10

3. The Salmonella TM677 "Forward Mutation" Assay 483

Airborne Polycyclic Aromatic Hydrocarbons and Their Derivatives: Atmospheric Chemistry and Toxicological Implications

4. Human Cell Mutagenicities of PAHs and PACs

D. BACTERIAL AND HUMAN CELL MUTAGENICITIES OF POLLUTED AMBIENT AIR 486

A. NOMENCLATURE AND SELECTED PHYSICAL AND SPECTROSCOPIC PROPERTIES OF POLYCYCLIC AROMATIC HYDROCARBONS (PAHs) AND POLYCYCLIC AROMATIC COMPOUNDS (PACs) 440 1. Combustion-Generated PAHs and PACs

440

2. Structures and IUPAC Rules for Nomenclature a. PAHs b. PACs

440

440

444

3. Solubilities and Vapor Pressures 451 4. G a s - Particle Partitioning, Sampling Techniques, and Ambient Levels of Selected PAHs and PACs 453 a. PAHs b. PACs

484

454 461

5. Absorption and Emission Spectra of Selected PAHs and PACs 461 B. BIOLOGICAL PROPERTIES OF PAHs AND PACs. I: CARCINOGENICITY 466

1. Bacterial Mutagenicity of Urban Air: A Worldwide Phenomenon 486 a. Background 486 b. Particle Size Distribution of PAHs and Mutagenicity c. Variables Influencing Mutagenicity Levels 488

487

2. Sources, Ambient Levels, Transport, and Transformation: Some Case Studies 491 a. Bacterial Mutagenicity 491 b. Human Cell Mutagenicity 497

3. Bioassay-Directed Chemical Analysis for PAHs and PACs in Fine Ambient Aerosols Using a Human Cell Assay 497 4. Bioassay-Directed Chemical Analysis for Vapor-Phase and Particle-Phase PAHs and PACs in Ambient Air Using Bacterial Assays 502 E. ATMOSPHERIC FATES OF PARTICLE-ASSOCIATED PAHs: HETEROGENEOUS REACTIONS 1. Background

504

504

1. Historical Perspective: Benzo[a]pyrene, the "Classic Chemical Carcinogen" 466

2. Theoretical and Experimental Structure- Reactivity Relationships 505

2. Carcinogenicity of PAHs, Cancer Potencies, and Potency Equivalence Factors 467

3. Field Studies of Atmospheric Reactions: Transport and Transformation 507

3. Carcinogenicity of Nitroarenes and Other Nitro-PACs 473 C. BIOLOGICAL PROPERTIES OF PAHs AND PACs.

II: MUTAGENICITY

475

1. Short-Term Tests for Genetic and Related Effects 475 2. The Ames Salmonella typhimurium Reversion Assay 475 a. Principle of Method: Direct vs Activatable Bacterial Mutagens 475 b. Assay Procedure 478 c. The "Microsuspension Modification" 478 d. The Salmonella typhimurium Reversion Assay for Gas-Phase Mutagens 479 e. Accuracy and Precision 480 f. Some Mutagenic Potencies in the Salmonella typhimurium Assay 480 g. Microbiology and Analytical Chemistry: Bioassay-Directed Fractionation and Chemical Analysis of Complex Mixtures of POM 482

a. Decay of Particle-Associated PAHs in Ambient Air 507 b. Time- Concentration Profiles of Particle-Phase PAHs during Transport across an Air Basin 508

4. Photochemical Reactions of Particle-Associated PAHs 510 a. Photooxidation in Solution and "Liquid-like" Surfaces of Organic Aerosols 510 b. Photooxidations on Inorganic Solid-Air Surfaces 512

5. Gas-Particle Reactions

513

a. Ozonolysis 513 b. Nitration 515

6. Atmospheric Fates of Particle-Associated Nitroarenes 518 a. Gas - Particle Reactions b. Photochemical Reactions

518 518

F. REACTIONS OF GAS-PHASE PAHs: ATMOSPHERIC FORMATION OF MUTAGENIC NITRoARENES 519

xiii

CONTENTS

1. Combustion-Generated Primary Emissions of Nitroarenes 519

b. Size 613 c. Typical Particle Concentrations in the Atmosphere

3. Measurement of Chemical Composition

2. Atmospheric Formation of Nitro-PAHs and Nitro-PACs 520 REFERENCES

a. b. c. d. e.

527

11 Analytical Methods and Typical Atmospheric Concentrations for Gases and Particles A. GASES

548

a. b. c. d.

Chemiluminescence 548 Fluorescence 548 Infrared Spectroscopy (IR) 549 DOAS ( U V - Visible Absorption Spectroscopy)

2. Mass Spectrometry a. b. c. d.

548

Sample Introduction Ionization 562 Mass Filters 564 Detectors 566

556

561

a. Single-Particle Laser Ionization Techniques 627 b. Alternate Potential Mass Spectrometric Methods for Sizing and Chemical Composition 630 c. Depth Profiling of Particle Composition 631

a. b. c. d. e. f.

a. b. c. d.

REFERENCES

Filters 567 Denuders 567 Transition Flow Reactors (TFRs) 568 Mist Chambers and Scrubbers 568

B. PARTICLES a. b. c. d.

607

608

1. Sampling and Collection of Particles

608

Filters 608 Impactors 610 Electrostatic Precipitators 611 Sedimentation Collectors 611

2. Measurement of Physical Characteristics: Mass and Size 612 a. Mass

612

636

12

A. CHEMISTRY OF THE UNPERTURBED STRATOSPHERE 657

NO, NO2, NOx, and NOy 569 03 583 CO 583 SO 2 584 N M H C and VOC 585 Aldehydes, Ketones, Alcohols, and Carboxylic Acids 589 PAN, Other Peroxynitrates, and Alkyl Nitrates 594 H 2 0 2 and Organic Peroxides 595 HO~ Free Radicals 598

5. Generation of Standard Gas Mixtures

635

Homogeneous and Heterogeneous Chemistry in the Stratosphere

567

4. Methods for, and Tropospheric Levels of, Specific Gases 569 a. b. c. d. e. f. g. h. i.

632

Atomizers and Nebulizers 633 Vibrating-Orifice Generator 634 Spinning-Disk Generator 634 Dry Powder Dispersion 634 Tube Furnaces 634 Condensation 635

3. Filters, Denuders, Transition Flow Reactors, Mist

Chambers, and Scrubbers

623

4. Real-Time Monitoring Techniques for Particles

C. PROBLEMS

561

619

Inorganic Elements 619 Inorganic Ions 622 Total Carbon: Organic versus Graphitic (Elemental) Speciation of Organics 625 Artifacts 626

5. Generation of Calibration Aerosols

1. Optical Spectroscopic Techniques

618

1. Stratosphere-Troposphere Exchange (STE) 2. Chapman Cycle and NO x Chemistry

658

660

B. HIGH-SPEED CIVIL TRANSPORT (HSCT), ROCKETS, AND THE SPACE SHUTTLE 662 1. HSCT

662

2. Space Shuttle and Solid Rocket Motors C. CHLOROFLUOROCARBONS 1. Types, Nomenclature, and Uses

667

669 669

2. Lifetimes and Atmospheric Fates of CFCs and Halons 670 3. Gas-Phase Chemistry in the Stratosphere 4. Antarctic "Ozone Hole"

673

675

5. Polar Stratospheric Clouds (PSCs) and Aerosols 680 a. Nature of Aerosols and PSCs 680 b. Uptake of HCl into PSCs 686 c. Heterogeneous Chemistry on PSCs and Aerosols

6. Effects of Volcanic Eruptions

690

688

626

xiv

CONTENTS

7. Ozone Depletion in the Arctic

696

8. Ozone Destruction in the Midlatitudes

1. Global Absorption and Emission of Radiation 700

2. Radiative Transfer Processes in the Atmosphere a. Macroscopic View 766 b. Molecular-Level View 767 c. From Molecules to the Global Atmosphere

D. CONTRIBUTION OF BROMINATED ORGANICS 701 1. Sources and Sinks of Brominated Organics 2. Bromine Chemistry in the Stratosphere

701

707

G. PROBLEMS

707

REFERENCES

768

B. CONTRIBUTION OF TRACE GASES TO THE GREENHOUSE EFFECT 770 1. Infrared Absorption by Trace Gases 2. Trends in Trace Gas Concentrations

708

a. b. c. d. e. f.

13 Scientific Basis for Control of Halogenated Organics A. INTERNATIONAL AGREEMENTS ON PHASEOUT OF HALOGENATED ORGANICS 727

CO 2 773 CH e 777 N 2 0 779 0 3 780 CFCs, HCFCs, and HFCs Other Gases 783

770 773

783

3. Radiative Forcing by Greenhouse Gases and Global Warming Potentials 783

B. OZONE DEPLETION POTENTIALS (ODP)

730

a. Instantaneous and Adjusted Radiative Forcing 783 b. Absolute and Relative Global Warming Potentials 784

C. TRENDS IN CFCs, THEIR REPLACEMENTS, STRATOSPHERIC 0 3, AND SURFACE UV 733

C. AEROSOL PARTICLES, ATMOSPHERIC RADIATION, AND CLIMATE CHANGE

1. Trends in CFCs and Their Replacements

1. Direct Effects

733

2. Trends in Stratospheric 0 3 736 3. Trends in Surface Ultraviolet Radiation

741

D. TROPOSPHERIC CHEMISTRY OF ALTERNATE CFCs 744 1. Kinetics of OH Reactions 2. Tropospheric Chemistry a. b. c. d. e.

744 746

Chemistry of HFC-134a (CH2FCF 3) 746 Chemistry HCFC-125 (CHF2CF 3) 748 Chemistry of HCFC-123 (CHCl2CF 3) 749 Chemistry of HCFC-141b (CH3CFCl 2) 750 Tropospheric Fates of Halogenated Products of HCFC Oxidation 750

a. Scattering of Solar Radiation 789 b. Absorption of Solar Radiation 796 c. Absorption of Long-Wavelength Infrared

2. Indirect Effects of Aerosol Particles

1. Absorption of Solar Radiation by Clouds

753

4. Volcanic Eruptions 5. Oceans

A. RADIATION BALANCE OF THE ATMOSPHERE: THE GREENHOUSE EFFECT 763

814

2. Feedbacks: Water Vapor, Clouds, and the "Supergreenhouse Effect" 819

F. PROBLEMS

Global Tropospheric Chemistry and Climate Change

799

D. SOME OTHER FACTORS AFFECTING GLOBAL CLIMATE 814

3. Solar Variability

14

798

a. Clouds 799 b. Heterogeneous Chemistry Involving Climate Species 814

753

753

788

789

E. SUMMARY

REFERENCES

766

3. Dependence of Net Infrared Absorption on Atmospheric Concentrations 769

702

E. CONTRIBUTION OF IODINE-CONTAINING ORGANICS 706 F. SUMMARY

763

821 822

822

E. OBSERVATIONS OF CLIMATE CHANGES 1. Observed Temperature Trends

823

a. Trends over the Past Century 823 b. Temperatures and Other Proxies for Climate Change over the Past ~ 105 Years 825

2. Other Climate Changes

828

823

CONTENTS

F. THE FUTURE G. PROBLEMS

V O C - NO x - Air Mixtures 878 e. Advantages and Limitations of Environmental Chambers

828 829

REFERENCES

3. Models

Indoor Air Pollution: Sources, Levels, Chemistry, and Fates 844

B. OXIDES OF NITROGEN

846

1. Levels of NO x 846

a. b. c. d. e.

886

Simple Models 886 Mathematical Models 887 Simple Mathematical Models 892 Grid Models: Urban to Regional Scales 893 Models Incorporating Particles 907

B. REACTIVITY OF VOC

907

1. Typical Reactivity Scales

907

2. Application to Control of Mobile Source Emissions 909

2. H O N O and HNO 3 847 C. CO AND SO 2 849 D. VOLATILE ORGANIC COMPOUNDS E. OZONE

850

859

F. INDOOR VOC - NO x - 0 3 CHEMISTRY

859

C. FIELD OBSERVATIONS OF VOC, NO x , AND 0 3 913 D. ALTERNATE FUELS

918

1. Reformulated Gasolines

918

2. Compressed Natural Gas (CNG)

919

G. PARTICLES

861

3. Liquefied Petroleum Gas (LPG)

H. PROBLEMS

865

4. Alcohol Fuels and Blends with Gasoline

REFERENCES

5. Hydrogen

865

16 Applications of Atmospheric Chemistry: Air Pollution Control Strategies and Risk Assessments for Tropospheric Ozone and Associated Photochemical Oxidants, Acids, Particles, and Hazardous Air Pollutants A. TROPOSPHERIC OZONE AND ASSOCIATED PHOTOCHEMICAL OXIDANTS 871 1. Environmental Chambers

872

Types of Chambers 872 Preparation of Reactants, Including "Clean Air" 876 Light Sources 876 Typical Time- Concentration Profiles of Irradiated

920 920

921

6. Electric Vehicles

a. b. c. d.

880

2. Isopleths for Ozone and Other Photochemically Derived Species 882

829

15

A. RADON

XV

921

E. CONTROL OF ACIDS

921

F. CONTROL OF PARTICLES

923

G. ATMOSPHERIC CHEMISTRY AND RISK ASSESSMENTS OF HAZARDOUS AIR POLLUTANTS 925 H. PROBLEMS

930

REFERENCES

932

Appendix I: Enthalpies of Formation of Some Gaseous Molecules, Atoms, and Free Radicals at 298 K 943 Appendix II: Bond Dissociation Energies 945 Appendix III: Running the OZIPR Model 947 Appendix IV: Some Relevant Web Sites 949 Appendix V: Pressures and Temperatures for Standard Atmosphere 951 Appendix VI: Answers to Selected Problems 952 Subject Index 957


Preface

What is written without effort is in general read without pleasure. Samuel Johnson ( 1709 - 1784) In William Seward Biographia

Given this admonition, written long ago by English poet and critic Samuel Johnson, we trust that this book may be read with some pleasure by a "spectrum" of readers. The field of atmospheric chemistry has undergone dramatic changes since our first book on this subject was published in 1986. Since then, a number of new, exciting, and highly relevant research areas have emerged. We treat these here, along with the fundamentals of spectroscopy, photochemistry, and reaction kinetics and mechanisms of atmospheric systems. For example, the discovery of the Antarctic ozone hole has left no doubt that chlorofluorocarbons (CFCs) have led to depletion of stratospheric ozone and has highlighted the importance of heterogeneous chemistry on polar stratospheric clouds. Atmospheric measurements of the CFCs have documented changes in their global concentrations in response to control measures. The importance of emissions and chemistry for climate changes on a global scale has become an area of intense popular interest and scientific research. Furthermore, the formation, chemistry, and fates of airborne particulate matter, particularly that less than 2.5 /xm in diameter (i.e., PM2.5), are now recognized not only as important to global climate issues but also as of concern for their toxicological effects. The similarity in much of the chemistry of indoor air pollutants and those outdoors is now evident, and the importance of understanding the chemistry of airborne toxic chemicals (also known as "hazardous air pollutants," HAPs) for the development of sound risk assessments is clear. A great deal more is known about the role of emissions and atmospheric reactions of polycyclic aromatic hydrocarbons in the mutagenicity and carcinogenicity of urban atmo-

spheres. In addition to these major developments in the field, there has been a significant increase in our understanding of the gas-phase atmospheric chemistry and photochemistry of organics, oxides of nitrogen, and SO 2 and rapidly increasing evidence for the importance of a wide range of heterogeneous reactions in the troposphere. We have attempted in this new book to present the current understanding of the chemistry of the natural and polluted upper and lower atmosphere in such a way that it will be useful to a range of atmospheric chemists as well as to atmospheric scientists and engineers working in this field. However, the fundamentals (e.g., theories, rates, and mechanisms of homogeneous and heterogeneous reactions and of spectroscopic and photochemical processes) are also emphasized. We believe that this approach is useful in providing the necessary background and tools for graduate students as well as for scientists and engineers in related fields who wish to enter this exciting and dynamic area. Problems are provided at the end of most chapters (and answers to selected problems are in Appendix VI) to enhance the book's use in teaching. The literature is assessed through the end of 1998 and, in some cases, into 1999. We have cited only papers in the peer-reviewed literature or, in a few instances, government agency reports that are readily available. Our approach has been to consider primarily typical examples of major papers in the refereed literature in the relevant areas. Because of the enormous breadth of the field today, we have not been able to reference all papers in all relevant areas, which will unavoidably lead to some omissions. We apologize in

xvii

xviii

PREFACE

advance to our colleagues whose work might not have been cited in the relevant area. We are deeply indebted to our many colleagues in the field whose outstanding work, generous sharing of results, and helpful discussions have made this work possible. It is our hope that this book does justice to

the current state of this exciting, rapidly maturing, and scientifically and societally relevant discipline. Barbara J. Finlayson-Pitts James N. Pitts, Jr. Fawnskin, California August 19, 1999

About the Authors

Barbara J. Finlayson-Pitts is Professor of Chemistry at the University of California, Irvine. Her research program focuses on laboratory studies of the kinetics and mechanisms of reactions in the atmosphere, especially those involving gases with liquids or solids of relevance in the troposphere. Reactions of sea salt particles to produce photochemically active halogen compounds and the subsequent fates of halogen atoms in the troposphere are particular areas of interest, as are reactions of oxides of nitrogen at aqueous and solid interfaces. Her research is currently supported by the National Science Foundation, the Department of Energy, the California Air Resources Board, the Dreyfus Foundation, and NATO. She has authored or coauthored more than 80 publications in this area, as well as a previous book, Atmospheric Chemistry: Fundamen-

cal Society, the American Geophysical Union, the American Association for the Advancement of Science, and Iota Sigma Pi. She is the recipient of a number of awards, including the Governor General's Medal at Trent University, a Woodrow Wilson Fellowship, a National Research Council of Canada Science Scholarship, Golden Key National Honor Society, and a Japan Society for the Promotion of Science Fellowship. She has been elected a Fellow of the American Association for the Advancement of Science and has several awards for undergraduate teaching.

James N. Pitts, Jr., is a Research Chemist at the University of California, Irvine, and Professor Emeritus from the University of California, Riverside. He was Professor of Chemistry (1954-1988) and cofounder (1961) and Director of the Statewide Air Pollution Research Center (1970-1988) at the University of California, Riverside. His research has focused on the spectroscopy, kinetics, mechanisms, and photochemistry of species involved in a variety of homogeneous and heterogeneous atmospheric reactions, including those associated with the formation and fate of mutagenic and carcinogenic polycyclic aromatic compounds. He is the author or coauthor of more than 300 research publications and three books: Atmospheric

tals and Experimental Techniques. At UCI, she teaches graduate-level courses in atmospheric chemistry on a regular basis. In addition, she teaches such classes as undergraduate instrumental analysis, in which she is developing a new laboratory curriculum centered around the analysis of complex environmental mixtures. This work has been supported by the Dreyfus Foundation and UCI. She received her undergraduate B.Sc. (Hons) in 1970 from Trent University in Peterborough, Ontario, Canada, where her interest in atmospheric chemistry was first sparked by discussions of the spectroscopy of auroras in a physical chemistry class taught by Professor R. E. March. She obtained her master's degree (1971) and Ph.D. (1973) from the University of California, Riverside. After a year's postdoctoral work at UCR, she joined the faculty of California State University, Fullerton, where she taught and carried out research in physical and atmospheric chemistry. In 1994, she joined the faculty at UCI. Dr. Finlayson-Pitts is a member of a number of professional societies, including the American Chemi-

Chemistry: Fundamentals and Experimental Techniques, Graduate School in the Sciences--Entrance, Survival and Careers, and Photochemistry. He has been coeditor of two series, Advances in Environmental Science and Technology and Advances in Photochemistry. He served on a number of panels in California, the United States, and internationally. These included several National Academy of Science panels and service as Chair of the State of California's Scientific Review Panel for Toxic Air Contaminants and as a member of the Scientific Advisory Committee on Acid Deposition.

xix

XX

ABOUT THE AUTHORS

He received his B.S. (1945) and Ph.D. (1949) from the University of California, Los Angeles; his research advisor was Professor Francis E. Blacet, who first identified the photolysis of NO 2 as the anthropogenic source of ozone in photochemical air pollution. From 1942 to 1945, he participated in laboratory and field studies in chemical warfare. He was on the faculty at Northwestern University from 1949 to 1954, leaving to join the faculty at the new University of California, Riverside, campus. He was a Guggenheim Fellow at University College, Oxford, in 1961 and a Research Fellow of Merton College, Oxford, in 1965. Dr. Pitts is a member of a number of professional societies, including the American Chemical Society, the American Geophysical Union, the American Association for the Advancement of Science, and the Ameri-

can Physical Society. He has received a number of awards for his research, including the Clean Air Award of the California Lung Association (1979), the Frank A. Chambers Award for "Outstanding Achievement in the Science and Art of Air Pollution Control" from the Air Pollution Control Association (1982), the Richard C. Tolman Medal (1983), the UCR Faculty Research Lectureship (1965), the F. J. Zimmerman Award in Environmental Science, and the Clean Air Award (1992) from the South Coast Air Management District. He is an elected Fellow of the American Association for the Advancement of Science. He has also received numerous commendations from local, state, and federal legislators for his application of fundamental atmospheric chemistry to air pollution problems.

Acknowledgments

Many individuals and organizations contributed to making this book a reality. First, our assistant, Mae Minnich, devoted innumerable long hours to the manuscript; her organization and professional talents were indispensable. Without her outstanding skills in word processing and manuscript preparation, combined with her patience, wonderful sense of humor, and tireless enthusiasm, this undertaking would not have been possible. Kelly Donovan took almost indecipherable drafts of figures and turned them into clear and attractive drawings, some of which reflect major creativity on her part. The cheerful dedication, professional skills, and willingness to "go the extra mile" on the part of both of these individuals helped the authors through some long days. The assistance of a number of undergraduate and graduate students, postdoctoral fellows, and colleagues was critical in this effort. Christopher Elliott and Ellen Fleyshman assembled and organized the thousands of references and, along with Lisa Wingen, Alisa Ezell, and John Elliott, also provided computer and technical assistance. Their unflagging assistance in the midst of the final fray and the moral and concrete support of others, including Stacie Tibbets and Mike and Connie Ezell, reenergized us during the final "countdown." The assistance of Bill Barney, Mike and Alisa Ezell, Krishna Foster, Michael Gebel, Matt Lakin, Lisa Wingen, and Weihong Wong with figures and final page proofing was very helpful, as was the able library assistance of Jean Miller. Sasha Madronich generously not only reviewed the section on atmospheric radiation, but provided his unpublished calculations of actinic fluxes at different altitudes in a form useful to the atmospheric chemistry community for estimates of photolysis rates from the troposphere through the stratosphere. A number of colleagues reviewed chapters or portions of chapters, and their insightful comments and suggestions are

greatly appreciated. They include Janet Arey, Roger Atkinson, Thorsten Benter, Theo Brauers, Carl Berkowitz, Don Blake, Chris Doran, Anders Feilberg, Mario Molina, Ole John Nielsen, Torben Nielsen, Joyce Penner, F. Sherwood Rowland, David Rusch, Stephen Schwartz, Chet Spicer, Jochen Stutz, Darin Toohey, Douglas Worsnop, Y. L. Yung, and Paul Ziemann. We are indebted to many colleagues who provided figures, data, and stimulating discussions, especially Ed Baum, John Barker, Thorsten Benter, William Carter, Glenn Cass, Ralph Cicerone, Bart Croes, David Crosely, Donald Dabdub, Paul Davidovits, Leon Dolislager, A1isa Ezell, Michael Ezell, Jerome Fast, Jeff Gaffney, David Golden, Les Grant, Bill Harger, John Holmes, John Jayne, John Johnson, Jake Hales, John Hemminger, Wolfgang Junkermann, George Kirk, Charles Kolb, William Lockett, Alan Lloyd, Nancy Marley, Mike Nicovich, Randy Pasek, Shankar Prasad, Michael Prather, Ralph Propper, Scott Samuelsen, Rolf Sander, Jim Seiber, Brian Toon, Ernie Tuazon, Ken Schere, Charles Weschler, Hal Westberg, Arthur Winer, Paul Wine, Ed Yotter, Mark Zahniser, and Barbara Zielinska. Michael Gery, who developed the OZIPR model, graciously provided advice on its use as well as electronic copies of the documentation. This model, which contains the two major chemical mechanism schemes for gas-phase, VOC-NO X chemistry in use in atmospheric chemistry, is available on the Academic Press Web site (http://www.academicpress.com/pecs/download). A number of problems using this model are included in the book, and it is a valuable teaching tool for assessing the effects of various model input parameters on predicted concentrations of a wide variety of gas-phase species. His assistance and that of Marcia Dodge of the U.S. EPA in making it available are appreciated.

xxi

xxii

ACKNOWLEDGMENTS

The authors have been very fortunate over the years to have a number of accomplished and creative mentors who generously shared their knowledge of, and enthusiasm for, the fields of fundamental photochemistry, spectroscopy, and kinetics; these include F. E. Blacet, E. J. Bowen, P. A. Leighton, W. A. Noyes, Jr., E. W. R. Steacie, and R. E. March. We hope that the central importance of these fundamentals to understanding atmospheric chemistry is evident in this book. We also appreciate the tremendously talented and enthusiastic researchers, too numerous to mention here, who have spent time in our research groups over the years. Many of them have gone on to careers in atmospheric chemistry; it has given us a great deal of satisfaction and delight to see the "academic lineage" being passed on from our mentors to them. The authors have also very much appreciated the scientific interactions with, and financial support of, a number of public and private agencies and foundations, which have allowed us to conduct research in various areas in atmospheric chemistry. Over several decades, key individuals at these organizations have been helpful in many ways, and we are grateful to them. Agencies and individuals include Jarvis Moyers, Anne-Marie Schmoltner, and the late Richard Carrigan at the National Science Foundation; David Ballantine, George Stapleton, Ari Patrinos, Michael Riches, Michelle Broido, Rickey Petty, and Peter Lunn at the Department of Energy; John Holmes, Bart Croes, Ralph Propper, Jack Suder, Randy Pasek, and Eileen McCauley at the California Air Resources Board; Ron Patterson and Marcia Dodge at the Environmental Protection Agency; Brian Andreen and the late Hal Ramsey at The Research Corporation; and Robert Lichter at the Dreyfus Foundation. We have been impressed indeed by the professional skills, patience, thoughtful and imaginative ideas, and

dedication of the staff at Academic Press. Special thanks and appreciation go to Executive Editor David Packer, as well as his colleagues at AP, including Cheryl Uppling (Senior Production Editor), Mike Early (Book Production Manager), Liz Novelozo and Kim Schettig (Marketing), Linda Klinger (Editorial Coordinator), design artist Amy Stirnkorb, and David Phanco. On a more personal note, the support of our colleagues at the University of California, Irvine, particularly the Chairs in the Department of Chemistry, John Hemminger and Richard Chamberlin, has been essential to the completion of this project. The hospitality and encouragement of the Science Education Project staff, especially Ann and George Miller, Lynne Davanzo, and Frank Potter, are appreciated. We are also grateful to Edward J. Mclntyre for his professional skills and wise and effective counsel and to Bill and Maura Dickerson for their timely advice during an interesting period in the genesis of this book. Successful completion of this project was only possible through the professional and personal support of our friends and neighbors, particularly in University Hills at UCI as well as in our Fawnskin mountain "hamlet" (population 360 and elevation 6500 ft), especially Doris Layne and Linda Neuman. Regular visits to our home in Fawnskin by a variety of wildlife of the nonhuman type, including eagles, coyotes, bear, quail, and wild turkeys, provided much sanity and perspective during the writing of this book. Finally, we are very grateful to all of our colleagues, families, and friends, as well as to our golden retriever contingent, Babe, Maj, and BR, who have been extremely patient with our being "missing in action" for too many years while finishing this book. They (and we, at times) wondered if it would ever e n d . . , but in large part through all of their patience, support, and encouragement, it finally has!

C H A P T E R

1 Overview of the Chemistry of Polluted and Remote Atmospheres Atmospheric chemistry is an exciting, relatively new field. It encompasses the chemistry of the globe, from polluted to "clean," remote regions and from the region closest to the earth's surface, the troposphere (< 10-15 kin), through the tropopause (~ 10-15 kin) into the upper atmosphere--which, for the purposes of this book, we restrict to the stratosphere (~ 10-50 km). Chemical and physical processes occurring at the earth's surface--emissions, transport, lifetimes, and fates of certain anthropogenic and biogenic/geogenic chemic a l s - c a n impact the stratosphere--and vice versa. Thus, even though in some early studies the tropopause was perceived as being a "barrier" between the lower and upper atmospheres, it has become increasingly clear that the troposphere and stratosphere are intimately connected. Witness the vertical transport of long-lived ozone-destroying anthropogenic emissions of chlorofluorocarbons (CFCs) and conversely the downward transport of stratospheric ozone into the troposphere. Hence, even though we devote separate chapters to the chemistry of the troposphere and stratosphere, the emphasis throughout is on one integrated "system" of global atmospheric chemistry. In this regard, there are several topics that fall outside the scope of this book, including the evolution of the earth's atmosphere. For reviews, the reader is referred to articles by Kasting (1993) and All~gre and Schneider (1994). Although we point out throughout this book the interconnectedness of the lower and upper atmospheres, practicality and length preclude extension to the obvious interrelationships and feedbacks with other components of the earth system, including the controversial Gaia Hypothesis (named after the Greek goddess of the earth). The reader is referred to articles by Lovelock (1989), Kirchner (1989), Schneider (1990), and Lemon (1998) for discussions of the latter. Although atmospheric chemistry is sometimes viewed as an "applied" science, its foundations rest on funda-

mental research in diverse areas of chemistry. These include theoretical and experimental aspects of spectroscopy, photochemistry, and the kinetics and mechanisms of homogeneous and heterogeneous organic and inorganic reactions. We believe it is useful for today's students, researchers, and educators to be aware that major resources for Leighton's masterful 1961 treatment of the newly emerging field of atmospheric chemistry, Photochemistry of Air Pollution, were, in fact, pioneering, basic research monographs published years earlier. They include Photochemistry of Gases by W. A. Noyes, Jr., and P. A. Leighton (1939), two editions of Atomic and Free Radical Reactions by E. W. R. Steacie (1946, 1954), and G. Herzberg's classics Atomic Spectra and Atomic Structure (1944); Molecular Spectra and

Molecular Structure I: Spectra of Diatomic Molecules (1950); and Infrared and Raman Spectra of Polyatomic Molecules (1945). Throughout the body of this book, we address the basic chemistry driving key atmospheric processes in the natural and polluted troposphere and stratosphere and illustrate their critical interactions on local, regional, and global scales. In so doing, our treatment overall reflects the message that Sam sings so eloquently to Bogart and Bergman in the classic movie Casablanca,... "The fundamental things apply, as time goes by." In this chapter we provide an overview of the chemistry of the lower and upper atmospheres. In Chapter 2, we illustrate how this chemistry plays a critical role in the concept of an integrated "atmospheric chemistry system"--a loop that starts with emissions (anthropogenic and natural) and ultimately closes with scientific health and environmental risk assessments and associated risk management decisions for the control of air pollutants. Chapters 3-12 present a detailed examination and explanation of how one applies the theoretical and

2

1. CHEMISTRY OF POLLUTED AND REMOTE ATMOSPHERES

experimental fundamentals of photochemistry, spectroscopy, and kinetics and mechanisms (structure and reactivity) to the most important homogeneous and heterogeneous processes that take place in our natural and polluted atmosphere. We conclude by illustrating how our understanding of these chemical processes in our clean and polluted troposphere and stratosphere plays a crucial role in generating the "exposure" portions of scientific health risk assessments. Such assessments provide the foundation for sound, health-protective and cost-effective strategies for the control of tropospheric ozone, particles, acids, and a spectrum of "hazardous air pollutants" (including carcinogens and pesticides)--as well as for the mitigation of stratospheric ozone depletion.

A. R E G I O N S A N D CHARACTERISTICS OF THE ATMOSPHERE

temperatures as a function of altitude.) In the troposphere, the temperature generally falls with increasing altitude (except in the presence of inversions; see Chapter 2.B.3). This is due to the strong heating effect at the surface from the absorption of radiation. Because hot air rises, this causes strong vertical mixing so that species emitted at the earth's surface can rise to the tropopause, the region separating the troposphere from the stratosphere, in a few days or less, depending on the meteorological conditions. Essentially all of the water vapor, clouds, and precipitation are found in the troposphere, which provides an important mechanism for scavenging pollutants from the atmosphere. However, at the tropopause the temperature profile changes, increasing with altitude throughout the stratosphere. The reason for this increase is a critical series of photochemical reactions involving ozone and molecular oxygen. The "Chapman cycle," reactions (1)-(4), hypothesized in the 1930's by Sir Sydney Chapman, 20,

(1)

M O -Jr- O 2 ---) 0 3 ,

(2)

O + 0 3 --> 2 0 2 ,

(3)

0 2 -Jr-hu ~

Figure 1.1 shows the different regions of the atmosphere. (See also Appendix V for typical pressures and

Pressure (Torr) 10-3 100

10 "2

t

10 1

I I Thermosphere

/

90 I\

80

0 3 +hv+

1

10

102

I //

I

I

Mesopause

-

70 ~"

~

-

-

\

Mesosphere -,,~e~-O~

,,%

~ X S ~X5 .~

60 --

k\

v

-o
2NO

2.

about 85% of the NO is oxidized in ~ 15 s. However, at 100 ppb NO, approximately 226 days would be required to achieve the same net oxidation! As a result, the so-called thermal (i.e., nonphotochemical)oxidation of NO by reaction (9) is generally too slow to be of importance in the atmosphere. One exception to this generalization is the case where high concentrations of NO (e.g., several thousand ppm) may be emitted from sources such as uncontrolled power plants. In the initial seconds as the plume enters the atmosphere before it has had a chance to become completely diluted with the surrounding air, the NO may be sufficiently concentrated that the oxidation (9) by 0 2 is significant. For example, at 2000 ppm NO, 90% of the reactant would be oxidized to NO 2 within 30 min if this high concentration were to be maintained for that long. The plume integrity is generally not maintained for this period of time; however, under some meteorological conditions, the plume can be sufficiently stable that a significant fraction of the NO can undergo thermal oxidation by 0 2, and NO 2 can be directly formed many kilometers from the stack. In summary, it soon became evident that in ambient photochemical smog, the thermal oxidation of NO could not explain the relatively rapid conversion of NO to NO 2. With respect to the role of the organics, it was suggested about 1969-1970 that the hydroxyl radical drives the daytime chemistry of both polluted and clean atmospheres (Heicklen et al., 1969; Weinstock, 1969; Stedman et al., 1970; Levy, 1971). Thus, OH initiates chain reactions by attack on V O C or CO. These chains are then propagated through reactions such as those in Fig. 1.4. In this cycle, the organic is oxidized to a ketone, two molecules of NO are converted to NO2, and OH is regenerated. Of course, the ketone can then photodissociate into free radicals or itself be attacked by OH, and a similar cycle occurs, leading to further NO oxidation.

(9)

Indeed, in the laboratory, at Torr concentrations, the clear, colorless NO is oxidized in air virtually instantaneously to dark, red-brown NO 2. However, the rate of this reaction is second order in NO; that is, the speed of oxidation increases as the square of the NO concentration. Thus when one lowers the NO from high (Torr) concentrations to ambient part per trillion (ppt) or part per billion (ppb) levels (ppt = parts in 1012; ppb = parts in 109), the speed of the oxidation drops to the point where the rate is very small. For example, at 100 Torr NO ( ~ 1.3 • 105 ppm),

OH + RCH2R NO2~

RCR '~

oo

I

~ H20 + RCHR /

NO

N II

02

NO

HO2 l

R --

!~

NO 2

O I CH

-- R

02

FIGURE 1.4 Typical sequence of elementary reactions in which OH initiates the oxidation of an alkane in the troposphere.

7

B. AIR POLLUTIONAND THE CHEMISTRYOF OUR TROPOSPHERE The chain reactions are eventually terminated by such reactions as M O H + NO 2 ~ HNO3,

(10)

M, H20

HO 2 -+- HO 2

> H 2 0 2 -+-O2,

R O 2 + H O 2 ~ R O O H + O 2.

(11) (12)

A major source of OH in both clean and polluted air is the photodissociation of O 3 by actinic UV radiation in sunlight to produce an electronically excited oxygen atom, O( ~D), 0 3 + hu(A < 340 nm) ~ O(1D) + O 2 ,

(13)

followed by a very rapid reaction, in competition with deactivation, of the excited oxygen atom with water vapor, which is always present in the atmosphere:

The relative importance of these sources of OH and HO 2 radicals depends on the species present in the air mass, and hence on location and time of day. Figure 1.5, for example, shows the relative contributions as a function of time of day of three sources of O H / H O 2 in an urban air mass. In this case, nitrous acid is predicted to be the major OH source in the early morning hours, H C H O in mid-morning, and O 3 later in the day when its concentration has built up significantly (Winer, 1985; Winer and Biermann, 1994). In summary, NO is now known to be converted to NO 2 during daylight hours in a reaction sequence initiated by OH attack on organics, and involving H O 2 and R O 2 free radicals. These peroxy radicals are the species that actually convert NO to NO 2 at ambient concentrations where the thermal oxidation of NO by 0 2 is negligible.

c. Nighttime Chemistry of NO 2 O(1D) + H 2 0 ~ 2OH.

(14)

In polluted airsheds, other direct sources also form O H through photodissociation, including nitrous acid: H O N O + hu(A < 370 nm) ~ O H + NO, and hydrogen peroxide (H202): H O O H + hu(A < 370 nm) ~ 2OH.

(16)

A very important "thermal" source of hydroxyl radicals, as well as NO2, is the reaction of H O 2 with NO: HO 2 + NO + O H + NO 2.

(17)

This is a major chain propagation step in the overall reaction mechanism for ozone formation in photochemical air pollution. Because HO 2 is intimately tied to OH through reaction (17) and cycles such as that in Fig. 1.4, when NO is present the sources and sinks of HO 2 are, in effect, sources or sinks of the O H radical. Sources of HO 2 include the reactions of 0 2 with hydrogen atoms and formyl radicals, both of which are produced, for example, by the photodissociation of gaseous formaldehyde following absorption of solar actinic UV radiation. HCHO+hv(A

2HNO 3. (25)

Furthermore, N205 plays an important role in key stratospheric heterogeneous processes (see Chapter 12). (2) Gas-phase nitrous acid Gaseous HONO plays an important role in the chemistry of irradiated mixtures of VOC and NOx in air, whether in smog chambers or in ambient atmospheres. Thus, it strongly absorbs actinic UV radiation and, at sunrise (Fig. 1.5), decomposes into OH radicals and NO with a high quantum efficiency: HONO+hv(A

HONO(g) § HNO3(ads?).

(27)

(24)

N O 3 § N O 2 N 2 0 5 .

NzO5(g ) § H 2 0

Use of the long-path length DOAS technique has confirmed its presence in polluted ambient as well as relatively clean continental air masses at levels ranging from ~ 15 ppb at night in a highly polluted air mass down to sub-ppb levels in remote regions (see Chapter 11). It is found at much higher levels in indoor air environments having combustion sources such as gas or propane stoves (see Chapter 15). Direct sources of ambient HONO, established unequivocally through use of the long-path length DOAS technique, include primary emissions, e.g., from lightduty motor vehicles having high levels of NO X in exhaust gases (Pitts et al., 1984). As discussed in Chapter 15, emissions from indoor combustion sources, e.g., gas-fired kitchen stoves and gas or propane heaters, can also produce high levels of HONO in poorly ventilated indoor air environments. Interestingly, heterogeneous processes appear to be involved in HONO formation, certainly in smog chambers and indoor air environments and most likely on a variety of surfaces outdoors. It is produced from gaseous NO 2 and adsorbed water in a heterogeneous reaction on surfaces (see Chapter 7):

~ OH+NO.

(26)

4. Acid Deposition a. Historical

The recognition of acid deposition, commonly called "acid rain," also has a long history. In England and Sweden, the presence of sulfur compounds and acids in polluted air and rain was recognized as early as the eighteenth century. Indeed, in 1692, Robert Boyle referred to "nitrous or salino-sulphurous spirits" in the air in his book A General History of the Air [see the excellent historical perspectives given by Brimblecombe (1978) and Cowling (1982)]. Remarkably, in 1872, a century before it became an international issue, a treatise on acid rain was published in England by Robert Angus Smith. Twenty years earlier, he had analyzed rain near Manchester and noted three types of areas as one moved from the city to the surrounding countryside: "that with carbonate of ammonia in the fields at a distance, that with sulfate of ammonia in the suburbs and that with sulphuric acid or acid sulphate, in the town."

In his 1872 book Air and Rain: The Beginnings of a Chemical Climatology, Smith coined the term acid rain and described many of the factors affecting it, such as coal combustion and the amount and frequency of

9

C. CHEMISTRYOF THE NATURALTROPOSPHERE:REMOTEATMOSPHERES precipitation. He also suggested experimental protocols to be followed in sample collection and analysis and described acid rain damage to plants and materials.

b. Overview of Acidic Rain and Fogs Acid rain arises from the oxidation of SO 2 and NO 2 in the troposphere to form sulfuric and nitric acids, as well as other species, which are subsequently deposited at the earth's surface, either in precipitation (wet deposition) or in dry form (dry deposition). The contribution of organic acids has also been recognized recently (see Chapter 8). These oxidation and deposition processes can occur over relatively short distances from the primary pollutant sources or at distances of a 1000 km or more. Thus both short-range and long-range transport must be considered. The gas-phase oxidation of both SO 2 and NO 2 is initiated by reaction with hydroxyl radicals:

HO 2 q- HO 2

M, H2O > H20 2 q-02,

(11)

HO 2 + RO 2 --+ R O O H + 02,

(12)

RO 2 + RO 2 --+ products.

(29)

Because reactions (11), (12), and (29) do not oxidize NO to NO2, they do not result in the formation of 0 3 (via NO 2 photolysis). Indeed, an additional reaction of HO2, that with 03, HO 2 + 0 3 --+ OH + 202,

H20 OH + 802 ---> HOSO 2 --+ > H2804,

(28)

OH + NO 2 M_~ HNO3"

(10)

M

tially the same V O C - N O X chemistry described earlier in this chapter. However, given that there are the significant biogenic sources of a wide variety of organics, the major difference lies in the NOx levels, which are much lower in remote regions. Because of these lower concentrations of NO, the self-reactions of HO 2 and RO 2 radicals can become competitive with NO reactions:

In the case of 802, oxidation in the aqueous phase, present in the atmosphere in the form of aerosol particles, clouds, and fogs, is also important. Thus SO 2 from the gas phase dissolves in these water droplets and may be oxidized within the droplet by such species as H202, 03, 02, and free radicals. Oxidation of SO 2 on the surfaces of solids either present in the air or suspended in the water droplets is also possible. On the other hand, it is believed that H N O 3 is formed primarily by reaction (10) in the gas phase and subsequently dissolves in droplets. These oxidation processes can lead to highly acidic fogs. For example, pH values as low as 1.69 have been measured in coastal regions of southern California (Jacob and Hoffmann, 1983). These high acidities, accompanied by high concentrations of other anions and cations, are likely due to evaporation of water from the fog droplets, leaving very high concentrations of ions in a strongly acidic liquid phase. Such acid fogs, whether in London or Los Angeles, are a major health concern because the droplets are sufficiently small to be efficiently inhaled (Hoffmann, 1984).

actually results in the destruction of ozone. As developed in more detail in Chapter 6, the NO concentration at which other reactions such as (11), (12), and (29) occur at approximately the same rate as the H O z / R O 2 + NO reaction is in the 10-50 ppt range. These concentrations are sufficiently low that they are encountered only in remote atmospheres, where the influence of anthropogenic emissions is minimized. However, it is noteworthy that on a global level, 0 3 levels have been increasing significantly over the past century, coinciding with the increase in fossil fuel usage and the associated increase in NO~ emissions (Bojkov, 1986; Volz and Kley, 1988; McKeen et al., 1989; Anfossi et al., 1991; Sandroni et al., 1992; Marenco et al., 1994; Anfossi and Sandroni, 1997). Figure 1.6 shows that 0 3 levels appear to have been about 10-15 ppb a century ago, compared to 30-40 ppb measured around the world today. This is consistent with a shift in the

0.03

E o.) tO

C. C H E M I S T R Y

OF THE NATURAL TROPOSPHERE: REMOTE ATMOSPHERES

Although there is sometimes a tendency to think of the chemistry of the "natural" troposphere as somehow different from that of more polluted areas, it is essen-

(30)

N o

0.02 -

0.01 -

0 1850

I

F,

1875

I

I

I

I

1900

1925

1950

1975

2000

Year

FIGURE 1.6 Typical tropospheric ozone concentrations in the 1800's and present values (adapted from Volz and Kley, 1988).

10

1. CHEMISTRYOF POLLUTED AND REMOTE ATMOSPHERES

chemistry on a global scale from the self-reactions of HO 2 and RO 2 [reactions (11), (12), (29), and (30)with 0 3] to their reaction with NO to generate NO 2 and subsequently 0 3. In short, the chemistry of remote regions differs primarily in the availability of NO~.

D. CHEMISTRY OF THE STRATOSPHERE The stratosphere is often referred to as the "ozone layer," because of the relatively high concentrations produced by photochemical reactions in this region of the atmosphere. Ozone, derived from the Greek word meaning "to smell," was first discovered by Sch6nbein in 1839. It has a rather pungent smell, which is sometimes noticeable around copy machines and laser printers that use discharge processes. As described earlier, in the stratosphere, a steadystate concentration of 0 3 is produced naturally by the Chapman cycle, reactions (1)-(4). Until about 1970, relatively little attention was paid to potential anthropogenic (i.e., man-made) perturbations of the stratosphere. At that time, Crutzen (1970) examined the potential role of NO and NO 2 formed in the stratosphere from reactions of N20 that was originally generated at the earth's surface. Because N20 is unreactive in the troposphere, it has a sufficiently long lifetime to end up in the stratosphere, where it can be converted into NO (see Chapter 12). Crutzen (1970) proposed that the NO and NO 2 formed from reactions of N20 can then participate in a chain reaction that destroys O3: NO + 0 3 ~ NO 2 + 0 2 ,

(8)

NO 2 + h v ~ NO + O,

(31)

NO 2 + O ~ NO

(32)

+ 0 2.

Subsequently, Johnston (1971) suggested that NO X emitted directly into the stratosphere from supersonic transport aircraft (SSTs)could decrease the steady-state concentration of 0 3 via reactions (8) and (32), leading to increased UV radiation at the earth's surface. Although the number of SSTs that ultimately were produced was far less than originally anticipated, the same issue of NOx destruction of 0 3 has again been revisited with respect to the possible future production and use of high-speed civil transport (HSCT) aircraft (see Chapter 12). In 1974, Cicerone and Stolarski suggested that if there were sources of atomic chlorine in the stratosphere, the following catalytic ozone destruction cycle

could occur: C1 + 0 3 ~ C10 + 0 2

(33)

C10 + O ~ C1 + 0 2

(34)

Net: 0 3 + O ~ 2 0 2 Shortly thereafter, Molina and Rowland (1974) published their seminal paper in which they showed that, because of the lack of removal mechanisms for these inert chemicals in the troposphere, chlorofluorocarbons (CFCs) are expected to reach the stratosphere. Once in the stratosphere, the CFCs are exposed to UV radiation in the region in which they absorb, leading to the production of chlorine atoms: CF 2C12 + h v ~ C1 + CF 2C1.

(35)

From the outset, their hypothesis caused worldwide concern because UV-B can damage DNA, cause skin cancer, including malignant melanoma in humans, and have long-term effects on the eye, in addition to its impact on a variety of ecosystems. This foundation of the interaction between tropospheric processes and emissions and stratospheric chemistry established by Crutzen, Molina, and Rowland led to their sharing the 1995 Nobel Prize in Chemistry. Until the mid-1980's, estimates of the reduction in the steady-state concentration of stratospheric 0 3 were based largely on the application of large computer models that incorporated the known mechanisms and kinetics of individual chemical reactions. Because of natural variations in stratospheric ozone, unequivocal detection of the expected change is quite difficult. However, in 1985 a remarkable observation was made by a group of British researchers (Farman et al., 1985) who had been monitoring total column ozone in the Antarctic. Total column ozone is the integrated ozone measured in a vertical column throughout the atmosphere and is often expressed in Dobson units (DU), where 1 D U is equivalent to a column of 0 3 of 10 .5 m height at STP (1 atm pressure and 273.15 K). Since approximately 85-90% of the total column ozone is in the stratosphere, changes in column ozone are particularly sensitive to stratospheric ozone levels. Figure 1.7 shows the total column ozone measured in October at one Antarctic location, Halley Bay, as a function of year. This includes the original Farman et al. (1985) data, as well as more recent data up to 1994 (Jones and Shanklin, 1995). It is clear that starting in the late 1970's, there was a dramatic drop in total column ozone at the end of the polar winter when sunrise occurs. Observation of such a rapid change is unprecedented and quite remarkable.

11

E. GLOBAL CLIMATE CHANGE 350

_

300

"-,.-r162,.

9

9 WV 9

o

250

c

oU 0

200

~6

150

m

_

VV 9

100

1950

I

1960

I

I

1970

1980

I

1990

2000

Year

F I G U R E 1.7 Average total column ozone measured in October at Halley Bay, Antarctica, from 1957 to 1994 (adapted from Jones and Shanklin, 1995).

cally and C10 rose simultaneously. The key to the dramatic changes in 0 3 occurring in this region now appears to be "heterogeneous" chemistry occurring on and in polar stratospheric clouds (PSC's), combined with the formation of a relatively well-contained air mass over the continent during the polar winter. This unique combination of chemistry and meteorology is discussed in more detail in Chapter 12. In short, although the history of anthropogenic perturbations to the stratosphere is much shorter, it is clear that these are also important. Indeed, such perturbations are expected to affect the chemistry of the troposphere as well; for example, increased UV radiation will alter photochemistry at the earth's surface.

E.

After the first reports of this phenomenon, major field campaigns were launched, which clearly established a relationship between ozone destruction and chlorine chemistry. For example, Fig. 1.8 shows simultaneous aircraft measurements of ozone and the free radical C10 as the plane flew toward the South Pole. As it entered the polar vortex, a relatively well-contained air mass over Antarctica, 0 3 dropped dramati-

1.20 - 2.50

0.84

1.75 .12} Q.

& o o

E v tO N

0.48

.

,~ I

v'~-.

,v,,, CIO

' 64

t

/

.~1 L/

| 0.75

It/

/

'

t

66

68 Latitude

O

'

!

t 0.25 70

~

F I G U R E 1.8 Measured concentrations of the chlorine monoxide free radical (C10) as well as 0 3 outside and inside the polar vortex on August 23, 1987 (adapted from Anderson, 1989).

GLOBAL

CLIMATE

CHANGE

Destruction of stratospheric ozone caused by relatively small atmospheric concentrations of chlorofluorocarbons has vividly illustrated the capacity of human activity to alter our atmosphere in a manner that has significant and far-ranging effects. There is similar concern for the effects of greenhouse gases on the earth's climate. Solar radiation provides energy to the earth. Radiative balance is achieved through a combination of processes summarized in Fig. 1.9. A major process is the re-emission of radiation in the infrared by the earth's surface, so-called "terrestrial radiation." This infrared radiation is absorbed and re-emitted by a number of atmospheric gases, during which part of the energy is converted to thermal energy. The result of the balance in the incoming and outgoing radiation is a relatively "warm" troposphere, compared to what it would be without this natural "greenhouse effect." It is known from measurements that the atmospheric concentrations of a number of greenhouse gases have increased significantly since the beginning of the industrial revolution. These include not only the wellrecognized CO 2 but also a host of other trace gases, including 03, as already discussed, and N20 and CH 4 whose major sources are biogenic. Increases in such gases alter the radiation balance by trapping more of the terrestrial infrared radiation, which results in a larger amount of energy converted to thermal energy in the troposphere. More complex are the effects of aerosol particles. These impact the radiative properties of the earth both through direct effects such as scattering and, to a lesser extent, absorbing solar radiation and through indirect effects, by altering cloud properties.

12

1. CHEMISTRY OF POLLUTED AND REMOTE ATMOSPHERES

n

I /Arela_ - ~ 4 ~ 7 ~ , ~r 2 / /

1368W m2

!

Incoming Solar Radiation

_•41368 = 342 W m2

'~'~,W'~..% Reflection in ~ Atmosphere ~

C

Surface Reflection (rim

Outgoing Radiation

. Heat Transfer by Thermals (24) ~J(~

~

Absorption in Atmosphere by CO 2, 0 3, H20 and 0 2 {1~7'~

~235

w m2

Upward Emission from Absorption by / ~ H20, CO2,03, clouds etc Release of H20, CO 2, 0 3 etc ,,/r" (195) Latent Heat of and clouds (350) ff/,J Direct Surface Evaporation A, ,,f~l,/ Radiation (40) (78) lib ~ ,,( r / " " 1~ ~ / / Downward Emission from /fir I F t( /J' r ] H20 CO 2 0 3 etc Condensation J j ~ -/ . . . . of Water I' /(390)1 (324) i i

bsorption by Surface F I G U R E 1.9 Global average mean radiation and energy balance per unit of earth's surface [adapted with permission from IPCC (1996)with numbers from Kiehl and Trenberth (1997)].

aEFEaENCES Significantly complicating predictions of the impact of human activities on global climate is incomplete understanding of feedbacks, which can be either positive or negative. As a result, there are currently large "error bars" on predictions of global climate change due to human activities, and this is a very active area of research. The chemical aspects of the atmosphere relevant to global climate issues are discussed in detail in Chapter 14.

F. I N D O O R

AIR POLLUTION

A major reason for understanding the chemistry of the atmosphere is the impact that changes can have on human health and well-being. With respect to effects due to direct inhalation of gases and particles, exposure occurs not only outdoors but also indoors as well. Indeed, the vast majority of time for most people is spent indoors. As a result, understanding the nature of the indoor atmosphere and human impacts on it is important as well. A number of studies have documented that concentrations of some of the directly emitted species found in outdoor atmospheres can be quite high indoors if there are emission sources present such as combustion heaters, gas stoves, or tobacco smoke. In addition, there is evidence for chemistry analogous to that occurring outdoors taking place in indoor air environments, with modifications for different light intensities and wavelength distributions, shorter residence times, and different relative concentrations of reactants. In Chapter 15, we briefly summarize what is known about the chemical composition and chemistry of indoor atmospheres.

13

Some Points for Discussion

a. What techniques were used for measuring ozone in the nineteenth century? What were the likely interferences a n d / o r parameters that could affect the results and how serious were they likely to be? b. At what locations were measurements made for a period of years? c. Given the potential interferences, what are the typical levels of tropospheric 0 3 in the nineteenth century? d. Given typical tropospheric 0 3 levels today, discuss possible causes for the increase. e. Name as many impacts as you can think of that might be associated with this increase in 0 3 on a global basis. Be sure to include effects on the chemistry of the atmosphere as well as effects on health, global change, etc. f. Some of the ozone levels reported in these papers are in mPa. What is 1 mPa in the more commonly used units of ppb? (See Chapter 2 for a discussion of units and conversions.) 2. O Z I P R M o d e l The OZIPR model is described in Appendix III. It is a box model that has comprehensive chemical submodels that lump organics in two different ways. It allows for variation of meteorology through time-dependent mixing heights, temperatures, etc. Familiarize yourself with this model and explore the effects of meteorology on the predicted peak ozone concentrations by varying the initial mixing height and its rate of change with time. Examine the effects of photochemistry by varying the date from June 21 to December 21 and by varying the latitude. References

G. D I S C U S S I O N T O P I C A N D OZIPR MODEL 1. D i s c u s s i o n T o p i c : " B a c k g r o u n d O z o n e " It was accepted for a number of years in the atmospheric chemistry community that so-called "background" ozone was typically around 30-40 ppb. However, starting in the mid-1980's, a number of researchers examined the literature of a century earlier, shortly after ozone was discovered by Sch6nbein in 1839, and discovered several series of ozone measurements that had been made at different locations in the troposphere. Some of the papers describing this research include the papers by Bojkov (1986), Volz and Kley (1988), and Anfossi et al. (1991).

Allbgre, C. J., and S. H. Schneider, "The Evolution of the Earth," Sci. Am., 66-75, October (1994). Anderson, J. G., "Free Radicals in the Earth's Atmosphere: Measurement and Interpretation," in Ozone Depletion, Greenhouse Gases and Climate Change, National Academy Press, Washington, DC, pp. 56-65, 1989. Anfossi, D., S. Sandroni, and S. Viarengo, "Tropospheric Ozone in the Nineteenth Century: The Moncalieri Series," J. Geophys. Res., 96, 17349-17352 (1991). Anfossi, D., and S. Sandroni, "Short Communication--Ozone Levels in Paris One Century Ago," Atmos. Environ., 31, 3481-3482 (1997). Atkinson, R., A. M. Winer, and J. N. Pitts, Jr., "Estimation of Nighttime N20 5 Concentrations from Ambient NO 2 and NO 3 Radical Concentrations and the Role of N20 5 in Nighttime Chemistry," Atmos. Environ., 20, 331-339 (1986). Blacet, F. E., "Photochemistry in the Lower Atmosphere," Ind. Eng. Chem., 44, 1339-1342 (1952). Bojkov, R. D., "Surface Ozone during the Second Half of the Nineteenth Century," J. Am. Meteorol. Soc., 25, 343-352 (1986).

14

1. CHEMISTRYOF POLLUTED AND REMOTE ATMOSPHERES

Brimblecombe, P., "Interest in Air Pollution among Early Fellows of the Royal Society," Notes Rec. R. Soc., 32, 123-129 (1978). Cicerone, R. J., and R. S. Stolarski, "Stratospheric Chlorine: A Possible Sink for Ozone," Can. J. Chem., 52, 1610-1615 (1974). Cochran, L. S., R. A. Pielke, and E. Kovfics, "Selected International Receptor-Based Air Quality Standards," J. Air Waste Manage. Assoc., 42, 1567-1572 (1992). Cowling, E. B., "Acid Precipitation in Historical Perspective," Environ. Sci. Technol., 16, 110A-123A (1982). Crutzen, P. J., "The Influence of Nitrogen Oxides on the Atmospheric Ozone Content," Q. J. R. Meteorol. Soc., 96, 320-325 (1970). Evelyn, J., Fumifugium: Or the Inconvenience of the Aer and Smoake of London Dissipated, Together with Some Remedies Humbly Proposed, Bedel and Collins, London, 1661. Farman, J. C., B. G. Gardiner, and J. D. Shanklin, "Large Losses of Total Ozone in Antarctica Reveal Seasonal C1Ox/NO X Interaction," Nature, 315, 207-210 (1985). Finlayson-Pitts, B. J., and J. N. Pitts, Jr., "The Chemical Basis of Air Quality: Kinetics and Mechanisms of Photochemical Air Pollution and Application to Control Strategies," Adv. Environ. Sci. Technol., 7, 76-162 (1977). Firket, J., "Fog along the Meuse Valley," Trans. Faraday Soc., 32, 1192-1197 (1936). Goodhill, V., "Maimonides--Modern Medical Relevance," XXVI Wherry Memorial Lecture, Transactions of the American Academy of Ophthalmology and Otolaryngology, p. 463, May-June 1971. Haagen-Smit, A. J., "Chemistry and Physiology of Los Angeles Smog," Ind. Eng. Chem., 44, 1342-1346 (1952). Haagen-Smit, A. J., and M. M. Fox, "Photochemical Ozone Formation with Hydrocarbons and Automobile Exhaust," J. Air Pollut. Control Assoc., 4, 105-108, 136 (1954). Haagen-Smit, A. J., and M. M. Fox, "Automobile Exhaust and Ozone Formation," SAE Trans., 63, 575-580 (1955). Haagen-Smit, A. J., and M. M. Fox, "Ozone Formation in Photochemical Oxidation of Organic Substances," Ind. Eng. Chem., 48, 1484-1487 (1956). Haagen-Smit, A. J., E. F. Darley, M. Zaitlin, H. Hull, and W. Noble, "Investigation on Injury to Plants from Air Pollution in the Los Angeles Area," Plant Physiol., 27, 18-34 (1952). Haagen-Smit, A. J., C. E. Bradley, and M. M. Fox, "Ozone Formation in Photochemical Oxidation of Organic Substances," Ind. Eng. Chem., 45, 2086-2089 (1953). Heicklen, J., K. Westberg, and N. Cohen, Center for Air Environmental Studies, Report No. 115-69, 1969. Hoffmann, M. R., "Response to Comment on 'Acid Fog,'" Environ. Sci. Technol., 18, 61-64 (1984). IPCC, Intergovernmental Panel on Climate Change, Contribution of Working Group I to the Second Assessment Report (J. T. Houghton, L. G. Meira Filho, B. A. Callander, N. Harris, A. Kattenberg, and K. Maskell, Eds.), Climate Change 1995, Cambridge Univ. Press, Cambridge, UK, 1996. Jacob, D. J., and M. R. Hoffmann, "A Dynamic Model for the Production of H+, NO3, and SO 2- in Urban Fog," J. Geophys. Res., 88, 6611-6621 (1983). Johnston, H. S., "Reduction of Stratospheric Ozone by Nitrogen Oxide Catalysts from Supersonic Transport Exhaust," Science, 173, 517-522 (1971). Jones, A. E., and J. D. Shanklin, "Continued Decline of Total Ozone over Halley, Antarctica, Since 1985," Nature, 376, 409-411 (1995). Kasting, J. F., "Earth's Early Atmosphere," Science, 259, 920-926 (1993). Kiehl, J. T., and K. E. Trenberth, "Earth's Annual Global Mean Energy Budget," Bull. Am. Meteorol. Soc., 78, 197-208 (1997).

Kirchner, J. W., "The Gaia Hypothesis: Can It Be Tested?", Rev. Geophys., 27, 223-235 (1989). Lenton, T. M., "Gaia and Natural Selection," Nature, 394, 439-447 (1998). Levy, H., "Normal Atmosphere: Large Radical and Formaldehyde Concentrations Predicted," Science, 173, 141-143 (1971). Lodge, J. P., Jr., Ed., The Smoake of London. Two Prophecies: Evelyn, J., 1661. Fumifugium: Or the Inconvenience of the Aer and Smoake of London Dissipated, and Barr, R., date unknown, The Doom of London, Maxwell Reprint Co., Elmsford, New York, 1969. Lovelock, J. E., "Geophysiology, the Science of Gaia," Rev. Geophys., 27, 215-222 (1989). Mamane, Y. "Air Pollution Control in Israel during the First and Second Century," Atmos. Environ., 21, 1861-1863 (1987). Marenco, A., H. Gouget, P. Nedelec, J.-P. Pages, and F. Karcher, "Evidence of a Long-Term Increase in Tropospheric Ozone from Pic du Midi Data Series: Consequences: Positive Radiative Forcing," J. Geophys. Res., 99, 16617-16632 (1994). McKeen, S., D. Kley, and A. Volz, "The Historical Trend of Tropospheric Ozone over Western Europe: A Model Perspective," in Ozone in the Atmosphere (R. D. Bojkov and P. Fabian, Eds.), A. Deepak Publishing, Hampton, VA, 1989. Middleton, J. T., J. B. Kendrick, Jr., and H. W. Schwalm, "Injury to Herbaceous Plants by Smog or Air Pollution," U.S.D.A. Plant Dis. Rep., 34, 245-252 (1950). Molina, M. J., and F. S. Rowland, "Stratospheric Sink for Chlorofluoromethanes: Chlorine Atom-Catalysed Destruction of Ozone," Nature, 249, 810-812 (1974). Perner, D., and U. Platt, "Detection of Nitrous Acid in the Atmosphere by Differential Optical Absorption," Geophys. Res. Lett., 6, 917-920 (1979). Platt, U., D. Perner, and H. W. P~itz, "Simultaneous Measurement of Atmospheric CH20 , 03, and NO 2 by Differential Optical Absorption," J. Geophys. Res., 84, 6329-6335 (1979). Platt, U., D. Perner, A. M. Winer, G. W. Harris, and J. N. Pitts, Jr., "Detection of NO 3 in the Polluted Troposphere by Differential Optical Absorption," Geophys. Res. Lett., 7, 89-92 (1980). Pitts, J. N., Jr., H. W. Biermann, A. M. Winer, and E. C. Tuazon, "Spectroscopic Identification and Measurement of Gaseous Nitrous Acid in Dilute Auto Exhaust," Atm. Environ., 18, 847-854 (1984). Rouech6, B., Eleven Blue Men, pp. 173-191, New Berkley Medallian Edition, New York, 1965. Sandroni, S., D. Anfossi, and S. Viarengo, "Surface Ozone Levels at the End of the Nineteenth Century in South America," J. Geophys. Res., 97, 2535-2539 (1992). Schneider, S. H., "Debating Gaia," Environment, 32, 5-9 & 29-32 (1990). Stedman, D. H., E. D. Morris, Jr., E. E. Daby, H. Niki, and B. Weinstock, "The Role of OH Radicals in Photochemical Smog Reactions," 160th National Meeting of the American Chemical Society, Chicago, IL, Sept. 14-18, 1970. Volz, A., and D. Kley, "Evaluation of the Montsouris Series of Ozone Measurements Made in the Nineteenth Century," Nature, 332, 240-242 (1988). Weinstock, B., "Carbon Monoxide: Residence Time in the Atmosphere," Science, 166, 224-225 (1969). Wilkins, E. T., "Air Pollution and the London Fog of December 1952," J. R. Sanitary Inst., 74, 1-21 (1954). Winer, A. M., "Air Pollution Chemistry," in HandbookofAir Pollution Analysis (R. M. Harrison and R. Perry, Eds.), 2nd ed., Chap. 3. Chapman and Hall, London, 1985. Winer, A. M., and H. W. Biermann, "Long Pathlength Differential Optical Absorption Spectroscopy (DOAS) Measurements of Gaseous HONO, NO2, and HCHO in the California South Coast Air Basin," Res. Chem. Intermed., 20, 423-445 (1994).

C H A P T E R

2 The A tmospheric Syst e r n As discussed in Chapter 1, much of our understanding of the chemistry of our atmosphere is based on early studies of air pollution; these are often treated in the context of an overall "system." This approach starts with the various sources of anthropogenic and natural emissions and tracks the resulting pollutants through their atmospheric transport, transformations, and ambient concentrations--on local, regional, and global scales--to their ultimate chemical and physical fates, including their impacts on our health and environment. Figure 2.1 is a simplified diagram illustrating the major elements. Primary pollutants are defined as those emitted directly into the air, e.g., SO 2, NO, CO, Pb, organics [including HAPS (hazardous air pollutants)], and combustion-generated particulate matter (PM). Sources may be anthropogenic, biogenic, geogenic, or some combination thereof. Once in the atmosphere, they are subjected to dispersion and transport, i.e., meteorology, and simultaneously to chemical and physical transformations into gaseous and particulate secondary pollutants; the latter are defined as those formed from reactions of the primary pollutants in air. Both primary and secondary pollutants are removed at the earth's surface via wet or dry deposition and, in the processes of transport, transformation, and deposition, can impact a variety of receptors, for example, humans, animals, aquatic ecosystems, forests and agricultural crops, and materials. From a detailed knowledge of the emissions, topography, meteorology, chemistry, and deposition processes, one can develop mathematical models that predict the concentrations of primary and secondary pollutants as a function of time at various locations. Depending on the particular model, these may describe pollutant concentrations over a variety of scales:

group of sources (long-range transport and regional models) Over the entire earth (global models) To test these models, their predictions must be compared to the observed concentrations of various species; model inputs are adjusted to obtain acceptable agreement between the observed and predicted values. These models can then be used, in combination with the documented impacts on receptors, to develop health a n d / o r environmental risk assessments and various control strategy options. Finally, through legislative and administrative action, health-protective and cost-effective risk-management decisions can be made, and regulatory actions implemented, that directly affect the starting point of our atmospheric system, that is, the primary emissions and their sources. To place the remainder of this book on atmospheric chemistry in perspective, the various components of our "atmospheric system" are treated briefly next.

A. EMISSIONS In describing a given air mass and the chemical reactions occurring therein, one must consider both natural and anthropogenic sources of primary emissions and evaluate their relative importance. Thus the impact on air quality of natural emissions can be an important issue because cost-effective control strategies must take into account the relative strengths of emissions from all sources, not just those of anthropogenic origin. However, it is not only the relative amounts of total emissions that must be considered but also the chemical nature of the emissions, e.g., their reactivities and their temporal and spatial distributions. Emissions inventories are typically obtained by combining the rate of emissions from various sources (the "emission factors")with the number of each type of source and the time over which the emissions occur.

9 In a plume from a specific point source (plume models) 9 In an air basin from a combination of diverse mobile and stationary sources (airshed models) 9 Over a large geographical area downwind from a

15

16

2. THE ATMOSPHERICSYSTEM

Inventories are compiled in various formats. For example, they can be assembled for various individual anthropogenic processes such as refining, or natural processes such as volcanic eruptions, in which emissions of all of the relevant species associated with that event are estimated. Alternatively, and more commonly, emissions inventories are compiled by species, showing the various sources that contribute to the total emissions of each. Emission factors for various sources in the United States have been published by the Environmental Protection Agency in the form of the document

AIRCHIEF, short for the Air Clearing House for Inventories and Emission Factors. Such data are available on C D - R O M as well as on-line through the EPA Web site (see Appendix IV). In Europe, the Commission of the European Communities has published a handbook of emission factors as well (e.g., see CEC (1988, 1989, 1991), McInnes (1996), and Web site in Appendix IV). Emissions inventories and emission factors for Europe are also found in the volume edited by Fenger et al. (1998). On a global scale, emissions inventories for a variety of species are currently under development under the

Sources ]

Emissions of Anthropogenic, Biogenic, Geogenic Primary Pollutants e.g. VOC, NOx, SO2, CO, PMlo ' 2.5, HAPs

Dispersion and Transport

Risk Management Decisions Air Pollution Control

Scientific Risk Assessment I

Chemical and Physical Transformations

!

AmbientAir Urban, Suburban, Rural, Remote, 03, Acids, Toxics, PMlo ' 2.5 etc.

j

L.

Monitoring

~

[

~

Models Local "Hot-Spot" Plume, Airshed, Long-range Transport, Global

FATES

1 Exposure1

f ,~

Long-Lived Species e.g. CFC, N20

I Wet and Dry Deposition

Transport to Stratosphere

Stratospheric Chemistry, Ozone Depletion

FIGURE 2.1 The atmospheric air pollution system.

Effects: Health and Environmental

l

Impactson Receptors (Humans, Animals, Agricultural Crops Forest and Aquatic Ecosystems, Visibility, Materials, etc.)

A. EMISSIONS

auspices of the International Global Atmospheric Chemistry Project (IGAC), and various available inventories are described by Graedel et al. (1993). Data for some of the major pollutants follow.

We shall follow here the convention in current use that defines NO~ as the sum of (NO + NO 2) and NOy as the sum of all reactive nitrogen-containing species, e.g., NOy = ( N O + N O 2 + H N O 3 + P A N + H O N O + NO 3 + N20 5 + organic nitrates etc.). By far the most significant species emitted by anthropogenic processes is nitric oxide, produced when N 2 and 0 2 in air react during high-temperature combustion processes. In addition, some NO~ is formed from nitrogen in the fuel. Smaller amounts of NO 2 are produced by the further oxidation of NO; trace amounts of other nitrogenous species such as HNO 3 are also formed. The fraction of the total that is emitted as NO clearly depends on the conditions associated with the specific combustion process. While most (typically > 9 0 % ) o f the NOx emitted is believed to be in the form of NO, the fraction of NO 2 can vary from less than 1% to more than 30% (e.g., Lenner, 1987). Figure 2.2 shows the contribution of various sources to the total anthropogenic NO~ emissions, 23 • 106 short tons, or 21 Tg (expressed as NO2), in the United States in 1996 (1 Tg = 1 teragram = 1012 g and one short ton = 0.907 • 106 g). This can be compared to total global anthropogenic emissions of approximately 72 Tg of NO~ (expressed as NO 2) (Miiller, 1992). Figure 2.3 shows the trend in NO~ emissions for North America, Europe, the USSR, and Asia from 1970 to 1986 (Hameed and Dignon, 1992). While those

Fuel Combustion Electrical U t i l i t y / ~ : ~ 28~

29.6 0 xZ 0 0o Z ~ "b ~o~ t'-(. o

1. O x i d e s of N i t r o g e n

On-Road Vehicles 30%

All Other 5~176 Non-Road~,'~i~' i:i~' i " , ~ // ~--~-"~ ! ~ Engines " ~ ~ / / y Fuel Combustion- Other 5% and Vehicles 19% Fuel Combustion Industrial 13% F I G U R E 2.2 Contribution of various sources to total anthropogenic NOx emissions in the United States in 1996 (from EPA, 1999).

17

/

23.0

North America

~~..-

~

~ ' - ----

19.7

.o o

16.4

~ o

13.1

._ C

f----~

26.3

-

USSR

9.9 [".......... I

6.0 1970

I

I

.,.... ~." :: :"':

"'")~sia l

l

l

I

I

l

1975

l

I

1980

l

I

I

I

l

1985

Year

FIGURE 2.3 NOXemissionsin million tons of equivalent N O 2 for the period 1970 to 1986 for Asia, Europe, North America, and the USSR (from Hameed and Dignon, 1992).

of North America and Europe have decreased or leveled off, those from Asia and the USSR increased significantly, a trend that has continued. Figure 2.4a shows the geographical pattern of the emission flux of NO X in Asia in 1987 (Akimoto and Narita, 1994). Clearly, Japan and China are major contributors to the flux of NOx in this region, with the City of Tokyo having the highest emission flux rate. There are also significant natural sources of oxides of nitrogen, in particular nitric oxide, which is produced by biomass burning as well as by soils where nitrification, denitrification, and the decomposition of nitrite (NO 2) contribute to NO production. Figure 2.4b, for example, shows the relative emission rates for biogenically produced NO in the United States in 1990 (EPA, 1995). Another important natural source is NO~ produced by lightning, with recent estimates in the range of 10-33 Tg yr -1 as NO 2 (Flat W and Hov, 1997; Price et al., 1997a, 1997b; Wang et al., 1998). By comparison to the estimated emissions from biomass burning and continental biogenic sources (Table 2.1), it is seen that lightning is quite important. There is also some NO produced from the oxidation of NH 3 by photochemical processes in oceans and by some terrestrial plants (e.g., Wildt et al., 1997). Table 2.1 gives an estimate for global-scale natural and anthropogenic emissions of NO~ as well as of CO, CH4, and VOC (Miiller, 1992). It is seen that biomass burning and biogenic emissions of NO are comparable and together equal to about half of the anthropogenic emissions. Nitrous oxide (N20, "laughing gas") is also produced by biological processes and, to a lesser extent, by anthropogenic processes (see Chapter 14.B.2c). While

18

2. THE ATMOSPHERIC SYSTEM

a

:FI

_[]

I~FI [] 1 (10-3 moIes as N m-2 yr-1) FFI3 N10 [] 20 [] 50

100 ~ 9200

b

Biogenic NO emissions ........... ,:____

.~, 5:

.............. , ....................... _

~::- Z -

_. : - 5=

:(.i

._

~

"

":~"~

~ ~

~

[] []

Above Average Average

[ ~ BelowAverage

r-i ~ow

FIGURE 2.4 (a) Pattern of 1987 annual emission flux of NO x in Asia (in units of millimoles as N per m2 per year) (from Akimoto and Narita, 1994). (b) Estimated relative rates of biogenic emissions of NO in the United States in 1990 (from EPA, 1995).

biomass burning is the major anthropogenic

source,

stratosphere, where it acts as a major source of NO X

there are a variety of smaller sources, including motor vehicles. Interestingly, the N 2 0 emissions from current catalyst-equipped cars appear to be higher than from noncatalyst-equipped vehicles (Berges et al., 1993), but

(see Chapters 1 and 12).

total NOx emissions are much lower. N 2 0 is inert in the troposphere and is ultimately transported to the

Historically, organics in the troposphere have been measured as non-methane hydrocarbons (NMHC). As

2. V o l a t i l e O r g a n i c C o m p o u n d s ( V O C )

A. EMISSIONS

Non-Road Engines and Vehicles

T A B L E 2.1 Global Emission Estimates for CO, N O x, C H 4, and V O C from Both Anthropogenic and N a t u r a l Sources (in T g / y r ) a

CO NOr CH 4

VOC

Anthropogenic sources

Biomass burning

Continental biogenic sources b

383 72 132 98

730 18 54 51

165 22 310 500

19

J ' ~ ~ - : : : : ~ Storage 7&oTransport~ ~ . : . : . : . : . : . : . : . : . : \

Oceans

Total

165 0.01 10 30-300

1440 122 506 750

Miscellaneous 3

On-Road Vehicles 29%

i tliiiiiiiiiiiiiiiiiiiiiiiiii ~

AII15Oto2er ~ : ~ / 2 . .

a Source: Miiller (1992). b Includes animal, microbial, and foliage emissions. c Expressed as N O 2.

~- ' ' - "

.--. " 2]

Solvent Utilization 33% FIGURE 2.5

discussed elsewhere in this book, methane oxidizes relatively slowly in the troposphere and hence on a short-term basis (e.g., hours) does not contribute significantly to ozone formation compared to other organics (with some exceptions, e.g., CFCs). As a result, in terms of smog formation in urban and regional areas, methane has been excluded from consideration and controls have focused on the suite of larger organics. However, with the recognition that a variety of organics (e.g., aldehydes) and not just hydrocarbons are important in the troposphere, alternative terminologies have been used to describe organics other than methane. These include the most commonly used term VOC (volatile organic compounds) as well as ROG (reactive organic gases) and NMOC (non-methane organic compounds). We use the term VOC, recognizing that certain volatile species such as CFCs are not included.

Contribution of various sources to total anthro-

pogenic VOC emissions in the United States in 1996 (from EPA, 1997).

Figure 2.5 shows the distribution of anthropogenic sources of VOCs in the United States in 1996, a total of 19 • 106 short tons, or 17 Tg (EPA, 1997). Figure 2.6 shows VOC emissions in Europe (Friedrich and Obermeier, 1999). Over the past decade, it has become clear that, on a global scale, biogenic processes also release substantial quantities of reactive hydrocarbons such as isoprene and a-pinene, in addition to methane and other organics, including oxygenated species such as methanol, 2-methyl-3-buten-2-ol, hexenol, acetone, and formic and acetic acids (see, for example, Fehsenfeld et al., 1992; Golden et al., 1993; Monson et al., 1995; Singh et al., 1995; Fall, 1999; Guenther, 1999); see Chapter 6.J.1.

3000 other sectors [-I road transport solvent use

2500 2000 o0

o=

1500

o O

,--~

1000

500

(D "t"

T5 Z ~-

-~

(/3

~

o a

o

F I G U R E 2.6 Annual 1990 VOC emissions in Europe [adapted from McInnes (1996) and Friedrich and Obermeier (1999)].

20

2. THE ATMOSPHERIC SYSTEM T A B L E 2.2 Estimates of Global V O C Emissions Rates (Tg of C yr - 1)a

Other more reactive Source Isoprene Monoterpene organics b

Woods Crops Shrubs Ocean Other All

372 24 103 0 4 503

95 6 25 0 1 127

177 45 33 2.5 2 260

On-Road V e h i c l e s / ~ ~ ! : : . ~ .

iiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiii 60

Other less reactive organics c Total VOC

177 45 33 2.5 2 260

821 120 194 5 9 1150

a From Guenther et al. (1995). b Defined as having lifetimes < 1 day under typical tropospheric conditions. c Defined as having lifetimes of > 1 day.

Table 2.2 shows one estimate of the global annual organic emissions from biogenic and oceanic sources and biomass burning. Such estimates are highly uncertain because the factors controlling biogenic emissions are complex, leading to large variations with time of day, season, geographical location, etc. (Lamb et al., 1993; Geron et al., 1994; Monson et al., 1995). Illustrative of this is the significant difference in the VOC emission estimates between Tables 2.2 and 2.1. Some sources are not included in these estimates, such as small amounts of isoprene produced by phytoplankton in oceans (e.g., Bonsang et al., 1992; Moore et al., 1994). A biogenics emission inventory for the United States is given by Pierce et al. (1998). The role of biogenic VOC in tropospheric chemistry is discussed in Chapter 6. For a review of VOC in natural and polluted atmospheres, see the book edited by Hewitt (1999). 3. C a r b o n M o n o x i d e Carbon monoxide is produced by the incomplete combustion of fossil fuels, and in major urban areas of developed nations a major source is the exhaust from light-duty motor vehicles (LDMV). Figure 2.7 shows the distribution of sources within the United States in 1996. A total of 89 • 106 short tons, or 81 Tg, of CO were emitted, about 60% of which comes from highway vehicles (EPA, 1997). Natural sources of CO include CO from biornass burning and the oxidation of organics such as methane and isoprene, CO from biological processes in soils, CO from vegetation and termites, and CO from the ocean.

4. Sulfur Compounds Ninety percent or more of the sulfur in fossil fuels is emitted in the form of sulfur dioxide (SO:) during

8% All Other 5%

Fuel Combustion - Other 5% Non-Road Engines and Vehicles 19%

F I G U R E 2.7 Contribution of various sources to the total anthropogenic CO emissions in the United States in 1996 (from EPA, 1997).

combustion, the remainder being primarily in the form of sulfates. Therefore, its emissions can be calculated for a given source with some accuracy from the rate of fuel consumption and the percentage of sulfur in the fuel. For example, Fig. 2.8 shows the SO 2 emissions from a number of countries as a function of total fuel consumption (Kato and Akimoto, 1992). Clearly, there is a relatively linear relationship between the two. In the United States in 1996, a total of 19 • 106 short tons, or 17 Tg, of SO 2 were emitted by anthropogenic sources, with the distribution of sources shown in Fig. 2.9 (EPA, 1997). Figure 2.10 shows the trend in SO 2 emissions from North America, Europe, the USSR, and Asia from 1970 to 1986 (Hameed and Dignon, 1992). While SO 2 emissions from North America have decreased significantly, those from Europe have remained about the same and in the USSR have increased somewhat. However, SO 2 emissions from Asia have increased dramatically due to the increased combustion of fossil fuels, particularly from the use of coal in China and from biomass burning in Southeast Asia (Arndt et al., 1997). For example, Fig. 2.11 shows the 1987 annual emission fluxes of SO 2 in Asia, where China is seen to be a major source (Akimoto and Narita, 1994). Emissions from ships are also thought to contribute significantly (e.g., Streets et al., 1997; Corbett and Fishbeck, 1997). Natural emissions of sulfur compounds to the atmosphere occur from a variety of sources, including volcanic eruptions, sea spray, and a host of biogenic processes (e.g., Aneja, 1990). Most of the volcanic sulfur is emitted as SO 2, with smaller and highly variable amounts of hydrogen sulfide and dimethyl sulfide (CH3SCH3). Sea spray contains sulfate, some of which is carried over land masses.

A. EMISSIONS

21

100 5

10

If"

g

s O

0.1 0

o~ O0

,//'"

0.01

0.001

.........

0.0001 0.01

I

.........

0.1

I

.........

1

1

.........

10

1

.........

100

I

.........

1000

10,000

Fuel C o n s u m p t i o n (106 t o n s yr -1)

F I G U R E 2 . 8 S O 2 emissions in m a n y different regions as a function of rate of fuel c o n s u m p t i o n . D a t a for E u r o p e and the U n i t e d States are for 1980, and t h o s e for Asia are for 1987 (from K a t e and A k i m o t o , 1992).

Biogenic processes, however, emit reduced forms of sulfur, including dimethyl sulfide and hydrogen sulfide, with lesser amounts of carbon disulfide (CS2), dimethyl disulfide (CH3SSCH3), carbonyl sulfide (COS), and methyl mercaptan (CH3SH). These reduced sulfur compounds are then oxidized in the atmosphere as described in detail in Chapter 8.E. One estimate of the global emissions of sulfur com-

, ~ + //-//

+

/ +/ +

/

I

[

!+

I

1+

l+ +

+

2 ~ + + + + + ~ // --.~ ~

ty EIect'ricut'"ili'--ty 67%

5. T o t a l S u s p e n d e d Particles ( T S P ) , P M 1 0 ,

and P M 2 . 5 Air quality standards for particulate matter in the United States were expressed some years ago in terms of the mass of total suspended particulate matter (TSP).

Fuel C o m b u s t i o n Industrial ,,x'~

v"/N 17%

+ + + + + + +\\,,'-//,j\//

.

+ + + + ~ _ ~ -i< : . ~ . +. . + . + ',', ~' ~ I~ ~ A +

+ +

+

+ +

+

+

+

+

+ +

~

+

~

.

.

.

.

.

.

I-ueluomousdon-

.

.

O t h e r 4 Yoo

~

+ ~ : : : : : : : l A l I O t h e r

~ : . : . i i i i l

~++++++++++++ Fuel \ Combustion-~

~

pounds from both anthropogenic and natural sources is given in Table 2.3 (Spiro et al., 1992).

+

~

~

7% N o n - R o a d E n g i n e s a n d V e h i c l e s 2%

+ + + + + + +"~-_/_ + " + " + " + " + " + " + " +7 Metals Processing 3%

"~"Nt- +" + +" + +" + + + +" + +" + /+ "

"~,~+++

++++

/

F I G U R E 2 . 9 C o n t r i b u t i o n of various sources to total a n t h r o p o g e n i c S O 2 emissions in the U n i t e d States in 1996 (from E P A , 1997).

22

2. THE A T M O S P H E R I C SYSTEM

~" t-

45

O

o

40

m ~

,

~

~

0

E :D

/

35

"~

." ."~,sia

~,~orth America

~

oJ

O

O3

30

- Europe

. . . " . ~ : . : .--4,,, ~

~

O C

o

25

..." .~'..-- "

,,:.i

''.~'

-

I

I

USSR

-

O9

E LU

20

,''I

1970

I

I

I

I

I

1975

I

I

1980

I

I

I

I

I

1985

Year

FIGURE 2.10 S O 2 emissions in million tons of equivalent S O 2 for the period 1970to 1986 for Asia, Europe, North America, and the USSR (from Hameed and Dignon, 1992).

The standard was then changed to mass of suspended particulate matter less than 10 /xm in size, commonly called PM10 or PM10, and more recently was modified to include particulate matter less than 2.5/xm in diameter, PM2.5 or PM2. 5. The rationale for basing air quality standards on smaller particles is evident from an examination of Fig. 2.12, a diagram of the human respiratory tract. Larger particles that are inhaled are removed in the head or upper respiratory tract. The respiratory system from the nose through the tracheobronchial region is covered with a layer of mucus that is continuously moved upward by the motion of small hairlike projections called cilia. Large particles deposit on the mucus, are moved up, and are ultimately swallowed. On the other hand, particles from fossil fuel combustion and gas-to-particle conversion are generally much smaller ( < 2.5-/xm diameter) and fall in the respirable size range. These particles can reach the alveolar region where gas exchange occurs. This region is not coated with a protective mucus layer, and here the clearance time for deposited particles is much greater than in the upper respiratory tract; hence the potential for health effects is much greater (Phalen, 1984). Figure 2.13 shows the deposition of particles in various regions of the respiratory tract as a function of particle diameter (Phalen, 1984; Phalen et al., 1991; Yeh et al., 1996). The deposition fraction of PM10 in the pulmonary and tracheobronchial regions can be quite large, so it is not surprising that health effects could be associated with these particles. Deposition in the upper portions of the respiratory system is dominated primarily by the large particles, which are readily taken out in the nose and upper airways.

The deposition curves shown in Fig. 2.13 can be understood in terms of the major mechanisms of removal of particles in the respiratory tract: diffusion, sedimentation, and inertial impaction (see Chapter 9). The smallest particles undergo rapid Brownian diffusion, which carries them to the lung surface where they can be taken up; this is responsible for the large deposition in the pulmonary region seen in Fig. 2.13 for particle sizes below about 0.5 /xm. Gravitational settling, i.e., sedimentation, is also an important mechanism of deposition both in the pulmonary region and in the tracheobronchial region. In both cases, the airways are relatively small so that the particle does not have large distances to travel before reaching a surface. The third mechanism, impaction, occurs when the airstream in which the particle is suspended changes direction due to a bifurcation in the lung, but the inertia of the particle carries it forward to impact the lung surface. There has been great interest in airborne particulate matter recently due to the results of a number of epidemiological studies showing a correlation between increased mortality and levels of airborne particles. Figure 2.14 shows one such correlation reported by Dockery et al. (1993). A clear relationship between mortality rates and the concentration of fine particles PM2. 5, as well as with particle sulfate, is seen. Since sulfate is found primarily in fine particles, these observations are not independent. Schwartz et al. (1996) report a 1.5% increase in total daily mortality with an increase of 10/xg m -3 in PM2. 5. Deaths due to chronic obstructive pulmonary disease increased by 3.3% and those to ischemic heart disease by 2.1%. What is somewhat puzzling, but certainly intriguing, is that the cities included in the studies in Fig. 2.14, as well as more recent ones where these findings have been corroborated, are quite disparate in terms of location and the types of air pollutants that would be expected to predominate in each region, yet a consistent relationship is found. Most such epidemiological studies to date are consistent with this finding. This suggests either that the health effects associated with particles are independent of their chemical composition or that there is some common chemical component. In addition, recent analysis of these studies also finds a correlation with other air pollutants as well and suggests that more than one pollutant may be involved (Lipfert and Wyzga, 1995). This issue is a fascinating one that clearly requires much more research on the formation, chemical composition, and effects of particles and associated air pollutants (e.g., see Phalen and McClellan, 1995; Dockery and Schwartz, 1995; Bascom et al., 1996a,b; Wilson and Suh, 1997; and the review by Vedal, 1997).

A. EMISSIONS

23

% m

[]

f

B3 Rq

__

w

W

ffl ffl Nq:t El

e#

NN

FI 1 (10 -3 moles as S m -2 yr -1) FR3 []10 I~ 20 I~ 50 I

i

200

N--[iI~l~tJ I

El

i,fi,~i~

~

;~:~:~,, i ~w~i i9,~-,44

,.,.,---.,

"

FIGURE 2.11 Patterns of 1987 annual emission flux of S O 2 in Asia (in units of millimoles as S per m2 per year) (from Akimoto and Narita, 1994).

Interestingly, these fine particles not only are of great concern from the point of view of health effects but also are responsible for most of the light scattering,

TABLE 2.3 Estimates of Global Emissions of Sulfur Compounds in 1980 (Tg y r - 1 a s S ) a Source

Sulfur e m i s s i o n s (Tg yr -1 )

Fuel combustion/industrial activities Biomass burning Volcanic eruptions Marine biosphere Terrestrial biosphere Total a Source: Spiro et

al.

(1992).

77.6 2.3 9.6 11.9 0.9 102.2

that is, visibility reduction. Thus an improvement in visibility in areas impacted by air pollution may be accompanied by a reduction in the total particle mass deposition in the alveolar region of the respiratory system as well. In the United States in 1996, the total emissions of PM10 were 31 • 106 short tons per year, or 28 Tg per year (EPA, 1997). Fugitive dust sources such as unpaved roads make up ~ 90% of the total PM m emissions. Figure 2.15 shows sources of PM m in the United States in 1996 split into (a) nonfugitive dust sources ( ~ 10% of the total) and (b) fugitive dust sources (EPA, 1997). Globally, anthropogenic emissions of PM m have been estimated to be 345 Tg yr -~ without including secondary nitrate and organics (Wolf and Hidy, 1997)

24

2. THE ATMOSPHERIC SYSTEM r,,.~o,,pharynx Turbinate~

3harynx

Vestibule

ttis

Sulfate (~tg m-3) 0

2

1.4

I

4

6

I

8

I

I

10 I

12 I

14

r

1.3 (13

Trachea

O

E

Lung

"0

Lung

1.2

.~-, 1.1

Bronchus

"O r "O r

E nducting Bronchiole

CO III

H

L

09

W pT

10 0

I

5

I

10

I

15

I

20

I

25

I

30

35

PM2.5 (Pg m 3)

erminal Bronchiole

F I G U R E 2.14 Estimated adjusted mortality rate ratios, taking the least polluted city, Portage, Wisconsin (P), as 1.0. T = Topeka, Kansas; W = Watertown, Massachusetts; L = St. Louis, Missouri; H = Harriman, Tennessee; S = Steubenville, Ohio. (Adapted from Dockery et al., 1993.)

Respiratory Bronchiole Alveolar Duct Alveolar Sac Mweolus

F I G U R E 2.12 Schematic diagram of human respiratory tract. (From Hinds, W. C. Aerosol Technology. Copyright 9 1982 John Wiley & Sons, Inc. Reprinted by permission of John Wiley & Sons, Inc.)

and are projected to increase by about a factor of two from 1990 to 2040. The composition and sources of particles are discussed in detail in Chapter 9. Major natural sources of particles include terrestrial dust caused by winds, sea spray, biogenic emissions, volcanic eruptions, and wild-

1i 0

0.8

F |

fires. As with anthropogenic particulate emissions, particle size is important. Operationally speaking, terrestrial dust is generally in the size range ~ > 10 /xm in diameter and is primarily composed of crustal elements, including silicon, aluminum, iron, sodium, potassium, calcium, and magnesium. Because particles from mechanical erosion processes tend to be quite large, they have been of less concern from the point of view of health effects since they tend to be removed in the upper respiratory system. Particles are generated at the surface of the ocean by the bursting of bubbles and some of these are

Naso-oro-pharyngo-laryngeal region ronchial region

g

0.6-

g O.4 s 0.2

0.0 0.001

0.01

0.1

1.0

10

Particle diameter (lam)

F I G U R E 2.13 Calculated deposition of particles in various regions of the lung for polydisperse aerosol (O-g = 2.5; see Chapter 9.A.2) (adapted from Yeh et al., 1996).

A. EMISSIONS

25

Sources excluding fugitive dust Non-Road Engines and Vehicles~,,., 15%

Fuel Combustion Other 15%

Other Industrial Processes 13%

Miscellaneous 19%

Fuel Combustion Industrial 7%

All Other 32%

Fugitive dust sources

Paved

Roads 9% All

Other~

Construction

~

~ , ~ URP:%ed

:' !_!.~_! ~ Ag ricultural Crops 17%

Natural

Sources 20% F I G U R E 2.15 Sources of PM10 emissions in the United States in 1996 (a) excluding fugitive dust sources; transportation sources accounted for 21% of the total, and (b) fugitive dust sources (adapted from EPA, 1997).

carried inland. While more than 50% tend to be > 3 /xm, in diameter, recent studies suggest that there are also a number of particles produced in the submicron size range as well (see Chapter 9). The chemical composition of sea-generated particles includes the elements found in seawater (primarily chlorine, sodium, sulfur, magnesium, potassium, and calcium) and organic materials, perhaps including viruses, bacteria, and so on. Biological emissions of particles may occur from plants and trees; additionally, volatile organics such as isoprene and terpenes can react in the air to form small particles. Volcanic eruptions are highly variable but can pro-

duce large amounts of particles. For example, for the St. Augustine (Alaska) eruption in 1976, particulate emissions over the period of 1 year were estimated to be ~ 6 • 106 metric tons (Stith et al., 1978). Wildfires and biomass burning also produce significant particulate matter, most of it in the respirable size range from 0.1 to 1 /xm. Elemental carbon and organics form the majority of these particles, with some minerals also being present.

6. Lead Prior to the introduction of unleaded fuels, gasoline combustion in motor vehicles was by far the greatest

26

2. THE ATMOSPHERIC SYSTEM Chemical and Allied Products 3% .........

B. METEOROLOGY ustion - Other 1% dl Other 4%

Non-Road Engines and Vehicles (Primarily Aircraft) 14%

Metals Processing 51%

.= Disposal 17%

300

lm 240

m

C 0 m 0 r 6O

C (,0

180

0 r

m

m

6O C

.o_

~0 ca0 E 0) "O _J0)

120

m

m

60

m

Clearly, the concentrations of pollutants in ambient air, and hence their impacts, are determined not only by their rates of emissions but also by the nature and efficiencies of their chemical and physical "sinks," e.g., chemical transformations, as well as wet and dry deposition to the earth's surface. To a large extent, these competing processes are affected not only by direct dispersion and transport but also by such meteorological factors as temperature, sunlight intensity, and the presence of temperature inversions as well as clouds and fogs. For example, during the severe air pollution episodes in London (Table 1.1), meteorological conditions were such that the pollutants were effectively contained in relatively small volumes, leading to high pollutant concentrations. Furthermore, fog water droplets provide an aqueous medium for the liquid phase of SO 2 to sulfate conversion. Additionally, the cycle of condensation of water vapor on aerosols at low temperatures, followed by evaporation during the day, is thought to be a major factor controlling the concentrations of pollutants within the droplets (Munger et al., 1983). Indeed, as noted in Chapter 1, fogs with pH values < 2 have been observed in heavily populated coastal cities in southern California (Jacob and Hoffmann, 1983; Hoffmann, 1984; Jacob et al., 1985), and a variety of studies have also shown that fogs in the pH range of 2-3 are common throughout the world. We discuss briefly in the following sections some meteorological parameters that are useful in the context of atmospheric chemistry.

m

0 1960

1. Lapse Rate: Temperature and Altitude

I 1970

1980

1990

Year

F I G U R E 2.16 (a) Contribution of various sources to total anthropogenic Pb emissions in the United States in 1996. (b) Trend in lead emissions in the United States (from EPA, 1995, 1997).

source of atmospheric lead. However, as seen in Fig. 2.16a, this is no longer the case in the United States and most likely in other countries where unleaded gasoline is in widespread use. Metals processing now accounts for more than half of the total. As seen in Fig. 2.16b, emissions of lead have dropped dramatically over the past two decades, due largely to the phase out of its use in gasoline (EPA, 1995, 1997).

As seen in Fig. 1.1, in the lowest 10 km of the earth's atmosphere, the air temperature generally decreases with altitude. The rate of this temperature change with altitude, the "lapse rate," is by definition the negative of the change in temperature with altitude, i.e., - d T / d z . Averaging over time and large geographic regions has shown that within the lowest 10 km of the atmosphere, the lapse rate is usually positive and is typically 6-7 degrees per km. A temperature inversion is said to exist when the lapse rate is negative. The origin of the lapse rate can be understood on the basis of fundamental thermodynamics. That is, under the assumptions of a dry air parcel rising adiabatically in the atmosphere, the temperature is expected to fall about 10 degrees per kilometer increase in altitude. This drop in temperature is defined as a positive lapse rate.

B. METEOROLOGY

27

B O X 2.1

DERIVATION

OF

LAPSE

To understand the origin of this lapse rate, it is convenient to consider first a volume (V) of dry air that is sufficiently large that exchange of molecules across its boundaries is negligible but sufficiently small that there is no significant heat exchange via entrainment/detrainment of large "blobs" of air from the surrounding environment. Air can be treated as an ideal gas and hence follows the ideal gas law: (A)

PV = nRr,

where P is the air pressure, n the number of moles of gas in the air parcel, and T the temperature (in K). The air pressure is due to the weight, i.e., gravitational force, of the column of air above it: P = F/A

= mg/A

= pVg/A,

(B)

where m is the mass of air of density p in a column of air of area A and volume V and g is the acceleration due to gravity (9.8 m s-2). Using the relationship V = A z , where z is the altitude, we may write the variation in pressure with z (expressed as a positive number), as dP/dz

= - d/dz

[ p Vg/A ]

= -d/dz

[p(Az)g/A]

= -pg.

(C)

However, using the ideal gas law and the definition of density, p = ( n / V ) M W . a , where MW.a is the molecular weight of air (29 • 10 -3 kg mol-1), we may rewrite Eq. (C) in the form dP/dz

= - (n/V)MWag

= - (P/RT)MWag.

(D)

This relationship for the variation of pressure with altitude can be converted into that between temperature and altitude as shown in Eq. (N) below for a parcel of dry air that rises without heat exchange occurring between it and the surrounding air; this is known as an adiabatic process. The first law of thermodynamics is dU = dq + dw,

(E)

where U is the internal energy of the gas, q is the heat going in or out of the system, and w is the work performed on or by the gas. Heat and work are defined such that heat going i n t o the system and work done o n the system are positive quantities. By

RATE

IN

TROPOSPHERE

definition, an adiabatic process is one for which d q = 0, so that d U = dw.

The change in internal energy, d U , of an ideal gas is given by (F)

dU = nc v dT,

where c v is the molar heat capacity at constant volume, i.e., the amount of heat required to raise the temperature of the gas I~ if it is heated at constant volume. (Under the latter conditions, any heat going into the gas must go into its internal energy since no work of expansion is done at constant volume.) Hence (G)

dw = nc v dT.

The work done by an expanding gas is defined as From the ideal gas law,

dw = -PdV.

d(PV)

(H)

= d(nRT)

and P dV + VdP = nR dT

(I)

dw = - P dV = VdP - nR dT.

(J)

so that

Combining Eqs. (G) and (J) and rearranging, one obtains" n(c v + R) dT = VdP.

(K)

For an ideal gas, however, ( C p - c v) = R, so that Eq. (K), with application of the ideal gas law for V, becomes nCp d T = ( n R T / P )

dP

(L)

or dT/dP

= RT/cpP.

(M)

Combining Eqs. (D) and (M), one obtains the temperature-altitude profile, or lapse rate, for a dry adiabatic gas" dT/dz

= [dT/dP]/[

dz/dP]

= - MW.a g / C p .

(N)

Substituting the value of MW a and Cp (29 J K -1 mo1-1) for air, the dry adiabatic lapse rate Fd is given by Fd = - d T / d z

= 9.8 • 10 -3 K m-1.

(O)

28

2. THE ATMOSPHERIC SYSTEM

In reality, measured lapse rates are ~ 6-7~ per km. This is due to the fact that atmospheric air is not dry but contains significant amounts of water vapor that also cools as the air parcel rises. When it reaches saturation, it condenses and releases its heat of vaporization, which warms the air somewhat, resulting in a less steep drop in temperature with altitude than expected for a dry air parcel.

2. Potential Temperature A concept that is very useful in relating meteorological conditions to the mixing and transport of air pollutants is that of potential temperature. Potential temperature (0) is defined as the temperature an air parcel of temperature T and pressure P would have if it were expanded or compressed under adiabatic conditions to some reference pressure P0. Adiabatic expansion or compression of an air mass maintains a constant potential temperature. From the definition of entropy, S, as dS = d q r e v / r , these processes are also constant-entropy processes since no heat is exchanged between the system (i.e., the air parcel) and its surroundings. Hence the term "isentropic" is used to describe processes that occur without a change in potential temperature. Potential temperature is a very useful parameter in several ways. First, air pollutants or trace gases within an air parcel having a constant value of 0 can be assumed to be well mixed within that air parcel. Thus only limited numbers of measurements within the parcel are necessary to characterize its composition.

Second, air parcels tend to conserve their potential temperature; i.e., air parcels tend to move along lines of constant 0. Thus the potential temperature becomes a sort of tracer for the history of air parcels. For example, a number of field campaigns have been carried out over the North Atlantic Ocean off the coast of Nova Scotia, Canada. Under some conditions, relatively high levels of 0 3 were observed whose source was unknown. However, in some instances the air parcel containing the higher ozone concentrations could be tracked upwind over northern Canada (Berkowitz et al., 1995). Figure 2.17 shows the isentropic surface (i.e., surface of constant potential temperature) leading back from the sampling point to higher altitudes over the Arctic region. This surface suggests that the source of the higher ozone over the North Atlantic on that particular occasion was upper tropospheric/lower stratospheric air that had traveled from the Arctic regions.

3. Temperature Inversions In a "normal" troposphere that has a positive lapse rate, i.e., where the temperature is falling with altitude, warm air close to the earth's surface, being less dense, rises and is replaced by cooler air from higher elevations. This results in mixing within the troposphere. In some situations, however, the temperature of the air, at some height within the troposphere, may start to rise with increasing altitude before reversing itself again; that is, the lapse rate changes from positive to negative to positive (Fig. 2.18). This region, with a

B O X 2.2

DERIVATION

OF

EXPRESSION TEMPERATURE

Recall (see any standard physical chemistry text) that for adiabatic expansions or compressions of an ideal gas, there are several relationships between P, V, and T that hold; e.g., P V ~ = constant, where y is the ratio of the heat capacities at constant pressure and volume, i.e., y = Cp/C v. Most useful in the context of potential temperature is TP ~/~- 1 = constant. Applying this latter relationship, T p ( r / r - ~) = O p ~ / ~ - ~);

i.e.,

(P)

FOR

POTENTIAL

0 = T ( P / P o ) ~//~-1= T ( P / P o ) cp/R,

(Q)

where c p / R = 3.5 for dry air. Rearranging Eq. (Q) gives T = 0 ( P / P o ) R/ C p = 0 ( P / P o ) 0.286.

(R)

A plot of p0.286 versus temperature for a given 0 is a straight line of constant potential temperature. Such plots are known as pseudoadiabatic charts when plotted with an inverted pressure scale so that pressure increases from top to bottom.

B. METEOROLOGY

q~

E

-

r

-@

"~6 ~

4

~

~o

"40 F I G U R E 2 . 1 7 Isentropic surface at 300 K (i.e., surface of constant potential temperature) over eastern North America on September 8, 1992 (from Berkowitz et al., 1995).

negative lapse rate, is known as an inversion layer. In effect, it acts as a "lid" on an air mass because the cooler air underneath it, being more dense, will not rise through it. In effect, pollutants trapped below the inversion layer are not mixed rapidly throughout the entire troposphere but are confined to the much smaller volume beneath the inversion layer; this generally results in much higher ground-level concentrations of species that are emitted at the surface. The formation of thermal inversions is one of the most important meteorological factors contributing to air pollution problems in urban areas. There are two major sources of thermal inversions. Radiation (or

ground) inversions are caused by the rapid cooling of the earth's surface, along with the layer of air immediately above it, by the emission of infrared radiation immediately after sunset. On calm nights, this cooling may be sufficiently rapid that the layer of air adjacent to the surface becomes cooler than the air above; that is, an inversion forms. This can persist until sufficient heating of the surface and the air above it occurs to "break" the inversion at dawn. With this type of inversion, the inversion height--the distance from the earth's surface to the point at which the lapse rate reverses--is often quite small. For example, in the 1952 London smog episode, inversion heights as low as 150 ft were observed in some locations. Overhead (or subsidence) inversions associated with photochemical air pollution are caused by the sinking motion of air masses as they pass over the continent. This leads to compression and heating of the air immediately below, resulting in a change in the lapse rate, that is, to the formation of an inversion layer. The inversion height is significantly higher than in the case of radiation inversions; for example, ~ 1500 ft would represent a relatively low subsidence inversion height. Interestingly, the vertical distribution of photochemical oxidant may not be such that it falls off rapidly at the inversion layer. In fact, in a classic series of experiments, Edinger and co-workers (1972, 1973) showed that oxidant concentrations in the Los Angeles air basin could be higher within the inversion layer than at ground level. Thus, Fig. 2.19, for example, shows one temperature and oxidant profile for June 20, 1970, over Santa Monica, California, a city adjacent to the Pacific Ocean. Several "layers" of oxidant (mainly 0 3) exist

0

~o"-,_ 4

b

~

29

Normposi al tive

_.~

~L

lnversion~'ayg

1)

50

3. SPECTROSCOPY AND PHOTOCHEMISTRY: FUNDAMENTALS

and combination bands (absorptions corresponding to the sum of two or more of the fundamental vibrations) are much more common. An atmospherically relevant example of just how complex infrared spectra of polyatomic molecules can be is water vapor. A multitude of observed transitions occurs in the regions ~ 1300-2000 cm-~ and 3000-4000 cm -~. This, combined with the relatively high, and often rapidly changing, concentrations of water vapor, makes long-path-length infrared spectroscopic determination of trace atmospheric species in ambient air very difficult in these wavelength regions. Similarly, absorption of CO 2 renders the region from 2230 to ~ 2390 cm -~ unusable for ambient air studies.

2. F a t e s of E l e c t r o n i c a l l y E x c i t e d M o l e c u l e s Once a molecule is excited into an electronically excited state by absorption of a photon, it can undergo a number of different primary processes. Photochemical processes are those in which the excited species dissociates, isomerizes, rearranges, or reacts with another molecule. Photophysical processes include radiative transitions in which the excited molecule emits light in the form of fluorescence or phosphorescence and returns to the ground state and nonradiative transitions in which some or all of the energy of the absorbed photon is ultimately converted to heat. a. Photophysical Processes Photophysical processes are often displayed in the form of the Jablonski-type energy level diagram shown in Fig. 3.9. The common convention is that singlet states are labeled So, S~, S 2, and so on and the triplets are labeled T~, T2, T3, and so on, in order of increasing energy. Vibrational and rotational states are shown as being approximately equally spaced only for clarity of presentation. Radiative transitions, for example, fluorescence (F) and phosphorescence (P), are shown as solid lines, and nonradiative transitions as wavy lines. Vertical distances between the vibrational-rotational levels of the singlet ground state, S 0, and the two electronically excited states, the first excited singlet, S~, and its triplet, T~, correspond to their energy gaps. Fluorescence is defined as the emission of light due to a transition between states of like multiplicity, for example, $1 ~ S o + he. This is an allowed transition, and hence the lifetime of the upper state with respect to fluorescence is usually short, typically 10-6-10 .9 s. For example, the fluorescence lifetime of OH in the electronically excited A 2s § state is ~ 0.7/xs (McDer-

1st excited singlet, S 1

Sv ''~ o '~"

J

=~

0c)1 F

v'= 1 (ISC)I V'=

A So -> $1

i So

So

1st triplet, T1

-f P

A

'=3

Sov'

(ISC)2 ! (

~ ---~T1

1

1 So

So

F I G U R E 3 . 9 Jablonski diagram illustrating photophysical radiative and nonradiative transitions. S o = ground singlet state, S 1 = first excited singlet state, T 1 = first triplet state, A = absorption of light, F = fluorescence, P = phosphorescence, IC = internal conversion, ISC = intersystem crossing. Radiative transitions are shown by solid lines, and nonradiative transitions by wavy lines. Photochemical processes are not indicated.

mid and Laudenslager, 1982; Crosley, 1989). Chemiluminescence is similar to fluorescence except that the excited state is generated in a chemical reaction. Phosphorescence is defined as the emission of light due to a transition between states of different spin multiplicities. Because this is theoretically not an allowed transition for an ideal unperturbed molecule, phosphorescence lifetimes tend to be relatively long, typically 10-3-10 -2 s. Intersystem crossing (ISC) is the intramolecular crossing from one state to another of different multiplicity without the emission of radiation. In Fig. 3.9 (ISC) 1 shows the transfer from the first excited singlet state $1 to the first excited triplet state T~. Since the process is horizontal, the total energy remains the same and the molecule initially is produced in upper vibrational and rotational levels of T~, from which it is deactivated as shown by the vertical wavy line. Similarly, (ISC) 2 shows the intersystem crossing from T~ to upper vibrational and rotational states of the ground state S 0, from which vibrational deactivation to v" = 0 then occurs. Internal conversion (IC) is the intramolecular crossing of an excited molecule from one state to another of the same multiplicity without the emission of radiation. As seen in Fig. 3.9, the horizontal wavy line (IC)~ represents internal conversion from the lowest excited singlet state $1 to high vibrational levels of the ground state So; this is generally followed by vibrational deactivation to v" = 0. Intramolecular photophysical processes available to an excited molecule shown in Fig. 3.9 predominate at

A. BASIC PRINCIPLES

low pressures where collisions with other molecules are relatively infrequent. However, at 1 atm pressure, or in the liquid state, the excited molecule can undergo many collisions with ground-state molecules; this can lead to collisional deactivation of the excited species by several paths. For example, an electronically excited molecule, A*, in the S~'>~ state could undergo a series of collisions, be vibrationally (and rotationally) deactivated, and fall into the S ~'=~ 1 state. The energy lost by A* is carried off as translational energy of the groundstate collision partner, B. From here, A ($1~' = o) can undergo photophysical or photochemical processes. Alternatively, energy transfer from A* to the collision partner can occur in which the excitation energy appears as excess vibrational, rotational, a n d / o r electronic energy of molecule B. Collisional deactivation and energy transfer play important roles in tropospheric chemistry. For example, electronically excited SO 2 in the 3B~ state can be deactivated by 02 (as well as by N 2 and H 2 0 ) to the ground (~A~) state, with part of this process occurring via triplet-triplet energy transfer to generate singlet electronically excited states of 02: 802(381)-1-O2(32~g)--->

8 0 2 ( 1 A 1 ) - 1 - O 2 ( 1 2 ~ ; ' 1Ag ) 9

(1)

51

b. Photochemical Processes

In contrast to the photophysical processes just described, photochemical processes produce new chemical species. Such processes can be characterized by the type of chemistry induced by light absorption: photodissociation, intramolecular rearrangements, photoisomerization, photodimerization, hydrogen atom abstraction, and photosensitized reactions. Of these, photodissociation is by far the most pervasive and important in atmospheric chemistry. For example, the photodissociation of NO 2 into ground-state oxygen atoms, NO2(X2A1) + h u (290 < a < 430 nm) --+ NO(X2[I) + O(3p),

(2)

followed by the reaction of O(3p) with 0 2 , is the sole known source of anthropogenically produced 03 in the troposphere. The reader will encounter numerous other examples of photodissociation throughout this text, so it will not be treated further here. However, as will become obvious in examining the chemistry of both the troposphere and stratosphere in later chapters, it is photochemistry that indeed drives the chemistry of the atmosphere. c. Quantum Yields

Note that in the transfer of electronic energy between an excited atom or molecule and a second atom or molecule, the Wigner spin conservation rule generally applies. This states that the overall spin angular momentum of the system should not change during the energy transfer (see Herzberg for details). Because 0 2 has the unusual property of having a ground triplet state, energy transfer from triplet collision partners can produce the reactive singlet states of molecular oxygen. Indeed this is the mode of action in some photodynamic therapies in medicine. Similarly, collisional deactivation is an important factor in trying to detect and measure various gaseous species in the troposphere using the technique of induced fluorescence. For example, as discussed in Chapter 11, induced fluorescence is one of the techniques applied to determine the concentration of OH free radicals in the troposphere. The OH is excited to the A 2s + state, from which it fluoresces as it returns to the ground state. However, collisional deactivation of excited OH by 0 2 and N 2 is significant at 1 atm pressure; this reduces the emitted light relative to interfering signals. Expansion of the air sample to lower pressures reduces this quenching and increases the overall sensitivity (see Chapter 11).

The relative efficiencies of the various photophysical and photochemical primary processes are described in terms of quantum yields, 4}. The primary quantum yield, qS, for the ith process, either photophysical or photochemical, is given by Eq. (I): Number of excited molecules proceeding by process i (hi =

Total number of photons absorbed

(I)

For example, the nitrate radical, which plays an important role in nighttime chemistry (see Chapter 6), absorbs light in the red region of the visible (600-700 nm). The electronically excited state formed on light absorption can dissociate into either NO 2 + O or into NO + O 2, or it can fluoresce: NO 3 + h u --+ NO~,

(3)

NO~ --+ NO 2 + O,

(4a)

--+ NO + 0 2 ,

(4b)

--+ NO 3 + h u.

(4c)

52


The primary quantum yields for each process are defined as follows: Number of NO 2 or O(3p) formed in the primary process ~4a =

Number of photons absorbed by NO 3

sorbed. Because of the potential contribution of secondary chemistry to the formation of stable products, the overall quantum yield of a particular product may exceed unity. Indeed, in chain reactions, overall quantum yields for some products may be of the order of 106 or more.

Number of NO or 0 2 formed in the primary process ~4b =

Number of photons emitted by NO 3 tJ)4c ~

B. A B S O R P T I O N OF LIGHT

Number of photons absorbed by NO 3 Number of photons absorbed by NO 3

~h4c is also known as the fluorescence quantum yield, ~hf. By definition, the sum of the primary quantum yields for all photochemical and photophysical processes taken together must add up to unity, i.e., ~(~f-Jr- ~p -Jr- ~deact-Jr-... ~a -Jr- ~b -Jr-... ) =

1.0,

where ~hf, ~hp, and ~deact are the primary quantum yields for the photophysical processes of fluorescence, phosphorescence, and collisional deactivation, respectively, and ~ha, ~hb, and so on are the primary quantum yields for the various possible photochemically reactive decomposition paths of the excited molecule. For example, in the case of NO 3, ~4a is, within experimental error, 1.0 up to 585 nm and then decreases to zero at 635 nm. As path (4a) falls off above 585 nm, path (4b), ~4b increases to a peak of 0.36 at approximately 595 nm and then also decreases at longer wavelengths (Orlando et al., 1993; Davis et al., 1993; Johnston et al., 1996). As the quantum yields for both (4a) and (4b) decline, fluorescence, (4c) is observed (Nelson et al., 1983; lshiwata et al., 1983), increasing toward unity at ~ 640 nm. In short, at different wavelengths the contribution of the various processes, (4a), (4b), and (4c), varies, but always consistent with their sum being unity. While the aim of photochemical studies is generally to measure primary quantum yields, this is not always experimentally feasible. For example, NO reacts rapidly with NO 3 to form NO 2. Thus determination of ~4a or ~4b by measuring the NO and NO 2 formed can be complicated by this secondary reaction of NO with NO3, and the measured yields of NO and NO 2 may not reflect the efficiency of the primary photochemical processes. In some cases, then, the overall quantum yield, rather than the p r i m a r y quantum yield, is reported. The overall quantum yield for a particular product A, usually denoted by ~a, is defined as the number of molecules of the product A formed per photon ab-

1. Basic Relationships Light has both wave-like and particle-like properties. As a wave, it is a combination of oscillating electric and magnetic fields perpendicular to each other and to the direction of propagation (Fig. 3.10). The distance between consecutive peaks is the wavelength, A, and the number of complete cycles passing a fixed point in 1 s is the frequency, v. They are inversely proportional through the relationship (J)

A = c/v,

where c is the speed of light in a vacuum, 2.9979 • 108 m s -1 .

Considered as a particle, the energy of a quantum of light E is (K)

E = h v = hc/A,

where h is Planck's constant, 6.6262 • 10 .34 J s per quantum, and the frequency v is in s -1. In the visible and ultraviolet regions of the spectrum, wavelength is commonly expressed in nanometer units, 1 n m = 10 .9 m. In the older literature, units of angstroms, 1 A = 10-1~ m, are also found. In the infrared region both microns [1 micron = 1 micrometer (/xm) = 10 -6 m] and wavenumbers ~o (in cm -~) are employed; ~o is the reciprocal of the wave-

EYI-"

Hz

I I

X

~

\

~I

_

/

I

Direction of

~

\

/

i

/i (Distance)

F I G U R E 3.10 The instantaneous electric (Ey) and magnetic (H z) field strength vectors of a plane-polarized light wave as a function of position along the a~s of propagation (x) (from Calvert and Pitts,

1966).

B. ABSORPTION OF LIGHT

length A expressed in centimeters. It is directly related to energy through the Planck relationship, E = hc o~,

kcal einstein -1), are sufficient to break chemical bonds ranging from, for example, the weak O 2 m O bond in ozone, ~ 100 kJ mo1-1 ( ~ 25 kcal mol-1), to the moderately strong C m H bond in formaldehyde, ~ 368 kJ mol- 1 ( ~ 88 kcal mol- 1). Other spectral regions are also important because the detection and quantification of small concentrations of labile molecular, free radical, and atomic species of tropospheric interest both in laboratory studies and in ambient air are based on a variety of spectroscopic techniques that cover a wide range of the electromagnetic spectrum. For example, the relevant region for infrared spectroscopy of stable molecules is generally from ~ 500 to 4000 cm -1 (20-2.5 /xm), whereas the detection of atoms and free radicals by resonance fluorescence employs radiation down to 121.6 nm, the Lyman c~ line of the H atom. Table 3.3 gives some relationships between commonly used energy units. Today the SI system of units is in general use, although much of the data in the literature is in the older units. Thus we use both types of units for energy, that is calories or kilocalories and joules or kilojoules, where 1 cal = 4.184 J.

(L)

and today is generally the unit of choice in infrared spectroscopy. Since chemists often deal experimentally with moles rather than molecules, a convenient unit is a mole of quanta, defined as 1 einstein. The energy of 1 einstein of light of wavelength A in nm is E = (6.02 • 1023)hv = 6.02 • 1 0 2 3 h c / A , = 1.196 • 105/A kJ einstein -1, = 2.859 • 104/A kcal einstein-1.

(M)

Another unit used in photochemistry to express the energy of a quantum of radiation is the electron volt; 1 eV = 96.49 kJ mo1-1 = 23.06 kcal mo1-1. Thus for A in nm E = hc/A

= 1.240 • 103/A eV.

(N)

To put these energies and wavelengths in perspective, Table 3.2 gives some typical wavelengths, frequencies, wavenumbers, and energies of various regions of the electromagnetic spectrum. The region of most direct interest in tropospheric photochemistry ranges from the visible at ~ 700 nm to the near-ultraviolet at 290 nm, the short-wavelength cutoff of the stratospheric ozone layer. The corresponding energies [Eq. (M)], 170.9 and 412.4 kJ einstein -1 (or 40.8 and 98.6

TABLE 3.2

Name Radiowave Microwave Far-infrared Near-infrared Visible Red Orange Yellow Green Blue Violet Near-ultraviolet Vacuum ultraviolet X-Ray y-Ray

53

2. T h e B e e r - L a m b e r t L a w The basis for the measurement of the strength of light absorption by a molecule at various wavelengths is shown in Fig. 3.11. A parallel monochromatic light beam of wavelength a and power P0 or intensity I0, defined as the energy per second striking a unit area,

T y p i c a l W a v e l e n g t h s , F r e q u e n c i e s , W a v e n u m b e r s , a n d E n e r g i e s of V a r i o u s R e g i o n s of t h e E l e c t r o m a g n e t i c S p e c t r u m

Typical wavelength or range of wavelengths (nm) ,-~ 108-1013 ,-~ 107-108 105-107 ,-~ 103-105 700 620 580 530 470 420 400-200 -~ 200-50 -~ 50-0.1 < 0.1

Typical range of frequencies v (s -1 ) ,-~ 3 • ,-~ 3 • ~ 3 • ,-~ 3 •

104-3 109-3 101~ 1012-3

• • • •

109 10 l~ 1012 10 TM

4.3 • 1014 4.8 • 1014 5.2 • 1014 5.7 • 1014 6.4 • 1014 7.1 • 1014 (7.5-15.0) • 1014 (1.5-6.0) • 1015 -~ (0.6-300) • 1 0 1 6 -~ 3 • 10 is

a For kcal einstein- 1, divide by 4.184 (1 cal = 4.184 J).

Typical range of wavenumbers co ( c m - l)

Typical range of energies (kJ e i n s t e i n - 1)a

10-6-0.1 0.1-1 1-100 102-104

,-~ 1 0 - 3 - 1 0 .8 ~ 1 0 - 2 - 1 0 .3 ~ 10-2-1 "-~ 1-102

1.4 • 104 1.6 • 104 1.7 • 104 1.9 • 104 2.1 • 104 2.4 • 104 (2.5-5) • 104 (5-20) • 104 (0.2-100) • 106 > 10 s

1.7 • 102 1.9 • 102 2.1 • 102 2.3 • 102 2.5 • 102 2.8 • 102 (3.0-6.0) • 102 ,-~ (6.0-24) • 102 -~ 103-106 > 106

54

3. SPECTROSCOPY A N D PHOTOCHEMISTRY: FUNDAMENTALS

of the absorbing substance, this integrates to

TABLE 3.3 Some Relationships between Commonly Used Energy Units

l n ( 1 / 1 o) = - kC1. (kJ m o l - 1) x 0.2390 = kcal m o l - 1

(P)

The most commonly used form of this B e e r - L a m b e r t law involves logarithms to the base 10:

x 0.0104 = e V • 83.59 = cm - 1

l o g ( I o / I ) = ~C1.

(Q)

(kcal m o l - 1) • 4.184 = kJ m o l - 1 • 0.04336 = e V • 349.8 = cm - 1 (cm - 1 ) • • 10 -2 = kJ m o l - 1 • • 10 -3 = kcal m o l - 1 • 1.240 • 10 -4 = e V

(eV) • 96.49 = kJ m o l - 1 • 23.06 = kcal m o l - 1 • 8.064 • 103 = c m - 1

passes through a sample of length 1 (cm) and concentration C (mol L - l ) . If this wavelength is absorbed by the sample, the power of the beam exiting the sample is reduced to P (or I). The transmittance (T) is defined as I / I o (or P/Po). As the light passes through a thickness of sample dl, it undergoes a fractional reduction in intensity proportional to the absorbing path length, i.e., ( d I / I (x ) - ( d l ) ,

(0)

where the negative sign reflects the reduction in intensity with an increase in path length. Since the constant of proportionality must involve the concentration (C)

--

l

I

>

i0 (Po)

(P)

C is in units of mol L-1, 1 is in cm, and the constant of proportionality ~ (L mo1-1 cm -1) is known as the molar absorptivity or molar extinction coefficient. The dimensionless quantity l o g ( I o / I ) is known as the absorbance, A, which is related to the transmittance by A = l o g ( I o / I ) = - l o g T. Most commercial spectrometers report absorbance, as defined in Eq. (Q), versus wavelength. This is very important to recognize, since as we will see later, calculations of the rate of light absorption in the atmosphere require the use of absorption coefficients to the base e rather than to the base 10. While the recent atmospheric chemistry literature reports absorption cross sections to the base e, most measurements of absorption coefficients reported in the general chemical literature are to the base 10. If these are to be used in calculating photolysis rates in the atmosphere, the factor of 2.303 must be taken into account. In gas-phase tropospheric chemistry, the most cornmon units for concentration, N, are molecules cm -3 and for path length, l, units of cm. The form of the B e e r - L a m b e r t law is then

[c] FIGURE 3.11 S c h e m a t i c d i a g r a m of e x p e r i m e n t a l a p p r o a c h to the B e e r - L a m b e r t law.

ln(Io/I)

=

oN1

(R)

or

I / I o = e x p ( - o-N1),

(S)

where it must again be emphasized that or, known as the absorption cross section, must have been measured with the appropriate form of the B e e r - L a m b e r t law to the base e. The dimensionless exponent o-N1 is often referred to as the "optical depth." In the past, some absorption coefficients for gases have been reported with concentrations in units of atmospheres, so that the absorption coefficient is in units of atm -1 cm -1. Since the pressure depends on temperature, the latter (usually 273 or 298 K ) w e r e also reported. For most tropospheric situations involving gaseous species, the B e e r - L a m b e r t law is an accurate method for treating light absorption; similar considerations apply to nonassociated molecules in dilute solution. However, under laboratory conditions with relatively high concentrations of the absorbing species, deviations may arise from a variety of factors, including concentration- and temperature-dependent association or dissociation reactions, deviations from the ideal gas law, and saturation of very narrow lines with increasing

C. ATMOSPHERIC PHOTOCHEMISTRY

55

TABLE 3.4 C o n v e r s i o n Factors for C h a n g i n g A b s o r p t i o n C o e f f i c i e n t s from One Set of U n i t s to A n o t h e r

Both units in either logarithmic base e or base 10 • 2.69 • 1019 = (atm at 273 K ) - 1 ( c m - 1 ) (cm 2 m o l e c u l e - 1) • 2.46 • 1019 = (atm at 298 K ) - 1 ( c m - 1) • 3.24 • 1016 = (Torr at 298 K ) - 1 ( c m - 1) • 6.02 • 102o = L m o l - 1 c m - 1 (atm at 298 K) -1 (cm - 1 ) • 4.06 • 10 .2o = cm 2 molecule -1 • 1.09 = (atm at 273 K ) - 1 ( c m - 1 ) • 4.46 • 10 .2 = (atm at 273 K ) - 1 ( c m - 1 ) (L m o l - 1 c m - 1 ) • 4.09 • 10-2 = (atm at 298 K ) - 1 ( c m - 1 ) • 5.38 • 10 -5 = (Torr at 298 K ) - 1 ( c m - 1) • 1.66 • 10 -21 = cm 2 m o l e c u l e - 1

Change of both logarithmic base and units

(cm 2 m o l e c u l e - 1 ) , base e

• • • •

1.17 1.07 1.41 2.62

• • • •

1019 1019 1016 102o

= = = =

(atm at 273 K ) - 1 ( c m - 1), base 10 (atm at 298 K ) - 1 ( c m - 1), base 10 (Torr at 298 K ) - 1 ( c m - 1), base 10 L m o l - 1 c m - 1, base 10

• 3.82 • 10 -21 = cm 2 m o l e c u l e - 1 base e (L m o l - 1 c m - 1), base 10 • 0.103 = (atm at 273 K ) - 1 ( c m - 1), base e • 9.42 • 10 -2 = (atm at 298 K ) - 1 ( c m - 1), base e (atm at 273 K) -1 ( c m - 1 ) , base 10

• 8.57 • 10 .20 = cm 2 m o l e c u l e - 1 , base e • 51.6 = L m o l - 1 c m - 1 base e • 7.11 • 10-17 cm 2 m o l e c u l e - 1 base e • 4.28 • 104 = L m o l - 1 c m - 1 base e =

(Torr at 298 K ) - 1 ( c m - 1), base 10 (atm at 298 K) -1 ( c m - 1 ) , base 10

concentrations, i.e., increasing pressures. Particularly important is the situation in which a "monochromatic" analyzer beam actually has a bandwidth that is broad relative to very narrow lines of an absorbing species. In this case, which is often encountered in the infrared, for example, the B e e r - L a m b e r t law is nonlinear. Clearly, to be on the safe side, it is good practice to verify the linearity of ln(Io/I) plots as a function of absorber concentration when experimentally determining absorption coefficients. Table 3.4 gives conversion factors for converting absorption coefficients from one set of units to another and for changing between logarithms to the base 10 and base e.

C. ATMOSPHERIC PHOTOCHEMISTRY 1. Solar Radiation and Its Transmission through the Atmosphere a. The Sun and Its Relationship to the Earth: Some Important Definitions for Atmospheric Chemistry The sun can be considered a spherical light source of diameter 1.4 • 10 6 km located 1.5 • 10 8 km from

• 9.35 • 10 .20 = cm 2 m o l e c u l e - 1 , base e • 2.51 = (atm at 273 K ) - 1 ( c m - 1), base e

the earth's surface. Incoming direct sunlight at the earth's surface is treated as a beam with an angle of collimation of ~ 0.5 ~ and thus is essentially parallel to _+0.25~. The total intensity of sunlight outside the earth's atmosphere is characterized by the solar constant, defined as the total amount of light received per unit area normal to the direction of propagation of the light; the mean value is 1368 W m -2, although variations from this mean are common (Lean, 1991). Of more direct interest for atmospheric photochemistry is the solar flux per unit interval of wavelength. Values up to approximately 400 nm are provided by Atlas 3 (see Web site in Appendix IV) and from 400 nm on by Neckel and Labs (1984). Figure 3.12 shows the solar flux as a function of wavelength outside the atmosphere and at sea level for a solar zenith angle of 0 ~ (Howard et al., 1960). Outside the atmosphere, the solar flux approximates blackbody emission at ~ 5770 K. However, light absorption or scattering by atmospheric constituents modifies the spectral distribution. The attenuation due to the presence of various naturally occurring atmospheric constituents is shown by the hatched areas in Fig. 3.12.

56


2.5 ,_,

2.0

~'E

1.5

i

1.0

~o

_

[~ !| I I~ I" I Ii

-

_

\

/ Energy curve for blackbody at 6000~ ): / /Solar irradiance curve outside atmosphere "//// ~/Solar irradiance curve at sea level for a zenith angle of 0~ ~ O~ / z / H20 H20

!l

~

t/

~'~" ,~ ~ ' ~

H20 H20

i I

0.5 _

H20 0

0

200 400 600 800 1000

1600

2000

H O 2600

3000

Wavelength (nm) F I G U R E 3.12 Solar flux outside the atmosphere and at sea level, respectively. The emission of a blackbody at 6000 K is also shown for comparison. The species responsible for light absorption in the various regions (03, H20, etc.) are also shown (from Howard et al., 1960).

Figure 3.13 shows the altitude corresponding to maximum light absorption by atomic and molecular oxygen and nitrogen and by 0 3 as a function of wavelength up to /~ = 300 nm with the sun directly overhead (Friedman, 1960). Because of the presence of these absorbing species in the upper atmosphere, only light of /~ > 290 nm is available for photochemical reactions in the troposphere. It is often expressed as the integrated radiation coming from all directions to a sphere and is referred to as actinic radiation, although in the strictest sense,

g

"actinic" means "capable of causing photochemical reactions." The ultraviolet region, A < 400 nm, is often divided into what is known as the UV-A region from 315 to 400 nm, the UV-B region from 280 to 315 nm, and the UV-C region from 200 to 280 nm. The effect of light scattering and absorption by atmospheric constituents on the intensity and wavelength distribution of sunlight at the earth's surface depends on both the nature and concentration of the gases and particles as well as the path length through

160

"-" 140

8

@:: 120 100 9

80

>

60 -_ _< 40 -

o "1:3

.g

m

by atomic and _ molecular > oxygen and nitrogen

\ _. . \ Absorption

20 0

20

40

60 80 100 120 140 160 180 200 220 240 260 280 300 Wavelength (nm)

F I G U R E 3.13 Approximate regions of maximum light absorption of solar radiation in the atmosphere by various atomic and molecular species as a function of altitude and wavelength with the sun overhead (from Friedman, 1960).

57


which the light passes, as expected from the B e e r Lambert law. The path length, that is, the distance from the outer reaches of the atmosphere to an observer on the earth's surface, is a function of the angle of the sun and hence time of day, latitude, and season. In addition, reflection of light from the earth's surface alters the light intensity at any given point in the atmosphere, as does the presence of clouds. The angle of the sun relative to a fixed point on the surface of the earth is characterized by the solar zenith angle 0, defined, as shown in Fig. 3.14, as the angle between the direction of the sun and the vertical. Thus a zenith angle of zero corresponds to an overhead, noonday sun, and a zenith angle of ~ 90 ~ approximates sunrise and sunset. The greater the zenith angle, the longer is the path length through the atmosphere and hence the greater the reduction in solar intensity by absorption and scattering processes. The path length L for direct solar radiation traveling through the earth's atmosphere to a fixed point on the earth's surface can be estimated geometrically using Fig. 3.14. This "flat earth" approximation is accurate for zenith angles < 60 ~ One can approximate L using cos 0 ~

(T)

h/L

or

L ~h/cos0~h

(u)

sec0.

A common term used to express the path length traversed by solar radiation to reach the earth's surface is the air mass, m, defined as Length of path of direct solar radiation through the atmosphere m

=

Length of vertical path through the atmosphere

.

(v)

Earth FIGURE 3.14 earth's surface.

Definition of solar zenith angle 0 at a point on the

in spectral distribution by absorption and scattering by gases as well as by particulate matter. As a result, the actual actinic flux to which a given volume of air is exposed is affected by the zenith angle (i.e., time of day, latitude, and season), by the extent of surface reflections, and by the presence of clouds. Madronich (1993) discusses these variables, with particular emphasis on the effects on the UV reaching the earth's surface. To estimate the solar flux available for photochemistry in the troposphere then, one needs to know not only the flux outside the atmosphere but also the extent of light absorption and scattering within the atmosphere. We discuss here the actinic flux F(A) at the earth's surface; the effects of elevation and of height above the surface are discussed in Sections C.2.d and C.2.e. The reduction in solar intensity due to scattering and absorption can be estimated using a form of the B e e r - L a m b e r t law: I/I o = e -tin.

With reference to Fig. 3.14, for zenith angles less than 60 ~, m ~- L / h

~-

sec 0.

(W)

At larger angles, corrections for curvature of the atmosphere and refraction must be made to L and m. Table 3.5 shows values of the air mass at various zenith angles 0, either estimated using m = sec 0 or corrected for curvature of the atmosphere and for refraction; it is seen that only for 0 > 60 ~ does this correction become significant.

b. Solar Spectral Distribution and Intensity in the Troposphere When the radiation from the sun passes through the earth's atmosphere, it is modified both in intensity and

(X)

In Eq. (X), I 0 is the light intensity at a given wavelength incident at the top of the atmosphere and I is the intensity of the light transmitted to the earth's surface; t is the total attenuation coefficient described below and m is the air mass as defined earlier. For the sun directly overhead (i.e., zenith angle 0 = 0) the air mass is unity (m = 1.0); the attenuation coefficient then reflects the minimum possible attenuation by the atmosphere. As 0 increases until the sun is on the horizon (i.e., sunset or sunrise), m also increases (Table 3.5); thus the attenuation of the sunlight increases due to the increased path length in the atmosphere through which the light must travel to reach the earth's surface. The attenuation coefficient, t, represents a combination of light scattering and absorption by gases and

58


B O X 3.1

CALCULATION

OF

SOLAR

The solar zenith angle can be calculated in the following manner for any particular location (i.e., latitude and longitude), day of the year (d,), and time of day as described by Spencer (1971) and Madronich (1993). First, one needs to calculate what is known as the local hour angle ( t h ) , which is defined as the angle (in radians) between the meridian of the observer and that of the sun: t h (in r a d i a n s ) = I r [ ( G M T / 1 2 ) -

1

where GMT is Greenwich mean time converted from the local time, longitudes (in degrees)west of the Greenwich meridian are negative, and E Q T is the "equation of time," given by

and place is the solar declination, 6, defined as the angle between the direction of the sun and the equatorial plane of the earth. The value of 6, which is 0 ~ at the spring and fall equinoxes and falls between + 23.45 ~ (June 21) and - 2 3 . 4 5 ~ (December 21), can be calculated from the following: 6 (in radians) = 6.918 • 10 .3 - 0.399912cos N 6.758 • 10 .3

• cos 2 N + 9.07 • 10 .4 • sin 2 N - 2.697 • 10 .3 cos 3 N + 1.480 • 10 .3 sin 3N. The solar zenith angle (0) for that particular time and place is then determined from: cos 0 = sin 6 sin(latitude)

E Q T = 7.5 • 10 -5 + 1.868 • 10 .3 cos N

+ cos 6 cos(latitude)cos th,

- 3.2077 • 10- 2 sin N - 1.4615 X 10 .2 c o s 2 N 4.0849 • 10 .2 sin 2N.

where N is defined as N (in radians) = 21rdn/365. The day of the year, dn, is defined as the day number (0-364), with 0 corresponding to January 1 and 364 to December 31. The second derived parameter that is needed for calculating the solar zenith angle at a particular time

particles and is actually a sum of four terms, t = tsg + tag + tsp + tap ,

ANGLE

+ 0.070257 sin N -

+ (longitude/180)] + EQT,

-

ZENITH

(Y)

where sg = light scattering by gases, ag = light absorption by gases, sp = light scattering by particles, and ap = light absorption by particles. Gases scatter light by molecular, or Rayleigh, scattering. The intensity, I(h, | of light of wavelength h scattered at an angle 0 to the direction of incident light is determined by a number of factors. These include the incident light intensity, the angle | the distance from the scattering molecule, and the index of refraction and size of the scattering molecule. In addition,

where 6 and t h are calculated as already described and latitudes north of the equator (expressed in radians) are positive and south are negative. If all of the input parameters are in radians, 0 is also obtained in radians and can be converted to degrees using 1 rad = 57.296 ~ For example, at Los Angeles, California (latitude = 34.03~ longitude = 118.14~ on September 21 at noon PST, GMT = 20.0, N = 4.53, E Q T = 0.0301, t h = 0.0626, 6 = 0.0179 rad, and cos 0 = 0.837, giving a solar zenith angle of 0.579 rad, or 33 ~

and most importantly, Rayleigh scattering varies inversely with the fourth power of the wavelength. Making the simplifying assumptions of a homogeneous atmosphere of fixed height of 7.996 • 105 cm and of uniform temperature and pressure throughout, Rayleigh scattering can be simplified for application to the atmosphere; as discussed in detail by Leighton (1961), the attenuation coefficient for scattering by gases, tsg, becomes tsg = 1.044 X 1 0 5 ( n 0 a - 1 ) 2 / A 4,

(Z)

where n0a is the index of refraction of air at wavelength h and the pressure and temperature of interest.

59


TABLE 3.5 Values of the Air Mass m at the Earth's Surface for Various Zenith Angles: (a) Calculated from m = sec 0 and (b) Corrected for Atmospheric Curvature and for Refraction Zenith angle 0 (deg)

m = see 0

Air mass (m)

0 10 20 30 40 50 60 70 78 86

1.00 1.02 1.06 1.15 1.31 1.56 2.00 2.92 4.81 14.3

1.00 1.02 1.06 1.15 1.31 1.56 2.00 2.90 4.72 12.4

In the atmosphere, light absorption in the ultraviolet region is predominantly due to 0 3 and this is predominantly in the stratosphere (Figs. 3.12 and 3.13). Since the absorption coefficients (or) of 0 3 are reasonably well established, a variant of the B e e r - L a m b e r t law can be applied to determine how much of the incident light is absorbed by 03: I --

= e -~Am.

(AA)

Io A is the effective column 0 3 (molecules cm-2), o- its absorption cross section at that wavelength, and rn the air mass. One needs to know, in addition to or, the 0 3 concentration as a function of altitude (z), that is,

Source: Demerjian et al. (1980). oo

A = f The dependence of Rayleigh scattering o n /~-4 is evident in Fig. 3.15, which shows the attenuation coefficient for Rayleigh scattering as a function of wavelength from 290 to 700 nm; shorter wavelengths (i.e., in the blue ultraviolet region) are scattered much more strongly than the longer wavelengths.

!

1.0

m

Aerosol extinction m

72. . . . .

gT~

0.1 m

I i

~

Rayleigh "~ttermg

m

Ozone _)~t 0.01 200

/

\

abs~176 I / 1

300

O3(z) dz. z=0

Using the published absorption coefficients (or) as a function of wavelength, one can then apply the B e e r - L a m b e r t law to calculate the intensity of light transmitted through such a vertical column to the earth's surface or to an altitude z. The resulting attenuation coefficients for 0 3 are shown in Fig. 3.15 for an overhead sun. Clearly, 0 3 is responsible for most of the attenuation of light directly from the sun of a < 310 nm reaching the earth's surface. This region of the spectrum around 300 nm is a crucial one for tropospheric photochemistry in both clean and polluted atmospheres. As we have indicated earlier, it is here that species such as ozone and aldehydes photolyze to produce atoms and free radicals critical to the chemistry of the troposphere. Scattering and absorption of light by particulate matter are much more complex and will not be treated in detail here. Clearly, the size distribution and chemical composition, as well as the concentration of the particles, are very important in determining the extent of light scattering and absorption. Since these parameters will vary significantly geographically, seasonally, and diurnally, accurately estimating their impact on light intensities at a particular location at the earth's surface is difficult. Simplifications for the attenuation coefficient for scattering by particles such as tsp = b / A n

\1

(BB)

[

400

500

600

700

800

Wavelength (nm) F I G U R E 3.15 Attenuation coefficients (t) for light scattering (Rayleigh scattering) and absorption (ozone absorption) by gases and for scattering and scattering plus absorption (aerosol extinction) by particles [from Peterson (1976) and Demerjian et al. (1980)].

are often made, where b depends on the concentration of particles and n on their size; for example, n decreases from ~ 4 to 0 as the particle size increases (Leighton, 1961). One estimate of the attenuation coefficients for light scattering by particles, tsp , is also given in Fig. 3.15

60

3. SPECTROSCOPYAND PHOTOCHEMISTRY:FUNDAMENTALS

(Demerjian et al., 1980). Also shown are these researchers' estimates of total scattering plus absorption due to particulate matter, known as the aerosol extinction:

the size distribution, concentration, and composition of particles, which to a large extent cause this scattering, are highly variable geographically and temporally and are not always well known for a particular point in space and time. The amount of light reflected from the earth's surface to a volume of air clearly depends on the type of surface, as well as the wavelength of light; thus snow is highly reflecting, whereas black lava rock reflects very little of the incident radiation. The term used to describe the extent of this reflection is the surface albedo, which is the fraction of light incident on the surface that is reflected. Reflection can be specular, in which the angles of incidence and reflection are equal (e.g., a water surface at large zenith angles), or diffuse, in which light is reflected equally in all directions regardless of the angle of incidence (e.g., white rocks or buildings); the latter is known as "Lambertian" reflection. Table 3.6 gives some reported values of surface albedos for different types of surfaces. It should be noted that, as expected, albedos are wavelength dependent (e.g., see McLinden et al. (1997) for wavelength dependence of ocean albedos and Herman and Celarier (1997) for albedos in the U V from 340 to 380 nm). One can thus estimate the total light intensity incident on a given volume of air in the troposphere due to direct solar radiation, scattering, and reflection. The light absorbed in that volume can then be calculated

tsp + tap = Aerosol extinction. In this case, the radii of the particles were assumed to fall between 0.01 and 2.0 /~m; the peak in the number versus size distribution was at 0.07/~m. Given estimated values for the attenuation coefficients for scattering and absorption of light by gases and particles (i.e., tsg , tag , tsp , and tap), one can calculate from Eq. (X) the fraction of the direct solar intensity incident on the top of the atmosphere that is transmitted to the earth's surface at any given wavelength. However, when one considers the actual light intensity that reaches a given volume of gas in the troposphere, one must take into account not only this direct solar radiation but also two other sources of indirect light: (1) light, either from the sun or reflected from the earth's surface, that is scattered to the volume by gases or particles, known as diffuse solar radiation or sky radiation, and (2) light that is reflected from the earth's surface. These are illustrated in Fig. 3.16. Estimating the intensity of the scattered light at a given point in the atmosphere is difficult because of the substantial uncertainties and variability involved in the factors that contribute to light scattering; for example,

\

Scattered sky radiation Direct solar radiation Reflected radiation reflected radiation

" Scattering by gases and particles

FIGURE 3.16 Different sources of radiation striking a volume of gas in the atmosphere. These sources are direction radiation from the sun, radiation scattered by gases and particles, and radiation reflected from the earth's surface.


61

TABLE 3.6 Some Typical Albedos for Various Types of Surfaces Type of surface

Albedo

Reference

Snow

0.69 0.93 b 0.9-1.0

Angle et al., 1992 Dickerson et al., 1982 Junkermann, 1994

Ocean

0.07 b 0.06-0.08 a

Dickerson et al., 1982 Eck et al., 1987

Forests

0.06-0.18 b 0.02 a 0.17

Dickerson et al., 1982 Eck et al., 1987 Angle et al., 1992

Fields and meadows

0.03-0.04 a

Eck et al., 1987

Desert

0.06-0.09 a

Eck et al., 1987

Salt flats

0.57-0.65 a

Eck et al., 1987

FIGURE 3.17 Typical device (Eppley Laboratories Model 8-48) used to measure solar irradiance. The detector consists of a differential thermopile with the hot junction receivers blackened with flat black coating and the cold junction receivers whitened with BaSO4 (photo supplied courtesy of G. L. Kirk, Eppley Laboratories).

a Minimum reflectivities at 370 nm. b M e a s u r e d with respect to N O 2 photolysis.

using the B e e r - L a m b e r t law, if the concentrations and absorption coefficients of all absorbing species are known.

2. Calculating Photolysis Rates in the Atmosphere a. Photolysis Rate Constant [ kp ], Radiance [ L ( h ) ], Actinic Flux [ F ( h ) ], and Irradiance [ E( h ) ] The rate of photodissociation of a molecule, A, upon light absorption, A+hv-~B+C,

(5)

can be described as a first-order process (see Chapter 5) with a rate constant, kp, known as the photolysis rate constant:

d[A]/dt = - k p [ A ] .

(cc)

shown in Fig. 3.17. Here light of a given wavelength that comes from all directions in a hemisphere and crosses the surface is measured. The net flux crossing the surface per unit area and time is known as the irradiance, E(A), and represents the flow of light across a fiat plane rather than the total light coming from all directions that a molecule actually encounters in the atmosphere. Although actinic flux and irradiance are clearly related, they are not identical; for a detailed treatment of the actinic flux, irradiance, and radiance (defined next), see Madronich (1987). An expression for the number of A molecules dissociating per unit volume per unit time is developed in Box 3.2. Comparing Eq. (NN) in Box 3.2 to Eq. (CC), the photolysis rate c o n s t a n t kp must be given by kp = f q ~ ( a ) o - ( a ) / L ( a , @ , "

In effect, k p t a k e s into account the intensity of available light that the molecule can absorb, the intrinsic strength of light absorption in that region by A, i.e., the absorption cross section o-, and the quantum yield for photodissociation, q~. The light available to a molecule in air for absorption and photodissociation includes both direct and scattered and reflected radiation coming from all directions as described earlier and depicted in Fig. 3.16. The term actinic flux or spherically integrated actinic flux, denoted by F(A), is used to describe the total intensity of this light and is the quantity of interest in calculating k p . However, in practice, available light intensity is often measured using fiat-plate devices such as the one

,

4

= s162

-

-

q~)do~da

(_o

da,

(OO)

where F(a) is the spherically integrated actinic flux, Eq. (FF). We again stress that o-(a) is the absorption cross section to the base e, arising from the use of the differential form of the B e e r - L a m b e r t law to obtain Eq. (KK). There are several approaches to measuring actinic fluxes and photolysis rate constants. One approach is to measure the rate of decay of a species such as NO 2 directly, so-called "chemical actinometry" (e.g. see Madronich et al., 1983). Another approach is to measure the light intensity and convert this to an actinic flux.

62


B O X 3.2

RELATIONSHIPS ACTINIC FLUX,

BETWEEN RADIANCE, AND PHOTOLYSIS RATE

To calculate kp, let us take the case shown in Fig. 3.18a of light striking the top of a very thin layer of air of thickness dz. The light originates in a solid angle d~o and strikes the top surface of the thin layer at an angle 0 to the vertical. The intensity of the incoming light at angle 0 shown in Fig. 3.18a is known as the radiance, L(A, 0, ~b). By definition, radiance is the number of photons (or energy) in the wavelength interval d h originating from a small solid angle d~o and striking a small surface area da in time dt at an angle 0 to the vertical. However, the net flux across the surface is, in effect, determined by the portion of the surface that is perpendicular to the incoming light beam. As seen in Fig. 3.18b, from simple geometric considerations

IRRADIANCE, CONSTANTS

the portion of the surface perpendicular to the incoming beam is 2x = (da)cos 0. Thus the net photons (or energy), dP in a wavelength interval d A that originates in a solid angle d~o and at an angle 0 to the normal and crosses a small surface area, da, in time dt is given by dP = L( A, O, 4))cos 0 d a d ~o dt d A.

(DD)

The irradiance E(A), which is directly measured by fiat-plate devices, is by definition the total number of photons per unit surface area, time, and wavelength. Thus

E(A) = f (dP/dadAdt) = f L(A, O,

0

(EE) The actinic flux F(A) is the total incident light intensity integrated over all solid angles, given by

do.)

0

F(A) =

,

d~o.

(FF)

Thus the irradiance, E(A), and the actinic flux, F(A), differ by the factor cos 0. Only for 0 = 0 ~ i.e., for a parallel beam of light perpendicular to the surface, are the irradiance and flux equal. We now need to convert d~o into terms involving the spherical coordinates 0 and (h. As shown in Fig. 3.19a, a given solid angle ~o traces an area a on the surface of a sphere of radius r. When a = r 2, the solid angle ~o is by definition 1 sr (st = steradian). For the more general case of a surface area a subtended by the solid angle ~o,

\,a '}' ......

f (a, 0,

---.

dz

~o (steradians) = a / r 2

L or

d ~o = d a / r 2 .

(GG)

As shown in Fig. 3.19b, for small changes in the angles 0 and ~h, there is a change in the surface area, da, on a sphere of radius r given by da = (r sin 0 dq~)(rdO) = r 2 sin 0 dO dq~,

i.e., z

= - 7b-gb

do~ = d a / r 2 = sin 0 dO dq~.

(HH)

Combining Eqs. (FF) and (HH), the actinic flux becomes F I G U R E 3.18 Typical light ray striking a thin layer of air in the atmosphere (adapted from Madronich, 1987).

F(A)=

f~ f 5 ( A, 0, q~)sin 0 dO dq~.

(II)

63


The absolute change in intensity at a particular wavelength is thus given by dI = cr(A)[A]I dl. This is just the number of photons absorbed as the light passes through the thin layer. I is the incident light, i.e., dP from Eq. (DD). If the quantum yield at this wavelength is 4~(A), then using I = dP and substituting in for dP from Eq. (DD), the total number of molecules of A that photodissociate is given by Number of A dissociating = q~(A) dI = q~(A) o- ( A)[A] dP dl = ~h(A) o- (A)[A] dl r sin 0

Z

• {L (A, 0, &)cos 0 da d o dt dA}.

i

r sin 0 de

(KK)

Since the total path length for light absorption is given by dl = d z / c o s 0 (Fig. 3.18b), this becomes Number of A dissociating = q~(A)o-(A)[A] dz • {L(A, O, 4~) dad~o dtdA}. (LL)

rdO I

/

o---// /

,s

.,~%.,5 /

//.." o I ,/':'S "/

/I

\

I

I

"",,.,.,.,,-"

I

\ I

'

~

dr

I

I

i

j,

•

t

,,."

d~

However, this represents the number of A molecules dissociating only due to light absorption in the wavelength interval d A and only for incident light over the solid angle d~o and surface area da. To obtain the total number of A dissociating, Eq. (LL) must be integrated over all wavelengths, solid angles, and surface areas Total number of A dissociating

FIGURE 3.19 coordinates.

Conversion

of solid angle

~o to spherical

Similarly, the irradiance E(A) is given by

= faf~s

O,dp)dzdad~odtdA

= [a]dt(fadzda )

E(A) = f L ( a , o, ,/,)cos o do, = Lfs(A,

0, ~h)sin 0 cos 0 dO dd~. (JJ)

Returning to Fig. 3.18a, as light passes through this thin layer of air, it can be absorbed by the molecules of A. From the differential form of the Beer-Lambert law, Eq. (O), the change in light intensity, dI, on passing through this small volume of air containing the absorber A is given by d I / I (x - (dl) = - k dl = - cr [A] dl,

where the negative sign indicates that the intensity decreases as the light passes through the sample. It is important to note that the application of this form of the Beer-Lambert law, inherent in which are natural logarithms, means that the absorption cross section, or, must be that to the base e.

(MM) The first integration over the surface area a is just the volume. Rearranging, Eq. (MM) becomes Total number of A dissociating/volume dt

d[A] dt

= [A]faL~h(A)cr(A)L(A, 0, ~h)d~o dA = [Alfx4~(A)o-(A)LL(A, 0, 4~) d~o dA. (NN)

64

3. SPECTROSCOPY A N D PHOTOCHEMISTRY: FUNDAMENTALS

Outer dome Quartz grains \

/

/ I n n e r dome

Coaxial quartz tubes

Radiation shield

Collimator

[ ~ ~ ~ ~ 1

Inte rfe re nce fllte r Photomultiplier

FIGURE 3.20 Schematic diagram of a 2~r radiometer used to measure actinic fluxes (adapted from Junkermann et al., 1989).

Light intensities can be measured using fiat-plate radiometers such as that in Fig. 3.17. As discussed earlier, this measures the flux through a horizontal plane, and as a result there is a difference of cos 0 between this measured irradiance and the actinic flux (Box 3.2). Another approach is to use what is sometimes referred to as a 2~r radiometer, which measures the light intensity striking half of a sphere. A typical example is shown in Fig. 3.20 (Junkermann et al., 1989). The light collector consists of coaxial quartz tubes located inside a quartz dome, such that light is collected from all directions in a hemisphere equally well (hence the "2~-" designation). Because light is collected from all directions within 2~r st, the cos 0 factor does not apply. The space between the quartz tubes and the dome is filled with quartz grains to scatter the light. The scattered light is transmitted by the quartz tubes through a set of interference filters to a detector. The filter-detector combination is chosen for the particular measurement of interest, e.g., O 3 or NO 2 photolysis. The shield is used to limit the field of view to exactly 2~r st. Two such detectors, one pointing up and one pointing down, can be used to cover the entire 4~r sphere. A similar detector for aircraft use is described by Volz-Thomas et al. (1996).

b. Estimates of the Actinic Flux, F(A), at the Earth's Surface

There are a number of estimates of the actinic flux at various wavelengths and solar zenith angles in the literature (e.g., see references in Madronich, 1987, 1993). Clearly, these all involve certain assumptions about the amounts and distribution of 0 3 and the concentration and nature (e.g., size distribution and composition) of particles which determine their light scattering and absorption properties. Historically, one of the most widely used data sets for actinic fluxes at the earth's surface is that of Peterson (1976), who recalculated these solar fluxes from 290 to 700 nm using a radiative transfer model developed by Dave (1972). Demerjian et al. (1980) then applied them to the photolysis of some important atmospheric species. In this model, molecular scattering, absorption due to 03, H 2 0 , 02, and CO2, and scattering and absorption by particles are taken into account. Madronich (1998) has calculated actinic fluxes using updated values of the extraterrestrial flux. In the 150to 400-nm region, values from Atlas are used (see Web site in Appendix IV)whereas from 400 to 700 nm those of Neckel and Labs (1984) are used. In addition, the

C. AT M O S P H E R I C PHOTOCHEMISTRY

ozone absorption cross sections of Molina and Molina (1986) and the radiation scheme of Stamnes et al. (1988) were used. Other assumptions, e.g., the particle concentration and distribution, are the same as those of Demerjian et al. (1980). In the "average" case for which the model calculations were carried out, absorption and scattering as the light traveled from the top of the atmosphere to the earth's surface were assumed to be due to 0 3 (UV light absorption with column 0 3 of 300 Dobson units), air molecules (scattering), and particles (scattering and absorption). The "best estimate" surface albedo varied from 0.05 in the 290- to 400-nm region to 0.15 in the 660- to 700-nm region. The surface was assumed to be what is known as an ideal "Lambert surface," meaning that it diffuses the incident light sufficiently well that it is reradiated equally in all upward directions, i.e., isotropically. Table 3.7 gives the calculated actinic fluxes at the earth's surface as a function of zenith angle assuming the "best estimate" surface albedo. These data are plotted for six wavelength intervals as a function of zenith angle in Fig. 3.21. The initially small change in actinic flux with zenith angle as it increases from 0 to 50 ~ at a given wavelength followed by the rapid drop of intensity from 50 to 90 ~ is due to the fact that the air mass m changes only gradually to ~ 50 ~ but then increases much more rapidly to 0 = 90 ~ [see Table 3.5 and Eq. (X)]. At the shorter wavelengths at a fixed zenith angle, the rapid increase in actinic flux with wavelength is primarily due to the strongly decreasing 0 3 absorption in this region. The actinic fluxes calculated by Madronich (1998) for altitudes of 15, 25, and 40 km are collected in Tables 3.15 to 3.17.

c. Effects of Latitude, Season, and Time of Day on F ( h ) To estimate photolysis rates for a given geographical location, one must take into account the latitude and season, as well as the time of day. The data in Table 3.7 are representative for the average e a r t h - s u n distance characteristic of early April and October. The orbit of the earth is slightly elliptical, so that there is a small change in the e a r t h - s u n distance, which causes a small change ( < 3%) in the solar flux with season. Correction factors for this seasonal variation for some dates from Demerjian et al. (1980) are given in Table 3.8. As discussed by Madronich (1993), the correction factors for solar intensity can be calculated for any other date using the following

( R o / R n ) 2 = 1.000110 + 0.034221 cos N + 1.280 • 10-3 sin N + 7.19 • 10 -4 cos 2 N + 7.7 • 10 -5 sin 2N,

65

where R 0 is the average e a r t h - s u n distance, R n is the e a r t h - s u n distance on day dn as defined earlier, N = 2 ~ r d , / 3 6 5 radians, and ( R o / R , ) 2 represents the correction factor. Table 3.9 summarizes the solar zenith angles at latitudes of 20, 30, 40, and 50~ as a function of month and true solar time. True solar time, also known as apparent solar time or apparent local solar time, is defined as the time scale referenced to the sun crossing the meridian at noon. For example, at a latitude of 50~ at the beginning of January, two hours before the sun crosses the meridian corresponds to a true solar time of 10 a.m.; from Table 3.9, the solar zenith angle at this time is 77.7 ~. To obtain the actinic flux at this time at any wavelength, one takes the fluxes in Table 3.7 listed under 78~ thus the flux in the 400- to 405-nm wavelength interval at 10 a.m. at 50~ latitude is 0.48 • 1015 photons c m - 2 s- 1. For other latitudes, dates, and times, the solar zenith angle can be calculated as described by Madronich (1993) and summarized earlier. Afternoon values of 0 are not given in Table 3.9 as the data are symmetrical about noon. Thus at a time of 2 p.m. at 50~ latitude, the flux would be the same as calculated for 10 a.m. Figure 3.22 shows the solar angle 0 as a function of true solar time for several latitudes and different times of the year. As expected, only for the lower latitudes at the summer solstice does the solar zenith angle approach 0 ~ at noon. For a latitude of 50~ even at the summer solstice, 0 is 27 ~ Figure 3.23 shows the diurnal variation of the solar zenith angle as a function of season for Los Angeles, which is located at a latitude of 34.1~ Clearly, the peak solar zenith angle varies dramatically with season. These differences in light intensity, and in its diurnal variation at different latitudes and seasons, are critical because they alter the atmospheric chemistry at various geographical locations due to the fact that photochemistry is the major source of the free radicals such as OH that drive the chemistry.

d. Effect of Surface Elevation on F( h ) The variation in the actinic flux with surface elevation is important because some of the world's major cities are located substantially above sea level. For example, Mexico City and Denver, Colorado, are at elevations of 2.2 and 1.6 kin, respectively. Table 3.10 shows the calculated percentage increase in the actinic flux at the earth's surface for an elevation of 1.5 km and atmospheric pressure of 0.84 atm (corresponding approximately to Denver) as a function of zenith angle for four wavelength intervals. In this calculation, it was assumed that the vertical 0 3 and particle

66

3. SPECTROSCOPY AND PHOTOCHEMISTRY: FUNDAMENTALS T A B L E 3.7 A c t i n i c F l u x V a l u e s F ( / t ) at the E a r t h ' s S u r f a c e as a F u n c t i o n of W a v e l e n g t h I n t e r v a l a n d Solar Z e n i t h A n g l e w i t h i n Specific W a v e l e n g t h I n t e r v a l s for B e s t E s t i m a t e S u r f a c e A l b e d o C a l c u l a t e d b y M a d r o n i c h ( 1 9 9 8 ) a

Wavelength interval (nm)

290-292 292-294 294-296 296-298 298-300 300-302 302-304 304-306 306-308 308-310 310-312 312-314 314-316 316-318 318-320 320- 325 325-330 330-335 335- 340 340-345 345- 350 350-355 355-360 360-365 365-370 370-375 375-380 380-385 385-390 390-395 395-400 400-405 405-410 410-415 415-420 420-430 430-440 440-450 450-460 460-470 470-480 480-490 490-500 500-510 510-520 520-530 530-540 540-550 550-560 560-570 570-580 580-600 600-620 620-640 640-660 660-680 680-700

Solar zenith angle (deg) Exponent b

0

10

20

14 14 14 14 14 14 14 14 14 14 14 14 14 14 14 14 14 14 14 14 14 14 14 14 15 15 15 15 15 15 15 15 15 15 15 15 15 15 15 15 15 15 15 15 15 15 15 15 15 15 15 16 16 16 16 16 16

0.00 0.00 0.00 0.01 0.03 0.07 0.18 0.33 0.51 0.66 0.99 1.22 1.37 1.67 1.70 5.30 7.72 8.26 7.98 8.64 8.73 10.00 8.98 9.97 1.24 1.10 1.26 1.06 1.17 1.17 1.43 2.02 1.97 2.06 2.09 4.13 4.26 5.05 5.66 5.75 5.86 5.74 5.94 6.10 5.98 6.20 6.38 6.37 6.55 6.61 6.69 1.35 1.36 1.37 1.38 1.43 1.40

0.00 0.00 0.00 0.01 0.03 0.07 0.18 0.32 0.49 0.65 0.97 1.19 1.34 1.64 1.68 5.24 7.63 8.17 7.91 8.57 8.65 9.92 8.91 9.90 1.23 1.09 1.25 1.06 1.16 1.17 1.42 2.01 1.96 2.04 2.08 4.11 4.24 5.03 5.64 5.72 5.83 5.72 5.91 6.07 5.95 6.17 6.38 6.34 6.52 6.58 6.66 1.34 1.35 1.37 1.37 1.42 1.40

0.00 0.00 0.00 0.01 0.02 0.06 0.15 0.29 0.45 0.60 0.90 1.12 1.27 1.56 1.60 5.03 7.36 7.91 7.67 8.32 8.42 9.67 8.69 9.67 1.20 1.07 1.22 1.04 1.14 1.14 1.40 1.98 1.92 2.01 2.05 4.04 4.18 4.96 5.56 5.64 5.75 5.64 5.83 5.99 5.87 6.09 6.27 6.26 6.43 6.49 6.57 1.32 1.34 1.35 1.35 1.40 1.38

30

40

50

60

Actinic fluxes (photons cm-2 S- 1) 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.01 0.00 0.00 0.00 0.02 0.01 0.00 0.00 0.04 0.03 0.01 0.00 0.12 0.08 0.04 0.01 0.23 0.16 0.09 0.04 0.37 0.28 0.17 0.08 0.51 0.40 0.27 0.14 0.79 0.64 0.45 0.25 1.00 0.82 0.61 0.36 1.14 0.96 0.73 0.46 1.42 1.22 0.95 0.62 1.47 1.27 1.01 0.69 4.66 4.10 3.34 2.36 6.88 6.15 5.12 3.75 7.44 6.70 5.65 4.23 7.24 6.56 5.59 4.25 7.88 7.17 6.15 4.73 7.99 7.30 6.30 4.88 9.20 8.43 7.31 5.71 8.28 7.62 6.64 5.22 9.23 8.52 7.46 5.91 1.15 1.07 0.94 0.75 1.02 0.95 0.84 0.67 1.17 1.09 0.97 0.78 1.00 0.93 0.82 0.67 1.10 1.03 0.92 0.75 1.10 1.03 0.92 0.76 1.35 1.26 1.13 0.93 1.91 1.79 1.61 1.33 1.86 1.75 1.57 1.30 1.94 1.83 1.65 1.37 1.98 1.87 1.69 1.41 3.92 3.70 3.36 2.82 4.05 3.84 3.50 2.96 4.82 4.57 4.18 3.56 5.39 5.12 4.67 3.98 5.48 5.21 4.77 4.08 5.60 5.32 4.89 4.19 5.49 5.23 4.81 4.14 5.68 5.42 5.00 4.31 5.82 5.54 5.09 4.38 5.71 5.44 5.00 4.32 5.92 5.64 5.20 4.49 6.10 5.81 5.35 4.63 6.10 5.81 5.35 4.63 6.26 5.96 5.49 4.74 6.31 6.01 5.53 4.78 6.39 6.09 5.60 4.84 1.29 1.23 1.13 0.98 1.30 1.24 1.14 0.99 1.31 1.26 1.16 1.01 1.32 1.26 1.17 1.02 1.37 1.31 1.21 1.06 1.34 1.28 1.19 1.05

70

78

86

0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.01 0.02 0.04 0.09 0.14 0.20 0.29 0.34 1.27 2.15 2.50 2.58 2.91 3.04 3.60 3.33 3.80 0.48 0.44 0.52 0.45 0.50 0.51 0.64 0.91 0.90 0.96 0.99 2.00 2.12 2.58 2.90 3.00 3.10 3.08 3.22 3.27 3.24 3.37 3.48 3.49 3.57 3.59 3.64 0.74 0.75 0.78 0.79 0.83 0.82

0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.01 0.02 0.04 0.06 0.10 0.13 0.52 0.96 1.16 1.23 1.40 1.47 1.76 1.64 1.88 0.24 0.22 0.26 0.23 0.26 0.27 0.33 0.48 0.48 0.51 0.54 1.09 1.17 1.45 1.65 1.72 1.80 1.80 1.90 1.92 1.92 2.00 2.07 2.08 2.13 2.13 2.17 0.45 0.46 0.48 0.50 0.53 0.54

0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.01 0.01 0.02 0.10 0.22 0.29 0.33 0.38 0.40 0.47 0.43 0.49 0.06 0.06 0.07 0.06 0.06 0.07 0.08 0.12 0.12 0.12 0.13 0.26 0.28 0.34 0.38 0.39 0.41 0.40 0.42 0.42 0.41 0.42 0.42 0.42 0.43 0.41 0.41 0.09 0.09 0.10 0.11 0.12 0.12

a The authors are grateful to Dr. Sasha Madronich for generously providing these calculations. b This column lists the power of 10 by which all entries should be multiplied. For example, at 0 = 0 ~ the total actinic flux in the wavelength interval from 306 to 308 nm is 0.51 • 1014 photons cm -2 s -1.

C. A T M O S P H E R I C

v

"-7 e0 |

E o

6

e. Effect of Height above Earth's Surface on F( h )

r

o

4

(t) o

0') (t)

=9

2

x :J M.--

o r

"=
26 are sufficiently energetic to allow reaction (3) to proceed: O2(u" > 26) + 0 2 ~ 0 3 + O(3p).

state, 22.5 kcal above the ground state, is ~ 67 min (Slanger and Cosby, 1988). As a result, when these electronic states are formed in the atmosphere, they are primarily collisionally deactivated to ground-state 0 29

Under low-pressure conditions in the laboratory or in the upper atmosphere where collisional deactivation is slow, weak radiative transitions from these two excited singlet states to the ground state are observed. The (0, 0) emission bands of the b 1 s X3s and the alZXg ~ X3s transitions occur at 761.9 and 1269 nm, respectively. The 761.9-nm band due to b1s is often observed in systems containing the a~ZXg state because of the energy pooling reaction: O2(alz~g) Jr- O 2 ( a l z ~ g ) ~

O2(X3~g)

(3)

Supporting this mechanism have been measurements of the production of O2(u" > 26) in the photolysis of 0 3 at shorter wavelengths (< 243 nm) (e.g., Miller et al., 1994; Syage, 1995, 1996a, 1996b; Stranges et al., 1995; Wilson et al., 1997). This is an important issue in the chemistry of the upper atmosphere, in that models have consistently underpredicted 0 3 at altitudes of 40-80 km, suggesting that there is an as yet unrecognized source of 0 3 at these altitudes. Although light absorption to excite 0 2 directly into the b ~s and a~ZXgstates is very weak, there are other sources of singlet oxygen, including energy transfer (e.g., from electronically excited NO~), 0 3 photolysis, and exothermic chemical reactions. When O2(a~ZXg) and O2(b1s are formed in such processes, they do not readily undergo radiative transitions to the ground state because, as in absorption, the processes are forbidden. Thus the radiative lifetime of the b~s state, which lies 37.5 kcal above the ground state, is ~ 12 s (Wallace and Hunten, 1968), whereas that for the a~zXg

(4)

B. OZONE 0 3 plays a central role in tropospheric chemistry. Not only is it a highly reactive and toxic species, but it absorbs both infrared and ultraviolet light, contributing to the "greenhouse effect" (see Chapter 14) and providing protection from exposure to damaging UV. In the process of absorbing in the UV, it generates electronically excited oxygen atoms that react to form OH, a ubiquitous atmospheric oxidant. 1. A b s o r p t i o n S p e c t r a Absorption spectra of 0 3 are shown in Figs. 4.7 and 4.8. The strongest absorption, in the 200- to 300-nm region, is known as the Hartley bands. It is this absorption that is responsible for the so-called "actinic cutoff"

..--..

r

1.2

..o "7 "5 0

1.0

6'3

Ix101

~x10 2

x10 3 I 1 0 4

0.8

G)

o E E 0 D ~. 04

0.6

0.4 0.2

CD ,r-

200

Jr- O 2 ( b l ~ ; ) .

1

I

250

300

I

350

;L (nm)

FIGURE 4.7 UV absorption of 0 3 at room temperature in the Hartley and Huggins bands. At the longer wavelengths, each curve has been expanded by the factor shown. (Adapted from Daumont et al., 1989.)

s. ozoNs 6.0 5.5 (13 t~

m

5.0 4.5 4.0

0

3.5

0

3.0

0,,1

2.5

0

2.0

E E

1.5 o,1 0

1.0 0.5 0.0 400 450 500 550 600 650 700 750 Wavelength (nm)

F I G U R E 4.8 Absorption cross sections of 0 3 in the Chappuis band at room temperature (adapted from Burkholder and Talukdar, 1994).

91

In addition to energetic considerations, however, there are other factors such as spin conservation that also determine the importance of various sets of products. As discussed in Chapter 3.A, since the ground state of 0 3 is a singlet, dissociation into either two singlet states (e.g., reaction (5)) or into two triplet states is expected to predominate. However, as discussed shortly, both hot-band absorption by rovibrationally excited 0 3 and by a spin-forbidden process are believed to contribute significantly to the atmospheric photochemistry of 0 3. The most important aspect of 0 3 photochemistry for the troposphere is the yield and wavelength dependence of O(1D) production in reaction (5) since it is a source of hydroxyl free radicals via its reaction with water vapor: O(1D) + H20(g ) ~ 2OH, k 6 =

in the troposphere at 290 nm; thus 0 3 in the stratosphere strongly absorbs A < 290 nm, limiting the UV reaching the earth's surface. The Huggins bands are in the 300- to 360-nm region, whereas the 440- to 850-nm region represents the Chappuis bands. As seen in Figs. 4.7 and 4.8, both absorptions are much weaker than the Hartley bands. Table 4.3 shows the ozone absorption cross sections (base e) at 298 and at 226 K measured by Molina and Molina (1986), with which recent studies are in agreement, particularly below 240 nm (e.g., see Malicet et al., 1995; DeMore et al., 1997). The cross sections below 275 nm are not very sensitive to temperature over the 226-298 K temperature range, typically < 1%, but the temperature dependence becomes significant above 275 nm. Table 4.4 shows the temperature dependence of these absorption cross sections at longer wavelengths, averaged over the spectral regions shown.

2. Photochemistry The photolysis of 0 3 produces molecular oxygen and atomic oxygen, either or both of which may be in electronically excited states, depending on the excitation energy. Table 4.5 shows the wavelength threshold below which each combination of products may be formed. For example, the primary photochemical process 0 3 + h u ~ O2(lZXg) + O(1D)

(5)

in principle requires light corresponding to A = 310 nm. (However, as discussed below, production of 0 ( 1D) has been observed experimentally out to at least 336 nm.)

(6)

2.2 • 10 -1~ cm 3 molecule -1 s -1

(DeMore et al., 1997). Only about 5% of the O ( 1 D ) - H 2 0 interaction results in deactivation of the O(1D) to the ground state O(3P), and < 1% in the alternate set of products H 2 + O 2 (DeMore et al., 1997). The very fast reaction of O(1D) with water vapor occurs in competition with its deactivation by air: O(1D) + M --. O(3p) + M,

(7)

k7 = 2.6 X 10 -11 cm 3 molecule -1 s -1 for M = N2, = 4.0 X 10 -11 cm 3 molecule -1 s -1 for M = O 2 (DeMore et al., 1997). In 1 atm of air at 50% relative humidity and 298 K, approximately 10% of the O(1D) produced via reaction (5) reacts with water vapor to form hydroxyl radicals. The wavelength dependence of the quantum yield for 0 ( 1D) production has been somewhat controversial. The recommended value between 290 and 305 nm is ~h[O(1D)] = 0.95, decreasing slightly at shorter wavelengths to values in the range 0.85-0.90 (see review in DeMore et al., 1997). Talukdar et al. (1998) reported yields of O(1D) of 0.89 _+ 0.04 (2o-), independent of temperature from 203 to 320 K, consistent with the production of small amounts of O(3p) in this region (e.g., Brock and Watson, 1980). The falloff in the quantum yield at shorter wavelengths is unusual and not well understood in terms of the spectroscopy and electronic states of 0 3 (e.g., Steinfeld et al., 1987; Wayne, 1987). Wayne (1987) suggests that the upper electronic states initially populated on light absorption have a curve crossing with a repulsive state that correlates with the ground-state products O(3p) + O2(3~g) and that the relative positions of the states are such

92

4. PHOTOCHEMISTRY OF IMPORTANT ATMOSPHERIC SPECIES TABLE 4.3

Ozone A b s o r p t i o n Cross Sections (Base e)a

1020 r ( c m 2 m o l e c u l e - l )

10 TMr ( c m 2 m o l e c u l e - l )

Wavelength (nm)

T = 226 K

T = 298 K

Wavelength (nm)

185.0 186.0 187.0 188.0 189.0 190.0 191.0 192.0 193.0 194.0 195.0 196.0 197.0 198.0 199.0 200.0 201.0 202.0 203.0 204.0 205.0 206.0 207.0 208.0 209.0 210.0 211.0 212.0 213.0 214.0 215.0 216.0 217.0 218.0 219.0 220.0 221.0 222.0 223.0 224.0 225.0 226.0 227.0 228.0 229.0 230.0 231.0 232.0 233.0 234.0 235.0 236.0 237.0 238.0 239.0 240.0 241.0 242.0 243.0 244.0 245.0 246.0 247.0 248.0 249.0 250.0 251.0 252.0 253.0 254.0 255.0 256.0 257.0 258.0 259.0 260.0 261.0

64.37 62.59 59.33 56.55 54.63 51.63 48.42 45.95 43.12 40.88 38.27 36.42 34.63 33.33 32.13 31.45 31.26 31.56 32.55 34.00 36.23 38.87 42.39 46.84 51.88 58.06 65.28 73.12 82.58 92.55 104.1 116.9 131.4 146.4 163.8 179.9 200.0 221.7 244.3 268.8 296.3 323.9 354.2 385.7 416.4 450.6 485.9 523.0 558.9 592.3 634.9 677.2 714.2 753.2 784.3 836.7 864.0 901.6 939.6 975.2 1007 1042 1058 1079 1124 1134 1123 1165 1149 1169 1174 1158 1147 1130 1151 1086 1095

65.37 61.87 59.41 56.59 54.24 51.14 48.80 46.06 43.36 40.66 38.64 36.73 35.00 33.49 32.09 31.54 31.15 31.79 32.51 33.65 35.85 38.55 42.00 46.40 51.18 57.16 64.02 71.94 81.04 90.96 102.3 114.6 128.7 143.9 160.1 178.5 198.2 220.0 242.9 268.4 294.3 322.6 351.3 382.9 414.1 447.6 481.4 518.1 554.9 589.1 631.8 672.2 709.4 748.6 781.9 831.4 860.3 897.1 933.3 971.7 993.2 1033 1047 1071 1112 1124 1114 1155 1140 1159 1161 1154 1139 1124 1145 1080 1094

262.0 263.0 264.0 265.0 266.0 267.0 268.0 269.0 270.0 271.0 272.0 273.0 274.0 275.0 276.0 277.0 278.0 279.0 280.0 281.0 282.0 283.0 284.0 285.0 286.0 287.0 288.0 289.0 290.0 291.0 292.0 293.0 294.0 295.0 296.0 297.0 298.0 299.0 300.0 301.0 302.0 303.0 304.0 305.0 306.0 307.0 308.0 309.0 310.0 311.0 312.0 313.0 314.0 315.0 316.0 317.0 318.0 319.0 320.0 321.0 322.0 323.0 324.0 325.0 326.0 328.0 330.0 332.0 334.0 336.0 338.0 340.0 342.0 344.0 346.0 348.0 350.0

a

F r o m Molina and Molina (1986).

T = 226 K

1064 1016 1013 961.2 946.7 877.9 872.2 803.8 796.1 736.7 710.7 666.0 601.3 587.1 537.9 504.0 461.7 424.5 398.3 364.2 322.0 299.0 264.4 239.8 219.5 197.8 168.7 153.5 136.5 122.5 104.3 93.75 81.48 72.70 62.02 54.69 47.14 42.31 36.16 31.28 27.96 23.25 21.56 17.71 16.03 13.67 12.02 10.64 8.637 7.925 6.697 5.691 5.334 4.186 3.896 3.351 3.063 1.999 2.859 1.368 2.117 1.529 0.7852 1.486 0.7276 1.158 0.2854 0.2415 0.3788 0.0871 0.1311 0.1549 0.0285 0.1271 0.0195 0.0181 0.0076

T = 298 K

1057 1022 1006 965.7 948.5 884.1 875.4 810.4 798.0 741.5 714.7 669.8 614.0 591.3 545.0 509.6 466.8 432.6 400.1 367.3 325.0 302.5 271.2 246.5 223.8 203.4 175.0 158.5 141.8 128.5 111.1 100.2 87.11 77.53 67.27 59.55 51.24 45.51 39.64 34.63 30.73 26.50 24.01 20.15 18.08 15.65 13.64 12.43 10.20 9.260 7.947 6.883 6.294 5.199 4.792 4.146 3.757 2.765 3.243 2.041 2.435 1.983 1.250 1.727 1.105 1.300 0.4923 0.4347 0.5343 0.1875 0.2086 0.2082 0.0776 0.1419 0.0641 0.0386 0.0294

S. OZONS TABLE 4.4

O z o n e A b s o r p t i o n C r o s s S e c t i o n s a'b as a F u n c t i o n of T e m p e r a t u r e A v e r a g e d over the Spectral Intervals Shown

Parameters

Wavelength range (nm)

a

277.778-281.690 281.690-285.714 285.714-289.855 289.855-294.118 294.118-298.507 298.507-303.030 303.030-307.692 307.692-312.5 312.5-317.5 317.5-322.5 322.5-327.5 327.5-332.5 332.5-337.5 337.5-342.5 342.5-347.5

b

4.0293 • 102 2.7776 • 102 1.8417 • 102 1.1300 • 102 6.5087 • 10 3.6161 • 10 1.9615 • 10 1.0459 • 10 5.4715 2.7569 1.3527 6.9373 • 10 -1 3.2091 • 10 -1 1.4484 • 10 -1 7.5780 x 10 .2

a O_(03, T ) = a + b ( T - 230) + C ( T cule-1 (base e). b F r o m Molina and Molina (1986).

+4.3819 + 6.3125 - 9.6665 - 1.0700 - 8.0018 - 6.7156 -4.4193 - 2.8831 - 2.0092 - 1.0067 - 5.7513 - 2.9792 - 1.9502 - 1.1025 - 5.7359

TABLE 4.5 Wavelength Threshold (nm) Below Which Indicated Reactions Are Energetically Possible i n t h e P h o t o l y s i s of O3 a

3p 1D 1S

Electronic state of molecular 0 2 3~g

lag

,Z+

3~ +

3Za-

1180 411 237

612 310 199

463 267 181

230 168 129

173 136 109

a F r o m Wayne (1987) and O k a b e (1978).

• • • • • • • • • • • • • • •

10 . 2 10 . 2 10 -2 10 -1 10 .2 10 -2 10 -2 10 -2 10 -2 10 -2 10 -3 10 -3 10 -3 10 -3 10 . 4

2.1026 3.2697 2.2679 3.3314 2.0338 1.3909 9.8870 2.9515 1.1088 3.1038 5.6456 2.8818 1.6055

0 0 • • • • • • • • • • • • •

10 -4 10 -4 10 . 4 10 -4 10 -4 10 -4 10 -5 10 -5 10 -5 10 -6 10 -6 10 -6 10 . 6

330 nm measured in three of these studies (Armerding et al., 1995; Ball et al., 1997; Talukdar et al., 1998). Silvente et al. (1997) observe substantial yields (~0.05) of O(1D) out to ~ 336 nm, well beyond the energetic threshold for reaction (5). It appears that there are two processes that contribute to this O(1D) production. The first is so-called "hot-band absorption" by O 3 in which the additional energy comes from internal vibrational and rotational energy (Michelsen et al., 1994), a phenomenon that is well established in the case of NO 2 photodissociation (see later). This is believed to be responsible for 0 ( 1D)

1.0

--

o

o

O O 9~ o

"t

0.8 ..-..

oxygen atom

c

230) 2. T is in K; ~r(O 3) is in units of 10 .20 cm 2 mole-

that higher energies are needed for this to occur. Theoretical studies by Banichevich et al. (1993) suggest that there are a number of excited singlet and triplet states that are only weakly bound or are repulsive and correlate with O(3p)4-O2(3s As discussed in Section 4.A.2, photolysis at shorter wavelengths, e.g., 226 nm, produces in part ground electronic state 0 2 that is in highly vibrationally excited states. As discussed earlier, the spin-allowed process that produces O(1D) a n d O2(1Ag) has an apparent energetic threshold corresponding to 310 nm. However, significant yields of O(1D) have been observed at wavelengths longer than this in a number of experiments (e.g., Arnold et al., 1977; Trolier and Wiesenfeld, 1988; Armerding et al., 1995; Takahashi et al., 1996a, 1996b, 1997; Ball et al., 1997; Silvente et al., 1997; Talukdar et al., 1997, 1998; Ravishankara et al., 1998). For example, Fig. 4.9 shows the quantum yields of 0 ( 1D) from 300 to

Electronic state of

93

O

9Amerding et al., 1995 o Ball et al., 1997 9Talukdar et al., I998

ii o II

0.6-

O

0.4-

~Q 0.2-

9 0.0 300

1

305

1

310

Ooo

[

1

[

315

320

325

9

I

330

X(nm)

F I G U R E 4 . 9 Some m e a s u r e m e n t s of the q u a n t u m yields for production of 0 ( 1D) in the photolysis of O 3 at 298 K.

94

4. P H O T O C H E M I S T R Y OF I M P O R T A N T A T M O S P H E R I C SPECIES

TABLE 4.6 Parameterization of Quantum Yields for O(1D) Production from 03 Photolysis in the 306- to 329-nm Region at Various Temperatures a

production in the region from approximately 306 to 324 nm. As expected, if this is the case, the quantum yield in this region decreases with temperature, since the Boltzmann population of the absorbing, vibrationally excited state of 0 3 decreases; for example, at 320 nm and 298 K, qS[O(1D)] = 0.15, but decreases to 0.06 at 203 K (Talukdar et al., 1998). The second process generating small amounts of O(1D) is thought to be the spin-forbidden reaction 0 3 + h p--* O(1D) + O2(3Eg).

O2(1Ag a n d / o r 3 s

O(3p).

306 307 308 309 310 311 312 313 314 315 316 317 318 319 320 321 322 323 324 325 326 327 328 329

(8)

This spin-forbidden process appears to be the major source of O(1D) beyond 325 nm, particularly at low temperatures where the contribution of vibrationally excited 0 3 is minimized. Thus, in contrast to the decreasing yields of O(1D)with temperature in the 306to 324-nm region, yields in the 325- to 329-nm region are relatively constant with temperature; at lower temperatures, the quantum yield in this region approaches 0.06, which has therefore been assigned by Talukdar et al. (1998) as the upper limit for the production of 0 ( 1D) in this region by reaction (8). Table 4.6 summarizes the wavelength and temperature dependence for O(1D) production recommended by Talukdar et al. (1998). Beyond 329 nm, the quantum yield from the spin-forbidden process appears to be 0.05-0.06 (Silvente et al., 1997; Talukdar et al., 1998). The absorption cross sections of 0 3 become sufficiently small beyond 360 nm that O(1D) production beyond this wavelength is not expected to be significant for atmospheric applications. Ball et al. (1993) and Ball and Hancock (1995) measured relative yields of O2(lZXg), the other product expected in the spin-allowed reaction (5), and found that it too was produced at significant yields. For example, Fig. 4.10 shows the measured yields of O2(lZXg) normalized to ~b[O(1D)] = 0.95 in the 290- to 300-nm region. The yields are similar in magnitude to those for 0(1D). If reaction (8) is " responsible " for O ( 1D ) production in the region beyond 325 nm, the yield of O2(lZXg) might also be expected to fall to zero. However, O2(lZXg) may also be generated by other paths, for example, that giving O(3p) + O2(lZXg), which requires light of 612 nm or shorter wavelengths (Table 4.5). This is consistent with studies of absorption in the Huggins bands, in which the products appear to be o2(l~g) and ground-state oxygen atoms (Okabe, 1978). Thus, the spin-forbidden process (9) to give electronically excited O 2 and ground-state oxygen atoms must be occurring: 0 3 + hp--,

Wavelength (nm)

(9)

In the Chappuis region (440-850 nm), the products

A

B

0.80 0.78 0.87 0.76 1.31 2.37 5.8 11.4 20.1 26.4 26.8 26.8 28.33 30.6 44.4 50.2 27.8 74.1 868 0.37 0.24 0.068 26.16 0.15

9.84 1.44 53.1 73.9 305.5 600 925.9 1191 1423 1514 1512 1542 1604 1604 1866 1931 1882 2329 3085 689 619 258 2131 470

a Using the quantum yield expression recommended by Talukdar et al., 1998: (f = 0.06 + Ae -B / r

appear to be the ground-state species (Okabe, 1978): 0 3 -Jr-hi;---> O 2 ( X 3 s

-Jr-O(3p).

(10)

In both these regions, the absorption coefficient is one to two orders of magnitude less than that at 300

1.0 o

.2 -~ 0.5

0.0

270

1

[

280

[

[

290

1

1

300

[

[

310

1

1

320

[

330

03 Photolysis wavelength (nm)

FIGURE 4.10 Relative yield of O2(1Ag)__ from the photolysis of 03 (adapted from Ball et al., 1993).

C. NITROGEN DIOXIDE

,,,-...., (b r,,O ,.Q

.o_ "7 "~

CO 69 CO

s r

0 0

E

O4

E o

03

"7,

c) T--

x

0 350

450

550

(nm)

F I G U R E 4.11 High-resolution (0.003 nm) absorption spectrum of NO 2 at room temperature (adapted from Harder et al., 1997).

nm. In addition, they produce O(3P) rather than 0(1D). Under conditions typical of urban and suburban areas where sufficient NO 2 is present, the production of O(3p) from photolysis of 03 will be much less than that from NO2; in addition, since the O(3p) reacts with 02 to regenerate 03, this leads to no net loss of ozone. As a result, the photolysis processes generating O(3p) are not significant in tropospheric chemistry.

C. N I T R O G E N

DIOXIDE

1. A b s o r p t i o n S p e c t r a

Tables 4.7 and 4.8 show the recommended absorption cross sections (DeMore et al., 1997). Those in the 202to 274-nm region are temperature independent. At the longer wavelengths there appears to be a small temperature dependence parameterized as shown in Table 4.8. As discussed in Chapter 11, NO 2 is frequently measured in the atmosphere using UV-visible absorption spectroscopy. Because of the large number of possible absorbing species in the atmosphere, NO 2 is usually measured using the peak-to-valley absorbance of selected peaks, rather than peak to zero absorbance. While reducing the temperature from 298 K only slightly affects the peak absorption cross sections, it does decrease the minima in the spectrum, thus increasing the peak-to-valley absorbance differences used for measuring NO 2 in the atmosphere by differential optical absorption spectrometry (e.g., Davidson et al., 1988; Harwood and Jones, 1994; Vandaele et al., 1996; Harder et al., 1997). This is expected since lowering the temperature will lower the population of the higher rotational energy levels in the ground state, leading to a smaller range of energies and wavelengths absorbed in a given transition. In addition to the temperature dependence, a pressure dependence of the absorption cross sections has been observed in high-resolution studies (Harder et al., 1997), which is important for application in the upper atmosphere. One of the complications in measuring NO 2 absorption cross sections, particularly at lower temperatures and higher concentrations, is the presence of the dimer N20 4 in equilibrium with it:

Figure 4.11 shows a high-resolution (0.003 nrn) absorption spectrum of NO: (Harder et al., 1997) and

T A B L E 4.7

95

2NO 2 N 2 0 4 .

Averaged A b s o r p t i o n Cross Sections (Base e) from 202 to 274 n m for N O 2 at 298 K a 10 20 o -b

k

1 0 TMo-b

(nm)

(cm 2 molecule - l)

(nm)

(cm 2 molecule- l)

202.02-204.08 204.08- 206.19 206.19-208.33 208.33-210.53 210.53-212.77 212.77-215.06 215.06-217.39 217.39-219.78 219.78-222.22 222.22-224.72 224.72-227.27 227.27-229.89 229.89-232.56

41.45 44.78 44.54 46.41 48.66 48.18 50.22 44.41 47.13 37.72 39.29 27.40 27.78

232.56-235.29 235.29- 238.09 238.09-240.96 240.96-243.90 243.90-246.91 246.91-250.00 250.00-253.17 253.17-256.41 256.41-259.74 259.74-263.16 263.16-266.67 266.67-270.27 270.27-273.97

16.89 16.18 8.812 7.472 3.909 2.753 2.007 1.973 2.111 2.357 2.698 3.247 3.785

a From DeMore et al. (1997) recommendation based on Schneider et al., J. Photochem. Photobiol. A., 40, 195, (1987). b Independent of temperature.

( 1 1 , - 11)

96

4. PHOTOCHEMISTRY OF I M P O R T A N T ATMOSPHERIC SPECIES 4.8 Recommended Averaged Absorption Cross Sections (Base e) for N O 2 from 2 7 4 to 4 2 2 nm at 0 ~ and Their Temperature D e p e n d e n c e a

TABLE

/k (nm) 273.97-277.78 277.78-281.69 281.69-285.71 285.71-289.85 289.85-294.12 294.12-298.51 298.51-303.03 303.03-307.69 307.69-312.50 312.5-317.5 317.5-322.5 322.5-327.5 327.5-332.5 332.5-337.5 337.5-342.5 342.5-347.5 347.5-352.5 352.5-357.5 357.5-362.5 362.5-367.5 367.5-372.5 372.5-377.5 377.5-382.5 382.5-387.5 387.5-392.5 392.5-397.5 397.5-402.5 402.5-407.5 407.5-412.5 412.5-417.5 417.5-422.5

10 20tr (cm 2 molecule - l)

1022a b (cm 2 molecule- 1 degree - l)

5.03 5.88 7.00 8.15 9.72 11.54 13.44 15.89 18.67 21.53 24.77 28.07 31.33 34.25 37.98 40.65 43.13 47.17 48.33 51.66 53.15 55.08 56.44 57.57 59.27 58.45 60.21 57.81 59.99 56.51 58.12

0.075 0.082 - 0.053 - 0.043 - 0.031 -0.162 - 0.284 - 0.357 - 0.536 - 0.686 - 0.786 - 1.105 -1.355 - 1.277 - 1.612 - 1.890 - 1.219 - 1.921 - 1.095 - 1.322 - 1.102 -0.806 -0.867 -0.945 -0.923 -0.738 -0.599 -0.545 - 1.129 0.001 - 1.208

a From DeMore et al. (1997) recommendation based on Davidson et al. (1988); see references. b This is the temperature coefficient of o-(t) = o-(0 ~ + at, where

t is in ~

As seen from the data in Table 4.9, N 2 0 4 also absorbs strongly in this region, so care must be taken in both spectroscopic and kinetic studies to take this into account. For low concentrations of NO 2 at room temperature, significant amounts of the dimer are not present, but this is not the case at higher concentrations. For example, using the recommended value of K = 2.5 • 10 -19 cm 3 molecule -1 at 298 K for reaction (11, - 1 1 ) (DeMore et al., 1997), the pressure of N 2 0 4 in equilibrium with 1.0 Torr of NO 2 (equivalent to 1318 ppm) is 8.1 • 10 .3 Torr, i.e., 0.81%, but for 1 ppb NO 2 (7.6 • 1 0 - 7 Torr), it is only 6 • 1 0 - 9 ppb, i.e., 6 • 1 0 - 7 % ! This equilibrium means that great caution must be exercised in extrapolating both kinetic and mechanistic studies from Torr concentrations of NO 2 to ambient ppt-ppb levels, since N 2 0 4 may contribute significantly

TABLE

4.9

Average Cross Sections (Base e ) for N 2 0 4 a'b 10 20o" (cm 2 molecule- l)

Wavelength (nm)

253 K

213 K

320 325 330 335 340 345 350 355 360 365 370 375 380 385 390 395 400 405 410 415 420 425 430

43.87 52.93 60.89 65.57 66.46 63.07 55.87 47.15 37.93 30.84 25.39 22.02 20.51 18.80 18.24 16.78 17.16 15.47 15.01 13.87 13.66 12.31 10.91

40.46 50.32 58.97 64.23 64.29 60.1 50.25 37.64 25.44 16.04 9.67 5.61 3.01 1.62 0.84 0.52 0.19

a Averaged over 5-nm intervals centered on the wavelength shown. b From Harwood and Jones (1994).

under some conditions, especially high concentrations and lower temperatures. In other cases, it is not a problem. For example, the reaction of hydrogen atoms with NO 2 is used as a source of OH in fast flow discharge systems. The reservoir holding NO 2 is often at relatively high concentrations and contains significant amounts of N204. However, the equilibrium (11, - 1 1 ) shifts sufficiently rapidly that the N204 dissociates to form NO 2 so that the initial presence of N204 is not a problem. 2. P h o t o c h e m i s t r y Nitrogen dioxide photodissociates at A < 420 nm to give nitric oxide and an oxygen atom: NO 2 + h v ~ NO + O.

(12)

The quantum yield for oxygen atom production in (12) has been studied extensively because of its role as the only significant anthropogenic source of 03 in the troposphere via (12) followed by (2): M O

-+- 0 2

--+ 0 3 .

(2)

C. NITROGENDIOXIDE TABLE 4.10 Calculated Wavelengths (nm) for NO 2 Photolysis Below Which the Fragments Shown Can Be Produced a'b

3p

Oxygen atoms 1D

1S

397.8 144.2

243.9 117.4

169.7 97

NO

1211 A2s +

a From Okabe (1978). b Assuming no contribution from internal energy of the molecule.

Table 4.10 gives the calculated wavelengths below which it is energetically possible to p r o d u c e the fragm e n t s in each of the electronic states shown if t h e r e is no contribution f r o m internal energy of the molecule. T h e threshold wavelength for p r o d u c t i o n of groundstate N O and O atoms is 397.8 nm. Figure 4.12 shows the primary q u a n t u m yields for the p r o d u c t i o n of N O in reaction (12). The q u a n t u m yield is within e x p e r i m e n t a l e r r o r of 1 up to 395 nm, declining slightly to 0.82 at the theoretical threshold for dissociation at 397.8 nm. This has b e e n attributed to the f o r m a t i o n of a nondissociative excited state of N O 2. In a m b i e n t air, electronically excited N O 2 which does not dissociate to f o r m O(3P) is collisionally deactivated. W h e n 0 2 is the collision partner, energy transfer may occur a fraction of the time to f o r m O2(1Ag) (Jones and Bayes, 1973a):

NO~

Jr- 0 2 ---->N O 2 Jr- o2(lZ~g).

(13)

Figure 4.12 shows that significant photodissociation continues at longer wavelengths ( R o e h l et al., 1994). T h e n o n z e r o q u a n t u m yield for reaction (12) b e t w e e n

1.1

10-

09

0.8 0.7

o

oo

ogSooo~o..

~o~/

7o 0.6

TOTAL

0.5 0.4 0.3

-. o

"'.. o

Internale n e r g ) ~ CollisionsN,_. . . . .

0.10.2

00

385

t

J

390

395

~ 400

i~

405

--- ~=-"'~ 410

415

;L(nm) FIGURE 4.12 Quantum yields for NO production in the photolysis of N O 2 a t 298 K. Calculated quantum yields due to internal energy (dotted line), the calculated dissociation due to collision (dashed line), and the sum of these two calculations (solid line) are also shown (adapted from Roehl et al., 1994).

97

397.8 and 420 n m is believed to be due to a contribution f r o m internal vibrational and rotational energy as well as to a small contribution f r o m collisions. Q u a n t u m yields calculated assuming that the v i b r a t i o n - r o t a tion internal energy of the molecule is available to o v e r c o m e the p h o t o n energy deficiency (dotted line in Fig. 4.12) as well as the energy acquired in collisions of the excited N O 2 with o t h e r molecules (dashed line) are shown ( R o e h l et al., 1994). A good fit to the data is obtained, supporting the i m p o r t a n c e of internal as well as collisional energy for photodissociation in the 398to 420-nm region. A t w o - p h o t o n process at 435 n m to g e n e r a t e O(1D) has b e e n observed in laboratory studies (Crowley and Carl, 1997). However, the c o m b i n a t i o n of relatively low light intensities and high pressures which quench excited N O 2 m a k e this u n i m p o r t a n t in the a t m o s p h e r e . TABLE 4.11

Quantum Yields for NO2 Photolysisa

Wavelength (nm)

~12

< 285 290 295 300 305 310 315 320 325 330 335 340 345 350 355 360 370 375 380 381 382 383 384 385 386 387 388 389 390 391 392

1.000 0.999 0.998 0.997 0.996 0.995 0.994 0.993 0.992 0.991 0.990 0.989 0.988 0.987 0.986 0.984 0.981 0.979 0.975 0.974 0.973 0.972 0.971 0.969 0.967 0.966 0.964 0.962 0.960 0.959 0.957

Wavelength (nm)

~12

393 394 395 396 397 398 399 400 401 402 403 404 405 406 407 408 409 410 411 412 413 414 415 416 417 418 419 420 421 422 423 424

0.953 0.950 0.942 0.922 0.870 0.820 0.760 0.695 0.635 0.560 0.485 0.425 0.350 0.290 0.225 0.185 0.153 0.130 0.110 0.094 0.083 0.070 0.059 0.048 0.039 0.030 0.023 0.018 0.012 0.008 0.004 0.00

a From DeMore et al. (1997) recommendations based on Gardner et al., J. Geophys. Res., 92, 6642 (1987); Harker et al., Chem. Phys. Lett., 50, 394 (1977); Jones and Bayes, 1973b and Davenport, Report # FAA-EQ-78-14, U.S. Federal Aviation Administration, Washington, D.C. (1978).

98

4. PHOTOCHEMISTRY OF IMPORTANT ATMOSPHERIC SPECIES

T a b l e 4.11 s u m m a r i z e s t h e r e c o m m e n d e d q u a n t u m yields in t h e 285- to 4 2 4 - n m r e g i o n ( D e M o r e et al., 1997).

10-17 (~ 10 -18 G)

D. NITRIC ACID

.s "7

H N O 3 + h v - - + O H + N O 2.

r "5

10-20

oE E o~ b

10_21

O

F i g u r e 4.13 shows t h e a b s o r p t i o n s p e c t r u m of g a s e o u s H N O 3 at r o o m t e m p e r a t u r e ( B u r k h o l d e r et al., 1993) a n d T a b l e 4.12 shows t h e r e c o m m e n d e d a b s o r p t i o n cross s e c t i o n s at 298 K ( D e M o r e et al., 1997). T h e q u a n t u m yield for p h o t o d i s s o c i a t i o n , r e a c t i o n (14), is a p p r o x i m a t e l y 1 f r o m 200 to 315 n m ( D e M o r e et al., 1997):

r

0-19

0_22 1

= _

10 -23

10 -24

I

I

I

I

I

I

I

I

200 220 240 260 280 300 320 340

(14)

Wavelength (nm) A t s h o r t e r w a v e l e n g t h s in t h e v a c u u m u l t r a v i o l e t , t h e p a t h to give O + H O N O a p p e a r s to b e c o m e i m p o r t a n t . D o n a l d s o n et al. (1997) h a v e p r o p o s e d t h a t a b s o r p t i o n in t h e visible d u e to O H v i b r a t i o n a l o v e r t o n e s c o u l d b e i m p o r t a n t in t h e l o w e r s t r a t o s p h e r e at l a r g e s o l a r z e n i t h angles. T r a n s f e r of e n e r g y f r o m t h e a b s o r b ing m o d e to t h e H O m N O 2 b o n d m a y t h e n c a u s e d i s s o c i a t i o n , as o b s e r v e d , for e x a m p l e , in H O C 1 (e.g., TABLE 4.12 Wavelength (nm)

190 192 194 196 198 200 202 204 206 208 210 212 214 216 218 220 222 224 226 228 230 232 234 236 238 240 242

10 2o (r

(cm 2

molecule- l) 1360 1225 1095 940 770 588 447 328 231 156 104 67.5 43.9 29.2 20.0 14.9 11.8 9.61 8.02 6.82 5.75 4.87 4.14 3.36 2.93 2.58 2.34

F I G U R E 4.13 Absorption spectrum of HNO 3 (solid line) at 298 K. Symbols represent data reported in previous studies from other laboratories. (Adapted from Burkholder et al., 1993.)

B a r n e s et al., 1997). T h e e n e r g y n e e d e d for H O m N O 2 b o n d c l e a v a g e is slightly l a r g e r t h a n 5 y o n , b u t a c o n t r i b u t i o n f r o m r o t a t i o n a l e n e r g y m a y m a k e this feasible.

Absorption Cross Sections (Base e) for H N O 3 Vapor at 298 K a Wavelength

(nm) 244 246 248 250 252 254 256 258 260 262 264 266 268 270 272 274 276 278 280 282 284 286 288 290 292 294 296

10 2o (r

(cm 2

molecule - l) 2.16 2.06 2.00 1.97 1.96 1.95 1.95 1.93 1.91 1.87 1.83 1.77 1.70 1.62 1.53 1.44 1.33 1.23 1.12 1.01 0.909 0.807 0.709 0.615 0.532 0.453 0.381

Wavelength

(nm) 298 300 302 304 306 308 310 312 314 316 318 320 322 324 326 328 330 332 334 336 338 340 342 344 346 348 350

10 2o o"

(cm 2

molecule- l) 0.316 0.263 0.208 0.167 0.133 0.105 0.0814 0.0628 0.0468 0.0362 0.0271 0.0197 0.0154 0.0108 0.00820 0.00613 0.00431 0.00319 0.00243 0.00196 0.00142 0.00103 0.00086 0.00069 0.00050 0.00042 0.00042

a From DeMore et al. (1997) recommendation based on Burkholder et al. (1993); see these for the temperature dependence, which can be significant in the atmosphere.

E. NITROUS ACID

At present, the quantitative data on absorption strengths etc. needed to fully evaluate this possibility are not available.

99

Cb (1)

o3 ._Q

5.85 _

"7,

4.55 -

t.,.)

E. N I T R O U S

__r 0 E

ACID

3.25

O4

As discussed in Chapter 1, nitrous acid is important in tropospheric chemistry because it photolyzes to form OH: HONO + h v ~ OH + NO. (15) Measuring experimentally the quantum yields for reaction (15) as well as the absorption cross sections has been difficult because of contamination of the HONO by other, strongly absorbing species, particularly NO2, necessitating corrections for their contributions to the absorption. However, it is believed, based on studies by Cox and Derwent (1976/1977) that the quantum yield for reaction (15) is 1.0 at A < 400 nm. Figure 4.14 shows the absorption spectrum of gasphase nitrous acid and Table 4.13 gives absorption

T A B L E 4.13

Wavelength (nm) 310 311 312 313 314 315 316 317 318 319 320 321 322 323 324 325 326 327 328 329 330 331 332 333 334 335 336 337 338

10 2o(r

b

2 o

1.95 .650 I

310

330

350

I

370

390

;L (nm)

F I G U R E 4 , 1 4 Absorption spectrum of HONO at 277 K (adapted from Bongartz et al., 1991). Note that the absolute values of the cross sections shown here should be multiplied by 0.855 as recommended by Bongartz et al. (1994).

cross sections (Bongartz et al., 1991) corrected by the factor of 0.855 as recommended by Bongartz et al. (1994).

Recommended Absorption Cross Sections (Base e) for H O N O a

(cm 2 molecule- l)

Wavelength (nm)

(cm 2 molecule - l)

10 2o(r

Wavelength (nm)

1.1 1.6 2.4 1.9 3.1 2.6 1.2 2.7 4.8 3.1 4.2 6.7 4.2 4.4 6.1 4.3 2.5 5.6 10.0 5.2 9.5 15.3 7.4 6.5 8.2 8.2 6.2 4.5 8.6

339 340 341 342 343 344 345 346 347 348 349 350 351 352 353 354 355 356 357 358 359 360 361 362 363 364 365 366 367

16.1 8.6 14.5 33.0 12.7 8.3 9.3 10.5 8.9 7.8 6.8 9.6 18.1 13.3 16.3 49.7 31.1 12.1 10.0 10.3 8.9 7.7 7.1 6.8 8.2 12.5 14.4 15.6 25.8

368 369 370 371 372 373 374 375 376 377 378 379 380 381 382 383 384 385 386 387 388 389 390 391 392 393 394 395 396

10 2oo" (cm 2 molecule - l)

44.5 33.2 15.2 9.7 8.6 6.6 5.3 4.5 4.5 4.3 5.0 6.8 8.2 9.7 13.6 18.0 20.6 17.4 11.5 7.7 4.8 2.9 2.3 1.7 1.3 0.94 0.5 0.86 0.3

a From DeMore et al. (1997) recommendation based on Bongartz et al. (1991), but multiplied by 0.855 as recommended by Bongartz et al. (1994). The value at 354 nm is in excellent agreement with the measured value of (5.02 + 0.76) • 10 -19 cm 2 molecule- 1 of Pagsberg et al. (1997).

100


F. P E R O X Y N I T R I C

TABLE 4.14 Recommended Absorption Cross Sections (Base e) for Gaseous H O 2 N O 2 a

ACID

Peroxynitric acid (HO2NO 2) is a relatively unstable species that is important as a reservoir for NO 2 at lower temperatures via the reversible reaction (16, - 16): HO2NO

(16,- 16)

2 H O 2 + N O 2 .

Figure 4.15 and Table 4.14 give the absorption cross sections. They do not show a significant temperature dependence down to 253 K (Singer et al., 1989). There are a variety of potential photodissociation pathways, including the following:

(17a)

HO2NO 2 + h v ~ HO 2 + NO2, HO 2 + NO + O,

(17b)

OH + NO 2 + O,

(17c)

OH + NO + 02,

(17d) (17e)

-~ OH + NO 3.

G) oO

3.0

-Q "7ID = O

2.4

E

1.2

G) O

O4

E

o 1o

o0

O

_

0 210

11.25

-Q `7 G)

9.00

1

240

270

l

1

300

1

1

1

330

6.75 4.50

E

o t3 c~ O4

o .r---

10 20o" (cm 2 molecule- l)

190 195 200 205 210 215 220 225 230 235 240 245 250 255

1010 816 563 367 239 161 118 93.5 79.2 68.2 58.1 48.9 41.2 35.0

260 265 270 275 280 285 290 295 300 305 310 315 320 325

28.5 23.0 18.1 13.4 9.3 6.2 3.9 2.4 1.4 0.9 0.5 0.3 0.2 0.1

a From DeMore et al. (1997) recommendation based on Molina and Molina (1981) and Singer et al. (1989).

At 248 nm, the quantum yield for OH is 0.34 _+ 0.16, and the formation of electronically excited NO~ with yields of < 30% was observed (MacLeod et al., 1988). Reaction (17a), where some of the NO~ is formed in an electronically excited state, and reaction (17e), which is the most direct path to OH, are thought to be the important paths. In the absence of additional studies, MacLeod et al. (1988) recommend for application to qD17a = 0 . 6 5 a n d

G. N I T R A T E

o9

E

Wavelength (nm)

qD17e = 0 . 3 5 .

RADICAL

0.6

~" G)

O4

10 20o" (cm 2 molecule- l)

atmospheric s i t u a t i o n s

1.8

Wavelength (nm)

ro G) O

Wavelength (nm)

2.25 0

280.0

292.5

305.0

317.5

330.0

Wavelength (nm)

F I G U R E 4 . 1 5 Absorption spectra of HO2NO 2 at 298 K in the (a) 210- to 300-nm and (b) 280- to 330-nm regions (adapted from Singer et al., 1989).

The nitrate radical, NO3, is an important intermediate in nightime chemistry. Its spectroscopy, photochemistry, and chemistry are reviewed in detail by Wayne et al. (1991) and by Atkinson (1991). As shown in Figure 4.16, NO 3 is unusual in that it absorbs strongly in the red region (620-670 nm) of the visible spectrum, unlike most atmospherically important species whose absorptions typically fall in the UV. Its absorption in this region is banded, which allows its detection and measurement using spectroscopic techniques (see Sections A.ld and A.4a in Chapter 11). Table 4.15 gives the absorption cross sections and quantum yields at 298 K. A number of studies report increased values at 662 nm at lower temperatures (e.g., Ravishankara and Mauldin, 1986; Sander, 1986; Yokelson et al., 1994), while one (Cantrell et al., 1987) finds no change. This is important, since these cross sections are used to derive absolute concentrations of NO 3 in the atmosphere, where the temperature during measurement can vary considerably.

I. NITROUS OXIDE

There are two possible decomposition pathways for NO 3 9 NO 3 + hv--+ NO 2 + O(3p),

(18a)

---> NO + O 2.

(18b)

Figure 4.17 shows the energetics of these pathways, including the possible formation of electronically excited singlet states of 0 2. The threshold for (18a) is 585.5 nm (Johnston et al., 1996). While reaction (18b) is close to thermoneutral overall, there is a substantial energy barrier to the dissociation, 47.3 _+ 0.8 kcal mol-1; the threshold observed for this reaction is 594.5 nm (Johnston et al., 1996). Figure 4.18 shows the results of a reevaluation of the quantum yields as a function of wavelength based on the observed energy thresholds for each channel and on consideration of the potential contributions of rotational and vibrational energy to the dissociation of NO 3 (Davis et al., 1993; Johnston et al., 1996). These are in excellent agreement with the experiments of Orlando et al. (1993), except for the region from 595 to 635 nm, where the experimental quantum yields for (18a) are larger than those predicted based on the model of Johnston et al. (1996). The reasons for the discrepancy in this region are not clear, but Johnston and co-workers suggest several possibilities, including contributions from an as yet unrecognized low-lying electronic state. q~lsa is ~ 1 from 570 to 585 nm and then decreases gradually to zero at 635 nm. On the other hand, as (18a) decreases at longer wavelengths, q~lSb first in-

2500

~"

101

creases to a peak value of ~0.36 at 595 nm and then drops off at longer wavelengths. Also shown is the increase in fluorescence quantum yield from zero at 595 nm. As expected (Chapter 3.A.2), as the quantum yields for photochemical channels decline, that for the photophysical process of fluorescence increases. For a solar zenith angle of 0 ~ the photolysis rate constants are estimated to be in the range 0.17-0.19 s -1 for (18a) and 0.016-0.020 s -1 for (18b) at the earth's surface in the absence of clouds (Orlando et al., 1993; Johnston et al., 1996). The O(3p) that is formed in the predominant path will add to 0 2 to generate 03, which can then react with NO 2 to regenerate NO 3.

H. D I N I T R O G E N P E N T O X I D E Figure 4.19 and Table 4.16 show the absorption spectrum and cross sections for N20 5 at room temperature (Harwood et al., 1993, 1998). It should be noted that while these are in excellent agreement with the recommendations of DeMore et al. (1997) above 260 nm, the Harwood et al. (1998)values are smaller than the recommended ones below this. The photolysis produces NO 3 with a quantum yield of approximately unity above 300 nm (DeMore et al., 1997). Harwood et al. (1998) report quantum yields for NO 3 of 0.96 _ 0.15 at 308 nm and 1.03 _ 0.15 at 352.5 nm. However, at 248 nm, the quantum yield for NO 3 drops to 0.64 _ 0.10; O(3p) production also occurs at 248 nm, with a quantum yield of 0 . 7 2 _ 0.17 (Ravishankara et al., 1986). The O(3p) generation could be due to the subsequent dissociation of an initially produced electronically excited NO~ (Oh et al., 1986), to reaction (19b), or to the dissociation of vibrationally excited NO 3 (Harwood et al., 1998):

2000

(I) o')

N205 + hv--+ NO 3 + NO~,

(19a)

..Q "7"

r

---> NO 3 + NO + O.

1500

(19b)

O O

E E

1000

I. N I T R O U S OXIDE

O

(D Od (D -r--

500

0 600

i

1 620

1

1 640

660

;L (nm)

FIGURE 4.16 Absorption spectrum of NO3 at 298 K [adapted from DeMore et al., 1997 based on data from Ravishankara and Mauldin (1986), Sander (1986), and Canosa-Mas et al. (1987)].

Figure 4.20 shows the absorption spectrum of N20 at room temperature and Table 4.17 the absorption cross sections. It does not absorb significantly in the "actinic" region above 290 nm, but at much shorter wavelengths found in the stratosphere, absorbs and dissociates with unit quantum yield to N 2 and electronically excited 0 ( 1D): N20 + hv--+ N 2 + O(1D).

(20)

102


TABLE 4.15

A b s o r p t i o n Cross Sections (Base e) and Q u a n t u m Yields for N O 3 at 2 9 8 K a

10 20 tr

Wavelength

10 20 tr

(nm)

(cm 2 molecule- l)

NO 2 + O

NO + 0 2

(nm)

(cm 2 molecule- l)

NO 2 + O

NO + 0 2

0.0 0.015 0.039 0.097 0.128 0.190 0.220 0.249 0.303 0.328 0.359 0.357 0.318 0.323 0.314 0.291 0.296 0.291 0.283 0.280 0.264 0.271 0.268 0.250 0.248 0.236 0.205 0.200 0.190 0.166 0.166 0.160 0.141 0.143 0.139 0.131 0.127 0.122 0.117 0.106 0.0985 0.092 0.085

628 629 630 631 632 633 634 635 636 637 638 639 640 641 642 643 644 645 646 647 648 649 650 651 652 653 654 655 656 657 658 659 660 661 662 663 664 665 666 667 668 669 670

702 672 638 470 344 194 142 128 159 191 193 162 121 99 91 93 92 85 72 69 60 51 49 52 55 61 76 93 131 172 222 356 658 1308 2000 1742 1110 752 463 254 163 113 85

0.0339 0.0294 0.0264 0.0236 0.0195 0.0177 0.0161 0.0146 0.0119 0.0107 0.0096 0.0086 0.0072

0.074 0.070 0.065 0.058 0.051 0.047 0.043 0.037 0.032 0.029 0.027 0.024 0.020

258 263 302 351 413 415 322 225 170 153 192 171 202 241 242 210 190 189 208 229 292 450 941 1407 1139 796 703 715

0.983 0.967 0.943 0.885 0.854 0.793 0.763 0.734 0.680 0.654 0.608 0.587 0.567 0.531 0.509 0.472 0.438 0.415 0.371 0.351 0.323 0.296 0.280 0.259 0.238 0.226 0.210 0.193 0.181 0.166 0.147 0.137 0.124 0.108 0.0993 0.0897 0.0769 0.0704 0.0643 0.0552 0.0487 0.0442 0.0393

585 586 587 588 589 590 591 592 593 594 595 596 597 598 599 600 601 602 603 604 605 606 607 608 609 610 611 612 613 614 615 616 617 618 619 620 621 622 623 624 625 626 627

Quantum yield

Quantum yield

Wavelength

a Absorption cross sections from DeMore et al. (1997) recommendation based on Ravishankara and Mauldin (1986), Sander (1986), and Canosa-Mas et al. (1987); quantum yields from Johnston et al. (1996).

It is this process, followed by reaction (21a) O(1D) + N 2 0 --+ 2NO, --+ N 2 + 02,

j. O R G A N I C N I T R A T E S A N D PEROXYACETYL NITRATE (21a) (21b)

that is primarily responsible for the production of reactive oxides of nitrogen in the stratosphere in a "natural" atmosphere (see Chapter 12).

1. Organic Nitrates Figure 4.21 shows the absorption spectrum of some simple alkyl nitrates and Tables 4.18, 4.19, and 4.20 give some typical absorption cross sections (Roberts and Fajer, 1989; Turberg et al., 1990; Clemitshaw et al.,

103

K. SULFUR DIOXIDE AND SULFUR TRIOXIDE

NO2(2A1) + O(3p)

50 ..-..

7

,

40

O E -~ 30

No(2n) + 02 (~z~)

E =

NO(2U) + 02 (lAg)

O

0 '8 -

--~ (l) ~_

" ..... .NO 2+O .~

0 .6 -

9

0

t'-

B m

I~

.~"

,**- , ~

....---

Fluorescence

., "

0.2-

>- 20

590

600

610

620

630

640

X(nm)

L NO(2H) + 02 (3Zg)

m

"

*~176176

0 4 - NO + O2 ""...."

r x_

ku 10

".~

F I G U R E 4 . 1 8 Quantum yields for NO 3 photolysis: dotted line, NO 3 ~ NO 2 + O; solid line, NO 3 ~ NO + O2; dashed line, fluorescence quantum yields (adapted from Johnston et al., 1996).

NO3(2A'2)

F I G U R E 4 . 1 7 Energetics of NO 3 photodissociation (adapted from Davis et al., 1993).

Photolysis by several pathways is possible: 1997). The absorption drops off strongly in the actinic region, with the result that they are not photodissociated rapidly in the troposphere (Roberts and Fajer, 1989; Turberg et al., 1990; Clemitshaw et al., 1997). There are several different possible primary processes, depending on the photolysis wavelength. For example, ethyl nitrate in principle has three paths: C2HsONO 2 + hv ~ C2H50 + NO2,

(22a)

CH3CHO + HONO, (22b) C2HsONO + O.

CH3C(O)OONO

2

CH3C(O)O

(23a)

+ N O 3,

CH3C(O)OO CH3C(O)

+ N O 2, + 0 2 nt- N O 2.

(23b) (23c)

The wavelength thresholds for these at 298 K are 1025, 990, and 445 nm, respectively. Mazely et al. (1995, 1997) have measured the quantum yields for production of NO 2 and NO 3 at 248 nm and find ~h(NO2) = 0.83 _+ 0.09 and ~h(NO3) = 0.3 _+ 0.1. The quantum yields at longer wavelengths have not been reported.

(22c)

While the relative importance of the various paths is not well established, it is expected that dissociation to the alkoxy radical, RO, and NO 2 will predominate. Luke et al. (1989) experimentally measured rates of photolysis of simple alkyl nitrates and compared them to rates calculated using the procedures outlined in Chapter 3.C.2. Figure 4.22 compares the experimentally determined values of the photolysis rate constants (kp) for ethyl and n-propyl nitrate with the values calculated assuming a quantum yield for photodissociation of unity. The good agreement suggests that the quantum yield for photodissociation of the alkyl nitrates indeed approaches 1.0.

K. S U L F U R D I O X I D E A N D SULFUR TRIOXIDE 1. S O 2 As seen in Fig. 4.24, SO2 absorbs light strongly up to 300 nm, with a much weaker absorption from 340 to

10-17 10-18 g9 03

a,-

2. P e r o x y a c e t y l N i t r a t e Peroxyacetyl nitrate, CH3C(O)OONO2, commonly referred to as PAN, is important as a means of transporting NO X over large distances. It is quite stable at low temperatures but decomposes at higher temperatures to release N O 2. Hence whether photolysis can compete with the thermal decomposition is of great interest. Figure 4.23 shows the absorption spectrum of PAN, CH3CO3NO2, and Table 4.21 the absorption cross sections (Talukdar et al., 1995).

10-19

O O

E E O

10-20

Od v

13

10-21 10-22 200

240

280

320

360

400

X(nm) F I G U R E 4 . 1 9 Absorption cross sections for N205 at room temperature (adapted from Harwood et al., 1998).

104

4. PHOTOCHEMISTRY OF IMPORTANT ATMOSPHERIC SPECIES TABLE 4.16 Wavelength (nm)

208 210 212 214 216 218 220 222 224 226 228 230 232 234 236 238 240 242 244 246 248 250 252 254 256 258 260 262 264 266 268 270

A b s o r p t i o n C r o s s S e c t i o n s (Base e) for N 2 0 5 a

10 20 o" (cm 2 m o l e c u l e - l)

Wavelength (nm)

418 380 335 285 236 196 165 140 119 105 92.6 83.8 76.9 70.8 65.8 61.4 57.1 53.1 49.3 45.6 41.9 38.6 35.5 32.6 29.9 27.5 25.2 23.1 21.1 19.4 17.8 16.2

272 274 276 278 280 282 284 286 288 290 292 294 296 298 300 302 304 306 308 310 312 314 316 318 320 322 324 326 328 330 332 334

10 20 o" (cm 2 molecule - l)

Wavelength (nm)

10 20 o" (cm 2 m o l e c u l e - l)

336 338 340 342 344 346 348 350 352 354 356 358 360 362 364 366 368 370 372 374 376 378 380 382 384 386 388 390 392 394 396 398

0.462 0.412 0.368 0.328 0.293 0.262 0.234 0.210 0.188 0.167 0.149 0.133 0.120 0.107 0.0958 0.0852 0.0763 0.0685 0.0613 0.0545 0.0484 0.0431 0.0383 0.0341 0.0305 0.0273 0.0242 0.0215 0.0193 0.0172 0.0150 0.0134

14.9 13.7 12.4 11.4 10.5 9.59 8.74 7.94 7.20 6.52 5.88 5.29 4.75 4.26 3.81 3.40 3.03 2.70 2.40 2.13 1.90 1.68 1.49 1.33 1.18 1.05 0.930 0.826 0.735 0.654 0.582 0.518

a From Harwood et al. (1998).

400 nm. T h e strong a b s o r p t i o n p e a k i n g at forms two singlet excited states of S O 2,

30 28

~ 290 n m

26 %" G)

-Q "7

o o

E c~E o D ~N o T-

24

S O 2 ( X 1 A 1 ) -}- hv(240 < A < 330 n m ) --+

22 802(1A2,1B1),

20 18 16

while the w e a k a b s o r p t i o n transition to a triplet state:

14 12 10

involves

(24)

a spin-forbidden

S O 2 ( X l A 1 ) nt- hv(340 < A < 400 nm) --+ 802(3B1).

(25)

8

6 4--

0

SO 2 can J I J I J I J I J IiI

iI

iI

202 204 206 208 210 212 214 216 218

dissociate

to SO + O only at w a v e l e n g t h s

below 218 nm"

SO2(Xlal )

+

hl,,(l~

0.95 at 363 nm (Benter et al., 1995).

Ingham et al. (1998) calculate that at large solar zenith angles, this weak absorption could decrease the lifetime with respect to photolysis by almost a factor of two.

2.5 .,..-..., (,9 .1:2

T

r

2.0 0 07

1.5

o

0,11

o

E

E

0,08 O, 06

1.0

o

A

0,05

0,04

b--

~o

0.5

0 310

1 320

1 330

340

350

360

370

380

X(nm)

4.36 Absorption spectrum of BrO at room temperature (adapted from Wahner et al., 1988).

FIGURE

w. paOBLEVS %" G) 6'3

"T

"5

1.5

O4

E o

1.0

t3

0.5

v-o ,i--

Absorption Cross Sections (Base e) Averaged over 5-nm Intervals in BrO and IO a

4-0

2.5

2.0

E

T A B L E 4.31

3.0

----r o G)

0.0

340

Wavelength region

6-0

frill

360

380

400

420

440

X(nm)

Absorption spectrum of IO at room temperature (adapted from Laszlo et al., 1995).

FIGURE 4.37

Figure 4.41 tions for H O I O H + I, with 1 at 355 nm

117

shows one set of absorption cross sec(Bauer et al., 1998). This dissociates to a q u a n t u m yield for O H production of (Bauer et al., 1998).

U. NITROSYL CHLORIDE (C1NO) A N D NITRYL CHLORIDE (C1NO 2) Figures 4.42 and 4.43 show the absorption spectra of nitrosyl chloride (C1NO) and nitryl chloride (C1NO2), respectively (Roehl et al., 1992; Ganske et al., 1992), and Tables 4.35 and 4.36 list the r e c o m m e n d e d room t e m p e r a t u r e absorption cross sections ( D e M o r e et al., 1997). In both cases, chlorine atoms are produced with unit q u a n t u m yield: C1NO+hv~

CI+NO,

C1NO 2 + h v ~ C1 + N O 2.

(37) (38)

V. HALOGENATED METHANES A N D ETHANES Figures 4.44 and 4.45 show absorption spectra of some simple chlorofluoro-methanes and ethanes, respectively (Hubrich and Stuhl, 1980). Tables 4.37 and 4.38 give the r e c o m m e n d e d absorption cross sections for some of these compounds ( D e M o r e et al., 1997). None of these compounds absorb in the actinic region above 290 nm, but do around 180-200 nm, wavelengths only found in the stratosphere. As discussed in Chapter 12, it is photolysis at these short wavelengths to generate atomic chlorine that is responsible, along with bromine and perhaps in some cases, iodine atoms, for the chain destruction of stratospheric ozone.

1018 O.av (cm z molecule- l)

(nm)

BrO

300-305 305-310 310-315 315-320 320-325 325-330 330-335 335-340 340-345 345-350 350-355 355-360 360-365 365-370 370-375 375-380 380-385 385-390 390-395 395-400 400-405 405-410 410-415 415-420 420-425 425-430 430-435 435-440 440-445 445-450 450-455 455-460 460-465 465-470

2.00 2.59 4.54 3.91 6.00 7.53 6.28 5.89 5.15 3.99 2.28 1.72 1.61 0.92 0.51

IO

1.5 2.9 3.3 3.7 4.2 4.9 3.8 5.6 6.0 6.7 7.6 7.8 9.4 9.3 11.1 8.3 11.4 6.1 11.7 3.0 14.1 4.0 10.0 4.2 2.8

a From DeMore et al. (1997) recommendation based on Laszlo et al. (1995) and Cox and Coker (1983).

Tables 4.39 and 4.40 give the r e c o m m e n d e d absorption cross sections of some hydrochlorofluorocarbons at room temperature, and those of some brominated organics are found in Table 4.41 ( D e M o r e et al., 1997). Finally, alkyl iodides, some of which have natural sources, are of interest for both the troposphere and stratosphere. Figure 4.46 shows the absorption spectra of some simple alkyl iodides and Table 4.42 the absorption cross sections for CH3I (Roehl et al., 1997).

W. PROBLEMS 1. Small amounts of 0 3 are generated by photolysis of 0 2 at 248 nm, even though this is beyond the

118

4. PHOTOCHEMISTRY OF IMPORTANT ATMOSPHERIC SPECIES TABLE 4.32

R e c o m m e n d e d A b s o r p t i o n C r o s s Sections (Base e) of C1OOC1 at 2 0 0 - 2 5 0 K a 10

TMo"

,k

10

TMo"

,k

10

TMo"

,k

10 20 o"

,k

(rim)

(cm 2 molecule - 1)

(nm)

(cm 2 molecule - 1)

(nm)

(cm 2 molecule - 1)

(nm)

(cm 2 molecule - 1)

190 192 194 196 198 200 202 204 206 208 210 212 214 216 218 220 222 224 226 228 230 232 234 236 238 240 242 244 246 248 250 252 254

565.0 526.0 489.0 450.0 413.0 383.5 352.9 325.3 298.6 274.6 251.3 231.7 217.0 207.6 206.1 212.1 227.1 249.4 280.2 319.5 365.0 415.4 467.5 517.5 563.0 600.3 625.7 639.4 642.6 631.5 609.3 580.1 544.5

256 258 260 262 264 266 268 270 272 274 276 278 280 282 284 286 288 290 292 294 296 298 300 302 304 306 308 310 312 314 316 318 320

505.4 463.1 422.0 381.4 344.6 311.6 283.3 258.4 237.3 218.3 201.6 186.4 172.5 159.6 147.3 136.1 125.2 114.6 104.6 95.4 87.1 79.0 72.2 65.8 59.9 54.1 48.6 43.3 38.5 34.6 30.7 28.0 25.6

322 324 326 328 330 332 334 336 338 340 342 344 346 348 350 352 354 356 358 360 362 364 366 368 370 372 374 376 378 380 382 384 386

23.4 21.4 19.2 17.8 16.7 15.6 14.4 13.3 13.1 12.1 11.5 10.9 10.1 9.0 8.2 7.9 6.8 6.1 5.8 5.5 4.5 4.1 3.8 3.5 3.2 2.9 2.7 2.4 2.2 2.1 1.9 1.7 1.6

388 390 392 394 396 398 400 402 404 406 408 410 412 414 416 418 420 422 424 426 428 430 432 434 436 438 440 442 444 446 448 450

1.4 1.3 1.2 1.1

1.0 0.92 0.85 0.78 0.71 0.65 0.60 0.54 0.50 0.46 0.42 0.38 0.35 0.32 0.29 0.27 0.25 0.23 0.21 0.19 0.17 0.16 0.15 0.13 0.12 0.11 0.10 0.09

a From D e M o r e et al. (1997) recommendation based on Cox and Hayman, Nature, 332, 796 (1988), D e M o r e and Tschvikow-Roux, J. Phys. Chem., 94, 5856 (1990), Permien et al., Air Pollution Report #17, Environmental Research Program of the CEC, Brussels (1988) and Burkholder et al., J. Phys. Chem., 94, 687 (1990).

1.5

m

"6"

g

1.0

% o

0.5

~o

0.0

m

I

250

I

I

I

I

I

300

I

I

I

I

I

I

I

350

I

I

I

400

I

I

I

I

I

I

I

450

X(nm) FIGURE 1987).

4.38

Absorption spectrum of OC10 at 204 K (adapted from Wahner et al.,

w. PaOBLEMS

03 (1) 09 e~

._-.. (b 9 20

m

O

E

.llr%

10

~ lngham et al., 1998 --.--Orlando and Burkholder, 1995 --Benteretal .. 1995 ---- Rattigan et al., 1996 ....... B a r n e s

e t al., 1 9 9 6

..... Deters et al., 1996

- lilIi

--r=1761761.52~[:t

O O4

- ,~ 3 _ / ~ - ~,,. ~.il 2.5 --,,, ,,,,l~,~,!,

"7 -

"7 (1)

E

1 19

1

O4

E

o.9.1

vo 19 o o4 o

1

t:>

0.5 0 200

I

I

I

250

300

350

;L (nm) F I G U R E 4 , 3 9 Absorption spectrum of HOC1 (adapted from Burkholder, 1993).

T A B L E 4.33 Wavelength (nm)

200 202 204 206 208 210 212 214 216 218 220 222 224 226 228 230 232 234 236 238 240 242 244 246 248 250 252 254 256 258 260

300

400

350

400

450

500

550

X(nm) F I G U R E 4 , 4 0 Absorption spectra of H O B r (adapted from Ingham et al., 1998).

A b s o r p t i o n Cross Sections of H O C 1 at R o o m T e m p e r a t u r e a

10 2o (r (cm 2 m o l e c u l e - l)

Wavelength (nm)

10 2o (r (cm 2 m o l e c u l e - l)

Wavelength (nm)

10 2o o" (cm 2 m o l e c u l e - l)

7.1 6.1 5.6 5.4 5.5 5.7 6.1 6.6 7.5 8.4 9.7 10.9 12.2 13.5 15.0 16.4 17.7 18.7 19.7 20.3 20.7 21.0 20.5 19.6 18.6 17.3 15.9 14.6 13.2 11.8 10.5

262 264 266 268 270 272 274 276 278 280 282 284 286 288 290 292 294 296 298 300 302 304 306 308 310 312 314 316 318 320

9.3 8.3 7.4 6.6 6.0 5.5 5.2 4.9 4.8 4.7 4.8 4.8 4.9 5.1 5.3 5.4 5.6 5.8 5.9 6.0 6.0 6.1 6.0 6.0 5.9 5.7 5.6 5.4 5.1 4.9

322 324 326 328 330 332 334 336 338 340 342 344 346 348 350 352 354 356 358 360 362 364 366 368 370 372 374 376 378 380

4.6 4.3 4.2 3.8 3.5 3.3 3.1 2.7 2.5 2.4 2.1 1.8 1.8 1.7 1.5 1.3 1.3 1.2 1.0 0.8 1.0 1.0 0.9 0.8 0.8 1.0 0.8 0.8 0.6 0.8

a From DeMore et al. (1997) recommendation based on Burkholder (1993).

120

4. PHOTOCHEMISTRY OF IMPORTANT ATMOSPHERIC SPECIES TABLE 4.34 ,k (rim)

A b s o r p t i o n Cross Sections of H O B r

10 20o" (cm 2 molecule - 1)

,k (nm)

10 TMo" (cm 2 molecule - 1)

,k (nm)

10 TMo" (cm 2 molecule - 1)

6.7 5.2 6.7 9.9 14.1 (10.5) 18.9 (14.6) 23.9 (18.7) 28.0 (22.1) 30.4 (24.3) 30.8 (25.0) 28.7 (24.0) 25.2 (21.9) 20.9 (19.1) 16.8 (16.2) 13.8 (13.6) 11.8 (11.8) 10.8 (10.8) 10.6 (10.5)

330 335 340 345 350 355 360 365 370 375 380 385 390 395 400 405 410 415

11.0 (10.8) 11.5 (11.3) 12.0 (11.9) 12.3 (12.3) 12.5 (12.4) 12.2 (12.2) 11.6 (11.5) 10.7 (10.5) 9.6 (9.32) 8.4 (8.00) 7.4 (6.66) 6.2 (5.38) 5.1 (4.22) 4.1 (3.24) 3.3 (2.43) 2.6 (1.80) 2.0 (1.36) 1.6 (1.08)

420 425 430 435 440 445 450 455 460 465 470 475 480 485 490 495 500 505

1.3 (0.967) 1.1 (1.00) 0.92 (1.15) 0.84 (1.40) 0.74 (1.68) 0.71 (1.96) 0.67 (2.18) 0.65 (2.28) 0.61 (2.28) 0.53 (2.14) 0.49 (1.91) 0.40 (1.62) 0.34 (1.30) 0.28 (0.993) 0.21 (0.723) 0.14 (0.502) 0.09 (0.334) 0.05 (0.212)

240 245 250 255

260 265 270 275 280 285 290 295 300 305 310 315 320 325

a From DeMore et al. (1997) recommendation based on Rattigan et al. (1996); values in parentheses are from Ingham et al. (1998).

..--....

0 o9

4

'r

10 -17 _

3

0

.__r o E

2

O4

E 10 -18

tO

~

0

300

350

400

450

X (nm)

F I G U R E 4.41 et al., 1998).

Absorption spectrum of HOI (adapted from Bauer

10 -18

'~

"~

0_19

lo .2o

...,-...

1 0-19 '7 (D o

10 -21 200

10 -20

1

0

E

oJ

E

v

_

I

I

I

I

I

I

350

400

450

500

550

600

650

X(nm)

FIGURE 4.42 et al., 1992).

1

1

[

1

1

[

280

[

1

[

1

320

1

,

1

1

360

[

1

[

400

F I G U R E 4.43 Absorption spectrum of C1NO 2 at 298 K (adapted from Ganske et al., 1992).

_

10 -22 300

[

240

;L(rim)

10 -21 _

0

tO

1

Absorption spectrum of C1NO (adapted from Roehl

w. PaOBLEMS TABLE 4.35

121

R e c o m m e n d e d A b s o r p t i o n C r o s s Sections (Base e) of C 1 N O at R o o m T e m p e r a t u r e a

Wavelength (nm)

10 20 o" (cm 2 m o l e c u l e - 1 )

Wavelength (nm)

10 20 o" (cm 2 m o l e c u l e - 1 )

Wavelength (nm)

10 20 o" (cm 2 m o l e c u l e - 1 )

Wavelength (nm)

10 20 o" (cm 2 m o l e c u l e - 1 )

190 192 194 196 198 200 202 204 206 208 210 212 214 216 218 220 222 224 226 228 230 232 234 236 238 240 242 244

4320 5340 6150 6480 6310 5860 5250 4540 3840 3210 2630 2180 1760 1400 1110 896 707 552 436 339 266 212 164 120 101 82.5 67.2 55.2

246 248 250 252 254 256 258 260 262 264 266 268 270 272 274 276 278 280 282 284 286 288 290 292 294 296 298 300

45.2 37.7 31.7 27.4 23.7 21.3 19.0 17.5 16.5 15.3 14.4 13.6 12.9 12.3 11.8 11.3 10.7 10.6 10.2 9.99 9.84 9.71 9.64 9.63 9.69 9.71 9.89 10.0

302 304 306 308 310 312 314 316 318 320 322 324 326 328 330 332 334 336 338 340 342 344 346 348 350 355 360 365

10.3 10.5 10.8 11.1 11.5 11.9 12.2 12.5 13.0 13.4 13.6 14.0 14.3 14.6 14.7 14.9 15.1 15.3 15.3 15.2 15.3 15.1 15.1 14.9 14.2 13.6 12.9 12.0

370 375 380 385 390 395 400 405 410 415 420 425 430 435 440 445 450 455 460 465 470 475 480 485 490 495 500

11.0 9.95 8.86 7.82 6.86 5.97 5.13 4.40 3.83 3.38 2.89 2.45 2.21 2.20 2.20 2.07 1.87 1.79 1.95 2.25 2.50 2.61 2.53 2.33 2.07 1.78 1.50

a From DeMore et al. (1997) recommendation based on Roehl et al. (1992) from 350-500 nm and Tyndall et al., J. Photochem., 36, 133 (1987) from 190-350 nm.

TABLE 4.36

R e c o m m e n d e d A b s o r p t i o n C r o s s Sections (Base e) of G a s e o u s C 1 N O 2 at 2 9 8 K a

Wavelength (nm)

10 20 o" (cm 2 m o l e c u l e - l )

Wavelength (nm)

10 20 o" (cm 2 m o l e c u l e - l )

200 210 216 220 230 240 250 260 270

468 320 348 339 226 133 90.6 61.3 35.3

280 290 300 310 320 330 340 350 360 370

22.0 17.3 14.9 12.1 8.87 5.84 3.54 2.04 1.15 0.69

a From DeMore et al. (1997) recommendation based on Ganske et al. (1992).

122


10-17 o3 c~

10-19 0 0

E

10-21

Ckl

E

0 v

1023 I

160

180

200

I

I

I

220

I

I

240

[

I

I

I

260

I

I

I

280

Z(nm)

FIGURE 4.44

Absorption cross sections of the halogenated methanes at 298 K: +, CHC13; zx, CHC12F; • CHC1F2; 0 , CH2C12; ~, CH2C1F; O, CH3C1; D, CC14; A, CC13F (F-11); v , CC12F2 (F-12); m, CC1F3 (adapted from Hubrich and Stuhl, 1980).

R e c o m m e n d e d A b s o r p t i o n Cross Sections of Some Chlorofluorocarbons (Units of 10 - 20 cm 2 m o l e c u l e - 1, Base e) at R o o m T e m p e r a t u r e a

T A B L E 4.37

Wavelength (nm)

CCI 4

CCI 3 F

170 180 190 200 210 220 230 240 250 260 270

772 144 64.8 46.6 17.0 4.07 0.784 0.183 0.0253 0.0061

( CFC- 11) 316 308 178 64.7 15.4 2.42 0.35 0.0464 0.00661 0.00147

CCI 2 F2 (CFC- 12)

CHCIF2 (HCFC-22)

CH 3CI

124 173 62.8 8.84 0.80 0.068 0.0055 0.00029

1.91 0.245 0.032

12.7 1.76 0.206

CH 3CCI 3

192 81.0 24.0 4.15 0.700 0.102

a From DeMore et al. (1997) recommendations based on Hubrich and Stuhl (1980), Hubrich et al., Ber. Bunsenges. Phys. Chem., 81, 437 (1977), Vanlgethem-Meur6e et al., Bull. Cl. Sci., Acad. R. Belg., 64, 31, 42 (1978); Geophys. Res. Lett., 6, 451 (1979), Green and Wayne, J. Photochem., 6, 375 (1976/77) and Simon et al., J. Atmos. Chem., 7, 107 (1988).

10-17

o3

T A B L E 4 . 3 8 R e c o m m e n d e d A b s o r p t i o n Cross Sections of Some Chlorofluorocarbons (Units of 10 - 2o cm2 m o l e c u l e - 1, Base e ) at R o o m T e m p e r a t u r e a

%

10-19

"-7 (3 0

E

10-21

E

0 v

L

5 10-23

~l

160

i

I

i

I

180

i

I

i

I

200

i

I

i

I

220

i

\lllll 240

?~(nm) F I G U R E 4.45 Absorption spectra of some chlorinated ethanes vs the wavelength at 298 K: ~, CH3CC13; v , CF3CH2C1; zx, CH3CH2C1; D, CF3CC1F2; v , CH3CC1F2; O, CC1F2CC1F2; m, CC12FCC1F2 (adapted from Hubrich and Stuhl, 1980).

Wavelength (nm)

CF2CICFCI 2 (CFC-113)

180 190 200 210 220 230

64.5 12.5 1.8 0.220 0.027

CF2CICF2CI (CFC-114)

CF3CF2CI (CFC-115)

26 5.35 0.80 0.104 0.012

1.45 0.27 0.044

a From DeMore et al. (1997) recommendations based on Simon et al., Annales Geophysicae, 6, 239 (1988).

W. PROBLEMS T A B L E 4.39

123

R e c o m m e n d e d A b s o r p t i o n Cross Sections (Base e) of H y d r o c h l o r o f l u o r o e t h a n e s at 2 9 8 K a 10 20 o" (cm 2 molecule - l)

Wavelength (nm)

CH 3CFCI 2 (HCFC- 14 lb )

CH 3CF2 CI (HCFC- 142b )

CF 3CHCI 2 (HCFC- 123)

CF 3CHFCI (HCFC- 124)

190 192 194 196 198 200 202 204 206 208 210 212 214 216 218 220

83.8 64.1 47.4 34.0 23.8 16.4 11.1 7.4 4.9 3.2 2.1 1.4 0.89 0.57 0.37 0.24

0.94 0.66 0.46 0.31 0.21 0.14 0.09 0.061 0.039 0.026 0.017 0.010 0.007 0.004 0.003 0.002

59.0 44.5 32.9 23.6 16.9 11.9 8.3 5.7 4.0 2.7 1.8 1.3 0.87 0.61 0.40 0.28

0.77 0.55 0.39 0.27 0.18 0.13 0.086 0.060 0.040 0.027 0.019 0.012 0.008 0.006 0.004 0.003

a From DeMore et al. (1997) recommendations based on Fahr et al., J. Geophys. Res., 98, 20467 (1993) for -141b, Gillotay and Simon, J. Atmos. Chem., 12, 269 (1991); ibid, 13, 289 (1991), and Orlando et al., J. Geophys. Res., 96, 5013 (1991) for -142b, -123, and -124.

T A B L E 4 . 4 0 R e c o m m e n d e d A b s o r p t i o n C r o s s Sections (Base e ) of C F 3 C F 2 C H C 1 2 ( H C F C - 2 2 5 c a ) and CF2C1CF2CHFC1 ( H C F C - 2 2 5 c b ) a 10 20 o" (cm 2 m o l e c u l e - l)

Wavelength (nm)

CF 3CF 2CHCI 2 (HCFC-225ca)

CF 2CICF2CHFCI (HCFC-225cb)

160 165 170 175 180 185 190 195 200 205 210 215 220 225 230 235 239

269 197 183 191 177 129 74 37 16 6.9 2.9 1.2 0.46 0.17 0.065 0.025 0.011

188 145 91 47 21 9.1 3.5 1.4 0.63 0.33 0.25

a From DeMore et al. (1997) recommendations based on Braun et al., J. Geophys. Res., 96, 13009 (1991).

124

4. PHOTOCHEMISTRY OF IMPORTANT ATMOSPHERIC SPECIES TABLE 4.41 ,k

R e c o m m e n d e d A b s o r p t i o n C r o s s Sections (Base e) for Some B r o m i n e C o m p o u n d s at R o o m T e m p e r a t u r e a 102~ or (cm2 molecule - 1)

(nm)

CH3Br

CHBr 3

CF2CIBr (Halon-1211)

190 192 194 196 198 200 202 204 206 208 210 212 214 216 218 220 222 224 226 228 230 232 234 236 238 240 242 244 246 248 250 252 254 256 258 260 262 264 266 268 270 272 274 276

44 53 62 69 76 79 80 79 77 73 67 61 56 49 44 38 32 28 23 19 15 12 9.9 7.6 5.9 4.5 3.3 2.5 1.8 1.3 0.96 0.69 0.49 0.34 0.23 0.16

399 360 351 366 393 416 433 440 445 451 468 493 524 553 574 582 578 558 527 487 441 397 362 324 295 273 253 234 214 194 174 158 136 116 99 83 69 57 47 38 31 25 20 16

47 58 70 83 96 112 118 121 122 121 117 112 106 98 90 81 72 64 56 49 42 36 31 26 22 18 15 12 10 8.0 6.5 5.1 4.0 3.2 2.4 1.9 1.4 1.1 0.84 0.63 0.48 0.36 0.27 0.20

CF3Br (Halon-1301) 6.4 7.5 8.5 9.5 10.4 11.2 11.8 12.2 12.4 12.4 12.0 11.4 10.7 9.8 8.8 7.7 6.7 5.7 4.7 3.8 3.1 2.4 1.9 1.4 1.1 0.81 0.59 0.43 0.31 0.22 0.16 0.11 0.076 0.053 0.037 0.026 0.018 0.012 0.009 0.006

h (nm)

278 280 282 284 286 288 290 292 294 296 298 300 302 304 306 308 310 312 314 316 318 320 322 324 326 328 330 332 334 336 338 340 342 344 346 348 350 352 354 356 358 360 362

10 TMor (cm 2 molecule - 1) CH3Br

CHBr 3

CF2CIBr (Halon-1211)

12 9.9 7.8 6.1 4.81 3.75 2.88 2.22 1.70 1.28 0.951 0.719 0.530 0.394 0.298 0.226 0.171 0.127 0.0952 0.0712 0.0529 0.0390 0.0289 0.0215 0.0162 0.0121 0.0092 0.0069 0.0052 0.0040 0.0031 0.0024 0.0018 0.0013 0.0010 0.00080 0.00064 0.00054 0.00046 0.00032 0.00024 0.00017 0.00013

0.15 0.1 0.079 0.058 0.043 0.031

CF3Br (Halon-1301)

a From D e M o r e et al. (1997) recommendations based on Gillotay and Simon, Annales Geophysicae, 6, 211 (1988) for CH3Br, Gillotay and Simon, J. Atmos. Chem., 8, 41 (1989) and Burkholder et al., J. Geophys. Res., 96, 5025 (1991) for Halon-1211 and -1301.

w. pt~OBLEVS

10 -18 ..Q

"Tm0~ 23

0

10 -19 1 0-20

0

E

eq

10 -21

E O t~ 10-22

CH 311~~C3H-~7I

B

I

I

I

I

I

I

I

200 220 240 260 280 300 320 340 360 380 X(nm) F I G U R E 4 , 4 6 Absorption spectra of some simple alkyl iodides (adapted from Roehl et al., 1997).

energetic threshold for 0 2 dissociation. (a) Using the heats of formation in Appendix I, calculate the maximum wavelength at which O 2 can photodissociate. (b) Given that the force constant for the bond in O 2 is k = 1140 N m -1, show that if ground-state O 2 has one vibrational quantum of energy, this will be sufficient to cause dissociation on absorption of 248-nm light. Assume O 2 behaves like a simple harmonic oscillator. 2. Using the force constant for O 2 given in Problem 1 and again assuming a simple harmonic oscillator behavior, calculate the highest possible vibrational level that ground electronic state O 2 could be produced upon photolysis of 0 3 at 290 nm if the oxygen atom generated is (a) ground-state O(3p) or (b) electronically excited O(1D). Using your calculations in Problem lb, calculate the vibrational energy level in which 0 2 can

TABLE 4.42

Absorption Cross Sections (Base e) of CH3I at 298 K a

Wavelength (nm)

10 2ocr (cm z molecule- l)

Wavelength (nm)

10 2ocr (cm z molecule - l)

200 205 210 215 220 225 230 235 240 245 250 255 260 265 270 275

181 9.05 4.27 5.35 6.92 9.05 12.4 20.8 37.9 64.3 93.0 112 113 96.6 70.9 45.9

280 285 290 295 300 305 310 315 320 325 330 335 340 345 350

26.9 14.6 7.54 3.79 1.98 1.07 0.603 0.352 0.212 0.122 0.0724 0.0415 0.0225 0.0131 0.00738

a From Roehl et al. (1997).

125

be generated by 0 3 photolysis at 243 nm. Compare to u' = 26 discussed in the text and comment on potential reasons for any differences. 3. Methyl isothiocyanate (CH3NCS) is used as a soil fumigant. It is also the toxic substance generated when metam sodium, C H 3 N H C S 2 N a + , reacts with water, as was the case near Dunsmuir, California, in the early 1990s, when a train carrying metam sodium derailed into a river. Alvarez and Moore (1994) have measured its absorption cross sections in the actinic region above 290 nm: Wavelength range (nm)

(units of 10 -2o cm 2 molecule 1, base e)

295-300 300-305 305-310 310-315 315-320 320-325 325-330 330-335 335-340

1.70 1.23 0.852 0.563 0.348 0.196 0.096 0.040 0.017

-

Using these data and assuming a quantum yield of unity for the subsequent process forming methyl isocyanide CH3NCS + hu ~ CH3NC + S, estimate the photolysis rate constant and lifetime for this compound at a latitude of 40~ on July 1 two hours before noon. 4. The absorption spectra of bromine compounds are generally shifted toward longer wavelengths compared to those of chlorine compounds. (a) Use the absorption cross sections for C12, BrC1, and Br 2 in Table 4.30 to calculate the factors by which the photolysis rate increases for BrC1 and Br 2 compared to C12 at a solar zenith angle of 20 ~ at the earth's surface on December 15. (b) Do the same calculation for a solar zenith angle of 86 ~. Comment on possible reasons for any differences. 5. The absorption spectra of iodine compounds are generally shifted toward longer wavelengths compared to those of both bromine and chlorine compounds. (a) Use the absorption cross sections for IO and BrO in Table 4.31 to calculate the factor by which the photolysis rate increases for IO compared to BrO at a solar zenith angle of 20 ~ at the earth's surface on July 1. Assume quantum yields of unity. (b) Do the same calculation for a solar zenith angle of 86 ~. Comment on possible reasons for the differences. 6. Burkholder and McKeen (1997) report, based on their measurements of the absorption cross sections for

126


SO B (Table 4.22), that photolysis of this c o m p o u n d will b e c o m e important at higher altitudes, sufficiently so to compete with its reaction with water to form sulfuric acid. Compare the SO B photolysis rate for an overhead sun for an altitude of 0 km to its rate of hydrolysis assuming a collision-controlled value for the effective bimolecular rate constant of 10 - l ~ cm 3 molecule -1 s -1 and 50% R H at 298 K at the earth's surface.

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aEFEaENCES Photolysis of Ozone between 301 and 336 nm: Evidence for the Participation of a Spin-Forbidden Channel," Chem. Phys. Lett., 264, 309-315 (1997). Singer, R. J., J. N. Crowley, J. P. Burrows, W. Schneider, and G. K. Moortgat, "Measurement of the Absorption Cross-Section of Peroxynitric Acid between 210 and 330 nm in the Range 253-298 K," J. Photochem. Photobiol. A: Chem., 48, 17-32 (1989). Slanger, T. G., and P. C. Cosby, "O 2 Spectroscopy below 5.1 eV," J. Phys. Chem., 92, 267-282 (1988). Slanger, T. G., L. E. Jusinski, G. Black, and G. E. Gadd, "A New Laboratory Source of Ozone and Its Potential Atmospheric Implications," Science, 241, 945-950 (1988). Solomon, S., R. W. Portmann, R. W. Sanders, and J. S. Daniel, "Absorption of Solar Radiation by Water Vapor, Oxygen, and Related Collision Pairs in the Earth's Atmosphere," J. Geophys. Res., 103, 3847-3858 (1998). Spicer, C. W., E. G. Chapman, B. J. Finlayson-Pitts, R. A. Plastridge, J. M. Hubbe, J. D. Fast, and C. M. Berkowitz, "Unexpectedly High Concentrations of Molecular Chlorine in Coastal Air," Nature, 394, 353-356 (1998). Steinfeld, J. I., S. M. Adler-Golden, and J. W. Gallagher, "Critical Survey of Data on the Spectroscopy and Kinetics of Ozone in the Mesosphere and Thermosphere," J. Phys. Chem. Ref. Data, 16, 911-951 (1987). Stranges, D., X. Yang, J. D. Chesko, and A. G. Suits, "Photodissociation of Ozone at 193 nm by High-Resolution Photofragment Translational Spectroscopy," J. Chem. Phys., 102, 6067-6077 (1995). Syage, J. A., "Photofragment Imaging of Ozone Photodissociation: 0 3 ~ O(3Pj)+ O2(X,v) at 226 nm," J. Phys. Chem., 99, 16530-16533 (1995). Syage, J. A., "Photofragment Imaging of Ozone Photodissociation: 0 3 ~ O(3Pj) + O2(X, v)," J. Phys. Chem., 100, 13885 (1996a). Syage, J. A., "Photofragment Imaging by Sections for Measuring State-Resolved Angle-Velocity Differential Cross Sections," J. Chem. Phys., 105, 1007-1022 (1996b). Takahashi, K., M. Kishigami, Y. Matsumi, M. Kawasaki, and A. J. Orr-Ewing, "Observation of the Spin-Forbidden O(1D) + O2(X3s Channel in the 317-327 nm Photolysis of Ozone," J. Chem. Phys., 105, 5290-5293 (1996a). Takahashi, K., Y. Matsumi, and M. Kawasaki, "Photodissociation Processes of Ozone in the Huggins Band at 308-326 nm: Direct Observation of 0(102) and O(3Pj) Products," J. Phys. Chem., 100, 4084-4089 (1996b). Takahashi, K., M. Kishigami, N. Taniguchi, Y. Matsumi, and M. Kawasaki, "Photofragment Excitation Spectrum for O(1D) from the Photodissociation of Jet-Cooled Ozone in the Wavelength Range 305-329 nm," J. Chem. Phys., 106, 6390-6397 (1997). Talukdar, R. K., J. B. Burkholder, A.-M. Schmoltner, J. M. Roberts, R. R. Wilson, and A. R. Ravishankara, "Investigation of the Loss Processes for Peroxyacetyl Nitrate in the Atmosphere: UV Photolysis and Reaction with OH," J. Geophys. Res., 100, 14163-14173 (1995). Talukdar, R. K., M. K. Gilles, F. Battin-Leclerc, A. R. Ravishankara, J.-M. Fracheboud, J. J. Orlando, and G. S. Tyndall, "Photolysis of Ozone at 308 and 248 nm: Quantum Yield of O(1D) as a Function of Temperature," Geophys. Res. Lett., 24, 1091-1094 (1997). Talukdar, R. K., C. A. Longfellow, M. K. Gilles, and A. R. Ravishankara, "Quantum Yields of 0 ( 1D) in the Photolysis of Ozone between 289 and 329 nm as a Function of Temperature," Geophys. Res. Lett., 25, 143-146 (1998). Thelen, M.-A., P. Felder, and J. R. Huber, "The Photofragmentation of Methyl Hydroperoxide CH3OOH at 193 and 248 nm in a Cold Molecular Beam," Chem. Phys. Lett., 213, 275-281 (1993).

129

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C H A P T E R

5 Kinetics and Atmospheric Chemis try Understanding the kinetics of reactions of various species in the atmosphere is critical for assessing their atmospheric fates. As such, reaction kinetics form a key component of risk assessments of airborne toxic chemicals, for example (see Chapter 16). In addition, it greatly simplifies the number of reactions that must be considered in assessing the atmospheric fates of a particular compound of interest. In the case of organics, for example, there are, in principle, many potential reactions that could occur in the atmosphere. Given the thousands of organics found in air, the total number of reactions that must be considered quickly becomes intractable unless there is some means to pare the list down. Knowing the kinetics of reaction of various classes of compounds with OH, 03, etc. allows one to estimate lifetimes under typical atmospheric conditions and hence to rule out those reactions that are too slow to be significant, allowing one to concentrate on the most important reactions. In this chapter we briefly review the fundamental kinetics needed for application to atmospheric systems and discuss some of the most common methods for determining rate constants in the laboratory. This includes so-called "heterogeneous" reactions, whose importance in the stratosphere is now well established and which are increasingly being recognized as important in tropospheric systems. For a review of these areas, see Molina et al. (1996).

number of important gas-phase processes in which three different species participate; these are termolecular reactions. In the troposphere they usually involve N 2 a n d / o r 0 2 as one of the three participants; the role of the third molecule is generally to act as an "inert gas" that stabilizes the energy-rich intermediate formed from the recombination of two species by siphoning off the excess energy, thus preventing dissociation back into the reactants. In such cases, rather than being specific as to the colliding third body, the symbol "M" is used. Examples of these three classes of gas-phase reactions are The thermal decomposition of PAN:

Unimolecular.

O

O

II

CH3COONO 2 Bimolecular.

II

> CH3COO + NO 2 .

(1)

Formation of the gaseous nitrate radical: 0 3 +

N O 2 --+

NO 3

+

0 2.

(2)

Terrnolecular. The formation of ozone by the reaction of a ground-state oxygen atom, O(3P), with O 2"

O(3p) + O 2 + M + O

3 +M.

(3)

Because M does not enter into the reaction chemically, such reactions are usually written with "M" above the arrow, although in some cases it is omitted entirely and simply understood to be present:

O(3p) +

A. F U N D A M E N T A L PRINCIPLES OF GAS.PHASE KINETICS

M 0 2 ---> 0 3 .

Each of these types of elementary processes will be treated in some detail in subsequent sections of this chapter. While two-body collisions are common in the gas phase, three-body collisions are much less probable and four-body collisions can essentially be ignored because of their low probability. Thus the majority of the reactions we deal with in the atmosphere are bimolecular, with a lesser number being termolecular or unimolecular.

1. D e f i n i t i o n s a. Elementary vs Overall Reactions Elementary reactions are defined as those that cannot be broken down into two or more simpler reactions. Generally, they consist of one or two reactant species and are referred to as unimolecular and bimolecular processes, respectively. However, there are a

130

A. FUNDAMENTAL PRINCIPLES OF GAS-PHASE KINETICS

An overall reaction includes two or more elementary reactions; indeed there is no limit to the number of reactants or elementary reactions comprising an overall reaction. Thus if a single reaction step as written has four or more reactants, it cannot be an elementary process, and it must occur via two or more consecutive steps. If a reaction step contains two to three reactants, it may, or may not, be an elementary reaction.

b. Rate Laws, Reaction Order, and the Rate Constant In studying any particular reaction, one does not really know a priori if it is an elementary reaction or not, unless it involves four or more species, in which case it cannot be elementary. Determination of the rate law for the reaction is the first step in assessing whether it could be elementary. The rate of a reaction is defined as the change in the concentration of a reactant or product with time. For simple reactions occurring with unit stoichiometry, the rate expressed in terms of reactant disappearance is the same as the rate in terms of product formation. For example, for reaction (4), the reaction of ozone with nitric oxide, NO + 0 3 --+ NO 2 + 02,

(4)

the rate is defined as Rate =

-d[NO]

- d [ O 3]

dt

dt

+ d [ N O 2]

+ d [ O 2]

dt

dt

(A)

For reactions of the more general form

(5)

a A + bB --+ cC + d D ,

where the stoichiometric coefficients a, b, c, and d are not all unity, the rate in terms of disappearance of A may not be equal to the rate in terms of disappearance of B or the appearance of C or D. To take such differences in stoichiometry into account, the rate of the generalized reaction (5) is defined by convention as

131

0 2 reacted, and the rate of loss of NO is twice that of 0 2"

Rate-

1 d[NO]

2

dt

=

-d[O2] dt

= +-

1 d[NO 2] 2

dt

While this convention is now widely used, it was not in some early kinetic studies. Thus one must be careful to note exactly how the rate is defined so that the reported rate constants are interpreted and applied correctly. In systems of atmospheric interest, the rate law or rate expression for a reaction, either elementary or overall, is the equation expressing the dependence of the rate on the concentrations of reactants. In a few reactions (mainly those in solution), products may also appear in the rate law. For the general overall reaction (5), the rate law has the form Rate = k[A]m[B]n[c]P[D]q, where, depending on the mechanism of the reaction, m, n, p, and q may be zero, integers, or fractions. As noted earlier, in most gas-phase atmospheric reactions, the exponents of the product concentration (i.e., p and q) are zero and the rate laws involve only the reactant species. It is important to stress here that in contrast to elementary reactions, in overall reactions the exponents in the rate laws (e.g., m, n, p, and q) do not necessarily bear a relationship to the stoichiometric coefficients of the reaction (e.g., a, b, c, and d). The importance of distinguishing between elementary and overall reactions comes in formulating rate laws. For elementary reactions only, the rate law may be written directly from the stoichiometric equation. Thus for the general elementary gas-phase reaction a A + bB --+ cC + d D ,

(7)

Rate = k[A]a[B] b, where (a + b) < 3 by definition of an elementary reaction. For example, the rate expression for the elementary reaction (4) is given by Rate = k4[NO][O3].

Rate =

1 d[A]

1 d[B]

a

b

dt

1 d[C] =

t

c

dt

dt

1 d[D] =

For example, in the thermal oxygen,

t

d

dt

.

(B)

oxidation of NO by

2NO + 0 2 --+ 2NO2,

(6)

two molecules of NO disappear for each molecule of

The rate constant, k, is simply the constant of proportionality in the expression relating the rate of a reaction to the concentrations of reactants a n d / o r products, each expressed with the appropriate exponent. The order of a reaction is defined as the sum of the exponents in the rate law. Thus reaction (4) is (1 + 1) = second order. The order with respect to each species appearing in the rate law is the exponent of the concentration of that species; thus reaction (4) is first order in both 0 3 and NO.

13 2

5. KINETICS AND ATMOSPHERIC CHEMISTRY

T h e basis of predicting rate laws for e l e m e n t a r y reactions f r o m the stoichiometric e q u a t i o n lies in the fact that they must occur during a single collision (although the probability of reaction during any one collision is equal to or less than unity). Thus doubling the c o n c e n t r a t i o n of 0 3 in reaction (4)will double the n u m b e r of collisions per second of 0 3 with NO. Assuming the probability of reaction per collision remains constant, then the n u m b e r of 0 3 and N O molecules reacting, and 0 2 and N O 2 f o r m e d per unit time (i.e., the rate), must double. T h e t h e r m a l oxidation of N O by m o l e c u l a r oxygen, reaction (6), is a n o t h e r example w h e r e the stoichiometry and the molecularity of the reaction are directly related, and the rate law is Rate

TABLE 5.1

Some Common Conversion Factors for Gas-Phase Reactions

Concentrations a 1 mol L-1 = 6.02 • 1020 molecules cm-3 1 ppm = 2.46 • 1013 molecules cm-3 1 ppb = 2.46 • 101~molecules cm-3 I ppt = 2.46 • 107 molecules cm-3 1 atm = 760 Torr = 4.09 • 10 - 2 mol L- 1 = 2.46 • 1019 molecules cm-3 Second-order rate constants cm3 molecule- 1 s- 1 • 6.02 • 1020 = L mol- 1 s- 1 ppm-1 min- 1 • 4.08 • 105 = L mol- 1 s- 1 ppm-1 min- 1 • 6.77 • 10-16 = cm3 molecule- 1 s- 1 atm -1 s-1 • 4.06 • 10-20 = cm3 molecule-1 s-1 cm 6

Third-order rate constants molecule-2 S-1 X 3.63 • 1 0 41 = L 2 mol-2 - 2 m i n - 1 • 9.97 • 1012 = L 2 m o l - 2 S-1

S-1

ppm p p m - 2 m i n - 1 • 2.75 • 10 -29 = c m 6 m o l e c u l e - 2 s - 1

= k~II[NO]2102].

Thus the rate is p r o p o r t i o n a l to the first power of the oxygen c o n c e n t r a t i o n and the square of the nitric oxide c o n c e n t r a t i o n and the reaction o r d e r is 1 + 2 = 3. However, in the t r o p o s p h e r e , the O 2 c o n c e n t r a t i o n is always so large relative to N O that it is effectively constant and thus can be i n c o r p o r a t e d into the rate constant k~Ix. T h e rate law is now written R a t e = kbi[NO] 2, and the reaction is r e f e r r e d to as pseudo-second-order. W e adopt the convention of writing a t h i r d - o r d e r rate constant a s k Ill, and a p s e u d o - s e c o n d - o r d e r rate constant as k bi, as illustrated in the preceding equations. T h e rate law and the reaction o r d e r can often be used to show that a reaction cannot be an e l e m e n t a r y reaction since, in the latter case, the exponents must be integers and the overall reaction o r d e r must be < 3. However, it should be n o t e d that these kinetic p a r a m e ters cannot be used to confirm that a particular reaction is elementary; they can only indicate that the kinetic data do not rule out the possibility that the reaction is elementary. In gas-phase reactions, concentrations are usually expressed in molecules cm -3 and time in seconds, the convention we employ in this book. Thus the units of k are as follows: first order, s-1; second order, cm 3 molec u l e - 1 s - 1. third order, cm 6 m o l e c u l e - 2 s - 1o C o n c e n t r a t i o n s of gaseous pollutants are often expressed in terms of parts per million (ppm) by volume, and time is expressed in minutes. Use of these concentration units must be reflected in the units used for the rate constants as well; for example, s e c o n d - o r d e r rate constants are in units of p p m -1 min -1. Occasionally, gas concentrations are given in units of mol L-1 or in units of pressure such as Torr, a t m o s p h e r e s , or Pascals; these can be converted to the m o r e conventional units

a The concentrations ppm, ppb, and ppt are relative to air at 1 atm and 25~ where 1 atm = 760 Torr total pressure.

in t r o p o s p h e r i c chemistry using the ideal gas law. Table 5.1 gives s o m e c o m m o n conversion factors for gas-phase concentrations and rate constants at 1 atm pressure (760 T o r r total pressure) and 25~ For solution-phase reactions, we use c o n c e n t r a t i o n units of mol L -1, with units for the c o r r e s p o n d i n g rate constants of L m o l - 1 s-1 (second order) and L 2 m o l - 2 s - 1 (third order). c. Half-Lives and

Lifetimes

A rate constant is a quantitative m e a s u r e of how fast reactions p r o c e e d and t h e r e f o r e is an indicator of how long a given set of reactants will survive in the atmosphere u n d e r a particular set of r e a c t a n t concentrations. However, the rate constant per se is not a p a r a m eter that by itself is readily related to the average length of time a species will survive in the a t m o s p h e r e b e f o r e reacting. M o r e intuitively meaningful p a r a m e ters are the half-life (tl/2) or the natural lifetime (r), the latter usually r e f e r r e d to simply as "lifetime," of a pollutant with respect to reaction with a labile species such as O H or N O 3 radicals. T h e half-life (tl/2) is defined as the time r e q u i r e d for the c o n c e n t r a t i o n of a r e a c t a n t to fall to one-half of its initial value, w h e r e a s the lifetime is defined as the time it takes for the r e a c t a n t c o n c e n t r a t i o n to fall to 1/e of its initial value (e is the base of natural logarithms, 2.718). Both tl/2 and r are directly related to the rate constant and to the concentrations of any o t h e r reactants involved in the reactions. T h e s e relationships are given in general f o r m in Table 5.2 for first-, second-, and t h i r d - o r d e r reactions and are derived in Box 5.1.

133

A. F U N D A M E N T A L PRINCIPLES OF GAS-PHASE KINETICS

TABLE 5.2

Relationships between the Rate Constant, Half-Lives, and Lifetimes for First-, Second-, and Third-Order Reactions

Reaction order

Reaction

Half-life of A

Lifetime of A

kl

First

(1) A

Second

(2) A + B

> Products

> Products k3 (3) A + B + C > Products

Third

A relevant example is the use of lifetimes to characterize the reactivity of organics. Compressed natural gas (CNG), for example, is a widely used fuel whose major component is methane, CU 4. The only known significant chemical loss process for CH 4 is reaction with OH: CH 4 +

OH

~

CH 3 + H20

,

k298 K = 6.3 • 10- 15 cm 3 molecule- 1 s- 1.

(8)

Taking a typical average, daytime OH concentration of 1 • 106 radicals cm -3, the lifetime of CH 4 with respect to this removal process is CH4 =

TOH

ks[OH]

[6.3 • 10 -15 cm 3 molecule -1 s -1 • • 10 6 radicals cm-3)] = 1.59 • l0 s s = 5 years. Propane is another organic used widely as a fuel. It also reacts with OH: C3H s +

t~/2 = 0 . 6 9 3 / k 1

r A = 1/k 1

t~/2 = 0 . 6 9 3 / k z [ B ]

r A=

t~/2 = 0.693/k3[B][C]

r A = 1/k3[B][C]

k2

OH

k298 K = 1.1 • 10-

1/kz[B]

gional photochemical smog formation, as suggested in Mexico City (Blake and Rowland, 1995). Two points should be made about such calculations of tropospheric lifetimes. First, they are valid only for the specified reaction; if there are other competing loss processes such as photolysis, the actual overall lifetime will be shortened accordingly. On the other hand, for a species such as CH4, which does not photolyze or react significantly with other atmospheric species such as O 3 or NO3, roncu4 is indeed close to the overall lifetime of CH 4 .

Second, in bi- and termolecular reactions, tl/2 and r depend on the concentration of other reactants; this is particularly important when interpreting atmospheric lifetimes. For example, as discussed earlier, reaction with the OH radical is a major fate of most organics during daylight in both the clean and polluted troposphere. However, the actual concentrations of OH at various geographical locations and under a variety of conditions are highly variable; for example, its concentration varies diurnally since it is produced primarily by photochemical processes. Finally, the concentration of OH varies with altitude as well, so the lifetime will depend on where in the troposphere the reaction Occurs.

+

C3H 7 + H20

12 c m 3

,

molecule- 1 s- 1.

(9)

Assuming the same OH concentration, one calculates rC3ns = 9.1 • 10 5 S = 10.5 days. OH These calculations illustrate why regulatory agencies have concentrated on controlling "non-methane hydrocarbons." Because methane reacts so slowly in the troposphere, it is generally not of concern from the point of view of ozone formation in urban areas, i.e., over the time scale of hours to a few days. It is also the reason that CH 4 is the only organic to survive long enough in the troposphere to cross the tropopause and enter the stratosphere in significant concentrations (see Chapter 12). On the other hand, propane reacts sufficiently quickly that it can contribute to local and re-

Thus when a lifetime of an organic in the atmosphere is cited with respect to OH attack, one should examine carefully the concentration of OH that was assumed in arriving at that lifetime; the substantial uncertainties in these estimated lifetimes that arise from the uncertainties in the estimated atmospheric OH concentrations should be clearly recognized.

2. Termolecular Reactions and Pressure Dependence of Rate Constants Termolecular elementary reactions, whose rates depend on the total pressure, are important in the atmosphere. Examples include the formation of O 3, O(3p)

+

0 2 nt-

M ~ O 3 + M,

(3)

134


BOX 5.1

DERIVATION

OF

HALF-LIVES

FROM

The expressions for half-lives and lifetimes in Table 5.2 can be readily derived from the rate laws. For a first-order reaction of a pollutant species A, the rate law for the reaction kl

[A] = 0.5[A] 0. Substituting into the integrated rate expression, one obtains In 0.5 0.693 tl/2 =

kl

-

kl

9

For second- and third-order reactions, if one ass u m e s the concentrations o f the reactants other than A are constant with time, the derivation is the same

-d[A]

except that k is replaced by k[B] (second order) or k[B][C] (third order). In most practical situations, however, the concentration of at least one of the other reactants is not constant but changes with time due to reactions, fresh injections of pollutants, and so on. As a result, using half-lives (or lifetimes) of a pollutant with respect to second- or third-order reactions is an approximation that involves a s s u m e d constant concentrations of the other reactants. These half-lives for bimolecular and termolecular reactions are thus directly affected by the concentrations of the other reactant. Derivation of the relationship between the rate constant k and the lifetime r follows that for tl/2, except that, from the definition of r, at t = r, [A] = [A]o/e.

dt

= kl[A]"

Rearranging, this becomes -d[A] = k l dt.

[A]

Integrating from time t = 0 when the initial concentration of A is [A] 0 to time t when the concentration is [A], one obtains [A] In

[A]0

= -kit.

After one half-life (i.e., at t = t l / 2 ) by definition

OH

LIFETIMES

A ~ Products is given by

and the oxidations of reactions:

AND

KINETICS

SO 2

and NO 2 via gas-phase O H

M + S O 2 --->

HOSO2,

(10)

M

O H + NO 2 ~ H O N O 2.

(11)

The reason for the pressure dependence of termolecular reactions can be seen by taking reaction (3) as an example. The exothermic bond formation between O(3p) and 0 2 releases energy that must be removed to form a stable O 3 molecule; if the energy remains as internal energy, the O 3 will quickly fly apart to re-form O + 0 2. The third molecule, M, is any molecule that stabilizes the excited (O3)* intermediate by colliding with it and removing some of its excess internal energy. Treating reaction (3) as an elementary reaction

d[O3] Rate = k3[O][O 2][M] = +

d-----7-'

one might expect the rate to increase with the concentration or pressure of the third body M. However, there clearly must be some limit since the rate cannot increase to infinity but only to some upper limit determined by how fast the two reactive species can combine chemically. As a result, one might intuitively expect the rates of reactions such as (3), (10), and (11) to increase initially as the pressure of M is increased from zero and then to plateau at some limiting value at high pressures. Let us take the reaction (10) of O H with SO2 as an example of a termolecular reaction of atmospheric interest and examine how its pressure dependence is established. It is common in kinetic studies to follow the decay of one reactant in an excess of the second reactant. In the case of reaction (10), the decay of O H is followed in the presence of excess SO 2 and the third body M, where M is an inert " b a t h " gas such as He,

135


Ar, or N 2. Since it is assumed to be an elementary reaction, the rate law for reaction (10) can be written for low pressures" -d[OH] dt

= /~III[OH][SO2

,~1o

][M]

z~III and [M] can be combined to If [M] is constant, ,~10 form an effective bimolecular rate constant, kb~ = klSI[M]:

/,7 III the To obtain the termolecular rate constant ,~10, effective bimolecular rate constant kb~ = kin[ M]10 is plotted in Fig. 5.2 as a function of total pressure (i.e., of [M]). As expected from the earlier discussion, kb~ increases with [M] at low pressures but approaches a plateau at higher pressures. Termolecular reactions can be treated, as a first approximation, as if they consist of several elementary steps, for example, for reaction (10),

-d[On] = /'-III[OH][SO2][M]

dt

'~10

= klb~[OH][SO2]

[OH] In [OH]0

kb~[SO2 ]~

Since the initial concentration of OH, [OH]0, is a constant, a plot of ln[OH] against reaction time t should be a straight line with slope or decay rate given by Decay rate (s- 1) = _ kb~[so2 ]0.

o

H O S O 2 + M.

(13)

HOSO~ is the exited O H - S O 2 adduct that contains the excess internal energy from bond formation in (12), and H O S O 2 is the stabilized adduct resulting when some of this internal energy is removed by a collision with M. If the system is treated as if the concentration of the energized adduct (HOSO~) remains constant with time, then its rates of formation and loss are equal. These rates can be written from Eqs. (12), ( - 1 2 ) , and (13) since these are assumed to be elementary reactions. Thus

= 0 = k a [ O n ] [ s o 2] - k b [ n O S O ~ ]

dt

- kc[HOSO ~ ][M]. This is an example of the steady-state approximation, widely employed in gas-phase kinetics and mechanistic studies.

'~

402 Torr 202Torr 100 Torr

"~ 300--

HOSO~ + M ~

d[HOSO~ ]

A plot of these decay rates against [SO2] 0 should thus be linear, with the slopes increasing with pressure since kb~ depends on [M]. Figure 5.1 shows such a plot of the absolute values of the observed OH decay rates against [SO2] 0 at total pressures of Ar from 50 to 402 Torr (Atkinson et al., 1976). As expected, the decay rates are linear with [SO2] 0 and increase with the pressure of M.

IEl

(12, - 1 2 )

HOSO~,

kb

Since SO2 is in great excess, its concentration does not change significantly even when all the OH has reacted and hence it remains approximately constant throughout the reaction at its initial value, [SO2 ]0. Rearranging the rate law and integrating from time t = 0 when the initial concentration of OH is [OH] 0 to time t when the OH concentration is [OH], one obtains

400~-

OH + SO 2 ~

6

50 Torr

200

%

100

55

o

00

I

0.4

I

I

0.8

I

I

1.2

~o f

10-15[SO2]0 (molecules cm -3) F I G U R E 5.1 Plots of the O H decay rates against the initial SO 2 concentration at total pressures of Ar from 50 to 402 Torr (adapted from Atkinson et al., 1976).

0

200

400

600

800

Total pressure of argon (Torr) F I G U R E 5 . 2 Plot of kb~ against total pressure for M = Ar for the reaction of O H with SO 2 (adapted from Atkinson et al., 1976).

13 6

5. KINETICSAND ATMOSPHERICCHEMISTRY

Rearranging, an expression for [HOSO~ ] is obtained in terms of the reactants OH and SO 2" ka[OH][SO 2]

[HOSO~ ] =

k b -+- kc[M ]

limiting low- and high-pressure rate constants from experimental data as well as to report the temperature and pressure dependence of termolecular reactions in compilations of kinetic data. Equation (C) gives the most commonly used form of the rate constant expression of Troe and co-workers:

The rate of the reaction in terms of product formation is given by

k =

= kc[M][HOSO ~ ], ka[OH][SO 2]

= kc[M ]

=

k b -+- kc[M ]

k~kc[M] } kb -+- kc[M ] [OH][SO2]"

In this form, kb~ can be rationalized by the combination of rate constants and [M] given in brackets. Alternatively, (kb~) -1 is given by 1

kab~

F(cl+[l~176176

(C)

1 + ko[M]/ks

d[HOSO 2]

dt

k0[M]

k b + kc[M]

kb

1

k~kc[M]

kakc[M]

k~

At "infinite" pressure, where 1/[M] = 0, the rate constant should have its high-pressure limiting value. It is seen that this high-pressure limiting value, k~0, is equal to k a. One would also qualitatively expect k~0 = ka from the reaction scheme consisting of (12), ( - 1 2 ) , and (13); thus in the limit of infinite pressure, all the energized adducts formed in (12)will be stabilized in (13) and none will have a chance to decompose back to reactants via ( - 1 2 ) . In this case, the rate constant will just be that for formation of HOSO~, that is, k~. This approximate treatment of termolecular reactions can also be used to examine how the third-order, low-pressure rate constant k III relates to the rate constants k~, k b, and k c for the elementary reactions assumed to be involved. As [M] approaches zero, kb~ approaches k~kc[M]/k b, so that ,~10/zIIIis given by iii k 10

kakc

In equation (C), k 0 (or k III as used earlier) is the low-pressure limiting rate constant and ks is the highpressure limiting rate constant. F c is known as the broadening factor of the falloff curve; its actual value depends on the particular reaction and can be calculated theoretically. Troe (1979) suggests that for reactions under atmospheric conditions, the value of F c will be ~ 0.7-0.9, independent of temperature. However, values as low as 0.4 are often observed. The NASA evaluations of stratospheric reactions (DeMore et al., 1997) take F c = 0.6 for all reactions. The IUPAC evaluation (Atkinson et al., 1997a,b) does not restrict F c to 0.6. However, it is important to note that the values of k 0 and ks will depend on the value of F c used to match the experimental data. For example, for reaction (11) M OH + NO 2 ~ HONO2,

(11)

the JPL evaluation recommends, at 300 K, k 0 = 2.5 • 10 .3o cm 6 molecule -2 s -1 and ks = 1.6 • 10 -11 cm 2 molecule- 1 s- 1 with F c = 0.6, whereas the IUPAC recommendation is k 0 = 2.6 • 10 .3o and ks = 7.5 • 10 -11 with F c = 0.41 (Atkinson et al., 1997b). (Note, however, that more recent studies discussed in Chapter 7.B.1 suggest both of these recommendations may give systematically high rate constants, especially below 240 K.) The temperature dependence of k arises primarily in the temperature dependence of k o and ks, which are expressed in terms of their values at 300 K, ko3~176 and k2 ~176 Thus --/7

kb

Many addition reactions such as the O H - S O 2 reaction are in the falloff region between second and third order in the range of total pressures encountered from the troposphere through the stratosphere. Troe and co-workers have carried out extensive theoretical studies of addition reactions and their reverse unimolecular decompositions as a function of pressure (e.g., see Troe, 1979, 1983). In this work they have developed expressions for the rate constants in the falloff region; these are now most commonly used to derive the

=

3-6-6

and

k2 = k 3~176

3-6-6

Thus, there are a number of parameters (k03~176 k 3~176n m, and F c) that must be known in order to calculate the rate constant throughout the falloff region. Experimentally, these are determined by obtaining a best fit to the experimentally measured values of k using a ,

"vS

,

'

13 7


calculated or assumed value of F c = 0.6, so that rate constants for the termolecular reactions can be calculated as a function of pressure and temperature. For example, for the reaction (10) of O H with SO 2 discussed earlier, k03~176 = (3.0 _+ 1.0) • 10 -31 cm 6 molecule -2 s-1, n = (3.3 _+ 1.5), k 3~176 = (1.5 _+ 0.5) • 10 -12 cm-3 molecule-1 s - l , and m = 0 ( D e M o r e et al., 1997). At 300 K and 760 Torr pressure, [M] = 2.46 • 1019 molecules cm -3 and k0[M] = 7.4 • 10 -12 cm 3 molecule -1 s -1. Thus the value of the rate constant under these conditions is calculated as 7.4 x 10 -12

k =

function of temperature, that is, if k03~176 k 3~176 n, m, and F c are known. At first glance, it might appear that the vast majority of the bimolecular reactions with which one deals in the troposphere are simple concerted reactions, that is, during the collision of the reactants there is a reorganization of the atoms, leading directly to the formation of the products. However, it has become increasingly apparent in recent years that some important reactions that appeared to be concerted exhibit characteristics such as pressure dependencies that are not consistent with a direct concerted process. A classic case is the reaction of O H with CO:

1 + (7.4 • 1 0 - 1 2 ) / ( 1 . 5 X 10 -12) O H + CO --+ H + C O 2 . •

: [1.25 • 10-12

]0.60"68

= 8.8 • 10-13 cm 3 molecule-1 s-1. However, under conditions typical of the lower stratosphere ( ~ 20 km), the t e m p e r a t u r e and pressure are much lower. Let us calculate kl0 for conditions where the t e m p e r a t u r e is 219 K and the total pressure is 39 Torr. First we need to recognize that [M] changes not only with pressure but also with t e m p e r a t u r e according to the ideal gas law: [M] = P / R T . U n d e r these conditions, [M] = 1.7 • 10 is molecules cm-3. The lowand high-pressure limiting rate constants at this temperature are given by Eqs. (D) and (E): ( 2 1 9 ) -33 k g19= 3.0 • 10 -31

376

= 8 . 5 X 10 -31 c m 6 m o l e c u l e - 2 s - 1 k219 oo = 1 . 5 X 1 0 - 1 2 ( 2 139~)

(D)

.0

= 1.5 • 10-12 cm 3 molecule- 1 s- 1.

This reaction appears to be an elementary bimolecular reaction involving a simple transfer of an oxygen atom from O H to CO. In accord with the definition of an elementary reaction, one can imagine that it occurs during one collision of an O H radical with a CO molecule. A n u m b e r of studies of the kinetics of this reaction were carried out in the 1960s and the early 1970s, and the room t e m p e r a t u r e rate constants, measured at total pressures up to ~ 200 Torr in inert gases such as He, Ar, and N2, were generally in good agreement with k14 ~ 1.5 X 10-13 cm 3 molecule-1 s-1 at room temperature. In fact, this reaction was often used to test whether a newly constructed kinetic apparatus was functioning properly. However, a variety of studies since the mid-1970s has established that it is not, in fact, a simple bimolecular reaction as implied by reaction (14) but rather involves the formation of an excited H O C O * intermediate (e.g., see Fulle et al., 1996; Golden et al., 1998; and references therein):

(E)

From Eq. (C), the rate constant at 219 K and 39 Torr pressure is 1.5 • 10 -12

k =

(14)

0.6{1 +(log[(7.4• 10-12)/(1.5• 10-12)])2} -1

OH+CO

".

ka

k_ a

M

" (HOCO)* ~ kb kc

HOCO,

> OH+CO, > H + CO 2.

(15)

(1 + 1.5 • 1 0 - 1 2 ) / ( 1 . 5 • 10 -12) •

0,611 q- (log[(1.4 ?K10-12)/(1.5 ?K10-12)])2] -1

= [7,5 •

1 0 - - 1 3 1 0 , 6 0.999

= 4.5 • 10-13 cm 3 molecule-1 s-1. This is almost a factor of two smaller than at 300 K and 760 Torr pressure. In summary, rate constants for addition reactions in the atmosphere can be estimated as a function of t e m p e r a t u r e and pressure if values are available for the low- and high-pressure limiting rate constants as a

In (15), H O C O is the radical adduct of O H + CO, and H O C O * is the adduct containing excess internal energy resulting from the energy released by bond formation between O H and CO. As described earlier, M is any molecule or atom that collides with the H O C O * , removing some of its excess energy; in practice, it is usually an inert bath gas such as He or Ar that is present in great excess over the reactants. Reactions such as (15), which proceed with the formation of a bound adduct between the reactants, are known as indirect or nonconcerted reactions. The

13 8


"OCOa

HOCO

uJ H + CO 2

HOCO Reaction coordinate F I G U R E 5.3 Typical potential energy diagram proposed for reaction of O H with CO (adapted from Mozurkewich et al., 1984). Note the well corresponding to formation of the (HOCO) intermediate.

adduct is "stable" in the sense that it corresponds to a well on the potential energy surface connecting the reactants and products (Fig. 5.3); as such it has a finite lifetime and should be capable of being detected using appropriate techniques. Because of the complex nature of the mechanism, such reactions can exhibit a relatively complex temperature dependence. In addition, if the rate of collisional stabilization of the excited adduct is comparable to its rate of decomposition, a pressure dependence may result, as in the O H + CO reaction. The distinction between bimolecular and termolecular reactions blurs in such cases. In any event, the OH + CO rate constant at 1 atm in air is now thought to be 2.4 • 10-13 cm 3 molecule- 1 s- 1, significantly greater than the low-pressure value, 1.5 • 10 -13 cm 3 molecule- 1 s-1, that had been widely accepted at one time.

3. Temperature Dependence of Rate Constants a. A r r h e n i u s Expression

The temperature dependence of many rate constants can be fit over a relatively narrow temperature range by the exponential A r r h e n i u s equation k = A e -Ea/RT

(F)

where R is the gas constant and the temperature T is in kelvin (K = ~ + 273.15). A, the preexponential factor, and E~, the activation energy, are parameters characteristic of the particular reaction. To a first approximation over the relatively small temperature range encountered in the troposphere, A is found to be independent of temperature for many reactions, so that a plot of In k versus T -1 gives a straight line of slope - E ~ / R and intercept equal to In A. However, the Arrhenius expression for the temperature dependence of the rate constant is empirically based. As the temperature range over which experiments could be carried out was extended, nonlinear Arrhenius plots of I n k against T -1 were observed for

some reactions. This is not unexpected when the predictions of the two major kinetic theories in common use today, collision theory and transition state theory, are considered. A brief summary of the essential elements of these is found in the following sections, as we refer to them periodically throughout the text. For many reactions, the temperature dependence of A is small (e.g., varies with T 1/2) compared to the exponential term so that Eq. (F) is a good approximation, at least over a limited temperature range. For some reactions encountered in tropospheric chemistry, however, this is not the case. For example, for reactions in which the activation energy is small or zero, the temperature dependence of A can become significant. As a result, the Arrhenius expression (F) is not appropriate to describe the temperature dependence, and the form (G)

k = BTne -G/R~

is frequently used, where B is a temperature-independent constant characteristic of the reaction and n is a number adjusted to provide a best fit to the data. While most reactions with which we deal in atmospheric chemistry increase in rate as the temperature increases, there are several notable exceptions. The first is the case of termolecular reactions, which generally slow down as the temperature increases. This can be rationalized qualitatively on the basis that the lifetime of the excited bimolecular complex formed by two of the reactants with respect to decomposition back to reactants decreases as the temperature increases, so that the probability of the excited complex being stabilized by a collision with a third body falls with increasing temperature. An alternate explanation can be seen by treating termolecular reactions as the sum of bimolecular reactions, as was illustrated in Section A.2 for the OH + SO 2 + M reaction. Recall that the third-order, lowpressure rate constant k Ill c a n be expressed as the product of the three rate constants ka, kb, and k c for the three individual reaction steps (12), ( - 1 2 ) , and

(13): iii _ k lO

kakr kb

Expressing each of the component rate constants in the Arrhenius form, k 10 III becomes III=Aloe-El~ lo

-(Ea+Ec-Eb)/RT

Ab

e

.

(H)

Thus the activation energy for the reaction, E~0, is a combination of the activation energies for the individual steps, (Ea + E c - Eb). If E b > ( E a + E c), that is, if

139


the activation energy for decomposition of the energized adduct (HOSO2)* back to reactants is greater than the sum of E a + E c, then the effective activation energy El0 for the termolecular reaction becomes negative and the rate constant decreases as the temperature increases. We now briefly consider the predictions of collision theory and transition state theory for the expected form of the rate constant and its temperature dependence.

end of the OH, the orientation is correct for an overall reaction to 02 + H" O + OmH

However, if it collides with the hydrogen atom end, no net chemical change will result, although, in fact, an exchange reaction (i.e., exchange of the oxygen atoms) may occur: O+H

m O - - + [ O ..... H ..... O]--+

b. Predictions of Collision Theory and Transition State Theory

OH + O

Collision theory is based on the concept that molecules behave like hard spheres; during a collision of two species, a reaction may occur. To estimate a rate constant for a bimolecular reaction between reactants A and B based on this theory, one needs first to calculate the number of collisions occurring in a unit volume per second (ZAB) when the two species, A and B, having radii rA and rB, are present in concentrations NA and NB, respectively. From gas kinetic theory, this can be shown to be given by Eq. (I): )2( 8 ~ ' k T ) 1/2

= ZAB

(r A + r B

/x

--+ [O ..... O ..... H] ---> 0 2 + H.

NAN

B.

(I)

/x is the reduced mass of A and B [/x = m A m B / ( m A + roB)], where m is the mass of the individual meucles, k is the Boltzmann constant (1.38 • 10 -23 J K-~), and T is the temperature in kelvin. In deriving Eq. (I), it is assumed that A and B are hard spheres and only collide when their centers come within a distance (r A + r B) of each other. In considering reactions between colliding molecules, one must take into account two additional factors. First, different collisions will occur with different amounts of energy depending on the speed of the molecules as they collide. Most reactions are expected to have an energy barrier that must be surmounted for a reaction to occur. This energy barrier arises from the net effect of simultaneous bond breaking and formation; at the transition state in the reaction, the energy released from bond formation is generally less than that required for bond breaking. The difference, which is the energy barrier, must be supplied in other ways if reaction is to take place. Second, even if the reactants collide with enough energy to surmount the energy barrier, they may not react if they are not in the proper orientation with respect to each other. The importance of this so-called steric factor can be illustrated using the reaction of ground-state oxygen atoms with the hydroxyl free radical. If the oxygen atom collides with the oxygen atom

(no net chemical change).

To take into account the energy requirement, one can modify the result in Eq. (I) by calculating only the number of collisions between A and B that have a certain minimum energy, E 0. In the simplest approach, it is assumed that no reaction occurs if the energy of the colliding pair is less than E0, and reaction occurs 100% of the time for energies > E 0. Alternatively, it can be assumed that for energies > E 0 the probability of reaction increases as the relative collision energy increases. To take into account this dependence on energy, the concept of A and B being hard spheres with radii rA and r B can be modified. Let CrAB= "n(rA + rB )2 be the collisional cross section for hard-sphere collisions between A and B and crR the reaction cross section for reaction between A and B. At energies < E0, crR = 0; that is, no reaction occurs. For energies > E0, CrR could be taken as a constant (e.g., equal to CRAB),which would correspond to assuming all collisions having energies above the threshold energy leading to reaction. The alternate approach of an increase in reaction probability with increasing energy above the threshold corresponds to assuming an expression for crR that is a function of total energy E, for example,

( E0)

crR = CrA~ 1 -- ~ = 0

for E _> Eo,

(J)

for E < E o.

For the form of crR shown in Eq. (J), integrating over all total energies from 0 to infinity, the rate of reactive collisions becomes 8kT ) 1/2 Z R = CrAB

e-E~ /krNa N~ .

(K)

The rate constant k can thus be identified with Eq. (L): 8kT k = craB ~r/x

1/2 e -E~

(L)

140


Equation (L), however, does not take into account the need for proper orientation of the colliding molecules in order for a reaction to occur. This is commonly done by including an extra factor, P (the steric factor), which is the probability that the colliding molecules will have the correct orientation. This leads to Eq. (M): k = P~rAB ~

e-E~

(M)

By comparison of (M) and (F), it can be seen that the preexponential factor A in the Arrhenius equation can be identified with P~raB(8kT/rrlx) 1/2 and the activation energy, Ea, with the threshold energy E 0. It is important to note that collision theory predicts that the preexponential factor should indeed be dependent on temperature (T1/2). The reason so many reactions appear to follow the Arrhenius equation with A being temperature independent is that the temperature dependence contained in the exponential term normally swamps the smaller T 1/2 dependence. However, for reactions where Ea approaches zero, the temperature dependence of the preexponential factor can be significant. Collision theory is used mainly as a reference for the efficiency of reactions. Thus, at a temperature of 25~ (298 K), the rate constant for a reaction between two molecules each having a radius of 0.2 nm and a molecular weight of 50 g mol-1 would, according to Eq. (M), be 2 . 5 • 10 -1~ cm 3 molecule -1 s -1 for P = 1 and E 0 = 0. That is, when there are no steric or energy barriers to the reaction, the rate constant should be of the order of 10 -l~ cm 3 molecule -1 s -1. A reaction with a rate constant of 10 -15 cm 3 molecule -1 s -1 is therefore one that goes in approximately every 105 collisions. Transition state theory is more commonly applied today than collision theory. It is especially useful in examining reactions in solution and avoids the problem of introducing arbitrary factors such as the steric factor to take into account steric requirements. Transition state theory treats a reacting system thermodynamically. Let us again take a bimolecular reaction between A and B. Transition state theory assumes that as A and B collide and start to react, they form a species called the activated complex, which corresponds to the A - B adduct at the peak of the energy hill lying between reactants and products. This activated complex is thus in a "transition state" and can either fall back to reactants or go on to form products. The activated complex is normally indicated with a double dagger symbol, AB 4. The reaction can thus be given

Assuming the activated complex is in equilibrium with reactants, one can define the equilibrium constant K 4 :

kf

K 4-

-

kr

[AB 4 ] [a][g]

(N) "

K 4 is related to the standard free energy change in going from reactants to the transition state, AG ~ by the usual thermodynamic relationship: K ~ = e -ac~

= (e-aH~176

(0)

A H ~ and AS ~ are the standard enthalpy and entropy changes in going from reactants to the transition state (not to products). The reaction rate is determined by the rate at which AB ~ forms products Rate of product formation = v[AB ~ ]. where v is the frequency with which AB ~ breaks up into products. Substituting for [AB ~] from Eq. (N) and for K ~ from (O), one obtains Rate - v K ~ [A][B], (p)

= u(eaS~176

The enthalpy change A H ~ is related to the energy change in going from reactants to the transition state, that is, to the activation energy. The frequency p of breakup of the activated complex into products is often approximated by p = k T / h , where k and h are the Boltzmann and Planck constants, respectively. Comparison of Eq. (P) to the rate equation shows that kT

(Q)

k = --=--(eAS~176

h Again the preexponential factor is seen to be temperature dependent, but for large activation energies, the exponential term dominates the temperature dependence of the rate constant. The preexponential factor involves the entropy change in going from reactants to the transition state; the more highly ordered and tightly bound is the transition state, the more negative AS ~ will be and the lower the preexponential factor will be. Transition state theory thus automatically takes into account the effect of steric factors on rate constants, in contrast to collision theory. An alternate form of the rate constant predicted by transition state theory using a statistical mechanical approach for the equilibrium constant K ~ is Eq. (R):

as kf

A +B

kp

>AB ~

< kr

> Products.

k=

--7[-

qaqB

e-

.

(R)

B. LABORATORYTECHNIQUES FOR DETERMINING ABSOLUTE RATE CONSTANTS FOR GAS-PHASE REACTIONS

Here the q's are partition functions for the reactants (qA and qB) and the transition state (qs and E * is the energy difference between the reactants and the transition state. The partition functions include contributions from translational, rotational, vibrational, and electronic degrees of freedom, qtrans qrot qvib qelect 9

Those for the reactants can be evaluated using conventional techniques discussed in physical chemistry texts; estimating the partition functions for the transition state requires making assumptions concerning the nature of the transition state. Transition state theory can be used to test reaction dynamics on a molecular scale. Thus one can hypothesize a spatial configuration of the atoms in the transition state and from this calculate AS ~ ; the predicted rate constant can then be compared to that observed. If the agreement is not acceptable, the molecular configuration of the transition state can be adjusted until such agreement is obtained. Assuming this molecular configuration approximates the actual form of the intermediate in the reaction, one can learn something about the chemical dynamics of the reaction. An example of the application of transition state theory to atmospheric reactions is the reaction of OH with CO. As discussed earlier, this reaction is now believed to proceed by the formation of a radical adduct HOCO, which can decompose back to reactants or go on to form the products H + CO 2. For complex reactions such as this, transition state theory can be applied to the individual reaction steps, that is, to the steps shown in reaction (15). Figure 5.3 shows schematically the potential energy surface proposed for this reaction (Mozurkewich et al., 1984). The adduct HOCO, corresponding to a well on the potential energy surface, can either decompose back to reactants via the transition state shown as HOCO~* or form products via transition state H O C O ~ .

c. Example of Importance of Temperature Dependence: P A N Decomposition As we have seen in Chapter 1, peroxyacetyl nitrate (PAN) is a powerful lachrymator and severe plant phytotoxicant formed in irradiated V O C - N O x mixtures from the reaction of peroxyacetyl radicals with NO2:

O

O

II CH3COO

II + NO 2

> CH3COONO

2.

(16)

PAN is thermally unstable, however, and decomposes at higher temperatures to reform peroxyacetyl radicals

141

and NO2, that is, the reverse of reaction (16): O

O

II

II

CH3COONO

2

> CH3COO

+ N O 2.

(-16)

The rate constant for PAN decomposition (k_~6) is strongly temperature dependent. As a result, this is important in the atmosphere because PAN acts as a reservoir for NO 2. At low temperatures, much of the total NO X may be tied up in PAN; as the atmosphere warms, due to either diurnal temperature variations or the air mass being transported into warmer regions, the rate of PAN decomposition increases. This releases NO2, which can then form other secondary pollutants such as 0 3 and H N O 3. The rate constant k_~6 can be fit to the Arrhenius form (see, for example, Grosjean et al., 1994) k_16 (s - 1 ) = 1.58 • 1016e -(112"5 kJ/mol)/RT

over the temperature range normally encountered in the troposphere. At room temperature (25~ = 77~ k_16 is 3.0 • 10 -4 S-1, giving a natural lifetime with respect to decomposition of 55 min. However, at 0~ (32~ k_16 is 4.8 • 10 -6 s -1, corresponding to a lifetime of ~ 57 h. At 35~ (95~ the rate constant k_~6 (1.3 • 10 -3 s -~) is larger than that at 0~ by almost three orders of magnitude and the lifetime for PAN is correspondingly shorter, only ~ 12 min! Clearly, the temperature dependence of rate constants can be extremely important in determining the lifetimes and fates of certain species in the atmosphere, as well as their contribution to secondary pollutant formation.

B. LABORATORY TECHNIQUES FOR DETERMINING ABSOLUTE RATE C O N S T A N T S FOR GAS.PHASE REACTIONS In this section we discuss the major experimental methods used to determine absolute rate constants for gas-phase reactions relevant to atmospheric chemistry. These include fast-flow systems (FFS), flash photolysis (FP), static reaction systems, and pulse radiolysis. The determination of relative rate constants is discussed in Section C. In general, we use simple bimolecular reactions of the type A + B -+ Products

(17)

as illustrations. However, the techniques can be modified to study termolecular reactions, as discussed earlier, as well as unimolecular reactions.

142

5. KINETICS A N D A T M O S P H E R I C CHEMISTRY

To study the reaction kinetics of a relatively reactive species A with a reactant B, one normally follows the loss of small amounts of A in the presence of a great excess of B. This requires then that one be able first to generate A and second to monitor its concentration as a function of time. Ideally, to fully elucidate the reaction mechanism, one would also monitor the concentrations of intermediates and products. As we shall see, in practice, for many reactions this proves to be much more difficult than to simply determine the rate constant itself. 1. K i n e t i c A n a l y s i s The rate law for a simple bimolecular reaction such as (17) is given by -d[A]

dt

= k17[A][B].

(S)

If a small concentration of A is generated in a great excess of B, then even if (17) is allowed to go to completion, the concentration of B will remain essentially constant at its initial concentration [B] 0. Integrating (S) and treating [B] 0 as constant, one obtains [A] In [A]o

- k17[B]ot.

(T)

That is, A decays exponentially with time determined by (k17[B]o), as if it were a first-order reaction. Thus under these so-called pseudo-first-order conditions, a plot of ln[A] against time for a given value of [B] 0 should be linear with a slope equal to (-k17[B]o). These plots are carried out for a series of concentrations of [B] 0 and the values of the corresponding decays determined. Finally, the absolute rate constant of interest, k17 , is the slope of a plot of the absolute values of these decay rates against the corresponding values of [B] o. Some examples are discussed below. As we have seen earlier, even third-order reactions can be reduced to pseudo-first-order reactions by keeping the concentrations of all species except A constant and in great excess compared to A. This technique of using pseudo-first-order conditions is by far the most common technique for determining rate constants. Not only does it require monitoring only one species, A, as a function of time, but even absolute concentrations of A need not be measured. Because the ratio [A]/[A] 0 appears in Eq. (T), the measurement of any parameter that is proportional to the concentration of A will suffice in determining k17 , since the proportionality constant between the parameter and [A] cancels out in Eq. (T). For example, if A absorbs light in a convenient

spectral region and Beer's law is obeyed, then the absorbance (Abs) of a given concentration of A, N (number cm-3), is given by In -I~- = Abs = crN1, I

(U)

where I 0 and I are the intensities of the incident and transmitted light, respectively, 1 is the optical path length, and cr is the absorption cross section of A (to the base e) Substituting into Eq. ( T ) f o r [A] = N = Abs/crl, one obtains (Abs/crl) In

(Abs/crl)0

(Abs) = In

(Abs)0

= -kly[B]0t,

(V)

where (Abs) and (Abs) 0 are the absorbance of the light by A at times t and t = 0, respectively. For example, O 3 has a strong absorption at 254 nm, which can be used to monitor its concentration. This ability to monitor a parameter that is proportional to concentration, rather than the absolute concentration itself, affords a substantial experimental advantage in most kinetic studies, since determining absolute concentrations of atoms and free radicals is often difficult. This pseudo-first-order kinetic analysis is generally applied regardless of the experimental system used. 2. F a s t - F l o w S y s t e m s Fast-flow systems (FFS) consist of a flow tube typically 2- to 5-cm in diameter in which the reactants A and B are mixed in the presence of a large amount of an inert "bath gas" such as He or Ar. As the mixture travels down the flow tube at relatively high linear flow speeds (typically 1000 cm s-l), A and B react. The decay of A along the length of the flow tube, that is, with time, is followed and Eq. (T) applied to obtain the rate constant of interest. The term fast flow comes from the high flow speeds. In most of these systems, discharges are used to generate A or another species that is a precursor to A; hence the term fast-flow discharge system (FFDS) is also commonly applied. Since fast-flow discharge systems have been applied in many kinetic and mechanistic studies relevant to tropospheric chemistry (e.g., see Howard, 1979; Kaufman, 1984), we concentrate on them. However, all fast-flow systems rely on the same experimental and theoretical principles. Figure 5.4 is a schematic diagram of a typical fastflow discharge system. The reactive species is generated in a microwave discharge and enters at the upstream

B. LABORATORY TECHNIQUES FOR DETERMINING ABSOLUTE RATE CONSTANTS FOR GAS-PHASE REACTIONS

143

Precursor to Reactant A

Resonance Lamp

I "vv~I

t

I

Pressure

D /

I

]MWDI

I Carrier

I

Gauge

2.5cm Reactant

/

-" gas ~ ~

+

In f Coolant .,(_/ In ]

~"-~G,,..

~

~

oolant ~Fte~c2ant

out/

T

TOP VIEW

Movable Inlet

To Pump

FIGURE 5.4 Schematic diagram of a fast-flow discharge system (adapted from Beichert et al., 1995). MWD = microwave discharge used to generate the reaction species from a precursor and to generate light of the appropriate wavelength to measure it using the resonance lamp.

end of the flow tube, and the second reactant is added through one port, C or D, of a movable inlet. The detector is fixed at the downstream end of the flow tube and the reaction time is varied by moving the mixing point for C and D (i.e., the movable inlet) relative to the fixed detector. For example, OH can be generated by the reaction of H with NO2: H + NO 2 ~ OH + NO.

d -.

u

v=

A.t'

7-

"

(x)

(18)

Atomic hydrogen is formed by discharging dilute H 2 / H e mixtures. Then these H atoms are converted to OH by the addition of excess NO2, e.g., through port C. By adjusting the concentration of NO 2 appropriately, essentially all of the H can be converted to OH before the second reactant is added through port D. Table 5.3 shows some typical sources of reactive species of atmospheric interest used in FFDS, while Table 5.4 shows some of the methods used to detect them. Under conditions where the plug flow assumption is valid, that is, concentration gradients are negligible so that the linear flow velocity of the carrier gas is the same as that of the reactants, the time (t) for A and B to travel a distance d along the flow tube is given by t =

Here u is the linear flow speed, which can be calculated from the cross-sectional area of the flow tube (AR), the total pressure (P) in the flow tube, the temperature (T), and the molar flow rates (tin /dt) of the reactants and the diluent gas:

T A B L E 5.3

Reactive species OH

Source H2

MWD> 2H

H + NO 2 ~ OH + NO F2 MWD> 2F F+H20 C1

NO 3

~ OH+HF

F2 MWD>2F F+HCI~

CI+HF

F2 MWD> 2F F+HNO 3 ~NO 3 +HF

HO 2

F2 MWD> 2F F + H202 --->HO 2 + HF

O(3p ) (w)

Some Typical Sources of Reactive Species in Fast-Flow Discharge Systems

N2 MWD> 2N N + NO ~ O(3P) + N 2

144


T A B L E 5.4 Some Typical Detection Systems U s e d for Reactive Species in Fast-Flow Systems

Technique

Typical application

Resonance fluorescence or light absorption

OH, 308 nm C1, 135 nm O(3p), 130.5 nm H(2S), 121.6 nm CH30, 303.9 nm NO3,662.9 nm

Chemiluminescence

O(3p):

O +NO

H(2S):

H+NO

Mass spectrometry

OH R HO 2 Stable products

Electron paramagnetic resonance (EPR) Laser magnetic resonance (LMR)

OH OH, H O

M ~ NO~ + M M ~ HNO* + M

2

walls of the flow tube. Unfortunately, OH radicals are particularly sensitive to removal by wall reactions. While the mechanism and products of these wall reactions are unknown, it has been established that the rate of loss at the walls can be minimized by using various flow tube wall coatings or treatments. These include substances such as teflon or halocarbon waxes, which simply cover the entire surface so the incoming reactive species are only exposed to relatively unreactive carbon-halogen bonds, or treatment with boric or phosphoric acids. While such treatments have been shown to lower the rates of removal at the walls, why they do so is not clear. Fortunately, the kinetics of the wall loss, measured from the decay of the reactive species in the absence of added reactant, are generally observed to be first order, so that corrections for these processes can be readily incorporated into the kinetic analyses. When these wall losses are significant, the integrated form of the rate expression (T) for reaction (17) of A + B becomes

[A] At typical linear flow speeds of 1000 cm s -1, 1 cm along the tube corresponds to ~ 1-ms reaction time. Thus a flow tube of length 1 m can be used to study reactions at reaction times up to 100 ms. Total pressures in most FFDS have typically been in the range 0.5-10 Torr where rapid diffusion across the flow tube ensures a relatively flat concentration profile of the reactants, so that Eq. (W) is valid. Maintaining the discharge used to generate atoms and free radicals is also difficult above a few Torr total pressure. The lower end of the pressure range is determined by the need to maintain viscous flow and to avoid significant axial concentration gradients. The latter may arise because of the lower concentrations of reactants at the downstream end of the flow tube; these can cause the true flow speed of the reactants to be greater than the calculated linear flow speed, due to their axial diffusion. Techniques for estimating errors due to such factors and for correcting measured rate constants for them are discussed in detail by Mulcahy (1973), Brown (1978), and Lambert et al. (1985). However, flow tube systems for use at much higher pressures, up to several hundred Torr, have also been designed and applied to reactions of atmospheric interest (e.g., see Keyser, 1984; Abbatt et al., 1990, 1992; Seeley et al., 1993; and Donahue et al., 1996a). At these higher pressures, the velocity and radical axial and radial concentration profiles are experimentally determined and the full continuity equation describing the concentration profiles is solved. A major factor in many FFDS studies is diffusion of the reactive species accompanied by their loss at the

In [A]0 - -(k17[B]~ + kw)t'

(Y)

where k w is walls of the constant k17 plots of the

the observed first-order loss of A at the flow tube in the absence of B. The rate can then be extracted from the slopes of pseudo-first-order rates of decay, R = ( k l 7 [ B ] 0 + k w ) , against [B] 0. An example is shown in Figs. 5.5 and 5.6 for the reaction of OH with nitrosyl chloride, C1NO. Figure 5.5 shows the decay of OH resonance fluorescence emission intensity (proportional to the OH concentration) as a function of reaction time in a fast-flow discharge system at ~ 1 Torr total pressure as the concentration of C1NO is increased from 0 to 14.1 • 1013 molecules cm -3. As expected from Eq. (Y), the absolute value of the slope of the lines increases as [C1NO] 0 increases. Figure 5.6 shows the plot of the absolute values of these slopes against [C1NO] 0. The slope of this plot gives the rate constant for the reaction of OH with C1NO, OH + C1NO ~ Products,

(19)

under these conditions, which in this case gives k19 = 5.6 • 10 -13 cm 3 molecule -1 s -1. The nonzero decay of OH when the C1NO concentration is zero is due to loss of OH at the walls of the flow tube. These wall reactions can be a problem in FFDS studies. To avoid unrecognized interferences in the data associated with these heterogeneous reactions, as well as other secondary reactions, it is generally recommended that flow tube studies of a particular reaction be carried out using as many different wall coatings as possible. In addition, the use of different carrier gases


have been used with modulation of the reactants to obtain mechanistic information. In this approach, the radical source is modulated, and changes in the spectra with the source on and off are used to identify and quantify products (Donahue et al., 1996b).

50 .m

tag t-

C .m 0 C

2.29

o

O m: 3

3.87

10

5.78

o E E O

8.40

-r O 0) > .m 4-., t~

14.1

rr

1I

I

I

I

10

I

I

20

I

30

Reaction time (ms) F I G U R E 5.5 Typical plot of O H resonance fluorescence intensity as a function of reaction time in the presence of increasing concentrations of C1NO (in units of 1013 molecules c m - 3) at 373 K (adapted from Finlayson-Pitts et al., 1986).

and flow tubes of different diameters is recommended. Flow tubes have also been used in combination with such techniques as mass spectrometry and FTIR for product studies. For example, high-pressure flow tubes with a White cell and FTIR at the downstream end

100 't,~

90

"6

80 7O 60

r

5o oo 0~

145

4o

30

Y_ 0

1

2

3

4

5

6

7

8

9

[CINO] (1013 molecules cm -3) F I G U R E 5.6 Typical plot of observed first-order rate constants for the decay of O H as a function of the initial C1NO concentration at 373 K (adapted from Finlayson-Pitts et al., 1986).

3. Flash Photolysis Systems As the name implies, this technique relies on flash photolysis to generate the reactive species A. In one of the most common configurations, resonance or induced fluorescence is used to monitor the decay of A - - h e n c e the name flash photolysis-resonance fluorescence (FPRF). Since lasers are now frequently used as the photolysis source, the term laser ]lash photolysis-resonance fluorescence (LFP-RF) is also used. Figure 5.7 is a schematic diagram of a typical FP-RF apparatus used to study chlorine atom reactions (Nicovich and Wine, 1996). For example, the fourth harmonic at 266 nm from an Nd:YAG laser can be used to generate chlorine atoms from the dissociation of phosgene, COC12. After a preset time following the photolytic flash, the time decay of the reactive species is monitored using the fluorescence excited by a resonance lamp. Since B is present in concentrations in great excess compared to A, care must be taken to avoid impurities that may react with A or photolyze to produce reactive species that do. A restriction on the nature of B is that it must not photolyze significantly itself; reactions of such species as NO 2 and 03, which dissociate to produce highly reactive oxygen atoms, are often difficult to study with this technique. In addition, care must be taken to avoid the buildup of reaction products or of photolysis products in the photolysis cell, since some of these can photolyze and produce interfering secondary reactions. This is usually accomplished by using a slow flow of gas through the cell. The limitations on the total pressure in the FP-RF cell are far less severe than those for FFDS. The lower end of the pressure range that can be used is determined by the need to minimize diffusion of the reactants out of the viewing zone. The upper end is determined primarily by the need to minimize both the absorption of the flash lamp radiation by the carrier gas and the quenching of the excited species being monitored by RF. In practice, pressures of ~ 5 Torr up to several atmospheres are used. The kinetic analysis is again typically pseudo-first-order with the "stable" reactant molecule B in great excess over the reactive species as outlined earlier. Table 5.5 gives some typical sources of reactive species used in FP-RF systems. An example of the use of FP-RF to study the kinetics of an atmospherically relevant reaction is found in Fig. 5.8 (Stickel et al., 1992). Chlorine atoms were

146


insertable

to pump

manometer

gwave power supply

PDL ~

I

;;=X2

t

gwave cavity

_

~

atomic line sourc~j~

. I

to pump

energy

I 1

filtergas

Pulsed Photolysis "~ Laser

'

'

MgF2t

~J~'/ j'~' "'~, '" ~...,

lenses

PG

thermostated 1 fluid

,.", :,, e"-/,"

i

to

/N.//..x\N~/./AGasU pump ,/~\\,NN ~ MCS )

t

.

DDG

~ PM

r _

filtergas sig

FIGURE 5.7 Laser photolysis resonance fluorescence apparatus for studying the kinetics of gas-phase reactions of H, O, C1, and Br atoms with atmospheric trace gases. A/D, amplifier/discriminator; DDG, digital delay generator; FM, flow meter; IF, interference filter; MCS, multichannel scaler; PD, photodiode array detector; PG, pulse generator; PMT, photomultiplier. (Graciously provided by J. M. Nicovich and P. H. Wine, Georgia Institute of Technology.).

formed by laser photolysis of COG12 at 266 nm and detected using resonance fluorescence in the 135- to 140-nm region. As expected, the decay of C1 in the presence of a great excess of CHBSCH 3 (DMS) is exponential (Fig. 5.8a), and slopes of such decays are linear with the concentration of DMS (Fig. 5.8b). From the slope of the line in Fig. 5.8b, the rate constant at this temperature and pressure was determined to be (k = 2.71 _+ 0.09) • 10 -1~ cm 3 molecule -1 s -1.

4. Pulse Radiolysis Pulse radiolysis has been used in a number of kinetic studies, for example by the group at RISO National Laboratory (Denmark) using the Febetron field emission accelerator facilities (e.g., see Nielsen and Sehested, 1993; Pagsberg et al., 1995; and Wallington et al., 1998). A short (30 ns) pulse of high-energy ( ~ 2 MeV) electrons impacts a reaction cell containing an


TABLE 5.5

Some Typical Sources of Reactive Species in F P - RF Systems

Reactive species

Source

OH

H20

Reactions of 0(1D), e.g., 03, N20 e.g., 03 + hv ~ O(1D) + 0 2 O(1D) + H 2 ~ OH + H or O(1D) + H20 --+ 2OH HNO 3 H202 C1

COC12

O(3p)

0 2

H RO

Alkanes, e.g., C3H s RONO

8.5

ii

rr

6.5 o o

4.5

0

50

100

150

Time (gs)

m

5-

o

ated by the electron b o m b a r d m e n t of H 2, H, and O H using H 2 0 and O from CO 2 radiolysis. The reaction cell has a White cell optical system (see Chapter l l . A . l c ) w i t h a pulsed xenon lamp light source. Once the radicals are formed, they are detected by their absorptions in the U V using the Xe lamp and a m o n o c h r o m a t o r - p h o t o m u l t i p l i e r or photodiode array detector. Thus the absorption spectra of the free radicals generated in the system can be measured and the absorption at a particular wavelength used to follow their reaction kinetics.

5. Cavity Ring Down Method

atom or free radical source and reactant of interest. For example, if SF6 is used as the bath gas, F atoms are generated, which can then undergo secondary reactions to form other reactant free radicals. For example, alkyl radicals can be generated from the F + R H reaction, and if carried out in the presence of 0 2, R O 2 is formed (e.g., see Nielsen and Sehested, 1993; and Wallington et al., 1998). Hydrogen atoms can be gener-

" _=

2

r

"t,

The cavity ring down method was first described by O'Keefe and Deacon (1988) and has been reviewed by Paul and Saykally (1997) and Scherer et al. (1997). This technique is based on the sequential loss of light intensity as a light pulse repeatedly traverses the length of a cell during multiple reflections between two mirrors. Loss of intensity occurs both during reflection at the mirrors and, if an absorbing gas is present, by its absorption as well. As will be seen shortly, the change in the time profile of the light transmitted through mirror B when an absorber is present can be used to follow the concentration of the absorber and hence to carry out kinetic studies. Figure 5.9 is a schematic diagram of a typical cavity ring down apparatus. A laser pulse enters a reaction cell that has two highly reflecting mirrors. The distance between the mirrors, L, must be large compared to the pulse width to avoid multibeam interference in the cell. After traveling through the length of the cell the first time, the laser pulse strikes mirror B. If the reflectivity of the mirrors is R (defined as the fraction of light reflected), then (1 - R ) is the fraction of light lost by reflection at this surface. If the incident light intensity is I', then the intensity of light lost by the reflection is (1 - R ) I ' . Assuming that both mirrors A and B have the same reflectivity, then in one round trip in the cell, from mirror B to mirror A and back, the lost intensity is d I ~ 2 ( 1 - R ) I ' . The time to make this one round trip is t r (where t r = 2 L / c , and c is the speed of light). U n d e r these conditions,

O

0

I

7

I

dt

14

[CH3SCH3] (1013molecule s cm-3)

FIGURE 5.8 (a) Typical decay of resonance fluorescence from atomic chlorine in the presence of CH3SCH3 (8.6 • 1013 molecules cm-3) at 297 K and in 50 Torr N2 as the carrier gas (adapted from Stickel et al., 1992). (b) Typical pseudo-first-order plot of slopes of plots such as those in part (a) against the initial concentration of CH3SCH 3 (adapted from Stickel et al., 1992).

2(1 -

-

dI 0

147

R)I'

-(1 - R ) I '

tr

tr/2

or

(1 - R )

I

m = exp

-

-

I0

[

~

tr/2

-t = exp Lfl[

c(1 - R ) ] (z)

= exp 7"0

148


Laser pulse

I--7

Tunable laser

Detector

Reaction cell Sample

Av

~--B

~E Time FIGURE

5.9

Schematic diagram of a cavity ring down apparatus (adapted from Paul and

Saykally, 1997).

where r 0 is a constant known as the cavity ring down time in the absence of an absorbing gas. Figure 5.9 also illustrates the resulting measured signal intensity as a function of time, from which the cavity ring down time is obtained. In the presence of an absorbing gas with absorption cross section cr and concentration N, the fractional change in light intensity is given by A = ( 1 - e -crNL) (see Problem 10). This change in intensity occurs in addition to the (1 - R) loss on the mirror and Eq. (Z) becomes

i

I0 = e x p

{

-t

t~2

R, AI} + t~2

=exp

it l ~r

'

(AA) where r is the cavity ring down time in the presence of the gas. It can be shown (see Problem 10) that the absorption (A) of the gas is therefore related to r and %, by tr[1 A=

--

2

r

1 r0

.

(BB)

Thus measurement of the cavity ring down times in the presence and absence of the absorbing gas allows the absorption of the gas to be obtained, and from A = (1 - e - ~ N L ) ~ crNL, the concentration of the gas determined. These concentrations as a function of time therefore allow the reaction kinetics of the absorbing species to be determined. The advantages of the cavity ring down technique are its high sensitivity (fractional absorptions per pass of ~ 10 -6) and good time resolution ( ~ 10 .5 s). In addition, it has the usual advantages associated with absorption spectrometry--e.g., most species can be detected through light absorption at specific wavelengths - - a n d is experimentally relatively simple to carry out.

Typical examples of its application to processes of interest for atmospheric chemistry are the measurement of the kinetics of the reaction of the vinoxy radical with 0 2 (Zhu and Johnston, 1995) and the kinetics of the C2H 5 + C2H 5 and C 2 H 5 0 2 + C 2 H 5 0 2 reactions (Atkinson and Hudgens, 1997). In addition, it has been shown to be useful for probing surface processes as well when combined with total internal reflection techniques (Pipino et al., 1997).

6. Static Techniques Some reactive species of atmospheric interest such as 0 3 are not so highly reactive that they must be generated and studied on subsecond time scales, as is true for atoms and free radicals. In these cases, where the reactions are generally much slower than those of atoms and free radicals such as OH, the experimental techniques used for determining absolute rate constants are often static in nature and do not require in situ generation of the reactive species. Taking 0 3 as an example, mixtures of 0 3 in 0 2 can be generated using commercial ozonizers, which are either electrical discharges or UV lights. This mixture is injected into a cell or chamber with the reactant of interest, and its decay followed. Monitoring 0 3 accurately and inexpensively in situ is reasonably straightforward because it absorbs strongly at 253.7 nm, a stable, inexpensive, and readily available mercury resonance line. Alternatively, samples of the reaction mixture can be withdrawn and the 0 3 analyzed with a commercial 0 3 monitoring instrument. Ozone decomposes on surfaces, the rate depending on the nature of the particular surface and whether it has been previously "conditioned" by exposure to 0 3. While this heterogeneous decomposition is much slower than the wall loss of OH, the homogeneous gas-phase

C. LABORATORY TECHNIQUES FOR DETERMINING RELATIVE RATE CONSTANTS FOR GAS-PHASE REACTIONS

"7 =-

g

0.03

.

o F

oID

0

//

0.04

~176 0.00 r I 0 20

/

independently, the second one can then be calculated from the experimentally determined ratio. In the simplest case, determining relative rate constants for a reactive species A is based on a competition between two reactions"

I 40

I 60

kl

A + X1 I I I > 80 100 120

[o-Cresol] (ppm) F I G U R E 5 . 1 0 Plots of ozone pseudo-first-order decay rate constant as a function of the o-cresol concentration using U.S. E P A protocol for determining O 3 rate constants (adapted from Pitts et al.,

1981).

reactions of O 3 of interest are also slower than those of OH. However, the kinetics are essentially the same, i.e., Eq. (Y) applies where A is O 3 and B the stable reactant of interest. Figure 5.10 shows the results of a typical experiment. An FEP Teflon reaction chamber was subdivided into two chambers using metal rods, and O 3 was then introduced into one side of the chamber at concentrations that gave ~ 1 ppm in the entire chamber when mixed. The organic in excess concentration was injected into the other side, both the organic and O 3 being in ultrahigh-purity air. After removal of the metal barriers and mixing, O 3 was monitored as a function of time using a commercial chemiluminescence ozone analyzer. In this particular case, a rate constant for the reaction of 0 3 with o-cresol of 2.55 • 10 -19 cm 3 molecule -~ s-~ was obtained (Pitts et al., 1981). Such experiments can also be carried out with 0 3 in great excess; however, a technique must be available for following the concentration of the reactant X with time, and corrections may have to be made for changing ozone concentrations due to wall losses during the experiment. In addition, interferences from secondary reactions are more likely under these conditions.

Many of the rate constants for gas-phase reactions of atmospheric interest reported in the literature were actually determined not as absolute values but rather as a ratio of rate constants. Thus if the absolute value for one of the rate constants has been determined

k2

A 4- X 2

, P1,

(20)

> P2"

(21)

X~ and X 2 are the reactants that compete for A, and P~ and P2 are the respective products of these reactions. By monitoring the concentrations of the reactants X~ and X2, the concentrations of the products P~ and P2, or the change in the concentration of one of these as the second reactant is added with time, one can obtain the rate constant ratio k l / k 2. For example, if X~ and X 2 are monitored, the relevant rate laws are as follows: -d[X~] dt = kl[A][Xl ]' -d[X

2]

= k 2 [ a ] [ x 2 ]"

dt

Rearranging one obtains - d ln[X 1] dt = k~[A],

(CC)

- d ln[X 2 ] dt = k2[A].

(DD)

Combining (CC) and (DD) to eliminate [A] yields Eq. (EE):

1 d ln[X1 ]

[A] =

k1

dt

=

1 d ln[X2 ] k2

dt

.

(EE)

Integrating from time t = 0 when the initial concentrations are [X1] 0 and [X2]0, respectively, to time t when the concentrations are [X1] t and [X2] t gives Eq. (FF) In

[X1]0 [Xl] t

C. LABORATORY TECHNIQUES FOR DETERMINING RELATIVE RATE C O N S T A N T S FOR GAS.PHASE REACTIONS

149

-

kl k2

[X2] 0

In ~ .

(FF)

[X2] t

Thus the concentrations of X~ and X 2 as a function of reaction time, plotted as given by Eq. (FF) (i.e., ln([X 1]0/[ X 1],) versus ln([X 2]0/[ X 2],)), can be used to derive the rate constant ratio k l / k 2. If an absolute value is known for one of the two rate constants from independent studies, then an absolute value for the second one can be obtained. For example, Figure 5.11 shows typical results from a relative rate experiment on the reaction of chlorine atoms with some simple alkanes (Beichert et al., 1995). The chlorine atoms in this case were produced by the

15 0


0.8 "~.

Ethane

Propane

o

.

f

,

/

0.6 >-

m

~o >" v

0.4

t-m

0.2

0

0

0.2

0.4

0.6

0.8

1

1.2

1.4

1.6

In([X]o / [X]t)

FIGURE 5.11 Plotsof the relative decays of pairs of organics (Eq. (FF)) in the presence of chlorine atoms at room temperature (adapted from Beichert et al., 1995).

blacklamp photolysis of C12. The reaction vessel was a 50-L Teflon collapsible chamber into which C12 and the alkanes were introduced as a dilute mixture in air. The mixture was sampled periodically into a gassampling valve interfaced to a gas chromatograph so that the concentrations of the alkanes were measured as a function of photolysis time. The most common sources of OH used for relative rate studies include the photolysis of H O N O (e.g., see Cox, 1975) or alternatively methyl nitrite (CH3ONO) in air in the presence of NO (Atkinson et al., 1981): C H s O N O + h u ~ C H 3 0 + NO, CH30

+ 0 2 --+

H C H O + HO2,

HO 2 + NO + OH + NO 2.

(22) (23) (24)

Determination of OH relative rate constants for compounds that photolyze significantly in actinic radiation requires a nonphotolytic source of OH. Three such OH sources are H 2 0 2 - N O 2 - C O mixtures (Campbell et al., 1975, 1979; Audley et al., 1982), the thermal decomposition of H O 2 N O 2 in the presence of NO (Barnes et al., 1982), and O3-hydrazine reactions (Tuazon et al., 1983) or O3-alkane reactions in the dark (Finlayson-Pitts et al., 1993). However, in these cases, the reactant must not react with 03, HO2, or H202, and care must be taken in interpreting the data since these systems have the potential of being rather complex. Indeed, the rate constants derived have not always agreed well with literature values. Until the

general features of the mechanisms involved in the production of OH in these systems have been fully elucidated, the simultaneous production of other highly reactive species, and hence possible interfering secondary reactions, cannot be firmly ruled out. Relative rate techniques have the advantage that such relative measurements can be made with greater precision than absolute rate constant measurements because only relative, not absolute, concentrations of X 1 and X 2 need be measured. Indeed, precisions of 5% or better are common using these techniques. Note, however, that increased precision does not necessarily imply increased accuracy. Another advantage is that the species A, which is frequently a highly reactive free radical such as OH which is difficult to measure, need not be monitored in such experiments; only X 1 and X2, which are usually stable and easily measured molecules, such as hydrocarbons, are followed. Finally, relative rate experiments can often be carried out under conditions directly relevant to the atmosphere, e.g., low concentrations of the reactants in high pressures of air. The accuracy of the results, however, depends critically on knowing enough of the mechanistic details of the reaction system to be sure that the kinetic analysis, which is not always straightforward in complex systems, is valid. Furthermore, obtaining an accurate rate constant from the rate constant ratio ( k ~ / k 2) requires accurate knowledge of the second, reference rate constant.

D. REACTIONS IN SOLUTION


151

(3) Diffusion of the solvated species into the bulk phase of the droplet (4) Reaction of the species in the aqueous phase or at the interface itself (see Section E.1).

1. Interactions of Gaseous Air Pollutants with Atmospheric Aqueous Solutions Because of the gaseous nature of many of the important primary and secondary pollutants, the emphasis in kinetic studies of atmospheric reactions historically has been on gas-phase systems. However, it is now clear that reactions that occur in the liquid phase and on the surfaces of solids and liquids play important roles in such problems as stratospheric ozone depletion (Chapters 12 and 13), acid rain, and fogs (Chapters 7 and 8) and in the growth and properties of aerosol particles (Chapter 9). We therefore briefly discuss reaction kinetics in solution in this section and "heterogeneous" kinetics in Section E. The aqueous phase that serves as a reaction medium in the atmosphere is present in the form of clouds, fogs, rain, and particulate matter consisting of either an aqueous solution containing pollutants or a film of water surrounding an insoluble core (see Chapter 9). For example, at typical relative humidities, ~ 30-50% of the aerosol mass is due to water (Graedel and Weschler, 1981). However, many of the species that are believed to react in such atmospheric solutions, for example, SO2, 03, H202, and NOx, are emitted or formed in the gas phase. Before reactions can occur in solution, then, several steps illustrated in Fig. 5.12 must first take place: (1) Diffusion of the gas to the surface of the droplet (2) Transport of the gas across the air-water interface

Diffusion of gases is fast relative to diffusion in the aqueous phase; i.e., step 1 is fast relative to step 3. Thus diffusion coefficients for gases at 1 atm pressure are ~ 0.1-1 cm 2 s -~, whereas in liquids they are 10 -5 cm 2 s -~ for small molecules. As discussed in detail by Schwartz and Freiberg (1981), gas-phase diffusion, in most (but not all) cases, will not be the slowest (i.e., rate-determining) step. Gases dissolve in aqueous solution to various extents, depending on the nature of the gas. At sufficiently long times, an equilibrium can be established between the gas- and liquid-phase concentrations, which is described by Henry's law: [X] = H x P x ,

where [X] is the equilibrium concentration of X in solution (in M = mol L - l ) , Px is the gas-phase equilibrium pressure (in atm), and H x is the Henry's law constant (in M atm-1). Table 5.6 shows some values of H for some species of interest dissolving in aqueous solutions at 25~ (Schwartz, 1984a). They range from 10 .3 M atm -1 for relatively insoluble gases such as 0 2 to ~ 105 M atm -~ for highly soluble gases such as H 2 0 2 and H N O 3. Henry's law can be applied to predict solution concentrations only if certain conditions are met. Thus it

Gas Phase

\

Ga

^

Diffusion

>

,-"

(GG)

Evaporation

Products

FIGURE 5.12 Schematic diagram of uptake and reaction of gases in liquids.

15 2


TABLE 5.6 Henry's Law Coefficients (H) of Some Atmospheric Gases Dissolving in Liquid Water at 25~ Gas 02 NO C2H 4 NO2 a 0 3 N20 CO2 b SO2 b HONO b NH 3b H2CO H20 2

HNO 3 HO 2 OH PAN CH3SCH 3

H (mol L -1 a t m - l)

Reference e

1.3 • 10 .3 1.9 • 10- 3 4.9 X 10 .3 1 X 10-2 (0.82-1.3) • 10-2 2.5 • 10 .2 3.4 • 10 -2 1.22 49 62 6.3 • 103 (0.7-1.0) • 105 1.4 • 105c 6.9 • 104 d 2.1 • 105 (1-3) • 103 30

Loomis, 1928 Loomis, 1928 Loomis, 1928 Schwartz and White, 1983 Briner and Perrottet, 1939 Loomis, 1928 Loomis, 1928 Maahs, 1982 Schwartz and White, 1981 Van Krevelen et al., 1949 Blair and Ledbury, 1925 Martin and Damschen, 1981 Yoshizumi et al., 1984 Hwang and Dasgupta, 1985 Schwartz and White, 1981 Schwartz, 1984b Golden et al., 1990; Hanson et al., 1992 Holdren et al., 1984 Dacey et al., 1984

5 0.48-0.56

Source: Adapted from Schwartz (1984a). a Physical solubility; reacts with liquid water. b Physical solubility exclusive of acid-base equilibria. CAt 20~ d Temperature dependence also reported as H = exp[7.92 • 1 0 3 / r (K) - 15.44]. e See Table 8.1 for references and additional data.

assumes that there are no irreversible chemical reactions that are so fast that the equilibrium cannot be established. It also assumes that the surface of the droplet is an unimpeded air-water interface; as discussed in more detail in Chapter 9, some atmospheric aerosols may have an organic surface film (e.g., Husar and Shu, 1975; Chang and Hill, 1980; Graedel and Weschler, 1981; Gill et al., 1983; Pankow et al., 1994a,b, 1997; Goss and Schwarzenbach, 1998), which could alter the establishment of the equilibrium anticipated by Henry's law. In addition, the high concentrations of ions in solutions of high ionic strength such as sea salt particles (especially near their deliquescence point) can alter gas solubility. In this case, the Henry's law constants must be modified using Setchenow coefficients to take this effect into account (e.g., Kolb et al., 1997). Ionized a n d / o r hydrolyzed species may be formed from the dissolved gas in some cases. An important example is SO2, which dissolves to set up equilibria involving HSO 3 and SO 2- in a manner similar to CO 2 (see Chapter 8). One major difference between reactions in the gas phase and in solution is the presence of solvent molecules in the latter case. In the liquid phase,

Gas Phase

IIIIIII II I IIIII IIIII IIIII III IIII IIIIII IIIII II II -[]me

Solution

IIIIIIII

IIIIIIIIIII

-[]me

IIIIH

F I G U R E 5,13 Patterns of A - B collisions expected in the gas phase and in solution (adapted from Adamson, 1973).

molecules are in close contact, with the space between molecules being ~ 10% of the distance between their centers. Reactants thus have a number of nearest neighbors, in the range of ~ 4-12, with which they can collide. The reactants can then be thought of as existing in a solvent "cage," in which they undergo many collisions before breaking out of that particular environment. If two participants required for a chemical reaction diffuse into such a solvent cage, they will then be held together for a period of time and undergo a number of collisions with each other; such a series of collisons is known as an encounter. Because of this cage effect, highly reactive species such as atoms and free radicals that are formed in the cage, for example, by photolysis, have a much higher efficiency of recombination than if they were in the gas phase. Compared to the gas phase, then, reactants take longer to diffuse together, but once they find themselves as nearest neighbors, they undergo a series of collisions rather than separating after one collision; this difference is illustrated in Fig. 5.13 (Adamson, 1973). As a result, for neutral nonpolar reactants (as opposed to ions; see later), the rate constants in solution are expected to be approximately equal to those in the gas phase. Treatment of systems in which gas-phase diffusion, mass accommodation, liquid phase diffusion, and reaction both in the bulk and at the interface must be taken into account is discussed in Section E.1.

2. D i f f u s i o n - C o n t r o l l e d R e a c t i o n s of U n c h a r g e d N o n p o l a r S p e c i e s in S o l u t i o n Let us first consider a very fast reaction between uncharged nonpolar reactants in solution. In this case, the rate is controlled by the number of encounters. Once A and B diffuse into the same solvent cage, they will react; hence the rate of these diffusion-controlled reactions is determined by how fast A and B diffuse together in solution. Fick's first law describes the rate of diffusion of a species A in solution across an area E in the direction


of the x axis, for example. The rate of diffusion, J = d n / d t (in molecules s - l ) , is given, according to Fick's first law, by

J =

dn dt

6 [ NA] = - D E ~ 6x

where dn is the amount of A crossing the area E (cm 2) in time dr, D is the diffusion coefficient (in units of cm 2 s - l ) , and 6 [ N a ] / 6 x is the gradient (in units of molecules cm -4) of the concentration of A in the x direction. Starting with Fick's first law, one can calculate for a solution of two reactants A and B the frequency of A - B encounters, which is in effect the reaction rate constant for diffusion-controlled reactions. This is given by the following, in units of L m o l - 1 s- 1: k = 4vrrABDAB(6.02

•

1020).

(HH)

tAB is the distance between the centers of the molecules when they react and DAB = (D A + DB) , where D A and D B are the diffusion coefficients of A and B, respectively. For typical values of tAB 0.4 nm and D A = D B =2• 10 -5 cm 2 s -1 a rate constant of ~101~ L mo1-1 s -1 is obtained. In solution then, diffusion-controlled bimolecular reactions between uncharged species occur with rate constants ~ 10 l~ L mo1-1 s -1. For reactions with significant activation energies a n d / o r steric requirements, the rate constants are correspondingly lower. In this case of uncharged, nonpolar reactions, there is little interaction between the reactants and the solvent. As a result, the solvent does not play an important role in the kinetics per se, except through its role in determining the solubility of reactive species and cage effects. The rate constants for such reactions therefore tend to be similar to those for the same reactions occurring in the gas phase. Thus, as we saw earlier, diffusion-controlled reactions in the gas phase have rate constants of ~ 1 0 -1~ cm 3 molecule -1 s -1, which in units of L mo1-1 s -1 corresponds to ~ 6 • 10 l~ L mo1-1 s -1, about equal to (usually slightly greater than) that for diffusion-controlled reactions in solution.

153

For ions and polar molecules, the nature of the solvent is an important factor in solution-phase reactions. Following the derivation of Laidler and Meiser (1982), we first consider the reaction between two ions A and B with charges Zae and ZBe, respectively, where e is unit electronic charge and Z a and Z B are the number of unit charges on the ions, i.e., are whole positive or negative numbers. The electrostatic force ( F ) between these two ions separated by a distance r in a vacuum is given by Coulomb's law,

ZAZB e2

F =

(II)

4~.a0r2 '

where e 0 = 8.85 • 10 -12 C 2 N -1 m -2, the permittivity of a vacuum. However, if the ions are immersed in a solvent, having a dielectric constant e, the electrostatic force between them is modified by the properties of the solvent. Equation (II) thus becomes F =

ZAZB e2 4~rg0g r 2 "

(JJ)

=

The higher the solvent dielectric constant e, the more the electrostatic force between the ions is reduced. From this expression for the force between two ions, one can calculate the work done to bring the two ions from infinite distance to the distance necessary to react, r = daB; this is equal to the change in free energy due to the electrostatic forces as the ions approach each other, zXGos. The total free energy change in bringing the ions together is the sum of this electrostatic term and a nonelectrostatic one, AG 0" Z~GTo T =

AG O + AGes, (6.02 • 1023)ZAZB e2

= AG O +

. (KK) 4~-a0 adAB

(Avogadro's number is included in the electrostatic term to convert to units of per mole rather than per molecule.) As seen in Section A.3.b, the transition state form of the rate constant is given by kT

k = - ~ e -AG~ 3. R e a c t i o n s of C h a r g e d S p e c i e s i n S o l u t i o n If the reactants are ionic with opposite charges, the rate constant can be greater than 10 l~ L mol-1 s-1 due to the favorable attractive forces. For example, the rate constant for the reaction of H § with O H - in aqueous solutions at 25~ is ~ 1011 L mo1-1 s -1. On the other hand, the electrostatic repulsion between ions of like sign can significantly slow their reaction. Similarly, if the reactants are polar molecules, electrostatic forces between them and the solvent may come into play.

The free energy of activation zXG~ for bringing two ions to the necessary distance daB in order to react is given by equation (KK). Thus the natural logarithm of the rate constant becomes I n k = In

kT

AG O

h

RT

(6.02

X

1023)ZAZBe2

4~reo edABRT

ZAZB e2 In k = I n k o -

4~'~0 ~dABkT '

(LL)

154


where In k 0 = In

kT

AG o

h

RT

transition state. The concentration of the activated complexes can be obtained from the equilibrium assumed between the reactants and the transition state"

and the Boltzmann constant k has been substituted in the electrostatic term for R/(6.02 • 10 23) = k. The term k 0 is the rate constant for the ion reactions in a medium where a = % that is, when the electrostatic forces have become zero. Equation (LL) predicts that the rate constant for the reaction between two ions in solution will depend on the dielectric constant a, and hence the nature of the solvent. A plot of In k against 1 / ~ should be a straight line with slope of --(ZAZBe2/4~oOodABkT). From the slope, dAB can be obtained. Experimentally, it is found that Eq. (LL) is indeed followed in many cases. In the atmosphere, most solutions available for reaction are aqueous. The dielectric constant for water at 25~ is e = 78.3. A second important factor for reactions in solution between ions or polar molecules is the ionic strength ( I ) of the solution. I is defined as

where C i is the molar concentration of the ith ion and Z i is its charge. For a 1 M solution of NaNO 3, for example, I = 1 M, whereas for a 1 M solution of NazSO4, I = 3 M. (CNa+= 2, ZNa+= +1, Cso4~-= 1, Zso 4~-= - 2 . ) In solution thermodynamics, the concentration (C) of ions is replaced by their activity, a, where a = Cy and 3' is the activity coefficient that takes into account nonideal behavior due to ion-solvent and ion-ion interactions. The Debye-Hiickel limiting law predicts the relationship between the ionic strength of a solution and 3' for an ion of charge Z in dilute solutions: (MM)

B is a constant that depends on the properties of the solution, for example, on its dielectric constant, and on the temperature. For water at 25~ B = 0.51 L ~/2 mo1-1/2. The Debye-Hiickel limiting law applies only for solutions of low ionic strength, for example, below 0.01 M for 1"1 electrolytes, such as NaNO 3, and below ~ 0.001 M for electrolytes of higher charge. The influence of ionic strength on solution rate constants can be anticipated by again resorting to transition state theory. For the reaction kp

A+B~

~ X 4=

aAaB

y:~ [X4=] TA[A]TB[B] "

Thus Rate = kpK 4=

TATB[A][B]

= k[A][B], where the reaction rate constant is given by YAYB

k= kpK4=~

4=

Yx Thus

log k = log(kpK 4) + log

YAYB 4=

Yx

Using the Debye-Hiickel limiting law for the relationship between the activity coefficients y and the ionic strength of the solution, one finds

1

I = -~ ~_, Ci Z2,

log y = - B Z 2 I 1/2.

ax.

~ Products,

the rate of the reaction is given by kp[X 4=], where [X 4=] is the concentration of the activated complex in the

logk=logk

0 +lOgyA + l o g y B - l o g y x

*,

= log k o + 2BZAZB I1/2. With B = 0.51 L 1/2 mo1-1/2 for aqueous solutions at 25~ this becomes log k = log k o + 1.02ZAZB I1/2.

(NN)

Thus a plot of log k against 11/2 should give straight lines of slope 1.02ZAZ B and intercepts of log k 0. The constant k 0 is seen to be the rate constant in a solution of zero ionic strength, that is, at infinite dilution. Equation (NN) also predicts that reactions between ions of the same sign should speed up as the ionic strength increases, whereas reactions between oppositely charged ions should slow down with increasing ionic strength. For reactions between an ion and an uncharged molecule (Z B = 0), ionic strength should not alter the rate constant. These relationships have been confirmed for solutions that are sufficiently dilute so that the Debye-Hiickel law is applicable (Fig. 5.14). As might be expected, deviations are observed at higher ionic strengths (e.g., see the text by Benson, 1960, for a more detailed discussion). This effect of ionic strength on solution rate constants is very important in studying reactions relevant to atmospheric chemistry. Thus care must be taken to study the effects of ionic strength over a range that


0.6 0.4

t

B

"

C

~"~'l o 0.2

-0.2 -0.4 -0.6

0

0.1

0.2

0.3

11/2

F I G U R E 5.14 Variation of rate constant with ionic strength (I) of the solution for reactants having different charges. Reactions: (A) [Co(NH3)sBr] 2+ + Hg 2+ + H 2 0 -+ [Co(NH3)5(H20)] 3+ +(HgBr)+; (B) $202- + I - - + (intermediates)-+ I j + 2SO 2- (not balanced); (C) [O2NNCOOEt]-+ O H - - + N20 + C O 2 - + EtOH; (D) cane sugar + O H - - + invert sugar (hydrolysis reaction); (E) H202 + H + + B r -+ H 2 0 + 1/2Br 2 (not balanced); (F) [Co(NH3)sBr] 2+ + O H - - + [Co(NH3)5(OH)] 2+ + Br-" (G) Fe 2+ q- C0(C204 )3--+ Fe 3+ + [Co(C204)3 ]4-. (Adapted from Benson, 1960.)

approximates those found in the atmosphere. Aerosols in polluted urban areas can be highly concentrated solutions with ionic strengths in the range of 8-19 M (Stelson and Seinfeld, 1981); reactions in these solutions will not follow the "ideal" relationships discussed earlier. On the other hand, cloud water and rainwater in clean areas contain much lower solute concentrations; for example, from the ionic composition of precipitation samples in the maritime area of Cape Grim, Australia (Ayers, 1982), the ionic strength can be calculated to be ~10-3 M.

4. E x p e r i m e n t a l T e c h n i q u e s U s e d for S t u d y i n g Solution Reactions The approaches to studying reaction kinetics in the liquid phase are analogous to those in the gas phase, that is, the use of various spectroscopic techniques to follow the loss of one reactant in the presence of a large excess of the second reactant. UV-visible spectroscopy is a primary tool for following both stable species and radicals in solution. As discussed in Chapters 7, 8, and 9, there are a number of free radical species whose reactions in the aqueous phase drive the chemistry of clouds and fogs. These include OH, HO2, NO3, halogen radicals such as Cly, sulfur oxide radicals, and RO 2. Generation of these radicals in the liquid phase for use in kinetic

155

studies is typically carried out using either flash photolysis or pulse radiolysis. Flash photolysis can be carried out using either broadband light sources in the 200- to 300-nm range, with pulse durations of the order of microseconds, or lasers with specific wavelengths and pulse durations of nanoseconds to femtoseconds. The advantages of lasers lie in the use of specific wavelengths, which minimizes the simultaneous photolysis of reactants or products at other wavelengths that can occur with broadband light sources, and in the availability of higher light intensities to generate larger radical concentrations. Excimer lasers have proven especially useful: ArF at 193 nm, KrF at 248 nm, XeC1 at 308 nm, and XeF at 351 nm. Pulse radiolysis relies on the interaction of highenergy ionizing radiation to generate free radicals during a short (/xs to ns) radiation pulse (see Pagsberg et al., 1995). The hydroxyl radical can be easily generated through processes described shortly, and other species by secondary reactions of OH. With water as a solvent, bombardment with ionizing radiation generates electrons (e-), H2 O§ and excited water molecules, H20*. The electrons either form hydrated electrons, eaq , or ionize additional water molecules. The hydroxyl radical is generated from reactions of H2 O§ and H20*: H 2 0 + + H 2 0 --+ OH + H 3 0 +, H 2O*

--->

H + OH,

(25) (26a)

--+ H 2 + O.

(26b)

Higher OH yields can be obtained if the solution contains N 20: eaq q- N 2 0

--+ N 2 + O

,

O - + H 2 0 --+ OH + O H - .

(27) (28)

To generate other free radicals, OH can be reacted with other species (e.g., see Zellner and Herrmann, 1995). For example, OH + HSOf + H 2 0 + 803-,

(29)

OH + HCO~ + H 2 0 + CO~-,

(30)

OH + C I - ~ C1OH-,

(31,-31)

C 1 O H - + H + ~ C1 + H 2 0 ,

(32,-32)

C1 + C1- + C12.

(33)

In the case of the chlorine reactions, the involvement of H + in the second step means that these reactions are efficient sources of chlorine atoms only at a pH less than about 4. Br- can be converted to Br atoms in reactions analogous to those for chlorine, but in this case, generation of atomic bromine occurs up to a pH of about 11 (Zellner and Herrmann, 1995).

156


Photodiode array detector

Illllllllllllll

I Fast pulser I

Spectrograph

Controller

Lens

.a,ooen, amo | Computer

____--

lamp

Lens -- Mirror Mirror with hole

~~s White ~

cell mirrors~~=V

Pulse-delay/ trigger unit

~n n7n~ ull[~

Lens

Excimer laser

Cell

F I G U R E 5.15 Schematic of typical apparatus used to study kinetics of reaction in the liquid phase (adapted from Zellner and Herrmann, 1995).

The reaction of OH with nitrate ion is endothermic so OH cannot be used to generate NO 3 in solution. It is often generated by reactions such as 82 O 2 -

-Jr- h v --+

2SO4,

SO 4 + NO~- ~ SO 2- + NO3,

(34)

(35, -35)

or the photolysis of ceric ammonium nitrate solutions at 350 nm (e.g., Alfassi et al., 1993): 2

CeIV(NO3)6

m

m

+ hv ~

NO 3 + CeIII(NO3)~ . (36)

Figure 5.15 is a schematic diagram of a typical apparatus used to study aqueous-phase kinetics (Zellner and Herrmann, 1995). An excimer laser is used to generate the free radical species of interest which is located in a cavity with White cell optics used to obtain long total path lengths (e.g., 60 cm) using a shorter base path (see Chapter ll.A.lc). The time dependence of the concentrations of the free radical reactant can be followed using its absorption of light from a halogen or D 2 lamp. In this particular apparatus, a photodiode array detector is used so that a range of wavlengths can be followed, rather than a single wavelength. This allows several different reactants and products to be monitored simultaneously, or if only one absorbing species is present, its absorption spectrum can be obtained.

E. LABORATORY TECHNIQUES FOR STUDYING HETEROGENEOUS REACTIONS "Heterogeneous" reactions in the context of atmospheric chemistry are defined as those between gases and either solids or liquids. Heterogeneous reactions have been suggested for decades as being important in the atmosphere. However, historically, forays into this area by a variety of scientists usually concluded in a rapid retreat, due to highly variable results that were (and, unfortunately, continue to be) common. The recognition of the importance of condensed phases in the oxidation of SO 2 to sulfuric acid (see Chapter 8.C.3) and the discovery of the dramatic Antarctic ozone "hole" (see Chapter 12.C.4) reinforced the key role of heterogeneous reactions, and the atmospheric chemistry community has tackled this area anew. As a result, while we are still in the "dark ages" in terms of understanding kinetics and mechanisms of these processes on a molecular level, a great deal of progress has been made in the past decade. There are many different types of surfaces available for reactions in the atmosphere. In the stratosphere, these include ice crystals, some containing nitric acid, liquid sulfuric acid-water mixtures, and ternary solutions of nitric and sulfuric acids and water. In the troposphere, liquid particles containing sulfate, nitrate, organics, trace metals, and carbon are common. Sea

E. LABORATORY TECHNIQUES FOR STUDYING HETEROGENEOUS REACTIONS

salt particles dominate in marine areas. In addition, there are large episodic sources of particles emitted directly into both the troposphere and stratosphere, such as rocket exhausts where particles containing carbon soot, alumina, and metal oxides can be emitted in large quantities. Before we describe some of the common techniques used to study the kinetics and mechanisms of heterogeneous reactions, a few words regarding the difficulties in this area are appropriate. To put these in perspective, consider first the current state of understanding of gas-phase kinetics. There are a number of both absolute and relative rate techniques available for studying gas-phase reactions, and the methodologies for preparing reactants and measuring products are generally quite well developed. As a result, agreement to within 15% on gas phase reaction rate constants measured using different techniques and by different laboratories is now common. In addition, understanding of the reactions at a molecular level is generally good. In contrast, in most heterogeneous reactions, we really do not even understand what one of the reactants, the surface, looks like on a molecular level; i.e., the condensed-phase molecule and its environment that the incoming gaseous reactant encounters is not well characterized. An example of our incomplete understanding of the nature of surfaces is controversy about effective surface areas for ice and whether ice surfaces prepared in the laboratory for example, are porous or not (e.g., see Keyser et al., 1993; and Hanson and Ravishankara, 1993a). In addition, how the surface changes during reaction and how such changes affect both the reactivity and mechanism are unclear. One well-recognized phenomenon is that of saturation of the surface, i.e., complete reaction of the surface species so that the reaction comes to a halt. However, there are other potentially important aspects as well. For example, it is well known in surface science that adsorbates can lead to restructuring of the surface (Somorjai, 1994), yet this phenomenon has yet to be demonstrated on surfaces of atmospheric interest. This is in part due to the incompatibility between atmospheric conditions (air at pressures up to 1 atm and containing significant amounts of water vapor) and the ultrahigh-vacuum conditions typical of surface science studies. However, marriage of these two fields will be ultimately needed for a complete understanding of heterogeneous processes in the atmosphere, and progress is being made with some systems (e.g., Hemminger, 1999). Different terms and symbols have been used in the literature to differentiate reversible, physical uptake from irreversible uptake via chemical reactions. Addi-

15 7

tional confusion arises from the fact that the observed uptake of a species from the gas phase is usually a net uptake affected by a number of factors such as changes in the surface during the uptake (e.g., due to saturation) and reevaporation into the gas phase due to limited solubility of the species. However, the most common terminology now in use is the following:

Surface reaction probability ('Yrxn) is the net fraction of gas-condensed phase collisions that leads to the irreversible uptake of the gas due to chemical reaction. The symbol "Yrxn (or sometimes q~) is most commonly used for reaction probabilities. Mass accommodation coefficient ( a ) is the fraction of gas-condensed phase collisions that result in uptake of the gas by the condensed phase: Number of gas molecules taken up by the surface OL

~

o

Number of gas-surface collisions

(oo) This is the probability that a molecule that strikes the surface will cross the interface into the condensed phase. It does not represent the net uptake; i.e., it does not include the reverse effect of reevaporation from the condensed phase into the gas phase, a is also sometimes referred to as a sticking coefficient for uptake on solid surfaces. Net collisional uptake probability (')/net) is the net rate of uptake of the gas normalized to the rate of gas-surface collisions and since this is what is measured, is also often referred to as 7moas. The net loss of the gaseous species in the presence of a condensed phase having known volume and surface area is what is typically measured in laboratory studies. To obtain information on the fundamental processes contributing to this net measured uptake, all of the processes shown schematically in Fig. 5.12 must be taken into account: diffusion of the gas to the surface, uptake, diffusion in the liquid phase, and reaction either in the bulk or at the interface itself. We therefore first briefly discuss the analysis of systems that involve diffusion in the gas and liquid phases, uptake, and reaction in the bulk liquid or at the interface. Following that, we give a brief description of some of the most common methods used to measure mass accommodation coefficients and reaction kinetics for heterogeneous atmospheric reactions. Included are some new approaches that appear to be especially promising. For a review of this area, see Kolb et al. (1995, 1997).

15 8


1. Analysis of Systems w i t h G a s - and

Liquid-Phase Diffusion, Mass A c c o m m o d a t i o n , and Reactions in the Liquid Phase or at the Interface Both for laboratory measurements and for atmospheric processes, the uptake of a gas into a liquid followed by reaction involves a number of different physical (e.g., diffusion and uptake at the interface) as well as chemical processes. These were depicted in Fig. 5.12. We treat here in more detail the individual steps and how the net uptake of a gas into solution is determined by these steps.

1. Transport of the gas to the surface and the initial interaction. The first step in heterogeneous reactions involving the uptake and reaction of gases into the liquid phase is diffusion of the gas to the interface. At the interface, the gas molecule either bounces off or is taken up at the surface. These steps involve, then, gaseous diffusion, which is determined by the gas-phase diffusion coefficient (D g) and the gas-surface collision frequency given by kinetic molecular theory. 2. Uptake at the interface. If the gas molecule is taken up at the surface, it enters the interface region and then the bulk. The efficiency of uptake involving crossing the interface is described by the mass accommodation coefficient (c~) defined earlier. Molecular-level mechanisms by which gas molecules are taken up into liquids are discussed elsewhere (e.g., see Davidovits et al., 1995; Taylor et al., 1996, 1997; and Nathanson et al., 1996). 3. Diffusion into the bulk. This is determined by the diffusion coefficient in the liquid (D1). Diffusion within the bulk aqueous phase is much slower than gas-phase diffusion and can be rate-limiting under conditions of high reactant concentrations where the rate of the chemical reaction is high. This appears to have been a problem in some experimental studies of some aqueous-phase reactions relevant to the atmosphere where either bulk solutions or large droplets and reactant concentrations higher than atmospheric were used (Freiberg and Schwartz, 1981). 4. Henry's law equilibrium. If there is no reaction in the liquid phase (or it is slow relative to uptake and diffusion), the gas-liquid system eventually comes to equilibrium, which can usually be described by Henry's law discussed earlier. This does not reflect a lack of uptake of the gas at equilibrium but rather equal rates of uptake and evaporation; i.e., it is a dynamic equilibrium (see Problem 12). The equilibrium between the gas-and liquid-phase concentrations is characterized by the Henry's law constant, H (mol L -1 atm-1), where H = [X]/P X.

5. Reaction in the bulk. Reaction can occur in solution close to the surface or throughout the bulk of the liquid phase, depending on the speed of the reaction compared to diffusion. We shall see that whether the reaction occurs close to the surface or throughout the bulk has important implications for the kinetics, since in the former case, the reaction depends on the particle surface area, whereas in the latter it depends on the particle volume. We shall treat here reactions occurring in the condensed phase as if they are first-order, irreversible reactions with a rate constant k. Of course, this also applies to second-order reactions with the second reactant, B, in great excess, since they are then pseudofirst-order with k' = k[B].

6. Reactions at the interface. There has been increasing recognition that reactions may also occur at the interface itself. That is, species such as SO 2, NH 3, and organics do not simply cross the interface by physical transport but rather form unique chemical species at the interface (e.g., Donaldson et al., 1995; Allen et al., 1999; Donaldson, 1999; Donaldson and Anderson, 1999). These unique interface species can then react at the surface without actually being taken up into the bulk of the solution. Although relatively little is currently known at a molecular level about such processes, reactions in this "fourth phase" may prove to be very important in atmospheric processes, for example in the generation of HONO in the NO 2 reaction with water at surfaces (see Chapter 7.B.3b). The measured, net uptake of a gas into a liquid phase can be related to these various processes, i.e., to Dg, c~, D1, H, and k. However, concentrations in both the gas and liquid phases as well as the volumes and surface areas available for reaction are quite different in laboratory studies compared to the atmospheric situation. We shall therefore also examine how net gas uptake measured in laboratory studies can be related to uptake under atmospheric conditions. The physical and chemical processes occurring in a gas-liquid system are often treated in terms of a resistance model described in Box 5.2. As discussed there, the net uptake of gas (Ynot) can be treated under some conditions in terms of "conductances," F, normalized to the rate of gas-surface collisions. Individual conductances are associated with gas-phase diffusion to the surface (Fg), mass accommodation across the interface (c~), solubility (F~ol), and finally, reaction in the bulk aqueous phase (Frxn). This leads to Eq. (QQ):

1

-

1

1 + -- +

1

rrxn+r o,

(QQ)

15 9


BOX 5.2 MODEL OF GAS-DROPLET INTERACTIONS

RESISTANCE

We shall treat the individual processes in terms of the rate of transfer of gas across a surface of unit area per second. However, this rate will be expressed relative to the number of gas-surface collisions per second, given according to kinetic molecular theory by Number of collisions per unit area second of gas with surface

Nguav 4

Ng ~ 8RT -

4

~ RT

vrM = Ng

2~//

. (PP)

In Eq. (PP), Ng is the gas concentration (molecules cm-3), U~v is the average molecular speed in the gas phase, R is the gas constant (J K-1 mol-1), T the temperature (K), and M is the molecular weight (kg) of the gas. The normalized rates, i.e., divided by the rate of gas-surface collisions in Eq. (PP), will be referred to as "conductances," F, for reasons that will become apparent shortly. However, the reader should keep in mind that these "conductances" just reflect the speeds of the individual processes. The uptake of a gas into a liquid followed by its reaction can be described by a series of coupled differential equations (Danckwerts, 1951, 1970), which can be solved exactly only for some specific cases. However, under many conditions the individual processes can be treated as if they are not coupled. In this case, an approximation that has found widespread use (e.g., see Schwartz and Freiberg, 1981; Schwartz, 1986; and Kolb et al., 1995, 1997), and that helps to assess the relative importance of each of the terms, is to treat the individual processes in terms of an electrical circuit (Fig. 5.16). Dimensionless "conductances," F [where conductance = (resistance) -1 ], associated with each process reflect rates normalized to the rate of gas-surface collisions, and the corresponding "resistances" are given by 1 / F . The net, overall measured resistance, (Ynot) -1, is then related to the individual resistances (see Problem 7) by 1

")/net

_

1

I-'g

1 +m+

O/

1

['rxn @ ['sol

(QQ)

As already discussed, ~net is a net probability normalized to the number of gas-surface collisions and is the parameter actually measured in experiments (and hence also often referred to as Ymoas)" In Eq. (QQ), each conductance represents one of the processes involved; i.e., Fg involves the conductance for gas-phase diffusion, ['rxn that for reaction in the aqueous phase, and Fso1 that for solubility and diffusion into the bulk. Each of the terms has been normalized and made unitless by dividing by the rate of gas-surface collisions, Eq. (PP), except for a, which by definition is already normalized to this parameter. Let us now examine each of these terms individually. Diffusion of the gas to the surface (Fg). As described by Fuchs and Sutugin (1970, 1971) in their comprehensive t r e a t m e n t of highly dispersed aerosols, the rate of transfer of mass to the surface of a spherical particle by diffusion of a gas is described by

dN clx

X2---7 - = constant,

(RR)

where N is the gas concentration a distance x from the center of the particle and the constant depends on the boundary conditions at the surface of the particle. Take the case where the Knudsen number, Kn, defined as the ratio of the mean free path of the gas to the radius of the particle (a), approaches 0 (i.e., the mean free path is small compared to the particle size); all molecules colliding with the surface are taken up at the surface and the boundary condition is (N)x= a = 0. Fuchs and Sutugin show that Eq. (RR) can be integrated to obtain the gas concentration as a function of x, N = N~(1 - a/x), where N~ is the gas concentration at x = % which can be taken as Ng for a constant gas-phase concentration. They also show that for small Knudsen numbers where Kn ~ 0, the rate of diffusion of the gas to the surface of the particle of radius a is given by Number of molecules per second diffusing to the surface = 4~-Dg aNg.

(ss)

The surface area of a spherical particle of radius a is 4~-a 2. Thus, the rate of diffusion to the surface

160


Gas Phase

Resistance = !

Fg --,/N/N/N/N/"

Transport to

across \

I

surface

interface~

'

Chemica! reaction ~

/X/X/X ]

I '

; / i /

I I

-J

Finterface F I G U R E 5.16 Schematic of resistance model for diffusion, uptake, and reaction of gases with liquids. Fg represents the transport of gases to the surface of the particle, c~ the mass accommodation coefficient for transfer across the interface, Fso1 the solubilization and diffusion in the liquid phase, Frx n the bulk liquid-phase reaction, and Finterfac e the reaction of the gas at the interface.

per unit surface area is given by Rate of diffusion to surface per cm 2 per second

= 4rrDgaNg/4rra 2= 2DgNg/d,

(TT)

where d = 2a is the particle diameter. Normalizing to the rate of collisions, Fg becomes

['g = 8 D g / u a v d .

(VV)

As discussed by Fuchs and Sutugin (1970, 1971) and Motz and Wise (1960), in this continuous regime, distortion of the Boltzmann velocity distribution in the region close to the surface occurs if there is rapid uptake. In effect, the normal thermal velocity distribution is distorted so that the effective speed toward the surface is higher. In the case of a surface where the uptake occurs on every collision, the net speed toward the surface is effectively doubled. This adds an additional term to the rate of transfer of the gas to the surface, which when normalized using Eq. (PP), gives an additional "resistance" term of - 1 , / 2 . The overall normalized conductance is therefore given by

1 Fg

=

u~vd

1

8Dg

2

.

(VV)

This applies to diffusion to a planar surface in the continuous regime, when the Knudsen number is small, and is the expression for gas diffusion most

often encountered in the atmospheric chemistry literature. For the other extreme of the free molecular regime where K, --+ % the particle radius is small compared to the mean free path. In this case, the thermal velocity distribution of the gas is not distorted by uptake at the surface. In effect, the gas molecules do not "see" the small particles. For this case, Fuchs and Sutugin (1970, 1971)show that for diffusion to a spherical particle of radius a Number of molecules diffusing to surface per second

= rra2oZcuavNg,

(WW)

where a c is the probability of uptake at the surface and U~v is the mean thermal velocity. For intermediate regimes of K,, which are common both in the atmosphere and in many laboratory studies, exact calculations are not readily carried out. Fuchs and Sutugin (1970, 1971) suggest the form Number of molecules diffusing to surface per second

= 4rrDgaNg/(1

+ hKn).

(XX)

As K n + 0, Eq. (XX) approaches Eq. (SS), as expected. Values of h are provided in the literature (Fuchs and Sutugin, 1970, 1971); as an approximation, 1.33 + 0.71/Kn a -. (YY) 1 + 1/K

161


Interface

I

--

.....

is the combination of the net flux in and out, i.e., ~C 7net 6C 62C ] 6c + d x ~ y , (BBB) = -D1--2- + D1 ox 6x 6x

....... ---)'-Fin ......... --),- Fout

\x /

slope

t,,.~ ~

~

which reduces to 6C

62C

6t - D1 6x 2 "

/ d x

> ~$ope=8_~_~+dx(~f~2x~)

x

x+dx

(CCC)

The rate of transfer (Rt, in units of molecules or moles per cm 2 per second) of the species across a plane at x = 0, is given by

X

F I G U R E 5.17 Schematic diagram for treatment of diffusion of a species in one dimension in a liquid.

Uptake across the interface into solution (c~). By definition, this is described by the mass accommodation coefficient, c~, and 1/c~ is the "interfacial resistance." Solubility and diffusion in the liquid phase (Fsol). Consider the diffusion of a dissolved species in one dimension into the bulk solution from the interface region. The concentration of the species in the liquid will depend on time as well as on the distance from the interface. It is assumed first that no reaction is taking place in the aqueous phase. Figure 5.17 shows a general case of diffusion in the x direction across a plane of unit area. The concentration of the diffusing species is taken as c. The net flux in at x, Fin (molecules cm 2 s-1), is given by ~C Fin = - D l ~ ,

(ZZ)

6x

where D 1 is the diffusion coefficient in the liquid phase (units of cm 2 s-1) and ( 6 c / 6 x ) is the concentration gradient at x. The concentration gradient (i.e., slope of the concentration versus distance) is also changing with distance. At a distance dr from the position x, this concentration gradient is given by [ ( 6 c / 6 x ) + dx(62C/6X2)] and the net flux across the plane at dx, i.e., out of the volume bounded by x and (x + dx), is given by 6C

F~

62C

+D1 6x + dx 6x 2

(AAA)

The net flux in the region bounded by x and (x + dx)

Rt=-D

1 -~x x=0 9

(DDD)

The concentration gradient (6c/6x)~= o at the surface depends on time because as uptake occurs, reevaporation back to the gas phase becomes increasingly important and the magnitude of the concentration gradient decreases. Equation (CCC) can be solved to obtain the rate, Rt, under certain boundary conditions. Take the case where the concentration in the bulk liquid is given by Cl,bulk at time t = 0 as well as at x = oc for times t > 0. It is assumed that there is a thin layer at the surface that contains the dissolved species in equilibrium with the gas immediately adjacent to the surface. This interface concentration is denoted as Cl,interfac e. Under these conditions, Eq. ( C C C ) c a n be solved (see Danckwerts, 1970, pp. 31-33) to obtain the rate of transfer per unit surface area after exposure time t, as Rate = (Cl,interfac e -- Cl,bulk)

~

~rt

.

(EEE)

As intuitively expected, the rate of transfer across the interface depends on the difference in the liquid-phase concentrations at the interface and in the bulk and on the diffusion coefficient in the liquid. In addition, it depends inversely on the time of exposure of the liquid to the gas because of the increasing importance of reevaporation back to the gas phase at longer times. When Cl,bulk = 0, Eq. (EEE) becomes Rate = (Cl,interface)

~

D1 7rt

9

(FFF)

The dissolved species at the interface is considered to occur in a thin layer of thickness (Dlt) 1/2

162


and is also assumed to be in Henry's law equilibrium with the gas immediately adjacent to the interface. While Henry's law is commonly expressed as H = [X]/Px, with the Henry's law constant H in units of mol L -~ atm -~, it can also be expressed in unitless form if the gas-phase concentration is expressed in units of mol L -~ instead of in units of pressure. With this conversion (see Problem 5), the liquidphase concentration in the interface region, N 1 = Cl,interface, is related to the gas-phase concentration immediately adjacent to the interface, Ng, by N 1 = N g H R T (where R is the gas constant). Substituting into Eq. (FFF), the rate of transfer of the gas is given by Rate = N ~ H R T

~/

D1 7rt

.

(GGG)

Normalizing the rate of gas-surface collisions using Eq. (PP), one obtains Fs~ =

4HRT / D V 1. Uav 1rt

(HUH)

Note that in this case, Fso1 decreases with increasing time of exposure of the liquid to the gas. This reflects reevaporation from the liquid becoming increasingly important as the concentration of the dissolved species increases. If the liquid layer is very thin, as is the case for some particles in the atmosphere, the interface layer with thickness (Dlt) 1/2 may comprise the entire particle, in which case the liquid is in Henry's law equilibrium with the gas, there is no net uptake, and Eq. (HHH) is not applicable. Similarly, at very long reaction times, i.e., as t ~ % Fso1 ~ 0. That is, at very long exposure times, there is no net uptake because the system has come to equilibrium and the rates of uptake and reevaporation are equal. Reaction in the liquid phase (Frxn). Now consider the case where an irreversible, first-order reaction with rate constant k (s -1) takes place, in addition to diffusion and solubilization. Equation (CCC) becomes

at

-

D

1

-

kc.

(III)

This can be solved assuming that the concentration in the bulk liquid is 0 inside the drop at time t = 0 as well as at x = oc at longer times, and that the concentration in the thin layer at the interface is in Henry's law equilibrium with the adjacent gas (Danckwerts, 1970, pp. 33-37). Under these conditions, the rate of transfer of molecules across a

plane under conditions where kt >> 1 becomes independent of time and is given by Rate = Cl,interfac e ~ / D l k .

(JJJ)

The normalized rate for reaction is then given by

4HRT ~~/D,k

rrx n =

(KKK)

Uav

(see Problem 6). This applies to irreversible reactions or to those where the solubility of the reaction product is very large. Let us now return briefly to the question of the relationship between net reactive uptake coefficients measured in laboratory systems, where the liquid films are generally quite thick, and particles in the atmosphere, which can be quite small and hence effectively have thin liquid films. A measure of the distance from the interface in which the reaction occurs is the diffuso-reactiue length, 1, which is defined as 1=

~

D1 k "

(LLL)

Associated with this is the diffuso-reactiue parameter, q, where ~k

q=a

a

D1

(MMM)

1

and a is the particle radius. Using an approach similar to that discussed earlier, Hanson et al. (1994) showed that the effective reactive uptake coefficient for small drops, %, is related to that measured in thick laboratory films, Ymoas, by 1

1 -

'Ye

1 t

O/

1

")/meas (coth q -

l/q) "

(NNN)

For large values of c~, then, ")/e ~

Ymeas(cothq

- 1/q)

(OOO)

and the factor (coth q - 1 / q ) is the correction factor that must be applied in extrapolating laboratory measurements to particles in the atmosphere. When q is large, i.e., the diffuso-reactive length is small compared to the size of the particle, the correction factor is ~ 1 and the values for the reactive uptake coefficient measured in the laboratory can be applied to atmospheric particles. However, when q is small, i.e., the diffuso-reactive length is about the same as, or greater than, the particle radius, reac-

163


,9o 0 IJ_ t.0

0.1

0

0

0.01 0.01

I

I

I

I I IIII

I

0.1

1

10

//a F I G U R E 5 . 1 8 Correction factors for the measured uptake coefficient, Tmeas, as a function of the ratio of the diffuso-reactive length (1) to the droplet radius (a) (adapted from Hanson et al., 1994).

tion occurs throughout the particle volume. In this case, the correction factor becomes significant and yo is less than Ymo~s"Figure 5.18 shows the values of

As seen in Box 5.2, Eq. (QQ) can be shown to be equivalent to Eq. (PPP): 1

1 _

Ynot

1 + m +

a

Fg

4HRT

v/D1 m + Dv~-fk "a-t

U av

, (PPP)

where the symbols are as defined earlier, R is the gas constant and U~v is the average molecular speed in the gas phase. There are several limiting cases of Eq. (PPP) of interest in atmospheric systems:

1. Fast gas transport, high solubility, and / or fast reaction. In this case, 1/Yno t approaches 1 / a ; i.e., the maximum value for the measured uptake approaches the mass accommodation coefficient. 2. Fast gas transport, low solubility, and fast reaction. In this case, Fso1 (C3H702)* --+ C3H 6 -+- HO 2 at 1 atm pressure and room temperature but becomes important at higher temperatures, e.g., in combustion systems (Kaiser, 1995; Benson, 1996; Kaiser and Wallington, 1996a).

(23a)

There is a second path for the larger R O 2 radicals corresponding to addition of the NO and followed by isomerization to form an alkyl nitrate" RO 2 + NO --+ R O N O 2 .

R + O2

185

(23b)

The mechanism is postulated to be (Darnall et al., 1976) RO 2 + NO ~ (ROONO)* ~ RO + N 0 2 ,

(23a)

1 (RONO2). M RONO2.

(23b)

Hence one would expect the second channel to depend on the size of the radical and to show some pressure dependence as well as a negative temperature dependence, all of which are found to be the case. Table 6.5 shows the fraction of the total reaction forming the stable alkyl nitrate, i.e., the branching ratio k z 3 b / ( k 2 3 a -+- k23 b) for some simple alkylperoxy radicals at 1 atm and room temperature. The ratio is smaller for primary and tertiary R 0 2 ; in general, the ratio kz3b/k23 a for primary and tertiary alkylperoxy radicals is about 40 and 30%, respectively, of that for secondary R O 2 (Atkinson, 1997a). It can be seen from the data in Table 6.5 that the fraction of the total R O 2 + N O reaction that forms the stable alkyl nitrate at 1 atm can be substantial, as much as 35% for the larger radicals. (However, more recent

186

6. GAS-PHASE REACTIONS IN IRRADIATED ORGANIC-NOx-AIR MIXTURES

measurements by Aschmann et al. (1999) give lower nitrate yields for the hexyl, heptyl, and octyl alkylperoxy radicals, ranging from 15% for the sum of (2-hexyl + 3-hexyl) to 24% for the sum of (2-octyl + 3-octyl + 4-octyl RO2.) It is thought that such reactions may be significant sources of such alkyl nitrates and perhaps form part of the "missing NOy" (see Section E.3).

TABLE 6.5

Yieldsof RONO 2 in RO 2 + NO Reactions at Room Temperature and 1 atma Branching ratio = k23b / (k23 a -I- k23b)

Ethane Ethyl Propane 1-Propyl 2-Propyl n-Butane 1-Butyl 2-Butyl

O - - C : ~

(27)

This reaction may proceed via the initial addition of 0 2 to give a trioxy intermediate ( - C - O O O ' ) followed by rearrangement and elimination of HO 2 (e.g., Jungkamp et al., 1997). Atkinson (1997a, 1997b) recommends rate constants of 9.5 • 10-15 and 8 • 10-15 cm 3 molecule- 1 s- 1 at 298 K for primary R C H 20 and secondary R 1R 2CHO radicals, respectively. Scission of a C - C bond:

a

H

1b 0 CH3C~

+ CH2CH 3 H

0% /CCH2CH3 >CH 3 + H (28)

D. REACTIONS OF ALKYL (R), ALKYLPEROXY (RO 2), A N D ALKOXY (RO) RADICALS IN AIR M

Scission to produce the larger alkyl radical dominates. For example, in the case of the 2-butoxy radical in Eq. (28), path (b) dominates over path (a) by more than two orders of magnitude (see Atkinson, 1997b, and references therein). 9 If the alkoxy radical has four or more carbon atoms and can form a 6-membered transition state via the alkoxy oxygen abstracting H from a C - H bond, an intramolecular isomerization may occur: .H H "O ......... C."" I~CH3 CHBm~ ~c/CH2 H

R O + NO --+ R O N O , M

RO + NO 2 --+ R O N O 2.

H2

(29)

H Structure-reactivity relationships have also been developed for these isomerizations. For H abstraction from a CH3-X, CH2-XY, or C H - X Y Z , respectively, through a 6-membered transition state, rate constants of 1.6 • 105 1.6 • 106 and 4 • 106 s -1 respectively, are recommended (Atkinson, 1997a, 1997b). These are corrected for neighboring group effects using F ( - C H 3 ) = 1.0, F ( - C H 2 - ) = F ( > C H - ) = F ( > C < ) = 1.27, and F ( - O H ) = 4.3. 9 Reaction with NO or NO 2"

-}-02 ---> CH3(CH2)3CHO + H O 2

Pentanal

CH 3CH CH 2CH 2CH 20 H MIO2 OO"

NO

I

k..r

NO 2

CH 3CH CH 2CH 2CH 20 H

O ~

>CH3CHCH2CH2CH2OH

[ OH

(31)

Table 6.7 compares the estimated rates of reaction with 02, decomposition, and isomerization for some alkoxy radicals with different structures. It is important to recognize that there is a great deal of uncertainty in many of these estimates, and this is an area that clearly requires more research. However, given these caveats, it is clear that where isomerization is possible, it usually predominates at room temperature. (Note, however, that this will be slower at the lower temperatures found at higher altitudes.) When isomerization is not feasible, e.g., for the smaller alkoxy radicals or for branched species, reaction with 0 2 is always significant and usually predominates. As the size of the alkane increases, so does the relative importance of isomerization. For example, consider the competition between isomerization and reaction with 0 2 for the n-pentoxy radical:

OH

CH3CH2CH2CH2CH20" [ Isomerization 1,6-H shift

(30)

The latter reactions with NO and NO 2 are less likely than the first three possibilities. With k30 k31 ~ 3 • 10 -1~ cm 3 molecule -~ s -~ at 298 K and 1 atm pressure, the pseudo-first-order rate of reaction, k30[NO] or k31[NO2] , is ~ 75 s -1 at 100 ppb of NO or NO 2. For a species such as (CH3)3CO" that can only decompose (and then only relatively slowly), these reactions could be responsible for ~ 10% of the alkoxy radical reaction at the relatively high NO and NO 2 concentrations of 100 ppb.

--+

CH3mCmCH2mCH2mCHCH3

189

HO2~j

CH3CHCH2CH2CHO < 4-Hydroxypentanal

02

Isomerization

OH

CH3CHCH2CH2CHOH

190

6. GAS-PHASE REACTIONS IN IRRADIATED ORGANIC-NOx-AIR MIXTURES T A B L E 6.7

Rates

RO

CH30" C2H50" n-C4H90.

(S- 1) of

A l k o x y Radical R e a c t i o n s at 2 9 8 K and 1 atm A i r a

Decomposition

Reaction with 0 2 b

Isomerization

5.3 x 10-2 0.3 5.8 x 102

1 x 104 5 x 104 5 x 104

2.0 x 105

2.3 • 104

4 x 104

H

I CH3~C~CH2CH

3

I

O

1 • 103

(CH3)3CO" H

I CH3 ~ C m C H 2 C H 2 C H

3

1.7 x 104

4 x 104

2 x 105

3

1.6 • 104

4 • 104

2.8 • 104

4 • 104

2•

3.4 x 104

4 x 104

2 x 105

1.1 X 10 .3

3.8 X 106

2.0 • 105

9.8 • 103

2.4 • 104

_>7 • 104

I

O

H

I

CH3CH 2 m C ~ C H 2 C H

I

O H

I

CH 3 ~ C ~ ( C H 2 ) B C H

3

106

I

O

H

I

CH3CH 2 m C m ( C H 2 ) 2 C H 3

I

O

(CH3)3COCH20" CH 3

I

CH 3 ~ C ~ C H 2 0 .

I CH 3 a Adapted from Atkinson (1994, 1997a, 1997b) and Atkinson et al. (1995b). b Shown as k[O 2 ] ( s - l ) , based on recommended rate constants for RO + 0 2 (Atkinson, 1997a).

The last step, forming 4-hydroxypentanal, which overall is net hydrogen abstraction from an c~-hydroxy radical to form a carbonyl compound and HO2, is characteristic of this type of radical; it is discussed later in more detail with respect to the reaction of the smallest c~-hydroxy radical, CH2OH, with 0 2 to form HCHO + H O 2.

The relative importance of the two pathways is reflected in the ratio of the yields of the hydroxycarbowls compared to the carbonyl products with the same number of carbons as the parent compound as well as in the absolute yields of the latter products. Thus the carbonyl product yields decrease from 70% for the n-butane reaction to 0.7% for the n-octane reaction; at the same time, the ratio of hydroxycarbonyl products to carbonyls increases from 0.14 to 50 (Kwok

et al., 1996a). In short, isomerization becomes increas-

ingly more important for the larger alkoxy radicals. A study by Eberhard et al. (1995) illustrates the differences in the importance of the various paths that are associated with what may appear at first glance to be small changes in structure. They assessed the relative importance of the various reaction paths for the 2-hexoxy and 3-hexoxy radicals expected from the abstraction of a hydrogen in either the 2- or 3-position, respectively (e.g., by OH or NO3), followed by addition of 0 2 and reaction with NO. Based on the final product yields, they estimated the branching ratios for each possible reaction path. Figure 6.1 shows the possible reactions of each of these alkoxy radicals and the branching ratios they calculate, which are in reasonable agreement with the estimates shown in Table 6.7.

191

E. REACTIONS OF ALKENES (INCLUDING BIOGENICS)

C-C-C-C-C-C O, |

0210003 Io

10

10.99710

C-C-C-C-C-C C+C-C-C-C-CHO C-C-C-C+C-CHOC-C-C-C-(~-CC-C-C-(~-C-C ii

i

i

OH

o

OH

C-C-C-C-C-C i

O.

02

,

~

0.02

C-C-C-C-C-C C-C+C-C-C-CHO C-C-C+C-C-CHOC-C-C-C-C-~; C-C-C-C-I~-C ii

i

i

0 OH OH FIGURE 6.1 Possible fates of the 2-hexoxy and 3-hexoxy free radicals (adapted from Eberhard et al., 1995).

It should be noted that some alkoxy radicals derived from oxygenated compounds may have some unique chemistry in addition to these reactions. For example, Tuazon et al. (1998b) have shown that alkoxy radicals derived from O H + acetate reactions, i.e., R1C(O)OCH(O')R2, rearrange and decompose to R C O and the acid RC(O)OH.

4. Summary of R, R O 2, and RO Radical Reactions in the Troposphere In summary, alkyl radicals, R, are all converted to alkylperoxy radicals, RO2, in the troposphere. R O 2 reacts with NO, except in very remote regions where the concentration of NO is of the order of ~ 40 ppt or less. The reaction converts NO to NO 2 and hence acts as a source of 0 3 via the NO 2 photolysis. For the larger alkylperoxy radicals, a substantial portion of the R O 2 + NO reaction also gives the stable alkyl nitrate, R O N O 2. As discussed in more detail in Section J.2.b, in remote regions where the NO concentrations are very low, reactions of R O 2 with other peroxy radicals such

as H O 2 and R O 2 occur. Finally, under very high concentrations of NO 2, R O 2 may react with NO 2. The alkoxy radical, RO, generated by R O 2 reactions can react with 02, decompose, or undergo an isomerization, depending on the structure of the radical. At very high NO X concentrations for a few alkoxy radicals, reactions with NO and NO 2 may also occur to a small extent. Figure 6.2 summarizes these possibilities.

E. REACTIONS OF ALKENES (INCLUDING BIOGENICS) 1. Hydroxyl Radical (OH) As for other organics in the atmosphere, the O H radical is a major oxidant for alkenes. Table 6.8 gives the rate constants for some O H - a l k e n e reactions as well as their temperature dependence in Arrhenius form. Several points are noteworthy: (1) the reactions are very fast, approaching 10 -l~ cm 3 molecule -1 s -1 for the larger alkenes; (2) the rate constants have a pressure dependence; (3) the apparent Arrhenius activation energies are "negative."

192

6. GAS-PHASE REACTIONS IN IRRADIATED O R G A N I C - N O x - A I R MIXTURES RCH 3 + OH

9

RCH2+ H20 0 ,M

RCH2ONO2 9 minor

NO

RCH202 R'O

RCH20+NO2 9 major

I

HO2 ~ RCH2OOH+ 02

Products O21 ~ RCH

Decomposition Isomerization FIGURE

TABLE 6.8

6,2

Summary of alkane oxidation by O H in air.

R a t e C o n s t a n t s a n d T e m p e r a t u r e D e p e n d e n c e a for t h e R e a c t i o n s of O H R a d i c a l s w i t h A l k e n e s e at 1 a t m T o t a l P r e s s u r e of A i r b

Alkene Ethene Propene 1-Butene cis-2-Butene trans-2-Butene 2-Methylpropene 1-Pentene cis-2-Pentene trans-2-Pentene Cyclopentene 3-Methyl-l-butene 2-Methyl-l-butene 2-Methyl-2-butene 1-Hexene Cyclohexene 1-Heptene trans-2-Heptene Cycloheptene 1,3-Butadiene 2-Methyl-l,3-butadiene (isoprene) Camphene 2-Carene Limonene cr-Phellandrene /3-Phellandrene cr-Pinene /3-Pinene cr-Terpinene y-Terpinene Terpinolene Methyl vinyl ketone Methacrolein

k c (10 -12 c m 3 molecule - 1 s - 1) 8.52 26.3 31.4 56.4 64.0 51.4 31.4 65 67 67 31.8 61 86.9 37 67.7 40 68 74 66.6 101 53 80 171 313 168 53.7 78.9 363 177 225 18.8 d 33.5 d

A (10 -12 c m 3

molecule - 1 s - 1)

E a/R (K)

1.96 4.85 6.55 11.0 10.1 9.47

-438 - 504 -467 - 487 - 550 - 504

5.32

- 533

19.2

- 450

14.8 25.4

- 448 - 410

12.1 23.8

-444 - 357

a k = A e - E a / R v ; valid only for the 250-425 K range. b F r o m Atkinson (1997a). c High-pressure limiting rate constants (k~) except for C z H 4 and C 3 H 6. d F r o m Atkinson (1994). e See Fig. 6.22 for structures of biogenics.

193


These observations are consistent with the primary reaction path being addition of OH to the double bond to form an adduct or intermediate that can decompose back to the reactants, or be stabilized:

Because the /3-hydroxyalkyl radicals formed are substituted alkyl radicals, they react with 0 2 to form alkylperoxy radicals, e.g., OO"

OH + / C

I o. (1 o.c (1 o.I C~.

\cl /

/C

C

( 3 2 , - 32)

(33)

CH3CH~CH2OH

I

+ 0 2

CH3CH~CH2OH. (35)

These /3-hydroxyalkyperoxy radicals undergo the same reaction discussed earlier for RO 2 radicals, i.e., reaction primarily with NO,

I OO"

This mechanism has been confirmed by the mass spectrometric observation of the OH-alkene adducts themselves (e.g., Morris et al., 1971; Hoyermann and Sievert, 1983). Only at low pressures for the smaller alkenes is decomposition back to reactants significant. However, this does illustrate the importance of understanding the fundamental mechanisms in order to extrapolate to atmospheric conditions reliably. A number of experimental techniques used for studying gasphase kinetics and mechanisms require low pressures and, under these conditions, decomposition of the OH-alkene adduct can predominate. As long as the fundamental mechanisms are understood and the kinetics determined as a function of pressure, extrapolation to atmospheric conditions is possible. Clearly, confirmation using studies at atmospheric pressure is also important. Hydrogen atom abstraction can occur to a small extent, particularly with larger and more highly branched compounds. However, the contribution of this path is, overall, relatively small. For example, for the reaction with 3-methyl-l-butene, where there is a weaker allylic C - H bond, ~ 5 - 1 0 % of the reaction proceeds by abstraction at 1 atm in air (Atkinson et al., 1998). For all but the two smallest alkenes, ethene and propene, the rate constants are at their high-pressure limits at 1 atm, and even for these two compounds, the effective rate constant is within ~ 10% of ks. In the case of unsymmetrical alkenes, the OH radical can add to either end of the double bond. There is evidence that, as expected, it preferentially adds to form the secondary radical. For example, for the propene reaction (Cvetanovic, 1976), ~ 6 5 % of the adducts formed correspond to (34a) and 35% to (34b):

C H 3 C H m C H 2 O H + NO O"

I

C H 3 C H m C H z O H + NO 2

(36a)

ONO 2

I

CH3CH~CHzOH,

(36b)

with rate constants that are about the same as those for simple alkylperoxy radicals. The hydroxy nitrate yields for some simple C4-C 6 alkenes are ~ 2 - 6 % , about half those for simple alkyl radicals (Table 6.5) (e.g., O'Brien et al., 1998; Chen et al., 1998). At low concentrations of NO, these RO 2 radicals also react with HO 2 and other RO 2 radicals (e.g., see Hatakeyama et al., 1995; and Tuazon et al., 1998a). There is increasing evidence that for some reactions, the RO 2 + NO reaction produces a fraction of the alkoxy radicals with sufficient energy that they decompose immediately. For example, Orlando et al. (1998) observed that in the reaction of OH with C 2 H 4 , approximately 25% of the HOCH 2CH 20 radicals generated in the reaction of H O C H z C H z O O with NO decomposed before they could be collisionally stabilized. Similar observations have been made for RO 2 from the reactions of alternate CFCs (see Chapter 13). The /3-hydroxyalkoxy radicals such as that formed in (36a) can, in principle, react as discussed earlier, i.e., with O2, decompose, or, for the larger radicals, undergo intramolecular isomerization. For the small alkenes (< C4), decomposition appears to be the most important process (e.g., Tuazon et al., 1998a). For example, for the /3-hydroxyalkoxy radical formed in (36a), the primary fate is O"

OH + C H 3 C H : C H 2

C H 3 C H m C H 2 O H , (34a) C H 3 C H ( O H ) ~ CH 2 . (34b)

C H 3 C H ~ C H 2 0 H ---o CH3CHO + ~ H z O H .

(37)

In the case of the stabilized alkoxy radical formed in the ethene reaction, decomposition and reaction with

194


0 2 are both important at 1 atm in air (Niki et al., 1981; Orlando et al., 1998): CH2(OH)CH20 + 0 2 ~

characteristic manner with 0 2" 9CH2OH

CH2(OH)CHO + HO 2,

CH2(OH)CH2 6 ~ HCHO + CH 2OH.

+ 0 2 ---> HCHO

+ HO2,

k = 9.4 • 10-12 c m 3 molecule-1 s-1

(38)

(Atkinson et al., 1997a).

(39)

The larger c~-hydroxy radicals react in a similar manner:

For the /3-hydroxyalkoxy radicals formed from alkenes C 5 and larger, there is experimental evidence that isomerization starts to dominate (e.g., see Atkinson et al., 1995d; Kwok et al., 1996b). Thus, isomerization followed by reaction with 02, NO, etc., ultimately leads to the formation of dihydroxycarbonyl compounds. For the reaction of OH with 1-butene, for example, isomerization of one of the alkoxy radicals ultimately leads to 3,4-dihydroxybutanal in competition with its decomposition and reaction with 02:

CR1R2OH + 0 2 ---> R 1 R 2 C ~ O + HO 2.

(40)

(41)

Interestingly, the mechanism appears to involve the initial addition of O 2 to the carbon radical, followed by isomerization and decomposition: 9CHzOH + O 2 ---> . O O C H z O H HOOCH20--+ HO 2 + HCHO.

(40)

OH 02

C H 3 C H 2 C H ~ C H 2 + OH - ~

CH3CH2CHmCH2OO"

IC

OH

NO2

OH

I "CH2CH2CH~ CH2OH

Isomerization < 1,6-H shift

CH3CH2CH~CH20" 07

M~O2

~Decomposition

OH

I

" O O C H 2 C H 2 C H ~ CH2OH

IC

NO2

OH

I

02 ~ , ~ HO2

"OCH2CH2CH m C H 2 O H

U ~ >O / / C

m CH 2CH(O H)CH 20 H 3,4-Dihydroxybutanal

The yields of such dihydroxycarbonyl products have been measured to increase continuously from 0.04 for 1-butene to 0.6 for the 1-octene reaction (Kwok et al., 1996b). The low yield for 1-butene reflects the fact that only one of the two possible alkoxy radicals formed can undergo isomerization via a 6-membered transition state. The carbon-centered hydroxy-containing radical CH2OH formed in reactions (37) and (39) is encountered frequently as an intermediate in tropospheric organic oxidations. It is commonly referred to as an c~-hydroxy radical. These types of radicals react in a

For example, no deuterium isotope effect is observed on the rate constant when CH2OD is substituted for CH2OH (Grotheer et al., 1988; Pagsberg et al., 1989), as would be expected if a direct hydrogen abstraction were occurring. Similar principles apply to more complex alkenes. One such case is that of isoprene, an important biogenically produced hydrocarbon:

Isoprene Figure 6.3 shows some of the major pathways in the

H CH,

I

I

H-C-;-CH=CH,

-66%

I

cp

]

,.

I

cp

OH

HOCH,-C-CH=CH, ON02 I

H CH3 I I ~-&PC=CH-CH

tOH

OH

(and CH,-C-CHXH,) I OH

3 or 4 Addition

CH H

I

I

major

?

l3 I

4

0,

l3 I

H ,=C-C-CH,OO*

I

I

OH

OH

-

CH3 H

I

NO

HOCH,-C=CHCH,ONO,

I

major

?

HOCH,-C=CH-CH,O*tNO,

I

v

I

OCHC=CHCH

,OH

CH CH,=C-C

methacrolein CH,OO*

I

C-C=CHCH,OH H I

HOCH,X=CHCH,OH

JNO

t

CH,OH

*O HCHOtHO,

I HOCH24=CHCH20HtN02 lsomerization

-

/ O\

CH,OH

I

HOCH,-CXHCH

t

I

0C-H

I CH3

CH,O*

0,

H

major

HC,

l 3 /H

0

I

c/-HzO

Decomposition

Methyl vinyl ketone

HOCH,X=CH$HOH

c

I

I

00.

I

CH,=C-C-CH,OHtNO,

lsomerization CH,. I HOCH,-C=CH-CH,OH

I

I

*C --OH

HOW-C=CH-CH,OO*

CH,OH

I

/ O\

y 3

Decomposition *CH,OHtCH,fCH=CH,

CH H

CH H (and CH,=C-C-CH,~)

H,C

CH3 H

HOCH,-C-CH=CH,tNO,

I

CH,=C--G-CH,OH

cp

minor

cp

HCHOtHO,

-

HO,

II

0

FIGURE 6.3 Some major pathways in the reaction of OH with isoprene (adapted from Carter and Atkinson, 1996).

II-

0,

nt.0,

'

CH3 3-methylfuran

196


OH-isoprene reaction. Detailed mechanisms have been developed by several groups (e.g., Paulson and Seinfeld, 1992b; Carter and Atkinson, 1996; Zimmermann and Poppe, 1996). Addition to the 1- (or 2-) position gives methyl vinyl ketone, whereas OH addition to the 3- (or 4-) position yields methacrolein. (Note that the 0 3 reaction also generates these products, e.g., Aschmann and Atkinson, 1994.) Small yields of 3-methylfuran have also been reported (e.g., Atkinson et al., 1989; Tuazon and Atkinson, 1990; Paulson et al., 1992a). However, together these account for only about 60% of the total reaction (Tuazon and Atkinson, 1990; Paulson et al., 1992a; Carter and Atkinson, 1996). There is experimental evidence for another path in which isomerization of the double bond to the 2,3-position occurs after the initial addition of the OH; subsequent addition of 0 2 and the usual reactions of RO 2 and RO give the hydroxycarbonyls C H O C ( C H 3 ) = C H C H 2 O H and HOCH 2C(CH 3) = CHCHO (Kwok et al., 1995). As shown in Fig. 6.3 and as expected by analogy to the chemistry of simple alkenes, organic nitrates are formed in small yields from the reaction of the peroxy radicals with NO 2. For example, Chen et al. (1998) report the formation of seven different organic nitrates in this system, with a total yield of 4.4%. This formation of organic nitrates may be important as a minor NO X sink in forested regions, however. For example, Starn et al. (1998a,b)estimated that the isoprene nitrates comprise 0.5-2% of the NOy during the day but up to 9% at night in a rural area in the state of Tennessee. 2. O z o n e ( O 3 ) The room temperature rate constants for the reactions of 0 3 with some alkenes are given in Table 6.9. While the values are many orders of magnitude smaller than those for the corresponding OH reactions, the fact that tropospheric ozone concentrations are so much larger makes these reactions a significant removal process for the alkenes. Despite many decades of research, the mechanisms of ozone-alkene reactions in the gas phase are still not well understood, certainly not as well understood as the corresponding reactions in the condensed phases. Figure 6.4 compares the overall characteristics of these reactions in the gas and condensed phases. It is well known that the initial step in the reaction is the addition of 0 3 across the double bond to form what is known as a primary ozonide, or rnolozonide:

R1NC / R3 0 3+ / mCN R2

R4

O R1)]C R2

J

O O C~I/R3

Primary

R4

ozonide

(42)

This primary ozonide is not stable. One of the two peroxy O - O bonds (marked a or b in Eq. 43) and the C - C bond (marked c) cleave simultaneously to give an aldehyde or ketone and an intermediate called the Criegee intermediate, named after the German chemist who originally proposed this mechanism (e.g., see Criegee, 1975): a

(+)

b R1R2C:O

Rl\l /C

R2

~

I/ R3

a/

+ R3R4COO" (-)

Criegee intermediate

C~

c

R4 b ~

(+) R3R4CmO

+ R1R2COO" (-)

Criegee intermediate

(43)

T A B L E 6.9 Rate C o n s t a n t s and T e m p e r a t u r e D e p e n d e n c e a for the G a s - P h a s e Reactions of 0 3 with Some A l k e n e s b

Alkene

Ethene Propene 1-Butene 2-Methylpropene cis-2-Butene trans-2-Butene

1-Pentene Cyclopentene 2-Methyl-2-butene 1-Hexene Cyclohexene cis-3-Methyl-2-pentene trans-3-Methyl-2-pentene

2,3-Dimethyl-2-butene 1,3-Butadiene 2-Methyl-l,3-butadiene Myrcene 2-Carene 3-Carene Limonene c~-Phellandrene /3-Phellandrene c~-Pinene /3-Pinene c~-Terpinene y-Terpinene Terpinolene Methyl vinyl ketone Methacrolein

k (10 - i s cm 3 A (10 -15 cm 3 E a / R molecule - 1 s - 1) molecule - 1 s - l) (K) 1.6 10.1 9.64 11.3 125 190 10.0 570 403 11.0 81.4 450 560 1130 6.3 12.8 470 230 37 200 2980 47 86.6 15 2.1 • 104 140 1880 5.6 c 1.2 c

9.14 5.51 3.36 2.70 3.22 6.64

2580 1878 1744 1632 968 1059

1.8 6.51

350 829

2.88

1063

3.03 13.4 7.86

294 2283 1913

1.01

732

a k = Ae -Ea/Rr

b From Atkinson (1997a) and Atkinson et al. (1997a); for structures of biogenics, see Fig. 6.22. c Average of Grosjean and Grosjean (1998a) and Neeb et al. (1998b).

197

E. REACTIONSOF ALKENES(INCLUDINGBIOGENICS)

Decomposition

M

~/ R3

l!

a

F \ / 0 3 + C=C / N R2 H

9 L

b

~03~

R

q

C=O

1CH2\ +

Criegee

R2/ ~ ' \ H_ _ l ' ~ Primaryozonide or molozonide

Solution or Intermediate / c~ (CI) J phase

\

R1CH2\ . J f C-CO, / I R2 Ester StabilizedCI channel

>/

Other products

II R1CH2COR

R1CH2 \ C/ R2/

C H 4 + C O 2 , 0

, CH3OH + CO.

(46e) (46f)

Table 6.10 gives the ranges of observed yields of the stabilized Criegee intermediates at 1 atm pressure in air and at room temperature. Clearly, significant decomposition of the intermediates occurs under typical tropospheric conditions. The fraction of the primary ozonide decomposition that goes by path a in reaction (43)vs path b for unsymmetrical alkenes has been determined from the product yields in a number of studies (e.g., see series of papers by Grosjean and co-workers and Atkinson and co-workers). One might expect, based on thermochemical arguments, that the decomposition giving the more stable, i.e., more highly substituted, biradical would be preferred, and this is indeed what has been observed (e.g., see Horie and Moortgat, 1991; Grosjean et al., 1994c; Atkinson et al., 1995c; and Grosjean and Grosjean, 1997, 1998a). Atkinson (1997a) recommends, based on the literature, the following branching ratios: O 3 + R 1CH = CH 2

0.5>

(R 1CHOO. )* + HCHO, o.5 R1CHO + ( H C H O O . ) * ,

TABLE 6.10 Yieldsof Stabilized Criegee Intermediates at Room Temperature and 1 atm Aira Alkene Ethene Propene 2-Methylpropene cis-2-Butene trans-2-Butene

2,3-Dimethyl-2-butene 1-Octene Cyclopentene Cyclohexene Cycloheptene 1-Methylcyclohexene Methylenecyclohexane c~-Pinene /3-Pinene

Yield 0.35-0.47 0.25-0.44 0.17 0.18 0.19-0.42 0.30 0.22 0.05 0.03 0.03 0.10 0.22 0.13 0.25

a From Hatakeyama and Akimoto (1994) and Atkinson (1994, 1997a); see these for original references.

0 3 + R1R2C=CH

2

0.65

(R1R2COO~

q-

HCHO,

0.35 , R1C(O)R 2

+ (H(;HO0.)*, 0 3 + R 1 R 2 C = C H R 3 0.65 (R1R2~OO.) , +R3CHO, o.35 R1C(O)R2 + (R3CHO0.)*. While free radicals were reported in ozone-alkene reactions as long ago as 1958 by Saltzman, the importance and magnitude of these processes have only relatively recently been appreciated by the atmospheric chemistry community. Of greatest significance is the production of OH in these reactions, which was first observed in these systems in a vibrationally excited state, OH ~, probably from the initial production of H atoms, followed by the well-known reaction H + 0 3 --+ OH ~ + 0 2 (Finlayson et al., 1972). Rapid secondary reactions of OH with the parent alkene complicates interpretation of the original O3-alkene mechanism and products. In addition, in kinetic studies where the loss of the alkene is followed, the derived rate constants may be too large unless the contribution of OH secondary reactions is taken into account or it is scavenged from the system. However, more important than the implications for laboratory studies are those for the role of ozone-alkene reactions in tropospheric chemistry. Thus, in addition to producing carbonyl compounds that can photolyze to generate free radicals, ozone-alkene reactions contribute directly to the generation of free radicals via reactions such as (45) and (46). This will be especially important at night when there are no photolytic sources of OH. Table 6.11 gives some measured yields of OH in some ozone-alkene reactions. These have been determined by trapping the OH, e.g., using a large excess of cyclohexane (e.g., Grosjean et al., 1994c; Grosjean and Grosjean, 1997) or 2-butanol (e.g., Chew and Atkinson, 1996), and measuring the major "OH-trap" products such as cyclohexanone and cyclohexanol (Atkinson et al., 1992; Aschmann et al., 1997a) and butanone, respectively (e.g., Chew and Atkinson, 1996), under conditions where the formation of OH via the HO 2 + 0 3 reaction is minimized. Other approaches include trapping by CO to form CO 2 (e.g., Gutbrod et al., 1997b) or following the loss of pairs of tracer organics that react with OH (e.g., Schiller et al., 1997; Paulson et al., 1997; Marston et al., 1998). In addition to the indirect measurements of OH, a direct measurement of its production using laser-induced fluorescence has been carried out for the reactions of 0 3 with ethene, trans-2-butene, 2,3-dimethyl-2-butene, and isoprene at 4- to 6-Torr

199


pressures in N 2. Yields of OH from 0.4 to 0.8 were observed by Donahue et al. (1998b). In short, O3-alkene reactions generate OH, with yields approaching unity in the case of highly branched alkenes. The mechanism of decomposition of the Criegee intermediates is believed to occur via several reaction channels shown for the [(R1CH2)(R2)CHOO]* Criegee intermediate in Fig. 6.4. The oxygen-atom elimination channel for simple alkenes is not believed to be important. However, the ester and hydroperoxide channels are important and explain the production of free radicals such as OH. Theoretical calculations have shed some light on this (e.g., Gutbrod et al., 1996, 1997a;

Anglada et al., 1996). Figure 6.5, for example, shows the calculations of Gutbrod et al. (1996) for the reaction of the Criegee intermediate C(CH3)2OO, formed in the reaction of 0 3 with 2,3-dimethyl-2-butene. The path that forms the dioxinane intermediate has a higher activation barrier than that forming the unsaturated hydroperoxide, and the latter channel that leads to OH formation is predicted to predominate. The key to OH production is predicted to be an alkyl group in a syn position so that it can interact with the terminal oxygen (Gutbrod et al., 1996, 1997a): H2 C/H

\

/ T A B L E 6.11

Yields of O H from G a s - P h a s e O 3 - A l k e n e Reactions at 1 atm P r e s s u r e a

Alkene Ethene Propene 1-Butene 1-Pentene 1-Hexene 1-Heptene 1-Octene cis-2-Butene trans-2-Butene

Cyclopentene Cyclohexene 1-Methylcyclohexene 2-Methylpropene 2- Methyl- 1-butene 2-Methyl-2-butene 2,3-dimethyl-2-butene Limonene Myrcene c~-Pinene /3-Pinene Terpinolene Camphene 1,3-Butadiene Isoprene

OH yield 0.12, b 0.08 h 0.33, a 0.18 h 0.41 a 0.377 0.327 0.27 f 0.18f-0.45 c 0.41, a 0.17 h 0.64, a 0.24 h 0.617 0.68 a 0.90 f 0.84a 0.83 a 0.89 a 0.5 h- 1.0 a'd 0.86 b 1.15 b 0 . 7 0 - 0 . 8 5 b'g'i

0.35 b 1.03 b _ R1CHO + SO 3.

(49)

However, as discussed in detail by Hatakeyama and Akimoto (1994), there is increasing evidence for the formation of an addition complex such as that proposed by Martinez and Herron (1981) and observed earlier by mass spectrometry by Schulten and Schurath (1975): RI~

R~(2HOO. + SO 2 ~

/

OmO

C H/

\

/,SmO

(50)

~O

In the presence of water vapor, the adduct then reacts with H 2 0 to generate sulfuric acid. Other reactions proposed for the Criegee intermediate include R 1 C H O O . + NO ---> R1CHO + NO2,

(51)

R1CHOO.

+ NO 2 --> R1CHO + NO 3,

(52)

R1CHOO.

+CO~R1CHO+CO

(53)

2.

However, the relative importance of these reactions is highly uncertain. Table 6.12 gives one estimate of the range of possible rate constants for these reactions of the Criegee intermediate and the calculated lifetimes of the intermediate under some typical atmospheric conditions. Reaction with water vapor is expected to be the major fate of the Criegee intermediate, with per-

201

TABLE 6.12 Range of Reported Rate Constants for the Reactions of the Criegee Intermediate with Some Gasesa and Associated Lifetimes of the Criegee Intermediate under Polluted Tropospheric Conditions Reactant

H20

CO HCHO NO NO2 SO2

k (cm 3 molecule - 1 s - l)

'I'CIb (S)

1 • 10-15-2 • 10-19 1 • 10-14 2 • 10-12-2 • 10-16 7 • 10-12 7 • 10-13-1 • 10-17 1.7 • 10-11-3 • 10-15

0.003-400 8 20-(2 • 105) 0.1 1_105 0.05-270

a Adapted from Hatakeyama and Akimoto (1994). b Lifetime of Criegee intermediate assuming the following concentrations: H20, 4 • 1017 molecules cm-3, corresponding to a relative humidity of ~ 50% at 298 K; CO, 0.5 ppm; HCHO, 1 ppb; NO, NO2, SO2, 50 ppb.

haps some contribution by reaction with NO a n d / o r SO 2 in more polluted atmospheres. It should be noted that while the mechanism outlined in this section describes the overall features of O3-alkene chemistry, there are also other minor paths as well. For example, small yields of epoxides that appear to be formed in the primary reaction have been observed as products of the reactions of some dienes and cycloalkenes (e.g., see Paulson et al., 1992b; and Atkinson et al., 1994a, 1994b). The reader should consult the rather extensive ozone literature for further details on both the condensed- and gas-phase reactions. As discussed in Chapter 9, there are a variety of low-volatility organics, some of which are unsaturated, which are found in or on particles in the atmosphere, and these would be expected to undergo heterogeneous reactions with O 3. Consistent with this, de Gouw and Lovejoy (1998) report reaction probabilities of 10-2-10 .3 for O 3 on organic liquids containing unsaturated sites.

3. Nitrate Radical ( N O 3 ) Table 6.13 gives the rate constants and temperature dependence for the reactions of the nitrate radical, NOB, with some alkenes. The room temperature rate constants span about six orders of magnitude, approaching diffusion controlled for some of the biogenics. Those activation energies that have been measured are relatively small (or negative by a small amount), and the rate constants increase with increasing alkyl substitution at the double bond. This, plus the much larger values compared to the rate constants for reactions with alkanes where H-abstraction occurs, suggests

202


R o o m T e m p e r a t u r e R a t e C o n s t a n t s and T e m p e r a t u r e D e p e n d e n c e a for the G a s - P h a s e R e a c t i o n s of the N O 3 Radical w i t h Some A l k e n e s b k (cm 3 molecule-1 s - 1) at 298 K

Alkene

Ethene Propene 1-Butene 2-Methylpropene cis-2-Butene trans-2-Butene

2-Methyl-2-butene 2,3-Dimethyl-2-butene 1,3-Butadiene 2-Methyl-l,3-butadiene (isoprene) Cyclopentene Cyclohexene Cycloheptene Camphene 2-Carene 3-Carene Limonene a-Pinene /3-Pinene a-Phellandrene /3-Phellandrene a-Terpinene y-Terpinene Terpinolene Methyl vinyl ketone Methacrolein

2.1 • 9.5 • 1.4 • 3.3 • 3.5 • 3.9 • 9.4 • 5.7 • 1.0 • 6.8 • 5.3 • 5.9 • 4.8 • 6.2 • 1.9 • 9.1 • 1.2 • 5.9 • 2.1 • 7.3 • 8.0 • 1.4 • 2.9 • 9.7 • < 6 • 3.3 X

10 -16 10 -15 10 -14 10 -13 10-13 10 - 13 10 -12 10-11 10-13 10 -13 10-13 10 -13 10-13 10-13f 10-11 10 -12 10-11 10-12f 10-12f 10-11 10-12 10-10 10-11 10-11 10-16 d

A (cm 3 molecule - 1 s - 1)

Ea /R (K)

c 4.6 • 10-13 3.14 • 10-13

c 1156 938

d

d

3.03 • 10-12

446

1.05 • 10-12

174

3.1 • 10-12

481

3.5 • 10-13 1.6 • 10-10

- 841 1248

10-- 15 e

a k = Ae-Ea/Rr

b From Atkinson (1994) and Atkinson et al. (1997a). c Recommended temperature dependence is k = 4.88 • l O - l S T 2 e - 2 2 s 2 / ~ cm 3 molecule -1 s -1 from 290 to 523 K. d Dlugokencky and Howard (1989) report a curved Arrhenius plot that can be matched from 204 to 378 K by k = (1.78 _+ 0.36) • 10 -12 e x p [ ( - 5 3 0 _+ 100)/T] + (1.28 _+ 0.26) • 10 -14 exp[(570 _+ l l 0 ) / T ] ; Atkinson (1997a) recommends k = 1.22 • 10-1Se 3s2/~ cm 3 molecule-1 s-1 from 204 to 378 K. e From Chew et al. (1998). f From Martinez et al. (1998).

that NO B undergoes electrophilic addition to the double bond in much the same fashion as OH: ONO 2 J

NO B + / C ~ C ~ - - ~

I

(54)

I However, there is no pressure dependence of the rate constants over the range from about 1 Torr up to 1 atm, suggesting that the adduct does not decompose significantly back to reactants under atmospheric conditions. This mechanism is consistent with the observation of significant yields of epoxide products and NO 2 for some alkenes (Olzmann et al., 1994). For example, Fig. 6.7 shows the infrared spectrum of the minor products from the reaction of NO Bwith 2,3-dimethyl-2-butene at

740 Torr in air, after the spectra of the reactants and the major product, acetone, as well as of NO 2 and H N O 3 have been subtracted out. Bands due to the epoxide (in ~ 17% yield) are clearly seen, as well as bands due to 2,3-dinitroxy-2,3-dimethyl-2-butane. Together with acetone, these compounds account for 75% of the reacted alkene (Skov et al., 1994). Figure 6.8 summarizes the mechanism of the reaction of NO B with 2,3-dimethyl-2-butene (Skov et al., 1994; Olzmann et al., 1994). The addition of the NO B to a double bond is about 20 kcal mol-1 exothermic, so that the initially formed adduct is excited and either can be stabilized or decompose by breaking the weaker C O - N O 2 bond to form the epoxide and NO 2. The stabilized adduct can also decompose in a similar manner, in competition with its reaction with 0 2. Under


1325

ONO,

/ I

700

I

I

1000

i

I

i

i

1300

1600

Wavenumber (cm "1) FIGURE 6.7 Infrared spectrum of products of the reaction of 2,3-dimethyl-2-butene with NO 3 (spectra of acetone, NO2, and HNO 3 have been subtracted out) (adapted from Skov et al., 1994).

same time as NO 3. However, there may still be sufficient amounts for this to be a significant fate of the peroxy radical. Figure 6.9 traces the expected fates of this radical through to stable organic products. The dinitroxybutane is formed from the alkoxy radical-NO 2 reaction and acetone from the alkoxy radical decomposition. The observation of smaller amounts of other aldehydes and formaldehyde also suggests a small contribution from a 1,4-H shift and subsequent reactions of the alkyl radical thus formed (Hjorth et al., 1990). The results of laboratory studies of the mechanism of NO 3 radical reactions with alkenes tend to be slanted toward observing polyfunctional organics containing more than one nitro group. The reason for this is that the thermal decomposition of N20 5, N20 5 ~ NO 3 + NO 2,

surface-level tropospheric conditions, 1 atm total pressure in air, the epoxide yields are small but not insignificant. For example, epoxide yields of 7, 12, 18, and 28% from the reactions of isobutene, trans-2-butene, 1-butene, and propene, respectively, have been reported (Berndt and B6ge, 1995). The lifetime of the excited NO3-alkene adducts is sufficiently long that rotation about the C - C bond leads to the same yields of trans- and cis-epoxides regardless of the configuration of the reactant alkene; for example, the reactions of both cis- and trans-2butene give about 80% of the trans form of the product epoxide and 20% of the cis form (Benter et al., 1994). The peroxy radical formed in Fig. 6.8 is expected to react in air as discussed earlier, i.e., with NO, HO2, RO2, or NO 2. Since NO 3 itself reacts rapidly with NO, large concentrations of NO will not be present at the

C~3

/

C=C

CH3

/ \

cH~*

CH3 + NO3

>

CH3

y

C H 3 \ ? NO2 /9 CH3 C--C CH3 / 02 ~

C H 3 \ / O\ :~ C - - C

\ CH3

CH3

/

/ CH3

\

CH3

u

NO2 CH3

/

C--C

\

CH3

FIGURE 6.8 Mechanism of the NO 3 reaction with 2,3-dimethyl2-butene (adapted from Skov et al., 1994).

203

(55,-55)

is often used as the source of NO 3. As a result, relatively high concentrations of NO 2 may also be present, and depending on the concentration regimes employed, these can trap the alkylperoxy and alkoxy radicals more efficiently than would normally be the case in the atmosphere. On the other hand, sources such as the reaction of fluorine atoms with HNO3, F + H N O 3 ~ HF + NO3,

(56)

are normally used at low pressures ( ~ Torr) and in the absence of air because such conditions are needed to generate the halogen atoms. Under these conditions, the products may be quite different from those at 1 atm in air. For example, the yields of epoxides are much greater, typically approaching 100%, under these low-pressure, l o w - 0 2 conditions (Skov et al., 1994; Berndt and B6ge, 1995). However, it is interesting that vicinal dinitrates have been measured in air with higher concentrations at night, as would be expected if nitrate radical reactions were the source. For example, Schneider et al. (1998) measured the diurnal variation of a series of alkyl nitrates in rural air in Germany and found higher nighttime levels of vicinal dinitrates. They attributed these to the addition of NO 3 to double bonds, formation of the alkylperoxy radical, and then reaction with NO, with the minor channel in the RO 2 + NO reaction leading to a second nitrate group on the adjacent carbon. As seen from the rate constants in Table 6.13, the reactions of biogenic hydrocarbons with NO 3 are quite fast, so this reaction is expected to be a major fate of these biogenics at night and to also contribute to the removal of NO X. The lifetime of isoprene with respect to reaction with NO 3 at 50 ppt, for example, is only about 20 min.. Addition of NO 3 occurs primarily at the 1-position and leads in the absence of NO to the

204

6. GAS-PHASEREACTIONS IN IRRADIATEDORGANIC-NOx-AIR MIXTURES .NO 2 -ONO 2 OO-

I

CH 3 ~ C ~

I

CH 3

M

?NO2 ? ONO2 ~ CH 3 ~ C ~ C ~ C H

I

CH 3

I

3

I

CH 3

C m CH 3

I

CH 3

ao:

ONO 2 O~

I

CH 3 -

I

C--

C - - CH 3 + RO + 0 2

CH 3

CH 3

I

I

Other Products

HO 2

Dinitrate

ONO 2 0 0 H CH 3 -

NO

9

I

I

C--

C - - CH 3 + 0 2

CH 3

CH 3

I

I

ONO 2 0 .

I

CH 3 m C

I

CH 3

I

C - - CH 3 + NO 2

I

CH 3 osition

M

NO

?NO2 ?NO2 CH 3 - - C - -

I

CH 3

C - - CH 3

I

H3 m C~

I

CH 3

CH 3

+ CH 3 CCH3

II

O

CH3 ICICH3 + NO 2 ii

O

FIGURE 6.9 Expected atmospheric fates of peroxy radical formed in NO 3 addition to 2,3-dimethyl-2-butene.

formation of 4-nitroxy-3-methyl-2-butenal as a major product (e.g., Jay and Stieglitz, 1989; Barnes et al., 1990; Skov et al., 1992; Kwok et al., 1996c; Berndt and B6ge, 1997)" O 2 N O ~ ~

methyl-2-butenal decreases with increasing NO concentrations, while the yields of methyl vinyl ketone and methacrolein increase (Berndt and B6ge, 1997). This has been attributed to the decomposition of the alkoxy radicals formed after 1,2-addition, e.g.,

O

4-Nitroxy-3-methyl-2-butenal

Also formed are a variety of hydroxynitrates, nitrooxyhydroperoxides, and hydroxycarbonyls anticipated from the reactions of alkylperoxy and alkoxy radicals formed after the initial addition of NO 3 to the double bond (e.g., Kwok et al., 1996c). The yield of 4-nitroxy-3-

O2NO~--..~

0 O + HCHO Methyl vinyl ketone

+ NO 2

205


TABLE 6.14 Some R e p o r t e d V a l u e s of the R a t e C o n s t a n t s for the R e a c t i o n s of C h l o r i n e A t o m s w i t h Some A l k e n e s at 1 atm T o t a l P r e s s u r e and 2 9 8 K

and after 3,4-addition, respectively:

@

O N O

O

2 ---+

~ O Methacrolein

+ HCHO + NO 2 .

In addition, in the nighttime atmosphere, reaction of the RO 2 radicals with NO 3 may occur as discussed earlier. As a result, the products of the isoprene-NO 3 reaction in the atmosphere will depend on the concentrations of NO, NO3, HO2, and RO 2. Field studies suggest that the nitrate radical reaction can also be a major contributor to isoprene decay at night, as well as contributing to the formation of organic nitrates in air. For example, Starn et al. (1998b) found that when the product of NO 2 and 0 3 (which form NO 3) was high in a forested region in the southeastern United States, isoprene often decayed rapidly at dusk. This reaction of NO 3 with isoprene was estimated to be the major sink for NO 3 under some conditions in this area. The reader is referred to two excellent reviews of NO 3 chemistry by Wayne et al. (1991) and Atkinson (1991) for further details. The plethora of multifunctional products anticipated from nitrate radical reactions with alkenes (Fig. 6.8) may well be in part responsible for what is commonly referred to as "missing NOy." In brief, one can measure total NOy using techniques such as the "master blaster" described in Chapter 11, in which one converts all oxygenated nitrogen-containing species into NO using a A u / C O converter and then measures the total NO. At the same time, one can measure various individual components of NOy (NO, NO2, HNO3, PAN, etc.) separately and determine whether their sum accounts for the observed NOy. As discussed in detail in Chapter 11.A.4a, the sum of the individual compounds is often less than the measured NOy, and this difference, which is not accounted for, is termed "missing NOy." Given the mixture of potential nitrogen-containing organics formed in the NO 3 radical reactions (e.g., Figs. 6.8 and 6.9), as well as in NO 2 reactions with organic free radicals (see earlier), it is perhaps not surprising that in many air masses species other than NOx, HNO3, PAN, and particulate nitrate are present at significant concentrations. It remains a challenge to identify and measure such species in a specific and sensitive manner. 4. C h l o r i n e A t o m s (C1) Table 6.14 gives the rate constants for the reactions of chlorine atoms with some simple alkenes and some biogenic hydrocarbons. As expected, the reactions are

Alkene

k (10 - l ~ cm 3 molecule- 1 s - l )

Ethene

0.99 a 1.2 b 1.2 c

Propene

2.3 a 2.7 b 2.5 d

1-Butene

2.2 a

1,3-Butadiene

4.2 a 5.6 e

Isoprene

4.6 h 5.5 e 2.8 f

(1R)-( + )- c~-Pinene

4.8 g

(1S)-(-)-c~-Pinene

4.6 g

3-Carene

5.6 g

Myrcene

6.6 g

p-Cymene

2.1 g

Methyl vinyl ketone

2.0 g

a Stutz et al. (1998). b Atkinson and Aschmann (1985), corrected to k(C1 + n-butane) = 2.18 • 10 - l ~ cm 3 molecule -1 s -1 c Kaiser and Wallington (1998), based on reported values of k 0 and k~. d Kaiser and Wallington (1996). e Bierbach et al. (1996). f Bedjanian et al. (1998); this is the value of k~. g Finlayson-Pitts et al. (1999). h Ragains and Finlayson-Pitts (1997).

all very fast, approaching the collision-controlled regime. The reaction proceeds primarily by addition to a double bond to form a chlorine-containing alkyl radical, which then adds 0 2 to ultimately generate characteristic chlorine-containing oxygenated products. For example, 1-chloro-3-methyl-3-buten-2-one is formed in the absence of NO from the addition of chlorine atoms to the 4-position of isoprene, followed by secondary reactions of the alkyl radical with 0 2 etc. This product might therefore be expected to be a reaction product in the low-NO X environment of the marine boundary layer (Nordmeyer et al., 1997). A small portion of the reaction proceeds by what appears to be abstraction of the allylic hydrogen atom from the weaker C - H bond (e.g., Lee and Rowland, 1977). In the case of the isoprene reaction, for example, about 15% of the overall reaction at 1 atm pressure and room temperature proceeds by abstraction (Ragains and Finlayson-Pitts, 1997). It should be noted,

206


however, that this net hydrogen abstraction may not be a simple abstraction per se but rather proceed via an a d d i t i o n - e l i m i n a t i o n reaction (e.g., Kaiser and Wallington, 1996b; Ragains and Finlayson-Pitts, 1997).

5. Nitrogen Dioxide

(NO 2)

NO 2 can add to the double bond of conjugated dienes, but the reaction is generally slow, ranging from 1 X 10 -20 c m 3 molecule -1 s -1 for 2,3-dimethyl-2butene to 1.3 • 10 -17 cm 3 molecule -1 s -1 for a-phellandrene (see Atkinson, 1997a, for a review). Addition of NO 2 to one of the double bonds generates an alkyl radical that reacts as expected, i.e., adds 0 2 to form RO 2 etc. This reaction is not likely to be significant under most atmospheric conditions; for example, at 0.1 ppm NO2, the lifetime of a-phellandrene is about 9 h, much longer than its lifetimes with respect to reaction with OH and 0 3 (see Problem 9). However, it may be important in some laboratory systems where high NO~ concentrations are used. For example, it may be responsible for the observed generation of OH in the dark reaction of isoprene and other conjugated dienes (e.g., Atkinson et al., 1984) with NO 2 in the presence of NO in the dark (Tuazon and Atkinson, 1990) (see Problem 10). Harrison et al. (1998) have also proposed that it could be responsible for some of the observed NO to NO 2 conversion during some air pollution episodes at night in London.

F. REACTIONS OF ALKYNES 1. Hydroxyl Radical (OH) The only significant loss of alkynes is reaction with OH, for which a pressure dependence is observed. Table 6.15 gives the high-pressure limiting rate constants for the OH reactions with acetylene, propyne, 1-butyne, and 2-butyne. The reaction of acetylene approaches the high-pressure limit at several thousand Torr (see Problem 5). However, for the larger alkynes, the reactions are essentially at the high-pressure limit at 1 atm (and room temperature).

The magnitude of the rate constants, their observed pressure dependence, and the products of the reactions are consistent with the mechanism involving the initial addition of OH to the triple bond. For example, the O H - l - b u t y n e reaction at 298 K is about a factor of three faster than the reaction with n-butane (see Table 6.2), despite the fact that it has fewer abstractable hydrogens and the = C - - H bond is much stronger than a primary - C - H bond ( ~ 125 vs ~ 100 kcal mol-1). In addition, a pressure dependence is not consistent with a simple hydrogen atom abstraction (see Chapter 5.A.2). The reactions give as major products the corresponding dicarbonyls; i.e., acetylene gives glyoxal [(CHO)2] , propyne gives methylglyoxal [CH3COCHO], and 2-butyne gives biacetyl [(CH3CO) 2] (Schmidt et al., 1985; Hatakeyama et al., 1986). This is consistent with the following reaction sequence: OH HC~CH

OH

Alkyne Acetylene Propyne 1-Butyne 2-Butyne

0.90 5.9 8.0 27.4

a From Atkinson (1994) and Atkinson et al. (1997a).

NO ~

0 2

'

NO 2

OH

I

HC:CHO0"

H

I

>HCmCO 9

II H

]

, H(;(OH)--C--O

o2

, (CHO) 2 + HO 2 (57)

Formic acid is another product, possibly from the rearrangement and decomposition of the intermediate II (Hatakeyama et al., 1986). Siese and Zetzsch (1995) and Bohn and Zetzsch (1998) have studied the O H - C 2 H 2 reaction using FPRF (see Chapter 5.B.3) and observed biexponential decays of OH. They propose that the adduct I has two channels in its reaction with 02, rather than one as shown above, and that one of the two generates OH and glyoxal, a small portion of which is excited and decomposes to HCO. The vinoxy radical has also been observed as an intermediate (Schmidt et al., 1985), suggesting that a portion of the initial adduct isomerizes:

TABLE 6.15 High-Pressure Limiting Rate Constants (ks) for the Reaction of OH Radicals with Alkynes at 298 K a k~(OH) (10 -12 cm 3 molecule- 1 S - 1)

I

+ OH -> H C m C H

OH C 2 H 2 -+-

OH

, [HC=CH]*

O --+ --+ H C - - CH2 Vinoxy radical

(58) The vinoxy radical reacts rapidly with 0 2 with a highpressure limiting rate constant ks = (1.9 _ 0.2) • 10 -13 cm 3 molecule -1 s -1 at P > 400 Torr (Zhu and Johnston, 1995). While the reaction leads to glyoxal

207

G. REACTIONS OF SIMPLE AROMATIC HYDROCARBONS

formation, the production of glyoxal is delayed compared to the disappearance of the vinoxy radical, suggesting that a long-lived adduct with 0 2 is first formed and then subsequently decomposes, in part to glyoxal,

10 =

o CH(O)CH 2

+ 0 2 --->

[CH(O)CH20(5 ] --+ __+ (CHO)2

+

OH,

+ 0 2 --+

OH + HCHO + CO.

424.4 K

t-

(60)

In short, while the overall features of OH-alkyne reactions are understood, more research needs to be done, especially on the alkynes larger than acetylene.

2

O cO

(59)

perhaps along with the regeneration of OH, which Schmidt et al. (1985) also observed in the presence of 0 2. However, Zhu and Johnston (1995) observed the yield of glyoxal from reaction (59) to be only 15%, whereas Hatakeyama et al. (1986)measured glyoxal yields in the O H - C 2 H 2 reaction of ~70%. Thus reaction (59) cannot be the major path for glyoxal formation in the OH + C2H 2 reaction. Another reaction proposed for the vinoxy radical (Gutman and Nelson, 1983) is reaction (60): CH(O)CH 2.

5

-!O

._~

K

1

,,i-,

rr

0.5

0.2

297.9 K I

0

I

Reaction

I 8

I

I 12

I ~.

time (ms)

F I G U R E 6.10 Semilogarithmic plots of OH decays as a function of reaction time in the presence of a great excess of toluene (-~ 5 • 1013 cm- 1) at temperatures from 297.9 to 424.4 K and in -~ 100 Torr argon (adapted from Perry et al., 1977).

straction of a hydrogen atom occurs, in the case of the substituted benzenes from the alkyl side chains, where the C - H bonds are weaker ( ~ 85 kcal mo1-1 vs ~ 110 kcal mo1-1 for C - H bonds in benzene):

G. REACTIONS OF SIMPLE AROMATIC HYDROCARBONS

CH3

CH 2

OH + ] ~ , )

1. Hydroxyl Radical (OH) The reactions of simple aromatic hydrocarbons with OH provide a classic example of how kinetics can be used to elucidate reaction mechanisms. Figure 6.10 shows a semilogarithmic plot of the decay of OH in the presence of a great excess of toluene from 298 to 424 K at ~ 100 Torr total pressure in argon. While one would expect such plots to be linear (Chapter 5.B.1), this is only observed to be the case at temperatures below 325 K and above about 380 K; at the intermediate temperatures, the plots are clearly curved. Even more unusual behavior is observed for the temperature dependence of the rate constant. Figure 6.11 shows these data in Arrhenius form for the reactions of toluene and 1,2,3-trimethylbenzene. At the higher temperatures, the Arrhenius plot is linear with a normal activation energy (i.e., the rate constant increases with increasing temperature). However, as the temperature is lowered, there is a sharp discontinuity in the plot and at lower temperatures the temperature dependence is reversed; i.e., the rate constants decrease with increasing temperature. These results are now known to reflect the occurrence of two, quite different mechanisms over this temperature range. At the higher temperatures, ab-

I 4

338.5 K

+ H20

(61)

5 x i 0 -11

1,2, 3-mrimethylbenzene~ I x I0 -11 _

//

_

/

I

,f

I

i

Addition

/

.=~

0

5

O

E

I

Abstraction

O3

E

I I Toluene

r

t

2 1 x 10 12 _

_

_

_

5x10

13

2.0

I

I 2.4

I

I 2.8

I

I 3.2

1000 / T (K)

F I G U R E 6.11 Arrhenius plots of log k vs 1000/T for the reaction of OH with toluene and 1,2,3-trimethylbenzene: (O) exponential OH decays observed; ( • nonexponential OH decays observed (adapted from Perry et al., 1977).

208


The temperature dependence is that expected for a hydrogen abstraction; i.e., the rate constant increases with temperature (left side of Fig. 6.11). As the temperature is lowered, the abstraction slows down but another reaction, addition of OH to the aromatic ring, takes place: CH 3

OH+~['-

;

OH

(62,-62)

Theoretical calculations support the expectation that the preferred site of initial OH attack is ortho to the methyl group (Andino et al., 1996), but addition to the other positions also occurs. If the O H - a r o m a t i c adduct, which contains ~ 18 kcal mo1-1 excess energy, is not stabilized, it decomposes back to reactants, reaction ( - 6 2 ) . The existence of the adduct in the case of the O H - b e n z e n e reaction has been observed spectroscopically (Fritz et al., 1985; Knispel et al., 1990; Markert and Pagsberg, 1993; Bjergbakke et al., 1996). As expected for such a mechanism, the rate constants at temperatures below ~ 300 K exhibit a pressure dependence at lower pressures. At higher temperatures, the rate of decomposition of the excited adduct back to reactants is higher, so the net contribution of adduct formation to the overall reaction is small compared to H-abstraction. Conversely, at the lower temperatures, the rate constant for H-abstraction is small while, at the same time, the rate of adduct decomposition is lowered. As a result, at the lower temperatures (right side of Fig. 6.11), adduct formation predominates and a "negative" temperature dependence, as well as a dependence on pressure is observed for the overall rate constant. In the intermediate region, both addition and abstraction are occurring at significant rates, leading to the curved OH decay plots in Fig. 6.10 and the discontinuities in the Arrhenius plots of Fig. 6.11. Table 6.16 shows the room temperature rate constants for the reactions of OH with some simple aromatics as well as the branching ratio for abstraction, i.e., the ratio k61/(k61 + k62). Abstraction accounts for less than about 10% of the reaction at room temperature for those alkylbenzenes studied to date. It is noteworthy that the reactions are all quite fast, even that for benzene being within approximately two orders of magnitude of diffusion controlled. The products of the abstraction path are easily predictable, based on our understanding of the fates of alkyl radicals produced in alkane reactions (see Sections C and D). For example, in the case of toluene, the

TABLE 6.16

R a t e C o n s t a n t s at R o o m T e m p e r a t u r e for

O H - Aromatic Hydrocarbon Reactions and Branching Ratios for the Abstraction Reaction a Aromatic Benzene Toluene Ethylbenzene o-Xylene rn-Xylene p-Xylene n-Propylbenzene Isopropylbenzene o-Ethyltoluene rn-Ethyltoluene p-Ethyltoluene 1,2,3-Trimethylbenzene 1,3,5-Trimethylbenzene tert-Butylbenzene

k (10 -12 cm 3 molecule -1 s - l)

Branching ratio for abstraction

1.2 6.0 (5.5) b 7.1 13.7 23.6 (22.0) b 14.3 6.0 6.5 12.3 19.2 12.1 32.7 57.5 (57.3) b 4.6

0.05 0.07-0.12 0.05-0.10 0.04 0.08

0.06 0.03

a From Atkinson (1994). b Numbers in parentheses from Kramp and Paulson (1998).

alkyl radical formed in reaction (61) adds 0 2 and ultimately forms benzaldehyde and smaller amounts of benzyl nitrate, as shown in Fig. 6.12. The greatest uncertainty in the mechanism is the fate of the stabilized O H - a r o m a t i c adduct. It had been assumed that it would react with 0 2 and ultimately lead to the formation of oxygenated compounds. However, there is also some evidence that the kinetics of the reaction of the adduct with 0 2 may be sufficiently slow at room temperature that the reaction with NO 2 may be competitive at the relatively high concentrations of NO 2 used in laboratory studies of these reactions. The rate constants at room temperature for the reactions with 0 2 and NO 2 have been reported to be ~ 5 • 10 -16 and ~ 3 • 10 -11 cm 3 molecule -1 s -1 respectively (Knispel et al., 1990; Goumri et al., 1991), in which case the removal of the adduct by these two pathways at 1 atm in air would be equal only at the very high NO 2 concentration of ~3.5 ppm. While these kinetic data suggest that the adduct will react with 0 2 under atmospheric conditions, in many labora-

d,c,o, c.oo

OH +

~

N; ~

~

+ HO2

NO$9 20N02 F I G U R E 6 . 1 2 Mechanism of formation of benzyl nitrate and benzaldehyde from reaction from the methyl side group of toluene.

209

G. REACTIONS OF SIMPLEAROMATIC HYDROCARBONS

tory studies where higher NO X levels are used, a significant portion of the O H - a r o m a t i c adducts may be trapped by NO 2, rather than reacting with 0 2. As a result, the products from some laboratory studies may not be directly applicable to ambient air. More recent measurements (Bjergbakke et al., 1996) report the observation of the H O - C 6 H 6 0 2 peroxy radical by UV absorption (which has been controversial; see Koch, 1997, and Pagsberg, 1997). This study reports rate constants for the O H - C 6 H 6 adduct of 5 • 10-13 cm 3 molecule-1 s-1 with 0 2 and 1 • 10 -11 cm 3 molecule -1 s -1 with NO2, in which case reaction of the adduct with NO 2 is not expected to be important. This is clearly an area that warrants further study in order to be able to extrapolate reliably the results of laboratory studies to atmospheric conditions. Species such as cresol are known to be products of the toluene oxidation and have commonly been assumed to be formed via reactions such as (63) CH 3

CH 3 -}-02 9

>

+ H O 2 (63)

The actual mechanism may not be as simple as implied by this equation, however. For example, Narita and Tezuka (1982) have shown that, in the solution phase oxidation at least, some of the cresol product contains an oxygen atom from the 0 2. In addition, in the case of reaction of benzene, phenol may be formed by direct displacement of a hydrogen atom by OH (Bjergbakke et al., 1996; Koch, 1997; Pagsberg, 1997). There are also data suggesting that 0 2 reacts with the O H - a r o m a t i c adduct by abstraction of the hydroxyl hydrogen to give an epoxide that photolyzes to the alcohol. Such epoxides are known to be in equilibrium with the isomeric oxepins. For the benzene reaction, for example, the formation of phenol may occur via the

following mechanism: OH

O2~...~ HO2

(;>O CO

O h ~ h v Benzene oxide

For example, Klotz et al. (1997, 1998) have shown that benzene oxide/oxepin photolyzes in sunlight to give phenol with a yield of 43.2 _+ 4.5%. This reaction mechanism is therefore feasible for the formation of phenol in the b e n z e n e - O H reaction. However, photolysis of toluene 1,2-oxide/2-methyloxepin gave o-cresol only in small yields, 2.7 _+ 2.2% (Klotz et al., 1998); this suggests that cresols formed in the O H - t o l u e n e reaction come primarily from the direct reaction (63)of the OH adduct with 02, in contrast to the conclusions of Moschonas et al. (1999). A variety of smaller multifunctional oxygenated compounds are also found as products of the gas-phase O H - a r o m a t i c reactions. Table 6.17 shows the yields of the smallest dicarbonyl compounds from these reactions, which, while small, are not insignificant. In addition to these products, a variety of other multifunctional compounds are typically found, the numbers, types, and concentrations of these products depending on the analytical methodologies used, the reaction conditions, and the skill and imagination of the experimentalist! Table 6.18, for example, shows some products observed in the photooxidation of toluene in air where the loss is due to attack by OH (Dumdei et al., 1988). In this particular study, ~ 44% of the reacted toluene could be accounted for by the products shown in Table 6.18. Similarly, Yu et al. (1997) and Yu and Jeffries (1997) report a total of 50 products from the oxidation of

TABLE 6.17 Fractional Yields of Dicarbonyl Compounds from the OH Reaction with Some Aromatic Hydrocarbons at 1 atm Total Pressure and 298 Ka Aromatic

Benzene Toluene o-Xylene rn-Xylene p-Xylene 1,2,3-Trimethylbenzene 1,2,4-Trimethylbenzene 1,3,5-Trimethylbenzene a From Atkinson (1994).

Glyoxal (CHO)2

0.21 0.08-0.15 0.03-0.09 0.09-0.13 0.12-0.24 0.06-0.07 0.05-0.08

Oxepin

Methylglyoxal CH3COCHO

0.08-0.15 0.12-0.25 0.27-0.42 0.11 0.15-0.18 0.36 0.60-0.64

Biacetyl (CH3CO)2

0.09-0.26 0.32-0.45 0.05-0.11

210

6. GAS-PHASE REACTIONS IN IRRADIATED ORGANIC-NOx-AIR MIXTURES T A B L E 6.18 Some Products Observed in the Photooxidation of Toluenea Species

study (Bartolotti and Edney, 1995). Kwok et al. (1997) have studied the reactions of OH with o-, m-, and p-xylene and identified a variety of open-chain multifunctional products, including

Percentage yield

CH3C(O)CHO (CHO) 2 Methylbutenedial C6HsCHO (Hydroxymethyl)butenedial Peroxyacetyl nitrate Oxoheptadienal CH 3COOH HCHO Hexadienal Hydroxyoxoheptadienal Hydroxydioxohexenal Dioxohexenal Hydroxyoxohexenal Hydroxyhexadienal Methyl vinyl ketone Methylfuran Hydroxyoxobutanal Hexadienedial Hydroxymethyl vinyl ketone Butenedial Hydroxybutenedial

7.7 5.8 5.8 5.0 3.4 3.3 2.7 1.8 1.0 1.0 1.0 0.8 0.7 0.6 0.6 0.5 0.5 0.5 0.3 0.3 0.3 0.2

C H 3C H O

0.2

Pentadienal Acrolein Pyruvic acid Total

HC(O)CH= CHCHO, CH3C(O)CH= C H C H O , CH 3C(O)CH = CH C(O)CH 3, CH 3C(O)C(CH 3) = CH CH = CH CH O,

0.1 0.06 0.06 44.2

a From Dumdei et al. (1988); see also Yu et al. (1997).

toluene in air in the presence of NOx. These include additional compounds from those in Table 6.18, such as benzoquinones and epoxides, the latter hypothesized to be formed from an epoxide type of structure predicted theoretically for the O H - a r o m a t i c - O 2 adduct in one

and their isomers. It is typical of such aromatic oxidations that only ~ 4 0 - 7 0 % of the reacted parent organic can be accounted for in measured products. Figure 6.13 shows one postulated reaction sequence taking the adduct to methylglyoxal and butenedial. This represents just one of many possible reaction paths. OH can also add to the ring at the meta or para positions, in addition to the ortho position shown in Eq. (62). The addition of 0 2, shown in Fig. 6.13 to occur at the 1-position relative to the OH group, could also occur at the 3- or 5-positions. Ring-cleavage products are also observed in the absence of NO X(Atkinson and Aschmann, 1994; Seuwen and Warneck, 1996). Subsequent cyclization of the peroxy radical to form allylically stabilized five-membered bicyclic radicals is expected to be the most energetically favored, with the formation of nonallylically stabilized biradicals being endothermic (Andino et al., 1996). Theoretical studies also suggest that some of the peroxy radicals formed by addition of 0 2 to the OH-aromatic adducts may react with NO in competition with cyclization (Andino et al., 1996), generating NO 2 and an aromatic peroxy radical that may subsequently form phenolic and unsaturated derivatives through reactions with O2:

R

R

I

I

R

. ~

+ N~

> N O 2 +l~

I

R

R

~OH H

HXo:

00" H

R

H

O2

R

~ O

H

R

O

+ HO2

R

H

R

OH H

+ HO 2

211

G. REACTIONS OF SIMPLE A R O M A T I C HYDROCARBONS

@OH

CH 3

CH 3

+ OH

CH 3

CH 3

H

OH H

O a-~ H

-oo.@oH CH 3

CH 3

0

+

H TM

OO 9 H

o

\\

HC H,LI

~

o

,~c

O~

o o. CH 3

\

OH

13scission

C/ / \ H

CH/CH

NO CH 3 OH H

+

NO 2

O

o II H C - - CH = C H - - C H

OH

O II CCH3

I

Cm

I

H

1 O O II II H C - - CH = C H - - CH

OH

0 II 9 - - CCH 3

I

+

I

H

Butenedial

O

O

"~c ii 3 - - CCH /

+

HO 2

H

Methyl glyoxal

FIGURE 6.13

O n e p o s t u l a t e d fate of t h e O H - t o l u e n e

The simplest phenoxy radical, C6H50, does not react with 0 2 (k < 5 • 10 -21 cm 3 molecule-1 s-1) but does react with NO (k = 1.9 • 10 -12 cm 3 molecule -1 s -1) and with NO 2 (k = 2.1 • 10 -12 cm 3 molecule -1 s-l), suggesting that reactions with NO X will be its primary fate in the troposphere (Platz et al., 1998b). However, this may not be the case for the larger, hydroxylated phenoxy radicals from the O H - a r o m a t i c - O 2 - N O reaction sequence. In short, there are a multitude of potential reaction pathways that must be considered in O H - a r o m a t i c reactions, and the details of the mechanism remain to be elucidated. The formation of multifunctional, highly reactive "products" such as butenedial may be responsible for

a d d u c t in air.

much of the missing carbon in these reactions. Such products will also react rapidly with OH, with 03, and, when present, with NO3, as well as photolyze. For example, Fig. 6.14 shows the infrared spectrum obtained after a mixture of the cis and trans forms of butenedial were photolyzed using fluorescent lamps with wavelengths 3208 A < 480 nm and the absorptions due to the products CO, CO 2, HCHO, and H C O O H were subtracted out (Bierbach et al., 1994). Most of the remaining bands are due to the product 3H-furan-2-one formed from an intramolecular rearrangement of butenedial. Reaction of butenedial with OH was also shown to give maleic anhydride as a major product (along with glyoxal). Figure 6.15 shows possible mechanisms for

212


OH +

9o-Z o.

H O

8 + 02

cyclisation

oZ- . 1900

1850

1800

1750

1700

1650

1600

O~.H NO

+ 021

Wavenumber (cm "1)

FIGURE 6.14 Infraredspectrum due primarilyto 3-H-furan-2-one formed in the photolysis of butenedial after subtraction of bands due to CO, CO2, HCHO, and HCOOH (adapted from Bierbach et al., 1994).

O

O~O~ ~'- H

OO,

Decomposition

"~ NO2+ O

formation of maleic anhydride from butenedial, which is itself formed in the t o l u e n e - O H oxidation. Further complicating mechanistic understanding is the possibility of different isomers of some of the multifunctional products. For example, one of the products of the p - x y l e n e - O H reaction is 3-hexene-2,5dione:

+NO2

CO2

and otherproducts

O~

NO2+ O ' @ O maleic anhydride

O

-~

1"O

O

+ OH

maleic anhydride

O As shown in Fig. 6.16, it is expected to react with OH to form 4-hydroxyhexane-2,3,5-trione, which, as Wiesen et al. (1995) have shown, exists in equilibrium with the enediol form. Further reaction of the enediol with OH forms the trione shown in Fig. 6.16. Wiesen et al. (1995) suggest that such polyketones could account for 25-30% of the carbon, which would improve the carbon balance in O H - a r o m a t i c reactions considerably. In short, the mechanism of O H - a r o m a t i c reactions remains today one of the least understood areas in tropospheric chemistry.

2. Nitrate Radical ( N O 3 ) As seen in Table 6.1, the reactions of the nitrate radical with the simple aromatic hydrocarbons are generally too slow to be important in the tropospheric decay of the organic. However, one of the products of the aromatic reactions, the cresols, reacts quite rapidly with NO 3. o-Cresol, for example, reacts with NO 3 with a room temperature rate constant of 1.4 • 10 -11 cm 3 molecule-1 s-1, giving a lifetime for the cresol of only 1 min at 50 ppt NO 3. This rapid reaction is effectively an overall hydrogen abstraction from the pheno-

F I G U R E 6.15 Postulated reaction scheme from the reaction of cis-butenedial with OH (adapted from Bierbach et al., 1994).

lic OH (although the mechanism is likely not a direct abstraction process; e.g., Atkinson, 1994): CH3 I~oH I~J

+ NO 3

CH 3 ~~j/O" >

+ H N O 3 (64)

3. C h l o r i n e A t o m s (C1) Chlorine atoms react with aromatic hydrocarbons, but only at a significant rate with those having saturated side chains from which the chlorine atom can abstract a hydrogen or unsaturated side chains to which it can add. For example, the rate constant for the C1 atom reaction with benzene is 1.3 • 10 -15 cm 3 molecule -1 s -1 (Shi and Bernhard, 1997). On the other hand, the rate constants for the reactions with toluene and p-xylene are 0.59 • 10 -l~ and 1.5 • 10 -l~ cm 3 molecule -1 s-1, respectively (Shi and Bernhard, 1997), and that for reaction with p-cymene is 2.1 • 10 -l~ cm 3 molecule -1 s -1 (Finlayson-Pitts et al., 1999). Hence

213

H. REACTIONS OF OXYGEN-CONTAINING ORGANICS

[•3+ OH

02 ~ N ~ _ ~ .

2

O H H O ii i i ii CH 3 - C - C = C - CCH 3

ring

I

opening

OH,

02

CH 3

O H H O ii i i ii CH 3 - C - C - - C - C - CH 3 i

OH

O H O ii i ii OH 3 - C - C - - C - C - OH 3 i ii OH O

~

i

O-O

NO2 O H H O ii i i ii CH 3 - C - C - - C - C - OH 3

2

i

OH

i

O

4 - hydroxyhexane - 2, 3, 5 - trione

O

O

ii

ii

CH 3 - C - C = C - C - CH 3 I

OH

OH

I

0

II

I

I

OH

Enediol

OH.

0

I

II

CH 3 - C - C - C - C - CH 3 OH ((x-Hydroxy

I

OH radical)

O OH O ii i ii CH 3 - C - C - - C - C - CH 3 i ii OH O 3, 3 - Dihydroxyhexane

- 2, 4, 5 - trione

F I G U R E 6.16 Formation of polyketones in the OH reaction with p-xylene (adapted from Wiesen et al., 1995).

reactions of aromatic hydrocarbons with chlorine atoms will be significant primarily for those species having reactive groups attached to the ring.

Aldehydes. Reactions with aldehydes occur by abstraction of the relatively weak ( ~ 87 kcal mo1-1) aldehydic hydrogen: OH(NO3, C1) + R C H O

H. R E A C T I O N S OF OXYGEN.CONTAINING ORGANICS 1. R e a c t i o n s of O H , N O 3, a n d C1 Table 6.19 gives the room temperature rate constants for the reactions of some oxygen-containing organics with OH as well as with NO 3 and with chlorine atoms. As expected, the OH reactions are reasonably fast, ranging from ~ 10-13 cm 3 molecule- 1 s- 1 with acetone to more than 10-1~ cm 3 molecule-~ s-~ with the aldehydes, furan, and, in general, compounds having an alkyl group larger than - C H 3. Chlorine atoms also react very quickly. The NO 3 reactions in most cases are sufficiently slow that OH is the only significant tropospheric oxidant on a global basis for these oxygenated compounds. However, NO 3 can play a role in aldehyde oxidations, especially at night when OH levels are very low.

(65)

RCO + H 2 0 (HNO3, HC1).

The RCO radical produced then adds 0 2 a s expected. For example, for acetaldehyde, reaction (65) is followed by O

O

II

II

C H 3 C + O2 ~

O

II

CH3CmO

O NO> C H 3 C ~ O

+ NO 2 , (66)

O

~o2

II > CH3CmOONO2

,

PAN

(67) O

II CH3CO

> CH 3 + CO 2 .

(68)

214


R o o m T e m p e r a t u r e Rate Constants (cm 3 m o l e c u l e - 1 s - 1) for the Reactions of Some O x y g e n - C o n t a i n i n g Organicsa

Organic

OH

Aldehydes HCHO CH3CHO CH3CHzCHO (CHO) 2 CH 3

9.2 1.6 2.0 1.1

I

CH2=C~CHO

• • • •

10-12 10 -11 10-11 10-11

NO 3

CI

5.8 • 10-16 2.7 • 10-15

7.3 • 10-11 w 7.2 • 10-11 1.2 • 10- lo

3.4 • 10-11

NO ~

OOCH2NHCH 3

(84)

NO 2 > O C H 2N H C H 3

~

HO2

O% C N/H H / v m * "~,CH 3"

(86)

Most of the products predicted by this reaction scheme have been observed in laboratory studies, including dimethylnitramine [(CH3)2 NNO2 ], HONO, tetramethylhydrazine [(CH 3)2 NN(CH 3)2 ], and small amounts of methylformamide as products of the photooxidation of dimethylamine (e.g., see Pitts et al., 1978; and Tuazon et al., 1978). In addition, HCHO, CO, and small amounts of dimethylformamide were identified in these studies. It is noteworthy that the dimethylamino radical reaction with 0 2 is about a factor of ~ 106-107 slower than its reactions with NO and NO 2. For example, Lindley et al. (1979) measured the ratio of rate constants k s l / k 8 2 = 1.5 • 10 .6 and k s l / k 8 3 a = 3.9 • 10 -7. Thus, at ~ 10 ppb NO X, reactions of the nitrogen-centered radical with NO and NO2, in addition to 02, become important. This is perhaps not surprising, given that NH 2 radicals also react extremely slowly with 02, with an upper limit of 6 • 10 -21 cm 3 molecule- 1 s- 1 (Tyndall et al., 1991). There are few kinetic or product studies of the reactions of 0 3 with simple amines. Figure 6.20 shows the infrared spectrum of the products of the reaction of dimethylamine with 0 3 (Tuazon et al., 1994). The major products observed are C H 3 N = C H 2 , CH3NO2, HCHO, CH3NHCHO, (CH3)2NCHO, (CH 3)2 NCH 2N(CH 3)2, HCOOH, and CO 2. The major nitrogen-containing products were explained by Tuazon and co-workers by the following mechanism:The

I. REACTIONS OF NITROGENOUS ORGANICS

223

(CH3)2NH + 0 3 (CH3)2NH (87)

l [(CH3)2NOH]*

(CH3)2NO + 0 2 + O H

CH3N~CH 2

CH3NO 2 + 0 2 + CH 3

>

C H 3 N H C H O arises from the reaction of the O H generated in the system with the parent amine to form the C H z N H C H 3 alkyl radical, followed by reaction (86). N , N , N ' , N ' - T e t r a m e t h y l d i a m i n o m e t h a n e , (CH3)zNCHzN(CH3)2, seen in Fig. 6.20C, was shown to be formed in the reaction of dimethylamine with the product HCHO: 2(CH3)zNH + H C H O (CH3)zNCHzN(CH3) 2 + H 2 0 .

(88)

Further reaction of this product with 0 3 was shown to give (CH 3)2NCHO, observed as a product in Fig. 6.20A. It should be noted, however, that in studies of amine photooxidations, it is generally true that a significant fraction of the reacted parent amine remains unaccounted for in the identified products. Clearly, the mechanisms and products are complex and warrant further investigation. b. Nitrosamines

One difficulty in studying the photooxidation of amines is the rapid reaction in the dark with nitrous acid to form nitrosamines (Hanst et al., 1977; Pitts et al., 1978): R2NH + HONO ~ R2NNO + H20.

(89)

As discussed in Chapter 7.B.3, NO 2 undergoes a surface reaction with water, which is perhaps enhanced at the air-water interface, forming HONO: 2NO 2 + H20

surface "; HONO

+ H N O 3.

> (CH3)zN + OH

(90)

Thus, in the course of preparing reactant mixtures for photooxidation studies under typical atmospheric conditions where both NO 2 and water vapor are present, it is essentially impossible to avoid the production of some HONO, and in the case of studies of amine reactions, some nitrosamines. However, this too is quite relevant, since nitrosamines are carcinogenic in experimental animals. In addition, there are a number of sources that emit nitrosamines directly into the air,

HO2

>H C H O including leather tanneries, rocket fuels, tire and amine factories, and tobacco smoke (e.g., see Fine, 1980). The major atmospheric fate of the N-nitrosamines such as N-nitrosodimethylamine is photolysis (Tuazon et al., 1984): (CH3)2NNO + h v ~ (CH3)2N + NO.

(91)

The dimethylamino radical then reacts as described earlier. c. Hydrazines

Hydrazines see widespread use as fuels, for example, in the space shuttle and as a source of emergency power in the F-16 fighter plane. As a result of the industrial and fuel uses of hydrazines, with their accompanying transport and storage, some emissions to the atmosphere occur and hence there is interest in their atmospheric reactions. Hydrazines do not photolyze in the actinic UV, but reactions with O H and 0 3 must be considered. The rate constants for reaction of O H with N2H 4 and C H 3 N H N H 2 are (6.1 _+ 1 . 0 ) • 10 -11 and (6.5 _+ 1.3) • 10 -11 cm 3 molecule -1 s -1, respectively, essentially independent of temperature over the range 298-424 K (Harris et al., 1979). At an O H concentration of 1 • 106 cm -3, the lifetimes of both N2H 4 and C H 3 N H N H 2 will be ~ 4 - 5 h. Harris and co-workers (1979) estimate that the rate constant for the reaction of OH with 1,1-dimethylhydrazine is ~ ( 5 _+ 2) • 10 -11 cm 3 molecule- 1 s-1, so that its lifetime with respect to O H will be similar, ~ 6 h. Reaction with 0 3 is also relatively fast. Tuazon et al. (1981) estimate that the rate constant for the N 2 H 4 - O 3 reaction at 294-297 K is ~ 1 • 10 -16 cm 3 molecule -1 s -1. This corresponds to a lifetime of about 1 h at an 0 3 concentration of 0.1 ppm. The rate constants for the O 3 - C H 3 N H N H 2 and O3-(CH3)2NNH 2 reactions were too fast to measure under their experimental conditions; the reactions of ~ 1-3 ppm 0 3 with 0.4-4 ppm C H 3 N H N H 2 and ~ 2 ppm 0 3 with ~ 0.2-2 ppm (CH3)2NNH 2 were complete in less than 2-3 min.

224

6. GAS-PHASE REACTIONS IN IRRADIATED O R G A N I C - N O x - A I R MIXTURES

From these data, the rate constants must be > 10 -15 cm 3 molecule- 1 s- ~ and the lifetimes of these two hydrazines must be less than 7 min at 0.1 ppm 0 3. The mechanism of the reaction of hydrazines with 0 3 has been investigated using FTIR (Tuazon et al., 1981; Carter et al., 1981b). In the case of N2H 4, the major product was H 2 0 2 , and N2O appeared as a minor product; these are consistent with the following mechanism:

CH3N=NH

H202

__m o ~

Initiation

H 2 N N H 2 + 0 3 --+ H 2 N - - l q H + OH

+ 0 2

O v~

o~ + 0 2 --->

HN:NH

+ 0 3 + HNIlq

+ HO 2

+ OH

H 2 N N H 2 + OH + H 2 N / l q H

+ 0 2

+ H20

(93)

CH3ONO2 /~.

0 C

a~rlllllP [[ II II /

H N = I ~ I --> H + N 2 M

H + 0 2 + HO 2 M, H 2 0

HO 2 + HO 2

> H 2 0 2 + 02

(95)

(96)

I

I

1000

I

I

1200

I

Wavenumber (cm"1)

I

2000

I

I

2200

FIGURE 6.21 FTIR spectra taken during reaction of CH3NHNH 2 with 03: (a) 3.68 ppm CH3NHNH2 before reaction; (b) 2 min after injection of 2.8 ppm 03; (c) 2.8 ppm 03 injected 38 min after first injection (spectrum taken 2 min after second injection). NH 3 absorptions have been subtracted from (a) and (b), and both NH 3 and 03 absorptions from (c) (adapted from Tuazon et al., 1981).

(97) (98)

According to this mechanism, most of the nitrogen in the hydrazine would form N2, which would not have been detected in this system. The diazene H N = N H would be expected to react with OH radicals: + OH --> H N : I ~ I + H 2 0 ,

I

800

(94)

Product Formation

HN:NH

III

CH2N2 + OH + 0 2 nt- HO2, (lOOa)

CH3NmNH

+ 0 3 --+ C H 2 N 2 + H 2 0

+ 0 2.

(100b)

In excess O3, CH 2N2 can react via (100c): CH2N 2 + 0 3 --+ H C H O + 0 2 nt- N 2.

(100c)

The formation of C H 3 O O H , H C H O , C H 3 O H , and C H 3 O N O 2 is expected from secondary reactions of the methyl radical as discussed earlier in this chapter. The reaction of 1,1-dimethylhydrazine with 0 3 gave the carcinogen N,N-dimethylnitrosamine as the major product in ~ 60% yields within 2- to 3-min reaction time (Tuazon et al., 1981). Minor products were H C H O , H 2 0 2 , H O N O , and perhaps NO X. For a discussion of the complex mechanisms, the reader should consult the original reference.

J. CHEMISTRY OF REMOTE REGIONS

NO 2 has also been shown to react with 1,1-dimethylhydrazine in air, forming HONO and tetramethyltetrazine-2, (CH3)2NN=NN(CH3)2 (Tuazon et al., 1983b). The reaction is also proposed to involve abstraction of a hydrogen from the weak N - H bond by NO2, forming HONO. The tetramethyltetrazine-2 is hypothesized to be formed by the addition of NO 2 to the (CH3)2NNH radical, followed by decomposition to (CH3)2N 2 + HONO and the self-recombination of the (CH3)2N 2 radicals (Tuazon et al., 1982). The apparent overall rate constant for the reaction was 2.3 • 10 -17 cm 3 molecule -1 s -1 so that the lifetime of 1,1-dimethylhydrazine at an NO 2 concentration of 0.1 ppm would be 5 h. Since the lifetimes with respect to 0.1 ppm 0 3 or ] X 10 6 OH radicals cm -3 are ~ 7 min and 6 h, respectively, the reaction of NO 2 can contribute to the atmospheric reactions of the hydrazine only at low 0 3 levels. Hydrazine, monomethylhydrazine, and asymmetrical dimethylhydrazine have also been shown to react rapidly (k > 10 -15 cm 3 molecule -1 s - l ) with HNO 3 in the gas phase to form the corresponding hydrazinium nitrate aerosols (Tuazon et al., 1982).

J. C H E M I S T R Y O F R E M O T E REGIONS Gas-phase chemistry in remote areas is, in most cases, analogous to that in more polluted regions. The major difference is in lower NO X emissions and hence concentrations. In addition, in continental regions, there are substantial emissions of biogenic organics, many of which are highly reactive toward OH, 03, NO3, and C1 atoms and in oceanic regions, dimethyl sulfide (DMS), which reacts with OH, NO3, and C1 atoms. As discussed briefly in Chapter 1 and in more detail in Chapter 14, it is unlikely that there are any regions at the earth's surface that have not been significantly impacted by anthropogenic emissions. Even over the central Atlantic and Pacific Oceans, for example, significant contributions to the chemistry from anthropogenic emissions are often observed (e.g., Parrish et al., 1993b; Dickerson et al., 1995). One means of testing for anthropogenic emissions is the use of a correlation between CO and 0 3. CO is generated primarily from combustion processes in vehicles and industrial processes (e.g., Fig. 2.7 and Table 2.1), and its chemistry is the relatively slow reaction with OH (lifetime of ~ 90 days at an OH concentration of 5 • 105 cm-3). As a result, as an air mass containing VOC and NOx ages and undergoes the photochemical reactions discussed earlier, 0 3 is formed. Such air masses therefore often

225

have both increased 0 3 and CO, with zX[O3]/zX[CO] 0.3-0.4 being typical (e.g., Dickerson et al., 1989; Chin et al., 1994; Wang et al., 1996; Harris et al., 1998; Parrish et al., 1993b, 1998; Kajii et al., 1998). This relationship is impacted, of course, by other factors as well such as the production of CO in the oxidation of VOC and by the deposition of 0 3 during transport, for which corrections can be estimated (e.g., Chin et al., 1994). In this section, we discuss briefly the distinguishing chemistry associated with remote regions of the troposphere, including continental and marine areas, focusing primarily on regions of minimal anthropogenic influence. There is increasing evidence of interesting chemistry in the upper troposphere, which is also discussed. Finally, we briefly treat some unusual chemistry occurring in the Arctic, which is characterized by long periods of darkness and low temperatures during winter followed by long periods of sunlight.

1. Emissions of Biogenic Organics As discussed briefly in Chapter 2, there are substantial biogenic emissions of VOC. (We exclude a discussion of methane, since it is of low reactivity and its chemistry is quite straightforward; see Chapter 14 for a discussion of its implications for global climate change.) On a global basis, the total may exceed anthropogenic emissions by as much as an order of magnitude (see Tables 2.1 and 2.2 and the volume edited by Hewitt, 1999). In urban areas, of course, their contributions are relatively less important. For example, Benjamin et al. (1997) report that biogenic hydrocarbons in the South Coast Air Basin of southern California are ~ 10% of the total VOC emission inventory on a typical summer day. The contribution of biogenic organics to urban 0 3 formation is of course variable, depending on the particular locations (e.g., Chameides et al., 1988; Roselle et al., 1991; Pierce et al., 1998; see also Chapter 16). Several thousand different biogenic VOCs have been identified (e.g., Isidorov et al., 1985; Graedel et al., 1986; Puxbaum, 1997; Helas et al., 1997; Fall, 1999). The most well known are ethene, isoprene, and the monoterpenes emitted by terrestrial plants. As discussed in this section, it has been increasingly recognized that there are biogenic emissions of oxygencontaining organics, from small alcohols such as methanol to larger aldehydes, ketones, and alcohols such as 2-methyl-3-buten-2-ol. Figures 6.22 and 6.23 summarize some of the VOC that have been observed and possible mechanisms of production (Fall, 1999). Different compounds are produced in different parts of the plant and by different physiological processes; the

226

Isoprene


o~-Pinene

J3-Pinene

I

Limonene

Terpinolene

Myrcene

FIGURE 6.22 hydrocarbons.

o~-Terpinene

y-Terpinene

Camphene

o~-Phellandrene ~-Phellandrene

A3-Carene

Ocimene

p-Cymene

Chemical structures of some biogenically emitted

present understanding of these processes is reviewed by Fall (1999). Ethene is a plant hormone emitted at a rate of several teragrams per year from plants; smaller amounts are emitted by soils and oceans (Rudolph, 1997). Interestingly, the emissions of ethene as well as some other VOCs (see later) increase significantly (by as much as factors of one to two orders of magnitude) when the plant is stressed, for example by mechanical means, high temperatures, or lack of water. Isoprene is the major single, non-methane VOC emitted by plants:

)

#

Isoprene

Significant amounts of larger hydrocarbons are also generated by plants and emitted to the atmosphere. The larger hydrocarbon compounds generally fall under the classification of isoprenoids, or terpenoids, con-

sisting of groups of 5-carbon isoprene type units (although they are not formed from isoprene). The monoterpenes are the C10 compounds, sesquiterpenes the C15 compounds, diterpenes the C20 compounds, triterpenes the C30 compounds, and tetraterpenes the C40 compounds. Table 6.24, for example, shows one estimate of the annual global emissions of isoprene, other monoterpenes, and VOCs as well as methane (Guenther, 1999). Emissions of isoprene are believed to be about four times those of the other monoterpenes and about equal to all other VOCs. Hardwood species such as oaks, poplars, aspen, and ironwood are generally isoprene emitters. However, even within plant families, not all species are isoprene emitters. For example, while North American oaks emit isoprene, many European oak species do not. For example, Steinbrecher et al. (1997) measured the emissions of isoprene and monoterpenes from five species of Mediterranean oak; two of them were strong isoprene emitters, whereas the other three did not emit significant amounts. Similarly, Kesselmeier et al. (1998) measured emissions of isoprene and monoterpenes from a Holm oak and a white oak growing side by side; the white oak was a strong isoprene emitter, whereas the Holm oak was a strong monoterpene emitter. In addition to deciduous trees such as some oaks, isoprene is also emitted from other plants, including shrubs, gorse, vines, ferns, and plants characteristic of tropical savannas and peatlands characteristic of boreal regions (e.g., Guenther et al., 1996a,b; Cao et al., 1997; Owen et al., 1998; Janson and De Serves, 1998; Guenther, 1999; Fall, 1999). Isoprene emissions have been reported to be correlated with successional patterns; for example, Klinger et al. (1998) report higher isoprene emissions in the early to mid-successional savanna ecosystems, compared to the later, rainforest development. Other sources (all believed to be relatively small) include marine phytoplankton (e.g., Bonsang et al., 1992; Moore et al., 1994; Milne et al., 1995; Broadgate et al., 1997), bacteria and fungi (e.g., Kuzma et al., 1995), the breath of humans and other mammals (e.g., Gelmont et al., 1981; Mendis et al., 1994; Phillips et al., 1994; Jones et al., 1995; Foster et al., 1996; Sharkey, 1996; Fenske and Paulson, 1999), and some industrial processes (e.g., Guenther, 1999). Isoprene production by plants is very sensitive to light as well as temperature (e.g., Sanadze and Kalandadze, 1966; Tingey et al., 1979). Figure 6.24a shows the production of isoprene from an aspen leaf when light in the photosynthetically active range is turned on and off (Monson et al., 1991; Fall, 1999). When the light is switched on, isoprene emissions rise and when it is turned off, fall even more rapidly. Figure 6.24b shows


[" manytissues ~ ETHYLENE, ._rL. , ~,/ v / "'[~ PHYTOHORMONESj cellwalls 99 METHANOL E~CT ~" !?-)DEPOSI P TIN ION [ flo;ers ~ 100sof VOCS~.~.._ FLORAL I~

~

227

LIGHT (

chloroplasts ] ~, (?)// .~f2-METHYL-3-BUTEN-2-OL HEMITERPENES:f:.~l/ISOPRENE MONOTERPENESy. \ resinducts 1 or glands |

, ~/ A J 1 0' / MONOTERPENES STORED -~--'fl ~ . ~ ~ ~~/~~..~l~q~ I MONOTERPENESj FORMALDEHYDE ~~1 ~ leaves, ~,7stems' """ FORMICACID C6ALDEHYDES [ cell me;b ranes"~ C6ALCOHOLS r~u,s ACETALDEHYDE ~ACETICACID [ PEROXIDATION. f L OTHERCl-C3 ETHANOL METABOLITES ACETONE / / / / / /

FIGURE 6.23 Schematic diagram of mechanisms of production of some biogenically emitted VOCs (adapted with permission from Fall, 1999).

the emissions as a function of this light intensity; the emissions rise rapidly and approach (but do not reach) a plateau (Monson et al., 1992; Fall, 1999). Finally, Fig. 6.25 shows the effect of temperature on the isoprene emission rate; emissions rise to a temperature of 40-45 ~ and then rapidly fall (e.g., Monson et al., 1992; Sharkey and Singsaas, 1995; Fall and Wilder-

TABLE 6.24 Estimated Global Annual Biogenic VOC Emissions (Tg y r - 1)a Source Canopy foliage Terrestrial ground cover and soils Flowers Ocean and freshwater Animals, humans, and insects Anthropogenic (including biomass burning) Total

Other Isoprene Monoterpenes VOCsb 460 40

115 13

500 50

0 1 0.003 0.01

2 < 0.001 < 0.001 1

2 10 0.003 93

~ 500

~ 130

~ 650

a From Guenther (1999) and references therein. b Other VOCs include all volatile organic compounds other than methane, isoprene, and monoterpenes.

muth, 1998; Fall, 1999). These influences of light and temperature are believed to be major factors in the variable emissions of isoprene measured in field studies (e.g., Lamb et al., 1996; Guenther et al., 1996b; Geron et al., 1997; Owen et al., 1997; Guenther and Hills, 1998; Drewitt et al., 1998). This dependence on light levels and temperature is believed to be due to the mechanism of production of isoprene in the plant, which involves the enzyme isop r e n e synthetase and dimethylallyl d i p h o s p h a t e (DMAPP) as a precursor to isoprene (e.g., see Silver and Fall, 1995; and Monson et al., 1995). Either the enzyme, the formation of DMAPP, or both may be light sensitive (Wildermuth and Fall, 1996). The temperature effect has been attributed to effects on the enzyme, increasing its activity initially and then leading to irreversible denaturation ( a n d / o r possibly membrane d a m a g e ) ( F a l l and Wildermuth, 1998). Conifers tend to be sources of a variety of larger terpenoids, a major component of which are the C10 monoterpenes (Table 6.24). The structures of some of these are shown in Fig. 6.22. The mechanisms of the formation of these hydrocarbons in plants are closely linked, so that emissions of more than one monoterpene often occur together. For example, geranyl

228


9 10o nt-

80

m

Leaf

_

.mO

~- "r.

400

-~9 E oO

300

.O

Light on

off

on

off

on

60

9

40

LU > .m

-~ rr

_

20 _ _

0 i 10

9

~ v O

._~ E

100

i

i

i

I

20

i

i

i

i

I

i

i

i

30

i

I

40

i

i

i

i

I

50

i

i

i

i

60

T (~

0 I , 0

, ,

,

,

I , , 60

, ,

,

I , 120

,~,

F I G U R E 6.25 Effect of willow leaf temperature on isoprene emission rate (adapted from Fall and Wildermuth, 1998; and Fall, 1999).

T i m e (min)

40 35

~r

m9 ~

9 ~

30 25

20

~

15

x

10

d

5 0

0

200

400

600

800 1000120014001600

Intensity of p h o t o s y n t h e t i c a l l y active radiation (gmol m -2 s-l)

F I G U R E 6 . 2 4 Effect of light on isoprene emission rate from (a) aspen leaf and (b) velvet bean leaf (adapted from Monson et al., 1991, 1992; and Fall, 1999).

diphosphate is a precursor for a- and /3-pinene, limonene, and myrcene, whose generation involves the enzyme limonene synthase (e.g., see Fall, 1999). In contrast to isoprene, which does not have a reservoir in leaves in many cases, monoterpenes are generated and stored in the plant prior to emission so that in general, their emission to the atmosphere is not as closely tied to short-term controls over their biosynthesis (e.g., Monson et al., 1995; Fall, 1999). Emissions of monoterpenes have been observed from a variety of plants, including pines (e.g., Juuti et al., 1990; Guenther et al., 1994; Street et al., 1997; Staudt et al., 1997), resin in pine forests (e.g., Pio and Valente, 1998), spruce (Street et al., 1996), some deciduous trees such as oaks (e.g., Benjamin et al., 1996; Street et al., 1997; Kesselmeier et al., 1998), and gorse (e.g., Cao et al., 1997). Interestingly, as for ethene, increased emissions have been observed when plants are stressed (Fall, 1999). For example, Juuti et al. (1990) report that

the monoterpene emission rates from a Monterey pine increased by factors of 10-50 during rough handling. As for isoprene, emission rates of the monoterpenes increase with temperature, although different plant species exhibit different temperature sensitivities and different compounds can also show different dependencies on temperature (e.g., Tingey et al., 1980; Loreto et al., 1996; Owen et al., 1997; Schween et al., 1997; Drewitt et al., 1998). The temperature dependence of monoterpene emissions is often taken into account by multiplying the base emission rate at a reference temperature ~ by the factor e [/~(;r_ ;rs)], where T is the leaf temperature and /3 is a coefficient that reflects the temperature sensitivity of emissions (e.g., Guenther et al., 1993). Light also affects monoterpene emissions but does not appear to be as significant as for isoprene (e.g., Tingey et al., 1980; Loreto et al., 1996; Guenther et al., 1996b). Larger hydrocarbons such as the C15 sesquiterpenes also have biogenic sources such as sage (Arey et al., 1995). As shown in Table 6.24, oceans and freshwater are not believed to be major sources of VOC to the atmosphere. As indicated earlier, isoprene is thought to be generated in small amounts in the oceans by marine phytoplankton. In addition, a variety of small hydrocarbons have been identified both in seawater and in the air above it, including alkanes (ethane, propane, nbutane, isobutane, n-pentane, isopentane, and nhexane), alkenes (ethene, propene, 1- and 2-butene, isobutene, 1-pentene, and 1-hexene), and acetylene (e.g., Arlander et al., 1990; Rudolph and Johnen, 1990; Bonsang et al., 1991; Plass-Diilmer et al., 1993). While some of these may be due to long-range transport from the continents (Rudolph and Johnen, 1990), it appears that the ocean is indeed a source of most, if not all, of these light hydrocarbons. A major organic found in ocean areas is dimethyl sulfide (DMS), whose oxidation

J. CHEMISTRYOF REMOTEREGIONS products are believed to play a significant role in particle formation and hence radiative properties in the marine boundary layer (see Chapters 8.E.1 and 14.C). DMS also plays a major role in determining the lifetime and fate of NO 3 (Carslaw et al., 1997). In addition to hydrocarbons, biogenic processes also produce a number of oxygen-containing organics. One of the most important appears to be 2-methyl-3-buten2-ol (MBO), first identified in a forested area by Goldan et al. (1993):

Direct emissions of this compound from loblolly pine as well as from lodgepole and ponderosa pines were subsequently demonstrated (Guenther et al., 1996b; Harley et al., 1998). The concentrations of MBO measured by Goldan et al. (1993) were factors of 4-7 times those of isoprene, indicating the potential importance of this compound in the chemistry of remote regions. Its emissions appear to be regulated by temperature and light, similar to that of isoprene. This may be due to formation from a common precursor such as DMAPP (Fall, 1999). Figure 6.26 shows the structures of some oxygencontaining compounds for which there is evidence of direct biogenic emissions. (3Z)-Hexenol and (3Z)-

no) / 2-Methyl-3-buten-2-ol

HCHO Formaldehyde

HCOOH Formic Acid CH3COOH

CH3CHO

HO

Acetaldehyde

2-Methyl-3-buten-2-ol

CH3OH Methanol

C2H5OH Ethanol

~

CH2OH

229

Acetic Acid

O II CH3CCH3 Acetone O

Z,/

Butanone

n-Hexanol

CHO n-Hexanal

(3Z)-Hexenol O

O II

6-Methyl-5-hepten-2-one ~.=~~

Camphor

~ O - C C H 3 (3Z)-Hexenyl acetate

CHO

(3Z)-Hexenal / " , N / / / ' ~ ~ CHO

OH

I

(2E)-Hexenal 1,8-Cineol

Linalool

Thujone FIGURE 6.26 Structuresof some oxygen-containingorganicswith biogenic emission sources.

230


hexenyl acetate, for example, are emitted by a number of plant species. For example, Arey et al. (1991a) identified emissions of these compounds from more than a dozen different agricultural plants in California as well as from Valley Oak and Whitethorn, and Kirstine et al. (1998) reported emissions from grass and clover. Such emissions are enhanced by mechanical damage. For example, Kirstine et al. (1998) reported that the emissions of (3Z)-hexenol and (3Z)-hexenyl acetate increased by three orders of magnitude when the grass and clover were mowed! A second, even larger, emission of these and other organics during the subsequent drying of the plants has been observed (de Gouw et al., 1999).These compounds as well as other volatile C 6 aldehydes and alcohols are generated by the biochemical degradation of c~-linolenic acid in the plants (Fall, 1999). Similarly, the breakdown of linoleic acid gives n-hexanal and n-hexanol. Other examples of emissions of larger oxygen-containing organics include camphor, cineole, and thujone which are emitted by California sagebrush (Arey et al., 1995), cineole from pines and eucalyptus (Staudt et al., 1997), the unsaturated alcohol linalool from the blossoms of Valencia orange trees (Arey et al., 1991b) and from certain pines (Kesselmeier et al., 1997; Staudt et al., 1997), and eucalyptol from grass and clover (e.g., Kirstine et al., 1998). Ciccioli et al. (1997) measured emissions in a Mediterranean region known as the Mediterranean Pseudosteppe and found that not only isoprene but also a range of C6-C10 aldehydes, linalool, and acetic acid were emitted by the vegetation. In addition to these larger VOCs, there are biogenic sources of a wide variety of small alcohols, aldehydes, ketones, and acids. For example, emissions of methanol and acetone have been reported from plant leaves, grass, and clover (e.g., MacDonald and Fall, 1993; Nemecek-Marshall et al., 1995; Fall and Benson, 1996; Kirstine et al., 1998). Table 6.25 shows some of the compounds measured in grass and clover emissions (Kirstine et al., 1998). Clearly, a wide variety of oxygen-containing species are emitted from this one source alone. Direct emissions of formaldehyde, acetaldehye, and formic and acetic acids have been observed from oaks and pines (Kesselmeier et al., 1997). Consistent with the direct observation of the emissions are field measurements in remote areas. For example, Fehsenfeld et al. (1992) measured the composition of VOC at two rural locations in the United States, summarized in Fig. 6.27. Alcohols and carbonyl compounds comprise 40-70% of the total. Of these, a significant portion appear to be direct emissions, with methanol being a major contributor to the oxygencontaining portion. Similarly, Singh et al. (1995) re-

TABLE 6.25 SomeOxygen-Containing Organics Observed in Emissions from Grass and Clovera Percentage of total carbon Compound Methanol Ethanol Acetaldehyde Acetone Propanol 2-Methyl-2-propenal 2,3-Butanedione Butanone 3-Methylbutanal Pentanal 2-Pentanone Hexanal 4-Methyl-2-pentanone Benzaldehyde Nonanal Eucalyptol

Grass

Clover

11-15 16-21 13-16 11-16 2 1 1 0.4-2 1 1 0.2-0.8 4 0.5-1.1 0.9-1.9 0.3-1.1 1

15.1 0.4 0.3 22.9 < 0.1 < 0.1 < 0.1 50.0 < 0.1 < 0.1 < 0.1 < 0.1 < 0.1 < 0.1 < 0.1 0.1

a Adapted from Kirstine et al. (1998).

ported relatively high concentrations of methanol and acetone in the free troposphere, at least a portion of which may be due to biogenic emissions. For example, based on a comparison of measurements and model predictions, Wang et al. (1998)propose that biogenic emissions account for about 40% of the acetone on a global basis. As is expected, there are also a variety of oxygencontaining organics found in rural and remote areas, which are oxidation products of the directly emitted biogenics. For example, in areas with significant isoprene emissions, the oxidation products methyl vinyl ketone (MVK), methacrolein (MACR), and 3-methylfuran are also typically present (e.g., Yokouchi et al., 1993; Montzka et al., 1993, 1995; Biesenthal et al., 1998; Helmig et al., 1998b). Biesenthal and Shepson (1997) suggest that MVK and MACR may also be generated by automobile exhaust, based on the correlation of these compounds with CO in an urban area. Indeed, separating out direct emissions and the formation by oxidation in air of other biogenics is not straightforward. For example, 6-methyl-5-hepten-2-one (Fig. 6.26) has been reported in air in different locations by a number of groups (e.g., Ciccioli et al., 1993a, 1993b; K6nig et al., 1995; Helmig et al., 1996). However, the reaction of 0 3 with organics containing the structural group (CH3)2C=CHCH2CH2C(CH3)=C also gives this compound (e.g., Fruekilde et al., 1998), as expected from the earlier discussion of mechanisms of ozonolysis and studies of struc-

231


Alcohols 18%/

Alkanes 10% ~//Y~~/~ Aromatics

Alabama

Total carbon: 76 ppbC

Carbonyls 22% Natural NMHCs 48%

Alkanes 92 2 %

Alcohols 35% Aromatics 4%

Colorado

Total carbon: 29 ppbC

SIIII.IIII.IIII.IIII.IIII.IIII.IIII.IIII.III....I~... ~ii.iiii.iiii.iiii.iiii.iiii.iiii.iiii.iiii.iiii.iiii.iiii.iiii.iiii.iiii..i

Natural NMHCs 6%

j

Carbonyls 33%

F I G U R E 6.27 Distribution of organics observed in rural areas in Alabama and Colorado, respectively (adapted from Fehsenfeld et al., 1992).

turally similar compounds such as linalool, (CH 3)2C= CHCH 2CH 2C(CH 3)(OH)CH = CH 2 (e.g., Shu et al., 1997). For example, ozonolysis of squalene (Fig. 6.28)was demonstrated to form gaseous 6-methyl5-hepten-2-one, acetone, and geranyl acetone, respectively. 4-Oxopentanal was also formed from the further oxidation of 6-methyl-5-hepten-2-one (Grosjean et al., 1996; Smith et al., 1996; Fruekilde et al., 1998). These compounds were also observed when leaves of common vegetation found in the Mediterranean were exposed to 0 3. Furthermore, these products could be formed from the reaction of glass wool that had been in contact with human skin, which also contains squalene as a lipid; such observations suggest the importance of avoiding contamination of samples during measurements of biogenic organics. In short, while a variety of oxygen-containing biogenic organics have been observed to be generated from plants and most are likely direct emissions, care must be taken to distinguish such direct emissions from

possible formation from oxidation of larger biogenic hydrocarbons and/or, in some cases, contamination during sample handling.

2. Chemistry a. Biogenic Hydrocarbons

Although some of the biogenic VOCs are relatively simple compounds such as ethene, most are quite complex in structure (e.g., Figs. 6.22 and 6.26). Furthermore, they tend to be unsaturated, often with multiple double bonds. As a result, they are very reactive (see Chapter 16.B)with OH, 03, NO3, and C1 atoms (e.g., Atkinson et al., 1995a). In addition, because they are quite large and of relatively low volatility, their polar oxidation products are even less volatile. This makes elucidating reaction mechanisms and quantifying product yields quite difficult. For a review of this area, see Atkinson and Arey (1998).

232

6. GAS-PHASE REACTIONS IN IRRADIATED O R G A N I C - N O x - A I R MIXTURES

1

Squalene

03

3

O

O

O

II

CH3CCH 3 +

acetone

9

+ 9

v'v'v'v

vvvv

I

Geranyl acetone 6-methyl-5-hepten-2-one

03

O

O

II

II

H-C-CH2CH2CCH 3 + Criegee intermediate

4-oxopentanal

FIGURE 6.28 Oxidation of squalene to 6-methyl-5-hepten-2-one, acetone, geranyl acetone, and 4-oxopentanal (adapted from Fruekilde et al., 1998).

Some of the reactions, e.g., that of isoprene with OH and NO3, were discussed earlier in this chapter. Table 6.26 summarizes some of the major products observed in the gas-phase reactions of several other biogenic hydrocarbons with OH and 0 3 (Atkinson, 1997a). These products are anticipated, based on the mechanisms described earlier in this chapter. As also expected, the yields of these major products generally do not account for 100% of the reactant lost, and there are a number of other products, including multifunctional species, that are also formed. As an example, the formation of more than 30 individual products has been observed from the reaction of cr-pinene with 0 3 in air, some of which are unidentified, and the same is true for the zX3-carene reaction (Yu et al., 1998). Products included hydroxy oxoacids, hydroxy dicarbonyls, and dicarbonyls. The formation of low-volatility products that form particles (e.g., Hoffmann et al., 1998; Jang and Kamens, 1999) is likely responsible for a significant fraction of

the products missing from the gas phase. For example, Jang and Kamens (1999) have identified a variety of multifunctional oygenated products in aerosol particles from the cr-pinene reaction with 03, including diacids such as 2,2-dimethylcyclobutane-l,3-dicarboxylic acid and pinic acid. These diacids have sufficiently low vapor pressures that they are believed to contribute significantly to the formation of particles in this system. This is supported by studies of the aerosol composition in forested areas. For example, Kavouras et al. (1998) identified cis- and trans-pinonic acids as well as pinonaldehyde and nopinone in particles in a forest in Portugal. The diurnal variations of the pinonic acids and formic acid were similar, peaking in the afternoon as expected if they were formed by the reaction of 0 3 with cr-pinene. On the other hand, the concentrations of pinonaldehyde, expected from the oxidation of crpinene by OH, 03, and NO3, and nopinone, from the oxidation of /3-pinene, were the smallest in the after-

233


TABLE 6.26

Biogenic

Limonene

Some P r o d u c t s and T h e i r Yields O b s e r v e d in the G a s - P h a s e Reactions of Some Biogenic H y d r o c a r b o n s with O H and O3 a Structure

Product

"l"

HO

OH reaction yield

0 3 reaction yield

0.29 _+ 0.06 b

I 0.20 _+ 0.03 b

< 0.04 b

2 • 10 -3

Utter et al., 1992

OH

275

>4 • 10- 3

Hanson et al., 1992

HO 2

275

>0.02 >0.2

Hanson et al., 1992 Mozurkewich et al., 1987 Worsnop et al., 1989

H20 2

273

0.18

NO 2

298

> 1 • 10- 3 1.5 • 10- 3

Ponche et al., 1993 George et al., 1992

NO 3

-~290 273

>2.5 • 10 -3 2 • 10 -4

Mihelcic et al., 1993 Rudich et al., 1996

N20 5

271-282

0.04-0.06

HONO

~ 245,297

0.05 0.06

HNO 3

268-293

0.07-0.19 0.11

295

0.097

George et al., 1992

0.06

Ponche et al., 1993; Bongartz et al., 1995

NH 3

Van Doren et al., 1990 Bongartz et al., 1994 George et al., 1992

Van Doren et al., 1990 George et al., 1992

0.12-0.02

Jayne et al., 1991

260-270

0.04

Jayne et al., 1992

260-285

0.066-0.013

Duan et al., 1993

HCOOH

260-291

0.1-0.02

Jayne et al., 1991

CH3COOH

260-291

0.15-0.03

Jayne et al., 1991

HC1

274-294

0.06-0.18

Van Doren et al., 1990; George et al., 1992

273

0.10

De Bruyn et al., 1994

273

0.14


273

0.13


CH3OH

260-291

HCHO CH3COCH 3

O II CH3SCH 3 O II CH3SCH 3 II

O O II CH3SOH II

O a Source: DeMore et al. (1997).

308

8. INORGANIC AND ORGANIC ACIDS IN THE TROPOSPHERE

phase found in clouds and fogs, there are some intriguing hints that there may also be some chemistry that is unique to the interface which might be thought of as a "fourth phase," in addition to gases, liquids, and solids. This is an area that remains to be explored. As discussed in detail in Sections C.3.d and C.3.e, the fastest atmospheric reactions of SO 2 are believed to be with H 2 0 2 and perhaps with 0 3 at higher pH values. Under extreme conditions of large fog droplets ( ~ 10 /xm) and very high oxidant concentrations, the chemical reaction times may approach those of diffusion, particularly in the aqueous phase. In this case, mass transport may become limiting. However, it is believed that under most conditions typical of the troposphere, this will not be the case and the chemical reaction rate will be rate determining in the S(IV) aqueous-phase oxidation. Many experimental studies of the rates of oxidation of S(IV) in solution have used either bulk solutions or droplets that are very large compared to those found in the atmosphere. In addition, reactant concentrations in excess of atmospheric levels have often been used for analytical convenience. The use of large droplets increases the diffusion times, whereas higher reactant concentrations speed up the aqueous-phase chemical reaction rates. The combination of these two factors can lead to a situation where the rates of the diffusion processes, either of the gas to the droplet surface or more likely within the aqueous phase itself, become comparable to, or slower than, the chemical reaction rate. If this is not recognized, the observed rates may be attributed in error to the intrinsic chemical reaction rate. In the atmosphere, suspended aqueous solutions are present in the form of aerosols, clouds, fogs, and rain. However, these have different liquid water contents (i.e., grams of H20(1 ) per cubic meter of air). As discussed in detail in Chapter 9, fine particles ( < 2-/xm diameter) emitted directly into the air or formed by chemical reactions can remain suspended for long periods of time. Many of these particles contain water, either in the form of dilute aqueous solutions or as thin films covering an insoluble core; as much as 50% of the mass may be liquid water. Since the total particulate mass in this size range per cubic meter of air can be as high as 10 - 4 g m -3 or more, the liquid water content due to these small particles is also of this order of magnitude. Clouds, fogs, and rain, however, have much greater liquid water contents and thus have the potential for contributing more to atmospheric aqueous-phase oxidations. Clouds typically have liquid water contents of the order of ~ 1 g m -3, with droplet diameters of the order of 5-50 /xm; the number concentration and size distribution depend on the type of cloud. Fogs, on the

other hand, have smaller liquid water contents (~0.1 g m -3) and smaller droplet diameters, generally ~ 0 . 5 - 1 0 /xm (Pruppacher and Klett, 1978). Raindrops are, of course, much larger than cloud or fog droplets, with diameters of ~ 0 . 2 - 3 mm and correspondingly large liquid water contents. However, because of their size, they remain suspended in the atmosphere for only minutes en route to the earth's surface, and hence the potential for oxidation processes to occur in raindrops is minimized. While the volume of liquid water present is much larger in clouds and fogs than that in fine particles, the solute concentrations in the latter may be much higher, and this may serve to increase the rate of aqueous-phase oxidations. More importantly, these fine particles are believed to serve as sites for the condensation of water vapor, leading to the formation of fogs and clouds (Chapter 14.C.2). The liquid water content of an air mass plays a role in determining the oxidation rate of SO 2 in aqueous atmospheric droplets. This can be seen from the expression developed in Box 8.2 for the rate of oxidation of S O 2 (in % h -1) in the liquid phase. Figure 8.12 shows the projected conversion of SO 2 to sulfate as a function of the volume of water per cubic meter of air available for conversion in the aqueous phase, covering a range typical of "haze particles," fogs, and clouds for atmospheric lifetimes which are typical for each (Lamb et al., 1987). As expected from Eq. (M), the conversion increases with the water available in the atmosphere. As we shall see, the aqueousphase oxidation does indeed predominate in the atmosphere under many circumstances. Equations (G) and (M) apply as long as the partial pressure of SO 2 in the gas phase, P s o 2 , is measured simultaneously with the solution concentration of S(IV). With these comments regarding the characteristics of atmospheric aqueous-phase oxidations in mind, we summarize the kinetics of S(IV) oxidation in solution by a series of individual potential atmospheric oxidants. c. Oxidation by 0 2 (Catalyzed and Uncatalyzed) While oxidation of S(IV) in solution in the presence of O 2 has been known for many years, there has been considerable controversy concerning the rates, mechanisms, and effects of catalysts such as Fe 3§ and Mn 2§ particularly under atmospheric conditions. However, studies over the past decade carried out in a number of laboratories, particularly those of Hoffmann and coworkers (e.g., Hoffmann and Boyce, 1983 and references therein) Martin and co-workers (1994 and references therein), have identified the various parameters that determine the overall rate of oxidation. As we shall see, the mechanism and kinetics are so complex that past confusion is understandable.

C. OXIDATION OF SO 2

309

B O X 8.2

CALCULATING

AQUEOUS.PHASE IN PERCENT PER

The expression for the rate of oxidation in % h in the liquid phase can be developed from a knowledge of the gas- and aqueous-phase reactant concentrations, the solution rate constant, the Henry's law constants (Table 8.1), and the liquid water content of air. In 1 m 3 of air, the rate of formation of S(VI) in the aqueous phase is given by d[S(VI)]

dt

(moles per m 3 air) = k[X][S(IV)]V,

dissolved

OXIDATION HOUR SO 2

RATES

concentration is given by

(I)

[SO2]aq = H s o 2 P s 0 2 ,

where H is the Henry's law constant based on physical solubility (Table 8.1) and P is the gas-phase pressure of SO 2. The total concentration of S(IV) in solution, taking into account the acid-base equilibria reactions (11)-(13), is then given by

(G) [S(IV)]aq = 7/nso2Ps02 ,

where k is the solution-phase rate constant (L tools -1), [X] and [S(IV)] are the aqueous-phase concentrations of the oxidant and S(IV), respectively, in units of moles per liter of solution, and V is the volume (L) of liquid water, that is, of aqueous solution available, per cubic meter of air. The rate of S(VI) formation is then expressed in moles per cubic meter of air per second. To express this rate of oxidation in % h-~, consistent with the units in which the results of field studies are often reported, one needs to divide this rate by the total number of moles of S(IV) per cubic meter of air, convert the unit time from s -~ to h -~, and multiply by 100 to convert the fraction to percent. The gas-phase moles of SO 2 in a cubic meter of air is given, according to the ideal gas law, by n ]

Jso2

1000Pso 2

RT

'

(H)

where the factor of 1000 converts from L to m 3. The concentration of S(IV) in solution can be calculated using Henry's law in combination with a knowledge of the concentration of total dissolved S(IV) relative to dissolved SO 2 (i.e., SO 2 9H20). Thus the

Studying the oxidation of S(IV) by 0 2 in "pure" water without traces of catalysts or inhibitors has proven extremely difficult. Based on a compilation of many studies, Radojevic (1984) has recommended that the uncatalyzed rate of oxidation (in terms of the rate of sulfate formation) is given by

d[SO2-]/dt = 0.32[SO2-1[H+] 1/2 (in mole L -1 s - l ) .

(N)

(J)

where r/ is the ratio of the total dissolved S(IV) to that of dissolved SO 2. If there are V liters of liquid water per cubic meter of air, then the total number of moles of S(IV) contained in the atmospheric water droplets found in 1 m 3 of air becomes Moles of aqueous S(IV) = rlHso2Ps02V.

(K)

The total number of moles of S(IV) in a cubic meter of air, including both gas and aqueous phases, is thus given by Total S(IV) per m 3 of air 1000Ps02 =

RT

-1- 7]Hso2Ps02 V.

(L)

Thus, combining (G) and (L), the rate of oxidation of SO 2 in % h -1, which occurs in aqueous solution in the atmosphere, is given by 100k[X][S(IV)]V %h-1 =

]

1000Pso 2/RT 7 7]Hs------o-~Pso2V ] X 3600.

(m)

Since it varies with the square root of the hydrogen ion concentration, it has a weak pH dependence. However, as we shall see, this uncatalyzed reaction is too slow to be of importance under typical tropospheric conditions. Much more relevant to the aqueous phase in clouds and fogs in the atmosphere is the catalyzed oxidation of S(IV) by O 2. Both Fe 3+ and Mn 2§ catalyze the oxidation and as described in Chapter 9, both are common constituents of tropospheric aerosols even in remote

3 10


102

"131

the role of iron in the S(IV) oxidation has involved first understanding the nature of such complexes in solution. Further complicating the iron-catalyzed oxidation is that the mechanism changes from an ionic mechanism in the low-pH regime (0-3.6) to a free radical mechanism at higher pHs (4-7). Table 8.6 shows two proposed mechanisms for the iron-catalyzed reaction at high acidities, in the pH range from 0 to 3.6. The recommended rate expression is given by (Martin, 1994):

~o-

tO

o

1-

O m

1 0.1 -

0

10"2 10-6

I

105

I

104

I

I

103

102

I

101

I

1

10

Specific reaction volume (ml m3)

-d[S(IV)]/dt

F I G U R E 8 . 1 2 Percentage of S O 2 converted to sulfate after a time interval At in small "haze" particles, fogs, and clouds as a function of the aqueous reaction volume; note that the time intervals for each one are different, reflecting how long they typically last in the atmosphere (adapted from Lamb et al., 1987).

areas due to their generation from erosion of the earth's crust. Iron has a particularly rich chemistry because it forms a variety of complexes with O H - and with various S(IV) aqueous forms (see, for example, Betterton, 1993; Brandt et al., 1994; and Millero et al., 1995). Figure 8.13 shows the calculated concentrations of various iron complexes as a function of pH in a solution containing 1 • 10 -6 M Fe 3+ and 1 • 10 -5 M S(IV) at an ionic strength of 0.01. Hence elucidating

10"6 FFo3 ~ ~ ~ , "

C

............. "'. Fe(OH)(S03)

""",,/

e'~ I < Z?

10.8

10.9

...

= k[Fe(III)][S(IV)]/[H + ] (in mole L-1 s- 1).

Fe(III) refers to the sum of all three-valent iron in solution, i.e., Fe(III) = Fe 3+ + FeOH 2+ + Fe(OH)~- + FeOHSO 3 + FeSO~-+ etc. Thus the rate of the ironcatalyzed reaction in the low-pH region decreases with increasing [H +]. This means that it shows the behavior depicted in Fig. 8.9a; i.e., it is self-quenching. That is, as S(IV) is oxidized to the acid, the pH falls and the rate also decreases. Despite the simplicity of the rate law implied by Eq. (O), the behavior of the kinetics is very sensitive to a variety of factors. Thus the reaction is inhibited not only by [H +] but also by the ionic strength (I) of the solution, by both S(IV) and S(VI), and, at high pH, by organics. Martin (1994) gives expressions for the dependence of the rate constant k in Eq. (O) on I, S(IV), and S(VI). The effect of ionic strength may be due to effects on the stability of complexes, whereas the sulfate is thought to complex one or more of the catalytic species in the reaction. Because of these complexities, the rate expression in Eq. (O) only applies for [S(IV)] < 1 • 10 -5 [Fe(III)] > 1 • 10 -7 I < 0.01 and [S(VI)] < 1 • 10 -4 mol L -1 where k = 6.0 s -1 (Martin, 1994).

T A B L E 8 . 6 Some P r o p o s e d M e c h a n i s m s for t h e C a t a l y z e d O x i d a t i o n of S ( I V ) in A q u e o u s S o l u t i o n s Hoffmann and Jacob (1984) F e 3 + + HSO~ ~ FeSO~ + H + FeSO~ --+ Internal redox

g 101~

fast fast

10-11

fast

.i10n-12 0

1

2

3

4

5

6

I 7

pH F I G U R E 8 . 1 3 Calculated concentrations of iron species in aqueous solution for [Fe(III)] = 1 • 10 .6 M, [S(IV)] = 1 • 10 .5 M, and I = 0.01. The sulfur complexes are shown by the dotted lines (adapted from Martin, 1994).

(O)

> Addition of HSO~ > Addition of 0 2 > Formation of products

Conklin and Hoffmann (1988) HSO~- ~ SO~- + H + Fe3++ H20 ~ FeOH2++ H + F e O H 2+ + SO~- ~ H O F e O S O 2 complex H O F e O S O 2 complex + O 2 ~ O 2 adduct acid-catalyzed

O 2 adduct

rearrangement

> H O F e 2 + + SO 3 + H 2 0 2

SO 3 + H 2 0 ~ 2 H + + SO42-


In the high-pH (4-7) region, the reaction is first order in S(IV), but the dependence on the catalyst concentration is complex. Martin (1994) recommends the following rate expressions for pH around 4.0 and 5-6, respectively, where it is assumed that there is at least a trace of iron available for catalysis in the 5-6 pH range: pH 4.0

- d [ S ( I V ) ] / d t = 1 • 109[Fe 3+ ]2[S(IV)] (in mol L -

1

s -1)

(P)

pH 5.0-6.0

- d [ S ( I V ) ] / d t = 1 • 10-3[S(IV)] (in mol L -~ s -~) (Q)

At pH > 4, the oxidation is inhibited by organics, suggesting a free radical mechanism. One proposed mechanism, which originates in the work of B~ickstr6m (1934), is shown in Table 8.7. The inhibition occurs when the organics react with the sulfate radical ion, SO[. This inhibition has also been seen in laboratory experiments using fogwater collected in Diibendorf, Switzerland, where oxidation rates for S ( I V ) w e r e less than expected based on the kinetics of the ironcatalyzed oxidation (Kotronarou and Sigg, 1993). The Mn2+-catalyzed oxidation of SO 2 is also complex in both kinetics and mechanism. At pH 2, for example, the reaction is second order in Mn 2+ and zero order in S(IV) (i.e., is independent of the dissolved sulfur concentration) when [S(IV)] > 10 -4 tool L -1 but first order in both Mn 2+ and S(IV) at low concen-

TABLE 8.7 Proposed Mechanism for the Catalyzed Oxidation of S(IV) in Aqueous Solutions of pH Range 4 - 7 a 2Fe(OH)~- SO 2 - -+- SO 3 + F e 2 + - + F e 3 + + SO 2 + organic ~ SO 2 - + p r o d u c t s ~ products k a = 4.5 x 109 M -1 s -1 k c = 1.5 x 109 M -1 s -1 k d = 1.3 • 107 M - 1 s - 1 k e = 2 • 109 M -1 s -1 k f = 9 . 9 • 108 M - i s - 1

kh

kg = see N e t a et al. (1988) a n d W i n e et al. (1989) k h = 410 s

-1

a B~ickstr6m (1934), H u i e and N e t a (1987), and M a r t i n (1994). R a t e constants f r o m H u i e a n d N e t a (1987), N e t a et al. (1988), T a n g et al., (1988), and W i n e et al. (1989).

31 1

trations, [S(IV)] < 1 0 - 6 mol L -1. Furthermore, the rate decreases with ionic strength at all S(IV) concentrations. Finally, there is a synergistic effect in the presence of both Mn 2+ and Fe B+ (Martin, 1994). In summary, the uncatalyzed oxidation of S(IV)occurs in aqueous solution but is very slow. However, given the ubiquitous occurrence of Fe B+ and Mn 2+ (see Chapter 9), the uncatalyzed oxidation is likely irrelevant to atmospheric solutions. The catalyzed oxidations are complex in both kinetics and mechanism. We shall defer a comparison of their importance until other oxidation mechanisms are discussed. However, we shall see that the catalyzed oxidations are likely to contribute significantly to S(IV) oxidation in solution only at pH values near neutral, i.e., in the range of ~ 6 - 7 . As the oxidation occurs and acid forms, the pH falls. The rapid falloff in the rate of the catalyzed oxidation with increasing [H +] then results in a rapid quenching of this path, as expected from Fig. 8.9a. As discussed in Chapter 7, there is some evidence that freezing of aqueous solutions containing nitrite accelerates its oxidation to nitrate (Takenaka et al., 1992, 1996). A similar phenomenon has been reported in cloud chamber studies, where sulfide was observed to be oxidized to sulfate during ice crystal formation from expansion of droplets containing ionic salts (Finnegan et al., 1991; Finnegan and Pitter, 1991; Gross, 1991). Possible mechanisms are discussed by Finnegan and Pitter (1997).

d. Oxidation by 03 While the Henry's law constant for ozone is fairly small (Table 8.1), there is sufficient ozone present in the troposphere globally to dissolve in clouds and fogs, hence presenting the potential for it to act as a S(IV) oxidant. Kinetic and mechanistic studies for the O3-S(IV) reaction in aqueous solutions have been reviewed and evaluated by Hoffmann (1986), who shows that it can be treated in terms of individual reactions of the various forms of S(IV) in solution. That is, SO 2 9U 2 0 , USO3-, and SO 2- each react with O 3 by unique mechanisms and with unique rate constants, although in all cases the reactions can be considered to be a nucleophilic attack by the sulfur species on O 3. Figure 8.14 shows the proposed mechanisms of reaction for each species. The overall rate of the S(IV) oxidation can then be represented by

- d [ S ( I V ) ] / d t = {k0a 0 + kla 1 +

k20L2}[S(IV)][O3]

,

(R)

where a0, a~, and oL2 are the fractions of the total S(IV) in the form of SO 2 9H 2 0 , HSO3, and SO 2-, respectively. That is, a0 = [SO2" HzO]/[S(IV)] etc., where S(IV) is the sum of { S O 2 " H 2 0 q - H S O 3 +

312

s. I N O R G A N I C A N D O R G A N I C A C I D S IN T H E T R O P O S P H E R E

0 3 --

AquatedSO 2 M e c h a n i s m H

N

H20 9SO 2

H\

0

03

+

o / "'"o-- o oN o

o~s /

O--H

Ns/ "'"o o/ "'-.o_ o / 0 #'"~0

H- - 0

No

'NO/

I

H O

II

_

/

...

H.'~

O - - H ....

HO/~;NX o + o2 o H

..~ 03-

HSO3 M e c h a n i s m

H I

-o/~N o

-o/~N o

0

-0

0 H -0

N

O-- S

/

03

+

o\/

ON

-O

H

-o N

O--S--O

/

oN

N

HO

/o

. .o

O-~S / .o

....

.o,,,.

No

.

9. . . .

N~

k = k o + Fjlj,

s

o

-0

o

oTs/No

H .... O

o,

/o N

H ....0

2e-

N

N

O--S--O

o

/ \

O--O

0

0-

I

.o

+ 02

'o' o

0 3 -- SO 2- M e c h a n i s m

-ON

-O

-

0

/ -

S

+

__

-ON

/

0

/

0

_o o o Ns/No ....

-O = S \ _ ~

0

-o,,

o

/

r ~

/

203 + OH---, OH + 0 2 + 202, OH + HSO;- --, HSO 3 + O H - , followed by reactions of HSO 3 discussed later.

0

i

-

O o

-O S

T A B L E 8 . 8 R a t e C o n s t a n t s for R e a c t i o n s of 8 0 2 9H 2 0 ( a q ) , H S O 3 , a n d SO 2 - w i t h O 3 in A q u e o u s S o l u t i o n at R o o m T e m p e r a t u r e a Reacting species

I

=

where k 0 is the rate constant at zero ionic strength, Ij is the ionic strength, and Fj. is a parameter that reflects the nature of the electrolyte. Interestingly, the nature of the electrolyte was shown to be important, with Fj. for Na2SO 4 being about 4 times larger than that for NaC10 4 and that for NaC1 being about twice as large. Lagrange et al. suggest that a free radical mechanism suggested by Penkett et al. (1979) is consistent with their data, rather than the mechanisms proposed in Fig. 8.14, i.e.,

N~

O

/_~ j

No

0

0

=

/~

-O--S / No

=

~

2 SO 5

so4~--, so~- + so~

s o ; - + s o ~ - --, s o 4 +

SO~ + SO~ ~ 2SO 4 + 0 2 SO~- + HSO~- ~ HSO~- + SO~HSO 4 + SO 4 so4 + s o ~ - --, s o 3 + so4~S04 + H 2 0 ~ H S 0 4 + OH S 0 4 + HSO~- ~ H S 0 4 + S0~S04 + H202 ~ H S 0 4 + H02 SO 4 + NO~- ~ SO 2- + NO 3 HS04 + OH ~ H 2 0 + S04 SO 5 + HO 2 --+ HSO~- + 02 HSO~- + 0 2 ~ SO~- + HO 2

Termination s o 3 + s o 3 --, s 2 0 ~ s o 4 + so4 --, s 2 0 ~ SO 5 + SO 5 --> 82 O 2 - + 0 2

(1.5-2.5) • 109 5.5 • 105 2.1 • 105 (2.2-6) • 10 s (0.83-2.5) • 104 (0.034-7.5) • 104 5.5 • l0 s 440 (0.68-2) • 109 1.2 • 107 1.1 • 105 k = 4 . 5 • 105 Keq ~ 10 3 5 • 107 3 • 104

b Weinstein-Lloyd and Schwartz, 1991

1.8 • 10 s 4.5 • 10 s (0.48-1.4) • 10 s

Waygood and McElroy, 1992 b b

b b b b b b b b b

Wine et al., 1989 Wine et al., 1988 Huie, 1995

a All in L - m o l - s units. b See references and discussion in Sections B and C.3g.

Bromine undergoes reactions similar to those of chlorine in the aqueous phase (Table 8.15). However, there are some important differences. For example, the oxidation of bromide ion by 0 3, while not extremely fast, is still five orders of magnitude larger than that for the oxidation of CI-. As a result, 0 3 reacts both at room temperature and just below freezing with solutions containing bromide to generate Br 2 (e.g., Oum et al., 1998b; Hirokawa et al., 1998). However, no reaction occurs in the dark with chloride-containing aqueous particles (e.g., Oum et al., 1998a; Hirokawa et al., 1998). A second difference is in the reactions HOX-+

H+~

H 2 0 + C1.

(43, - 4 3 )

Although the forward reactions are very fast for both HOC1- and H O B r - , the reverse reaction for Br is very slow (1.4 s -1) compared to that of C1 (2.5 • 105 s-1). As a result, the equilibrium is far to the right for H O B r - even at high pH, whereas for HOCI-, it is to the left except at low pH. Hence generation of reactive chlorine in solution occurs only in acidic solu-

tions, whereas reactive bromine can be generated in neutral to somewhat basic solutions. This is important, because seawater and hence the sea salt droplet generated from it initially have a pH of approximately 8. A third factor comes into play in bromine chemistry, which is that atmospheric solutions containing bromide and chloride are most typically formed from seawater. Wave action generates small airborne droplets of seawater, which thus initially contain the elements in the ratios found in seawater. The molar ratio of Br- to CI- is ~ 1:650. However, despite the relatively small amounts of bromide relative to chloride, it plays a disproportionate role because of its reactivity and because its chemistry is closely intertwined with chloride ion chemistry. Table 8.16, for example, shows some of the interhalogen reactions of bromide and chloride. It can be seen that the chemistry preferentially generates Br 2 rather than C12. This is also true of the reactions of gaseous bromine compounds with solid chloride salts. For example, Mochida et al. (1998a) observed Br 2 and BrC1 as products of the reaction of gaseous HOBr with solid NaC1

320

s. INORGANIC AND ORGANIC ACIDS IN THE TROPOSPHERE TABLE 8.11

Some Aqueous-Phase HO x Chemistry k 29 8K

(L m o l - 1 s - l )

Reaction

Reference

O H + 0 3 --+ H O 2 + 0 2 O H + H O 2 --+ H 2 0 + 0 2 O H + 0 2 --+ O H - + 0 2 O H + H 2 0 2 --+ H 2 0 + H O 2 O H + O H --+ H 2 0 2 O H + OH---+ H 2 0 + O O H + H 0 2 --, H O 2 + 0 H 0H+H 2--, H 2 0 + H H O 2 + 0 3 --+ O H + 2 0 2 H O 2 + H O 2 ---> H 2 0 2 + 0 2 H O 2 + 0 2 --+ H O 2 + 0 2 H O 2 + H 2 0 2 --+ O H + 0 2 + H 2 0

1.1 • 10 s 1 • 101~ 1 • 10 l~ 2.7 • 107 5.5 • 109 1.3 • 101~ 7.5 • 109 4.2 • 107 < 1 • 104

N e t a et al., 1988 Elliott and Buxton, 1992 Elliott and Buxton, 1992 Buxton et al., 1988 Buxton et al., 1988 Buxton et al., 1988 Buxton et al., 1988 Buxton et al., 1988

8.3 • 10 5 9.7 • 107 0.5

Bielski, 1978 Bielski, 1978 Weinstein and Bielski, 1979

0 2 + 0 3 ---> 0 2 + 0 3

1.5 • 109

Sehested et al., 1983; Z e l l n e r and H e r r m a n n , 1995

O 3 + OH ~ 0 2 + HO 2 --+ O H - + O 3 O~ + 0 - ~ 20 2 0 3 ---> O - + 0 2 O f + H + --+ O H + 0 2 --+ H O 3 H O 3 --+ O H + 0 2 H++ 03

5.9 2.6 7.0 5.0 9.0 5.2 1.1 3.7

109 109 10 s 103 10 l~ 101~ 105 104

Buxton et al., 1988 Buxton et al., 1988 Buxton et al., 1988 N e t a et al., 1988 N e t a et al., 1988 Bfihler et al., 1984 Biihler et al., 1984 Bfihler et al., 1984

0 - + H 2 0 --+ O H - + O H 0 - + H z O 2 --+ 0 2 + H z O O-+HO 2 ~02 +OHH20 O- + O 2 >2 O H - + 0 2 0 - + H 2 --+ O H - + H

1.8 • 106 < 5 • 108 4 • 108

Buxton et al., 1988 Buxton et al., 1988 Buxton et al., 1988

6 • 108 8 • 107

Buxton et al., 1988 Buxton et al., 1988

OH-+

48 70 4.2 • 108

N e t a et al., 1988 N e t a et al., 1988 Sauer et al., 1984

H 0 2 + O 3 --+ O H + 0 2 + 0 2 H O 2 + O(3p) --+ O H + O 2

5.5 • 106 5.3 • 109

N e t a et al., 1988 Sauer et al., 1984

H + H O 2 --+ H z O 2

1.0 • 10 l~ 2.1 • 10 l~ 3.7 • 10 l~ 10 4 • 109 1.6 • 109

Buxton et al., 1988 Buxton et al., 1988 N e t a et al., 1988 Hartig and Getoff, 1982 Kl~ining et al., 1984 Sauer et al., 1984

03

+ H20

OH-+

fast> O H - +

OH + 0 2

O 3 --+ H O 2 + 0 2 --+ H O 2 + 0 2 O(3p) --+ H O 2

H + O 2 --+ H O 2 H+O 3--+OH+O 2 H + H 2 0 --+ H 2 + O H O(3p) + 0 2 ---> 0 3 O(3p) + H 2 O 2 --+ O H + H O 2

TABLE 8.12

Some Aqueous.Phase Carbonate/Bicarbonate k (L m o l - 1 s - l )

Reaction

H C O 3 + O H --+ H 2 0 + CO 3 O H + CO 2- --+ O H - + C O ~ CO 3 CO 3 COB COB

+ + + +

• • • • • • • •

0 2 ---> CO 2- + 0 2 H 2 0 2 --+ H C O ~ + H O 2 H 0 2 --* HC0~- + 0 2 O H ~ product

COB + O3 -+ CO 2- + 03

8.5 • 3.9 • 6.5 • 8.0 • 5.6 • 3 • 6 •

106 10 s l0 s 105 107 109 107

Chemistry

Reference

Buxton et al., 1988 Buxton et al., 1988 Eriksen et al., 1985 B e h a r et al., 1970 B e h a r et al., 1970 H o l c m a n et al., 1987 H o l c m a n et al., 1987

C. OXIDATION OF SO 2 T A B L E 8.13

32 1

Some A q u e o u s - P h a s e Chlorine C h e m i s t r y Rate c o n s t a n t k 29sK o r equilibrium constant

Reaction

C I - + 03 --+ C 1 0 - + 0 2 C I - + N O 3 --+ NO B + C1 C I - + OH --+ HOC1HOC1---+ C I - + OH H O C I - + H+--+ H 2 0 + C1 C1 + H 2 0 --+ H O C I - + H + C1 + C1 ~ C12 C1 + C I - ~ C12 C12 ~ C1 + C1C12 + C12 ~ C12 + 2C1C12 + H O 2 ~ 2C1-+ 0 2 + H + C12 + H 2 0 2 ~ 2C1-+ HO 2 + H + C12 + H 2 0 ~ 2C1-+ H + + OH C12 + OH ~ HOC1 + CIC12 + H ~ H + + 2 C I C12 + OH---+ H O C I - + CIC12 + H 2 0 ~ HOC1 + C I - + H + C12 + HO 2 + H + + C12 + 0 2 C12 + C1 - , C13 HOC1 + C I - + H + ~ C12 + H 2 0 HOCI+ O2 ~ CI-+ OH+ O2

and Hirokawa reaction

(L mol -1 s - l)

Reference

2 . 0 • 10 .3

Hoign6 et al., 1985; Neta et al., 1988 Exner et al., 1992 Jayson et al., 1973 Jayson et al., 1973 Neta et al., 1988; de Violet, 1981 Buxton et al., 1998 Neta et al., 1988 Buxton et al., 1998 Buxton et al., 1998 Neta et al., 1988 Neta et al., 1988 Neta et al., 1988 Buxton et al., 1998 Wagner et al., 1986 Buxton et al., 1988 Grigor'ev et al., 1987 Wang and Margerum, 1994 Bjergbakke et al., 1981 Bunce et al., 1985 Eigen and Kustin, 1962 Long and Bielski, 1980

1 • 10 7

4.3 • 10 9 6.1 • 10 9 1.45 • 101~ 2.5 • 105 8.8 X 10 7 8.5 • 10 9 6.0 • 10 4 2.0 • 10 9 (1-4.5) • 10 9 1.4 • 105 1.3 • 10 3 1.0 • 10 9 7.0 • 10 9 4.5 • 10 7 22 1.0 • 10 9 5.3 • lO s 1.8 • 10 4 7.5 • 10 6

et al. ( 1 9 9 8 ) o b s e r v e d o n l y B r 2 f r o m t h e

of 0 3 with a mixture

of NaBr

a n d N a C 1 at

Extensive

data

bases

for

kinetics

high relative humidity. Similarly, the reaction of gaseous

Standards

C12 w i t h n a t u r a l o r s y n t h e t i c s e a s a l t g a v e o n l y B r 2 as a

and by the University of Notre Dame

product (Mochida

ratory (see Appendix

et al., 1 9 9 8 b ) .

Excellent summaries

of aqueous-phase

chemistry are

f o u n d in H u i e ( 1 9 9 5 ) a n d Z e l l n e r a n d H e r r m a n n

T A B L E 8.14

(1995).

of

aqueous-phase

reactions are provided by the U.S. National Institute of a n d T e c h n o l o g y ( s e e N I S T , R o s s et al., 1 9 9 4 ) Radiation

Labo-

I V f o r W e b site).

A variety of models incorporating

the reactions of a

v a r i e t y o f i n o r g a n i c s a n d o r g a n i c s , in a d d i t i o n t o s u l f u r

A q u e o u s - P h a s e C h e m i s t r y Involving S(IV) and Reactive Chlorine and B r o m i n e Ions Rate c o n s t a n t k 29sK

Reaction

Chlorine chemistry C1 + C l - ~ C12 C12 + HSO~- ~ 2 C l - + SOB + H + C l - + SOg ~ SO 2- + C1 C I - + HSO~- -~ HOC1 + SO 2HOC1 + SO 2- -~ HSOg + C1Bromine chemistry b Br + B r - - , Br 2 Br 2 + HSO~- -~ 2 B r - + SOB + H + Br 2 + SO 2- -~ 2 B r - + SOB B r - + HSO~- -~ HOBr + SO 2HOBr + SO 2- + HeO -~ S 0 2 - + B r - + 2H +

(L mol -1 s - l)

Reference

8.5 • 109 3.4 • 108 2.5 • 10 s 1.4 • 10 .3 > 7.6 • 10 8a

Buxton et al., 1998 Huie and Neta, 1987 Wine et al., 1989; Huie and Clifton, 1990; Huie, 1995 Fortnum et al., 1960; Clarke and Radojevic, 1983 Fogelman et al., 1989

1.6 • 101~ 6.3 • 10 7 2.2 • 108 1 >5 • 109 a

Scaiano et al., 1992 Shoute et al., 1991 Shoute et al., 1991 Fortnum et al., 1960 Troy and Margerum, 1991

a Measured at pH of -~9-10; increases as pH decreases. See Vogt et al. (1996) for application to sea salt particles. b See also Mozurkewich (1995).

322

s. INORGANIC AND ORGANIC ACIDS IN THE TROPOSPHERE T A B L E 8.15

Reaction

B r - + 03 --+ B r O - + 0 2 BrO- + H + ~ HOBr H O B r + B r - + H+--+ Br 2 + H 2 0 B r - + O H --+ H O B r HOBr---+ B r - + O H --+ Br + O H H O B r - + H+--+ H 2 0 + Br Br + H 2 0 --+ H O B r - + H + H O B r + Br---+ Br 2 + O H Br + Br---+ Br 2 Br 2 + Br 2 --+ Br 2 + 2 B r Br 2 + H 2 0 --+ H O B r + B r - + H +

Some A q u e o u s - P h a s e B r o m i n e C h e m i s t r y Rate c o n s t a n t k (L mol - 1 s - l) or equilibrium c o n s t a n t Keq

Reference

1.6 • 102 Keq = 10 s's 1.6 • 10 l~ L 2 mo1-2 s-1 1.1 • 10 l~ 3.3 • 107 4.2 • 106 1.1 • 10 l~ 1.4s -1 Keq = 70 9 • 109 Keq = 1.1 • 105 2 • 109 1.1 • 102

Haag and Hoign6, 1983 Haag and Hoign6, 1983 Eigen and Kustin, 1962 Kl~ining and Wolff, 1985 Kl~ining and Wolff, 1985 Kl~ining and Wolff, 1985 Kl~ining and Wolff, 1985 Kl~ining and Wolff, 1985 Mamou et al., 1977 Nagarajan and Fessenden, 1985 Mamou et al., 1977 D'Angelantonio et al., 1988 Eigen and Kusten, 1962

compounds, have been developed for reactions in clouds and fogs. The reader is referred to papers by Seigneur and Saxena (1988), Pandis and Seinfeld (1989a, 1989b), Jacob et al. (1989), Lelieveld and Crutzen (1990), Walcek et al. (1990, 1997), Dennis et al., (1993), McHenry and Dennis (1994), Sander and Crutzen (1996), and Liang and Jacob (1997) for some examples of models used to examine the relative importance of various oxidation mechanisms. h. Effect of Droplet Size on S(1V) Oxidation

The chemical composition of fogs, clouds, and particles (see Chapter 9) varies as a function of particle size. For example, Figure 8.19 shows the concentrations of the major cations and anions measured in small and large cloud droplets at La Jolla peak in southern California (Collett et al., 1994, 1999). The large drops are enriched in soil and sea salt derived species such as Mg 2+, Ca 2+, Na+, and C1- whereas the smaller particles contain higher concentrations of sulfate and H +, i.e., have a smaller pH. In a given sample, the difference in pH between large and small drops was usually

TABLE 8.16

Reaction

H O B r + C1---+ BrC1 + O H BrC1 + H 2 0 H O B r + C I - + H + BrC1 + Br---+ Br2C1B r 2 C I - ~ Br 2 + C1BrC1 + C I - ~ BrCly BrC12 + H 2 0 H O B r + H + + 2C1-

less than one pH unit, although in some samples, the difference in pH was as much as 2. Similar results have been obtained at a variety of sampling sites and with different types of samplers (e.g., Bator and Collett, 1997), indicating this is a general phenomenon. These data suggest that the larger fog and cloud droplets are formed on the larger, coarse particles, which contain largely elements associated with mechanical processes such as erosion and sea spray (see Chapter 9.A). Similarly, the smaller fog and cloud droplets arise from the smaller, accumulation mode particles, which contain species such as sulfate and ammonium ions. It is somewhat surprising that the drops retain in part the chemical signatures of the nuclei on which they were formed, given the number of chemical and physical interactions associated with fogs and clouds as they form and dissipate. For oxidation of S(IV) by H 2 0 2 , which does not vary greatly over the pH range of atmospheric interest, such differences in droplet composition are not important. However, for the other oxidation mechanisms that do depend significantly on pH a n d / o r other solute con-

Some A q u e o u s - P h a s e B r o m i n e / C h l o r i n e C h e m i s t r y Rate c o n s t a n t k (L m o 1 - 1 s - l) or equilibrium c o n s t a n t Keq

Reference

")/icea = 0.001--0.25 k > 105s -1 Keq = 1.8 • 10-5 M 2 k > 10 s Keq = 1.8 X 104 M - 1 Keq = 1.3 M Keq = 6.0 M - 1 Keq = 3.0 X 10 -6 M 3

Abbatt, 1994; Kirchner et al., 1997 Wang et al., 1994 Wang et al., 1994 Wang et al., 1994 Wang et al., 1994 Wang et al., 1994 Wang et al., 1994 Wang et al., 1994

a Reaction probabilities for H O B r with chloride ion in ice; the larger values are for reaction with HC1 due to acid catalysis.


2000 ~" ,~

1500

"~

1000

g

8

NH 3. Winiwarter et al. (1992) show that deviations in the opposite direction can result for highly soluble species in a closed system. In short, the use of average bulk aqueous-phase composition to assess the importance of various oxidation pathways can lead to erroneous results if the composition of drops varies significantly with particle size.

I I Large drops (D > 23 gm) I-I Small drops (4 < D < 23 gm)

500 0

Na + NH4+ Mg2+ Ca 2+

Cl

NO 3" SO42

3 23

H+

F I G U R E 8.19 Concentrations of major anions and cations in coastal stratus clouds at La Jolla Peak, California, on July 6, 1993 (adapted from Collett et al., 1994).

centrations, e.g., by 0 2 and 03, these variations in composition with particle size can be important. Thus oxidation of S(IV) by O 2 and 0 3 would be expected to be inhibited in the more acidic droplets. In the vast majority of studies, the chemical composition of fogs and clouds is measured on bulk samples, i.e., collected without discrimination by size, and the relative importance of various oxidation mechanisms assessed based on this bulk composition. As discussed by Collett and co-workers (Collett et al., 1994; Bator and Collett, 1997), the use of the average cloudwater composition without taking into account the differences in chemistry in different drop sizes leads to an underestimation of the rate of S(IV) oxidation if the pH-dependent oxidation mechanisms contribute significantly. For example, model predictions by Gurciullo and Pandis (1997) suggest that the sulfate production could be underestimated by a factor of ~30 in the early stages of cloud formation. The size of the effect depends not only on the time but also especially on the concentrations of the oxidants H202 and O3. The reason is that at higher pH values, 0 3 can contribute significantly to the oxidation of S(IV) and this will therefore occur preferentially in the larger drops having lower concentrations of H +. On the other hand, if the H 2 0 2 concentration is large relative to 0 3, the oxidation by the peroxide overwhelms the ozone reaction; in this case, there is little effect of particle size since the H 2 0 2 oxidation is not highly dependent on pH. Using bulk average composition to assess the chemistry of fogs and clouds can be misleading in other respects as well. For example, Pandis and Seinfeld (1991, 1992) show that the bulk mixture formed from drops with different pH values that are each in equilibrium with the gas phase does not itself conform to Henry's law, with the bulk mixture being supersaturated with respect to species such as weak acids and

i. F o g - S m o g - F o g

Cycles

Studies of the chemical composition of fogs have shown that they can be quite acidic and contain high concentrations of other cations and anions as well. For example, pH values in the range of 2-4 are not uncommon, and values as low as 1.69 have been observed (e.g., see Munger et al., 1983; Hoffmann and Jacob, 1984; Jacob et al., 1985; Fuzzi et al., 1984; Muir et al., 1986; Hering et al., 1987; and Erel et al., 1993). The reason that the aqueous-phase concentrations in fogs can be so high is related in part to the liquid water content (LWC), which is a major difference between clouds and fogs. The liquid water content for fogs is typically of the order of 0.1 g m -3 of air, whereas that for clouds is about an order of magnitude higher. This small LWC in fogs corresponds to increased solute concentrations. A second reason for the higher concentrations in fogwater is the relatively rapid condensation and evaporation of water vapor onto preexisting particles as the fog forms and then dissipates. Figure 8.20 is a schematic of this process. Aerosol particles, which typically contain sulfate and nitrate as well as a variety of other species, act as condensation nuclei for the condensation of water as the relative humidity increases during fog formation. As more water condenses, the particle components become increasingly diluted. Simultaneously, the solution can take up gases such as SO 2, HNO3, and NH3, and chemical reactions such as the oxidation of S(IV) to sulfate can occur in the droplet. Both the initial sulfate in the particle on which water condensed and sulfate formed by oxidation of S(IV) in the fog aqueous phase can contribute significantly to the total sulfate deposition. For example, Bergin et al. (1996) estimate that in a typical Arctic radiation fog, preexisting sulfate aerosol accounts for about 60% of the sulfate deposition and oxidation of S(IV) in the aqueous phase the remaining 40%. Because these fog drops are much larger than the initial particles, gravitational settling is more important and acts to remove some of the particles (Pandis and Seinfeld, 1990). When the temperature rises, water begins to evaporate out of the fog droplet, concentrating the solutes. It is this process that likely led to the

324

s. I N O R G A N I C A N D O R G A N I C A C I D S IN THE T R O P O S P H E R E

Further ~-A..,.,.,~,.,l~k Growthby water evaporation ~ / ~--'~" ~ " ~ condensation f ret, t,u,~e~i ~on existing nuclei

/

_

.

) Di,ution

\

~/

Dissolution

of

.ConcentratlI~ ~ Fog Droplet ~l~gaseous components Dy evaporation ' ~ e.g., NOx, SO2, NH3

Temperature

Chemical conversions and deposition to surface

~( Relative humidity

FIGURE 8.20 Schematic of role of atmospheric aerosols in fog formation and evaporation (adapted from Munger et al., 1983).

extremely low observed pH value of 1.69. This mechanism is consistent with fog-smog-fog cycles that are commonly observed in urban areas (Munger et al., 1983; Pandis et al., 1990). [However, it is interesting to note that in some cities with relatively good air quality, reduced fog formation has been observed and attributed to increased temperatures via the heat island effect (Sachweh and Koepke, 1995).] Because such fogs can be so concentrated, there has been a concern regarding their effects on both human health and ecosystems. For example, Hoffmann (1984) has estimated the sulfate ion concentrations in the London fog during the 1952 smog episode. While the pollutants that were present and their concentrations are not well known, 4000 excess deaths occurred in a 5-day period (Chapter 1.B). Hoffmann and co-workers estimate that the sulfate concentration was approximately 11-46 milliequivalents per liter (mequiv L -1) during this air pollution episode. This is of the same order of magnitude as those measured in acid fogs in other areas. Hence the health effects of acids and associated species in fogs continue to be of interest.

4. Oxidation on Surfaces Adsorption of SO 2 on the surfaces of solids, followed by its oxidation on the surface, may provide a third route for the formation of sulfuric acid. The solid surfaces may be suspended in the gas phase or in atmospheric droplets, as discussed in detail by Chang and Novakov (1983). It is known that SO2 in air is oxidized on both graphite and soot particles, with water vapor enhancing

the sulfate formation on the surfaces (Hulett et al., 1972; Novakov et al., 1974; Halstead et al., 1990). Whether this surface oxidation is enhanced by the presence of oxidants other than air (e.g., 0 3 or NO 2) is not clear. For example, under certain conditions some researchers report an enhancement of the rate of uptake of SO 2 due to the presence of the copollutants 0 3 and NO 2 (Cofer et al., 1981; Britton and Clarke, 1980); however, others have not observed such an effect (Baldwin, 1982; Halstead et al., 1990), particularly at low gas concentrations more closely approximating those in the atmosphere (Cofer et al., 1984). Carbon surfaces of various types have been the subject of most studies. However, other types of surfaces, including alumina, fly ash, dust, MgO, V205, Fe203, and MnO2, have also been shown to oxidize SO 2 a n d / o r remove it from the gas phase (Hulett et al., 1972; Judeikis et al., 1978; Liberti et al., 1978; Barbaray et al., 1977, 1978; Halstead et al., 1990). As expected, the rate of removal depends on the nature of the particular surface, the presence of copollutants such as NO 2, and, as in the case of carbonaceous surfaces, the relative humidity. The increase with increasing water vapor suggests that oxidation of the SO 2 may occur in a thin film of water on the surface of the solid. Carbonaceous particles suspended in aqueous solutions can also act as sites for efficient SO 2 oxidation. For example, Novakov and co-workers have studied the kinetics of such processes with respect to the concentrations of 0 2, S(IV), and carbon particles as well as the pH and temperature dependencies (Brodzinsky et al., 1980; Chang et al., 1981; Benner et al., 1982).

C. O X I D A T I O N

Extrapolation of their results to atmospheric conditions led them to suggest that such heterogeneous reactions could be important in the aqueous-phase oxidation of S O 2.

It is difficult to evaluate quantitatively the importance of such heterogeneous reactions in the overall oxidation of S(IV). Their rates depend on the physical and chemical natures of the surfaces involved, including specific surface areas, the presence of defects and surface adsorbed water, etc., yet these are not well understood, especially for highly complex environmental gas-liquid-solid systems. For example, the rates of oxidation of SO 2 at 80% relative humidity on two different samples of fly ash obtained from two coal-fired power plants differed by more than an order of magnitude (Dlugi and Gusten, 1983). Even in laboratory systems the nature of relatively simple surfaces such as carbon depends on the history of the material. However, the available evidence suggests that such heterogeneous reactions should be considered as potential contributors to the overall oxidation of S(IV), especially close to sources where particle concentrations and hence available surface areas are relatively high (see Chapter 9). Baldwin (1982), for example, estimates that the loss of gaseous SO 2 due to interactions with particle surfaces at a particle density of 100 /xg m -3 could be as high as 1% h -1. Such rates may not be sustained for long periods of time due to saturation of the surface, if the surface is dry. However, Chang and Novakov (1983) point out that when the surface is wet, the active sites are constantly regenerated as the sulfate formed on the surface dissolves in the surrounding aqueous solution; given the availability of water in the atmosphere, it seems unlikely that any surfaces are really "dry." Herring et al. (1996) have reported evidence for the heterogeneous oxidation of both SO 2 and NO X on the surfaces of soil dust particles in the smoke plume from the 1991 Kuwait oil fires. The rate of SO 2 oxidation was estimated to be 6.5% h-1. In addition to the dark oxidation of S(IV) on surfaces, there may be photochemically induced processes as well. For example, irradiation of aqueous suspensions of solid ce-Fe20 3 (hematite) containing S(IV) with light of A > 295 nm resulted in the production of Fe(II) in solution (Faust and Hoffmann, 1986; Faust et al., 1989; Hoffmann et al., 1995). This reductive dissolution of the hematite has been attributed to the absorption of light by surface Fe(III)-S(IV) complexes, which leads to the generation of electron-hole pairs, followed by an electron transfer in which the adsorbed S(IV) is oxidized to the SO~" radical anion. This initiates the free radical chemistry described earlier.

OF SO 2

325

The photochemical reactions of such semiconductors in aqueous solutions may also influence S(IV) oxidation via the production of H 2 0 2. The possibility of such photoassisted surface reactions in the atmosphere was first examined by Calvert in 1956. For example, F e 2 0 3 particles suspended in a bisulfite solution in the presence of 0 2 rapidly oxidize the S(IV) to S(VI) in the presence of light, whereas no oxidation occurs in the dark (Frank and Bard, 1977). It has been suggested that this is due to the absorption of light by F e 2 0 3 and the migration of electrons in the conduction band to the particle surface where they react with 0 2 and H § to form H202: 02

-Jr-

2e~rface

-Jr- 2H +

surface>

H 202.

(44)

Similar reactions have been observed more recently on a variety of solids, including ZnO, TiO2, and desert sand; in the presence of organics, organic peroxides have also been observed (e.g., see Kormann et al., 1988). For a review of this area, the reader is referred to the review by Hoffmann et al. (1995). 5. R e l a t i v e I m p o r t a n c e P a t h w a y s for SO 2

of V a r i o u s O x i d a t i o n

Figure 8.21 shows one estimate of the relative importance of the oxidation of S(IV) by 0 3 and H202, the Fe- and Mn-catalyzed 0 2 oxidation, and the oxidation on a carbon surface at concentrations typically found in the atmosphere (Martin, 1984; Martin et al.,

103

_ ~

03' 50ppb

/

102 "7

o~ tO -.,,~ o ." m X

10 1

o

~10-1 m O rr

10-2

10-3 10-4

0

/

I /

1

I

2

I

I

3

4

I

5

I

6

I

7

pH FIGURE 8.21 Estimated rates of oxidation of S(IV) in a hypothetical cloud with liquid water content of I mL m-3 (i.e., LWC of 1 g m -3) based on 5 ppb gaseous SO2 as a function of pH (adapted from Martin, 1984; Martin et al., 1991).

326

s. INORGANIC AND ORGANIC ACIDS IN THE TROPOSPHERE

1991). It was assumed that there were no limitations on the rates of oxidation due to mass transport; as discussed in detail by Schwartz and Freiberg (1981), this assumption is justified except for very large droplets (> 10 /xm) and high pollutant concentrations (e.g., 0 3 at 0.5 ppm)where the aqueous-phase reactions are very fast. It was also assumed that the aqueous phase present in the atmosphere was a cloud with a liquid water content (V) of 1 g m -3 of air. As seen earlier, the latter factor is important in the aqueous-phase rates of conversion of S(IV); thus the actual concentrations of iron, manganese, and so on in the liquid phase and hence the kinetics of the reactions depend on the liquid water content. Only the oxidation by H 2 0 2 is relatively independent of pH. This arises because the rate coefficient for the reaction and the solubility of S(IV)show opposite trends with pH (Fig. 8.9b). For the other species, the effects of the S(IV)solubility and the pH dependence of the kinetics work in the same direction (Fig. 8.9a), leading to a strong overall dependence on pH. The uncatalyzed oxidation of 0 2 is not shown, because it is generally believed to be unimportant compared to the other mechanisms in real atmospheric droplets containing "impurities" such as metals that will act as catalysts. The estimates in Fig. 8.21 show that H 2 0 2 is expected to be the most important oxidant for S(IV) in clouds and fogs at pH < 4.5. At higher pH values, both O3 and the iron-catalyzed O2 oxidation can compete. Figure 8.22 shows an estimate of the contributions to the oxidation of S(IV) by H 2 0 2, by the iron-catalyzed O 2 oxidation, and by OH in both the gas and aqueous phases of a cloud (Jacob et al., 1989). It is seen that H 2 0 2 and the iron-catalyzed process predominate at night, but the gas-phase oxidation by OH becomes significant during the day when it is formed by photochemical processes. On the other hand, the contribur 0 .m 4-. r ::3

~,

8-

.

.

Fe(lll)\

.

.

.

.

0 Q..

..o m g "6 =,E. tl:i

13:

2 -

f"_'~.__ "'"/.

00-- ........ ~ ~ - ~ . : t 6

12

OH (aq) < ......""-

18

24

Time (h) F I G U R E 8.22 Calculated contributions to S(IV) oxidation in a cloud of the iron-catalyzed oxidation by O2/Fe(III), by H 2 0 2 and OH in solution, and by OH in the gas phase, OH(g), expressed in terms of rate of production of column S(IV) (adapted from Jacob et al., 1989).

tion from oxidation by OH in the aqueous phase is relatively small.

D. ORGANIC

ACIDS

While the focus in terms of acid deposition has been on sulfuric and nitric acids, it has been increasingly recognized that organic acids can also contribute significantly to the acidity of both the gas and aqueous phases in both urban and remote regions. A review of carboxylic acids in the atmosphere is given by Chebbi and Carlier (1996). The major organic acids found in the gas phase are formic acid (HCOOH) and acetic acid (CH3COOH), with smaller contributions from larger aliphatic acids and multifunctional acids such as pyruvic acid (CH 3COCOOH) and glyoxalic acid (COOHCHO) (e.g., see Kawamura et al., 1985; and Khwaja, 1995). Formic and acetic acids have sufficiently high vapor pressures that they are found almost totally in the gas phase rather than in particles (e.g., see Kawamura et al., 1985; Talbot et al., 1988; Grosjean, 1989; and Khwaja, 1995). Concentrations of HCOOH in urban areas are typically a few ppb (e.g., see Dawson et al., 1980; and summaries by Lawrence and Koutrakis, 1994; and Khwaja, 1995) although levels up to 45 ppb have been observed (Lawrence and Koutrakis, 1994). CH3COOH is also present in urban areas in the low-ppb range, with reported concentrations as high as 15 ppb in the Los Angeles area (Grosjean, 1989). In semirural and rural areas, the concentrations of HCOOH and CH3COOH are somewhat smaller but still tend to be in the range around a ppb (e.g., Sanhueza et al., 1996; Kumar et al., 1996; Granby et al., 1997a, 1997b). For example, at a rural site in Virginia, concentrations of formic and acetic acids are typically ~1 and 0.5 ppb, respectively, although peak formic acid concentrations as high as 10 ppb have been measured (Talbot et al., 1995; Keene et al., 1995). In the northern Congo, ground-level concentrations of ~0.5 ppb each have been measured for these two acids, but the levels were as high as ~ 3 - 4 ppb in the boundary layer above the surface (Helas et al., 1992). In marine remote regions, concentrations of formic and acetic acids are typically about 0.1-0.3 ppb, although much higher levels have been observed in stable layers of air at higher altitudes above the ocean (e.g., Chapman et al., 1995). Formic and acetic acids may constitute a large fraction of the gas-phase acidity. For example, Grosjean (1990) measured the concentrations of these two organic acids as well as the inorganic acids HNO 3 and HC1 in southern California during a smog episode in

D. ORGANIC ACIDS

[] HNO 3 9 HCI [] HCOOH [] CH3COOH

..El EL vo_ 8 - (O

O C O

6-

o

~ I

i --

I/I I/I

I1)

I/I

_>, 2 t"

I/I

"4-'-' C O

May

dun

Jul

Aug

Sep

Oct

Nov

Dec

Month

F I G U R E 8.23 Average concentrations of gas-phase acids at eight sites in southern California in 1986 (adapted from Nolte et al., 1997).

August 1986. The two organic acids taken together represented 44-93% of the total gas-phase acids. Similarly, Fig. 8.23 shows the monthly average concentrations of HCOOH, CHsCOOH, HC1, and HNO 3 measured at eight sites in southern California in 1986 (Nolte et al., 1997). Formic and acetic acids are seen to be the major acids, consistently exceeding HNO 3. There are a number of potential sources, both primary and secondary, of formic and acetic acids in urban areas (e.g., see review by Chebbi and Carlier, 1996). Clearly the oxidation of organics can lead to the formation of acids. For example, as discussed in Chapter 6.E.2, the reaction of 0 3 with alkenes generates a Criegee intermediate that can react with water vapor to generate a carboxylic acid. Aqueous-phase oxidations, e.g., of HCHO to HCOOH, followed by evaporation from the condensed phase may also potentially contribute. Direct emissions (i.e., primary sources) include automobile exhaust, biomass combustion, stationary source emissions, for example, from vinegar manufacturing and cooking of food, and possibly natural emissions from vegetation (e.g., see Kawamura et al., 1985; Talbot et al., 1988, 1995; Dawson and Farmer, 1988; and Keene and Galloway, 1988). The relative contributions of these sources depends, not surprisingly, on the particular location. For example, Granby et al. (1997a, 1997b) found that the concentrations of H C O O H and C H s C O O H were similar in central Copenhagen and at a semirural site 30 km away, where the NOy concentrations were an order of magnitude smaller. In addition, they correlated with the photochemical species such as 0 3, suggesting that the acids were formed by chemical reactions (e.g., of 0 3 with alkenes) during long-range transport. In southern California, these two acids were correlated with elemental carbon, used as an indicator of primary emissions, at upwind site close to the coast but not at

327

downwind sites, suggesting that local primary emissions dominated upwind (Nolte et al., 1997). Conversely, the correlation between formic and acetic acids with both 0 3 and particle sulfate increased downwind, suggesting a significant contribution to their formation from chemical reactions both in the gas phase and in solution (where most of the sulfate is formed). Similarly, in Venezuela, atmospheric oxidation of organics appears to be the major source of these acids, perhaps with some contribution from emissions from soils during the dry season (Sanhueza et al., 1996). Grosjean (1989, 1992) estimated that about 40% of the H C O O H and ~ 75% of the C H s C O O H measured at Claremont, just east of Los Angeles, are due to direct emissions. This is qualitatively consistent with carbon isotope measurements of is C/12C of formic and acetic acids in rainwater collected in this area, which show H C O O H is from a combination of direct emissions and secondary oxidation chemistry and C H s C O O H is primarily from direct emissions (Sakugawa and Kaplan, 1995). A variety of dicarboxylic acids have been measured in air, including, for example, oxalic acid [(COOH)2], succinic acid [HOOCCH2CH2COOH], and malonic acid [HOOCCH2COOH], as well as larger straightand branched-chain carboxylic acids; unsaturated and aromatic acids such as phthalic acid are also observed in smaller concentrations (e.g., see Kawamura et al., 1996a, 1996b). Because of their lower vapor pressures, they are found predominantly in particles (see Chapter 9). As discussed in Chapters 6 and 9, these dicarboxylic acids are believed to result in part from the oxidation of organics in air. However, Kawamura and Kaplan (1987) have also shown that automobile exhaust can be a significant source as well. Organic acids can be removed by reaction with OH (see Chapter 6.H) as well as by wet or dry deposition. As a result, these acids are a common component of rain, clouds, fogs, and dews as would be expected from their large Henry's law constants, ~103-104 mol L -1 atm -1 (see Keene et al., 1995), and are found in the condensed phase from remote to highly polluted urban areas (e.g., see Norton, 1985; Keene and Galloway, 1986, 1988; Likens et al., 1987; Weathers et al., 1988; Muir, 1991; Sakugawa et al., 1993; Keene et al., 1995; and Khwaja et al., 1995). In short, while the focus has been primarily on sulfuric and nitric acids as a source of acid deposition, it is clear that organic acids can also contribute significantly. The gas-phase concentrations of the simplest carboxylic acids, formic acid and acetic acid, are relatively high even in remote regions, of the order of a ppb. Both natural and anthropogenic sources have been

328

s. INORGANICAND ORGANICACIDS IN THE TROPOSPHERE

proposed, but the nature of the individual sources and their relative contributions are not well established.

E. O X I D A T I O N OF SULFUR COMPOUNDS OTHER THAN

SO

2

As discussed in Chapter 2 and in more detail in Chapter 11, a variety of organic sulfur compounds in addition to inorganics such as H2S and COS are emitted by biological sources. In the troposphere, they may ultimately be oxidized to SO 2 and H2SO 4. However, the chemistry of these compounds tends to be complex, and a variety of partially oxidized sulfur compounds is formed first. Table 8.17 summarizes the rate constants and estimated tropospheric lifetimes of some of these sulfur compounds with respect to reaction with O H or NO 3. The assumed concentrations of these oxidants chosen for the calculations are those characteristic of more remote regions, which are major sources of reduced sulfur compounds such as dimethyl sulfide (DMS). It is seen that O H is expected to be the most important sink for these compounds and that NO 3 may also be important, for example, for DMS oxidation (see also Chapter 6.J). As discussed in Chapters 6 and 7, NaC1 and other chloride salts in airborne sea salt particles are now believed to react to generate photochemically active compounds such as C12 and C1NO 2 that photolyze to form chlorine atoms. Peak concentrations of atomic chlorine as high as 104-105 cm -3 in the early morning hours have been predicted (Keene et al., 1993, 1996; Pszenny et al., 1993; Wingenter et al., 1996; Singh et al., 1996; Tuckermann et al., 1997; Spicer et al., 1998). Chlorine atoms also react rapidly with reduced sulfur compounds and hence are potential oxidants for these organics in the early morning hours in coastal regions.

TABLE 8.17

Finally, there has been a great deal of interest in the halogen oxides IO, BrO, and C10 as potential oxidants for organic sulfur compounds such as dimethyl sulfide. We therefore also discuss the current status of these reactions. Dimethyl sulfide (DMS) is of particular interest since it is the major reduced sulfur compound produced in oceanic areas. In addition, as discussed in Chapter 14, it is believed to play a significant role in global climate issues; thus, its oxidation to SO 2 is expected to lead to the subsequent formation of sulfuric acid and sulfate particles by the chemistry that is the focus of this chapter. In addition, as we shall see, another oxidation product of DMS is methanesulfonic acid (CH3SO3H), which is highly soluble and has a sufficiently low vapor pressure that, like sulfuric acid, it exists primarily in the condensed phase (e.g., Ayers et al., 1980; Clegg and Brimblecombe, 1985; Kreidenweis and Seinfeld, 1988a, 1988b). As a result, oxidation of DMS can generate condensation nuclei on which water condenses to form clouds. These clouds can scatter radiation back to space, lowering the surface temperature and hence the production of DMS, forming a feedback loop (Charlson et al., 1987). Because of this potential feedback, there have been many recent studies examining the relationship between DMS, its oxidation products, and condensation nuclei (e.g., see Kreidenweis et al., 1991; Berresheim et al., 1993; Pandis et al., 1994; Ferek et al., 1995; Clarke et al., 1996; and Bandy et al., 1992, 1996). However, it is clear that understanding such feedbacks requires understanding the chemistry of DMS and, in particular, its role as a source of low-volatility products that can act as condensation nuclei. As a result, we focus here on what is known about the tropospheric chemistry of DMS. As we shall see, the chemistry of even this relatively simple compound is complex, and much remains to be learned about its reaction mechanisms. For larger reduced sulfur com-

Rate Constants and Lifetimes at Room Temperature for the Reactions of OH and N O with Some Reduced Sulfur Compounds Emitted Biogenically O H (day)

Compound

CH3SCH 3 CH3SSCH3 CH3SH H2S CS2 COS

k (cm 3 m o l e c u l e - 1 s - 1)a

6.5 • 2.3 • 3.3 • 4.8 • 4.7 • 2.0 •

10-12 10-10 10-11 10-12 10-12 10-15

N O 3 (night)

'rOHb 2 days 1.2 h 8h 2.4 days 2.5 days 16 yr

k

(cm 3 m o l e c u l e

1.1 • 0.7 • 0.92 • 5.4 days (Patroescu et al., 1996). Thus, in low-NO X environments, the CH3SCH2OO radical will react with HO 2 or other R O 2 radicals:

The fate of the CH3S radical in the atmosphere is not clear but may include reaction with 02, 03, or NO 2. Table 8.18 summarizes the rate constants for these reactions at 298 K and the corresponding lifetimes under typical tropospheric conditions. Although only an upper limit can be placed on the rate constant for the 0 2 reaction, it may still be the predominant reaction of CH3S because of the large oxygen concentration in the atmosphere. However, this process forms a weakly bound adduct with a C H 3 S - O O bond energy of about 11 kcal mo1-1, and decomposition back to reactants also occurs: CH3S+O 2+M~CH3SOO+M.

(53,-53)

Evidence for formation of this adduct has been obtained in laboratory studies between 216 and 258 K, where CH3S is observed to come to equilibrium in the presence of 0 2 (Turnipseed et al., 1992). A contribution from the back reaction is difficult to avoid in experimental systems, making measurements of the true forward rate constant somewhat uncertain. Extrapolation of the measured kinetics to 298 K suggests that approximately 30-75% of the CH3S would be in the form of the adduct at typical tropospheric temperatures of 298-275 K and 1 atm pressure (Turnipseed et al., 1992).

Removal Rates of C H 3 S in the Troposphere at 298 K a k ( 2 9 8 K) (cm 3 m o l e c u l e - 1 s - l)

[Reactant] (molecules c m - 3)

x(s)

< 6.0 • 10-18 5.4 • 10-12 5.8 • 10-11

5.2 • 1018 9.8 • 1011(40 ppb) 2.5 • 10l~ (1 ppb) 2.5 • 108 (10 ppt)

>0.03 0.19 0.69 69

a Rate constants from Tyndall and Ravishankara (1991), Domin6 et al. (1992), and Turnipseed et al. (1993).

331

E. OXIDATION OF SULFUR COMPOUNDS OTHER THAN SO 2

Other products?

(57d)

The fate of the CH3SOO adduct is not known but by analogy to other peroxy radical reactions is expected to include reactions with NO and NO 2 (Turnipseed et al., 1993):

k57(298 K ) = 6.0 • 10 -13 cm 3 molecule -1 s -1

CH3SOO + NO ~ CH3SO + N 0 2 ,

CH3SO + NO 2 ~ CH3SO 2 (or CH 3 + 802) -Jr-NO,

(54)

(Domin6 et al., 1992)

(58)

k54 (227-256 K) = 1.1 • 10-11 cm 3 molecule- 1 s- 1

kss (298 K) = 1.2 • 10-11 cm 3 molecule- 1 s- 1

CH3SOO + NO 2 C H 3 S O O N O 2 (or CH3SO + NO3),

(55)

k55 (227-246 K) = 2.2 • 10-11 cm 3 molecule- 1 s- 1. Another alternative is isomerization to CH3802 (Turnipseed and Ravishankara, 1993). The data in Table 8.18 show that reactions of CH3S with NO 2 and O 3 may also be important. This also appears to be the case with larger thio radicals such as C2H58 (Black et al., 1988). The reaction with NO 2 produces primarily CH3SO + NO (Barnes et al., 1987; Hatakeyama, 1989; Tyndall and Ravishankara, 1989; Domin6 et al., 1990); a minor addition channel produces CH3SNO2, which has been observed in laboratory systems using FTIR (Barnes et al., 1987). CH3SO has been observed as a product of the reaction of CH3S with O 3 (Domin6 et al., 1992), suggesting that one channel is an oxygen atom transfer: CH3S + O 3 ~ CH3SO + O 2.

(56a)

The yield of CH3SO at low (Torr) pressures is only 15%. However, the yield has not been determined at 1 atm. Since (56a) is highly exothermic ( ~ 59 kcal tool-l), the CH3SO may contain sufficient energy to decompose at low pressures to CH 3 + SO; quenching of an excited CH3SO at 1 atm may lead to much greater yields of this radical. Other potential channels that are exothermic are the following: C H 3 S + O 3 ~ CH 3 + S O +

O 2,

(56b)

CHzSO + H + O2,

(56c)

CHzSO

(56d)

CHzS

+ HO2, + OH

+ 02,

C H 3 0 + SO 2.

(56e) (56f)

Although Domin6 et al. (1992) place an upper limit of 4% on the contribution of (56e), Barnes et al. (1996) point out that even a small production of CHzS followed by its oxidation could be responsible for the small yields of COS they observe by FTIR. The CH3SO radical will be further oxidized under tropospheric conditions" CH3S O + O3 13%> CH3S + 202,

(57a)

CH3SO 2 (or CH 3 + 802) -Jr-02, (57b) CH2SO 2 + H

+ 02,

(57c)

(Mellouki et al., 1988; Hatakeyama, 1989; Domin6 et al., 1990). There is also evidence for an addition reaction of CH3SO with O 2 (e.g., Hatakeyama et al., 1989; Barone et al., 1995): CH3SO+O 2+M~CH3S(O)OO+M.

(59)

For example, methylsulfinyl peroxynitrate, CH3S(O)OONO2, has been observed by FTIR (e.g., see Barnes et al., 1987; Hatakeyama, 1989; and Jensen et al., 1992), presumably from reaction (59) followed by reaction of the peroxy radical formed with NO 2. By analogy to other reactions, there are several other fates of the CH3S(O)OO radical formed in (59) that might be expected to be important. These include decomposition or reaction with NO or other peroxy radicals (Barnes et al., 1987): C H 3 S ( O ) O O ~ CH 3 + SO 3,

(60)

C H 3 S ( O ) O O + NO ~ CH3S(O)O + NO 2. (61) The SO 3 formed in (60) will react with H 2 0 to form H z S O 4 as described in detail earlier in this chapter, whereas CH3S(O)O formed in (61) may decompose or react further. The CH3SO 2 radical has a number of potential fates, including decomposition to CH 3 + SO 2 and reactions with NO2, O2, and 0 3" CH3SO 2 + M ~ CH 3 + SO 2 + M,

(62)

k62(298 K) - 510 s -1 at 1 Torr pressure (Ray et al., 1996) CH3SO 2 + NO 2 ~

CH3SO 3 + NO,

(63)

k63(298 K) = 2.2 • 10 -12 cm 3 molecule -1 s -1 (Ray et al., 1996). Only upper limits for rate constants for the reaction of CH3SO 2 with O 2 and 0 3 of < 6 • 10 -18 and < 8 • 10-13 cm 3 molecule-1 s-1 have been measured (Turnipseed and Ravishankara, 1993). I f the rate constant for decomposition measured at 1 Torr is the same at 1 atm, the decomposition is expected to predominate at 298 K. As discussed by Ray et al. (1996), as the

332


temperature decreases, the rate of the thermal decomposition will also decrease and some of the other reactions may become competitive. However, Patroescu et al. (1999) observed in FTIR experiments in 1 atm air that the yield of methanesulfonic acid increased as the NO X concentration increased. They attributed this observation to the addition of 0 2 to CH3SO 2, with secondary reactions involving NO leading to the formation of methanesulfonic acid, CH3SO2OH. Thus, at 1 atm in air, the reaction of CH3SO 2 appeared to be much faster than its t h e r m a l decomposition. R e a c t i o n of the CH3SO2(OO) adduct with NO 2 was postulated to give C H 3 S O 2 O O N O 2 , methanesulfonyl peroxynitrate, which was observed by FTIR (Patroescu et al., 1999). The reaction of OH with dimethyl sulfoxide also appears to occur primarily by addition. Urbanski et al. (1998) have reported yields of CH 3 in this reaction of unity and suggest that the reaction is O H + CH 3S(O)CH 3

[ C H 3 S ( O ) ( O H ) C H 3 ]* methanesulfinic acid CH 3 + CHBS(O)OH.

(64)

The rate constant at 298 K, based on the production of CH 3, was measured to be k64 = (8.7 ___ 1.6) • 10 -11 cm 3 molecule-1 s-1 (Urbanski et al., 1998). The formation of methanesulfinic acid in the OH-dimethyl sulfide system has been reported by SOrensen et al. (1996), who proposed that it resulted from secondary reactions of the O H - D M S adduct rather than reaction (64). Further addition of OH to methanesulfinic acid followed by reaction with 0 2 may then form methanesulfonic acid, CH3S(O)(O)OH, which in the atmosphere is rapidly scavenged into particles (e.g., Clegg and Brimblecombe, 1985). The adduct may also react with 0 2 and, by analogy to the O H - C H 3 S C H 3 adduct, is expected to give HO 2 + CH3S(O)(O)CH3, dimethyl sulfone. The oxidation of dimethyl sulfide (DMS) to dimethyl sulfoxide (DMSO) and the subsequent oxidation of the latter to methanesulfonic acid (MSA) have been observed in field studies. For example, one study in Antarctica which focused on the chemistry of dimethyl sulfide (Berresheim and Eisele, 1998) measured not only DMS but also a variety of its oxidation products, including DMSO, MSA, and dimethyl sulfone (Berresheim et al., 1998). The measured concentrations of DMSO were in agreement with model results if 80-100% of the OH + DMS reaction gave DMSO; as discussed earlier, the addition channel that leads to DMSO becomes relatively more important at the lower temperatures found in Antarctica. Furthermore, the

MSA concentrations were generally consistent with its formation from the O H - D M S O reaction if 80% of the reaction generated methanesulfinic acid and 50% of this compound was oxidized to MSA (Davis et al., 1998). Figure 8.25 summarizes the current understanding of the mechanism of DMS oxidation in the troposphere. For reviews, see Plane (1989), Turnipseed and Ravishankara (1993), Barnes et al. (1993, 1996), Barone et al. (1995), and Berresheim et al. (1995). b. Reaction with the Nitrate Radical (NO 3)

The nitrate radical is also known to react rapidly with DMS: CH3SCH 3 + NO 3 ~ CH3SCH 2 + HNO 3, (65) k65 = 1.1 • 10-12 cm 3 molecule-1 s-1 (Atkinson et al., 1997a). In contrast to the OH reaction, which occurs during the day due to the photolytic sources of OH, reaction (65) is a nighttime reaction due to the rapid photolysis of NO 3 at dawn. It is interesting that while the ouerall reaction appears to correspond to a hydrogen atom abstraction rather than addition to the sulfur atom, the mechanism is believed to involve the initial formation of an adduct, followed by its decomposition to HNO 3 and the alkyl radical shown in reaction (65) (Jensen et al., 1992; Daykin and Wine, 1990; Butkovskaya and Le Bras, 1994). NO 2 is not formed in the NO3-DMS reaction (Dlugokencky and Howard, 1988), ruling out an oxygen atom transfer either directly or via decomposition of the adduct. c. Oxidation by Chlorine Atoms

Atomic chlorine reacts rapidly with DMS, with an overall rate constant of (3.3 _+ 0.5) • 10 -l~ cm 3 molecule -1 s -1 at 298 K and 700 Torr total pressure (Stickel et al., 1992). As is the case for the OH reaction, the chlorine atom reaction proceeds by two reaction channels, one an abstraction and the other addition to the sulfur atom: C1 + CH3SCH 3 ~ HC1 + CH3SCH2, M

CH3S(C1)CH 3.

(66a) (66b)

A potential reaction path producing CH3S + CH3C1 can be ruled out based on the observation of very small yields of CH3C1, only 0.13% at 1 atm pressure (Langer et al., 1996), and Zhao et al. (1996) show a path producing CH 3 is also not important. Stickel et al. (1992) suggest that at 298 K and 1 atm pressure, the


two paths, (66a) and (66b), are about equally important. The fate of the C1-DMS adduct is not known.

concentration of IO radicals is not well known, as it has only just been directly detected and measured in the troposphere at concentrations up to 6 ppt at a coastal site (Alicke et al., 1999). Using 108 cm -3, one can calculate a lifetime for DMS with respect to reaction (68) of about 10 days, too slow to compete with OH, NO3, and C1 atoms. However, with the recent recognition of the potential importance of atomic chlorine and bromine under certain conditions in the Arctic at polar sunrise (e.g., see Barrie et al., 1988; and Niki and Becker, 1993), the potential for BrO and C10 chemistry has been reconsidered. As described in Chapter 6.J.4, at polar sunrise there is a rapid loss of ground-level 0 3 that appears to be associated with reaction with atomic bromine and at the same time, there is evidence that chlorine atoms are playing a major role in the organic removal (Jobson et al., 1994). This is consistent with reactions of sea salt particles generating atomic bromine and chlorine, although the exact nature of the reactions and halogen atom precursors remains unknown.

d. Oxidation by Halogen Oxides: I0, BrO, and ClO

Halogen oxides are also potential reactants with DMS in the marine boundary layer (e.g., see Barnes et al., 1989). As discussed in Chapter 12, CH3I as well as CH 3C1 and CH 3Br have natural oceanic sources. While CH3C1 and CH3Br do not absorb light significantly in the actinic region, CH 3I and other alkyl halides do (see Chapter 4.V), photolyzing to form iodine atoms: CH3I + h u ~ CH 3 + I.

(67)

The iodine atoms react with 03, forming the IO radical, which can potentially oxidize DMS to DMSO, regenerating iodine atoms in a chain process: IO + CH3SCH 3 ~ I + C H 3 S ( O ) C H 3.

333

(68)

The current recommendation for k68 at 298 K is 1.2 • 10 -14 c m 3 molecule- 1 s- 1 (Atkinson et al., 1997a). The

O

II CH3-1SlOH3 + HO2 o

dimethyl l sulfone

o ,

OH

CH3SCH3 ~

02

o "

;

o ,

(CH3-S-CH3)

dimethyl sulfoxide

9 CH3 + CH3 SOH

OH

HO2-~.%1

CH3S~O%2

OH

OH

02 ~CH H,.O

"

CH3+ SO2

~

v

1"O2

3

~ ~

~

cmscmooH CmSOHO + o=

sulfonic acid

~

IO2, M

~ r

+ H=O + O=

"

O

.. ~.N (slow)

U2 > 2

CH3SC. O

NO2 NO

~

~

J

'"-'2

~

/

~O 3

methyl thioformate

~

O

.,~H3S~ + NO2

CH3SOO

2

II

methane

p,.r~ r,.,-~ 4 HCHO + CH3S /'-N. 99CH3SO ") " ~ "~2 /' / ~ SCH CO ~ OH SCH O ~NO2 ~10 /

3 I

o

9 CH3 SOH.

methane sulfinic acid

/ ~(cH3S-OH3)

/ CH3SCH3 _

OH, 02

"~

t ~ ~

/ O3

/

"IX, /

1

~

NO2 NO

t

~ O2,0ther

t

90H3SO2

~~ / ,y 0302

X

products/

\

channels

/

\ 02

X

I ~' CH3SO+ 02 Other products

OH3 + SO3 .{

NO/

~2 M " '

O

/

/

9CH3SO3

~ /~'

CH3;OI~ tt

O

CH3SOO II O

NO2 ~r~

/

o

II

CHalSI

/

9 OH3 + SO2 etc

o

F I G U R E 8.25 Overview of oxidation of DMS by OH in the troposphere (note that many of the reactions after the first step are the same in DMS reactions with SO3, C1, etc.).

/

IfHO2 I(

NO2 I ''~ 02

NO

O

,,

CH3-S-OH

-I~

methane sulfonic acid

334


The C10 + DMS reaction is quite slow, with a rate constant at 298 K of only 9.5 • 10 -~5 cm 3 molecule -~ s -~ (Barnes et al., 1991). Combined with the fact that C1 reacts with organics at essentially a diffusion-controlled rate, while the rate constant for reaction with 0 3 is about an order of magnitude slower, the C10 + DMS reaction is not expected to be important in the troposphere. The reaction with BrO, however, is potentially important under some conditions. As discussed in Chapter 11, UV-visible differential optical absorption spectrometry (DOAS) has been used to measure BrO under these conditions (e.g., see Hausmann et al., 1993; Platt and Hausmann, 1994; and Tuckermann et al., 1997). Concentrations up to about 30 ppt have been measured. The rate constant for the B r O - D M S reaction is 2.6 • 10 -13 cm 3 molecule-1 s -~ at 298 K (Bedjanian et al., 1996) and gives with essentially unit yield DMSO and a bromine atom (Barnes et al., 1993; Bedjanian et al., 1996): BrO + CH3SCH 3 --~ CH3S(O)CH 3 + Br.

(69)

At 30 ppt BrO (7 • 10 s cm-3), the lifetime for the reaction of DMS with BrO is about 1.5 h and hence under these conditions can be significant. Obviously the situation in the Arctic at polar sunrise is unusual so that this represents the top end of estimates of the importance of BrO. However, model calculations suggest that it should be considered in other regions as well (Toumi, 1994). e. Oxidation by Ozone in the Aqueous Phase

high reactivity, even small concentrations of DMDS may be significant. For example, Jefferson et al. (1998) measured H2SO 4 and DMS in Antarctica and found that the DMS was insufficient to produce the levels of H 2 S O 4 that were observed; they postulated that the oxidation of sulfur compounds such as DMDS may be responsible for the additional sulfuric acid formation. The reaction of OH with dimethyl disulfide proceeds primarily by addition to a sulfur atom. As discussed by Abbatt et al. (1992), although one would expect abstraction of an H atom from the CH 3 group to occur as well at about the same rate as for DMS, such a channel is overwhelmed by the fast addition. The initially formed adduct is thought to decompose: OH + CH 3SSCH3

[CH3S(OH)SCH3 ] --~ CH3SOH + CH3S.

(70)

The CH3S and CH3SOH then react in a manner analogous to the reactions discussed for the DMS oxidation, ultimately forming SO 2 and methanesulfonic acid (e.g., Hatakeyama and Akimoto, 1983; Barnes et al., 1994a). The mechanism of the NO 3 radical reaction with DMDS is complex. The rate constants for the NO 3 reaction with CH3SCH3, CH3SH, and CH3SSCH 3 (Table 8.17) are similar and much larger than that for the reaction with H2S. This suggests that by analogy with the dimethyl sulfide reaction, the initial step is addition to a sulfur atom to form an adduct that then reacts further:

While the emphasis has been on oxidation of DMS and other reduced sulfur compounds in the gas phase, there is some indication that oxidation in the aqueous phase in clouds and fogs should also be considered. For example, Lee and Zhou (1994) have shown that DMS reacts with 0 3 in aqueous solutions quite rapidly, with a rate constant at 288 K of 4 • l0 s L mo1-1 s -1. They estimate that at 30 ppb 0 3, a level found globally, the lifetime for in-cloud oxidation of DMS is about 3 days, of the same order of magnitude as that for the gas-phase oxidation by OH (see Table 8.17). Given the moderately high solubility of not only DMS but other sulfur compounds as well (see Henry's law constants in Table 8.1), this is clearly an area that warrants further research.

FTIR studies of the N O 3 - D M D S reaction (Jensen et al., 1992; MacLeod et al., 1986) identified HCHO, SO2, CH3ONO2, CH3SO3H, HNO3, and CH3SNO 2 as products of the reaction with NO 3. In addition, there were a number of FTIR bands that were unidentified. Presumably the decomposition of the adduct gives at least in part intermediates such as CH3S and CH3SO, which then go on to react as described earlier (see Fig. 8.25).

2. D i m e t h y l D i s u l f i d e ( C H 3 S S C H 3 )

3. M e t h y l M e r c a p t a n ( C H 3 S H )

As seen from the data in Table 8.17, the OH reaction with CH3SSCH 3 (DMDS) is almost two orders of magnitude faster than that with DMS. Because of its

The OH radical also reacts with CH3SH by addition. The adduct appears to decompose at least in part to CH3S (Hatakeyama and Akimoto, 1983; Grosjean,

CH3SSCH

3 + N O 3 --~ [ C H 3 S ( N O 3 ) S C H

3 ] --~ P r o d u c t s .

(71)


1984), which then reacts as described earlier:

Possible atmospheric fates of the SHO radical reactions include

OH + CH3SH -~ [CH3S(OH)H] -~ CH3S + H 2 0 . (72) The NO 3 radical reaction with CH3SH gives a variety of products, including CH3SO3H, SO2, CH3ONO2, HCHO, CH3SNO2, HNO3, and CH3SSCH 3 (Jensen et al., 1992; MacLeod et al., 1986). The mechanism proposed is addition to the sulfur atom, followed by decomposition: NO 3 + CH3SH --~ [CH3S(NO3)H ] --~ CH3S + H N O 3, (73a) SO 2 -~-Other

4. Hydrogen Sulfide ( H 2 S ) The reactions of both OH and NO 3 with H2S appear to proceed by hydrogen atom abstraction: + H2S

NO 3 + H2S

+

(74)

SH,

+

H20

+

HNO 3 + SH.

(75)

The SH radical then reacts further: (76)

SH + 0 2 + Products, k76 ~

4 • 10 -19 cm 3 molecule -1 s - 1

(Stachnik and Molina, 1987) (77)

SH + 0 3 -0 HSO + 0 2 , k77 =

3.7 • 10-12

cm 3

molecule-1

-1

molecule -1

(79)

2 • 10-17 cm 3 molecule-1 s-1 (Lovejoy et al., 1987) HSO + 0 3 ~ Products,

(80)

k80 = 1.1 • 10-13 cm 3 molecule-1 s-1 (Atkinson et al., 1997a) HSO

+ NO 2 ~

HSO 2 +

NO,

(81)

k81 = 9.6 • 10-12 c m 3 molecule-1 s-1 (Lovejoy et al., 1987) followed by HSO 2

+ 0 2 --+

Products,

(82)

k82 = 3.0 • 10-13 cm 3 molecule-1 s-1 (Atkinson et al., 1997a). The reaction (80) of HSO with 0 3 has one channel generating HS + 2 0 2 as well as others giving further oxidation of the HSO (Wang and Howard, 1990). The ultimate oxidation product in the atmosphere is SO 2.

5. Carbon Disulfide

(CS 2)

The atmospheric sources and sinks of C S 2 and COS are reviewed by Chin and Davis (1993). Oxidation of CS2 generates COS (discussed in the following section). COS is also generated by natural emissions, including photochemical production from organosulfur compounds in seawater (e.g., see Zepp and Andreae, 1994). The reaction of OH with CS 2 is similar to the O H - C O reaction (see Chapter 5.A) in that the reaction proceeds by addition (e.g., Hynes et al., 1988; Diau and Lee, 1991a,b). An intermediate adduct is formed that can react with O2:

HOCS 2 + 0 2 ~

(78)

SH + NO 2 + HSO + NO, 10 -11 c m 3

k79 ~

OH + C S 2 -~- M ~ HOCS 2 + M,

(Atkinson et al., 1997a)

k78 = 5 . 8 •

HSO + 0 2 ~ Products,

(73b)

products.

The direct production of S O 2 in a reaction such as (73b) was suggested by the independence of the measured SO 2 yields on the oxygen concentration. NO 2 is not produced in significant yields ( < 5 % ) in the N O 3 - C H 3 S H reaction (Dlugokencky and Howard, 1988), ruling out a simple oxygen atom transfer or decomposition of the adduct to NO 2 and other products.

OH

335

Products,

(83) (84)

k84(298 K) = 3 • 10 -14 cm 3 molecule-1 s-1

-1

(Atkinson et al., 1997a). At concentrations of 0 3 of 30 ppb and NO 2 of 0.1 ppb, the lifetimes of the SH radical with respect to reaction with 02, 03, and NO 2 are > 0.5, 0.4, and 7 s, respectively. Thus all three species may be involved in the atmospheric reactions of SH, with the latter two giving the SHO radical.

(Atkinson et al., 1997a) The recommended (Atkinson et al., 1997a, 1997b) lowand high-pressure limiting rate constants for reaction (83) are k 0 = 8 • 10-31[N2] cm 6 molecule -2 s -1 and ks=8• 10-12 cm 3 molecule-1 s - l , with F c = 0 . 8 (see Chapter 5.A.2), giving an effective bimolecular rate constant at 298 K and 1 atm of 4.7 • 10 -12 cm 3 molecule-1 s-1. There is no experimental evidence for

336


a direct channel forming HS + COS, and if it occurs, the rate constant must be less than 2 • 10 -15 cm 3 molecule-1 s-1 (Murrells et al., 1990). The equilibrium constant for reaction (83) at 299 K is Ks3 = 1.7 • 10-17 cm 3 molecule-1 (Murrells et al., 1990). As discussed in detail by Stickel et al. (1993), there are 26 potential sets of products of reaction (84) of the adduct with 0 2 that are exothermic! They have measured the yield of COS to be 0.83 _+ 0.08, of CO to be 0.16 _+ 0.03, and of SO 2 to be 1.15 _+ 0.10. Both the COS and about 75% of the SO 2 are produced rapidly, indicating that they either are formed directly from the reaction of the adduct with 02, e.g., via Hoes

2 + 0 2 ---> C O S

-Jr- HOSO,

COS + H + SO2,

(85a) (85b)

--+ C O S

+ HSO2,

(85c)

--+ C O S

+ H O 2 + S,

(85d)

--+ H O 2 + C O

+ 82,

(85e)

or, alternatively, are formed from very fast secondary reactions of other primary products. CO is also produced rapidly, perhaps from such pathways as Hoes

2 + 0 2 ---> C O

-Jr- HOSO

+ S,

--+ CO + H + SO 2 + S.

(85f) (85g)

The formation of HO 2 in high yield, 95 _+ 15% of the OH reacted, has also been observed using laser magnetic resonance (Lovejoy et al., 1990), indicating the importance of channels such as (85d) and (85e).

6. Carbonyl Sulfide (COS) The reaction of OH with COS is significantly slower than the CS 2 reaction, with a rate constant at 298 K of 2.0 • 10 -15 cm 3 molecule-1 s-1 (Atkinson et al., 1997a, 1997b). At an OH concentration of 1 • 106 radicals cm -3, this gives a lifetime with respect to this reaction of about 16 years, far too long to be a significant removal path in the troposphere. Unlike the CS 2 reaction, the reaction of OH with COS is independent of pressure and of the presence of 0 2 (e.g., see Cheng and Lee, 1986; and Wahner and Ravishankara, 1987). The reaction products are not known.

F. PROBLEMS 1. Derive Eq. (B), relating the pseudo-Henry's law constant to the true Henry's law constant for the physical solubility of CO 2 exclusive of other reactions and to

the pH and equilibrium constants for reactions (2) and (3). 2. Using the kinetic and thermodynamic data given in the text, calculate the rates of sulfate formation in mol L-1 s-1 from the reaction of 0 3 with each of the S(IV) species, i.e., SO2. H 2 0 , H S O f , and SO 2-, and the total rate of formation of S(VI) from these reactions at a pH of 3.0 and with a constant gas-phase concentration of SO 2 of 20 ppb and 0 3 of 50 ppb (you can compare to the data in Fig. 8.15). 3. (a) What is the effective Henry's law constant for SO 2 for an aqueous solution of pH 5.0? (b) What is the total concentration of S(IV) in solution under these conditions if the gas-phase SO 2 concentration is 10 ppb? 4. Repeat Problem 3 for pH 2.0. Compare your answers to those for Problem 3 and explain based on the chemistry of S(IV) in solution. 5. Calculate the mole fractions of SO2(aq), HSO;-, and SO 2- in solution at a pH of 3.0, characteristic of many fogs, as well as the total concentration of S(IV) in solution for 20 ppb SO 2 in the gas phase. 6. Calculate the mole fractions of SO2(aq), HSO;-, and SO 2- in solution at a pH of 8.0, equal to that of seawater and therefore expected for airborne particles formed from seawater in clean regions as well as the total concentration of S(IV) in solution for 20 ppb SO 2 in the gas phase. Comment on the implications for the potential role of sea salt particles in the oxidation of SO 2 in the marine boundary layer. 7. How fast is the oxidation of SO 2, in percent per hour, in the gas phase for an OH concentration of 1 • 107 radicals cm -3, a typical peak concentration in the lower troposphere? Comment on the significance of your calculated rate in light of observed rates of SO 2 oxidation of up to 30% h -1 in the troposphere. 8. Calculate the rate of oxidation of SO 2 in percent per hour by H 2 0 2 in the aqueous phase for an SO 2 concentration of 10 ppb, a pH of 3.0, a temperature of 298 K, a constant gas-phase H 2 0 2 concentration of 1 ppb, and a liquid water content (LWC) of (a) 0.1 g of liquid water per m -3 of air, typical of fogs, and (b) 1 g m -3, typical of clouds. Why does the LWC affect the rate of oxidation of SO2? 9. An air mass has an 0 3 concentration of 60 ppb and cloud drops with pH 4.0. What must the concentration of gaseous H 2 0 2 be so that the rate of oxidation of S(IV) in the cloud by these two oxidants is equal? 10. The acid dissociation constants for H C O O H , CH3COOH, and H O N O are 1.77 • 10 -4, 1.76 • 10 -5, and 4.6 • 10 -4, respectively. Assume that the liquidphase acidity is provided by dissolution of gas-phase H N O 3. What would the gas-phase concentrations of H N O 3 need to be to give equal concentrations of the

aEFEaENCES

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348


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Bulk Cloudwater or Fogwater Samples Obey Henry's Law?' by S. N. Pandis and J. H. Seinfeld," J. Geophys. Res., 97, 6075-6078 (1992). Worsnop, D. R., M. S. Zahniser, C. E. Kolb, J. A. Gardner, L. R. Watson, J. M. Van Doren, J. T. Jayne, and P. Davidovits, "Temperature Dependence of Mass Accommodation of SO 2 and H 2 0 2 on Aqueous Surfaces," J. Phys. Chem., 93, 1159-1172 (1989). Wurzler, S., A. I. Flossmann, H. R. Pruppacher, and S. E. Schwartz, "The Scavenging of Nitrate by Clouds and Precipitation," J. Atmos. Chem., 20, 259-280 (1995). Xu, H., P. J. Wentworth, N. W. Howell, and J. A. Joens, "Temperature Dependent Near-UV Molar Absorptivities of Aliphatic Aldehydes and Ketones in Aqueous Solution," Spectrochim. Acta, Part A, 49A, 1171-1178 (1993). Yermakov, A. N., B. M. Zhitomirsky, G. A. Poskrebyshev, and S. I. Stoliarov, "Kinetic Study of SO~- and HO 2 Radical Reactivity in Aqueous Phase Bisulfite Oxidation," J. Phys. Chem., 99, 3120-3127 (1995). Zafiriou, O. C., "Natural Water Photochemistry," in Chemical Oceanography (J. P. Riley and R. Chester, Eds.), Vol. 8, pp. 339-379, Academic Press, London, 1983. Zellner, R., M. Exner, and H. Herrmann, "Absolute OH Quantum Yields in the Laser Photolysis of Nitrate, Nitrite, and Dissolved H 2 0 2 at 308 and 351 nm in the Temperature Range 278-353 K," J. Atmos. Chem., 10, 411-425 (1990). Zellner, R., and H. Herrmann, "Free Radical Chemistry of the Aqueous Atmospheric Phase," in Spectroscopy in Environmental Science (R. J. H. Clark and R. E. Hester, Eds.), Chap. 9, pp. 381-451, Wiley, New York, 1995. Zepp, R. G., J. Hoign6, and H. Bader, "Nitrate-Induced Photooxidation of Trace Organic Chemicals in Water," Environ. Sci. Technol., 21, 443-450 (1987). Zepp, R. G., B. C. Faust, and J. Hoign6, "Hydroxyl Radical Formation in Aqueous Reactions (pH 3-8) of Iron(II) with Hydrogen Peroxide: The Photo-Fenton Reaction," Environ. Sci. Technol., 26, 313-319 (1992). Zepp, R. G., and M. O. Andreae, "Factors Affecting the Photochemical Production of Carbonyl Sulfide in Seawater," Geophys. Res. Lett., 21, 2813-2816 (1994). Zhao, Z., R. E. Stickel, and P. H. Wine, "Branching Ratios for Methyl Elimination in the Reactions of OD Radicals and C1 Atoms with CH3SCH3," Chem. Phys. Lett., 251, 59-66 (1996). Zhou, X., and Y.-N. Lee, "Aqueous Solubility and Reaction Kinetics of Hydroxymethyl Hydroperoxide," J. Phys. Chem., 96, 265-272 (1992). Zhu, X., J. M. Prospero, D. L. Savoie, F. J. Millero, R. G. Zika, and E. S. Saltzman, "Photoreduction of Iron(III) in Marine Mineral Aerosol Solutions," J. Geophys. Res., 98, 9039-9046 (1993). Zhu, X. R., J. M. Prospero, and F. J. Millero, "Diel Variability of Soluble Fe(II) and Soluble Total Fe in North African Dust in the Trade Winds at Barados," J. Geophys. Res., 102, 21297-21305 (1997). Zhuang, G., Z. Yi, R. A. Duce, and P. R. Brown, "Link between Iron and Sulphur Cycles Suggested by Detection of Fe(II) in Remote Marine Aerosols," Nature, 355, 537-539 (1992). Zuo, Y., and J. Hoign6, "Formation of Hydrogen Peroxide and Depletion of Oxalic Acid in Atmospheric Water by Photolysis of Iron(III)-Oxalato Complexes," Environ. Sci. Technol., 26, 1014-1022 (1992). Zuo, Y., and J. Hoign6, "Evidence for Photochemical Formation of H 2 0 2 and Oxidation of SO 2 in Authentic Fog Water," Science, 260, 71-73 (1993).

C H A P T E R

9 Particles in the Troposphere The most obvious characteristic of air pollution is the loss of visibility. This is primarily due to suspended airborne particles, which scatter light efficiently, giving the atmosphere a "hazy" appearance. As discussed in Chapter 14, scattering of light has direct effects on climate, altering the amount of solar radiation reaching the earth's surface; these suspended particles also contribute indirectly via their effects on clouds by acting as cloud condensation nuclei. Particles may also lead to heating of the lower atmosphere if they contain light absorbers such as elemental carbon, which strongly absorbs visible light, or mineral dusts, which absorb long-wavelength infrared (see Chapter 14.C.1). (For a series of papers on particle characteristics and their relationship to radiative forcing and climate, see the July 20, 1998, issue of Journal of Geophysical Research.) In addition to visibility and climate implications, these particles have significant health impacts which have been increasingly recognized (see Chapter 2.A.5). Finally, particles may act as sinks of reactive species such as HO2, particularly in remote regions, hence affecting the chemistry of the gas phase as well (e.g., Saylor, 1997). As a result of these widespread implications, understanding their direct sources, their formation from chemical reactions in air, their fates, and how their physical and chemical properties determine health and visibility impacts is critical. While many of the overall chemical and physical characteristics of particles have been elucidated, as we shall see in this chapter, there remain large gaps in our knowledge in areas central to policy and regulatory issues (National Research Council, 1998). As a result, this is a particularly active and rapidly evolving area of research in atmospheric chemistry. A. P H Y S I C A L

liquid, with diameters between ~0.002 and ~ 100/xm. The lower end of the size range is not sharply defined because there is no accepted criterion at which a cluster of molecules becomes a particle. The upper end corresponds to the size of fine drizzle or very fine sand; these particles are so large that they quickly fall out of the atmosphere and hence do not remain suspended for significant periods of time. There are, of course, larger particles produced in the atmosphere (e.g., raindrops, ~ 1 mm, and hail, ~ 1-20 mm), but their rapid fallout precludes, for all practical purposes, their inclusion in the definition of atmospheric particles. As we shall see, the most important particles with respect to atmospheric chemistry and physics are in the 0.002- to 10-/xm range. Aerosols are defined as relatively stable suspensions of solid or liquid particles in a gas. Thus aerosols differ from particles in that an aerosol includes both the particles and the gas in which they are suspended. However, while this is the rigorous definition of aerosols, one should note that the term is often used in the atmospheric chemistry literature to denote just the particles. Particles may be either directly emitted into the atmosphere or formed there by chemical reactions; we refer to these as primary and secondary particles, respectively. The relative importance of primary and secondary particles will clearly depend on the phenomena examined, the geographical location with its particular mix of emissions, and the atmospheric chemistry. There are a number of properties of particles that are important for their role in atmospheric processes. These include, in addition to their number concentration, their mass, size, chemical composition, and aerodynamic and optical properties. Of these, size is the most important; it is related not only to the source of the particles (see later) but also to their effects on health, visibility, and climate. Atmospheric particles are usually referred to as having a radius or a diameter, implying they are spherical. However, many particles in the atmosphere have quite

PROPERTIES

1. S o m e D e f i n i t i o n s Particles, or particulate matter, may be solid or

349

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A. PHYSICALPROPERTIES

irregular shapes for which geometrical radii and diameters are not meaningful. Some means of expressing the size of such particles is essential since many important properties of the particle such as volume, mass, and settling velocity depend on the size. In practice, the size of such irregularly shaped particles is expressed in terms of some sort of equiualent or effectiue diameter that depends on a physical, rather than a geometrical, property. There are several different types of effective diameters. One of the most commonly used is the aerodynamic diameter, D~, which is defined as the diameter of a sphere of unit density (1 g cm 3) that has the same terminal falling speed in air as the particle under consideration. This effective diameter is particularly useful because it determines the residence time in the air and it reflects the various regions of the respiratory system in which particles of different sizes become deposited. D~ is given by Eq. (A): Oa

Dgkl/PP . V P0

(a)

Dg is the geometric diameter, pp is the density of the particle, neglecting the buoyancy effects of air, P0 is the reference density (1 g cm 3), and k is a shape factor, which is 1.0 in the case of a sphere. Because of the effect of particle density on the aerodynamic diameter, a spherical particle of high density will have a larger aerodynamic diameter than its geometric diameter. However, for most substances, pp 10 so that the difference is less than a factor of ~3 (Lawrence Berkeley Laboratory, 1979). Particle densities are often lower than bulk densities of pure substances due to voids, pores, and cracks in the particles. Throughout this chapter, we use the term diameter with the understanding that it is the aerodynamic diameter of the particle unless stated otherwise. Another type of diameter commonly used is the Stokes diameter, D s. This is defined as the diameter of a sphere that has the same density and settling velocity as the particle. Thus Stokes diameters are all based on settling velocities, whereas the aerodynamic diameter (D~) also includes a standardized density of unity. Figure 9.1 summarizes some of the characteristics of particles and aerosols encountered in both environmental and industrial atmospheres (Lapple, 1961). This chapter treats those aerosol phenomena that are known or believed to be important in atmospheric chemistry. For treatment of related, but specialized, topics, a number of excellent references are available. The classic works on aerosol physics are The Mechanics of Aerosols by the late N. A. Fuchs (1964) and Highly Dispersed Aerosols (Fuchs and Sutugin, 1970). The

351

chemical engineering aspects of aerosols are emphasized in S. K. Friedlander's book Smoke, Dust, and Haze: Fundamentals of Aerosol Behauior (1977); several topics, including photophoresis, thermophoresis, and coagulation, are included in Aerosol Science, edited by C. N. Davies (1966); and health considerations are treated in Inhalation Studies by R. F. Phalen (1984) and in Pulmonary Toxicology of Respirable Panicles (Sanders et al., 1980). Finally, several works deal with laboratory generation and characterization of aerosols, including those by Hinds (1982), Mercer (1973), Willeke (1980), and Liu (1976).

2. Size D i s t r i b u t i o n s

a. Number, Mass, Surface, and Volume Distributions The atmosphere, whether in remote or urban areas, always contains significant concentrations of particles, up to 10 s cm 3. These may have diameters anywhere within the entire range from molecular clusters to ~100 /xm. Because the size of atmospheric particles plays such an important role in both their chemistry and physics in the atmosphere as well as in their effects, it is important to know the distribution of sizes. We thus consider first how these size distributions are characterized. An obvious way to express the distribution of particle sizes found in the atmosphere would be to plot, in the form of a histogram, the number (AN) of particles found in certain arbitrarily chosen intervals of diameter (zXD), for example, from 0.002 to 0.01 /xm, 1-10 /xm, and so on. However, since in the atmosphere there tend to be a much greater number of small particles relative to large particles, a linear plot of AN against zXD would give what would appear to be a narrow spike at the origin whose details could not be distinguished. This can be seen in Fig. 9.2, where typical impactor data from Wesolowski et al. (1980) for a sample consisting of Arizona road dust have been plotted; the original data in the form of a mass distribution have been converted to a number distribution using simplifying assumptions. Figure 9.3 shows the same data but with the horizontal axis plotted as log D. This allows one to show a much larger particle size range than that which can be shown using a linear scale for the diameter. A second problem in expressing the size distribution of aerosols is that the intervals of diameter over which it is experimentally convenient to measure the number of particles are not constant in terms of either D or log D. For example, as discussed in Chapter 11, a multistage impactor is often used to measure the number of particles in certain size ranges. The size intervals are

352

9. PARTICLES IN THE TROPOSPHERE

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characterized by the 50% cutoff point for each stage, where the 50% cutoff point is defined as the diameter of spheres of unit density, 50% of which are collected by that stage of the impactor. For example, a typical set of 50% cutoff points is 8.0 (stage 1), 4.0 (stage 2), 1.5 (stage 3), and 0.5/xm (stage 4), respectively, with particles smaller than 0.5/xm being collected on an afterfilter. One might then use the approximation that each stage captures particles ranging from a diameter corresponding to midway between its cut point and that of the next higher and lower stages. With these assumptions, the ranges of particle diameters captured by each

stage for a typical impactor are given in Table 9.1; also given are the diameter intervals AD as well as h log D for each stage. It is seen that neither the intervals in terms of diameter nor those in terms of the logarithm of the diameter are equal. As a result, a plot of AN against D or log D will look something like those shown in Figs. 9.2 and 9.3. These give a somewhat distorted picture of the size distribution, however, because the height of any bar, that is, the number of particles, depends on the width of the interval taken, that is, AD or h log D. To give a more physically descriptive picture of the size distribution, a modified plot of the number of particles normalized for the width of the diameter interval is used; that is, the number of particles per unit size interval is plotted on the vertical axis. With log D as the horizontal axis then, a normalized plot is one of AN h log D against log D, where AN is the number of particles in that interval of h log D; Fig. 9.4 is an example of this type of plot using the data in Figs. 9.2 and 9.3. The area under each rectangle then gives the number of particles in that size range. Rather than showing histograms, one usually draws a smooth curve through the data. Figure 9.5, for example, shows one such curve of AN A log D against log D for a typical urban model aerosol; to emphasize the wide range of numbers involved, a logarithmic scale has also been used for the vertical axis (Whitby and Sverdrup, 1980). However, it is not only the number of particles in each size interval that is of interest but also how other properties such as mass, volume, and surface area are distributed among the various size ranges. For example, the U.S. Environmental Protection Agency's air quality

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102

log D

F I G U R E 9.3 Plot of number of particles (N) against log D (D diameter). Data same as those in Fig. 9.2.

aerodynamic

A. PHYSICAL PROPERTIES

TABLE 9.1

353

Typical 5 0 % Cutoff Points and Approximate Range of Particle Diameters Captured by Each Stage in a Four-Stage Impactor a

50% cutoff point (Ixm)

Stage 1 2 3 4 Afterillter

Approximate range of particle diameters captured by that stage (Ixm)

AD

A log D

(Ixm)

(Ixm)

14.0b-6.0 6.0-2.8 2.8-1.0 1.0-0.3 c 0.3 -0.00 lC

8.0 3.2 1.8 0.7 0.3

0.37 0.33 0.45 0.52 2.48

8.0 4.0 1.5 0.5 0.5

a From Wesolowski et al. (1980). b Assuming a cutoff of 20 /xm for the sampling probe. c Assuming the afterfilter collects all particles >0.001-/xm diameter.

standards for particles are expressed in terms of mass of particulate matter, with diameters 10/xm per unit volume of air, PM10, and those with diameters 2.5 /xm, PM2. 5 (see Chapter 2.D). It is thus important to know the mass distribution of atmospheric particulate matter. Similarly, surface and volume distributions are important when considering reactions of gases at the surface of particles or reactions occurring within the particles themselves, for example, the oxidation of SO 2 to sulfate. Because of this need to know how the mass, surface, and volume are distributed among the various particle sizes, distribution functions for these parameters (i.e., mass, surface, and volume) are also commonly used for atmospheric aerosols in a manner analogous to the number distribution. That is, Am A log D, AS A log D, or AV A log D is plotted against D on a logarithmic scale, where Am, AS, and AV are the mass, surface area, and volume, respectively, found in a given size interval; again the area under these curves gives

the total mass, surface, or volume in the interval considered. Figure 9.6 shows the surface and volume distributions for the number distribution shown in Fig. 9.5; also shown is the same number distribution for comparison, where the vertical axis is now linear (rather than logarithmic as in Fig. 9.5). When the particle data are plotted as mass, surface, or volume distributions, an important characteristic of typical urban aerosols emerges clearly. As seen in the number distribution of Fig. 9.6a, there is a large peak at ~0.02 /xm and a slight "knee" in the curve around 0.1 /xm. The volume distribution for the same aerosol (Fig. 9.6c) shows two strong peaks, one in the 0.1- to 1.0-/xm range and the second in the 1- to 10-/xm region, with a minimum in the 1- to 2-/xm region. The surface distribution (Fig. 9.6b) shows a major peak in the vicinity of 0.1 /xm, with smaller peaks in the region between 0.01 and 0.1 /xm and between 1 and 10 /xm. The "knee" observed in the number distribution

~176 t z

o b .v-

Ici

40

8'

30 St4ge

10 0 10 -3

10 -2

10 1

I

Stage 2

Stage

1

,/ 10

Stage 1 ]

102

.~ f

log D F I G U R E 9.4 "Normalized" plot of AN A log D versus log D for data shown in Figs. 9.2 and 9.3.

354


%

106 105

E

104

8,

103

m'E

15 10 5 0

I

102

O

g,

10~

?-

,

o E

g,,

10-1

600

b m

400 200 I III

10-2

C

10-3 10-4 0.001 0.01

0.1

1.0

10

100

log D (gm)

F I G U R E 9.5 Plot of number distribution AN iX log D versus log D for a typical urban model aerosol (adapted from Whitby and Sverdrup, 1980).

suggests that the observed curve may be a combination of two different distributions, shown by the dashed lines in Fig. 9.6a. While the multimodal nature of atmospheric aerosols had been suggested earlier (e.g., Junge, 1963), Whitby and co-workers, in a classic set of papers (Whitby et al., 1972a, 1972b; Husar et al., 1972), were the first to establish and explore in detail the significance of such a bimodal distribution in terms of the origins, chemical characteristics, and removal processes of the two groups of particles of different sizes. Although instrument artifacts can be a confounding factor, it is now widely accepted that atmospheric aerosols usually occur in specific size groupings that are different in their origins and properties. Indeed, based on the number, surface, and volume distributions shown in Fig. 9.6, Whitby and co-workers suggested that there were three distinct groups of particles contributing to this atmospheric aerosol. Particles with diameters > 2.5/xm are identified as coarse particles and those with diameters 2.5/xm are called fine panicles. The fine particle mode typically includes most of the total number of particles and a large fraction of the mass, for example, about one-third of the mass in nonurban areas and about one-half in urban areas. The fine particle mode can be further broken down into particles with diameters between ~0.08 and 1-2 /xm, known as the accumulation range, and those with diam-

o

30-

vGI

20

o

100

.001

I

IIIII 0.01

0.1

1.0

10

100

D (gm)

F I G U R E 9.6 Number, surface, and volume distributions for a typical urban model aerosol (adapted from Whitby and Sverdrup, 1980).

eters between 0.01 and 0.08 /xm, known as the transient or Aitken nuclei range. As the technology for measuring small particles has improved (see Chapter ll.B), ultrafine panicles have also been increasingly studied. While there is no fixed definition of these particles, they are usually taken to mean those with diameters less than 0.01 /xm, i.e., 10 nm. In short, particles in the atmosphere are now frequently treated in terms of the four modes summarized in Fig. 9.7, which also shows the major sources and removal processes for each one. Although the vertical axis is not shown, it could in theory be any of the distributions discussed, that is, number, mass, surface, or volume. Particles in the coarse particle range are usually produced by mechanical processes such as grinding, wind, or erosion. As a result, they are relatively large and hence settle out of the atmosphere by sedimentation, except on windy days, where fallout is balanced by reentrainment. Schmidt et al. (1998) also point out that particles generated in blizzards and sandstorms develop an electrostatic charge and that the electrostatic forces between the particles and the surface may be similar in magnitude to gravitational forces. These large particles


I Hot vapor I

Chemical conversion of gases to low volatility vapors

Condensation

Lowvapor volatility

I Primary particles I Coagulation I Chain aggregates

355

f-- Homogeneous nucleation Condensation growth 9 of nuclei" /

Wind blown dust +

Chemical conversion of gases to low volatility vapors

Emissions

+ Sea Spray + Volcanos + Plant particles

lation

Lowvapor volatility]

%/.-',

Homogeneous nucleation

"-"~i/

7

/

/

washoutl ~ \ \

,,,~edimentation~

__K

0.001

0.01

0.1 1 2 Particle diameter (gm)

Transient nuclei Ultra fine orAitken nuclei particles ~]_~ range ~]_~

Accumulation range

10

100

Mechanically generated aerosol range

~: Fine particles Coarse particles FIGURE 9.7 Schematic of an atmospheric aerosol size distribution showing four modes. The original hypothesis of Whitby and co-workers is shown by the solid, trimodal curves, and the fourth, ultrafine particle mode, as well as the two peaks sometimes observed in the accumulation mode are shown by the dashed lines (adapted from Whitby and Sverdrup, 1980).

can also be removed by washout. In the atmosphere, transport of coarse particles over long distances can occur, however, by convective processes. Chemically, their composition reflects their source, and hence one finds predominantly inorganics such as sand, sea salt, and so on in this range, although significant amounts of organics have also been reported associated with dust particles (e.g., Boon et al., 1998). Because the sources and sinks are different from those of the smaller modes, the occurrence of particles in this mode tends to be only weakly associated with the fine particle mode. The majority of biological particles, spores, pollens, and so on, tend to be in the coarse particle range. It should be noted, however, that this by no means indicates that compounds and elements associated with such mechanical processes are always found exclusively in the coarse particle range. For example, O'Dowd and Smith (O'Dowd and Smith, 1993b; Smith and O'Dowd, 1996)

report that sea salt was the major component of all particles with radii above 0.05 /xm over the northeast Atlantic Ocean. While particles in the coarse particle mode are generally sufficiently large that they are removed relatively rapidly by gravitational settling, there are largescale mechanisms of transport that can carry them long distances during some episodes. For example, there are many studies showing the transport of dust in larger particles from the Sahara Desert to the northwestern Mediterranean, Atlantic Ocean, Bermuda, the Barbados, Ireland, the Amazon Basin, and the United States (see, for example, papers by Prospero and Nees, 1977, 1986; Prospero et al., 1981; Talbot et al., 1986; Wolff et al., 1986; Savoie et al., 1989; Mateu et al., 1994; Artaxo and Hansson, 1995; Dentener et al., 1996; Jennings et al., 1996; Gatz and Prospero, 1996; Chiapello et al., 1997; Perry et al., 1997; Moulin et al., 1997; and Li-Jones

356

9. P A R T I C L E S IN T H E T R O P O S P H E R E

and Prospero, 1998). Similarly, dust transported from Asia has been recorded on a regular basis over the Pacific (e.g., see Duce et al., 1980; Uematsu et al., 1985; and Zhang et al., 1997) and in the Canadian Arctic (e.g., Kawamura et al., 1996b). Asian dust has been observed, usually during the spring, at the Mauna Loa Observatory in Hawaii (e.g., see Shaw, 1980; Braaten and Cahill, 1986; Zieman et al., 1995; and Holmes et al., 1997); the elemental signature in particles in the size range 0.5- to 3.0-/xm, particularly in terms of the Si Fe and Ti Fe ratios, is very similar to those measured during dust storms in Beijing, consistent with long-range transport of these particles (Braaten and Cahill, 1986). The mineralogy of the dust has proven useful in source determination of particles after longrange transport (e.g., see Leinen et al., 1994; Merrill et al., 1994; Avila et al., 1997; and Caquineau et al., 1998). While these "dust storms" are episodic in nature, such long-range transport of dust has been proposed to play a significant role in SO 2 and NO X heterogeneous chemistry, in the photochemical cycles leading to 0 3 formation (Dentener et al., 1996), and, as discussed in Chapter 14, possibly in the radiative balance of the atmosphere as well. Particles in the accumulation range with diameters from ~0.08 to ~ 1-2 /xm typically arise from condensation of low-volatility vapors (e.g., following combustion) and from coagulation of smaller particles in the nuclei range either with themselves or, more likely, with the larger particles in the accumulation range. The coagulation rates for particles in the nuclei range with the larger particles in the accumulation range are usually much larger than for self-coagulation of the small particles; this occurs because of the high mobility of the smaller particles combined with the larger target area of the bigger particles. Because of the nature of their sources, particles in the accumulation range generally contain far more organics than the coarse particles (other than biologically derived particles) as well as soluble inorganics such as NH 4, NO 3 , and SO 2 . While many particle distributions show one peak in the accumulation range, many instances have been observed in which there are two peaks. For example, as seen in Fig. 9.8, John and co-workers (1990) observed two peaks within the traditional accumulation mode, one at 0.2 0.1 and one at 0.7 0.2/xm, in studies of particles carried out in a relatively polluted urban area. This bimodal character of particles appears to occur quite frequently; for example, Hering et al. (1997) showed that during the summer in the Los Angeles area, two modes often occur, having diameters of about 0.26 /xm (range from 0.10 to 0.39 /xm) and 0.65 /xm (range from 0.46 to 0.90/xm), respectively.

E 0

300 -

250

-

E 200 =,S, o 150 100 _o "o9 50 0 I C "o9 0.01

I

0.1

1

10

100

D a (gm)

FIGURE 9.8 Typical size distribution of nitrate in southern California in 1987 fitted by the sum of three log-normal distributions with peaks at 0.2, 0.7, and 4.4/xm (adapted from John et al., 1990).

Similarly, Fig. 9.9 shows a particle number size distribution for measurements made at a relatively clean site, Cheeka Peak, which is located approximately 2 km from the coast in Washington State at an altitude of 480 m (Quinn et al., 1993). The data in Fig. 9.9 represent air masses arriving at the sampling site from aloft over Canada. Again two peaks are seen in the number distribution in the 0.1- to 1-/xm range, one peaking at 0.15 and the other at 0.45/xm. A bimodal distribution of sulfate with mass mean diameters of 0.25 and 0.55 /xm, respectively, has also been observed during the summer in Greenland (Kerminen et al., 1998). Hering and Friedlander (1982) made similar observations for particle sulfate and attributed the smallest mode (referred to as the "condensation" mode) to formation from gas-phase SO 2 oxidation and the larger modes (the "droplet" mode) to oxidation in the condensed phase. Meng and Seinfeld (1994) have shown that the droplet mode particles cannot arise from growth of the smaller, condensation mode particles and propose that the condensation mode particles are activated to form fog or cloud droplets, followed by chemical reactions and subsequent evaporation to form the "droplet" particles.

/

1000 -I a

o.m, E

0

i_

-o

~E -n

Z

-o

F

500 f 0

II

0.01

J

. . . . . . . . . .

0.1

I

I . . . . . .

10

D (gm)

FIGURE 9.9 Particle number size distributions measured at Cheeka Peak, Washington, in 1991 (adapted from Quinn et al., 1993).

A. PHYSICALPROPERTIES Based on such observations and proposals, Ondov and Wexler (1998) have offered a modification of the Whitby accumulation mode, summarized in Fig. 9.10. High-temperature combustion sources produce particles in the Aitken nuclei and the accumulation mode. For example, the mass median diameters of particles emitted from a variety of combustion sources, including incinerators, coal- and oil-fired boilers, and automobiles and trucks, typically fall in the range of 0.05-0.35 /xm (Hildemann et al., 1991a, 1991b; Ondov and Wexler, 1998). The relative numbers of particles produced in the Aitken nuclei range compared to the accumulation range depend on the nature of the combustion process (e.g., fuel and operating conditions) as well as the conditions of dilution. Figure 9.11, for example, shows the surface distribution of particles produced by the combustion of several organic compounds as well as by automobiles and by a burning candle. The "dirtier" flames (e.g., the candle and the acetone flame) produced significant numbers of particles in the accumulation mode, whereas the cleaner flames produce more Aitken nuclei (National Research Council, 1979). Because the particles generated in high-temperature combustion processes contain hygroscopic compounds, water vapor can be taken up or evaporate, depending on the atmospheric conditions. In addition, the particles can be taken up into clouds and fogs. The condensed-phase water provides a medium for atmospheric reactions that generate low-volatility species; the best known example is the oxidation of SO 2 to sulfate (see Chapter 8.C.3). When the water evaporates, the remaining particle contains this additional material and hence has grown to a larger size (Fig. 9.10; see also Section B-3). Aitken Nuclei

combusti temperature,,,,,~A on

Hygroscopic Growth (reversible)

Methanol

2

8'

\ 1-

I

/

Cand

I

I

I

I

0.1

I

I

\\Acetone

I I / I .."..'". '% ...." '%.

9 "'.. "'% ""%

\

\

\

ndle 0

i

0.005 0.01

0.05 0.1 D (l,tm)

0.5

FIGURE 9.11 Surface distribution of particles from the combustion of several organics and from automobiles and a candle (adapted from National Research Council, 1979).

Two peaks in the accumulation mode may also result if the particles are externally mixed rather than internally mixed (the latter meaning that all particles have the same composition corresponding to a mixture of the various components). That is, the composition of individual particles may not be the same as the overall bulk particle composition, with some having more hygroscopic components than others. These are referred to as "externally mixed" particles; indeed single-particle analyses suggest that externally mixed particles are common (see Chapter 11.B.4). In this case, a Coarse Particles

Aged

A

A

Fresh

//~~~]/ll~~~~'~'@ H20

Coagulation/condensation Conversion (nearplume) (irreversible) I

a

Aged

AA

I

Propane

3

Accumulation Mode Fresh

High

357

I

I

I t

/~P-- / /

I

I

D(l,tm)

I

I

I

// /

I

I

1.0

I

~,

I

I

I

FIGURE 9.10 Modifiedparticle modes and growth processes for sulfate particles involving aqueous-phase reactions in low altitude fogs and in higher altitude clouds upon advection of boundary-layer air upwards. (Adapted with permission from Ondov and Wexler, 1998. Copyright 1998 American Chemical Society.)

35S


monodisperse (i.e., single mode) particle size distribution can develop into a bimodal distribution when the more hygroscopic particles take up increased amounts of water (e.g., see modeling studies of Kleeman et al., 1997). In short, growth of sulfate particles at least in the accumulation mode and the presence of two peaks are both believed to be largely controlled by interactions with water in the atmosphere, including the aqueous phase oxidation of SO 2 to sulfate. Particles in the accumulation range tend to represent only a small portion of the total particle number (e.g., 5%) but a significant portion (e.g., 50%) of the aerosol mass. Because they are too small to settle out rapidly (see later), they are removed by incorporation into cloud droplets followed by rainout, or by washout during precipitation. Alternatively, they may be carried to surfaces by eddy diffusion and advection and undergo dry deposition. As a result, they have much longer lifetimes than coarse particles. This long lifetime, combined with their effects on visibility, cloud formation, and health, makes them of great importance in atmospheric chemistry. The Aitken nuclei, with 0.01 D 0.08 /~m, arise from ambient-temperature gas-to-particle conversion as well as combustion processes in which hot, supersaturated vapors are formed and subsequently undergo condensation. These particles act as nuclei for the condensation of low-vapor-pressure gaseous species, causing them to grow toward the accumulation range; alternatively, these nuclei may grow larger by coagulation. This range contains most of the total number of particles but relatively little of the total mass because of their small size. The lifetime of these particles is short, sometimes on the order of minutes, due to their rapid coagulation. Development of techniques to measure ultrafine particles with diameters less than 0.01 /~m (10 nm) has established the presence of what might be considered a fourth mode, sometimes referred to as the nucleation mode. As discussed later, these ultrafine particles are generated by gas-to-particle conversion processes that are not yet well understood on a molecular level. Although they are sufficiently small that they do not contribute a large fraction of the total mass, they can be present at significant number concentrations. For example, in the Los Angeles area, the 24-h average concentration of ultrafine particles (0.017 D 0.1 /xm) during the winter has been reported to be ~ 104 particles cm 3 (Hughes et al., 1998). As discussed by Hughes and co-workers, although these particles do not contribute a large amount (only ~ 1 /xg m 3) to the total particulate mass, about 1011 of these ultrafine particles will be deposited in one day in the respiratory

tract of a typical person breathing this air. Hence if the toxicological effects are determined primarily by the number of particles, rather than their mass, these ultrafine particles could ultimately prove to be quite important. A good example of the complexity of size distributions is found in the work of D o d d e t al. (1991). They measured the composition of various elements in particles collected in a rural area in Maryland as a function of size and observed in individual samples as many as four distinct modes between 0.09 and 1 /xm. For example, in one particular sample, there were multiple peaks at ~1 /xm containing Na, Fe, Ti, Mn, and A1, at 0.5-0.6 /xm containing As, Sb, Se, V, La, and Ce, and at 0.28/xm containing S and a single peak at ~ 0 . 1 / x m containing Na, Ga, Ti, A1, La, Ce, and Fe. Dodd and co-workers suggest the smallest particles represent primary emissions from high-temperature combustion sources, whereas the larger ones result from both secondary reactions (e.g., forming sulfate) and direct contributions from multiple sources. Based on numerous size distributions measured in air, various categories of tropospheric aerosols have been proposed. Table 9.2, for example, shows a typical set of categories and some of their associated characteristics. However, these should be taken merely as examples rather than as fixed categories since many aerosols will display characteristics of more than one category.

b. Atmospheric Aerosols and Log-Normal Distributions Ideally, one would like to describe various size distributions by some relatively simple mathematical function. Because there is no single theoretical basis for a particular function to describe atmospheric aerosols, various empirical matches have been carried out to the experimentally observed size distributions; some of these are discussed in detail elsewhere (e.g., see Hinds, 1982). Out of the various mathematical distribution functions for fitting aerosol data, the log-normal distribution (Aitchison and Brown, 1957; Patel et al., 1976) has emerged as the mathematical function that most frequently provides a sufficiently good fit, and hence we briefly discuss its application to the size distribution of atmospheric aerosols. Most readers will be familiar with the bell-shaped normal distribution plotted in Fig. 9.12. When applied to the size distribution of particles, for example, such a distribution is fully characterized by the arithmetic mean D and the standard deviation or, where cr is defined such that 68% of the particles have sizes in the range D cr In the log-normal distribution, the logarithm of the diameter D is assumed to have a normal distribution. (Either logarithms to the base 10 or loga-

A. PHYSICAL PROPERTIES T A B L E 9.2

359

Some Typical Tropospheric Aerosols and T h e i r Associated Properties a

Typical number concentration (cm 3)

Type Marine

Remote continental

Urban aerosol

100-400

Three modes: Dp 0.1, 0.1-0.6, and >0.6/xm; most (>95%) of particle mass but only 5-10% of total number in largest mode; large particles mainly sea salt aerosol, smallest are products (e.g., SO 2 ) of gas (e.g., DMS) to particle conversion

~ 104

Three modes centered at Dp ,-~ 0.02, 0.12, and 1.8/xm; includes products of gas-to-particle conversion and biological sources, e.g., pollens

105

Nonurban continental

Typical characteristics

~ 103

Desert

Variable depending on location

Polar

~15-150

Free troposphere

~30

Biomass burning

,-~104 close to source; 103 downwind

Three modes: nuclei, accumulation, and coarse; larger particles contain crustal elements (Fe, Si, etc.), smaller contain nitrate, sulfate, ammonium, and elemental and organic carbon and are formed by combustion processes and gas-to-particle conversion Three modes similar to urban aerosol Most are large, Dp > 1 /xm; contain desert soil elements such as Fe, Si, A1, and Ca Typical Dp ,-~ 0.15/xm; contain gas-to-particle conversion products such as sulfate and ammonium Monodisperse with Dp ,-~ 0.2-0.5/xm; nucleation major source; sulfate is major component Two modes often seen: Dp ,-~ 0.1-0.4 and Dp > 2/xm; smaller mode contains gas-to-particle conversion products (sulfate, nitrate, ammonium, organics)while larger mode has soil and ash particles

a From Pandis et al. (1995), Whitby and Sverdrup (1979), and Fitzgerald (1991).

rithms to the base e can be used, but since the latter is more common, we follow through the discussion using natural logarithms, ln.) This distribution is expressed as dN

d In D

NT

~

expl

In %

(lnD In DgN)2 1 2(ln %)2 , (B)

where N is the number of particles having diameters whose logarithms are between In D and In D d In D, N T is the total number of particles,__% is known as the geometric standard deviation, and Dg N is the geometric number mean diameter, which for this distribution is equal to the number median diameter, defined as the diameter for which half the number of particles are smaller and half are larger. A typical plot of Eq. (B) is shown in Fig. 9.13, where ( d N N T) d In D is plotted against the diameter on a logarithmic scale. The geometric number mean diameter, DgN, is related to the arithmetic mean of In (diameter): In Dg N

m

D

DI

F I G U R E 9.12 Meaning of standard deviation for a normal distribution. The hatched area represents 68% of total area under curve.

Enj In dj AfT .

(C)

Here nj is the number of particles in a group_whose diameters are centered around dj. Thus In Dg N is really a weighted value of In d, where the weighting is by the number of particles in that size interval.

360


the size distribution of the other properties of interest, that is, mass, surface, or volume. In these cases, the average size used to characterize the distribution is known as the geometric mass mean diameter, the geometric surface mean diameter, and geometric volume mean diameter, respectively. Like the geometric number mean size, the geometric mass mean diameter is given by

0.8

Cl c: "o

0.6

"-P 0.4 Z Z v

0.2

Emj In dj 0

In Dg M 1

10

F I G U R E 9.13 Frequency distribution scale) (adapted from Hinds, 1982).

curve (logarithmic size

The meaning of the geometric standard deviation o-g can be seen by referring to the meaning of standard deviation for a typical normal distribution. In this case, as shown in Fig. 9.12, the standard deviation is a measure of the spread about the mean (D) and is defined such that there is a 68% probability of a particle having a diameter in the range D o-, and a 95% probability of it having a diameter in the range D 2o-. o- can be calculated for the normal distribution using Eq. (D), m

D1

D,

(D)

m

where D is the diameter for which 50% of the particles have smaller diameters and 50% have larger diameters and DI( D o-) is the diameter for which there is an 84%( 50 68 2) probability that a particle will have a diameter equal to or less than this value. For a log-normal distribution, o-g is still a measure of spread of the distribution, but it has a slightly different definition because In D, rather than D, is assumed to have a normal distribution. For a log-normal distribution % is defined by Eq. (E): In o-g

In D 1

In D

D 1

In _ . D

(E)

Thus

~g

(O) ,

100

Particle diameter (gm)

o-

MT

D 1

(F)

and 68% of the distribution is between Dg o-g and Dgo-g. o-g is a dimensionless parameter that must be > 1.0 since D 1 > D. Although the log-normal distribution in Eq. (B)was given in terms of the distribution of the numbers of particles as a function of size, it can also be applied to

where mj is the mass of particles in group j having representative diameter dj and M T is the total mass in the sample, M T Emj. Assuming that the particles are all perfectly smooth and spherical with density p, then the volume of each particle in group j is vj ~-d~ 6 and the total volume of the group j is nj(rrd~ 6). The mass of particles in group j is then given by mj nj p(~rd~ 6). Thus, the total mass M T of all the particles in all groups is M T F,nj p(~r@ 6). Equation (G) becomes _ In Dg M

(~rp 6 ) E ( n j @ ) l n dj (~rp 6)Enjd~

Enjd~ In dj Enjd~ . (H)

Because mass and volume are directly related, with p as the constant of proportionality, the expression for the geometric volume mean diameter, Dgv, is the same as that for DgM given in Eq. (H). Similarly, the geometric surface mean diameter, Dgs, is defined as the average diameter weighted according to the surface areas sj in the various groups: Esj In dj In Dgs

ST

Enid 2 In dj F,njd2 ,

(I)

using area ~rd 2 and where S T is the total surface area of all particles in all groups. The surface mean diameter (as opposed to the geometric surface mean diameter) is also a very useful parameter, since

_ Ds

Esjdj ST

Enj@ ~njd 2

6 MT p ST

- ~.

(J)

Thus from the total mass ( g T) of the particle sample and its total surface area (S T) (which can be determined experimentally), the surface mean diameter of the particles can be calculated if their density ( p ) is known. An advantage of applying the log-normal distribution to atmospheric aerosols is that the value of the geometric standard deviation, o-g, is the same for a given sample for all types of distributions--count, mass, surface, and volume. It is only the value of the geometric mean diameter that changes, depending on the


TABLE 9.3 Values of the Constant b in the Hatch- Choate Equations for Converting the Count Geometric Mean Diameter to Mass, Surface, or Volume Diameters

0.8 m

o

0.6

C

Type of mean

0.4 ~-

diameter to be calculated

0.2 0

361

1

10

100

Particle diameter (gm) F I G U R E 9 . 1 4 Count and mass distributions for a hypothetical log-normal sample. The spread, O-g, of the two curves is seen to be the same, but the mean diameters associated with each are different (adapted from Hinds, 1982).

Mass Surface Volume

d (Number median diameter)

exp[ b(ln 2 Crg)], (K)

where d is the average diameter to be determined using a known value of the number median diameter, Crg is the geometric standard deviation, and b is a constant whose value is determined by the type of average diameter, d, which is to be calculated. Recall that for a log-normal distribution, the number median

mean

ba

ba

3 2 3

3.5 2.5 3.5

a See Eq. (K).

diameter is equal to the geometric number mean diameter, Dg N. Thus Eq. (K) becomes d DgN

property used for the distribution. This can be seen in Fig. 9.14, which shows a count and mass distribution for a hypothetical log-normal sample; the spread, O-g, is seen to be the same for each distribution, but the geometric mean diameters, Dg N and Dg M, are quite different. In practice, when one measures the size distributions of aerosols using techniques discussed in Chapter 11, one normally measures one parameter, for example, number or mass, as a function of size. For example, impactor data usually give the mass of particles by size interval. From such data, one can obtain the geometric mass mean diameter (which applies only to the mass distribution), and O-g, which, as discussed, is the same for all types of log-normal distributions for this one sample. Given the geometric mass mean diameter (Dg M) in this case and O-g, an important question is whether the other types of mean diameters (i.e., number, surface, and volume) can be determined from these data or if separate experimental measurements are required. The answer is that these other types of mean diameters can indeed be calculated for smooth spheres whose density is independent of diameter. The conversions are carried out using equations developed for fine-particle technology in 1929 by Hatch and Choate. These H a t c h - C h o a t e equations are of the form

Geometric mean

exp[b( ln20-g)].

(L)

Table 9.3 gives the values of b to be used for converting the count geometric mean diameter to the other types of geometric or mean diameters, respectively. While we have concentrated here on the geometric means weighted by number, mass, surface, or volume, respectively, other types of "average" diameters are also sometimes cited in the literature. One commonly used is the diameter of average mass, which is defined as the diameter such that the mass of this particle multiplied by the total number of particles gives the total mass. The geometric mean diameters discussed here can also be converted to these types of diameters using a form of the H a t c h - C h o a t e equations. The reader is referred to the discussion by Hinds (1982) for definition of these other types of diameters and their conversion to the geometric mean diameters discussed here. Whitby and co-workers have indicated that the many atmospheric aerosol size distributions that they have measured under a variety of conditions and at many locations can be fit reasonably well assuming three additive log-normal distributions corresponding to the Aitken nuclei range, the accumulation range, and the coarse particle range, respectively, as described earlier. Each of these log-normal distributions has its own characteristic value of Crg as well as, of course, average diameters. For example, Fig. 9.6 contains the number, surface, and volume distributions for a typical urban aerosol; these were calculated to be consistent with the sum of two, or in the case of the surface distribution, three, additive components. These components are shown by the dashed lines in Fig. 9.6. From these separate distributions, geometric mean diameters and Crg could be found assuming smooth spheres. Table 9.4 summarizes the parameters derived by Whitby and

362


TABLE 9.4

S u m m a r y of Parameters D g and o"g for the T h r e e Additive L o g - N o r m a l Distributions Characterizing Data in Fig. 9 . 6 "

Log-normal distribution Type of distribution used

Mode

crg

~b (Ixm) m

Number

DgN DgN

Aitken nuclei

0.013

1.7

0.069

2.03

m

Accumulation Surface

Aitken nuclei Accumulation Coarse particle

Volume

Aitken nuclei Accumulation Coarse particle

Dgs 0.023 Dgs 0.19 DgS 3.1

c c 2.15

Dgv 0.031 Dgv 0.31 Dgv 5.7

c c d

a From Whitby and Sverdrup (1980). b Note that the diameter--number, surface, or volume--changes depending on the type of distribution used, whereas O-g remains constant for each of the three modes. See text. c Same as derived from the number distribution. d Same as derived from the surface distribution.

Sverdrup (1980) based on the normal distributions in the three size ranges in Fig. 9.6. The examples used so far are generally based on surface measurements. The particle concentration and size distribution also depend on altitude. Typical vertical distributions are discussed, for example, by Jaenicke (1992).

then moves with a constant velocity known as the

terminal velocity. One can apply Stokes' law to atmospheric particles to calculate how fast they will settle out of the air when subjected to gravity alone. Thus the terminal settling velocity occurs when the frictional and gravitational forces are balanced, that is, FR

3. P a r t i c l e M o t i o n One of the important properties of particles that contributes to both the observed size distribution and the number concentration of aerosols in the atmosphere is the motion they undergo when suspended in air. This includes gravitational settling and Brownian diffusion.

rrD3pg

In the free troposphere, particles are subjected to gravitational forces. They also can be subjected to electrical forces in nature as well as in the course of detection and measurement. When such forces are applied, the particle moves relative to the gas and hence is subjected to a resistance force. Stokes' law gives the force (F R) acting on smooth spherical particles due to the laminar flow of air over them:

3rrrlvD.

(M)

rl is the gas viscosity, v is the particle velocity relative to the gas, and D is the particle diameter. When a force such as gravity is applied to the particle, it speeds up until the frictional force equals the applied force; it

mg,

(N)

where m is the mass of the particle and g is the acceleration due to gravity (9.8 m s 2 at sea level). One can apply the relationship between mass, volume (~rD 3 6), and density ( p ) of the particle. Equations (M) and (N) combine to give Eq. (O):

a. Gravitational Settling

FR

fgravity

3rrrlvD

or

D20g v

~(for 18~/

D > 1.5 /xm).

(O)

The settling velocity thus increases with the density of the particle and with the square of its diameter. In developing Eq. (O), the buoyancy effect of air, which tends to lower the effective particle density, has been ignored since it is much smaller than the particle density; it can be included if desired by replacing p by (Pp &ir), where pp is the particle density and &ir is the air density ( 1 . 2 • 10 3 g c m 3 a t 2 0 ~ latm pressure).


The expression given in Eq. (O) for the terminal settling velocity only applies to particles with diameters > 1.5 /xm because its derivation is based on the assumption that the relative speed of the air at the surface of the particle is zero. However, as the particle becomes smaller, the air molecules appear less as a continuous fluid and more as discrete molecules separated by space through which the particles can "slip." The net effect is that the particles can move faster than predicted by Eq. (O) due to this slipping between gas molecules. To correct for this effect, a correction factor must be applied to the resistance force predicted by Stokes' law, Eq. (M). The correction factor, C, is a number greater than 1. Thus Eq. (M) is modified to

3~luD FR

(p)

C

363 10 9 8 7

o

6

o (3

5-

m

m

9

-

9

m

v

.

to m

43m

o .

m

2

CO

0.01 0.1 1

and the settling velocity becomes u

m

0.02 0.2 2

0.04 0.4 4

0.06 0.6 6

0.08 0.1 0.8 1.0 8 10


D2pgC

~ . 18~

(Q)

F I G U R E 9.15 Cunningham correction factors to be applied in calculating terminal settling velocities of particles (adapted from Hinds, 1982).

The correction factor, first derived by Millikan (1923), is given by Eq. (R): C

1

'E

D 2.514

0.800exp

(

0.55--/-

. (R)

This is often called the Cunningham correction factor, although Cunningham's original correction was of the form 1 A(1 D) and did not include the exponential term. 1 is the mean free path between air molecule collisions, defined as the average distance traveled between collisions with another molecule; it can be easily calculated using simple kinetic molecular theory, and for air at 1 atm pressure and 20~ 1 is 0.066 /xm. For large particles, D > 1.5 /xm; the term in brackets in Eq. (R) is 0.10 and hence Eq. (O) is a good approximation. For particles between ~0.1 and 1.5 /xm, the third (exponential) term in Eq. (R) is small relative to 2.5, and hence the Cunningham correction factor can be approximated by 2.1541 C

1

D

(for0.1 /xm

D

1.5 /xm). (S)

Figure 9.15 shows the Cunningham slip correction factors for air at 1 atm pressure and 20~ for the smallest particles, there is a significant correction to the speeds calculated using Eq. (O). The assumptions inherent in the use of Stokes' law (e.g., relatively low speeds), which normally apply to atmospheric particles, are discussed in more detail by Hinds (1982) and Fuchs (1964). For our purposes, we

need only recognize that the settling velocities for atmospheric particles calculated using Eq. (O) or (Q) are large for particles > 10 /~m in diameter. Figure 9.16 shows the settling velocities of spherical particles with p 1 g cm 3 in still air at 0~ and 1 atm pressure as a function of particle diameter. It is seen that particles with diameters of the order of 10 /~m or greater have settling velocities >0.1 cm s 1 and hence will settle out of the atmosphere relatively rapidly. However, those particles with diameters 1 /~m will remain suspended for long periods of time and hence can participate in atmospheric transformations. b. Brownian

Diffusion

Small particles do not settle via gravity at a significant rate, but they do undergo Brownian diffusion. The classic example of Brownian diffusion is the random zigzag motion of smoke particles in air that can be observed because of light scattering by the particles. In fact, in the absence of convection, particles 0.1 /~m in diameter are transported largely by Brownian diffusion, which is primarily responsible for the rapid coagulation of particles in the Aitken nuclei range. The rates of Brownian diffusion can be quantified by considering a box of cross-sectional area 1 cm 2 having N particles cm 3 in one section (A) on the left, while the right section (B) is initially empty. Particles (or gas molecules) will always tend to diffuse from a region of higher concentrations to one of lower concentrations.

364

lo3p


1021

,o,p

/

,001 /

TABLE 9.5 Cumulative Deposition of Unit Density Particles onto a Horizontal Surface from Unit Aerosol Concentrations a during 100 s by Diffusion and Gravitational Settling b

/

Cumulative deposition Diameter (l~m) 0.001 0.01 0.1 1.0 10 100

F

Diffusion (number em 2) 2.5 0.26 2.9 • 10 5.9 • 10 1 . 7 • 10 5.5 • 10

2 3 3 4

Settling (number em 2)

Ratio diffusion settling

6.5 • 10 5 6.7 • 10 4 8.5 • 10 3 0.35 31 2500

3.8 • 104 390 3.4 1.7 • 10 2 5.5 • 10 5 2.2 • 10 7

a This assumes an aerosol concentration of 1 particle cm side the gradient region. b F r o m Hinds (1982).

10-1 100 101 102 103 104 Diameterof particle(l.tm) F I G U R E 9 . 1 6 Settling velocities in still air at 0~ and 760 Torr pressure for particles having a density of 1 g cm 3 as a function of particle diameter. For spherical particles of unit density suspended in air near sea level, the Stokes law applies over a considerable range of particle sizes, where the line is straight, but a correction is required at the particle size extremes (adapted from LBL, 1979).

The rate at which they diffuse depends on the concentration gradient, d N dr; the larger the gradient, the faster the rate of diffusion. This is the basis of the well-known Fick's first law of diffusion" dN

J

D ~.

dr

(T)

J is the flux of particles crossing a 1-cm 2 plane in 1 s (i.e., number cm 2 s 1). The constant D is known as the diffusion coefficient and is simply the proportionality constant relating the flux to the concentration gradient. (Fick's first law applies, of course, not only to particles but also to gas and liquid molecules.) Intuitively, one might expect that the rate of diffusion would increase with temperature and decrease with increasing gas viscosity and particle size. Indeed, this is observed to be the case because all these parameters contribute to the diffusion coefficient D, given by kTC o

~

.

3vr~/D

(u)

k is the Boltzmann constant (1.38 • 10 23 j deg 1), T the temperature (in K), C the Cunningham correction factor given by Eq. (R), r/ the viscosity of the gas in which the particles are suspended, and D the particle

3 out-

diameter. The reader is referred to Hinds (1982) for a derivation of Eq. (U). For larger particles where C ~ 1 (i.e., the slip correction is negligible), the rate of diffusion varies inversely with the particle diameter. For very small particles, the rate varies with 1 D 2, which contributes to making diffusion a major transport mechanism for particles 0.1 /xm. It is this Brownian diffusion that helps to carry small particles through the boundary layer to surfaces where they may stick on impact. The relative importance of Brownian diffusion and gravitational settling in the deposition of particles may be seen by calculating the total deposition of particles onto a horizontal surface by these two processes in a given period of time under certain conditions. Table 9.5 shows the results of such a calculation for particle diameters from 0.001 to 100 /xm, assuming spherical particles of unit density with a constant concentration of 1 particle cm 3 outside the gradient region; also shown is the ratio of the number of particles deposited by diffusion compared to the number deposited by gravitational settling. At a diameter of ~0.2 /xm, the two mechanisms become equal, with diffusion greatly exceeding gravitational settling for particles in the Aitken nuclei range. (Note that if the mass deposited were calculated, the results would be quite different.) Other factors also come into play in laboratory systems. For example, McMurry and Rader (1985) have shown that particle deposition at the walls of Teflon smog chambers is controlled by Brownian and turbulent diffusion for particles with Dp 0.05 /xm and by gravitational settling for particles with Dp > 1.0 /xm. However, in the 0.05- to 1.0-/xm range, the deposition is controlled by electrostatic effects; Teflon tends to


hold an electrostatic charge, which leads to rapid deposition of charged particles and deposition rates almost an order of magnitude faster than expected if these effects of charge are not taken into account. It should be kept in mind that these calculated rates of diffusion and gravitational settling are only applicable to still air. In fact, in the atmosphere the air is rarely still and is usually undergoing some degree of turbulent motion. In this case, the transport of particles becomes more complex and faster due to the velocity gradients and contorted patterns of air flow; however, a discussion of this is outside the scope of this book.

and Absorption and Relationship to Visibility Reduction 4. L i g h t S c a t t e r i n g

As discussed in Chapter 3, solar radiation passing through the atmosphere to the earth's surface is both scattered and absorbed by gases and particles. The intensity of radiation striking the surface can be expressed in the form of a Beer-Lambert law: e bextL

(V)

I0 where I 0 and I are the incident and transmitted light intensities, respectively, L is the path length of the light beam, and bext is known as the extinction coefficient and has units of (length) 1. This extinction coefficient, representing the total reduction in light intensity due to scattering and absorption of light by gases and particles, is the sum of two terms, bex t

bg

bp,

(W)

where bg is the extinction due to gases and bp that due to particles. Each of these terms can be broken down into contributions from light scattering and absorption so that Eq. (W) becomes bex t

bag

bsg

bap

bsp ,

12

~-E o

%

I~1 Scattering by Air Molecules [-I Absorption by NO 2 Gas [] Absorption by Particles m Scattering by Coarse Particles D Scattering by Fine Particles

~'E" 9 ~b "- 6 {/)

5 11 17 23 29 4 10 16 22 28 3 July

August

9 15 21 27

September, 1984

Individual contributions of scattering and absorption of light at 550 nm by gases and particles in Pasadena, California, during the summer of 1984 (adapted from Larson and Cass, 1989).

F I G U R E 9.17

Their

a. Light Scattering and Absorption

I --

3 65

(X)

where bag and bsg are the light extinctions due to absorption and scattering by gases, and bap and bsp are those due to absorption and scattering by particles. Scattering and absorption of light by gases have already been discussed in Chapters 3 and 4. In terms of the absorption of visible light in the troposphere by gases, only NO 2 is believed to contribute significantly. However, light absorption by NO 2 is usually much less than the total light scattering and absorption by particles. Thus, when NO 2 is present in sufficient concentrations to absorb light, the atmosphere usually also contains relatively high concentrations of other pollu-

tants, including particles. As a result, light scattering and absorption by particles usually exceed light absorption by NO 2 in most situations. For example, in studies of Denver's "brown cloud," it was shown that of the total extinction excluding Rayleigh scattering by gases, 7% was due to light absorption by NO 2 and 93% was due to light scattering and absorption by particulate matter (Groblicki et al., 1981; Waggoner et al., 1983b). Similarly, Fig. 9.17 shows the calculated individual contributions of various gases and particles to the scattering and absorption of light at 550 nm during a typical summer in Pasadena, California (Larson and Cass, 1989). Even in this area of relatively large NO X sources, the contribution of NO 2 to the total light attention (expressed as bext) is small. For reasons to be discussed in more detail, scattering by fine particles predominates, with a smaller contribution due to light absorption by the particles. Historically, the symbol b has been used both for the total extinction coefficient and for the individual contributions due to light absorption and scattering by particles and gases. However, the Radiation Commission of the International Association of Meteorology and Atmospheric Physics (IAMAP, 1978) has recommended use of the symbol o- instead. Thus the reader may find in the literature the terms (Tag, O[sg, (Tap, and ~p, respectively, rather than bag, bsg, bap, and bsp. However, in keeping with the common practice in the literature in this area and to avoid confusion with light absorption cross sections by gases defined in Chapter 3, we use the symbol b. The scattering of light by particles falls into three regions depending on the size of the particles relative to the wavelength ( a ) o f the light: (1)particle diameter D a, known as Rayleigh scattering, (2) D ~ a, known as Mie scattering, and (3) D >> a. Since we are concerned at the earth's surface primarily with visible and near-ultraviolet light in the actinic region, that is, 290

366


A 750 nm, the first case corresponds to particles with D 0.03 /xm and the third to particles with D 10 /xm. Particles with sizes between these two extremes fall in the second category where D ~ A; as we have seen, this is the most important size regime for atmospheric particles. A common convention used when discussing light scattering as a function of particle size is to define a dimensionless size parameter a, which is the ratio of the circumference of the particle to the wavelength of the incident light: oz

~-D ~. A

(Y)

Very small particles (c~ 3) behave like gaseous molecules in scattering light and hence produce Rayleigh scattering, described in Chapter 3. Because the particles or molecules undergoing this type of scattering are small relative to the incident wavelength, the entire species is subjected at any instant of time to what appears to be a uniform electromagnetic field; this creates a dipole that oscillates with the changing electromagnetic field of the light wave and reradiates the energy in all directions. Thus Rayleigh scattering is symmetric in the forward and backward directions relative to the incident light beam and, as we saw in Chapter 3, varies as a 4. Very large particles, on the other hand, that is, those with D >> a (c~ >> 3), undergo geometric scattering, where the light beam refracted through the particle can be treated using classical optics. Between these two regimes where D ~ a (c~ ~ 3), much more complex light scattering occurs, known as Mie scattering. Because particles undergoing Mie scattering have dimensions of the same order as the wavelength of the incident light, the electromagnetic field of the light wave is not uniform over the entire particle at one instant of time, and a three-dimensional charge distribution is set up in the scattering particle. In 1908 Mie developed the solutions for scattering of light of wavelength a by a homogeneous sphere of diameter D. As shown in Fig. 9.18, light is considered to be incident on the sphere and to be scattered at various angles 0 to the direction of the unscattered beam. The incident and scattered beams are shown as the combination of two independent polarized beams: one (I1) has its electric vector perpendicular to the scattering plane defined by the incident and the scattered beams, and the other (I~) is parallel to it. The intensity of light at a distance R and a scattering angle 0 from the particle is given by IoA2(il I ( 0 , R)

i11)

8,7r 2R 2

,

(Z)

Scatterin

Incident~ light,Io

~

Scattered light,I(e, R)

F I G U R E 9.18 Diagram showing scattering angle, scattering plane, and the polarized components of scattered light. (From Hinds, W. C. Aerosol Technology. Copyright 9 1982 John Wiley & Sons, Inc. Reprinted by permission of John Wiley & Sons, Inc.)

where I 0 is the intensity of the incident light beam (taken as unpolarized) and i~ and i~ are known as the Mie intensity parameters for the perpendicularpolarized and parallel-polarized components of the scattered light, respectively. The Mie intensity parameters are a complex function of the refractive index (m) of the scatterer, the size parameter (a), and the scattering angle (0). For further mathematical details and descriptions of Mie scattering, the reader is referred to the books by Van de Hulst (1957), Kerker (1969), and Bohren and Huffman (1983). The refractive index of a material, m, is defined as the ratio of the speed of light (c) in a vacuum to that (v) in the material, that is, m c v. Because light travels more slowly in materials than in air, m > 1. The index of refraction, rna, of materials that absorb light as well as scatter it is expressed in the form of a complex number (see Hinds, 1982) in a

mr(1

ai),

(AA)

where i ~/ 1, m r is the real refractive index, and a is a constant that depends on the absorption coefficient of the material as well as to the wavelength. Values of the index of refraction of A 589 nm of some materials either found in the atmosphere or used to calibrate instruments that measure particle size using light scattering (see later) are given in Table 9.6. A typical refractive index for a dry aerosol that absorbs light in the atmosphere is 1.5 0.02i (Covert et al., 1980). Unlike Rayleigh scattering, which occurs equally in the forward and backward directions, Mie scattering is predominantly in the forward direction, except for the smallest particles. This can be seen in Fig. 9.19, which shows the Mie intensity parameters i 1 and i1~ as a function of the scattering angle 0 for three different values of the size parameter c~ defined by Eq. (Y), assuming the droplets are composed entirely of liquid water (i.e., m 1.333). For c~ 2.0 and 10 (i.e., D A

A. PHYSICAL PROPERTIES T A B L E 9 . 6 I n d e x of R e f r a c t i o n at 5 8 9 n m for Some Species F o u n d in the A t m o s p h e r e or U s e d f o r Instrument Calibrationa Species

Vacuum Water vapor Air Water (liquid) Ice Rock salt Sodium chloride in aqueous solutions Sulfuric acid in aqueous solutions Benzene c~-Pinene d-Limonene Nitrobenzene Dioctyl phthalate Oleic acid Polystyrene latex Carbon Iron Magnetite ( F e 3 0 4) Copper

3 67

1041 103

I n d e x of r e f r a c t i o n

1.0 1.00025 1.00029 1.333 1.309 1.544 1.342-1.378 b 1.339-1.437 c 1.501 1.465 1.471 1.550 1.49 1.46 1.59 1.59-0.66i d' f 1.51-1.63i d 2.58-0.58i e 0.62-2.63i d

a Data from the Handbook of Chemistry and Physics, unless otherwise noted (~20-25~ b For solution densities from 1.035 to 1.189. c For solution densities from 1.028 to 1.811. d From Hinds (1982). e From Huffman and Stapp (1973). f A t A 491 nm.

0.64 or 3.2), both i~ and i~ fall from a maximum value at 0 0 ~ as 0 increases. For smaller particles with a 0.8 (i.e., D A 0.26), i1~ initially falls as 0 increases but then rises as 0 approaches 180 ~ corresponding to backscattering; for these particles i~ decreases only slightly from 0 ~ to 180 ~ However, in all cases, i~ and i~1, and hence the scattered light intensity Eq. (Z), show their maximum values at 0 0 ~ corresponding to forward light scattering. The variation of scattered light intensity with 0 as typified by Fig. 9.19 clearly becomes more complex as the particle size increases, with sharp oscillations seen at a 10. However, recall that this is for a spherical homogeneous particle of a fixed size and for monochromatic light (e.g., a laser); when the particle is irregular in shape, these oscillations are far less prominent. This is also true for a group of particles of various sizes, that is, a polydisperse aerosol, where the overall scattering observed is the sum of many different contributions from particles of various sizes. Finally, nonmonochromatic light and fluctuations in polarization also help to smooth out the oscillations.

i(D

E

102 ,,

..

,-,

10

Q. ffl

.~_

10 -1

-

10 -2

10-3 10-4

%%

f/ \

0

I

20

I

40

I

60

\

\

/

I \, /I

80

/

/

I

I

I

100 120 140 160 180

Scattering angle (deg) F I G U R E 9 . 1 9 Mie intensity parameters versus scattering angle for water droplets (rn 1.333) having c~ 0.8, 2.0, and 10.0. Solid lines are il, and dashed lines are ill (adapted from Hinds, 1982).

The dependence of Mie scattering on particle size can be seen in Fig. 9.20, which shows the sum of the Mie intensity parameters (i 1 i1~) as a function of the size parameter c~ for two scattering angles, 0 30 ~ and 0 90 ~ respectively. It is seen that Mie scattering generally increases with size over this range of values of c~ and, as seen in Fig. 9.19, scattering is more pronounced in the forward direction (i.e., at smaller values of 0). It is noteworthy that in Fig. 9.20 the function becomes smooth and approaches a variation with D 6 as the size parameter decreases toward small values. This is expected, since in the limit of very small particles or molecules, Mie theory reduces to Rayleigh scattering, which, as seen in Chapter 3, varies with D 6. It should be noted that scattering of light by particles can be measured using remote sensing techniques on satellites, from which such parameters as total aerosol optical thickness i.e., the exponent (bext L ) in I I 0 exp( boxtL), Eq. (V), albedo, etc. can be determined. However, as discussed in detail by Mishchenko et al. (1995), application of conventional Mie theory can lead to significant errors in the aerosol optical thickness if the particles are not spherical, as is assumed in development of Mie scattering theory. Particles can also absorb light in the atmosphere; the radiant energy absorbed is then converted to heat. As discussed later, graphitic carbon is believed to be the species responsible for most of the light absorption occurring in typical urban atmospheres, although there

368

9. PARTICLESIN THE TROPOSPHERE

1000 -

= 30 ~

.m --Ira .m

E

100

m

i-, ..m if} cQ)

10 v t.m

.i-,

if} Q) > .m

1.0

m Q)

rr

0.1 1.0

10

100

Size parameter, o~

I 0.17

I 1.7

I 17

Particle diameter for ~, = 0.52 gm

FIGURE 9.20 Relative scattering Mie intensity parameter (i 1 i11) versus size parameter for water droplets (rn 1.33)at scattering angles of 30~ and 90~ (adapted from Hinds, 1982).

is evidence that organics may also contribute. Iron oxides such as hematite and magnetite also absorb light strongly (Table 9.6), their concentrations are not believed to be sufficiently high to contribute significantly to light absorption in urban areas compared to carbon. However, as discussed in Chapter 14, on a global scale such particles can make significant contributions to light absorption. Light scattering by particles is usually measured with

I Power supply I

an integrating nephelometer (Fig. 9.21). Light scattering through all angles (commonly known as bscat) due to Rayleigh scattering by gases as well as scattering by particles is measured by this device (Ruby, 1985), i.e., bscat bsg bsp. However, in the atmosphere under many conditions, bsp >> bsg , so that bscat measured using a nephelometer effectively measures the light scattering due to particles. In this instrument, air flows slowly through the measuring chamber where a flash lamp produces a pulse of white light. A photomultiplier measures the intensity of the scattered light. Rayleigh scattering by gases can be used to calibrate the instrument. For example, the scattering coefficient of He is calculated to be 3.0 • 10 7 m 1 whereas that of clean air is 2.8 • 10 5 m 1 at 460 nm (Charlson et al., 1967). Values of bsp , the extinction coefficient for light scattering by particles, measured in ambient air using an integrating nephelometer (Fig. 9.21) range from 10 3 m ~ in highly polluted urban areas to ~ 10 7 m 1 in remote locations (e.g., Nyeki et al., 1998). The extinction coefficient for particulate absorption, bap , is generally from ~10 4 to 10 S m ~ (Bodhaine, 1995; Nyeki et al., 1998). However, the relative contributions of light scattering and absorption at any particular location will clearly depend on the nature of the sources; this is discussed in more detail later in this chapter.

b. Relationship of Light Scattering and Absorption to Visibility Reduction One of the most evident manifestations of anthropogenic air pollution is the production of a haze which causes a reduction in visibility, that is, in visual range. Visual range is defined as the distance at which a black object can just be distinguished against the horizon. Two factors enter into visual range: visual acuity and contrast. In the daytime atmosphere, particles reduce the contrast perceived by an observer by scattering light from the object out of the line of sight to the

Aerosol outlet

I tl

~lClean air Photo multiplier

- Opal glass

7' I I I I ~176 I .....

Scattering volume

,

Collimating disks I

Amplifier

II

Flash tube power supply

Recorder

J

9 I I I t

Aerosol inlet

Parson's optical black

Clean air purge

FIGURE 9.21 Schematic diagram of an integrating nephelometer (from Charlson et al., 1967).


observer's eyes; simultaneously, sunlight is scattered into the line of sight, making dark objects appear lighter. The result is a decrease in the contrast between the object and the horizon. At night, scattering of light out of the visual path decreases the contrast and hence the source intensity becomes a factor in visual range as well. The Koschmieder equation has been shown to approximate the change in contrast of an object with distance away from an observer (Middleton, 1952); note that it has a form similar to that of the B e e r - L a m b e r t law: C

(BB)

e bextL.

Co In Eq. (BB), C O is the contrast relative to the horizon (or background) of an object seen at the observation point itself, that is, at a distance L 0, and C is the contrast at the distance L. The contrast is defined as the ratio of the brightness of the object (B o) to that of the horizon or background (B n) minus one: c

B o

BH

1.

(cc)

For example, a black object at zero distance has a brightness of zero (e.g., it absorbs all the visible light) and hence has a C o of 1.0. bext in Eq. (BB) is the total extinction as defined in Eq. (V). Observers typically can differentiate objects on the horizon if C C o 0.02-0.05. A contrast of 0.02, corresponds, using Eq. (BB), to a visual range Va of ln(C o C) VR bext

3.9 ~. bext

(DD)

369

contribution from light absorption by suspended particulate matter. It should be noted that the definition of visual range in Eq. (DD) is not always in accord with visual ranges reported from qualitative sightings of surrounding landmarks, as is done, for example, at airport observation towers. There are a number of factors that might influence this, such as the targets not being black or there being differences between various observers. In general, the airport visual ranges are less than those predicted from Eq. (DD)(Stevens et al., 1983; Waggoner, 1983; Lodge, 1983). Indeed, on the basis of airport visual range observations, Ozkaynak et al. (1985) suggest that the use of the coefficient 3.9 in Eq. (DD) is optimistic and that a value less than half that, 1.8, may be more appropriate in urban areas. Other methods of assessing visibility in the atmosphere are discussed by Richards et al. (1988, 1989). Most of the light scattering by particles in the atmosphere is due to particles in the size range 0.1-1 /xm as shown by calculations carried out during World War II for screening smoke particle sizes (Sinclair, 1950). This can be seen in Fig. 9.22, which shows the scattering coefficient of a single particle per unit volume as a function of the particle diameter for spheres with a refractive index of 1.50 and light of wavelength 550 nm. The portion of the total extinction coefficient due to particle scattering, bsp, can be obtained by combining the curve in Fig. 9.22 with the particle volume size distribution, that is, by the curve of AV A log D vs log D.

1203

E

:=L 04

For a contrast of 0.05, Va 3.0 bext. In a clean, particle-flee atmosphere, some light scattering occurs due to the Rayleigh scattering by gases. For this scattering, bsg 1.5 • 10 5 m 1 integrated over the solar spectrum at sea level and 25~ (Ouimette et al., 1981). If light absorption by gases is negligible, as it usually is unless significant concentrations of NO 2 are present, then bext bsg 1.5 • 10 5 m 1, and the visual range is ~ 200-260 km for contrasts of 0.05-0.02. While this is an approximation that depends on the nature of the object and on the observer, it does give some idea of the visual range that can be expected in clean air. Visual ranges can vary from hundreds of kilometers in remote areas to only a few kilometers in heavily polluted urban areas. In the latter case, most of the loss in visibility is due to light scattering, with some

E

m

0

E

m

m

>o {.-

6-

i.0 m

0 .0 m

{3.

3-

0

Q.

0

0.01

I I

0.1

1.0

10


FIGURE 9.22 Scattering coefficient per particle divided by particle volume plotted as a function of diameter. The particles are assumed to be spheres of refractive index 1.50 and the light has A 550 nm (adapted from Waggoner and Charlson, 1976).

370


03

This is supported by the correlation between the total aerosol volume of particles with diameters in the 0.1- to 1-/zm range and the experimentally determined values of bsp obtained using a nephelometer in many studies (e.g., Fig. 9.24). The slopes of lines such as that in Fig. 9.24, however, depend critically on the nature and history of the air mass and can vary by more than a factor of 10 from clean, nonurban air to highly polluted air in the vicinity of sources. For example, Sverdrup and Whitby (1980a) have shown that the ratio of submicron aerosol volume to bsp, which corresponds to the slope of the line in Fig. 9.24, varies from 5 to 80, depending on the nature of the air mass. In addition, the correlation between bsp and fine particles is usually not as clear-cut as seen in Fig. 9.24. Consistent with the relationship between the aerosol fine particle volume and the particle scattering coefficient, a number of studies have shown that the fine particle mass and bsp are also related. Figure 9.25 shows the scattering coefficient observed in studies in Denver, Colorado, by Groblicki and co-workers (1981) as a function of the observed mass in the fine and coarse particle ranges, respectively. It is seen that a good linear relationship exists between bsp and the fine particle mass (FPM) but not between bsp and the coarse particle mass. This has been observed in a number of areas ranging from pristine to industrial, with the ratio of the scattering coefficient to the fine particle mass concentration (bsp FPM) being approximately 3 in many areas (Waggoner et al., 1981; Conner et al., 1991). Because of the dependence of Mie scattering on the refractive index and hence chemical composition of the particles, one would expect the light scattering coeffi-

E 121.

E

>o

._o Gi 13_

"4--'

t'-.m

if}

"O Gi O3 O .--I

I 0.01

0.1

1.0

10

I 0.01

0.1

1.0

10

Gi if)

Gi .ca m

CI.

o

g,

._1

Particle diameter (gm) 9 . 2 3 (a) Aerosol particle size distribution measured at Pomona during the 1972 State of California Air Resources Board A C H E X program. (b) Calculated optical scattering by particles, bsp , for measured size distribution (adapted from Waggoner and Charlson, 1976). FIGURE

For example, Fig. 9.23a shows the measured volume distribution of one ambient aerosol sample. When this volume distribution is multiplied by the size distribution of the scattering coefficient per unit volume in Fig. 9.22, one obtains the calculated curve for light scattering in Fig. 9.23b. It is seen that the particles in the 0.1to 1-/xm-diameter range, that is, in the accumulation mode, are clearly expected to predominate the light scattering.

200 m,

E

O 03

.~E

100

--

9 O0

O

00

~

000

o-io

0 aD~ 0

I 5

I 10

I 15

I 20

bsp (in units of 10 "4 ml)

F I G U R E 9 . 2 4 Plot of measured aerosol fine particle volume (including only those particles of 0.1- to 1.0-/zm diameter) versus measured bsp. Measurements were part of the State of California Air Resources Board A C H E X program (adapted from Waggoner and Charlson, 1976; data supplied by Dr. Clark of North American Rockwell).


5

E ~, 4

?

_

b "6

0

o 9

"6 3

9

--

9

9 9

--

if) ..4..-, m t-

~ 1 7~6

2

9 # 0 o~

9 ~176

1

0

50

9

9 9

9

0 9 o%0

..Q

0

3 71

9

o 9 ~176

9

"~'g" ";" . 9 .~.-.%.. ~:.-'.. ~,. 9 . 100

I

I I I I 150 200 250

0

0

50

100

150

I

I I 200 250

Coarse particle mass (gg m 3)

Fine particle mass (gg m 3 )

FIGURE 9.25 Correlation of bsp with fine and coarse particulate mass (adapted from Groblicki et al., 1981).

identifying certain species (e.g., sulfate, nitrate, and organics) as being important in terms of light scattering (e.g., see Groblicki et al., 1981; Appel et al., 1985; Vossier and Macias, 1986; Solomon and Moyers, 1986; McMurry et al., 1996; E a t 9 et al., 1996; Laulainen and Trexler, 1997; and Mclnnes et al., 1998). Values of a i are usually in the range between 1 and 20, with a typical value of ~ 5 m 2 g 1. Table 9.7 gives some values of a i reported for various c o m m o n aerosol constituents. However, as discussed in detail by Sloane (1986) and White (1986), Eq. (EE) only holds if all of the species act independently. There are several reasons to believe that in many cases they do not. For example, if one species is removed from a particle, not only will its contribution ( a i m i) disappear but the particle size may

cient bsp to depend on the particular particle composition. Many studies have probed this relationship, coming at it from several different directions. One common approach has been to measure the chemical composition of particles and then test for empirical relationships between light scattering and chemical composition. In this case, a best fit of the measured bsp to the chemical composition of the particles is often found using an equation of the form bsp

ao

~aiM

i.

(EE)

where M~ is the mass concentration of the ith chemical species and a i is the light scattering coefficient for that species per unit concentration. Such a relationship has been applied in many studies and has proven useful in

TABLE 9.7

Some Reported Values of the Light Scattering Coefficients (a i ) for Common Constituents of Atmospheric Particles a

Compound Elemental (black) carbon Sulfate

Nitrate Organics

a i (m

2

g

1)

0.45-1.4 5.2-13 b 3-7 2.1 2.2-3.2 3.2-13.5 1.8 0.5-6.0 3-7 1.1 0-4.6

a For dry particles unless otherwise specified. bAt 80% RH.

Reference

Petzold et al., 1997 Howell and Huebert, 1998 McMurry et al., 1996 Eat 9 et al., 1996 Hegg et al., 1995 White, 1986 Eat 9 et al., 1996 White, 1986 McMurry et al., 1996 Eat 9 et al., 1996 White, 1986

372

9. PARTICLES IN THE TROPOSPHERE 1

m

15 14 , -i

13

04

11 10

12

E v o o3 i i z

9 8 7

6 5 4

3

I I I I I I 20 30 40 50 60 70 80

RH (%) FIGURE 9 . 2 6 Calculated light scattering coefficient for NH 4 HSO 4 droplets as a function of relative humidity (adapted from Sloane, 1986).

also change. As seen from the earlier discussion of Mie scattering and from data such as those in Fig. 9.22, any such change in particle size will also change its contribution to light scattering. Another potential effect is if one species affects the concentration of others due to chemical interactions. If the concentration of a species that is hygroscopic changes, the water content and average index of refraction for the particle may change simultaneously. For example, a reduction of about 7% in the index of refraction due to increased hydration can lead to a change in the scattering efficiency of as much as 34% (Hegg et al., 1993). Figure 9.26 shows the calculated light scattering coefficient for N H 4 H S O 4 (aNn4nso4) , a common constitutent of atmospheric particles, as a function of relative humidity (Sloane, 1986). The contribution of this one compound to light scattering increases from 3.7 to 15.8 m 2 g 1 as the R H increases from 20 to 80%. This key role of water in visibility reduction has been increasingly recognized as being very important. It is not surprising, given the changes in both the size and refractive index that water uptake causes. Atmospheric aerosols are hygroscopic, taking up and releasing water as the R H changes (see also Section C.1) because some of the chemical components are themselves deliquescent in pure form. For example, sodium chloride, the major component of sea salt, deliquesces at 298 K at an R H of 75%, whereas ammonium sulfate, ( N H 4 ) 2 S 0 4 , and ammonium nitrate, NH4NO3, deliquesce at 80 and 62% RH, respectively. (See Table 9.16 for the deliquescence points of some common constituents of atmospheric particles.) De-

pending on their size and chemical composition, such particles can act as cloud condensation nuclei (see Chapter 14.C.2), taking up enough water to form cloud droplets. The water content of aerosols varies widely. In the Los Angeles area, for example, a range of water contents in particles from a few percent up to 90 wt% has been observed (Ho et al., 1974; Sverdrup and Whitby, 1980a; Witz et al., 1988; Li et al., 1997), whereas in Sapporo, Japan, 0.4-3.2% of the fine particles (D 2 / z m ) w a s reported to be water (Ohta et al., 1998). In accumulation mode particles over the North Atlantic Ocean, an average of 31% of the particle mass was water (Kim et al., 1995). The large amounts of water in particles in the high-humidity marine boundary layer have also been documented using a combination of atomic force microscopy (AFM) and transmission electron microscopy (TEM); thus, P6sfai et al. (1998) showed that the volume of particles measured under ambient conditions using A F M was up to four times that of the same particles under the vacuum conditions of the T E M measurements. The uptake of water with increasing R H causes an increase in both mass and radius and a decrease in refractive index; the net effect of all these factors is an increase in light scattering. This can be seen in Fig. 9.27, where the light scattering coefficient, bscat bsp , for some ambient aerosol particles measured using an integrating nephelometer was found to increase with the liquid water content of the aerosols. The effect of water on aerosol light scattering is sometimes expressed as the hygroscopic growth factor,

ZX I

Zk

I

6 E .TO .m r v=-9

g -Q

5

/ O

4

/YR"

/

O / f /

I

/

Zk

3 2

O1,

m

/

/

'-/

16

O, / 0

O

I

O P o m o n a (10/4-10/5) 9 P a s a d e n a (9/20) Zx P a s a d e n a (9/15) [] G o l d s t o n e (11/1)

O /

9

--

0,1"

, o/

9

/ /

0 :g----t I 0 10

I 20

I 30

I 40

I 50

I 60

I 70

I 80

I 90

I 100

Liquid water concentration (~g m "3)

F I G U R E 9.27 Comparison of light-scattering coefficient with the liquid water concentration in atmospheric aerosols at some locations in California in 1972 (adapted from Ho et al., 1974).

A. PHYSICAL PROPERTIES 150 Y / / 4.// /i/./

03|

E o

04

/ I / / /v./

100

m

RH = 8 0 %

--

RH < 3 0 %

IZ-/I Most Effective Light-scattering

./v./

Range

./vI

E

/ x / /At/ /.A./

a

/

0

./i ./

/. ~ , - / ! / / / I / / / I /

0 0.01

I 0.1

f ~

1

10

100

D geo (lzm)

F I G U R E 9.28 Size distribution of particle geometric cross section (A) as a function of geometric diameter for a typical rural aerosol (adapted from Hegg et al., 1993).

defined as the ratio of bsp at 80% RH to that at a low reference RH (e.g., Waggoner et al., 1983a; Hegg et al., 1993). Changes in light scattering due to changes in RH again reflect the effects on particle size and the index of refraction. Hegg and co-workers (1993) treat the effects on particle size in terms of two components, the first a change on the geometric cross section of the individual particles, and the second a change in the particle size distribution. Figure 9.28 shows the size distribution of the geometric cross section of particles under low RH conditions and at an RH of 80%. The initial distribution represents particles typical of a rural continental site having a number mean diameter of DN 0.076 /xm, a geometric standard deviation of 2, and a number concentration of 2300 particles cm 3. Increasing the relative humidity to 80% results in a larger fraction (41 vs 16%) of the particles being in the range that is most efficient for light scattering, giving a hygroscopic growth factor in this case of 4.2. Values in the range of 1-3 are typical (e.g., see Hegg et al., 1996; Svenningsson et al., 1992; Ten Brink et al., 1996; and Mclnnes et al., 1998). The increase in RH is thus seen in this case to have major effects primarily through this shift in the particle size distribution into a more effective range for light scattering. Hegg and co-workers (1993) point out that this effect is expected to be the largest for particles in rural areas, where the size distributions often peak below the efficient light scattering range and which can therefore take up water and grow into this range. On the other hand, aged urban particles often already exist with sizes in efficient light scattering ranges so that the uptake of water has relatively less effect. For example, Mclnnes et al. (1998) measured hygroscopic growth factors for particles (using a reference RH of 40%) at Sable Island, Nova Scotia, Canada. Particles in air masses from the northeastern

373

United States that had been impacted by anthropogenic emissions had a hygroscopic growth factor of 1.7 0.1, compared to 2.7 0.4 for particles in marine air from the open ocean. In short, the contribution of various components of particles to the scattering of light in the troposphere is complex, depending not only on the particular species of interest but also on its interactions with other constituents, on the initial particle size distribution, and on the relative humidity. Significant advances have also been made in illustrating the effects of visibility reduction by generating simulated photographs calculated to mimic the scattering and absorption of light by particles typically found in the particular area of interest. Both ground-based photography (see, for example, Williams et al., 1980; Maim et al., 1983; Larson and Cass, 1989; Eldering et al., 1993; and Molenar et al., 1994) and satellite images (Eldering et al., 1996) have been mimicked using this approach, both very effective ways to demonstrate the effects of particles on visibility reduction. Significant progress has also been made on developing grid-based models that incorporate both physical and chemical aerosol processes to predict effects on visibility (e.g., see Eldering and Cass, 1996). Not only light scattering but also light absorption by particles can occur. It is generally agreed that the major contributor to light absorption in the visible region is black or graphitic carbon, often referred to as elemental carbon. On a practical level, this is the aerosol component that is insoluble in organic solvents and is not oxidized at temperatures below 400~ (Penner and Novakov, 1996). However, as discussed by Chang et al. (1982), carbon particles in the air are made up of a number of crystallites 2-3 nm in diameter, with each crystallite consisting of several carbon layers having the hexagonal structure of graphite, Fig. 9.29. Because of the presence of defects, dislocations, and discontinuities, there are unpaired electrons that constitute active sites in the carbon; during the formation of the carbon particle in combustion processes, these active sites can react with gases to incorporate other elements such as oxygen, nitrogen, and hydrogen into the structure. Thus elemental, black, or graphitic carbon found in atmospheric particles is not chunks of highly structured pure graphite, but rather is a related, but more complex, three-dimensional array of carbon with small amounts of other elements. As discussed later, the presence of polar groups on the surface is believed to play an important role in determining their properties, such as uptake of water. Being black, this atmospheric carbon is a strong absorber of visible radiation. The specific mass absorption coefficient, bap , h a s been measured to be in the

374


FIGURE 9.29 Structure of elemental carbon.

range from 2 to 20 m 2 g 1 (e.g., see Horvath, 1993; Liousse et al., 1993; Hitzenberger et al., 1996; Petzold et al., 1997; Kuhlbusch et al., 1998; and MoosmiSller et al., 1998). Not surprisingly, the absorption depends on the wavelength of the light. MoosmiSller et al. (1998) report that bap varies with A 2.7, whereas Horvath et al. (1997) report that the absorption coefficient for aerosol in Santiago, Chile, assumed to be due to elemental carbon, varied with A o.92. Because of the nature of its sources, its contribution to light extinction varies geographically and temporally. For example, wood-burning fireplaces and diesels are major sources of elemental carbon, and areas with large numbers of these two sources generally have more graphitic carbon in the atmospheric aerosol, and hence more light absorption. Where wood burning is significant, more particulate graphite carbon would be expected in winter than in summer. Absorption of light by carbon is expected to lead to heating of the atmosphere since the light energy is converted into thermal energy (see Chapter 14.C.lb). This is the opposite effect from scattering of light by

TABLE 9.8

Contribution of Light Absorption to Total Light Extinction Due to Particles in the Troposphere a

Type of atmosphere

Remote Rural

Urban residential

particles back into the upper atmosphere. This heating effect would be expected to be most important in polluted urban areas (e.g., see Liu and Smith 1995; and Horvath, 1995). Table 9.8 summarizes some measurements of the contribution of light absorption by particles, bap , for different types of locations from urban residential to remote. In urban areas, bap varies from ~ 13 to 42% of bext, whereas in very remote areas, the contribution of absorption is much less. This is not surprising since the combustion sources producing graphitic carbon tend to be in urban-industrial areas. However, it is noteworthy that soot has been found associated with many sulfate particles even over the remote oceans (Buseck and Pdsfai, 1999). There are other visible light absorbers than elemental carbon present in particles. These include carbonyl compounds from the oxidation of larger organics in the atmosphere (see Section C.2) and some soil components such as hematite ( F e 2 0 3) (see Chapter 14.C.lb). Their contribution to total light extinction in various regions from urban to remote is not well understood, but it appears that it can be significant. For example, Malm et al. (1996) assessed the contributions of various particle components to scattering and absorption of light at the Lake Mead, Arizona, recreation area and at 18 sites in national parks in the western United States. The contributions of organic carbon, elemental carbon, and soil components to the total absorption of light by particles at Lake Mead were estimated to be approximately equal, at 18, 21, and 19% of the total absorption, respectively (the remainder was due to an unknown absorber or set of absorbers). At the national park sites, the contributions of organic and elemental carbon were again about equal, at 45 and 41% of the total absorption.

hap

Location

(m l)

Mauna Loa, Hawaii Allegheny Mts. Laurel Hill Shenandoah Valley Blue Ridge Mts. Off East Coast of U.S. (Virginia) Denver, Nov Dec. 1978 Detroit, July 1981 Houston, Sept. 1980 Los Angeles Basin, Oct. 1980

6 • 10 8 1.9 • 10 5 1.0 • 10 5 (0.01-1.0) • 10 5 6.6 • 10 5 2.4 • 10 5 3.0 • 10 5 6.4 • 10 5

a From Japar et al. (1986). bAt altitudes from 0.2 to 3 km.

bap (bap bsp) 0.069 0.13 0.051 0.029-0.24 0.42 0.13 0.18 0.25

Reference

Clarke and Charlson, 1985 Japar et al., 1986 Ferman et al., 1981 Novakovet al., 1997b Groblicki et al., 1981 Wolff et al., 1982a,b Dzubay et al., 1982 Pratsinis et al., 1984

3 75

B. REACTIONSINVOLVEDIN PARTICLEFORMATIONAND GROWTH Graphitic carbon also scatters light; in one study, its mass scattering coefficient was found to be approximately the same as that for sulfate (Appel et al., 1985). Because of its contribution to both scattering and absorption of light, graphitic carbon tends to play a proportionally much greater role in light extinction than its contribution to the particulate mass would suggest. For example, Horvath et al. (1997) report that 13% of the total light extinction in Santiago, Chile, was due to absorption by elemental carbon, but an additional 14% was due to scattering; as a result, more than a quarter of the total light extinction by aerosol particles could be attributed to elemental carbon. This disproportionate effect on light extinction of elemental carbon compared to its mass fraction has been reported in a number of areas. For example, in Denver, graphitic carbon was found to represent 15% of the fine particle mass but contribute ~ 3 5 % to the total light extinction (Groblicki et al., 1981); at Zilna Mesa, Arizona, during one sampling period, graphitic carbon comprised 1.4% of the total aerosol mass and 3.5% of the fine particle mass but was responsible for 15% of the light extinction (Ouimette and Flagan, 1982). Similarly, in one Los Angeles area study, graphitic carbon was found to represent ~ 8.5-10% of the fine particle mass but to account for ~ 14-21% of the total light extinction (Pratsinis et al., 1984). The absorption of light by carbon is sufficiently strong that it has been used as a marker of long-range transport in unusual circumstances. Thus, the plumes from the burning of oil wells in Kuwait were detected at the Mauna Loa Observatory in Hawaii in the form of increased hap values. Before the oil well burning, hap was in the (0.1-1) • 10 7 m 1 range, whereas during the burning, values as high as 8.6 • 10 7 m 1 were measured (Bodhaine et al., 1992). Table 9.9 summarizes the results of some recent studies of the contribution of various particle components to light scattering. The contribution of absorp-

TABLE 9.9

tion, when measured, is also shown. Such data suggest that sulfate and organics are major contributors to light scattering, with the contribution of nitrate being more variable. Light absorption, even in these non-urban areas, appears to be significant as well. (It should be noted, however, that apportionment of light scattering to various components of the particles is sensitive to whether it is assumed that the particles are internally or externally mixed, which is not well established in most cases; e.g., see Malm and Kreidenweis, 1997.)

B. REACTIONS INVOLVED PARTICLE FORMATION AND GROWTH

IN

The formation of secondary particulate matter in various size ranges by chemical reactions in the atmosphere may occur by a number of mechanisms. These include (1) reactions of gases to form low-vapor-pressure products (e.g., the oxidation of SO 2 to H2SO 4 or the reaction of aromatics with O H to form multifunctional oxygenated products) followed by nucleation to form new particles or condensation on preexisting particles, along with some coagulation between particles, (2) reaction of gases on the surfaces of existing particles to form condensed-phase products (e.g., the reaction of gaseous H N O 3 with sea salt particles to form NaNO3), and (3) chemical reactions within the aqueous phase in fogs, cloud, or aerosol particles (e.g., SO 2 oxidation to sulfate; see Chapter 8). 1. N u c l e a t i o n ,

Condensation,

and Coagulation

The condensation of a low-vapor-pressure species to form a new particle is known as h o m o g e n e o u s nucleation. Recall that the vapor pressure of a substance over the curved surface of a droplet is greater than over a flat surface of the same substance (e.g., see

Some Reported Contributions (%) of Various Particle Compounds to Light Scattering and Absorption by Particles in the Troposphere

Location

SO42

Meadview, Arizonaa 25 Meadview, Arizonaa 40 Bryce Canyon, Utah 31 San Gorgonio Wilderness, California 21 Fraser Valley, British Columbia, Canadaa 29-39 Canyonlands National Park, Utah 28 Hopi Point, Grand Canyona 39 a Percentage contributions to bsp only.

NO3 3 16 34 16-35 5

Scattering Organics Soil coarseparticles Absorption 32 48 20 25 22-33 19 50

12

28

10 9 1 12

22 21 35

Reference Sisler and Maim, 1997 McMurry et al., 1996 Sisler and Maim, 1997 Sisler and Malm, 1997 Pryor et al., 1997 Eatough et al., 1996 McMurry et al., 1996

376


Adamson, 1973, and Chapter 14.C.2). The smaller the radius of the droplet, the higher is the vapor pressure over the droplet surface. For example, for pure water at 25~ the vapor pressure is only 0.1% greater over a droplet of 1-/xm radius compared to that for a flat surface but is 11% greater if the radius is 0.01/xm. (see Problem 1). This raises the question as to how homogeneous nucleation of even a single species can occur at all, since the first very small droplets formed would tend to evaporate rapidly. The explanation lies in the formation of molecular clusters of molecules that occur as molecules collide in the gas phase. When the system becomes supersaturated, the concentration of the condensable species increases, as does that of the clusters. The clusters grow by the sequential attachment of molecules until they reach a critical diameter (D*) above which the droplets are stable and grow and below which they evaporate (Friedlander, 1977, 1983). The critical diameter is given by D*

4~ k T In s '

(FF)

where y is the surface tension of the chemical forming the particle, ~ is the molecular volume, k is the Boltzmann constant, T is the temperature (K), and s is the saturation ratio, defined as the ratio of the actual vapor pressure to the equilibrium vapor pressure at that temperature. The term binary h o m o g e n e o u s nucleation is used to describe the formation of particles from two different gas-phase compounds such as sulfuric acid and water; such nucleation can occur when their individual concentrations are significantly smaller than the saturation concentrations needed for nucleation of the pure compounds. It is believed that in the atmosphere, formation of particles from low-volatility gases occurs not by condensation of a single species but rather by the formation and growth of molecular clusters involving at least two, and as described shortly, probably three or more different species. A great deal of work on nucleation in the atmosphere has focused on sulfuric acid, since the oxidation of SO 2 is such a common atmospheric process and the acid product has low volatility. Similar considerations apply to methanesulfonic acid, which is formed by the oxidation of some organic sulfur compounds (see Chapter 8.E) and which is also found in particles because of its low vapor pressure (e.g., see Kreidenweis and Seinfeld, 1988a, 1988b). Even for the ostensibly simple case of sulfuric acid and water, the predicted and observed binary homogeneous nucleation rates are not in good agreement. For example, Fig. 9.30 shows both the theoretically predicted (Jaecker-Voirol and Mirabel, 1989) and the experimentally determined (Wyslouzil et al., 1991) concentrations of sulfuric acid needed to have

IO0 =-

80

--

60

40 2O

0.01

I

I

0.1

1.0

10

[H2S04] (ppb)

F I G U R E 9 . 3 0 Theoretically predicted and experimentally measured concentrations of H2SO 4 required for homogeneous nucleation of sulfuric acid at a rate of I particle cm 3 s 1 (adapted from Hoppel et al., 1994; based on theoretical calculations of JaeckerVoirol and Mirabel (1989) and experimental data of Wyslouzil et al. (1991).

homogeneous nucleation into particles at various relative humidities (Hoppel et al., 1994). Clearly, there is a significant discrepancy between the predicted and measured concentrations. As discussed in detail by Pandis et al. (1995), nucleation theory shows there is a critical sulfuric acid concentration above which binary nucleation of sulfuric acid and water should occur. Based on the work of Jaecker-Voirol and Mirabel (1989), they parameterize the theoretical dependence of this critical concentration needed to generate nuclei at a rate of 1 cm 3 s 1 on the relative humidity (RH, expressed in this case as a fraction, i.e., between 0 and 1) and temperature, T (in K): H 2804 crit (/~g m 3) 0.16 exp{0.10T

3.5RH

27.7}. (GG)

Note that this predicts critical concentrations that are more than an order of magnitude below the experimentally observed concentrations of Wyslouzil et al. (1991) shown in Fig. 9.30 (see Problem 3). For a recent treatment of the binary homogeneous nucleation of H2S04 and H 2 0 , see Kulmala et al. (1998). The atmospheric situation is complicated by varying conditions of temperature, relative humidity, and concentrations of other gases such as NH 3 which can enhance nucleation rates over those expected for a well-mixed air mass at a fixed temperature and RH (e.g., see Nilsson and Kulmala, 1998). However, there is a general consensus that the observed rates of nucleation of H2S04 often, indeed usually, exceed those expected from classical binary homogeneous nucleation theory. (Note that this is not always the case. For example, Pirjola et al. (1998) reported that the measured formation of nuclei in the Arctic boundary layer

3 77

B. REACTIONS INVOLVED IN PARTICLE FORMATION AND GROWTH

was consistent with homogeneous nucleation when the air was contaminated with SO 2 from upwind smelters; however, it was much greater than predicted theoretically for clean marine air.) As discussed by Hoppel et al. (1994), it is likely that under atmospheric conditions, nucleation involves condensation onto preexisting clusters of molecules, or "prenucleation embryos," which are so small in size and number that they have not been measured. Ionic clusters, or neutral clusters formed from the recombination of positively and negatively charged clusters, could serve as embryos for nucleation of sulfuric acid and other species. For example, Turco et al. (1998) propose that recombination of positive ion clusters such as H 3 0 (HzO)n with negative ion clusters such as H S O 4 ( H z S O 4 ) m ( H z O ) q (which are known from mass spectrometric measurements to be ubiquitous in the atmosphere; see Chapter 11.A.2) may lead to stable large neutral cluster embryos. Hoppel and co-workers predict, for example, that a cluster of radius 1 nm could act as a nucleus for the condensation of sulfuric acid and water at 60% RH and a n H z S O 4 concentration of only 1 ppt, far below the concentrations shown in Fig. 9.30 for homogeneous nucleation. The development of techniques to measure particles down to 3 nm in diameter (see Chapter 11.B) has provided some new insights into nucleation in the atmosphere. For example, McMurry, Eisele, and coworkers have measured simultaneously ultrafine particles and gaseous H z S O 4 in a number of locations (e.g., Weber et al., 1995, 1996, 1997, 1998; Eisele and McMurry, 1997). Formation of ultrafine particles occurs at much smaller concentrations of gaseous H z S O 4 than expected based on classical binary nucleation theory for sulfuric acid and water. In addition, the dependence on the concentration of gaseous H 2SO 4 is much smaller than expected theoretically. They propose that ammonia may assist in the nucleation process. Figure 9.31 shows the H2S04 vapor pressure above the pure liquid and above a 1:1 mixture of NH3 and H2S04 (Marti et al., 1997a; Eisele and McMurry, 1997). The vapor pressure of H z S O 4 is reduced by two to three orders of magnitude when NH 3 is present at equimolar levels. This large reduction in vapor pressure suggests that the reaction of NH 3, which is ubiquitous in the troposphere (see Chapter 11.A.4a), with H 2S04 on a molecular level may play a key role in the high observed rates of nucleation. This is consistent with the observation of high concentrations of the smallest measurable ultrafine particles downwind of a large penguin colony, which would be expected to be a significant source of NH 3 (Weber et al., 1998). Interestingly, the growth rate of ultrafine particles is often an order of magnitude

10 -5 10 -6 ~, n

10 7

0~r./) -1-

1 0.8

10 -9 10-10 I

I

I

5

i

i

i

i

I

10

i

i

i

i

I

15

Relative Humidity (%)

F I G U R E 9.31 Vapor pressure of H2SO4 as a function of relative humidity for pure H2SO 4 (O) and a 1:1 mixture of NH 3 and H2SO 4 ( ) (adapted from Marti et al., 1997a; and Eisele and McMurry, 1997).

larger than can be explained based solely on the uptake of gaseous H2SO 4 by the particles, suggesting that other species such as organics contribute to the growth of ultrafine particles once nucleation to form new particles has occurred (e.g., see Marti et al., 1997b; and Weber et al., 1997). This is consistent with observations that organics can contribute up to 80% of the concentrations of condensation nuclei (CN) under some conditions (e.g., Rivera-Carpio et al., 1996). The uptake of inorganics such as HNO 3 and HC1 combined with NH 3 has also been proposed to contribute to particle growth in some cases (e.g., Kerminen et al., 1997). Similar observations have been made in the marine boundary layer where biogenically emitted sulfur compounds such as dimethyl sulfide (DMS) may serve as the source of gaseous H2SO 4 and other low-volatility species, e.g., methanesulfonic acid, MSA (see Chapter 8.E). For example, Clarke et al. (1998) report measurements of ultrafine particle formation, DMS, SO2, and gaseous H2SO 4 in the tropical marine boundary layer that are consistent with the oxidation of DMS to SO 2 and then H2S04, followed by nucleation. As in many other studies, the nucleation rate of H2SO 4 was much larger than expected from classical nucleation theory. Interestingly, Weber et al. (1995) conclude from measurements of H2SO4, MSA, and ultrafine particles at the Mauna Loa Observatory that H 2SO 4 was the major precursor to ultrafine particles and that the contribution of MSA was small. For a review of nucleation in the atmosphere, the reader is referred to Nucleation and Atmospheric Aerosols (Fukuta and Wagner, 1992; Kulmala and Wagner, 1996) and Microphysics of Clouds and Precipitation (Pruppacher and Klett, 1997).

378


Heterogeneous condensation is secondary aerosol formation by the scavenging of the low-vapor-pressure products onto preexisting particles. If the concentration of particles is sufficiently high, this dominates over the formation of new nuclei via homogeneous nucleation (e.g., Friedlander, 1978, 1980). The condensation of low-volatility vapors on preexisting particles depends on a number of factors, including the rate of collisions of the gas with the surface, the probability of uptake per collision with the surface, i.e., the mass accommodation coefficient (see Chapter 5.E.1), the size of existing particles, and the difference in partial pressure of the condensing species between the air mass and the particle surface. While some of these parameters are reasonably well known, others are not. For example, mass accommodation coefficients for the complex surfaces found in the atmosphere are not well known. Indeed, the exact nature of the surfaces themselves, which determines the uptake and the partial pressures of gases at the surface, remains a research challenge. However, despite the uncertainties in quantifying homogeneous new particle formation compared to condensation on preexisting particles, there are a variety of field data supporting the occurrence of both in the troposphere. Figure 9.32, for example, shows some particle measurements made from a ship off the coast of Washington State by Covert and co-workers (1992). Figure 9.32a shows the total particle surface area for particles with diameters >0.02 /xm; this surface is needed for condensation of low-volatility vapors onto preexisting particles. It is seen that at about 1430 hours, the surface area decreased significantly, from about 22 to less than 5/~m 2 cm 3 of air. As the surface area fell to low values, the relative numbers of ultrafine particles, defined in this case as those with diameters 0.015 /xm, increased (Fig. 9.32b). This is consistent with the formation of new particles by homogeneous condensation, since it occurred in the smallest size range and when the available surface area for condensation on existing particles was small. The subsequent growth in the number of slightly larger particles, with diameters in the 0.02- to 0.024-/xm range (Fig. 9.32c) may then be due to heterogeneous condensation on the ultrafine particles. This causes an increase in these larger particles accompanied by a decrease in the ultrafine range, as seen in Fig. 9.32b. Smog chamber studies have documented similar aerosol growth mechanisms. For example, in the photochemical oxidation of dimethyl sulfide, the formation and growth of particles in an initially particle-free system was observed. However, if seed particles with 34-nm mean size were present, an oscillation in the

40

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30

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20

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10

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0 ~ o

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3

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0

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,

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,

I

,

I

300

100

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-9 o

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04

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F I G U R E 9 . 3 2 Some evidence from field studies along the coast of Washington State on April 22, 1991, both for new particle formation and for condensation on preexisting nuclei: (a) particle surface area for particles with D > 0.02 /xm; (b) ratio of number of particles with D 0.015 /xm to those with D > 0.015 /xm; (c) number concentration with diameters 0.02-0.024 /xm (adapted from Covert et al., 1992).

number of fine particles was observed, with bursts of nucleation occurring periodically (Flagan et al., 1991). This was attributed to the heterogeneous condensation on the seed particles initially, but at too slow a rate to remove the low-volatility products. As the latter accumulate, homogeneous condensation occurred, forming a burst of new particles. Condensation on these new particle surfaces then occurs until they coagulate with the seed particles and the process begins anew. For detailed discussions of the quantitative treatment of such condensation processes in the atmosphere, the reader is referred to articles by Pandis et al. (1995) and Kerminen and Wexler (1995). Coagulation refers to the formation of a single particle via collision and sticking of two smaller particles. Small particles undergo relatively rapid Brownian motion, which leads to sufficient particle-particle collisions to cause such coagulation. The coagulation of smaller particles with much larger ones is similar to condensation of a gas on the larger particle and acts primarily to reduce the number of small particles, adding relatively little to the mass or size of the larger particles. Hence the larger mode will not show significant growth by such a mechanism. As expected, the rate of such processes depends on the diameter of the

379

B. REACTIONS INVOLVED IN PARTICLE F O R M A T I O N A N D G R O W T H

large particle, how rapidly the smaller particle is carried to it (i.e., the diffusion of the smaller particle), and the concentrations of the particles. So-called self-coagulation, where the particles are approximately the same size, can, however, lead to changes in the size distribution of the aerosol particles. As one might expect, the rate of this process is a strong function of the particle concentration as well as the particle size. Table 9.10 shows an estimate (Pandis et al., 1995) of the time scales for coagulation of smaller particles onto larger ones characteristic of various types of air masses. For comparison, typical time scales for condensation, dry deposition of the particles, and transport are also shown. (For discussions of dry deposition of particles, see, for example, Slinn (1982, 1993), Arimoto et al. (1987), and Main and Friedlander (1990)). As expected, condensation is fast, but coagulation is also significant on these time scales in some situations. 2. R e a c t i o n s of G a s e s at P a r t i c l e S u r f a c e s There are some well-known examples of reactions of gases with solids at the interface that are potentially important in the atmosphere. For example, the reactions of 0 3 with adsorbed solid polycyclic aromatic hydrocarbons (PAH) are discussed in Chapter 10. Another example is the reaction of NaC1 and NaBr in sea salt particles with gaseous oxides of nitrogen such as H N O 3 (see Chapter 6.J): NaCl(s )

HNO3(g ) ~ HCI(g)

NaNO3(s).

(1)

The replacement of chloride in sea salt particles by nitrate as well as by sulfate has been observed in many measurements in coastal areas using bulk filter samples (e.g., see Section C.1 and Chapter 6.J.2b) and, more recently, in single particles (e.g., Murphy et al., 1997; Gard et al., 1998). Even in these relatively simple cases,

TABLE 9.10

however, the nature of the reactions at the interface and how they contribute to particle growth and or transformations is not clear. For example, in the reaction of NaC1 with H N O 3 and other oxides of nitrogen such as NO 2, the initial reaction forms one or more unique surface nitrate species which do not alter the particle morphology significantly (Vogt and FinlaysonPitts, 1994; Vogt et al., 1996; Allen et al., 1996). However, if exposed to water vapor even well below the deliquescence point and then dried, this surface nitrate reorganizes into small microcrystallites of NaNO 3 attached to the original salt surface (see Fig. 11.63 in Chapter 11). This generates very small particles of NaNO 3 and may be responsible for the observation of small particles in the marine boundary layer which are almost completely devoid of chloride (e.g., Mouri and Okada, 1993). A further complication is the recent indication that even small amounts of strongly bound surface-adsorbed water may play a critical, indeed determining, role in the interaction of gases with surfaces traditionally thought to be solids. For example, in the N a C I - H N O 3 reaction, there is evidence that the reaction even in laboratory vacuum systems occurs on sites holding adsorbed water. As a result, the surface does not become saturated as one would expect for a solid surface, since the underlying reactant salt continues to dissolve in the surface water (Beichert and Finlayson-Pitts, 1996). Another example of reactions at interfaces that is only now being recognized, due to the lack of suitable experimental techniques in the past, is that of species such as SO 2 and NO 2 at liquid interfaces. As discussed in Chapters 7 and 8, there is increasing evidence that the reactions of such species at the air-water interface can be fast relative to that in the bulk and may have unique reaction mechanisms compared to those in the bulk or gas phases. Given the paucity of data on such processes at the present time, they are generally not included in present models of aerosol growth. How-

Typical Time Scales for Various Aerosol Fates a Type of air mass

Fate

Condensation Coagulation of 0.03-/xm-particles with larger particles Deposition 0.03-/xm particles 0.3-/xm particles Transport a From Pandis et al. (1995).

Urban

Remote marine

Free troposphere

Nonurban continental

0.01-1 h 0.1-2 days

1-10 h 10-30 days

2-20 h -~50 days

0.5-20 h 1-5 days

0.5-10 days -~1 month 2-5 days

0.5-10 days -~1 month 1-2 weeks

3 days to 2 weeks

-~1 month -~1 month 1-2 weeks

380


ever, as our understanding of such reactions at interfaces expands, their implications for the growth of aerosols in the troposphere will need to be critically assessed.

3. Reactions in the Aqueous Phase We have seen in Chapter 8 that reactions in the aqueous phase present in the atmosphere in the form of clouds and fogs play a central role in the formation of sulfuric acid. Thus, an additional mechanism of particle formation and growth involves the oxidation of SO 2 (and other species as well) in such airborne aqueous media, followed by evaporation of the water to leave a suspended particle. As discussed earlier, this process is believed to be responsible for the presence of more than one peak in the accumulation mode. For example, Hoppel and coworkers (1985, 1986, 1994) measured particle size distributions in and around clouds. Figure 9.33 shows a typical size distribution for in-cloud particles (the socalled interstitial particles) and for particles in the air mass just below the cloud (Hoppel et al., 1994). The larger mode peaking at ~0.15 /xm below the cloud is attributed to processing of the smaller particle mode by clouds in two ways (Hoppel et al., 1986). First, the smaller particles are taken up by the cloud droplets. If the cloud evaporates before precipitating (which is most common), the agglomeration of smaller particles taken up into the droplet becomes a particle of larger size. Second, SO 2 can be absorbed into the droplet and oxidized to sulfate (see Chapter 8). When the cloud evaporates, it leaves a particle containing the newly formed sulfate. This has been proposed to explain the bimodal distribution of particles such as that in Fig. 9.33. In-cloud oxidation of SO 2 to sulfate has also been 104

E O i-

102

O

"O

Different mechanisms of aerosol growth give rise to different so-called growth laws, which are expressions relating the change in particle size (e.g., volume or diameter) with time to the particle diameter. Because different mechanisms of particle formation give rise to different growth laws, one can test experimental data to see which mechanism or combination of mechanisms is consistent with the observations. For a more detailed discussion of this approach, see Friedlander (1977), Heisler and Friedlander (1977), McMurry and Wilson (1982), Pandis et al. (1995), and Kerminen and Wexler (1995).

C. CHEMICAL COMPOSITION OF TROPOSPHERIC AEROSOLS

i-

z

4. Relative Importance of Various Aerosol Growth Mechanisms

103

_13

E -, E

invoked to explain particles at ~0.1 /xm in marine aerosols (e.g., Fitzgerald, 1991). Similarly, Kerminen and Wexler (1995) have examined potential sources of the bimodal particle distribution often observed in the accumulation mode and suggested that it results from aqueous-phase oxidation of SO 2 to sulfate. They assume that the larger particles in such bimodal distributions are formed by the growth of particles in the smaller mode. In this case, to be consistent with atmospheric observations, the mechanism responsible for the two modes must occur on a time scale of hours (a day or so maximum) and must involve growth of the smaller particles preferentially over the larger ones. In addition, because the smaller mode does not get completely converted into the larger mode, only some of the smaller particles must grow or there must be a continuous source of these particles to replace those that grow into the larger mode. They show that only aqueous-phase oxidation of SO 2 in clouds or fogs is consistent with these requirements; growth by condensation of gases on existing particles or by coagulation of particles is too slow to be compatible with the atmospheric observations.

101

-

/V

i i

"

,,

11

i

i

5,

-- - in cloud 100 , i i i l llll 0.001 0.01

i l i

\/below

",,, c'~ ",~

i

i i i l llll

0.1

i

I't

I l lll

1.0

Radius (gm)

F I G U R E 9.33 Size distribution of particles in clouds (solid line) and below the clouds (dashed line), showing two modes (adapted from Hoppel et al., 1994).

As we have seen in our earlier discussion of the size distribution of tropospheric particles, the chemical components are not generally distributed equally among all sizes but, rather, tend to be found in specific size ranges characteristic of their source. Generally, the smallest ultrafine particles are produced by homogeneous nucleation and hence tend to contain secondary species such as sulfate and likely organics (see Section A.2). Particles in the Aitken nuclei range are produced

381


by combustion processes, by coagulation of smaller particles, and by condensation of low-vapor-pressure products of gas-phase reactions. Hence these particles and accumulation mode particles tend to contain elements such as carbon and some trace metals such as V, which are characteristic of combustion, as well sulfates, nitrates, and polar organics. Finally, because mechanical processes are primarily responsible for coarse particles larger than about 2.5 /xm, these larger particles typically contain elements in soil, sea salt, etc. In the following, we illustrate these principles using data from studies in various locations and characteristic of various types of environments from polluted urban to the free troposphere. 1. I n o r g a n i c Species a. Size Distribution

Table 9.11 shows the aerodynamic mass median diameter (MMD) for some typical inorganics that are common components of tropospheric particles. Also shown are the calculated crystal enrichment factors, EFcrus t. These are a measure of the enrichment of the element in the airborne particles compared to that expected for the earth's crust, using aluminum as the reference element. Thus EFcrus t for a particular element X is defined as EFcrus t

{Xai r A l a i r}

{Xcrus t Alcrust} ,

(HH)

where "air" indicates the concentration in airborne particles and "crust" that in the earth's crust. A value of EFcrus t of approximately unity indicates that the composition of the particles is consistent with that expected if they were formed by erosion of the earth's surface. In this case, one would also expect that they will fall in the coarse particle range, having typical diameters above about 2.5/xm. The data in Table 9.11 are based on an extensive review of the literature through 1985 by Milford and Davidson (1985). Because they represent composites of many different studies carried out in many different locations by many different investigators, they will not match any particular sample of airborne particulate matter; indeed, for many of these elements, the size distribution is multimodal, which is not reflected in the median values shown in Table 9.11. On the other hand, such data demonstrate very clearly some characteristics of tropospheric particles that are common to many conditions. For example, the most common elements in the earth's crust (Table 9.12 and Fig. 9.34) are O, Si, A1, Fe, Mg, Ca, Na, K, and Ti. These elements have MMDs > 3/xm and enrichment factors that are generally less than three (Table 9.11). That the enrichment

T A B L E 9.11 A e r o d y n a m i c Mass Median D i a m e t e r s of T r o p o s p h e r i c Particles C o n t a i n i n g V a r i o u s Elements O b s e r v e d in a N u m b e r of Studies and E n r i c h m e n t Factors a Element W Pb Hg Se Cd Sb Br Ni I As Cr Zn Cu V U In Ta Cs Mn Eu Co Th Sm C1 Ba Fe K Na Si Sc A1 Ca Ce Ga Mg Ti Hf Sr

MMD (Ixm)

EFcrust b

0.43 0.55 0.61 0.68 0.84 0.86 0.89 0.98 1.03 1.11 1.11 1.13 1.29 1.44 1.56 1.75 1.77 1.89 2.06 2.59 2.63 2.73 2.76 3.04 3.31 3.42 3.76 3.78 3.90 4.39 4.54 4.64 5.10 6.00 6.34 6.52 7.65 11.9

19 1500 560 3100 1900 1400 1900 32 510 310 8.1 260 100 14 2.9 90 1.1 12 3.9 2.7 3.5 1.8 2.1 740 5.5 2.1 2.0 4.4 0.79 1.2 1.0 2.8 2.6 2.5 2.4 1.4 2.0 1.5

a From Milford and Davidson (1985). b Calculated crustal enrichment factors.

factors are not all unity suggests that there is some contribution of other sources to the airborne particles. For example, road dust is thought to be a major source of Ca since crushed limestone is often used in road paving materials. On the other hand, in the literature reviewed for Table 9.11, particulate lead has an MMD of only 0.55 /xm. This is consistent with the major source of lead to the atmosphere in the 1970s being combustion of leaded gasoline, the remainder being due mainly to smelting processes. As a result, Pb is primarily found in the accumulation mode and has a large enrichment factor

382


T A B L E 9.12 A b u n d a n c e in E a r t h ' s C r u s t of Some of the Most C o m m o n Crustal Elements and T h o s e C o m m o n l y F o u n d in Atmospheric Aerosols or U s e d as a T r a c e r a

Weight percent

Atom percent

46.60 27.72 8.13 5.00 2.09 3.63 2.83 2.59 0.44 0.105 0.0950 0.0260 0.0200 0.0135 0.0030 0.0020 0.0013 0.0006

62.55 21.22 6.47 1.92 1.84 1.94 2.64 1.42

O Si A1 Fe Mg Ca Na K Ti P Mn S C V La N Pb Sm

of these two processes will depend on the particular location and conditions. A word of caution is also in order with respect to assigning a particular particle to the fine or coarse particle modes. Since the size distributions can generally be described as log-normal, they do not have sharp cutoffs. A few particles at the top end of the fine mode distribution will have diameters larger than 2.5/xm and a few at the bottom end of the coarse mode will have diameters smaller than this. For example, as Lodge (1985) points out, for a coarse particle distribution with a geometric mean diameter of 15/xm and a geometric standard deviation of 3, about 5% of the particles will have diameters below the 2.5-/xm fine particle cutoff. This may be responsible for observations that while Si and Ca dominate the coarse particle mode, they are also often found at significant levels in fine particles (e.g., see Katrinak et al., 1995). In addition to crustal elements being found in airborne particles due to weathering processes, in marine areas one also finds particles characteristic of sea salt. Wave action entrains air and forms bubbles that rise to the surface. As they rise, dissolved organics may become adsorbed on them. The bubbles burst on reaching the surface, producing small droplets that are ejected into the air. Two types of drops have been distinguished--jet drops and film drops. Jet drops are produced from the jet of water that rises from the bottom of the collapsing bubble; film drops are produced from the bursting of the bubble water film. Particles with a wide range of sizes, from less than 0.1 /xm to greater than 100 /xm, are formed (Blanchard and Woodcock, 1980; Blanchard, 1985). It might be assumed that initially the composition of the liquid drops would approximate that of seawater, given in Table 9.13. As evaporation of water in the droplet into the surrounding air mass occurs, the salt

a From Mason and Moore (1982).

of 1500. (Similar considerations are expected to apply to manganese in some polluted urban areas where (methylcyclopentadienyl)manganese tricarbonyl, MMT, is added to gasoline to improve octane and act as an antiknock agent (e.g., see Loranger and Zayed, 1997; and Wallace and Slonecker, 1997). Intermediate between these two extremes is vanadium, with an MMD of 1.44/xm and EFcrus t 14. Vanadium is produced by the combustion of fuel oil, which is expected to generate small particles, but is also found in the earth's crust (Table 9.12), which should lead to its being found in large particles as well. Clearly, the relative contribution

102 o~

.~

1 10 .2

13r) .m 0-4 9

O

E
2 Ixm) 20.9 32.7 55.9 32.2 65.2 50.5 103 18.1 2.90 0.61 0.67 0.64 22.8 0.23 0.43 1.42 1.28 0.43 0.81 0.40 0.46 7.55

12.5 13.8 29.8 12.9 23.4 43.4 36 11.3 1.37 0.35 0.28 0.31 12.5 0.11 0.21 0.86 0.34 0.22 0.52 0.09 0.25 1.60

111

135

a From Artaxo et al. (1988).

nitrate are again seen to be much larger in the coarse particles. These data also illustrate again that soil components such as Si, A1, Mg, Fe, and Ca tend to be found mainly in the coarse particle mode, whereas sulfur (i.e., sulfate) is primarily in fine particles. Artaxo and co-workers have deconvoluted these elemental distributions using principal factor analysis into contributions from different sources, including soil, dust, and two other sources that appeared to be plant related. On the other hand, nitrate is often a major component of the fine particle mode as well, especially in more polluted urban areas. Figure 9.37, for example, shows the frequency of observations of sulfate and nitrate in various particle sizes during a study in southern California in the summer of 1987 (John et al., 1990). Three distinct peaks are seen for both sulfate and nitrate, at 0.2 0.1, 0.7 0.2, and 4.4 1.2 /xm, respectively. Table 9.15 shows the concentrations of nitrate, sulfate, and ammonium in each size range, from which the predominance of nitrate, as well as sulfate and ammonium, in the accumulation mode is evident.

386


35

so42-

30 25 o tO o" O L.. LL

20 15 10 5 0 0.1

1

10

3530 >, o t'-

NO 3-

2520-

00, 15-

U_

0.1

1

10

D a (gm)

F I G U R E 9 . 3 7 Frequency of observation of SO42 and NO 3 as a function of particle size during summer, 1987, in southern California (adapted from John et al., 1990).

b. Trace Elements as Tracers

In a number of studies, the elemental composition of aerosol particles has been related to their source. Schroeder and co-workers (1987) review the sources, sinks, analysis, deposition, chemical forms, and global cycles of various trace elements. Because the relative amounts of trace elements vary for coal compared to oil-fueled power plants, for example, it has been suggested that certain elements or ratios of elements may serve as tracers for various sources (e.g., see Gordon, 1988; and Rahn and Lowenthal, 1984). Thus V and Ni are indicative of oil combustion, and elevated concentrations of elements such as As and Se are usually

TABLE 9.15 Mean Concentrations of Ammonium, Nitrate, and Sulfate by Size in Particles in Southern California in 1987 a Peak size (p~m)

0.2 0.7 4.4

NH 4 75 352 64

84 283 50

a From John et al. (1990).

NO 3 50 219 128

90 238 75

SO42 24 129 29

18 66 14

associated with coal burning and smelter operations (e.g., see Mosher and Duce, 1987; and Rabano et al., 1989). The ratio of Mn to V and of Fe to Mg have been suggested as indicators of coal burning (Rahn, 1981; Lowenthal and Rahn, 1987; Parekh and Husain, 1987). The relative concentrations of rare earth elements in fine particles have also been suggested as being indicative of emissions from oil-fired power plants and from refineries. For example, the ratio of lanthanum to samarium is much higher in emissions from these sources than found in particles formed from the earth's crust (Olmez and Gordon, 1985). Thus the La Sm ratio was measured to be about 20 for refinery emissions and 28 for emissions from an oil-fired power plant compared to 5 from a coal-fired power plant and 9 from the earth's crust (Olmez and Gordon, 1985; Kitto et al., 1992). Rare earth isotopes have also been suggested as being useful for tracer studies. That is, they can be intentionally added to a source, such as flue gas or the fuel used in combustion, and their unique isotopic signature used to follow the plume as it disperses downwind. This approach, using the isotopes of Nd, for example, is discussed in more detail by Ondov and co-workers (Ondov and Kelly, 1991; Wight et al., 1991; Ondov et al., 1992; L i n e t al., 1992). While there has been a great deal of interest in using the elemental composition of particles as tracers of sulfate and nitrate from sources located large distances upwind, caution must be exercised in using elements or ratios of their concentrations as indicators of such long-range transport. For example, element concentrations and ratios may change from that at the source as the air mass travels over downwind sources. For a discussion of this and some of the other potential problems, see Thurston and Laird (1985), Rahn and Lowenthal (1985), Lowenthal et al. (1988), Lowenthal and Rahn (1988), and Gordon (1988). c. Source Apportionment Models

Because certain sources emit particles with characteristic elemental signatures, in principle one ought to be able to measure the composition of particles in the atmosphere and then "work backward" to calculate how much each source contributed to obtain the final, observed particle composition. This approach involves the use of receptor models, defined as "models that assess contributions from various sources based on observations at sampling sites (the receptors)" (Gordon, 1988). There are a number of different types of source apportionment models, including the chemical mass balance method, factor analysis, multiple linear regression analysis, and Lagrangian modeling. The chemical

387


mass balance method is based on the idea that the observed concentration of the ith element in the particles (OCi) is the sum of contributions from each source:

OCi

~cijs j.

(II)

J

In Eq. (II), C ij is the concentration of species i in the jth source and is the source strength for source j (e.g., as mass m sj3) contributed by particles from that particular source. This assumes there are no loss processes for the element between source and receptor, e.g., reactions and or volatilization. If the concentration profiles for the various contributing sources (i.e., the cij) are measured, as well as the total concentration OCi at the receptor, the set of simultaneous equations represented by (II) can be solved to obtain sj, the contribution of each source for that particular element. Figure 9.38, for example, shows the application of the chemical mass balance approach to the fine particle fraction of particles collected at a location in Philadelphia (Dzubay et al., 1988; Olmez et al., 1988; Gordon, 1988). If the set of equations (II) fitted the data perfectly, the sum of the contributions of the various sources would be 100% for each element. Clearly, from the top frame, this is not the case for a number of elements, and both positive and negative deviations from 100% can be seen. However, the contributions of several sources are clear: Si and Fe from soil, Ni, V, and Ca from oil-fired power plants, Ti from a paint pigment plant, La, Ce, and Sm from a catalytic cracker, K, Zn, and Sn from an incinerator, Sb from an antimony roaster, and Pb and Br from motor vehicles. This approach has been used to show, for example, that non-catalyst-equipped vehicles are the major source of fine particles in the area around Silo Paul, Brazil (Alonso et al., 1997), and in Mexico City (Vega et al., 1997). A second approach that is used extensively is factor analysis. This is based on examining the variations of the concentrations of various elements from their mean value. The data are then analyzed to find groups of elements whose changes as a group account for these variations. These groups of elements, which change in such a way as to indicate they are associated with each other, are known as "factors." In the factor analysis approach, the minimum number of factors or elemental groupings needed to explain the observed variations in the composition is sought. Figure 9.39, for example, shows the results of factor analysis applied to particles collected in the Amazon Basin (Artaxo and Hansson, 1995). Only two factors are needed to explain this data

set, one a mineral dust component containing such elements as Si, A1, Ti, and Fe that peaks at a size of 2.0 /xm, and the second a biological component containing P, S, K, Ca, Zn, and Sr. The biogenic-associated K, P, S, and Zn are distributed bimodally, as seen in Fig. 9.39. The biogenic component that has enhanced S, Sr, and Zn peaks at 0.25 /xm while that with enhanced P and K is found in the 2- to 4-/xm range. This approach has also been applied to the organic component of particles. For example, Veltkamp et al. (1996) measured 18 organics in particles at Niwot Ridge, Colorado, as well as a variety of organics and inorganics in the gas phase. They identified seven factors, which include gas-phase internal combustion products (e.g., benzene), particle-phase oxygenated biogenic hydrocarbons (e.g., camphor), particle-phase anthropogenic products (e.g., n-heptadecanal), and particlephase biogenic aldehydes (e.g., mid-molecular-weight n-aldehydes). A third common approach is multiple linear regression analysis. In this approach, multiple linear regressions of the form a (tracer of source 1) b (tracer of source 2)

z

(JJ)

are carried out, where Y/ is the measured species concentration and the tracers for various sources are measured simultaneously. The regression coefficients a, b , . . . are then determined from the data. For example, Zweidinger et al. (1990)used multiple regression analysis to examine the use of volatile organic hydrocarbons (VOC) as mobile source tracers. They showed that there was a high correlation between certain VOCs and Pb, used as a tracer of mobile source emissions, and proposed that as Pb levels drop due to the increased use of nonleaded gasoline, these could be used to calculate the contribution of mobile sources to organics in fine particles. Swietlicki et al. (1996) also used this approach to elucidate the source of air pollution in Lund, Sweden. Application of a combination of all of these approaches can be particularly powerful when the requisite input data are available. For example, Malm and Gebhart (1997)used a combination of these source apportionment techniques to show that in the Grand Canyon area, about 50% of the sulfur in particles was associated with sources that also emitted Se and hence was attributed to coal-fired power plants in the region. A strong association with bromine was also found, suggesting that wood smoke was also a significant source of sulfur in this region. Finally, Lagrangian approaches to source apportionment have been used in some airsheds. For example,

3 88


g

150

o

100

~,

so

"(3

r

8o~ 0

60 40 20 0 120 100 80

Ni

mm

60 20 ,o o

.m

-

100 80 60 40 20 0 100

Ti

.

Incinerator

60 o

m

m--__

m

. m_

.

Cemm m Sm

K

_

m

_

| _

i

Sn

Zn

m

8o ,o

100 80 60 40 20 0 120 100 80 60 40 20 0

m

m

80 ~ 60 40 20 0 120 100

20

Oil-fired power plants

Ca

m

II

II

m

m Sb

Antimony roaster

m Br/2

Motor vehicles

I AIIsil

S I KlCalTil

m.

m

Pb

/

V IMnlFe I N i l z n l A s l s e l B r / 2 1 M o l c d l s n l s b l k a l c e l s m l p b l

F I G U R E 9.38 Example of chemical mass balance approach applied to fine particle fraction in Philadelphia in 1982 (adapted from Gordon, 1988).

Cass and co-workers (e.g., Rogge et al., 1996; Gray and Cass, 1998; Kleeman and Cass, 1998) have used this method to assess the contribution of various sources of particles in the Los Angeles area. If the particle emissions from each source are known and assumed to remain in air as external mixtures, they can be tracked separately in the model as they are transported and react. Hence their contribution to the concentrations at a particular location can be calculated.

For further details of the application of receptor models, see Hopke (1985) and Gordon (1988).

d. Forms of Inorganics in Tropospheric Particles The discussion thus far has focused on the elemental composition of particles, which can be determined using such techniques as neutron activation or X-ray fluorescence analysis (see Chapter 11). One of the questions of increasing interest is the chemical form in

C. CHEMICALCOMPOSITIONOF TROPOSPHERICAEROSOLS 1.4 ID

1.2

o r if}

1

o

0.8

m O

-

0.08 0.06 0.04 0.02 0.00

0

50

1O0

150

200

250

300

350

400

AMo (gg m -a)

F I G U R E 9.53 Yield (Y) of secondary organic aerosol as a function of the amount of aerosol generated, zXM0, during the V O C - N O X oxidations in air of some aromatic hydrocarbons (adapted from Odum 1997b).

et al.,

(KK)

The yield of secondary organic aerosol depends on the organic particle mass concentration because of the gas-particle partioning of the semivolatile organic products (see later). Thus, Odum et al. (1996) showed that the yield of secondary organic aerosol, Y, is given by Y

0.14

(LL)

In Eq. (LL), M 0 is the concentration of the condensed-phase organic (in /xg m 3) available to absorb semivolatile organic products, c~i is a constant that relates the concentration of the ith secondary organic aerosol component formed, Ci, to the amount of parent precursor organic reacted i.e., C i (ng m 3) 1000c~izX(parent organic in /xg m 3), and Kom,i is the gas-particle partioning coefficient for the ith component. As discussed in more detail in Section D, Kom,i is in effect an equilibrium constant between the condensed- and gas-phase concentrations. Thus, if a particle secondary oxidation product does not get partitioned efficiently into the condensed phase (i.e., Kom,i is small) or the available organic condensed phase for uptake of the semivolatile product is small, Eq. (LL) reduces to Y Mo~oliKom,i and the secondary organic aerosol yield is proportional to the amount of condensed phase available for uptake of the low-volatily gaseous products. On the other hand, if Kom,i and M 0 are large, Eq. (LL) becomes Y E;c~i, independent of the amount of condensed phase available for product uptake. Figure 9.53, for example, shows a plot of the yield of secondary organic aerosol from the V O C - N O X oxidation in air of some aromatic compounds as a function

of zXM0 (Odum et al., 1997b). The rapid initial increase in the yield is due to the increasing availability of condensed phase for uptake of the semivolatile products. However, as the amount of available condensed phase continues to increase, this becomes less limiting, and the yields approach the final, limiting yield determined by c~i. In short, the same types of aerosol organic products have been identified both in model systems and in polluted urban ambient air and can generally be rationalized based on the oxidation of known constituents of air. The measured yields of organics in the particles can depend on the amount of particle phase available for uptake of the organic if it is semivolatile and partitions between the gas and condensed phases. This partitioning, and its dependence on the amount of condensed phase available, may be at least in part responsible for discrepancies in the yields of secondary organic aerosol reported in a number of studies. Grosjean (1989, 1992) has used estimates of the emissions of gaseous secondary organic aerosol precursors to estimate the amounts and types of organic compounds expected in particles in the Los Angeles area under typical moderately polluted conditions. These calculations suggest that a variety of organics in particles should be formed by reactions of the precursors and that carbonyl compounds, organic acids, nitrated derivatives, and phenols should all be formed. Gaseous aromatics were predicted to predominate in the formation of organic aerosols, which has been confirmed in smog chamber experiments (Odum et al., 1997a, 1997b). The complexity of organics found in particles as well as the possibility of multiple sources, both primary and


secondary, often makes identification of sources difficult. There are some organics that are "markers" for particular sources, e.g., diterpenoid acids and retene as a marker of wood smoke (Mazurek et al., 1991; Rogge et al., 1993d; Standley and Simoneit, 1994; Radzi bin Abas et al., 1995). Another approach is to use receptor models similar to those applied to inorganics (see Section C.l.c). For example, Greaves et al. (1987) applied factor analysis to particles collected at Boulder, Colorado, an urban site. They identified more than three dozen organic components and four factors: factor 1 was a photochemical factor that included oxygenated organics such as carboxylic acids, aldehydes, furans, lactones, ketones, and phthalates; factors 2 and 3 were different biogenic sources and included such compounds as terpenoids, bornyl acetate, camphor, phytone (6,10,14-trimethylpentadecan-2-one), and high molecular weight n-alkanes; and factor 4 was a vehicular source that included some saturated hydrocarbons and an unidentified branched-chain carboxylic acid. Similarly, Rogge et al. (1993d) and Schauer et al. (1996) analyzed particles collected in southern California and identified and measured more than 80 organic compounds. Some (e.g., the C 2 9 - C 3 3 n-alkanes) could be identified with biogenic sources, some with primary anthropogenic emissions (e.g., oleic acid from meat cooking), and some with secondary products from the oxidation of precursors (e.g., some aliphatic dicarboxylic acids). Figure 9.54, for example, shows the composition of particles as a percentage of the total mass at (a) the west end of the Los Angeles air basin and (b) downwind at Rubidoux. Evidence of secondary aerosol generation is found in the increased total particle mass at the downwind location, 42.1 compared to 24.5 /~g m 3, the increase in the relative amount of nitrate, and the increased aliphatic dicarboxylic acids. The relative importance of secondary organic particle formation from reactions of gaseous precursors has been examined in a number of studies, particularly in the Los Angeles area. For example, Gray and coworkers (1986) reported that under conditions of high photochemical activity, as much as 27-38% of the organic particulate carbon may be due to secondary aerosol formation at downwind locations in this area. Similarly, Pandis et al. (1992) estimate that 15-22% of the particulate organic carbon is secondary in nature at downwind locations with aged air masses, compared to only 5-8% at upwind locations. Turpin and Huntzicker (1995) have shown that at Claremont, California, located east of Los Angeles, under many conditions primary organic emissions such as those discussed by Hildemann et al. (1994a)predominate in the organic portion of the particles. However, during periods of smog formation, about 40% was typically secondary in

407

nature, with as much as 50% being secondary in Pasadena (Turpin et al., 1991). Similar conclusions have been reached by Hildemann et al. (1994b). The fraction of total carbon and nitrogen in air that exists in the condensed phase is highly variable. For example, the percentage of organics in particles has been reported to be as high as 12.5% of the total carbon in the Los Angeles area, while as much as 25% of the nitrogen was found in particles (Fraser et al., 1996). Most studies of the chemical composition of particles in the troposphere to date have used analysis of bulk samples, which are usually collected in the boundary layer close to the earth's surface. As discussed in Chapter 6. J.3, there is a great deal of interest in the chemistry of the upper troposphere. Much less is known about the chemical composition in this region, particularly of particles. However, it appears that organics are also important constituents of particles in this region as well. For example, Novakov et al. (1997) in studies of particles both onshore and offshore of the eastern United States found that the mass fraction of the particles due to carbon compounds increased as a function of altitude. In the boundary layer, the fraction was typically ~ 10-40%, increasing to ~ 50-90% at an altitude of 2-3 km. Figure 9.55a shows the results of single-particle analysis (see Chapter 11.B.4a)of a typical particle in the upper troposphere (Murphy et al., 1998). In the negative ion spectra, a variety of fragments due to organics are observed, along with sulfates and some halogens. In other particles, soot and minerals were also common constituents. For comparison, Fig. 9.55b shows that a typical particle in the stratosphere is primarily sulfate (see Chapter 12.C.5). (3) Surfactants in aerosol particles The presence of long-chain organics having one or more polar functional groups (e.g., the carboxylic acids and the nitrates) in the condensed phase suggests that these may act as surfactants in aqueous atmospheric aerosols, forming an organic coating over the surface of the aerosol. Molecules that have long-chain (> C5), nonpolar groups attached to polar tails, such as those in Table 9.23, can form a surface film on droplets by lining up with the polar ends in the water and nonpolar, hydrophobic ends projecting into air as shown schematically in Fig. 9.56. The degree of compression of the film would be expected to alter its effects on uptake and evaporation from the particle. Thus, as organic films of linear-chain surfactants are compressed, they go through different stages, from a disordered 2-D "gas-like" phase at low

408

9. PARTICLES IN THE TROPOSPHERE a

Los Angeles Total fine particle mass 24.5 gg m 3

100%

Others

80%

Resolvable elutable organics 910 ng m 3

Unidentified organics / Other

Ammonium Nitrate

60%

Sulfate

40%

PAHs Aromatic polycarboxylic acids

Diterpenoid Acids

Aliphatic dicarboxylic acids

Elemental carbon

n-Alkenoic Acids

n-Alkanoic acids

20% Organics

n-Alkanes

b

Rubidoux

100% m

(downwind)

Total Fine Particle Mass 42.1 gg m 3

Resolvable Elutable Organics 1070 ng m -3

Others

Unidentified organics

80% r' ~~

Ammonium

Aromatic polycarboxylic acids

Nitrate

Aliphatic dicarboxylic acids

60%

40%

n-Alkenoic acids

Sulfate

20%

Elemental carbon

/ Other PAHs Diterpenoid acids

n-Alkanoic acids

Organics n-Alkanes

F I G U R E 9.54 Composition of particles in Los Angeles, at west end of air basin, and in Rubidoux, at east end (adapted from Rogge et al., 1993d).

degrees of compression to a liquid-like state to finally a highly ordered, solid condensed state (Gaines, 1966; MacRitchie, 1990). The compressed films would be expected to be less permeable. There is some evidence for this. For example, Rubel and Gentry (1985) measured the accommodation coefficient (see Chapter 5.E.1) for water as well as ammonia on acid droplets coated with hexadecanol. The water accommodation

coefficient decreased from 8 • 10 3 as the alcohol coverage increased, i.e., as the degree of compression of the organic film increased, to 4 • 10 4, with a sharp change at the point that the film underwent a phase transition from the liquid to the solid condensed state. Similarly, D~iumer et al. (1992) showed that coating an H2SO4 aerosol with straight-chain organics retarded the rate of neutralization by ammonia, whereas

409

C. CHEMICALCOMPOSITION OF TROPOSPHERIC AEROSOLS 0.4-

a

0.3

o\c
.m

0.0

Class of compound

H

CHO2"

35CI .~, I

k~rl

0

I

/

_] _~

20

40

Hsoa-II a4 so -/ ii, so, \ 79Br-"

60

"SO4

"

80

b

0.3

Structure

HSO4k

O"

>,

TABLE 9.23 Structures of Some Classes of Surface-Active Molecules Found in the Atmospherea

C N & C2H2

100

120

i

R--CH 2OH

Acid

RmC ~O ~OH

Aldehyde

RmC ~O ~H

Ketone

R~C ~O ~R'

Ester

J

HSO4-

rv"

Alcohol

R~C//O ~OR' R--NH 2

Amine

0.2

a From Gill et al. (1983). HSO3-

0.1

0.0

O" 0

SO3

OH-

20

/

/

"

S-

/i

40

60

m/z

80

100

120

FIGURE 9.55 Typical negative ion mass spectra for a single particle in the (a) upper troposphere (14.6 km, 22~ and (b) stratosphere (19 km, 31~ (adapted from Murphy et al., 1998).

branched molecules did not, presumably because the permeability of the films was much larger. The representation in Fig. 9.56 is simplified in that such an orderly arrangement applies to surfactants having saturated hydrophobic chains that can compress to an ordered solid condensed phase. However, this is not the case for all potential surface-active compounds in the atmosphere; for example, the presence of a double bond as in the case of oleic acid, C H 3 ( C H 2 ) 7 C H = C H ( C H 2 ) 7 C O O H , gives the molecule a "crooked" shape, which requires more surface area per molecule and which does not lead to a well-ordered solid condensed phase at the interface. Thus, Xiong et al. (1998) measured effects of organic films on the uptake of water into particles of H zSO 4. The saturated straight-chain lauric and stearic acids significantly reduced the uptake of water when present at amounts equivalent to one monolayer, whereas oleic acid had no effect at this concentration. The possibility of such organic films being formed on aerosol particles in the atmosphere as well as on fog, cloud, and rain droplets and snowflakes has been discussed in detail by Gill and co-workers (1983). As seen from our earlier discussions on the types of organics that have been observed in both urban and nonurban aerosols, there is no question that surface-active species

that can form organic films on water do exist in the atmosphere. However, Gill and co-workers (1983) have estimated, based on the limited data available, that only in aerosol particles does sufficient surface-active organic clearly exist to form a film around the particle. They suggest that the presence of surface-active films on cloud droplets, snowflakes, and raindrops is unlikely. While the focus has been on long-chain surfactantlike organics, Donaldson and Anderson (1999) suggest, based on their measurements of the standard free energy of adsorption of gases onto water, that there may also be significant surface coverages of smaller organics under atmospheric conditions. There is some field evidence for the existence of organic films on the surfaces of particles. For example, Fig. 9.57 shows the results of electron microscopy of haze aerosol collected in Los Angeles (Husar and Shu, 1975). The droplets are "wrinkled" in appearance, and they suggest this is due to "haze aerosol" droplets being coated with an organic layer that collapsed when the water in the particle evaporated under vacuum. Husar and Shu propose that the wrinkled appearance is due to a nonvolatile layer of organics that shrunk after water evaporated from the particle during analysis;

/ / / / /HY~176176176176 Water surface Hydrophilic group

FIGURE 9.56 The orientation of surface-active organic molecules at the water surface (adapted from Gill et al., 1983).

410


B

,i :"

9

_...^~, . . "

G

:.-~.-...,. ".'/"; "5 It

.

.

.

.

.

.

9

9"" ,)v

.... :: i: :.,-i

:-~"Z

...~..:..)~,.. 9

9

9

.

- , . . - ?.-; - ~ . . . . . . .:, ....

".,." ..~.." ~._ 9 _"';" . " "_'.",

..""...'.':

9" ~ .

,

.

.

..,: .

9 .

!

..

. ... 9

'~.

"

...

.

~...;.....'..~.,.: 9 :":~ ::"-i

-" 11ll~-

,: .27

.

' .. -" .~!,~._..;-.~

-,

9 ~ :. :~j.....,.-.

'g".:2"5

."

%:-

~..'.. _ ."

".-~:'."i,.

{:

."

,..,,~.--.2.

.0.:-.-~-,_,

;d,.,.

....

".

i

~: .

-

9 9

"

9". %-' "....-~ .... . :

t

9 ,..:' ..!.

;i-.'.§

~-

.. ~

'

.

~.

.,

il.""

i~",

..... ,

9."

"-.".' "7. ,..

;i:'i;:: i " " z-:

. "...a,.

9 9

~

"...". "

:" r 0e'.>: ".:

.'~ ~:.'a',~.

",

".. ..~Q.~:.,. ...... 9 .

-;iiliP

7-.

~":'-'~". "'.'"

' '{i:~::

> b k P and (OO) reduces to S0 S

~L

bL P.

(PP)

However, as first developed by Yamasaki et al. (1982) and subsequently by Pankow (1987), the fraction of the total surface sites that hold adsorbed molecules, A L, is just proportional to the concentration of adsorbed species on the particles, F (ng m 3), di-

the temperature dependence of Pc as log Pc (Torr) c T d. Figure 9.61 shows some typical plots to test Eq. (VV) for some SOC in the form of higher molecular weight alkanes and some PAH (see Chapter 10) at various relative humidities (Storey et al., 1995). In this case, the gas-particle partitioning coefficient Kp has been normalized to take into account different surface areas by using a surface area normalized coefficient, Kp, s, defined as Kp, s

SOC

ON

Kp (Specific surface area of adsorbing substrate).

(WW)

Figure 9.61 shows that log Kp, s is indeed linear with

PARTICLES

vided by the total concentration of total suspended particles, TSP (ng m 3), where the proportionality constant M contains the conversion constants to convert the TSP concentration into the total surface area and F into the area of the adsorbed molecules: AL

kl k 1 SOC(ads ).

If the gas-phase pressure of SOC(g) is P, then

klP(S

OF

M ( F TSP).

(QQ)

Similarly, the right-hand side of Eq. (PP) can be converted into a term involving the gas-phase concentration A in Eq. (MM) using the ideal gas law, P (n V)RT. The number of moles per m 3, (n V), is converted to ng m 3 using the molecular weight (MW) of the SOC. Combining the MW, R, T, and other conversion factors into one constant N, P NA, and the right side of Eq. (PP) becomes b c P b LNA. Equation (PP) therefore becomes M ( F TSP)

b L NA

or

bL

( M N ) ( F TSP) A.

(RR)

However, as seen in Eq. (OO), b L k 1 k l; i.e., b L is in effect an equilibrium constant for the adsorption of SOC(g) on the solid and the reevaporation of SOC(ads) from the solid surface. Hence b L k 1 k_l K' and Eq. (RR) becomes K'

(Constant)(F TSP) A.

(SS)

In short, the form of the gas-particle partitioning defined in Eq. (MM) is consistent with Langmuir adsorption of the SOC on the surface of the TSP.

log Pc, and the slopes are typically close to 1 as expected. Fig. 9.61 also illustrates the effects of relative humidity ( R H ) o n the adsorption of these compounds on a quartz fiber filter (QFF); the values of Kp, s decreased by about an order of magnitude as the RH increased from ~30 to 70% which was attributed to changes in the properties of the surface as water adsorbed onto it. Furthermore, the values for adsorption of these compounds on urban particulate matter (UPM) that had been observed in other studies (Yamasaki et al., 1982; Foreman and Bidleman, 1990) are clearly much larger than on the QFF. This suggests that if mineral oxide particle surfaces in the atmosphere behave like the QFF, adsorption to such inorganic sur-

415

D. GAS-PARTICLE DISTRIBUTION OF SEMIVOLATILE ORGANICS

TABLE 9.24 Subcooled Liquid Vapor Pressures (pL) at 20~ for a Series of Polycyclic Aromatic Hydrocarbons" and Their Temperature Dependence b PAH

Fluorene Phenanthrene Anthracene Fluoranthene Pyrene Benzo a fluorene Benzo b fluorene Benz a anthracene Chrysene Triphenylene Benzo b fluoranthene Benzo k fluoranthene Benzo a pyrene Benzo e pyrene

log PL (20~ (Torr) 2.72 3.50 3.53 4.54 4.73 5.24 5.22 6.02 6.06 6.06 7.12 7.13 7.33 7.37

PL (Torr)

c

1.9 3.2 3.0 2.9 1.9 5.8 6.0 9.6 8.7 8.7 7.6 7.4 4.7 4.3

3632 3982 4004 4464 4529 4792 4814 5179 5200 5208 5711 5711 5777 5799

• 10 3 • 10 4 X 10 4 • 10 5 • 10 5 • 10 6 • 10 6 X 10 7 X 10 7 X 10 7 • 10 s • 10 s • 10 s • 10 s

d 9.68 10.09 10.14 10.70 10.73 11.11 11.21 11.66 11.69 11.71 12.37 12.36 12.39 12.42

a From Pankow and Bidleman (1992). Original data were for 25~ and published by Yamasaki et al. (1984). They were corrected to 20~ by Pankow and Bidleman (1992). b Temperature dependence given by log PL (Torr) c T d.

faces will not be significant in urban areas (but may be more important in remote areas). It should be noted that a similar trend in Kp with relative humidity appears to apply to the gas-to-particle distribution of PAH in urban areas (Pankow et al., 1993), although in studies by Cotham and Bidleman (1992), R H did not appear to significantly affect the adsorption of organochlorine pesticides, at least over the range of R H of 30-95%. While many sets of data appear to follow Eq. (VV) relatively well, with slopes of m r 1 as predicted, deviations in the values of m r and b r a r e often observed. There are a number of reasons for such deviations (e.g., see Pankow and Bidleman, 1992). For example, changes in temperature, concentrations of SOC, and relative humidity during sampling, nonattainment of equilibrium, and sampling artifacts can all lead to deviations from the predicted, equilibrium relationship. In addition, if ( A H d AHva p) in Eq. (UU) is not constant along the series, relationship ( V V ) w i l l not hold because the value of b r is changing. However, it is also the case that slopes different from m r 1 can occur even under equilibrium conditions, due to the nature of the molecular interactions involved in adsorption on a surface. For example, Goss and Schwarzenbach (1998) propose a modified formulation of Eq. (VV) that more explicitly takes into account the interactions between the SOC and the surface on a molecular level. Thus, based on work by Goss (1997), they express the relationship between a gas-particle partition coefficient K iads defined as the adsorbed con-

centration on the surface (in mg m 2) divided by the gas-phase concentration (in mg m 3), and pL in the following form" In

g i ads

[

0.133(3/vdw)0"5]In PL 2.08/3i(y )0.5

2.09(yvdw) ~

1.37c~i(y )o.5

19.5.

(xx) The first two terms represent van der Waals interactions between the adsorbed SOC and the surface, which would apply to all SOC. The second two terms represent Lewis acid-base interactions, which can be important for compounds containing O, N, or aromatic rings, for example, the adsorption of alkyl ethers on the polar surface of quartz. The y coefficients (in mJ m 2) describe the surface properties, where yvdw is associated with its van der Waals interactions with adsorbing gases, y describes its electron-acceptor interactions, and y describes the electron-donor interactions of the surface. On the other hand, the properties of the adsorbing species are described by In PL for the van der Waals interactions and by the dimensionless parameters /~i and c~i, which relate to the electron-donor and electron-acceptor properties (if any), respectively, of the adsorbing molecule. This formulation explicitly accounts for both the properties of the surface and those of the adsorbing SOC that determine on a molecular level the amount of adsorption of the gas on the surface, and hence the gas-particle partitioning. For example, Goss and Schwarzenbach (1998) describe the implications of three

416


a

given by 0.133(yvdw) ~ but there is an extra term in the intercept, corresponding to the y or y term in Eq. (XX); hence the line will parallel that for the first case but be shifted from it by a constant amount. For the third case, the terms y and y vary with the compound, but often in a manner that follows In PL as well, because of the common molecular properties that determine both terms. For example, the aromatic rings of PAH act as electron donors, with the electron-donor parameter /3i a measure of the Lewis acid-base interaction. However, as this interaction changes for a series of PAH, there is a concomitant change in their vapor pressures, PL. As a result, the term in Eq. (XX) expressing the Lewis acid-base interaction, 2.08/~i( T )0.5, rises proportionally to In PL. The slope of a plot of In Kiads against In PL is thus larger than for the first two cases. In short, the slope of plots of the natural logarithm of the gas-particle partitioning coefficient against In PL for a series of adsorbing SOC and surfaces can help to elucidate on a molecular level the types of interactions between the two. The partitioning of an SOC between the gas and particle would be expected to depend on temperature. Intuitively, one expects that an increase in temperature would result in less adsorption and a higher gas-phase concentration. Indeed, the temperature dependence of Kp in Eq. (MM) is usually (Pankow, 1987, 1991, 1992) expressed as

Alkanes

FF: ~30%~_x.~p M ~70%z}.

m

)

0

--

X

Od

E I

E

v

if}

d. O'1 O

--

6

I

-7

-8

I

-6

I

-5

b

-4 PAHs

5

m

32

~30%, QFF O ~ UPM ~70%,QFF

1 0-1

-8

I -7

I -6

I -5

I -4

log Kp -3

O

Log PL (torr) F I G U R E 9.61 Surface area normalized gas-particle partition coefficients as a function of liquid vapor pressure for some alkanes and PAHs on urban particulate matter (UPM) and quartz fiber filters (QFF) at 30% R H and 70% R H (adapted from Storey et al., 1995; data for U P M from Yamasaki et al. (1982) and Foreman and Bidleman (1990)).

different cases for the slopes of plots of In Kiads against lnpL" (1)where no Lewis acid-base interactions occur (e.g., /~i O/i 0 and or y y 0), for example, the adsorption of alkanes on Teflon; (2) compounds with the same Lewis acid-base interactions with the surface but differing van der Waals components, e.g., alkylbenzenes on quartz; and (3)compounds with different functionalities having different Lewis acid-base interactions. In the first case, slopes of In Ki~d~ against In PL should reflect the van der Waals interaction term 0.133(yvdw) ~ . In the second case, the slope is again

Cp T

dp.

(YY)

This expression is developed in detail by Pankow (1987), and its origin treated in Box 9.2. Figure 9.62 shows some typical plots by Pankow (1991) of log (F TSP) A against 1 T for some PAH measured by Yamasaki et al. (1982). The plots are reasonably linear, as expected from Eq. (YY). Pankow (1991, 1992) shows that since dp is expected to be similar for similar compounds Eq. (BBB), assuming a single value of dp for such a group seems to be justified. In summary, adsorption of semivolatile organic compounds (SOC) on solid particles in the atmosphere is expected to occur, leading to partitioning of such compounds between the gas and condensed phases. As expected, this partitioning is temperature dependent, with increasing amounts adsorbed on the particles as the temperature is lowered. The relationship between the logarithm of the measured gas-particle partitioning coefficient and the logarithm of the vapor pressure of the liquid SOC at that temperature (subcooled, if necessary) is expected to be linear, and a slope of 1 is common. However, this slope and deviations from 1

417

D. GAS- PARTICLEDISTRIBUTIONOF SEMIVOLATILEORGANICS

BOX 9.2 TEMPERATURE

DEPENDENCE

The origin of the relationship between Kp and temperature in Eq. (YY) can be seen by reexamining equations (RR) and (SS), which show that Kp is directly proportional to b c, i.e., to the ratio of the rate constants for adsorption and desorption, k 1 k_v The rate constant for desorption from the surface, k_l, can be expressed as a function of the heat of desorption, zXHd (Adamson, 1982): k 1

k 0 exp(

AH d R T ) ,

(ZZ)

where k 0 is approximately 1012-1013 s 1; i.e., if the heat of desorption is zero, the molecule desorbs within the period of one vibration. Since b L and hence Kp depend on 1 k_l, they should vary with exp( A H d RT); i.e., log Kp should vary with 1 T, with the slope related to the heat of desorption of the SOC from the solid surface. The rate constant for adsorption, k l, is also temperature dependent, but the dependence is small compared to that for k_ 1. The value and temperature dependence of k 1 are determined by the rate of gas-solid collisions, which from kinetic molecular theory is given by Rate of collisions of gas with surface G ( R T 2~-M) 1 2 (number per cm 2 s 1),

(AAA)

can provide useful insights into the nature of the interactions between the SOC and the solid surface (assuming that equilibrium between the gas and solid is really achieved).

2. Absorption into Liquids As discussed by Bidleman (1988), there is a variety of evidence from both laboratory and field studies that the gas-particle partitioning in many cases in the atmosphere is consistent with partitioning to a liquid rather than adsorption on a solid; i.e., the gaseous SOC is absorbed into a liquid particle or a liquid on the surface of a solid. For example, tetrachlorodibenzo-pdioxin (TCDD) is found in both the vapor and particle phases, whereas none would be found in the gas phase

OF

iKp

where G is the gas concentration (number cm 3), T is the temperature, and M is the molecular weight of the gas. Replacing G by P R T using the ideal gas law, the rate of collisions of the SOC with the surface, and hence k l, varies with T 1 2, which is usually small compared to the exponential dependence on temperature of k_ 1. In short, a plot of l o g ( F TSP) A against T 1 should, in principle, be linear over the relatively narrow range of temperatures commonly found in the troposphere, and hence be well-described by Eq. (YY). Thus such plots should give a straight line of slope Cp and intercept dp. Pankow (1987, 1991) shows that Cp and dp should be given by Cp

(AM d 2.303R)

Tamb 4.606

(BBB)

and dp

log{Nws p 275k0(MW ramb)l 2}

1 4.606,

(ccc) where Tamb (in K) is the middle of the temperature range over which Kp is studied, ATSp is the specific surface area for the particles (cm 2 /xg 1), k0 (s-l) is as defined in Eq. (ZZ), and MW is the molecular weight of the SOC.

if the solid vapor pressure controlled the distribution. Similarly, the distribution of PCBs between the gas and particle phases is consistent with absorption into a liquid (Falconer and Bidleman, 1994). In this case, the gas-phase concentration of the organic need not exceed its saturation vapor pressure to partition into the particles since the process, in effect, involves dissolving the gas into a solution. If an SOC is absorbed into a liquid organic layer on the particle, a relationship between Kp and Pc that is similar to that developed for adsorption onto a solid can be derived (Pankow et al., 1994a, 1994b). In this case, the gas-particle partitioning coefficient for the ith compound is defined as

Kp (Fi,om TSP)

A/,

(DDD)

418


t (~

40

0

a "1

--

"2

--

Phenanthrene + Anthracene

O.

--

-6

where fi, om is the concentration of i in air (ng m 3) that is particle associated and dissolved in the liquid organic material, "om," phase, TSP has units of /xg m 3, and A is the gas-phase concentration in units of ng m 3. As for adsorbed SOC, Eq. (MM), Kp has units of m 3 per /xg. The thermodynamics of absorption are described in Box 9.3. The expression for the gas-particle partitioning coefficient Kp given by Eq. (DDD) in units of m 3 per/xg is developed by Pankow (1994a, 1994b) as

O

f i,om T S P

o

3.2

Kp

I

I

I

I

I

3.3

3.4

3.5

3.6

3.7

3.8

I

b

Fluoranthene

_

-1

,,< co

106 MWom Tom p L ,

(JJJ)

where MWom (g mol 1) is the mean molecular weight of the species constituting the o m (liquid organic matter) phase, fom is the weight fraction of the total suspended particulate matter that is the absorbing o m phase (in addition to organics, this includes any inorganics and water that may be present), and the other symbols have their usual meaning discussed earlier. Taking the logarithm of (JJJ), one obtains Eq. (KKK):

--

-2

v

o

Ai

760Rrfo m

O

-3 -

log Kp

J

-4 -5 3.2

I

I

I

I

I

3.3

3.4

3.5

3.6

3.7

C

3.8

BaP + BeP

:f~ 3.2

I

I

I

I

I

3.3

3.4

3.5

3.6

3.7

3.8

1000/T (K 1 )

F I G U R E 9.62 Plots of log (F TSP) A against T 1 for some PAH from Pankow (1991) based on data of Yamasaki et al. (1982).

log P L

log ")/om

log

7 6 0 R T f om

106 MWom 9 (KKK)

Thus, Eq. (KKK) and the analogous logarithmic form of Eq. (III) in Box 9.3 predicts that a plot of log K P against log PL for the partitioning of a series of compounds into liquid particles or into a liquid layer on particles should be a straight line with a slope of 1 if the activity coefficients in the liquid phase, Yom, remain constant. Figure 9.63, for example, plots log Kp against log PL for the partitioning of a series of PAH (see Chapter 10) between the gas phase and particles of either dioctyl phthalate (DOP) or secondary organic aerosol (SOA) generated from the photooxidation of gasoline vapor (Liang et al., 1997). The slope of the plot for uptake into DOP is 1.09 and that for uptake into SOA is 1.05, in agreement with Eq. (KKK) if the activity coefficients are the same for these compounds. However, as pointed out earlier, a slope of 1 is only expected when the activity coefficients do not change along the series of SOC. In fact, Yi frequently does change. For example, Jang and Kamens (1998) reported that the activity coefficients for the partitioning of alkanes and PAH on wood soot particles increased with the relative humidity. Furthermore, changes in the activity coefficients are often correlated with the changes in PL (Goss and Schwarzenbach, 1998). In that case, since In Yi is proportional to In PL,

419

D. GAS - PARTICLE DISTRIBUTION OF SEMIVOLATILE ORGANICS

B O X 9.3

RELATIONSHIP SUBCOOLED

BETWEEN VAPOR

Following the approach of Goss and Schwarzenbach (1998), for equilibrium between a species, i, in the gas phase at pressure Pi and dissolved in a liquid organic material at a mole fraction Xi, om, the chemical potentials (/x) in the two phases must be equal, i.e., ]&i,g

]&i, om"

ix i

R T ln(P i p0),

(FFF)

where po is a standard-state pressure and ]&i,g is the associated chemical potential of the standard state. Taking the standard state as the pure liquid, P~ is the saturation vapor pressure over pure liquid (subcooled, where appropriate), i.e., P PL. Similarly, the chemical potential for the i th compound in

the slope becomes (1 s), where s reflects the simultaneous changes in In Yi when PL changes. They conclude that the slope should lie between about 0.2 and 1.0 when the possible changes in the activity coefficients are taken into account. For example, the slope of In K i defined by Eq. (III) as a function of In PL for the partitioning of a series of alkanes and alkylbenzenes between the gas phase and liquid octanol at 25~ is about 1; however, that for partitioning of chlorobenzenes into octanol is only 0.59. Goss and Schwarzenbach (1998) attribute this to the smaller attractive forces between the chlorobenzenes and octanol compared to those in the pure liquid chlorobenzenes. Because the slopes of log Kp against log PL can be 1 for both absorption into a liquid layer and adsorption onto a solid, the slope alone cannot be used to differentiate the mechanism of gas-particle partitioning. However, a combination of the slope and the absolute values of Kp can be used to test the two mechanisms for cases where only van der Waals interactions occur, since Kpads can be calculated from Eq. (XX). For example, Goss and Schwarzenbach (1998) show that the calculated values of Kpads for a series of alkanes and PCBs partitioning into typical urban

AND SOC

the liquid organic layer is given by /xi, om

/Xi

R r ( l n %Xi,om),

(GGG)

where the standard state is the pure liquid at 1 bar pressure. Combining Eqs. (EEE) through (GGG), one obtains

Pi PL

(EEE)

The chemical potential in the gas phase is given by iXi,g

ABSORPTION PRESSURE OF

")liSi,om 9

(HHH)

Goss and Schwarzenbach (1998) define a unitless gas-particle partitioning coefficient, Ki, as the ratio of molar concentrations in the condensed, om, phase (Ci, om) to that in the gas phase (Ci,g). They show that using Eq. (HHH), K i can be given by Ki

where layer.

gom

Ci,om Ci,g

RT ~ , %pLVom

(III)

is the molar volume of the liquid organic

aerosols are much smaller than reported in the literature, indicating that absorption into a liquid phase must be important. As is the case for adsorption, the trend in gas-solid partitioning coefficients for absorption can be used to probe for similarities and differences in interactions between the gaseous SOC and liquid on a molecular level. For example, Goss and Schwarzenbach (1998) point out that if two liquids and their interactions with SOC are the same, a plot of the gas-particle partitioning coefficients for a series of gases in one liquid, Ki,1, against the analogous coefficients in the second liquid, Ki,2, should be a straight line with a slope of 1.0. If the interactions are different, however, this correlation will not hold. For example, such a log-log plot of the gas-solid partitioning coefficients for the uptake of a series of alkanes and ethanol in n-heptane and dibutyl ether shows that the correlation does not hold for ethanol. This indicates that the two liquids are not chemically similar in terms of their interactions with the SOC, which is not surprising in this case because of the strong interactions between ethanol and the ether. If the liquids were ambient air particles of unknown surface composition, such a plot would show that the

420


is usually less sensitive to changes along a series of compounds than is a single value of the activity coefficient. In this approach, the parameter PL is replaced by a different descriptor of an SOC's volatility, its octanol-air partition coefficient, Koa (Finizio et al., 1997; Harner and Bidleman, 1998; Pankow, 1998). This is defined for a given SOC partitioned between liquid octanol and air as

AH -2

F' G~'

-3I~ mO

-4--

-5

m

-6

--

D

C'

KOA B "7

-8 -8

I

I

I

I

I

I

I

-7

-6

-5

-4

-3

-2

-1

0

log PLo (Torr)

FIGURE 9.63 Plots of om-phase-normalized gas-particle partitioning constant log Kp,om vs logarithm of the subcooled liquid vapor pressure, log PL, for a series of semivolatile PAHs partitioning on (o) dioctyl phthalate (DOP) or ( 9 secondary organic aerosol (SOA) from photooxidized gasoline vapor. PAHs are as follows: naphthalene, A; acenaphthalene, B; fluorene, C and C'; phenanthrene, D and D'; anthracene, E and E'; fluoranthene, F and F'; pyrene, G and G'; chrysene, H (adapted from Liang et al., 1997).

particles are quite different in their chemical properties and hence their effects on gas-particle partitioning. In short, differentiating between adsorption on a solid and absorption into a liquid for partitioning of semivolatile compounds in the atmosphere is often difficult to do in an unambiguous manner. However, the use of a combination of approaches can help to differentiate these two mechanisms and, perhaps more important, give some insight into the mechanisms of interaction of the SOC with the condensed-phase material.

3. O c t a n o l - Air Partitioning Coefficients One problem with the use of PL as a key parameter in both adsorption and absorption is the difficulty in obtaining accurate values for PL for solid SOCs, since they are not experimentally accessible and must be estimated (e.g., see Finizio et al., 1997, and references therein). In addition, as discussed in the preceding section with respect to absorption into a liquid phase, slopes of 1 for plots of log Kp against log PL are only expected if the activity coefficients, Yi, do not change along a series of compounds. An alternate approach has been proposed that avoids the use of PL and introduces a ratio of activity coefficients rather than an absolute, single value. The advantage in terms of the activity coefficients is that the ratio

C O CA,

(LLL)

where C O and CA are the concentrations of the SOC in octanol and in air, respectively, both in units of mol m 3. A distinct advantage of the use of KOA compared to P L is that it can be directly measured (e.g., see Harner and Mackay, 1995; Harner and Bidleman, 1996) or can be estimated from knowledge of two well-established and widely used partition coefficients, Kow for octanol-water and KAW for air-water (e.g., see Mackay et al., 1992; and Baum, 1998). The relationship between Kp and KOA c a n be readily developed based on thermodynamic principles (e.g., see Finizio et al., 1997); see Box 9.4. Thus, the relationship between Kp and KOA by Finizio et al. (1997) leads to Eq. (PPP): log Kp

log KOA

Yo] log ~

log

760fom(MW)~ 1 109MWom Po (PPP)

To test the validity of using log KOA rather than PL as a direct descriptor of the volatility of SOCs, Finizio and co-workers (1997) calculated values of log Kp (normalized to 25~ from published ambient air data on gas-to-particle partitioning of PAHs (as well as PCBs and organochlorine pesticides) at seven sampling sites ranging from urban to rural airsheds. Koa values were obtained from the literature or calculated from the ratio Kow Kaw, where Kow is the octanol-water partition coefficient (e.g., see Mackay et al., 1992; and Baum, 1998) and Kaw is the dimensionless air-water partition coefficient calculated from literature values of Henry's law constants. Figure 9.64 shows a plot of log Kp against log Koa for the PAH data. A linear relationship is observed, as expected from Eq. (PPP). The slope of 0.79 is below the theoretical value of 1.0, reflecting the fact that the slopes plotted in the conventional manner, i.e., log Kp vs log PL, were also significantly less than the theoretical value of m r 1.00. The authors cite several possible reasons, including nonequilibrium conditions as described by Kamens et al. (1995). Figure 9.65a shows the percentages of the 3- and 4-ring PAHs, as well as some PCBs and polychlorinated naphthalenes (PCNs), found on aerosol particles in the

421

D. GAS-PARTICLE DISTRIBUTION OF SEMIVOLATILE ORGANICS

BOX 9.4

BETWEEN Kp A N D KOA

RELATIONSHIP

First, using the ideal gas law, the partial pressure (which is equal to the fugacity) of the SOC in air is given by CART. Over the octanol solution, the fugacity (fi) of the SOC is given by Pi PL; i.e., its partial pressure over the octanol solution is given by Pi fiPL" Using fi Xi, o~,o, the partial pressure of the SOC over an octanol solution containing a mole fraction Xi, ~ of the SOC is given by P~ (Xi, o~,oPL), where the symbols have their usual meaning of mole fraction, activity coefficient, and liquid vapor pressure, respectively, all in octanol. The mole fraction can be replaced in this expression, since Xi, 0 Co(MW o) 10300, where C O is the concentration of the SOC in octanol in units of mol m 3, MW ~ is the molecular weight of octanol, and 00 is the density (kg m 3) of octanol. Thus, the fugacity of the SOC above the octanol solution must be Pi Co(MWo)~/oPL 1030o, and we know that P~ CART. Hence

o

-3 -4

CA

(MWo)ToPL

Kp

760RTfom

(MWo)Yo KoA

1030oRT

(To)[ K~

~

~

760fom(MW)~ 109MWom 0o

BeP @

9Green Bay (Wisconsin) 9Chicago (Illinois)

BkF

9Brazaville (Congo) [] Osaka (Japan)

CHR FL

x Portland (Oregon) O Lake Superior (USA/Canada)

AN

- - average

-5 -6

i 6

I 7

i

I 8

i

I 9 log

i

.

106MWom Tom

Denver (Colorado)

PH

FLE O

1030oRT (NNN)

This, then, gives the relationship between PL and the octanol-air partitioning coefficient, KoA. Rearranging Eq. (NNN) to obtain an expression for PL and substitution into Eq. (JJJ) give (Finizio et al., 1997)

PY m

CO KOA

[]

-2

(MMM)

Equation ( L L L ) t h e n becomes

BaP

BaA m

Co(MWo)ToPL 10300 .

CART

I 10

i

I 11

i 12

KOA

FIGURE 9.64 Plot of log Kp for 10 PAHs in ambient air samples collected at seven sites worldwide as a function of calculated values of log Koa. FLE fluorene, PH phenanthrene, AN anthracene, PY pyrene, FL fluoranthene, BaA benz a anthracene, CHR chrysene, BaP benzoa pyrene, BeP benzo e pyrene, BkF benzok fluoranthene (adapted from Finizio et al., 1997).

(ooo)

422

9. P A R T I C L E S IN T H E T R O P O S P H E R E

urban Chicago area in February and March (Harner and Bidleman, 1998). The solid line shows the predicted percentages calculated using the Junge-Pankow adsorption model, Eq. (TT). Figure 9.65b treats the same data set in terms of absorption into a liquid using the octanol-water partitioning coefficient assuming the fraction of organic matter, fom, in the particles is either 10 or 20%. Both models are in reasonable agreement with the PAH data for winds from the southwest. With winds from the northeast, enrichment of the PAH in particles is observed compared to model predictions using the absorption model (Fig. 9.65b), which the authors suggest could be due to nonattainment of equilibrium or to the trapping of nonexchangeable PAHs in the particles. The adsorption model overpredicts the percentages of PCBs expected to be found in the particles, but the KOA absorption model is effective in explaining the field partitioning data for PCBs and PCNs. In short, data such as those in Figs. 9.64 and 9.65 support the use of the octanol-air partitioning coefficient as a useful parameter for characterizing gas-particle partitioning of SOC into liquid particles or liquid layers on particles in air. As discussed elsewhere in this book, there is increasing evidence for reactions at the air-water interface in the atmosphere. Pankow (1997) has treated partitioning of gases to the interface as well and predicts that as for adsorption on a solid and absorption into a liquid, there should be a linear relationship between log Kp and In PL with a slope of approximately 1. This discussion of gas-particle partitioning has focused on the idea that equilibrium between the two phases is attained in the atmosphere. However, it should be noted that equilibrium cannot always be assumed in the atmosphere. For example, Wania et al. (1998) ex-

lOO

amined the exchange of SOC between the air and the surface under two assumptions, one being equilibrium and one that treats the kinetics of the exchange of the SOC with the soil and with incoming and outgoing air. They show that some atmospheric observations can be explained by the kinetic effects. For example, the concentration of y-hexachlorocyclohexane (HCH) in the gas phase has a significant temperature dependence while c~-HCH does not. This is unexpected for two such similar molecules if equilibrium gas-particle partitioning controls the gaseous concentrations. Wania et al. (1998) suggest that c~-HCH concentrations are determined largely by the kinetics of transport of air containing this compound to the measurement sites, since it is distributed globally and hence there are large reservoirs such as the oceans, y-HCH, on the other hand, has lower global "background" levels but has been used extensively in industrialized countries, leading to higher concentrations downwind. Hence they attribute the y-HCH levels, measured at sampling sites located downwind, to a shift in the equilibrium toward the gas phase as soil temperatures increased, leading to increased evaporation. Because of such "real-world" nonequilibrium situations, some efforts, for example, by Turco and coworkers (Jacobson et al., 1996), Kamens and co-workers (Odum et al., 1994), and Seinfeld and co-workers (Bowman et al., 1997), have focused on developing dynamic models to describe gas-particle distributions. In a similar vein, the time scales to achieve equilibrium for inorganics have been examined by Meng and Seinfeld (1996), who show that small (submicron) particles can come to equilibrium with the gas phase in less than a few hours typically but that larger particles may not. The major factors determining the time needed to reach equilibrium are the aerosol size distribution,

......~

b

a

80

i'-

60

oOoo

40

~ 1 7 6 zz

~Oo%~

20 n u

i

-6

I

-5

B

'~

i

I

-4

i

~

-a

"

TIIt~L~

-2

[]

I

-1

i

i

7

i~-

8

~lm~/i-",

9

i

10

11

,

i

12

,

13

FIGURE 9.65 Observed percentages in the particle phase of PAHs (( ) N-E air; (*) S-W air), PCBs (( ) nonortho, ( ) monoortho, and multiortho PCBs, respectively), and (o) PCNs in Chicago air compared to model predicted values. (a) Solid line is calculated with Junge-Pankow (J-P) adsorption model (Eq. (TT)). (b) Solid and dotted lines are calculated with absorption model for aerosols assumed to contain 10 and 20% organic matter (om), respectively (adapted from Harner and Bidleman, 1998).

aEFEaENCES

temperature, and accommodation coefficient for uptake of the gas into the particle (see Chapter 5). In short, traditional equilibrium theories of gas adsorption provide at least a qualitative framework for describing the partitioning of SOC between the gas and particle phases in the atmosphere, and more recent theories are providing further insight into these partitioning processes in the atmosphere.

E. P R O B L E M S 1. Use the Kelvin equation (Chapter 14.C.2) to show that it is true that the vapor pressure of pure water at 25~ is only 0.1% greater over a 1-/~m radius particle than over a fiat surface, but 11% greater over a 0.01-/~m radius particle. The surface tension of water is 72 dyn cm ~ at 25~ 2. At what particle radius would the vapor pressure of water be twice that over a flat surface at 25~ The surface tension of water is 72 dyn cm ~ at 25~ 3. Use Eq. (GG) to predict the critical concentration of gaseous H2SO 4 in ppb for nucleation at a relative humidity of 40% and a temperature of 25~ How does this compare to the theoretical predictions and experimental observations shown in Fig. 9.30? 4. Write reaction mechanisms for the oxidation of 1-octene in the atmosphere that will give the major observed condensed-phase products heptanal and heptanoic acid. 5. In smog chamber studies of the oxidation of 1-decene, octanal was the condensed-phase product formed in the next highest yield after nonanal, nonanoic acid, and dihydro-5-pentyl-2(3H)furanone (Forstner et al., 1997a). Write a plausible mechanism for its formation. 6. Show how m-toluic acid could be formed from the abstraction reaction of OH with m-xylene (Table 9.22). 7. An important aspect of Eq. (LL) is that it shows that the yield of secondary organic aerosol depends on the mass of the condensed-phase organic available to take up semivolatile organic products. Show that Eq. (LL) can be derived for a single species. Use the relationship between C i and A (organic)and Ki, om Fi, om ( A i M o ) . Fi, om is the concentration of the species (ng m 3) in the condensed ("ore") phase and M 0 is the mass of condensed-phase organic available for uptake of the semivolatile organic. You will also need to use the mass balance restraint C i A i F~,om and the definition of the yield of secondary organic aerosol, Y [ 103 VczX(organic)]

EF~Vc,

where Vc is the volume of air in which the reaction takes place.

423

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C H A P T E R

10 Airborne Polycyclic Aromatic Hydrocarbons and Their Derivatives Atmospheric Chemistry and Toxicological Implications

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Polycyclic aromatic hydrocarbons (PAHs) are members of a unique class of air pollutants relevant to many scientific and societal issues having a variety of aspects: chemical, toxicological, engineering, technological, public health, economic and regulatory, and legislative. They are products of incomplete combustion formed during the burning or pyrolysis of organic matter and are released into ambient air as constituents of highly complex mixtures of polycyclic organic matter, POM. As defined in the U.S. Clean Air Act Amendments of 1990 (CAAA, 1990), POM "includes organic compounds with more than one benzene ring, and which have a boiling point greater than or equal to 212~ (100~ '' While we generally refer to specific PAHs and polycyclic aromatic compounds (PACs) throughout this chapter, it should be noted that the complex mixture represented by the term POM is the subject of many studies and regulatory designations. Atmospheric PAHs occur in the form of gases (e.g., 2-ring, highly volatile naphthalene, III), solids adsorbed/absorbed to the surfaces of fine respirable aerosol particles (e.g., 5-ring benzo[a]pyrene, BaP, I), and 3- and 4-ring semivolatile compounds that are distributed between the gas- and the particle-phases (e.g., the semivolatile 3-ring phenanthrene, IV, and 4-ring pyrene, II, and fluoranthene, V); for gas-particle partitioning, see Chapter 9D and Section A.4 in this chapter. The ubiquitous nature of these airborne PAHs is evident from the fact that the 16 U.S. Environmental Protection Agency "Priority Polycyclic Aromatic Hydrocarbon Pollutants" shown in Table 10.1 (U.S. EPA, 1988) are found, as we shall see in this chapter, in urban airsheds throughout the world. Their widespread

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I

II

Benzo[ a ]pyrene

Pyrene 10

8 7 6

~

1 2

~ 3

5

1

9 4

8

4

7

III

IV

Naphthalene

Phenanthrene 2

10 9(

4

8

5 7

6

V Fluoranthene

presence is due to their emissions from a wide range of combustion sources, including diesel and gasoline engines, biomass burning of agricultural and forest fuels (Jenkins et al., 1996), and outdoor wood smoke (Watts et al., 1988) (for summaries of the extensive literature 436

43 7

A. NOMENCLATURE AND SELECTED PHYSICAL AND SPECTROSCOPIC PROPERTIES OF PAHs AND PACs

T A B L E 10.1

Structures, Common Names, Empirical Formulas, Molecular Weights, Melting Points, Boiling Points, and CAS Numbers for the 16 U.S. EPA "Priority P A H Pollutants ''a'b Empirical formula

MW (g mol- 1 )

Mp (oc)

Bp (oc)

CAS number

Naphthalene

C s H 10

128.18

80.5

218

91-20-3

Acenaphthylene

C12 H s

152.20

92

265-275

208-96-8

Acenaphthene

C12

Hlo

154.20

96.2

277.5

83-32-9

Fluorene

C 13H 10

166.23

116

295

86-73-7

Phenanthrene*

C 14H 10

178.2

101

339

85-01-8

Anthracene*

C14 H10

178.2

216.2

340

120-12-7

Pyrene

C16 H lo

202.3

156

360

129-00-0

Fluoranthene

C16 Hlo

202.3

111

375

206-44-0

Benz[ a ]anthracene

C is H 12

228.3

160

435

56-55-3

Chrysene

C18 H 12

228.3

255

448

218-01-9

Common name

Structure

7 552 6

5

4

1

2

7

3

4 6

5

1

2

I

I

7

4 6

5

8

9

1

6

3 5

4

10

1

3

8 7 6 8

9

1

6

3 5

10

10

4

1

3 8 7 6

5 2

lO 9 8

5 7

6 2

4 10 9 8

1123 7

6

12

9

1

8 7

6

(continues)

438

10. AIRBORNE PAHs: CHEMISTRY A N D T O X I C O L O G I C A L IMPLICATIONS

T A B L E 10.1 (continued)

Common name

Structure

12 11~

Empirical formula

MW (g m o l - 1 )

Mp (oc)

Bp (oc)

CAS number

1 32

9

Benzo[ a ]pyrene

C20H 12

252.3

175

495

50-32-8

Benzo[ b ]fluoranthene

C20H 12

252.32

168

481

205-99-2

Benzo[ k ]fluoranthene

C20H 12

252.32

217

481

207-08-9

Benzo[ghi ]perylene

C22 H 12

276.34

277

525

191-24-2

Indeno[ 1,2,3-cd]pyrene

C22 H 12

276.34

163

-

193-39-5

Dibenz[a,h]anthracene

C22 H14

278.35

267

524

53-70-3

8 7

6

5 2

1 1 ~ 5 10 9

9

8

8 11

6

12

7

6

5

1 2

10

~

3

9

4

6

12 9 8

~ 7

1

3

4 6

5 2

12

4

10 9

a Adapted from Mackay et al. (1992); data on indeno[1,2,3-cd]pyrene from Harvey (1997). Structures and their numbering are based on IUPAC recommendations as described by Loening et al. (1990). b U.S. EPA (1998) designation. Exceptions are noted by asterisks.

on combustion sources, see Bj6rseth (1983), Back et al. (1991a), Venkataraman and Friedlander (1994c), Schauer et al. (1996), Harrison et al. (1996), Howsam and Jones (1998), and Simoneit (1998); for diesel and gasoline engines, see IARC (1989), Benner et al. (1989), Westerholm et al. (1991), Bagley et al. (1992), Johnson et al. (1994), Lowenthal et al. (1994), Hammerle et al. (1994), WHO (1996), Miguel et al. (1998), and Schauer et al. (1999); for coal fly ash, see Gohda et al. (1993); and for municipal incinerators, see Lee et al. (1993). PAHs are also common constituents of air indoors,

arising from coal and wood combustion (Mumford et al., 1990), wood combustion (Alfheim and Ramdahl, 1984), and environmental tobacco smoke, ETS (Gundel et al., 1995b; and the California EPA, 1997).

The physical and chemical processes by which PAHs and PACs are formed in combustion are very complex and beyond the scope of this book; the reader is referred to articles in the literature such as those by Badger (1962), Haynes (1991), and Vander Wal and co-workers (1997). Only a relatively few PAHs (ca. 100) are stable enough to survive the combustion-pyrolysis

439


process and enter our air environment as primary pollutants in complex combustion-generated mixtures in amounts sufficient to be of concern. The toxicity of such combustion-generated mixtures is reflected in the fact that POM is one of the 189 hazardous air pollutants (HAPs) cited in the 1990 U.S. Clean Air Act Amendments (Kelly et al., 1994; Kao, 1994). Their complexity is evident from the fact that the term POM includes not only PAHs, the most abundant and intensively studied chemical class in POM emissions and ambient air, but also a wide range of N-, O-, and S-atom polycyclic aromatic compounds, including nitroarenes and azaarenes; PAH lactones, ketones, and quinones; and thioarenes. Structures and nomenclatures for selected PAHs and PACs relevant to atmospheric chemistry are discussed in Section A.2. Some of these N-, O-, and S-atom-functionalized polycyclic aromatic compounds (PACs) are powerful bacterial mutagens and animal, and possible human, carcinogens, e.g., the exocyclic nitro-substituted PAH 1-nitropyrene, VI, which is a primary pollutant in diesel exhaust (IARC, 1989). The endocyclic heterocyclic PAC dibenz[a,h]acridine, VII, is also classified as a "possible human carcinogen" (IARC, 1987; see Sections B and C). NO 2

10

]1 13

14

1

3 4

12 8 11

7 6

5

6

10

VI

VII

1-Nitropyrene

Dibenz[ a,h ]acridine

The physical and chemical complexity of primary combustion-generated POM is illustrated in Fig. 10.1 (Johnson et al., 1994), a schematic diagram of a diesel exhaust particle and associated copollutants. The gasphase regime contains volatile (2-ring) PAHs and a fraction of the semivolatile (3- and 4-ring) PAHs. The particle-phase contains the remainder of the semivolatile PAHs ("particle-associated") along with the 5- and 6-ring heavy PAHs adsorbed/absorbed to the surface of the elemental carbon spheres that constitute the "backbone" of the overall diesel soot particle. Also present is sulfate formed from oxidation of sulfur present in the diesel fuel and gas- and particlephase PACs. The elemental carbon core particles ("black carbon") range from ca. 0.01- to 0.08-/xm aerodynamic diameter. They not only add significantly to the total mass of the aerosol particle but, as seen in Chapter 9, are also a

Solid elemental '

-..-.:..~ e% ~ - . . "..-. Gas phase ~ . "

'leres

9 o eeee

ace sulfate

organics

Surface-absorbeu /

adsorbed organics

F I G U R E 10.1 Schematic of a diesel soot particle consisting of an agglomeration of elemental carbon spheres (0.01- to 0.08-/xm diameter). Its surface is covered with absorbed/adsorbed particle-phase organics, including 5-ring (e.g., BaP) and 6-ring PAHs. Gas-phase organics include all of the highly volatile 2-ring PAHs (e.g., naphthalene and methylnapthalenes). Semivolatile 3-ring (e.g., phenanthrene and anthracene) and 4-ring PAHs (e.g., pyrene (II) and fluoranthene (V)) are distributed between both phases. Sulfate is also associated with diesel particles. (Adapted with permission from Johnson et al., 1994, SAE Paper 940233 9 940233 Society of Automotive Engineers, Inc.; see also Schauer et al., 1999.)

significant cause of visibility degradation in polluted atmospheres (e.g., see Larson et al., 1989). Polycyclic aromatic hydrocarbons and their associated more polar (hence more water soluble) heteroatom derivatives (polycyclic aromatic compounds, PACs; see later) not only are present in air as gases and particles but also are present, for example, in urban "street dust" samples (Takada et al., 1990) and other soil and water environments throughout the world due to wet and dry deposition (see, for example, Schwarzenbach et al. (1993), Wild and Jones (1995), and Neilson (1998)). Furthermore, through long-range transport of polluted air masses, PAHs can be found in ambient air at receptor sites far from their original sources. For example, PAHs from the European continent and Great Britain have been observed at a "background site" at Birkenes, Norway, and in Copenhagen, Denmark (Bjarseth et al., 1979; Bj6rseth and Olufsen, 1983; Nielsen et al., 1999a, 1999b). PAHs from the former Soviet Union have been reported to reach the Norwegian Arctic (Pacgna and Oehme, 1988) and sources in Africa and in northern Europe (e.g., Germany, Belgium, and France) have been reported to impact Corsica, France (Masclet et al., 1988). Of particular interest in terms of atmospheric chemistry are reactions of certain PAHs in VOC-NO~-air atmospheres to form biologically active polycyclic aromatic compounds, PACs. Thus, not only is the fundamental chemistry of the formation and fates of these secondary air pollutants of interest, but it can also have major toxicological implications. For example, in some airsheds certain PACs that are reaction products (e.g., nitro-PAH and nitro-PAH lactones) contribute signifi-

440


cantly more to the overall, direct-acting bacterial mutagenicities of the gaseous and particle phases of ambient air than do the PACs in the primary emissions directly emitted by sources. Furthermore, certain of them, e.g., 2-nitrofluoranthene (XXVII), are human cell mutagens: NO 2 1

(

"~

2

3

//"

4

2-Nitrofluoranthene

XXVII Nitro-PAHs, regardless of whether directly emitted or formed in air, are of concern because many of them are animal, and possible human, carcinogens (IARC, 1989; see also review of environmental carcinogens by Tokiwa et al. (1998) and references therein). The goal of this chapter is to illustrate the atmospheric chemistry of PAHs and PACs. However, because of their unique biological properties, we also provide some perspective on their relevance to air pollution toxicology and the development of sound scientific health risk assessments for specific carcinogenic PAHs such as benzo[a]pyrene (California Air Resources Board, 1994) and for complex combustiongenerated emissions such as diesel exhaust (e.g., IARC, 1989; WHO, 1996; and California Air Resources Board, 1998). For discussions of such broad topics as research on analytical methods for the determination of PAHs in environmental samples, see, for example, the review of chromatographic methods by Poster, Sander, and Wise (1998) and references therein. Similarly, for reviews of their environmental chemistry and related carcinogenicities, see Neilson (1998) and Harvey (1997). The book Enuironmental Organic Chemistry by Schwarzenbach, Gschwend, and Imboden (1993)provides a useful perspective on the subject and contains helpful specific examples relating to PAHs.

A. N O M E N C L A T U R E A N D SELECTED PHYSICAL A N D SPECTROSCOPIC PROPERTIES OF POLYCYCLIC AROMATIC H Y D R O C A R B O N S (PAHs) A N D POLYCYCLIC AROMATIC C O M P O U N D S (PACs) 1. Combustion-Generated PAHs and PACs Historical, chemical, and toxicological interest in PAHs and PACs goes back over two centuries when Sir

Percival Pott proposed that the high rate of cancer of the scrotum incurred by London's chimney sweeps was due to the presence of certain chemicals in the fireplace soot (i.e., POM) to which they were heavily exposed (Pott, 1775). Some 150 years later, Passey (1925) reported that organic extracts of such "domestic soot" induced tumors in experimental animals. In the early 1940s, Leiter et al. (1942) demonstrated that a similar phenomenon occurred with organic extracts of ambient air particles--that is, injection of "tars extracted from atmospheric dusts" collected at locations throughout New York City produced subcutaneous sarcomas in mice. Shortly thereafter, Leiter and Shear (1943) reported that "marginal doses of 3,4benzpyrene" (known today as benzo[a]pyrene, BaP, I), the powerful carcinogen earlier isolated from coal tar and synthesized by Cook et al. (1933), also produced subcutaneous tumors in mice. These observations, coupled with the discoveries of BaP in chimney soot (Goulden and Tipler, 1949) and in ambient air particles collected at 10 stations throughout Great Britain (Waller, 1952), and the tumorigenic/carcinogenic properties of extracts of ambient particles collected during episodes of Los Angeles photochemical smog (Kotin et al., 1954) were key factors in establishing the atmospheric chemistry of PAHs and PACs as a new field in air pollution research. Since then, many monographs, handbooks, symposia proceedings, and specialized chapters (in addition to thousands of research papers) dealing with the chemistry and biological and toxicological aspects of PAHs, POM, and PACs have appeared in the scientific, engineering, and medical sciences literature. Examples are cited in Box 10.1.

2. Structures and IUPAC Rules for Nomenclature a. PAlls

Over the decades several significantly different PAH and PAC numbering/nomenclature systems have been proposed and widely used in the older literature, e.g., that of Clar (1964). Unfortunately, even today this can lead to confusion on the part of those unfamiliar with the history of different systems of nomenclature. We follow the 1979 IUPAC recommendations summarized in Polynuclear Aromatic Hydrocarbons: Nomenclature Guide (Loening et al. 1990). The American Chemical Society also publishes the Ring Systems Handbook, which, ca. 1990, contained structural diagrams for over 70,000 unique ring systems (American Chemical Society, 1977 to present). A detailed discussion of these rules and nomenclature is beyond the scope of this book. However, we

A. N O M E N C L A T U R E A N D S E L E C T E D P H Y S I C A L A N D S P E C T R O S C O P I C P R O P E R T I E S O F P A H s A N D P A C s

441

BOX 10.1 SELECTED REFERENCES TO THE CHEMISTRY TOXICOLOGY OF AIRBORNE PAHs AND Review/evaluations by government agencies: U.S. National Research Council/National Academy of Sciences (NRC, 1972, 1983, 1988); International Agency for Research on Cancer monographs (IARC, 1983, 1987, 1989); World Health Organization; Environmental Health Criteria (EHC) monographs (e.g., Diesel Fuel and Exhaust Emissions, EHC 171, WHO, 1996); U.S. Department of Health and Human Services, Public Health Services and National Toxicology Program, the Reports on Carcinogens (8th Report, U.S. DHHS, 1998); the California Air Resources Board (1994, 1997, 1998); the California EPA (1997); and the Danish Environmental Protection Agency (e.g., see Nielsen et al., 1994, 1997, 1998). Other publications include the following: handbooks and collections of chapters on PAHs and

outline them in Box 10.2 and give examples in Tables 10.1 and 10.2 of how they are generally employed with respect to atmospherically relevant PAHs. The nomenclature for selected N-, O-, and S-atom-functionalized PACs is described briefly in Section A.2.b and examples are found in Tables 10.3-10.5. The same molecule may appear with different orientations. For example, in the current literature, the structure for the 3-ring PAH phenanthrene is drawn as (IVa), (IVb), or (IVc). While one or more IUPAC rules may be broken, their ring numberings are internally consistent. 10

1 9

9

PACs (e.g., Bj0rseth, 1983; Grimmer, 1983b; Bj0rseth and Ramdahl, 1985; White, 1985; Howard et al., 1990; Neilson, 1998); the series PolynuclearAromatic Hydrocarbons published by Battelle Press (e.g., see Cooke et al., 1982); monographs on polycyclic aromatic hydrocarbons (e.g., Mackay et al. 1992; and Harvey, 1997); chapters dealing with PAHs and PACs in books on atmospheric chemistry and environmental organic chemistry (e.g., in Finlayson-Pitts and Pitts (1986), Schwarzenbach et al. (1993), and Graedel, Hawkins, and Claxton (1986)); review articles on atmospheric reactions and fates of PAHs by Nielsen et al. (1983), Nielsen (1984), Nikolaou et al. (1984), Van Cauwenberghe (1985), Pitts (1987, 1993a, 1993b), Baek et al. (1991), Atkinson and Arey (1994), and Arey (1998a).

Some PAHs may have two hydrogens at a specific carbon position, the indicated hydrogens. Such an indicated hydrogen should be mentioned by carbon number, even when it is further oxidized to a carbonyl group; consider, for example, 7H-benz[ de ]anthracene (XlI) and 7H-benz[de]anthracen-7-one (XlII, benzanthrone).

H / ~H

10

2

XII 8

~"

4

3

7

7

5

Ira

4

IVb 3

2 1

~-%/,- -,,,,~ 10 8

9

Ire

0 XIII

2 6

r

AND PACs

3

Similarly, the structure of a nitro-PAH lactone found in ambient aerosols and also formed in laboratory irradiations of phenathrene-NOx-air mixtures, 2-nitro6H-dibenzo[b,d]pyran-6-one, which is a powerful, direct-acting bacterial mutagen and potent human cell mutagen, is commonly shown as (XI) (Helmig et al., 1992a; Arey et al., 1992; Sasaki et al., 1995, 1997b; Arey, personal communication). Finally, a possible source of confusion when assigning structures to a given PAH is the "correct" number of the double bonds within the fused rings. The Fries Rule states that "the most stable form of a polynuclear

442

10. AIRBORNE PAHs: CHEMISTRY AND TOXICOLOGICAL IMPLICATIONS

B O X 10.2 SELECTED

IUPAC RULES AND NUMBERING

NAMING, SELECTED

One starts with the IUPAC "preferred name" (sometimes referred to as the "trivial name")used for certain of the simple PAHs, e.g., anthracene, phenanthrene, pyrene, and fluoranthene. More complicated structures are then built up by the addition of, for example, benzo, dibenzo, or naphtho groups to the skeleton of the simple PAH. The sequence of steps to follow when drawing a structure of a fused PAH ring system is the following: 1. Define the molecular structure for which a preferred name exists in the IUPAC list (1979). 2. Place the maximum number of rings in a horizontal row. 3. Place the greatest number of remaining rings above and to the right of a vertical axis drawn through the center of the molecule (i.e., in the upper right-hand quadrant). Structure I shown earlier for benzo[a]pyrene illustrates these points. 4. After properly orientating it, number the carbon atoms in the molecule in a clockwise direction starting with the carbon atom n o t common to another ring (i.e., not involved with fusion of two rings) that is in the most counter-clockwise position of the uppermost ring that is farthest to the right (see structures for PAHs I-V and Table 10.1). Carbon atoms c o m m o n to two or more rings are n o t numbered. 5. Letter the faces of the ring in alphabetical order. 6. Starting with "a," letter the faces of the rings in alphabetical order beginning with that side between carbon atoms 1 and 2. Continue clockwise around the molecule, noting, however, that ring faces common to two rings are n o t lettered. See fluoranthene, V, and benzo[k]fluoranthene, XXVI. 2

2

10

11

9 k 8

4 e 5 7

6

12

10

4

9

5 8

7

V

XXVI

Fluoranthene

Benzo[k ]fluoranthene

6

ORIENTING, PAHs

FOR

Since addition of another ring is defined by the bond(s) involved in the fusion, in this case the k face of fluoranthene, the name becomes benzo[k]fluoranthene. The numbering in the final molecule can be different from the one used in the original compound, since the addition of more rings may change its orientation in the frame of reference. For example, consider anthracene (VIII), 9,10-dimethylanthracene (IX), and 7,12-dimethylbenz[a]anthracene (X):

8

7

9

~

2

9

1

7

6

3 5

CH3

8

1

10

6

4

CH3 VIII

IX 2

C H ~ 10

"~

8

71

3 4

6

CH3 X

Note that the numbering of anthracene and 9,10dimethylanthracene does not conform to the foregoing IUPAC rules. This "common" numbering is derived from anthraquinone, in which the middle, or m e s o , positions are numbered 9 and 10.

443

A. NOMENCLATURE AND SELECTED PHYSICAL AND SPECTROSCOPIC PROPERTIES OF PAHs AND PACs TABLE 10.2

Structures, C o m m o n N a m e s , Empirical Formulas, Molecular W e i g h t s , Melting Points, Boiling Points, and C A S N u m b e r s for Selected P A H s of A t m o s p h e r i c Interest a'b Empirical formula

MW (g m o l - 1 )

Mp (oc)

Bp (oc)

CAS number

1-Methylnaphthalene

Cll H 10

142.2

- 22

244.6

90-12-0

2- Methylnaphthalene

Cll H 10

142.2

34.6

241.9

91-57-6

Biphenyl

C 12 H 10

154.21

71

256

92-52-4

1-Methylphenanthrene

C15 H 12

192.26

123

359

832-69-6

Cyclopenta[ cd ]pyrene

Cls H 10

226.28

175

-

27208-37-3

Benzo[ghi ]fluoranthene

Cls H 10

226.28

149

432 a

203-12-3

Triphenylene

C is H 12

228.3

199

438

217-59-4

Benzo[ e ]pyrene

C20H 12

252.3

178

493 b

192-97-2

Benzo[ j ]fluoranthene

C20H 12

252.3

166

-~480 b

205-82-3

Common name

Structure CH 3 ,

I

6

3 5

4

8 1 7C H~ 2 6 5

4

10 11 9

~ 8

2

3

4 7

6

5

CH 3

lo

I

3 8 7 6

10

1

7

4 6

5

10

1

7

4 6

5 2

14

1 ~/"

11 10~

56

7 2

1 //-.

12

3

11

4

10

5

7

11

2

9 4 5 7

6

(continues)

444


TABLE 10.2 (continued) Common name

Structure

Empirical formula

MW (g m o l - 1 )

Mp (oc)

Bp (oc)

CAS number

2 1

3

11

4

10

5

Perylene

C20 H 12

252.3

277

495

198-55-0

Coronene

C24 H 12

300.36

> 350

525

191-07-1

Dibenzo[ a,e ]pyrene

C24 H 14

302.28

234

-

192-65-4

9 ~'~,,,/7 8 12

1

3 4

7

6 14

12

~0 ~

1

~

s

3

l] ) 4 7 ,,,,,,~

5

6

a Adapted from Mackay et al. (1992) and Mackay and Callcott (1998). Structures and numbering are based on IUPAC recommendations; see Loening et al. (1990). b From Bjorseth (1983).

hydrocarbon is that in which the maximum number of rings have the benzenoid arrangement of three double bonds" (Fieser and Fieser, 1956). This statement is consistent with a subsequent proposal by Clar (1972) that for such polyarenes one should maximize the number of "aromatic sextets." NO 2

(W

O XI

2-Nitro- 6 H- dibenzo[ b,d ]pyran- 6-one

For example, of the three possible electronic structures for naphthalene, the symmetrical structure (a) with two benzenoid rings (aromatic sextets)would be the more stable, and hence the preferable representation.

a

b

c

[Note that when the Clar convention is used for Kekul~ structures, the benzenoid rings are often shown as circles (e.g., see Moyano and Paniagua, 1991).] Finally, since our focus is on the atmospheric chemistry of unsubstituted (or methyl substituted) PAHs, we do not discuss major classes of PAHs from natural sources that are not found at significant levels in ambient air but may be important in soil and water environments, e.g., alkylated PAHs (see the review by Simoneit, 1998). b. PACs

Despite the generally lower emission factors and ambient levels of PACs compared to PAHs (there are some exceptions), interest in certain airborne N-, O-, and S-atom PACs has increased significantly, in large part because of their mutagenic and carcinogenic properties. For example, in 1976, Cautreels and Van Cauwenberghe used GC-MS to identify more than 100 compounds in three different fractions (neutral, acidic,

445

A. NOMENCLATURE AND SELECTEDPHYSICAL AND SPECTROSCOPIC PROPERTIESOF PAHs AND PACs TABLE 10.3

Structures, Common Names, Empirical Formulas, and Molecular Weights of Selected Nitrogen-Containing Mono- and Polycyclic Aromatic Compounds (N-PACs) Identified in Ambient Air Empirical formula

MW (g mol- 1 )

Pyridine

C 5H 5N

79

Quinoline

C9H7N

129

a, b, c

Isoquinoline

C9H7N

129

b, c

Benzo[ f ]quinoline (5,6-Benzoquinoline)

C13H9N

179

a, b, c

4

Benzo[ h ]quinoline (7,8-Benzoquinoline)

C13H9N

179

a, b, c

4

Phenanthridine

C13H9N

179

a, b, c

Acridine*

C13H9N

179

b, c, d

Benz[a]acridine (1,2-Benzacridine)

C17 HllN

229

d

Benzo[c ]acridine (3,4- Benzacri dine)

C17 HllN

229

a

Common name

Structure

Reference

1

5

3 4 8

1

5

4

8

1

6

3

7~ ~ / N

2

6

3 5

4 2

10 9

8

1 ~ 3 ~

N

7

4

6 2

1N ~ 3 9 8 7

6 2

10

1

9 8 7

6

8

9

1

5

10

4

6

3

2

4 10 9 8

7

6 2

11

12

1 [~/

N~'

109 8

7

"~

3

"4

6

(continues)

446

10oAIRBORNEPAHs: CHEMISTRYAND TOXICOLOGICAL IMPLICATIONS TABLE 10.3 (continued)

Structure

Empirical formula

MW (g m o l - 1 )

Reference


C21H13N

279

a,d

Dibenz[ a,j ]acridine

C21H13N

279

10-Azabenzo[ a ]pyrene

C19HllN

253

C19H1002N2

298

3-Nitrobiphenyl

C12H9NO2

199

2-Nitronaphthalene

C10H7NO2

173

f,g,h,i,j,l,

rn

9-Nitroanthracene

C14H9NO2

223

f,g,h,i,k,l,

rn

2-Nitrofluoranthene

C16H9NO2

247

f,g,h,i,k,l,

rn

1-Nitropyrene

C16H9NO2

247

f,g,h,i,k,l,

rn

2-Nitropyrene

C16H9NO2

247

Common name 2

4

12

lO 9 12

2

11

3

lO

4 9 8

7

6

12 lO

1

11

2

8 7

6

5 12

1

3-Nitro-6-azabenzo[ a ]pyrene 7

6

5

NO2 ~

4 6

5

1

~

f,g,h,i,j,l

NO 2

4

NO 2

1

3

NO2 2 "~ 3~4

NO2 "~

1 2

( ~

3 a

23NO2 g,l

(continues)

447

A. NOMENCLATURE AND SELECTED PHYSICAL AND SPECTROSCOPIC PROPERTIES OF PAHs AND PACs T A B L E 10.3 (continued)

Structure

Common name

Empirical formula

MW (g m o l - 1 )

Reference

1,6-Dinitropyrene

C16H8N20 4

292

h,/77

Pyrrole

C4HsN

63

Indole

CsH7N

117

Carbazole*

C12 H9N

167

NO 2

]

(

4

3

5

NO2 H1 N 5(

~2

4

3 7

U 1

6

2

5

3 4 8

H 9

1

7 6

2 3 5

4

Note: Exceptions to I U P A C rules are designated with asterisks. a Nielsen et al. (1986).

b Chen and Preston (1998). c Dong et al. (1977). d Yamauchi and Handa (1987). e Sera et al. (1994). f Wilson et al. (1995). g Atkinson and Arey (1994) and references therein. h Lewtas and Nishioka (1990) and references therein. i Zielinska et al. (1990) and references therein. J Fraser et al. (1998). k Hannigan et al. (1998). 1Arey (1998a) and references therein. m Tokiwa and Ohnishi (1986), Tokiwa et al. (1987), and references therein.

and basic) of extracts of ambient aerosols; the compounds included not only PAHs and aliphatics but also polar oxygenated substances, including several O-PACs and N-PACs. Similarly, Gundel and co-workers (1993) identified O-PACs and organic nitrates and nitrites as well as amides, amines, and nitro compounds in polar fractions of extracts of airborne particulate matter from particles (Ds0 < 15 /xm) in ambient air in Elizabeth, New Jersey, and in the U.S. National Institute of Standards and Technology (NIST) Standard Reference Material SRM 1649, "Urban Dust/Organics" (see Box 10.3). Other investigators have identified S-PACs (e.g., dibenzothiophene) in the gas and particle phases of both direct emissions and ambient air. Hence, some understanding of the nomenclature used for O-, N-, and S-PACs is also useful. Given the presence of N-, O-, and S-heteroatoms as

members of fused PAH ring systems (endocyclic), attached to such fused PAH rings (exocyclic), or both, the nomenclature of PACs is more complex than for PAHs. However, Chapter 2 of Polynuclear Aromatic Hydrocarbons--Nomenclature Guide by Loening et al. (1990) and Systematic Nomenclature of the Nitrogen,

Oxygen, and Sulfur Functional Polycyclic Aromatic Compounds (Later et al., 1990) provide a useful synopsis of the IUPAC rules for monofunctional N-, O-, and SPACs (IUPAC, 1979). Many of the IUPAC rules for the parent, nonsubstituted PAHs described in Box 10.2 are applicable to these heteroatom derivatives. These rules are summarized in Box 10.4 (N-PAC), Box 10.5 (O-PAC), and Box 10.6 (S-PAC), respectively. As an example, both replacement (aza) and substitutive (-NO 2) nomenclature are used in naming the powerful bacterial mutagen 3-nitro-6-azabenzo-

448

10. AIRBORNE PAHs: CHEMISTRYAND TOXICOLOGICAL IMPLICATIONS TABLE 10.4

Structures, Common Names, Empirical Formulas, and Molecular Weights for Selected Oxygen-Containing Polycyclic Aromatic Compounds (O-PACs) Identified in Ambient Air Empirical

MW

formula

(g mol- 1)

Reference

1-Acenaphthenone

C 12H s O

168

a

Fluoren-9-one (9H-Fluoren-9-one)

C13HsO

180

a, b, d, f, g

Benzanthrone (7H-Benz[ de ]anthracen-7-one)

C17H100

230

a, b, f, g

11H-Benzo[b]fluoren-11-one

C17H100

230

a, b, g

6 H-Benzo[ cd ]pyren- 6-one

C19H100

254

a, f, g

9,10-Anthracenedione (Anthraquinone)

C14H802

208

b, d, f

9,10-Phenanthrenedione (Phenanthraquinone)

C14 HsO 2

208

a

7,12-Benz[ a ]anthracenedione (Benz[ a ]anthracene-7,12-quinone)

C18H100 2

258

b, c, d, f, g

5,12-Naphthacenedione ( 5,12-Naphthacenequinone )

C18H100 2

258

a, b, g

C o m m o n name

Structure

Ketones O

6

3 5

4

O

6

3 5

4 2 1

3

4 5

O O

7

3 6

5

4 1

2

3 4

O Quinones O

6

3

O O

6

O

5

4

3 2

O

1 ~

3 4

10 9 8

[[7

6

O O i

i

9 3

8

O

(continues)

449


TABLE 10.4 (continued) Empirical formula

MW (g m o l - 1 )

Reference

C12H603

198

a, b, c, f, g

C13H7NO 4

241

e

Furan

C4H40

68

Dibenzofuran

C 12 H s O

168

Common name

Structure Anhydrides O-

O

O

7

Naphthalene-l,8-dicarboxylic anhydride ( 1 , 8 - N a p h t h a l i c anhydride)

2

6

3 5

4

Lactones NO2

1

2 3

10

4

9(

2-Nitro- 6 H- dibenzo[ b,d ]pyran- 6-0ne

8

o

Endocyclic O - P A C s 5

2

4

3 9

1 2

7

3 6

5

d

4 e H e l m i g et al. (1992a). f F r a s e r et al. (1998). g H a n n i g a n et al. (1998).

a A l l e n et al. (1997). b K6nig et al. (1983). c Wilson et al. (1995). d Ligocki and Pankow (1989).

[a]pyrene (XXXI) found in diesel exhaust and ambient air (e.g., see Section C and Sera et al. (1994)). 12

1

11

3-nitrobenzanthrone (3-nitro-7H-benz[de]anthracen-7one, XXlV)was recently identified by Enya et al. (1997) in both diesel exhaust and ambient air; it is an extremely powerful, direct bacterial mutagen:

2 0 2

NO 2

9 8 7

6

5

XXXI 3-Nitro-6-azabenzo[a ]pyrene

O XXIV

Structures of several N-PACs identified in ambient air (and references) are shown in Table 10.3. As an example, we saw earlier, replacing the CH 2 group in the parent PAH 7H-benz[ de ]anthracene, XlI, by a keto group, C = O , gives 7H-benz[de]anthracen-7one, XlII (benzanthrone). It is well known as a major O-PAC in combustion-generated emissions and ambient air (see Table 10.4) and the 3-nitro derivative,

3-Nitro-7H-benz[ de ]anthracen-7-one

or 3-Nitrobenzanthrone

Structures of compounds representing several classes of O-PACs identified in ambient air, e.g., PAH ketones, quinones, anhydrides, and lactones, are shown in Table 10.4.

450

10. AIRBORNE PAHs: CHEMISTRYAND TOXICOLOGICAL IMPLICATIONS

TABLE 10.5

Structures, Common Names, Empirical Formulas, and Molecular Weights for Selected Sulfur-Containing Mono- and Polycyclic Aromatic Compounds (S-PACs) in Ambient A i r a'e'f and Combustion-Generated Emissions, c'a Including Cigarette Smokeb

Structure S 1

5r ~2 4

Empirical formula

Common name

Thiophene

C4H48

MW (g m o l - 1)

Reference

84

3 7 S 1

2 5

Benzo[b]thiophene

CsH6 S

134

Dibenzothiophene

C12

H sS

184

Benzonaphthothiophene

C16 H10 S

234

c

3-Nitrobenzothiophene

CsHsSNO 2

179

d

b

3 4

(7s 6

7 8

4 3

a,b,c,e,f

2

9

1

4

6

7

3

8

2

9 1

11

10

4

5 ~ 6

~/NO2 S"

7

2

1

a Ligocki and Pankow (1989); ambient air. b Schmid et al. (1985); cigarette smoke condensate. c Tong and Karasek (1984); diesel particulate matter (DPM). d Schuetzle and Daisey (1990) and references therein. e Fraser et al. (1998); ambient air. fAtkinson et al. (1988a); ambient air.

B O X 10.3 SOURCE

OF

NIST SRM

STANDARD 1649, "AIR

Throughout this chapter, we cite examples of the use of the NIST Standard Reference Material SRM 1649, which is referred to as "Air Particles" or "Urban Air Particulate Matter," (a) to validate analytical procedures for determination of PAHs and PACs in samples of complex mixtures of particulate matter in ambient air and (b) for laboratory intercomparisons of methodologies for bacterial bioassays and bioassay-directed fractionations of organic extracts of such mixtures (e.g., see Claxton et al., 1992a; Lewtas et al., 1990a, 1992; and May et al., 1992).

REFERENCE PARTICLES"

MATERIAL

SRM 1649 was originally collected in a "bag house" over a period of one year in the late 1970s at a site near Washington, DC, screened through a 200-mesh sieve (cutoff point 125 /xm), and since then apparently stored in bottles at room temperature (Lewtas et al., 1990a). Understandably, Claxton et al. (1992a)caution "While the air particles in SRM 1649 are similar to other air particulate sampies, they do not represent a typical air particulate sample as collected by most researchers--they are intended as reference materials and not as samples for assessment of levels of airborne toxicants."

451

A. N O M E N C L A T U R E A N D SELECTED P H Y S I C A L A N D S P E C T R O S C O P I C P R O P E R T I E S OF P A H s A N D P A C s

B O X 10.4 SELECTED N-ATOM

IUPAC NOMENCLATURE POLYCYCLIC AROMATIC

Endocyclic PACs. There are two approaches to naming endocyclic N-PACs, e.g., dibenz[a,h]acridine, VII (see Later, 1985; and Later et al., 1990). In one style, the ending ine is used for endocyclic 6-membered rings, e.g., pyridine and quinoline. As with PAHs, if a benzo group is fused to a given side, the letter for that side is placed in brackets; e.g., addition of a benzo group to the h side of quinoline forms benzo[h]quinoline (note letters and numbers within square brackets are italicized, as with PAHs). Alternatively, it is named relative to the carbon atoms involved in the benzo-ring fusion (that constitute the ends of side h), i.e., 7,8-benzoquinoline:

@

8

1

+

Benzo

> 6

~

3

5 4 Quinoline

3. Solubilities and Vapor Pressures PAHs have low solubilities in water (Table 10.6) as expected from their nonpolar character. These decrease dramatically in going from the 2- and 3-ring compounds (e.g., naphthalene, with a solubility of 31 mg L -1) through the 3- and 4-ring semivolatile organics (e.g., fluoranthene, with a solubility of 0.26 mg L -1) to 5-ring BaP, with a solubility of only 0.0038 mg L -~ (Mackay et al., 1992). However, reactions of PAHs in ambient air to form more polar species (e.g., nitro-PAHs, ketones, quinones, lactones, and dicarboxylic acids) greatly enhance their solubilities in aqueous systems. This has major implications when one considers the distribution of PAHs, and their atmospherically formed PAC derivatives, through the air, water, and soil environments. These increases in solubility upon reaction are important not only from an environmental chemistry perspective but also in terms of possible impacts on public health and ecosystems, e.g., in both the exposure and the health effect

RULES FOR COMPOUNDS 9

2 1

8

10

1

3

7

4

5

2

9

3

8

4

7

6

B e n z o [ h ]q u i n o l i n e 7,8- B e n z o q u i n o l i n e

Similarly, benzo[f]quinoline - 5,6-benzoquinoline. The second approach for endocyclic N-PACs is to use replacement nomenclature, in which the term aza is a prefix to the corresponding PAH. The position of the N-atom in the fused-ring system of the PAH precedes the term aza. Thus, for example, quinoline is 1-azanaphthalene. We generally use this convention for three or more ring N-PACs with endocyclic nitrogen heteroatoms. Substitutive nomenclature is employed for exocyclic N-PACs (and the O- and S-PACs). Functional groups such as amino (-NH2), cyano (- C = N ) , nitro (-NO2) , and hydroxyl ( - O H ) are attached as prefixes to the parent PAH, e.g., 1-nitropyrene, VI.

portions of risk assessments of PAHs (e.g., see Mackay et al., 1992; Schwarzenbach et al., 1993; Neilsen, 1998; Baum, 1998; the review by Mackay and Callcott, 1998; and references therein). Values recommended by Mackay et al. (1992) and Mackay and Callcott (1998) for the vapor pressures of a number of PAHs of atmospheric interest and several properties used in theoretical calculations of their gas-particle partitioning in ambient air are given in Table 10.6: for examples of their use, see discussion of gas-particle partitioning theory in Chapter 9.D and references cited above. Measured vapor densities of several volatile and semivolatile PAHs over a temperature range from 10 to 50~ have been used by Sonnefeld and co-workers (1983) to generate vapor pressures as a function of temperature: log P~ = - A / T

+ B

Values for A and B are given in Table 10.7.

452


B O X 10.5 SELECTED O.ATOM


The naming of endocyclic O-atom heterocycles is based on the same principles as for N-PACs; e.g., the structures of furan (XVlI) and dibenzofuran (XVIII) are /01

8< 2 9

)2

1

7

3

5

XVIII

Furan

Dibenzofuran

Similarly, with exocyclic O-PACs, substitutive nomenclature and common suffixes are used to describe the O-atom compounds formed when an Hatom on a PAH is replaced by another atom or functional group, e.g., 6-hydroxybenzo[a]pyrene, XlX, or 2-methoxychrysene, XX: 12

9(s

1

I

OH

XXI 2-Hydroxy- 1-nitropyrene

For ketones in which a CH 2 group is replaced by a C = O , one adds the suffix o n e . For example, 9Hfluorene (XXII) becomes 9H-fluoren-9-one or 9fluorenone (XXIII).

2

11

,%

O

3

8

H2

1

7

8 6

7

5

6

OH

5

4

8 2

7

3

6

II

1 2 3

5

4

XIX

XXII

XXIII

6-Hydroxybenzo[ a ]pyrene

9H-Fluorene

9H-Fluoren-9-one

or Fluorene

or 9-Fluorenone

12

1

11

8 7

NO 2

4

XVII

10

When PAHs are substituted with two exocyclic functional groups, e.g., the PAC XXI, one may employ either a prefix and a suffix, e.g., 1-nitropyren-2ol, or two prefixes, 2-hydroxy-l-nitropyrene (we use the latter):

3 6

10

RULES FOR COMPOUNDS

6

XX 2-Methoxychrysene

OCH 3

A. N O M E N C L A T U R E A N D SELECTED P H Y S I C A L A N D S P E C T R O S C O P I C PROPERTIES OF P A H s A N D PACs

453

BOX 10.6 SELECTED


The IUPAC nomenclature of S-atom PACs is also similar to that for N-PACs. Thus the "trivial" name thiophene (XlV) is a basis for the common names of S-heterocycles containing one endocyclic S-atom in a 5-membered fused ring. Addition of two benzo groups gives dibenzothiophene (XV): 1

9

5

2

4

3

structure, XVI (Later et al., 1990): 9

10

1

6

S 5

4

7

XVI Thianthrene

7 6

XIV Thiophene

1

S.ATOM

RULES FOR COMPOUNDS

S 5

4

XV Dibenzothiophene

Names of S-heterocycles with two S-atoms in a 6-membered ring are based on the thianthrene

4. Gas- Particle Partitioning, Sampling Techniques, and Ambient Levels of Selected PAHs and PACs Whether or not a given PAH exists virtually entirely in the gas phase or in the particle phase, or is partitioned between them, is a critical factor in determining its physical and chemical fates in ambient air and in subsequent intra- and intermedia transport through our air/water/soil environments. This is true not only for physical processes such as wet and dry deposition but also for their chemical reactivity, lifetimes, and fates in V O C - N O X systems characteristic of polluted airsheds. For example, the homogeneous gas-phase reactions of pyrene and fluoranthene differ dramatically from the rates, mechanisms, and products of their particle-associated heterogeneous reactions (Sections E and F). Similarly, the toxicological effects may depend on whether the compound is inhaled as a gas, inhaled as a particle, or adsorbed/absorbed on the surfaces of coexisting particles. The distribution between these forms varies with the particular compound as well as with a number of parameters such as temperature. For example, in hot weather in ambient air in Claremont and Riverside, California, up to 25% of the PAC 2-nitro6H-dibenzo[b,d]pyran-6-one (structure XI; vide supra)

Examples of common names and structures of several S-PACs and nitro-S-PACs identified in combustion-generated emissions (including tobacco smoke condensate) and ambient air are given in Table 10.5.

was in the gas phase, rather than being adsorbed/absorbed on surfaces (Arey et al., 1994). This nitrophenanthrene lactone, formed in a gas-phase atmospheric reaction of phenanthrene, is widely distributed in urban ambient air and is a powerful bacterial mutagen (Helmig et al., 1992a, 1992b; Arey et al., 1992), as well as being mutagenic in the MCL-5 human cell assay (Sasaki et al., 1997b). Gas-particle partitioning of semivolatile organics, including PAHs, is discussed in more detail in Chapter 9.D. The reader is also referred to books by Mackay et al. (1992), Schwarzenbach et al. (1993), and Baum (1998) and articles by Pupp et al. (1974), Junge (1977), Cautreels and Van Cauwenberghe (1978), Yamasaki et al. (1982), Pankow (1987, 1991, 1994a, 1994b, 1998), Pankow et al. (1993, 1994), Bidleman and Foreman (1987), Bidleman (1988), Ligocki and Pankow (1989), Foreman and Bidleman (1990), Pankow and Bidleman (1992), Storey et al. (1995), Wilson et al. (1995), Liang and Pankow (1996), Jang et al. (1997b), Finizio et al. (1997a, 1997b), Liang et al. (1997), Chen and Preston (1997), Goss (1997), Goss and Schwarzenbach (1998), Harner and Bidleman (1998), Jang and Kamens (1998), Mackay and Callcott (1998), and Feilberg et al. (1999a). We now consider sampling techniques and typical ambient concentrations of PAHs and PACs.

454


TABLE 10.6 V a p o r Pressures, Subcooled Liquid V a p o r P r e s s u r e s ( p ~ ) , Solubilities in W a t e r (S, c S ) , L o g a r i t h m of O c t a n o l / W a t e r P a r t i t i o n C o e f f i c i e n t (log K o w ) , C a l c u l a t e d H e n r y ' s L a w C o n s t a n t ( H ) , a n d C a l c u l a t e d L o g a r i t h m of t h e O c t a n o l - A i r P a r t i t i o n C o e f f i c i e n t (log K O A ) at 2 5 ~ f o r S e l e c t e d P A H s of A t m o s p h e r i c I n t e r e s t a'e

Vapor pressure P

PAH b

P~L

Pa

Torr

Naphthalene 10.4 1-Methylnaphthalene 8.84 2-Methylnaphthalene 9 Biphenyl 1.3 Acenaphthene 3 • 10 -1 Acenaphthylene 9 • 10 -1 Fluorene 9 • 10 .2 Phenanthrene 2 • 10 .2 Anthracene 1 • 10 .3 Pyrene 6.0 • 10 .4 Fluoranthene 1.23 • 10 .3 Chrysene 5.7 • 10 . 7 Triphenylene 2.3 • 10 .6 B[a]A 2.80 • 10 -5 Benzo[a]pyrene 7.0 • 10 .7 Benzo[e]pyrene 7.4 • 10 .7 Perylene 1.4 • 10 - s

7.8 6.63 7 9.8 2 6.8 7 1.5 8 4.5 9.2 4.3 1.7 2.1 5.3 5.6 1.1

B[b IF

.

.

B[k]F

5.2 • 10 - s 3.7 • 10 - l ~ 2.0 • 10 - l ~

B[ghi]P DB[a,h]A Coronene

• • • • • • • • • • • • • • • • •

Pa

10-2 10 .2 10 .2 10 -3 10-3 10-3 10 . 4 10 .4 10 . 6 10 . 6 10 . 6 10 . 9 10 - s 0 .7 10 . 9 10 . 9 10-10

. 3.9 • 10-1~ 2.8 • 10 -12 1.5 • 10 -12

36.8 8.8 11.2 3.7 1.5 4.1 7.2 • 1.1 • 7.78 • 1.19 • 8.72 • 1.07 • 1.21 • 6.06 • 2.13 • 2.41 • _

Calculated

Solubility

Torr

10 -1 10 -1 10 . 2 10-2 10 . 3 10 . 4 10 . 4 10 . 4 10 -5 10 -5

2.76 6.6 8.4 2.8 1.1 3.1 5.4 8 5.84 8.93 6.55 8.03 9.1 4.55 1.60 1.81

• • • • • • • • • • • • • • • • -

10-1 10 . 2 10 . 2 10 . 2 10-3 10 . 2 10 . 3 10 .4 10 . 4 10 -5 10 -5 10 . 7 10 . 7 10 . 6 10 . 7 10 . 7

. 4.12 • 10 . 6 3.09 • 10 - s 2.25 • 10 -5 1.69 • 10 . 7 9.2 • 10 - s 6.9 • 10 - l ~ -

S

Cs

(mg / L)

(mmol / m 3)

log Kow

31 28 25 7 3.80 16.1 1.90 1.10 0.045 0.132 0.26 2 • 10 -3e 0.043 0.011 0.0038 0.004 0.0004

241 197 176 45 24.6 107 11.4 6.17 0.253 0.652 1.19 0.188 0.048 0.0151 0.016 0.0016

3.37 3.87 3.86 3.90 4.00 e 4.00 4.18 4.57 4.54 5.18 5.22 5.75 e 5.49 5.91 6.04 6.25

0.0015

0.0060

5.80

0.0008 0.00026 0.0006 0.00014

0.0032 0.00097 0.0022 0.00047

6.00 6.50 6.75 -

a A d a p t e d from Mackay et al. (1992). T h e r e are their " s e l e c t e d " (i.e., " b e s t " ) values given in their Summary off). See also Sonnefeld et al. (1983) for " c o n v e n t i o n a l " v a p o r pressures. b Abbreviations: B[a]A, benz[a]anthracene; B[b]F, benzo[b]fluoranthene; B[k]F, benzo[k]fluoranthene; DB[ a,h ]A, dibenz[a,h]anthracene. c Source: Finizio et al. (1997). d This H is for H e n r y ' s law expressed as H = P / [ P A H ] , where P is the gas-phase concentration in pascals concentration in moles p e r cubic meter. Traditionally in atmospheric chemistry, H e n r y ' s law is expressed as H = e Source: Table 1, Mackay and Callcott (1998).

a. P A H s

Sampling quantitatively each compound in the spectrum of relevant PAHs present in urban ambient air is challenging because their conventional vapor pressures cover a range of some 10 l~ Torr, e.g., ~ 8 • 10 .2 for naphthalene vs 1.5 • 10 -12 for coronene (see Table

10.6). Furthermore, the concentrations of the gas phase 2and 3-ring PAHs are generally far higher than those of the 5- and 6-ring particle-phase species. Thus, as seen in Fig. 10.2, average concentrations in urban southern California air at four sites (shown in Figure 10.23)were ~6000, 30, and 50 ng m -3 for gas phase naphthalene, fluorene, and phenanthrene, compared to ~0.14, 0.29, and 0.77 ng m -3 for BaP, indeno[1,2,3-cd]pyrene and benzo[ghi]perylene, respectively in the particle phase (Fraser et al., 1998). Although the temperatures were

HU (Pa m 3 mol - l)

Calculated log KOAc

43 45 51 29 12.17 e

5.13 e 5.61 e

8.4 7.87 e 3.24 e 3.96 e 0.92 1.04 0.0122 e 0.012 0.581 e 0.046 0.020 0.003 0.016 0.075 0.00017 e 6.75

6.23 e 6.68 7.45 7.34 8.43 e 8.60 10.44 10.80 e 10.71 e 11.13 11.70 e 11.19 13.91 e

Table 2.2 (some were r o u n d e d

B[ghi]P, benzo[ghi]perylene;

and [PAH] is the liquid-phase [ X ] / P x.

quite high during this period (27~ day and 22~ night), similar relative concentrations of gas-phase vs particlephase PAHs have also been seen in other studies at lower temperatures. For example, Fig. 10.3 shows measured PAHs in Chicago, Illinois, during F e b r u a r y / March 1995 when the mean day/night temperatures were ~ I~ (Harner and Bidleman, 1998). As seen in Fig. 10.3, ~96% of the fluorene and phenanthrene were in the gas phase (naphthalene was not sampled) and accounted for ~56% of the total mass of the measured PAHs. Table 10.8 summarizes some measured concentrations of various PAHs in a variety of locations as well as the percentage of each found in the particle phase. Consistent with the data in Figs. 10.2 and 10.3, the percentage found in the particle phase increases with the size of the PAH. For example, phenanthrene con-

455


TABLE 10.7 Vapor Pressures (Pascals)a as a Function of Temperature between 283.15 and 323.15 K for Several Polycyclic Aromatic Hydrocarbons b logP~ Compound

o o o

-A/T+B

A

o

B o

Naphthalene Naphthalene-ds Acenaphthylene Acenaphthene Fluorene Phenanthrene Phenanthrene-d]0 Anthracene Fluoranthene Benz[a]anthracene Pyrene

3960.03 3689.52 3821.55 4535.39 4616.07 4962.77 4704.13 4791.87 4415.56 4246.51 4760.73

+ 58.9 + 27.6 _+ 23.7 _+ 47.3 _+ 30.3 _+ 32.8 _+ 31.5 _+ 50.3 _+ 46.2 _+ 132.2 _+ 26.2

14.299 13.392 12.768 14.669 14.385 14.852 14.060 12.977 11.901 9.684 12.748

+ 0.200 + 0.09 _+ 0.079 _+ 0.159 _+ 0.101 _+ 0.109 _+ 0.11 _+ 0.170 _+ 0.155 _+ 0.431 _+ 0.087

co

E r v t-

Vapor-phasei~

~ 9 10

Particle-phase

o 0 o

E

5

a l P a = 7.5 • 10 -~ mm. b Adapted from Sonnefeld et al. (1983).

centrations range from 5.1 to 78 ng m -~, with 0-4.5% in the particle phase, whereas for fluoranthene, they range from ~3.5 to 12.4 ng m -~, with 0.7-23% reported in the particle phase. As discussed in Chapter 9.D, another important factor determining the gas-particle partitioning is the mass of total suspended particulate matter (TSP) in the air parcel being sampled, and the size and chemical and physical properties of their surfaces. For example, Kamens and co-workers (1995) calculated the effects (Table 10.9) of PAH vapor pressure and the size and concentrations of coexisting aerosols on the gas-particle partitioning of phenanthrene (3 rings), pyrene (4 rings), and benzo[a]pyrene (5-rings). They concluded, for example, that the percentage of pyrene in the particle phase would increase from ~3 to 77% in going from an air mass at 25~ containing 10 /xg m - 3 of 0.5-/xg-size aerosols to one with 500 /xg m -~ of 0.25-/xm-size aerosols. There is an important caveat. These, and other experimental values cited for the gas- and particlephase concentrations of semivolatile PAHs, do not necessarily represent their actual equilibrium concentrations in the ambient air being sampled. Thus, there can be both positive and negative artifacts associated with the sampling methodologies. Furthermore, sorption equilibrium may not be achieved in the polluted ambient air. Let us now briefly consider experimental methods for sampling, beginning with the gas phase. Naphthalene, the methylnaphthalenes, and other abundant and highly volatile 2-ring PAHs are quantitatively trapped with Tenax-GC or Tenax-TA solid absorbents (e.g., see Arey et al., 1989a; Baek et al., 1992; and Zielinska et

m

m

rI1)

e~

e"

1.1_

C~

z

m eI1)

~

rI1)

e~

eI1)

rI1)

i_

i..,--'-,

"~

"~

=,T

e~

eI1)

rI1)

e~

eI1)

rI1)

e~

r

~

='~'~=

o o~0%

~:~omm N

,

~

~

N

~

0

oO~O>,~

oo

9 r m

-o c

t-

m

FIGURE 10.2

Average ambient concentrations (ng m - 3 ) of gasand particle-phase PAHs collected at four urban sampling sites across the Los Angeles Air Basin for eight 4-h sampling periods over two hot days, September 8-9, 1993, during a heavy smog episode (peak ozone 290 ppb at Claremont, California). For concentration data and site locations, see Table 10.8 and Fig. 10.23. Naphthalene was collected in a V O C stainless steel canister; 3- to 6-ring PAHs were sampled independently using a quartz fiber f i l t e r / P U F plug, Hi-Vol dichotomous virtual impactor system similar to that shown in Fig. 10.4 (adapted from Fraser et al., 1998).

al., 1996). For example, Arey and co-workers (1989a) have used Tenax-GC in "low-flow" ( ~ 1 L min -1) and "high-flow" cartridges (~10 L min-1). Naphthalene was sampled with the low-flow cartridge because of its great abundance (e.g., the highest naphthalene concentration in a 12-h sample was 6100 ng m -3 = 1.1 ppb), and the remaining seven PAHs with a high-flow cartridge. Table 10.10 gives the concentrations of volatile PAHs averaged over a 9-day event in August 1986, with nine 12-h daytime and nine 12-h nighttime sample collection periods. The sampling site was less than 1 km from a major freeway so that these data are characteristic of air parcels heavily impacted by motor vehicles. The nighttime average concentrations for naphthalene were

45 6


4O -1-

35

~6 80 ffJ

t~ E t~

25

m

.4--'

20 r

._o +.,

15

"r-

~0

0

=

m {-

c-

c-

c-

LL

~

~-,

~_

,-"

c-

c-

c-

c-

{-

L_

L_

c-

c-

i,_

i._

e-- r

+-,

ft.

o =

N I::=

~

c-

c-

c-

c-

L_

Q)

~

_~

+-,

=

o

~

.~

%

~

~

.~--

I~I

O

o ~

O N

O N

rn

rn

N

~

C ~~

c-

~::

'~' O

~ N

"o

.-

I::=N

Polycyclic aromatic hydrocarbon

FIGURE 10.3

Average mass percent contributions (gas + particle phase) of 13 PAHs collected in urban Chicago air during 15 sampling events in February/March 1995. Open bars, approximate percentage in gas phase; solid bars, approximate percentage in particle phase. 2-Ring PAHs not determined. Average percent of total mass in the particle phase was 4.4 and 42.9 for 3- and 4-ring PAHs, respectively. Averages: EPAHs = 58 ng m-3; TSP = 66 /xg m-3; temperature 1~ (range 0 to -4~ Sampling apparatus, double glass fiber filters and two PUF plugs (adapted from Harner and Bidleman, 1998).

consistently higher than for the daytime periods, ranging from 3000 to 6100 ng m -3 (overall average 4300) at night compared to 2000-4300 (overall average 3100 during the day). This was also true for the methylnaphthalenes and the six other volatile PAHs. Concentrations of the 2-methyl isomer were always significantly higher than those of the 1-methyl isomer, averaging 240 vs 130 ng m -3 daytime and 710 vs 370 ng m -3 nighttime. Concentrations of the gaseous nitroarenes, 1- and 2-nitronaphthalene, and 3-nitrobiphenyl in these 18 samples were determined concurrently with the PAHs, in this case sampling with a Teflon-impregnated glass fiber filter (TIGF)/polyurethane foam (PUF) apparatus (Fig. 10.4). The PUF plugs quantitatively trapped the gas-phase nitroaromatics, but not naphthalene and biphenyl and several 3-ring PAHs. Cass and Fraser (private communication) and Fraser et al. (1998) have used another technique, collection in an internally electropolished stainless steel canister, to

collect naphthalene and the 1- and 2-methylnaphthalenes and several monoaromatics, e.g., benzene and toluene. Semivolatile 3- and 4-ring PAHs were sampled with a glass fiber filter (GFF)/PUF system (see later). Naphthalene concentrations averaged 6000 ng m -3 and ranged from 0.0 to 22,600 ng m -3, consistent with the measurements of Arey et al. (1989a)discussed earlier (see Table 10.10). However, Fraser et al. (1998) noted that this canister sampling technique may underestimate the methylnaphthalenes. Thus, their methylnaphthalenes/naphthalene ratios were lower than those obtained by Arey et al. (1989a) using Tenax-GC solid adsorbent. This may be due to significant adsorption of the methyl isomers to the canister (Arey, personal communication). Zielinska and co-workers (1996) evaluated measurement methods for VOCs up to C20 emitted from motor vehicles and reported that Cs-C12 hydrocarbons were more stable on the Tenax cartridge than in canisters. Similar problems with canister sampling for organics are discussed in Chapter 11.A.4e. A variety of filters have been used to collect 5- and 6-ring particulate PAHs. These include glass and quartz fiber, GFF and QFF (e.g., Ligocki et al., 1989; Fraser et al., 1998), Teflon membrane, TM (e.g., Ligocki et al., 1989; Smith and Harrison, 1996), and Teflon-impregnated glass fiber filters, TIGF (e.g., Arey et al., 1987; Feilberg et al., 1999a). A common type of apparatus for determining total PAH and PAC concentrations and their gas-particle partitioning that has found widespread use is shown in Fig. 10.4 and is illustrated by studies such as those by Cautreels and Van Cauwenberghe (1978), Thrane and Mikalsen (1981), Yamasaki et al. (1982), Atkinson et al. (1988a), Arey et al. (1989a, 1992), Harger et al. (1992), Sasaki et al. (1995), Gupta et al. (1996), and Fraser et al. (1998). A Teflon-impregnated glass fiber filter that collects particle-associated PAHs is followed by three glass cylinders packed with polyurethane foam ("PUF plugs") that trap the gas-phase PAHs that pass through the particle filter. The filter and the PUF plugs are extracted by organic solvents (e.g., CHzC12) and the particle-phase and gas-phase PAHs (and PACs) identified and quantified by various techniques. Generally, a conventional Hi-Vol sampler draws the polluted ambient air through the apparatus; however, given the availability of highly sensitive analytical techniques and microbiological assays (e.g., the "microsuspension modification" to the Ames Salmonella t y p h i m u r i u m bacterial assay; see Section C.2.c), medium- and low-volume samplers are often used as alternatives. Two to four PUF plugs are typically used in the sampling train, depending on sampling conditions (e.g., ambient temperature) and the volatility of the PAHs

45 7

A. NOMENCLATURE AND SELECTED PHYSICAL AND SPECTROSCOPIC PROPERTIES OF PAHs AND PACs T A B L E 10.8 Sum of the Particle-Associated (Filter) and Gas-Phase (Solid Adsorbent, P U F Plugs f or T e n a x Cartridges g) Concentrations of E P A Priority P A H Pollutants and the Percentage of Each P A H in the Particle Phase Polycyclic aromatic hydrocarbon, formula, and molecular weight

Total concentration (filter + solid adsorbent, ng m - 3 ) and percentage (in parentheses) of specific PAHs in particle phase Birmingham University, U.K. a

London, U.K?

Naphthalene, C8Hlo, 128 Acenaphthylene, C12 H8,152 Acenaphthene, C12 H10,154 Fluorene, C13H10, 166 Phenanthrene, C14 H10,178 Anthracene, C14 Hlo, 178 Fluoranthene, C16 Hlo, 202 Pyrene, C16Hlo, 202

15.4 (4.0) 13.5 (11.9) 13.7 (7.7) 24.1 (4.5) 4.49 (8.7) 12.4 (9.4) 38.0 (6.2)

Cyclopenta[cd]pyrene/C18 H10,226 Benz[a]anthracene, C18 H12 , 228 Chrysene, C18 H12,228 Benzo[ b ]fluoranthene,

C2oH 12,252

Benzo[k ]fluoranthene, C2o H 12,252

5.59 (26) 6.49 (34) 2.15 (87) 1.20 (93)

Benzo[ e ]pyrene/C20 H12,252 Benzo[ a]pyrene, C20 H12,252

Benzo[ghi]perylene, C22H12 , 276 Indeno[1,2,3-cd]pyrene,

C22H 12,276

Dibenz[a,h]anthracene,

C22H 14,278

0.81 (90) 1.97 (97) 1.96 (100) 0.83 (94)

Los Angeles Air Basin c

Torrance, CAu

6000 b

3300 g

(o)

(0)

15.6

32.0 (0.1)

(o) -

29.8

-

(0) 5.12 (2.1) 2.84 (6.3) 3.46 (23) 3.79 (21) 4.08 (64) 1.41 (56) 1.62 (75) 1.78 (9O) 0.75 (91) 2.31 (87) 1.63 (88) 3.31 (100) 1.57 (100) 0.12 (100)

Portland, OR e

50.3 (0) 3.04 (0) 9.85 (0.7) 7.30 (1.0) 0.40 (35) 0.25 (60) 0.78 i (44) -

78 k (0.4) 6.1 (0.5) 8.0 (5.8) 8.0 (7.5) -

0.22 (100) 0.22 (100) 0.14 (100) 0.77 (100) 0.29 (100) -

-

-

11.1 (0.6) 26.3 (1.0) 3.44 (O.9) 8.42 (4.9) 7.33 (7.1) 1.57 (73) 2.06 (65) 3.71 j (96)

2.1 (100) 0.59 (98) -

a From Smith and Harrison (1996). Sampler elevation 15 m, situated ~300 m from a high traffic "spine roadway." Particle phase collected on Teflon membrane filter (TMF); gas-phase species trapped downstream on PUF plugs. Daily 24-h samples collected during February 1992; mean TSP during that period 60/xg m-3. b From Baek et al. (1992). Annual average concentrations derived from 48 week-long curbside samples collected throughout 1987 at an urban, traffic-dominated site in south Kensington, London. Sampler elevation 5 m; particle-associated PAHs collected on a PTFE filter (Millipore Ltd.) and gaseous PAHs on a cartridge containing Tenax-TA adsorbent. c From Fraser et al. (1998). Average concentrations of eight 4-h samples collected day and night at four sites in southern California (see Fig. 10.2) during an intense two-day photochemical smog episode, September 8-9, 1993 (1-h average O 3 peaked at 0.29 ppm at Claremont, California). Volatile PAHs collected in stainless steel canisters, with the remainder collected with a quartz fiber filter/PUF solid adsorbent system. Average daytime and nighttime temperatures (~ Long Beach, 22.2 (19.6); downtown Los Angeles, 25.8 (20.6); Azusa, 28.5 (22.9); Claremont, 30.4 (23.1) (Cass and Fraser, personal communication, 1999). d From Arey et al. (1987). Site, a roof (9-m elevation) at E1 Camino Community College in Torrance, California (20 km south of central Los Angeles). Daytime 12-h sampling period, 0600 to 1800 hours, February 25, 1986; maximum temperature 35~ at 1100 hours. Sampling systems: Tenax-GC cartridge for naphthalene and phenanthrene; Teflon-impregnated glass fiber filter (TIGF) and PUF plugs for phenanthrene and heavier PAHs (see diagram, Fig. 10.4). e From Ligocki and Pankow (1989). One sampling system had a glass fiber filter (GFF) and backup PUF plugs. Others employed either two GFFs or a single Teflon membrane filter (TMF); both had downstream PUF plugs. See original article for experimental details. Samples collected during February and April 1984 and February and April 1985; overall mean sampling temperatures were 8 and 5~ respectively. Sampling periods 5-30 h. For gas- and particle-phase concentrations of those PACs that were copollutants with these PAHs, see Table 10.11. 7 Gas-phase PAHs collected on PUF plugs unless otherwise noted (e.g., on Tenax-GC or Tenax-TA adsorbents). g Collected on a Tenax-filled cartridge. h Collected in a canister sampler. i Chrysene/triphenylene. J Benzol b + j + k]fluoranthene. k Same value for collection on Tenax-GC cartridge and the sum of the concentrations of phenanthrene trapped on three PUF plugs in series. 1Not EPA priority pollutants.

458


TABLE 10.9 Calculated Effect of Vapor Pressure and Aerosol Concentration on the Percentage in the Aerosol Phase of Phenanthrene, Pyrene, and Benzo[a]pyrene a Aerosol concentration c (ixgm -3)

10 100 500

Percentage of PAH in particle phase Aerosol size (ixm)

Phenanthrene (5.3 X 1 0 - 4 ) b

Pyrene (3.3 • 1 0 - 5 ) b

Benzo [ a ] pyrene (9 • 1 0 - s ) b

0.5 0.35 0.25

0.2 3.0 17

3 32 77

93 99 100

a Adapted from Kamens et al. (1995). b Subcooled liquid vapor pressures in Torr at 25~ calculated by Kamens et al. (1995) from Hawthorne et al. (1989). c Concentration of aerosol, /xg m-3; particle density 1.25.

and PACs being sampled. However, Arey and coworkers (1989a) found that even with four PUF plugs, fluorene (and certainly not naphthalene and the methylnaphthalenes) was not collected quantitatively and some breakthrough of phenanthrene also occurred, especially during daytime sampling periods. Thus seven of the volatile 2- and 3-ring PAHs were collected on Tenax-GC high-flow cartridges, and naphthalene, because of its very high concentrations (up to 6100 ng m-3), on a low-flow cartridge. Tenax-TA solid adsorbent also effectively traps these volatile PAHs (Zielinska et al., 1996). Because of these sampling difficulties, the reported gas-phase concentrations of the semivolatile 3- and 4-ring PAHs are "operationally defined"; i.e., they are the quantity of a given PAH extracted from the solid adsorbent PUF plugs (or other solid sorbents such as Tenax cartridges). Similarly, the 5- and 6-ring PAHs along with that fraction of the semivolatile PAHs collected on the "upstream" filter are "operationally" defined as being in the particle phase. We emphasize the term "operational" when describing the gas- and particle-phase concentrations of PAHs sampled using this popular filter/sorbent technique, because it has been recognized for some time that this and other methods are subject to both positive and negative artifacts (e.g., see Van Vaeck and Van Cauwenberghe, 1984; Fitz et al., 1984; Coutant et al., 1988; Back et al., 1991b; Kaupp and Umlauf, 1992; Hart and Pankow, 1994; Kamens and Coe, 1997; and Feilberg et al., 1999a). These artifacts may lead to significant deviations of the measured gas and particle concentrations of semivolatile organics (SOCs) from their "true" equilibrium concentrations used in theoretical treatments of the gas-particle partitioning of PAHs (see Chapter 9.D and references cited above). As examples, the term "blow off" refers to a phenomenon in which the pressure drops across a filter can cause particle-associated SOCs (e.g., 3- and 4-ring

PAHs) to be stripped from the collected particulate matter and trapped downstream on the PUF sorbent, leading to an underestimate of their true particle-phase concentrations and an overestimate of their gas-phase levels. However, several studies (e.g., McDow and Huntziker, 1990; Turpin et al., 1994) suggest that this is less of a problem than adsorption of organics; see following discussion). Conversely, gaseous semivolatile organics can adsorb to filter surfaces, resulting in artificially overestimating their particle-phase concentrations and underestimating their gas-phase levels. For example, Hart and Pankow (1994) conducted a study using two filter/PUF sorbent samplers similar to the apparatus in Fig. 10.4 but each incorporating a "backup" filter immediately downstream from the regular filters used to collect particle-phase PAHs. They found that for a typical sampling event in ambient air, gas adsorption to a single quartz fiber filter in a conventional Hi-Vol sampler could shift the apparent gas-particle partitioning parameter in favor of the particle phase, leading to overestimations by factors ranging from ~ 1.2 to 1.6. However, a second sampler operating in parallel and using a single Teflon membrane filter gave gas-particle distributions approximately the same as the corrected quartz fiber filter values. Gas-particle partitioning is also impacted by changes in ambient temperature and aerosol concentrations during sampling, resulting in artifacts in the relative particle- and gas-phase concentrations (see Hart and Pankow, 1994). A sampling system designed to reduce artifact problems was employed by Liang and co-workers (1997) in an experimental chamber study of the gas-particle partitioning of several PAHs (and n-alkanes) on three types of model aerosols and ambient urban particulate matter. The system consisted of a dual glass fiber filter system as described by Hart and Pankow (1994), followed by two parallel sampling trains for trapping


gas-phase species. One had two sequential PUF plugs and the other had two Tenax-GC cartridges in sequence to trap the more volatile species. An alternative approach is that using denuder-based samplers for determining gas-particle partitioning of semivolatile PAHs and other organics (see Feilberg et al., 1999a; Chapter ll.A.3b has a description of denuders.) In principle, gas-phase PAHs in ambient air are first trapped on a denuder surface coated with a sorbent material (e.g., a resin); particles pass through and are collected on a follow-up filter. In practice, experimental problems exist, for example, in extracting quantitatively species from the sorbent coating on the denuder walls. Gundel and co-workers (1995a) addressed

I

> ~

~d 9- uff

m

tt')

INTAKE POLLUTED AMBIENT AIR

tt')

g~ d

m

9 +.a

f

HI-VOL FILTER ~ ~_

~
,

12 I

I

400

"6

6

I 440

I

;L(nm)

I 480

Xl

o~

0

I

210

I

I

I

I

250

I

I

I

I

I

300

I

I

I

I

I

350

I

I

I

I

I

400

X (nm) F I G U R E 10.7 U V absorption and fluorescence spectra of anthracene in cyclohexane (adapted from Karcher et al., 1985).

Gundel et al., 1995b). Advantages of using this approach include high sensitivity (50 ng per gram of sample in these particular studies) and the ability to discriminate between compounds such as benzo[b]fluoranthene and benzo[k]fluoranthene, which is difficult by other techniques such as gas chromatography.

Solvents also affect the measured absorption spectra. Thus, the ~r,~r* bands shift to longer wavelengths (a "red" bathochromic shift) in polar solvents. For example, the long-wavelength band of anthracene shifts from ~375 nm in n-hexane to 381 nm in acetonitrile (Wehry, 1983).

i

Fluorescence

14 12

!

'O

'~ if}

6

.Xl0

/

~

/

4

I

i

,

360 __ m )

i

400

,

i

440

I

2 0

I

210

I

I

I

I

250

I

I

I

I

I

300

I

I

I

I

I

350

I

I

I

I

I

400

;L (nm) F I G U R E 1 0 . 8 U V absorption and fluorescence spectra of pyrene in cyclohexane (adapted from Karcher et al., 1985).

A. NOMENCLATUREAND SELECTEDPHYSICALAND SPECTROSCOPICPROPERTIESOF PAHs AND PACs

Absorption

463

Fluorescence

~- 7r

6=

E o

5-

~

4

E

I

A

I

i

400

I

i

440

I

i

480

I

i

520

560

i

I

600

i

640

J

~

T"

3

O

N ,m

2

t--

& ~

]

o I

I

I

I

210

I

250

I

I

I

I

I

300

I

I

I

I

I

350

I

I

I

I

I

400

X (nm)

FIGURE 10.9 UV absorption and fluorescence spectra of fluoranthene in cyclohexane (adapted from Karcher et al., 1985).

tions with respect to the atmospheric reactions, lifetimes, and fates of PACs. Thus, as discussed in Section F, photolysis significantly exceeds OH radical attack as an "efficient" loss process for some gas-phase nitroarenes such as 1-nitronaphthalene (e.g., see Atkin-

A d d i n g N-, O-, or S-atom functionalities to a P A H can cause m a j o r changes in its U V - v i s i b l e absorption spectrum. F o r example, as seen in Fig. 10.12, addition of N O 2 groups to BaP to f o r m the 1-, 3-, and 6-nitro isomers results in p r o n o u n c e d red shifts in their absorption spectra (Pitts et al., 1978). This e n h a n c e d ability to absorb solar radiation has significant implica-

son et al., 1989; and Feilberg et al., 1999a). T h e physical state of a P A H also can have a dra-

~"

6

7 -~

/

Fluorescence

i

A

4

420 X (nm)460 ~

.

.

I

I

500

o I

210

I

I

I

I

I

I

I

I

I

300

I

I

I

I

X (nm)

I

I

I

I

I

I

400

I

FIGURE 10.10 UV absorption and fluorescence spectra of benzo[a]pyrene in cyclohexane (adapted from Karcher et al., 1985).

464

10. AIRBORNE PAHs: CHEMISTRY AND TOXICOLOGICAL IMPLICATIONS 7

0

..Q

6 5

'7

E 0

'7

4

__m 0

E

_1

0 if)

3 2 1 0

I

210

I

I

I

I

250

I

I

I

I

I

300

I

I

I

I

I

350

I

I

I

I

I

400

I

I

I

I

I

450

I

I

I

I

500

~. (nm)

F I G U R E 10.11 UV-visible absorption spectrum of cyclopenta[cd]pyrene in cyclohexane (adapted from Karcher et al., 1985). This PAH does not fluoresce.

T A B L E 10.12 C o n c e n t r a t i o n s of 10 P A H s in N a t i o n a l I n s t i t u t e of Standards and T e c h n o l o g y ( N I S T ) Standard R e f e r e n c e Material S R M 1 6 4 9 A i r Particles D e t e r m i n e d by D u a l P r o g r a m m a b l e F l u o r e s c e n c e D e t e c t o r M e t h o d C o m p a r e d to T h e i r N I S T R e f e r e n c e C o n c e n t r a t i o n s a

Concentration (Ixg / g)

Wavelength change (at time in min)

Excitation wavelength (nm)

Emis sion wavelength (nm)

Detector I c Phenanthrene Pyrene Benz[a]anthracene Benzo[ a]pyrene Indeno[1,2,3-cd]pyrene

0.0 11.5 19.8 27.5 30.3

250 235 225 290 245

370 380 395 410 480

4.5 6.3 2.6 2.9 3.3

+ 0.3 + 0.4 _+ 0.3 _+ 0.5 + 0.5

7.3 5.6 2.8 2.8 4.0

_+ 0.6 _+ 1.0 _+ 0.1 _+ 0.2 _+ 0.1

Detector 2 c Fluoranthene Chrysene Benzo[ b]fluoranthene Benzo[k]fluoranthene Benzo[ghi]perylene

0.0 19.4 22.9 22.9 30.4

230 260 230 230 225

450 370 430 430 415

7.1 3.5 6.2 2.0 4.5

_+ 0.5 _+ 0.1 _+ 0.3 _+ 0.1 _+ 1.1

6.5 3.4 5.7 2.2 3.4

_+ 0.7 + 0.1 + 0.3 _+ 0.1 + 0.2

PAH

a Adapted from Mahanama et al. (1994). b Certificate of analysis, Standard Reference Material 1649, Urban Dust/Organics (see Box 10.3). c Detector 1 follows the column, detector 2 follows detector 1.

Reference b (NIST)

Measured


5.0 4.5 4.0 oJ o

I

~'~~."/.~.//1 " " ,

3.5

"J

O 3,0

~

1112 1

--

9 8 ~ 4

2.5-

"- ~ "~

k

,,

\

2 3

7 6 5

BaP

I

\

/~

\ \

I

300

\

\

\

6-NO2-BaP

2.0200

1 -NO2-BaP \\3NO2"BaP

I

400

I

500

Wavelength

600

(nm)

F I G U R E 10.12 Absorption spectra in methanol of benzo[a]pyrene and its 1-, 3-, and 6-nitro derivatives (adapted from Pitts et al., 1978).

cm

20000

-1

25000

I

I

Vapor I& ,

.i

--

30000

~ ~

I

I

FI

Abs

~.1

.

\

tt.O i O.

ti n

to O

Solution

~

J

/(V~l~,/'~J ~

FI

~

I

O .13

Abs / ~ ' ~

oJ C~ O _d

I.L O ._1

~~J 500

-

FI I

450

~1

Abs

I

400

~%%~ I

350

300

Wavelength(nm) F I G U R E 10.13 Absorption and fluorescence spectra of anthracene in three phases. Pronounced red shifts occur going from the vapor state to solution in dioxane to the solid state; separation of the 0-0 bands also increases (adapted from Bowen, 1946).

465

466


matic effect on its UV-visible absorption spectrum. Figure 10.13 shows the pronounced red shifts for anthracene for both absorption and fluorescence accompanying the phase changes from gas to solution to solid. This phenomenon, demonstrated by E. J. Bowen over a half century ago, is discussed in his classic monograph Chemical Aspects of Light (1946). Infrared spectra of PAHs and PACs are also available, in part because of interest in PAHs in interstellar space (e.g., see Hudgins and Sandford, 1998a, 1998b, 1998c; and Langhoff et al., 1998). Their theoretical treatment and experimental infrared data may also prove useful in studies related to the troposphere. For a detailed discussion of the fundamental aspects of the spectroscopy and photochemistry of PAHs and PACs, see monographs by Calvert and Pitts (1966), Turro (1978), Wayne (1988), and Gilbert and Baggott (1991).

B. BIOLOGICAL PROPERTIES OF PAHs A N D PACs. I: CARCINOGENICITY 1. Historical Perspective: Benzo[a]pyrene, the "Classic Chemical Carcinogen" In the latter part of the nineteenth century, workers in the paraffin refining, shale oil, and coal tar industries had high incidences of skin cancer. A possible cause emerged during the period 1915-1918 when Japanese scientists discovered that painting the ears of rabbits and mice with coal tar extracts produced tumors, some of which were malignant (Yamagiwa and Ichikawa, 1918). The subsequent search for the "coal tar carcinogen" by chemists and medical researchers is classic (Phillips, 1983; Osborne and Crosby, 1987). Starting in 1922, research by a team of British chemists at the Institute of Cancer Research in London culminated in the synthesis in 1930 of the first pure chemical compounds to demonstrate carcinogenic activity, dibenz[a,h]anthracene (DBA), XXV, and its 3-methyl derivative (Kenneway, 1930; Kenneway, 1955, references therein). After distilling and fractionating two tons of pitch from a "Gas, Light, and Coke Company," they ultimately isolated several grams of two C20H12 polycyclic aromatic hydrocarbons. Three years later, Cook and co-workers (1933) synthesized for the first time benzo[a]pyrene and benzo[e]pyrene and proved them identical to these two "coal tar PAHs." Ultimate confirmation of the carcinogenicity of BaP came when all five survivors of a group of ten mice whose backs had been painted with

synthetic BaP developed tumors; BeP was not carcinogenic (Cook et al., 1933). Subsequently, BaP was identified in carcinogenic extracts of ambient air particles (Leiter et al., 1942; Leiter and Shear, 1943; Waller, 1952; Kotin, 1954) and in chimney soot (Goulden and Tipper, 1949).

? XXV Dibenz[ a,h ]anthracene

Since then, long-term sampling data show that annual averages of BaP concentrations in major urban areas have dropped dramatically over a period of several decades. For example, at a roadside sampling site (1.5 m above ground level) located on Fleet Street, London, BaP fell from 39 ng m -3 in 1962-1963 to 10 ng m -3 in 1972-1973 (Commings and Hampton, 1976), similar to 2.0 ng m -3 in 1987 measured at a site 5 m above ground in Central London ( B a e k e t al., 1992). (Note that concentrations in urban areas tend to decrease with sampling height, i.e., with distance from mobile sources.) Key reasons for the major decline in BaP concentrations in London and other urban airsheds throughout the world have been the enactment and enforcement of "clean air" legislation and the trend to cleaner fuels (see, for example, NRC, 1972, 1983; Hoffman and Wynder, 1977; Bjorseth, 1983; Grimmer, 1983a, 1983b; Holmberg and Ahlborg, 1983; Osborne and Crosby, 1987; Baek et al., 1992; California Air Resources Board, 1994, 1998; and references therein). There is, however, a word of caution. Holmberg and Ahlborg point out in their 1983 Consensus Report: Mutagenicity and Carcinogenicity of Car Exhaust and Coal Combustion Emissions, "It should be stressed, however, that a reduction in the BaP level does not necessarily mean a reduction in the potential health hazards, since the spectrum of pollutants has also changed with time." For example, in their long-term study of BaP concentrations and the bacterial mutagenicity of ambient particles collected in Sapporo, Japan, from 1974 to 1992, Matsumoto and co-workers (1998) reported that BaP levels dropped 75-80%, but the level of overall bacterial mutagenicity remained "relatively unchanged." Similar concerns with focusing solely on BaP have been expressed by other researchers (e.g., see Lane and Katz, 1977; Pitts, 1983;

B. B I O L O G I C A L P R O P E R T I E S OF P A H s A N D PACs. I: C A R C I N O G E N I C I T Y

Rosenkranz and Mermelstein, 1985a; Tokiwa and Ohnishi, 1986; Lewtas, 1993a; Atkinson and Arey, 1994; Rosenkranz, 1996; Nielsen et al., 1996; Finlayson-Pitts and Pitts, 1997; and Collins et al., 1998). However, BaP has often been used as a "marker" for POM to set air quality and emission standards (e.g., see Nielsen et al., 1995). For example, the Netherlands Environmental Programme 1988-1991 draft document gives an "acceptable level" for the annual average concentration of BaP in ambient air of 0.5 ng m -3 and a "tolerable level" of 5 ng m -3 (Montizaan et al., 1989). Based on "technical and economic feasibility" as well as concentrations in other western European cities, Germany has an "orienting value" of an annual average of BaP of 10 ng m -3 (Montizaan et al., 1989); a basis for this approach, as stated by The Umwelt Bundesamt (Federal Environmental Agency), is that "dose-effect relationships for man do not exist" (Montizaan et al., 1989; Nielsen et al., 1995, 1999a). Such relationships have, however, been established for PAHs in occupational exposures (e.g., see Mastrangelo et al., 1996). Because certain PAHs are carcinogens, and thus there is no "safe level," the WHO (1987) does not recommend one. However, it has developed risk assessments based on BaP "as an indicator" (Nielsen et al., 1995). 2. C a r c i n o g e n i c i t y of P A H s , C a n c e r P o t e n c i e s , and Potency Equivalence Factors There are a variety of sources of data on the carcinogenicity of environmental PAHs and PACs. A key source is a series of monographs published by the International Agency for Research on Cancer (IARC), The Evaluation of Carcinogenic Risks to Humans. This series includes Polynuclear Aromatic Compounds, Vol. 32, Part 1, Chemical, Environmental, and Experimental Data (1983); Overall Evaluations of Carcinogenicity: A n Updating of I A R C Monographs, Volumes 1-42 (1987); and Diesel and Gasoline Engine Exhausts and Some Nitroarenes, Vol. 46 (1989). Other useful evaluations of the carcinogenicity of specific PAHs and PACs include, for example, those of the U.S. Environmental Protection Agency (1986), the California Air Resources Board (1994), and the U.S. Department of Health and Human Services, Public Health Service, National Toxicology Program (U.S. DHHS, 1998). Evaluations of the carcinogenicity of selected PAHs and PACs are summarized in Tables 10.13 and 10.14. Designations are defined in Box 10.7. Unless otherwise noted, we use IARC definitions and symbols for the carcinogenicity of a given agent; for example, BaP and benzo[b]fluoranthene (BbF) are

467

Class 2A and 2B animal carcinogens and "probable" and "possible" human carcinogens, respectively. A note of caution is appropriate, however, regarding differences in the classification schemes. For example, the EPA classifications for BaP and BbF are both B2, "sufficient evidence from animal studies." Because it is not only BaP but also a variety of other PAHs and PACs that are of concern in terms of the possible inhalation cancer risk to humans of complex mixtures of combustion-generated POM, a number of approaches have been developed for evaluating the potencies of various compounds (e.g., Holmberg and Ahlborg, 1983; IARC, 1989; Nisbet and La Goy, 1992; Lewtas, 1985b, 1993a, 1993b; 1994; Heinrich et al., 1994; CARB, 1994; Mastrangelo et al., 1996; Nielsen et al., 1996; OEHHA, 1998; CARB, 1998; Collins et al., 1998; Tokiwa et al., 1998). One approach is to calculate the inhalation "unit risks" for excess lung cancer for BaP and each of its copollutant carcinogenic PAHs and PACs in polluted ambient air. The latter values are divided by the unit risk for BaP to obtain their individual potency equivalence factors, PEFs, based on BaP = 1.00, e.g., Nielsen et al., 1996, and Collins et al., 1998. These PEFs are listed in Table 10.13 for PAHs and Table 10.14 for PACs. The PEFs range from 0.01 for chrysene and 2-nitrofluorene to 10 for each of the three 6-ring dibenzopyrenes (C24 H14 , MW 302) and for the N-PACs 6-nitrochrysene and 1,6-dinitropyrene. Interestingly, Cavalieri et al. (1989, 1994) reported that the dibenzo[a,/]pyrene isomer is actually 100-200 times more tumorigenic than BaP, and they termed it "the most potent carcinogenic aromatic hydrocarbon." To assess both the relative and absolute contributions of various PAHs and PACs to health impacts, the potencies must be combined with concentrations of the individual PAHs and PACs in air. The levels of compounds other than BaP can be quite substantial and hence contribute significantly to the overall carcinogenicity and mutagenicity. For example, Allen and co-workers (1998) reported that, while the individual concentrations of the biologically active 6-ring PAHs identified and quantified in urban Boston air are relatively small (e.g., the c o n c e n t r a t i o n of dibenzo[a,e]pyrene was 0.133 ng m-3), their total concentration of ~1.5 ng m -3 is comparable to the BaP present as a copollutant in the sample. Collins et al. (1998) applied the PEF data for PAHs and PACs (Tables 10.13 and 10.14) to daytime concentrations measured in ambient air in Riverside, California ( ~ 9 0 km east and downwind of Los Angeles; see Fig. 10.23), during the months of July and August 1994 (Atkinson and Arey, 1997; Krieger et al., 1997). As seen in Table 10.15, the "PEF adjusted concentration," 269 pg m -3, is ~ 40% of the total mass of PAHs and PACs

TABLE 10.13 Carcinogenicities of Selected PAHs in Ambient Air As Evaluated by IARC, U.S. EPA, and U.S. DHHS, Cancer Potency Equivalence Factors Relative to BaP = 1.00 from Nielsen et al. (1996), CARB (19941, and Collins et al. (19981, and Human Cell Mutagenicities Relative to BaP = 1.00

Structure

Name

IARC Class“ Animals Human

us.

EPA Class”

us.

DHHS ClassC

Potency equivalence factors Nielsend9’ CARB / OEHHAe-”

Relative human cell mutagenicity (BaP = l.OO)h-’

0.022

L

& /

/

/

S

B2

-

0.03

0.01

0.017

2B

NC

S RAHC

-

1.0

0.63

2A

B2

S RAHC

1.0

1.0

1.00

2B

B2

S

0.05

0.1

0.25

RAHC

/

& \

Benzo[b]fluoranthene

3

\

Benzo[k]fluoranthene

S

2B

NC

S RAHC

0.05

0.1

0.26

2B

B2

S RAHC

0.05

0.1

0.11

@ /

/

&

Benzo[ ghilperylene

0.02

-

0.19

0.1

0.1

0.31

0.4

0.29

\

\

/

/

S

2B

B2

S RAHC

Dibenz[a,h]anthracene

S

2A

B2

S RAHC

Cyclopenta[ cdlpyrene

L

3

Dibenzo[ a,e]pyrene

S

2B

NC

S RAHC

1.0

Dibenzo[ a,h]pyrene

S

2B

NC

S RAHC

10.0

Dibenzo[ a,i]pyrene

S

2B

NC

S RAHC

10.0

\

~

~

6.9

0.02

/

@ /

Indeno[l,2,3-cdlpyrene

/

/

2.9

(continues)

TABLE 10.13 (continued)

Structure

Name

Dibenzo[a,l]pyrene

IARC ClassU Animals Human

S

2B

us.

EPA Classh

DHHS Classc

NC

S RAHC

factors CARB / OEHHAe-”

Relative human cell mutagenicity (BaP = 1.00Ih-’

10.0

24

Potency equivalence

us.

Nielsend9’

-

“ From Supplement 7 (IARC, 1987), unless othenvise indicated: I, inadequate evidence; L, limited evidence; ND, no adequate data; S, sufficient evidence; 1, Group 1, the agent is carcinogenic to humans; 2A, Group 2A, the agent is probably carcinogenic to humans; 2B, Group 2B, the agent is possibly carcinogenic to humans; 3, Group 3, the agent is not classifiable as to its carcinogenicity to humans. ’’ Source: U.S. EPA Carcinogen Classifications (1986). Group A, human carcinogen; Group B, probable human carcinogen; Subgroup B1, “agents showing limited evidence of carcinogenicity from epidemiologic studies;” Subgroup B2, “agents for which there is ‘sufficient’evidence from animal studies and for which there is is ‘inadequate evidence’ or ‘no data’ from epidemiologic studies;” Group C, possible human carcinogen; Group D, not classifiable as to human carcinogenicity; NC, not classified. U.S. DHHS/PHS/NTP (1998). Class S is sufficient evidence of carcinogenicity in animals. All PAHs in this table that have been evaluated are classified “reasonably anticipated to be human carcinogens (RAHC).” Nielsen et al. (1996) “City Air Pollution of Polycyclic Aromatic Hydrocarbons and Other Mutagens: Occurrence, Sources, and Health Effects.” California Air Resources Board/Office of Environmental Health Hazard Assessment, Benzo[a]pyreneas a Toxic Air Contaminant (1994); Office of Environmental Health Hazard Assessment/California Environmental Protection Agency, Air Toxics Hot Spot Program Risk Assessment Guideline, Part II: Technical Support Document for Describing Ainilable Cancer Potency Factors (1998); Collins et al. (1998). For the actual values and a discussion on various approaches by a range of scientific groups to calculating such values of unit risks for cancer for the PAHs and PACs in Tables 10.14 and 10.15, see Nielsen et al. (1995, 1996, 1999b), OEHHA/Cal EPA (19981, Collins et al. (19981, and references therein. Collins et al. (1998). ’I Source: Durant et al. (1996). ‘ Recalculated from data in Durant et al. (1996). Several of these relative potencies differ from those given in Hannigan et al. (1998); see Table 10.25.

‘

’

’

’

471

B. BIOLOGICAL PROPERTIES OF PAHs AND PACs. I: CARCINOGENICITY T A B L E 10.14 Carcinogenicities of Selected PACs in Emissions and Ambient Air a As Evaluated by the I A R C , U.S. EPA, and U.S. D H H S , Cancer Potency Equivalence Factors As Evaluated by C A R B / O E H H A (BaP = 1.0), and H u m a n Cell Mutagenicities Relative to BaP = 1.0 CARB / OEHHA

Name

Animals

Humans

U.S. EPA Class c

1,6-Dinitropyrene

S

2B

NC

RAHC

10.0

0.20

O2N~

1,8-Dinitropyrene

S

2B

NC

RAHC

1.0

1.2

(

6-Nitrochrysene

S

2B

NC

-

10.0

2-Nitrofluorene

S

2B

NC

-

0.01

0.05

1-Nitropyrene

S

2B

NC

RAHC

0.1

0.025

4-Nitropyrene

S

2B

NC

RAHC

0.1

1-Nitronaphthalene

I

3

2-Nitronaphthalene

I

3

9-Nitroanthracene

N.D.

3

IARC Class b Structure

U.S. DHHS Class d

Potency equivalence factor e'f

Relative human cell mutagenicity (BaP = 1.0 g)

NO 2

NO2 NO 2

NO2 ~

~

C

~'NO2

NO 2

NO2

I

~

,,NO2

NO 2

(continues)

4 72

10. AIRBORNE PAHs: CHEMISTRY AND TOXICOLOGICAL IMPLICATIONS T A B L E 1 0 . 1 4 (continued)

IARC Class b Structure

Name

U.S. EPA Class c

U.S. DHHS Class d

CARB / OEHHA potency equivalence factor e'f

Relative human cell mutagenicity (BaP = 1.0 g)

0.0026

Animals

Humans

6-Nitrobenzo[ a ]pyrene

L

3

3-Nitrofluoranthene

I

3

-

-

-

Dibenz[ a,j ]acridine

S

2B

NC

RAHC

0.1


S

2B

NC

RAHC

0.1

7H-Dibenzo[ c,g ]carbazole

S

2B

NC

RAHC

1.0

NO 2

H a For definitions of the classifications by the various organizations, see Table 10.13. b From I A R C (Supplement 7, 1987; 1989). c From U.S. E P A (1986). NC, not classified. d From U.S. Department of Health and Human Services (1998); RAHC, reasonably anticipated to be a human carcinogen. e From California Air Resources Board (1994) and Office of Environmental Health Hazard Assessment, California Environmental Protection Agency (Cal EPA) (1998). f From Collins et al. (1998). g From Durant et al. (1996).

(651 pg/m3). That is, the "total inhalation risk" associated with these PAHs and PACs is ~40% of that predicted if one assumes each c o m p o u n d has the same carcinogenic potency as BaP. One important caveat was pointed out by Collins et al. (1998): since PAHs are "multipathway" carcinogens, their calculated inhalation risk may account for only 5-10% of the estimated total multipath exposure risk. Nielsen and co-workers (1996) employed a similar "potency equivalence factor" methodology in their risk assessment evaluation of the health impacts of PAHs and PACs for typical ambient levels encountered at several sites in Copenhagen, Denmark. For other ex-

amples of comparative risk evaluations, see, for example, the calculation by Collins et al. (1998) of the equivalent inhalation risk of ambient aerosols in London, using the 1991 data of Halsall et al. (1994), and Lewtas (1993a, 1993b, 1994, and references therein). Finally, while several volatile and semivolatile PAHs, e.g., naphthalene, the methylnaphthalenes, phenanthrene, pyrene, and fluoranthene, are not significant mutagens or carcinogens (hence not included in Table 10.13), they are precursors to powerful direct bacterial mutagens formed in gas-phase atmospheric reactions with hydroxyl during the day and nitrate radicals at night (see Section F). Furthermore, 2-nitrofluoranthene,

B. B I O L O G I C A L P R O P E R T I E S OF P A H s A N D PACs. I: C A R C I N O G E N I C I T Y

473

B O X 10.7

I A R C , U.S. E P A , A N D U.S. D E P A R T M E N T

OF HEALTH AND HUMAN SERVICES (DHHS) CLASSIFICATIONS FOR CARCINOGENICITY OF CHEMICAL AGENTS

IARC Ouerall Eualuations of Carcinogenicity: An Updating of IARC Monographs, Volumes 1-42, Supplement 7 (1987): I Inadequate evidence L Limited evidence ND No adequate data S Sufficient evidence 1, Group 1 The agent is carcinogenic to humans 2A, Group 2A The agent is probably carcinogenic to humans 2B, Group 2B The agent is possibly carcinogenic to humans 3, Group 3 The agent is not classifiable as to its carcinogenicity to humans U.S. EPA Carcinogen Classifications (1986):

an atmospheric reaction product of fluoranthene, is a potent human cell mutagen present in ambient urban/suburban air environments throughout the world (Ciccioli et al., 1996; Arey, 1998a, and references therein). Clearly, a sound evaluation of the total mutagenic/carcinogenic potencies of a complex mixture of POM emissions (e.g., diesel exhaust) should include not only the PEFs of the primary particle- and vapor-phase PAHs and PACs but also those of the mutagens formed in atmospheric reactions of precursor PAHs (see, for example, Arey et al. (1992), Lewtas (1993b), Atkinson and Arey (1994), Nielsen et al. (1996), Arey (1998a), and Section F). For examples of such formal scientific health risk assessments prepared by the State of California Air Resources Board and Office of Environmental Health Hazard Assessment, see "Benzo[a]pyrene as a Toxic Air Contaminant" (CARB, 1994) and "Identification of Diesel Exhaust as a Toxic Air Contaminant" (CARB, 1998).

3. Carcinogenicity of Nitroarenes and Other Nitro-PACs In the mid-1970s with the advent of a relatively simple short-term bacterial assay, the "Ames Test"

Group A Human carcinogen Group B Probably human carcinogen Subgroup B1 Agents showing limited evidence of carcinogenicity from epidemiologic studies

Subgroup B2 Agents for which there is 'sufficient' evidence from animal studies and for which there is 'inadequate evidence' or 'no data' from epidemiologic studies Group C Possible human carcinogen Group D Not classifiable as to human carcinogenic-

ity U.S. DHHS Report on Carcinogens (1998): Class S is sufficient evidence of carcinogenicity in animals. RAHC is the abbreviation for their listing "Reasonably Anticipated to be Human Carcinogens"

(Box 10.9), a new class of mutagenic compounds was discovered, the nitroarenes. Within a short time, they were identified as directly mutagenic copollutants present along with well-known promutagenic PAHs in combustion-generated emissions and distributed worldwide in urban ambient air (see Section D). Since then, mono- and dinitro-PAHs have been the subject of intense research interest in atmospheric chemistry and toxicology. This is in large part because a number of them have been evaluated by IARC (1989) and other evaluating bodies, e.g., the U.S. Department of Health and Human Services (U.S. DHHS, 1998) and the State of California Air Resources Board and Office of Environmental Health Hazard Assessment (CARB, 1994; OEHHA, 1998), and classified as being "possible" or "reasonably anticipated to be" human carcinogens. These are shown in Table 10.14, along with two dibenzacridines and a dibenzocarbazole. Note that the PEFs of the 1,6- and 1,8-dinitropyrenes are 10 times that of BaP and equal to BaP, respectively; 6-nitrochrysene is also estimated as 10 times as potent as BaP. We discuss the bacterial and human cell mutagenicities and related atmospheric chemistry of these biologically active N-PACs in subsequent sections. Box 10.8 contains examples of literature on this subject.

4 74


TABLE 10.15 Concentrations of Carcinogenic PAHs and PACs Measured in Ambient Air in Riverside, California, Their Cancer Potency Equivalence Factors (PEF), and Calculated PEF-Adjusted Concentrations a

PAH and PAC

Benzo[a ]pyrene Benz[a ]anthracene Benzo[b + j + k]fluoranthenes Dibenzo[a,e ]pyrene Dibenzo[a,h ]pyrene Dibenzo[a,i ]pyrene Dibenzo[a,l ]pyrene Indeno[1,2,3-cd]pyrene Chrysene Dibenz[a,h + a,c ]anthracene 1-Nitropyrene 2-Nitrofluoranthene 4-Nitropyrene 1,6-Dinitropyrene 1,8-Dinitropyrene 6-Nitrochrysene 2-Nitrofluorene Total carcinogenic PAHs

Measured concentration (pg m - 3 ) b

36 39 360 1.7 0.5 /~m for BaP and several other PAHs and two O-PACs in aerosol samples collected at five sites in, and near, Toronto, Canada. The size distribution of the particles was approximately log-normal, and the majority of the mass of the associated PAHs and PACs was found in particles with diameters below 3 /~m. Miguel and Friedlander (1978) used an eight-stage, low-pressure impactor with 50% cutoffs down to 0.05-/~m aerodynamic diameter (Hering et al., 1978, 1979) to determine the mass distributions of BaP and coronene in aerosols collected in Pasadena, California. Approximately 75% of the BaP and 85% of the coronene were associated with particles with diameters less than 0.26 /~m and ~50% of the mass fell in the narrow size range from 0.075 to 0.12/~m. Similar studies have been conducted at various locations around the world, including Belgium, The Netherlands, and the United Kingdom (e.g., see Van Vaeck and Van Cauwenberghe, 1978; and Baek et al., 1991). To better understand the effects of atmospheric processes (reactions, gas-particle partitioning, etc.) on the size distributions of PAHs in ambient aerosols, Venkataraman and Friedlander (1994b)carried out measurements of gases and particles during winter and

488


summer periods at several sites in the Los Angeles Air Basin. Concurrently, they determined the mass distributions of these PAHs in aerosol particles in vehicular emissions using the same instrumentation and analytical procedures (Venkataraman et al., 1994a). As seen in Fig. 10.15a, the distribution of fluoranthene in ambient aerosols collected during the winter is bimodal with a primary emissions mode at 0.05-0.5/xm (mode I) and an accumulation mode at 0.5-4.0 /xm (mode II). A significant fraction of the fluoranthene present in the total sample is present in the accumulation mode, II. Three other 4-ring, semivolatile PAHs sampled (pyrene, chrysene, and benz[a]anthracene) had similar mass distributions. In contrast, during the winter, the nonvolatile 5- and 6-ring PAHs BaP (Fig. 10.15b), benzo[b]fluoranthene, dibenzanthracene, benzo[ghi]perylene, benzo[k]fluoranthene, and indeno[cd]pyrene had ~ 6 3 - 8 2 % of their masses in the 0.05- to 0.5-/xm range, primary emission mode I. However, as seen in Fig. 10.15c, the pattern shifts significantly to larger sizes for aerosols sampled in the summer. In a similar study, Allen and co-workers (1996) determined the particle size distribution for 15 PAHs with molecular weights ranging from 178 (e.g., phenanthrene) to 300 (coronene) and associated with urban aerosols in Boston, Massachusetts. As for BaP in the winter (Venkataraman and Friedlander, 1994b), PAHs with MW >228 were primarily present in the fine aerosol fraction (Dp < 2 /xm). A study of 6-ring, MW 302 PAH at the same site showed bimodal distributions, with most of the mass in the 0.3- to 1.0-/xm particle size size range; a smaller fraction was in the "ultrafine" mode particles (0.09-0.14/xm) (Allen et al., 1998). For PAHs with MW 178--202, the compounds were approximately evenly distributed between the "fine" and "coarse" (Dp > 2 /xm) fractions. Polycyclic aromatic hydrocarbons in size-segregated aerosols col-

Q.

2.5

El 2.0 E~ O

.J

13

2.5

a Fluoranthene (Winter)

2.0

1.5

0~ 1"0 f 013 0.5 0.0 0.01

b

lected a month later at a rural site were present to a greater degree in the coarse fraction than those collected in urban Boston, consistent with other observations (e.g., Pierce and Katz, 1975; Van Vaeck and Van Cauwenberghe, 1985; Venkataraman and Friedlander, 1994b). These size distributions are consistent with the condensation of the large, nonvolatile PAHs on small particles during cooling of the exhaust. However, the smaller, semivolatile PAHs become distributed between the smaller and larger particles via continuing vaporization and condensation processes in the atmosphere. Furthermore, a number of studies have shown that the direct mutagenicity of particles is primarily associated with particles having Dp < 2.5 /xm. For example, Viras and co-workers (1990) report that at two sites in Athens, Greece, 81 and 92% of the total direct activities (TA98 - $ 9 ) w e r e associated with particles having Dp < 3.3 /xm. Furthermore, ~60 and 80%, respectively, were in submicron particles, Dp < 1.0 /xm. Similar results have been reported for the particle size distribution of promutagenic activity (TA98 + $9) of ambient particles collected by Pagano and co-workers (1996) near a busy road in Bologna, Italy. Figure 10.16 shows their data for fractions ranging from 33,000 pounds. e PM = particulate matter. f Two classes of on-road motor vehicles, light- and medium-duty trucks, that had relatively lower tons/day of exhaust emissions of PM are omitted for clarity. The three categories of light-duty trucks, noncatalyst, catalyst, and diesel, emitted a total of 1.43 and 0.99 tons/day of POM in 1982 and 1993, respectively. PM exhaust emissions totaled 0.04 and 0.05 tons/day in 1982 and 1993, respectively, for catalyst and noncatalyst, gasoline-fueled medium-duty trucks.

t h e w o r l d ( A t k i n s o n a n d A r e y , 1994; Ciccioli et al., 1996; A r e y , 1998a; s e e S e c t i o n F). Finally, f r o m the " e x p o s u r e " aspects of a h e a l t h risk assessment of ambient POM, recall that the data for c o n c e n t r a t i o n s o f P A H s a n d P A C s in t h e c o m p o s i t e

T A B L E 10.28

w h o l e s a m p l e a r e b a s i n w i d e a n n u a l a v e r a g e s f o r 1993. The actual short-term (and probably annual) averages o f c o n c e n t r a t i o n s o f h u m a n cell m u t a g e n s at e a c h o f t h e f o u r u r b a n s a m p l i n g sites l o c a t e d m a n y k i l o m e t e r s a p a r t (Fig. 10.23) c a n b e s i g n i f i c a n t l y d i f f e r e n t . F o r

Usage and Exhaust Emissions of Particulate Matter (PM) for Light-Duty Automobiles in California's South Air Basinb in 1982 and 1993 a Percentage of in-use light-duty automobiles

Gasoline light-duty automobiles Noncatalyst Catalyst Diesel light-duty automobiles

Percentage of vehicle miles traveled

Relative percentage of exhaust emissions of PM by light-duty automobiles

Emissions of PM by light-duty automobiles (tons / day)

1982

1993

1982

1993

1982

1993

1982

1993

44.2 53.9

8.0 90.5

33.1 63.9

4.3 94.4

28 27

12 48

1.46 1.41

0.27 1.06

1.9

1.5

3.0

1.3

45

40

2.33 E = 5.20

0.87 E = 2.20

a Calculated from data in California Air Resources Board "Motor Vehicle Inventory MVE17G Corrected" (1996); see data in Table 10.27. b Formal designation. For clarity and frame of reference, it is often referred to (e.g., in this chapter) as the Los Angeles Air Basin, or simply southern California Air Basin (see map Fig. 10.23).

502

10. AIRBORNEPAHs:CHEMISTRYANDTOXICOLOGICALIMPLICATIONS

example, as discussed in more detail shortly, Fraser et al. (1998) reported that the 48-h average concentrations of the reactive PAHs benzo[a]pyrene and cyclopenta[cd]pyrene in samples collected in September 1993 were much higher in central Los Angeles, 0.29 and 1.04 ng m -3, respectively, than in downwind Claremont, 0.03 and 0.03 ng m -3. Conversely, concentrations of 2-nitronaphthalene, a major atmospheric reaction product (see Section F.2) and mutagen in the MCL-5 human cell line (Sasaki et al., 1997b), averaged 3.81 ng m -3 in central Los Angeles but 8.11 ng m -3 at the downwind Claremont site.

4. B i o a s s a y - D i r e c t e d C h e m i c a l A n a l y s i s for Vapor-Phase and Particle-Phase PAHs and P A C s in A m b i e n t A i r U s i n g B a c t e r i a l Assays Since the initial discovery of the mutagenicity of air samples in 1977, most of the research has focused on genotoxicity of organic extracts of the airborne particulate matter. However, this emphasis has been shifting and there is increasing interest in, and concern over, the vapor-phase mutagenicity of urban atmospheres and its toxicological implications. We discuss here examples of some of the studies in which bacterial assays have been used to help in the identification of both gas- and particle-phase PAHs and PACs in ambient air. In one of the first studies of the vapor-phase mutagenicity of polluted urban air, Alfheim and co-workers (1985) collected both ambient particles and vapor-phase compounds and used the Salmonella typhirnurium reversion assay. The direct activities ( - $ 9 ) of the extracts generally exceeded the promutagenicities (+ $9), and furthermore, the vapor-phase mutagenicity ranged from 0 to 88% of the total activity. Subsequently, at several sites in Finland with different emission sources (using a similar sampling methodology) Pyysalo and co-workers (1987) found direct mutagenicity in both gas and particle phases with the S a l m o n e l l a / m i c r o s o m e test and activity in the sister chromatoid exchange (SCE) assay employing eukaryotic Chinese hamster ovary cells. In urban air in Helsinki, Tuominen and co-workers (1988) found the total direct and indirect activities to be about the same. Furthermore, the majority of the activity in each phase was due to polar compounds. Interestingly, the total SCE-inducing activity of unfractionated vapor-phase extracts was consistently higher than that of particlephase extracts. A major advance in this area came in 1992 when Harger and co-workers (1992) reported the direct mutagenicity of vapor- and particle-phase organics separated using HPLC in samples collected during the

summer of 1987 in Claremont, California (see map in Fig. 10.23). This HPLC procedure had several significant advantages over earlier studies. First, mutagenic compounds were separated from possibly interfering toxic compounds and hence were more likely to be detected. Second, a mutagen polarity profile (mutagram or "fingerprint") of the directly mutagenic gas-phase PAC was obtained that could be compared with the mutagram of the concurrently sampled particulate organics collected on the filter and with the chromatographic characteristics (e.g., polarity) of individual (or classes of) known or suspected direct mutagens. Figure 10.25 shows these mutagrams for the vapor and particle phases, respectively. Interestingly, the total direct mutagenicity of the vapor phase, 210 rev m -3, was actually greater than that of the particle phase, 160 rev m -3", furthermore, its mutagenicity profile was substantially different. Thus, fraction 4 is the major peak for the vapor-phase sample whereas most of the particle-phase mutagenicity is in the more polar peaks 6 and 7. Similar enhancements in the contributions of more polar species were reported for bioassay-directed fractionation of SRM 1649 urban air particulate matter (Schuetzle and Lewtas, 1986; Nishioka et al., 1988;

Vapor-phase

100

7_,= 210 rev m -3

80 6O {?, E

40

20 ,"4--' m (/) t-"(3

"-

100

~

80

c~r)

Particle-phase Z = 160 rev m 3

60 40 20

1

2

3

I~1 4

5

6

7

8

,

9

H P L C fraction n u m b e r

Increasing polarity

,~

FIGURE 10.25 Mutagen polarity profiles ("mutagrams") of the direct-acting mutagenicities of HPLC-separated fractions of dichloromethane extracts of polyurethane foam solid adsorbent (PUF, "vapor phase") and TIGF filter ("particle phase") samples collected concurrently in ambient air in Claremont, California, August 1987. Mutagen densities (rev m-3) on strain TA98, -$9, microsuspension modification of standard Salmonella typhimurium reversion assay (adapted from Harger et al., 1992).

D. BACTERIAL AND HUMAN CELL MUTAGENICITIES OF POLLUTED AMBIENT AIR

Lewtas et al., 1990a). Similarly, Greenberg and coworkers (1990, 1993) found that the polar neutral compounds in inhalable particle samples at Newark, New Jersey, were the greatest contributors to ambient mutagenicities, regardless of the season of the year. The importance of more polar compounds in the directacting TA98 mutagenicity has also been seen in the emissions of POM from heavy-duty diesel trucks and wood smoke (e.g., Schuetzle and Lewtas, 1986). (Note, however, that the structures, concentrations, and mutagenicity polarity profiles of diesel exhaust emissions can be strongly impacted by such parameters as fuel composition, engine type, age, operating conditions, control devices, etc.) A breakthrough in identifying possible particle-phase polar compounds responsible for the direct mutagenicity of fraction 6 in the Claremont, California, study of Harger et al. (Fig. 10.25) came with the isolation and identification by Helmig and co-workers (1992a, 1992b) of a new class of direct mutagens in particulate matter collected in Riverside, California, the nitrolactones of phenanthrene, specifically, the 2- and 4-nitro-6Hdibenzo[b,d]pyran-6-ones. The 2-nitro isomer (structure XI) is a very powerful, direct mutagen (240,000 rev /xg -1) and Harger et al. (1992) ascribe a major portion of the mutagenicity in their fraction 6 to this widely distributed product of the gas-phase OH radical attack on phenanthrene (see Atkinson and Arey (1994), Sasaki et al. (1995), Arey (1998a), and references therein; and Section F). Based on results from concurrent environmental chamber studies of the photooxidations of naphthalene, fluorene, and phenanthrene by Arey and coworkers (1992), Harger and co-workers attributed a significant fraction of the vapor-phase direct activity of fraction 4 (Fig. 10.25)to 2-nitronaphthalene (2-NN) present in these 12-h daytime samples as a result of an OH-initiated homogeneous gas-phase reaction of naphthalene in ambient air. Although 1-nitronaphthalene (1-NN) is also present at significant concentrations in both ambient and laboratory systems, it is a much weaker, direct mutagen (see Table 10.20; Arey et al., 1992). Harger and co-workers noted that methylnitronaphthalenes (MNN)were also abundant in southern California air, but their mutagenic activities in the microsuspension were unknown. Subsequently, Gupta and co-workers (1996) determined them for the 14 methylnitronaphthalenes (see Table 10.18) and used these data to quantify contributions of nitronaphthalenes and methylnitronaphthalenes to the vapor-phase mutagenicity of 12-h daytime and 12-h nighttime samples collected on PUF plugs at an air monitoring site in

100

~]

5 03

1- and 2-Nitronaphthalene

Daytime

Methylnitronaphthalenes

1~ = 23 rev m -3

80 60

40 E ~ID

20

ffl r

Nighttime

.i...,

100 rr

Z = 31 rev m -3

80 60

/// / / / / / /

40 20 1

2

3

4

5

6

7

8

9

H P L C Fraction n u m b e r Increasing polarity

F I G U R E 10.26 Mutagrams of direct mutagenicity on different HPLC fractions for composite 12-h daytime and nighttime vaporphase samples collected on PUF plugs in Redlands, California, in October 1994. Mutagen densities (rev m -3) on TA98, - $ 9 , microsuspension modification (adapted from Gupta et al., 1996).

Redlands, California, a receptor site ~ 110 km downwind from central Los Angeles. Bioassay-directed fractionations of the daytime and nighttime samples produced the mutagrams shown in Fig. 10.26. The ambient nitronaphthalenes and methylnitronaphthalenes in fraction 4 contribute ~18% of the total daytime mutagenic activity of 23 rev m-3; they are formed by OH radical attack on the parent PAHs (see Sasaki et al., 1995, and Section F). The total nighttime activity is higher than the daytime activity, 31 rev m -3, and was attributed to more efficient formation of the nitronaphthalenes and methylnitronaphthalenes in NO 3 radical initiated reactions (Atkinson and Arey, 1994). In summary, Gupta and co-workers (1996) reported that ambient concentrations of the vapor-phase nitronaphthalenes and methylnitronaphthalenes are generally greater than those of other particle-phase nitroPAHs present simultaneously (Arey et al., 1987; Wilson et al., 1995). This is particularly important in that, as noted earlier (Pitts et al., 1985c; Pitts, 1987), 2nitronaphthalene has been reported on intraperitoneal injection to be converted to 2-hydroxyaminonaphthalene, the same metabolite that 2-aminonaphthalene (Johnson and Cornish, 1978), an animal and human

5o4


carcinogen (IARC, 1987; U.S. DHHS, 1998), is converted to. Additionally, Tokiwa and Ornishi (1986) report that in feeding studies on animals, 2nitronaphthalene has shown carcinogenic activity. This would seem to raise relevant atmospheric chemistry/ inhalation toxicology questions. Finally, in the context of the overall vapor-phase mutagenicity of ambient air, we note that significant fractions of the two powerful human cell and bacterial mutagens discussed earlier, cyclopenta[cd]pyrene (XXVIII) and 2-nitrodibenzopyranone (XI), have been found in the gas phase (i.e., trapped on PUF plugs) in samples collected during hot weather at sites in southern California (Fraser et al., 1998, and Arey et al., 1994, respectively). Hence the contributions of such species, which are normally considered to be primarily in the particle phase, to the gas-phase mutagenicity at high ambient temperatures should also be considered.

E. ATMOSPHERIC FATES OF PARTICLE.ASSOCIATED PAHs: HETEROGENEOUS REACTIONS Atmospheric reactions of PAHs fall into two broad categories: (1) heterogeneous processes involving particle-associated compounds (including the semivolatile 4-ring and 5- and 6-ring PAHs), e.g., photolysis/photooxidation and gas-particle interactions, and (2) homogeneous gas-phase reactions of volatile 2- and 3-ring and semivolatile 4-ring PAHs initiated by OH (daytime) and NO 3 radicals (nighttime) and ozone (24 h). As we have seen, a great deal is known about emission sources and strengths, ambient levels, and mutagenic/carcinogenic properties of the particlephase PAHs in airborne POM. However, because of the tremendous physical and chemical complexity of the aerosol surfaces on which photolysis, photooxidations, and gas-particle interactions take place in "real" polluted ambient air, much less is known about the structures, yields, and absolute rates and mechanisms of formation of PAH and PAC reaction products, especially for the more polar PACs. This is one area in which there exists a major gap in our knowledge of their atmospheric chemistry and toxicology.

1. Background In 1954, Kotin and co-workers reported the "carcinogenicity of atmospheric extracts" of Los Angeles air. Subsequently, in 1956, they reported the carcinogenic activity of oxidation products of aliphatic hydrocarbons and, in 1958, of ozonized gasoline. Concurrently, Falk and co-workers (1956) published results

of their laboratory studies of photooxidations of PAHs. A decade later, Tebbens et al. (1966) and Thomas and co-workers (1968) reported that major decreases in the concentrations of BaP and perylene occurred in smoke that was irradiated while passing through flow chambers (e.g., a 60% decrease in the BaP content of soot). Based in part upon this pioneering research, a National Academy of Sciences Committee suggested in their 1972 document Particulate Polycyclic Organic Matter that photolysis, photooxidations, and gas-particle interactions might lead to significant degradation of PAHs in ambient particles and formation of products more polar than the parent PAHs (NRC, 1972). Shortly thereafter, laboratory studies of the photooxidation and ozonolysis of PAHs deposited on several substrates (e.g., glass surfaces and thin-layer chromatography plates) and exposed to sunlight a n d / o r various levels of ozone in air found significant differences in the reactivities of individual PAHs. Half-lives of BaP and benz[a]anthracene were short ( ~ 1 h or less), benzo[e]pyrene (BeP) and pyrene intermediate, and benzofluoranthenes long (Lane and Katz, 1977; Katz et al., 1979). Among the reaction products were three BaP diones that had previously been identified in ambient aerosols collected in Toronto, Canada (Pierce and Katz, 1976), and subsequently at Duisburg, West Germany (K6nig et al., 1983), and Elverum, Norway (Ramdahl, 1983b). During the same time frame, extracts of ambient aerosols collected throughout the world were found to be directly mutagenic in the Ames bacterial assay. Concurrently, BaP deposited on various substrates (e.g., glass fiber filters) and exposed to ppm and sub-ppm levels of ozone was shown to react readily to form direct-acting frameshift mutagens (Pitts et al., 1978). Furthermore, exposures of particle-associated 4- and 5-ring PAHs to similar levels of NO 2 (and a required "trace of gaseous HNO3") in air resulted in significant yields of nitro-PAHs; these products, which reflected a range of reactivities of the parent compounds, were also direct-acting frameshift mutagens (J~iger, 1978; J~iger and Haniis, 1980; Pitts et al., 1978; Pitts, 1979; Hughes et al., 1980; Tokiwa et al., 1981; Nielsen et al., 1983; vide infra). For example, an 8-h exposure of BaP deposited on a glass fiber filter to 250 ppb of NO 2 (the California air quality standard) and traces of HNO 3 gave an ~20% yield of products, predominantly 6NO2-BaP with lesser amounts of the 1- and 3-isomers. Soon thereafter, mutagenicity studies by Wang and co-workers (1980) implied the presence of nitroarenes in ambient air. Subsequently, several nitroarenes were isolated and positively identified by Ramdahl and coworkers (1982b). We note, however, that these initial researchers cautioned that their rate, mechanism,

E. ATMOSPHERICFATESOF PARTICLE-ASSOCIATEDPAHs:HETEROGENEOUSREACTIONS product, and mutagenic data from laboratory exposures of PAHs deposited in several ways on a variety of substrates (evaporation of a BaP solution on a Hi-Vol filter, vapor deposition of PAHs on synthetic and real particle substrates, etc.) and exposed to relatively high levels of simulated sunlight and 0 3 and NO 2 in clean air may not be truly representative of "real world" heterogeneous reaction systems. We discuss in this section four key aspects of heterogeneous reactions: (1) theoretical and experimental structure and reactivity relationships; (2) field measurements of relative and absolute PAH decay rates in near-source ambient air and during downwind transport; (3) laboratory studies of the photolysis/photooxidation and gas-particle interactions with 0 3 and NO 2 of key 4- and 6-ring PAHs adsorbed on model substrates or ambient aerosols; and (4) environmental chamber studies of the reactions of such PAHs associated with several physically and chemically different kinds of combustion-generated aerosols (e.g., diesel soot, wood smoke, and coal fly ash). Where such data are available, we also briefly consider some toxicological ramifications of these reactions. Space considerations preclude extensive discussions of these broad topics; for details, see reviews and critiques by Nielsen and co-workers (1983), Van Cauwenberghe and Van Vaeck (1983), Pitts (1983), Van Cauwenberghe (1985), Bje~rseth and Olufsen (1983), Valerio et al. (1984), Nikolaou et al. (1984), Finlayson-Pitts and Pitts (1986), Pitts (1987, 1993a, 1993b), Baek et al. (1991), Atkinson and Arey (1994) (gas-phase reactions only), and Arey (1998a).

2. Theoretical and Experimental Structure- Reactivity Relationships In their thoughtful 1983 review, Nielsen and coworkers noted that particles of diesel soot or wood smoke can absorb significant amounts of water. Thus, they suggested that the most plausible mechanism(s) for nitration (and possibly other electrophilic reactions) of particle-associated PAHs in ambient air may involve reactions both in a liquid film and on solid surfaces and that fundamental laboratory studies of the rates, products, and mechanisms of PAHs in polar solvents would be atmospherically relevant for reactions in the liquid films. Based on this, they proposed a classification scheme for the reactivities of key PAHs in electrophilic reactions, which was subsequently described in detail (Nielsen, 1984). Given the complexities of the system involved, it is interesting that this reactivity scale has proven to be a useful predictor not only for overall relative decay rates of PAHs associated with aerosols in ambient air but

5 05

also for the relative rates of their electrophilic heterogeneous reactions in simulated atmospheres. The original suggestion of Nielsen and co-workers (1983) that these heterogeneous processes involved particle-phase PAHs interacting with other compounds, organic and inorganic, present in "liquid" (or quasiliquid) layers on the surfaces of carbonaceous particles (e.g., wood smoke) is consistent with results from subsequent studies. For example, Strommen and Kamens (1997) cite as evidence for this phenomenon the fact that a viscous, oily slick, with little particle integrity, results when particles of diesel or wood soot impact a metal foil (McDow et al., 1994). Furthermore, Kamens, McDow, and co-workers have demonstrated the strong influence of aerosol water content (humidity), solvent effects, and dissolved aerosol constituents on such processes as the photodegradation of benz[a]anthracene (vide infra; also see Odum et al., 1994a; McDow et al., 1994, 1995, 1996; Jang and McDow, 1997; and Jang and Kamens, 1998). The diffusion model of Odum and co-workers (1994b), which describes the gas-particle partitioning of semivolatile organics (including PAHs), incorporates the idea of diffusion of the SOCs into viscous, liquid-like, organic layers on particles of diesel soot. Subsequently, Strommen and Kamens (1997) extended this concept with a dual-impedance model that incorporated rapid diffusion into an outer layer of liquid-like organic material and slower diffusion into an inner core of discrete solids. We discuss in the following the implications of such liquid-like organic layers to the rates and mechanisms of the photochemical and gas-particle reactions of particle-associated PAHs and nitro-PAHs (e.g., see Fan et al., 1996a, 1996b; and Feilberg and Nielsen, 1999b). In developing the reactivity scale, Nielsen first investigated the transformation rates of 25 PAHs and four derivatives of anthracene in water-methanol-dioxane solutions, taken as a model of "wet particles," and containing small amounts of dinitrogen tetroxide and nitric and nitrous acids. The measured half-lives and relative rates are shown in Table 10.29. The range of reactivities in solution for PAHs of different structures is remarkable, from 100,000 (arbitrarily set) for anthanthrene (XXXlI) to 300 days for benzo[b]fluor a n t h e n e and > 8 0 0 days for indeno[1,2,3-cd]pyrene, with the relative reactivities decreasing from 1100 to

BaP

XXXIII

This product suggests a singlet oxygen mediated mechanism for quinone formation (Lee-Ruff et al., 1986). Subsequently, Lee-Ruff and Wang (1991) conducted a similar study of 6-methylbenzo[a]pyrene, a methyl isomer whose carcinogenicity is approximately equal to that of BaP. It photooxidized ~20 times faster than BaP and, in addition to quinones, formed as a major product a seco ketone analogous to XXXIII from BaP. Its formation (as with BaP) is ascribed to a 102-mediated mechanism. Since the mid-1980s, Kamens, McDow, and coworkers have carried out a number of studies on factors influencing the rates, mechanisms, and products of

5 11

the photodegradation of PAHs. They employed both model laboratory systems in which solutions containing various organics found in ambient aerosols are irradiated in photoreactors and environmental chamber studies in which real organic combustion aerosols relevant to ambient atmospheres, e.g., diesel soot and wood smoke, are irradiated with natural or simulated sunlight under a variety of ambient conditions. For example, Kamens and co-workers (1988) reported that daytime photodegradation rates of PAHs on wood smoke and gasoline combustion soot irradiated by natural sunlight in 25-m 3 outdoor Teflon film chambers were clearly influenced by such variables as solar intensity, relative humidity, and ambient temperature. Thus, under "moderate conditions" (e.g., T = 20~ "moderate" humidity, and a solar intensity of 1 cal cm -2 min-1), PAH half-lives for the most reactive PAHs (Class II, Table 10.30), cyclopenta[cd]pyrene and BaP, were 0.3 and 0.5 h, respectively, compared to 0.8 and 1.3 h for the low-reactivity (Class V) benzofluoranthenes (k and b). At very low values of sunlight intensity, water vapor concentrations, and ambient temperatures (-10~ half-lives increased significantly; e.g., for cyclopenta[cd]pyrene and BaP, the half-lives increased to 2 and 6 h, respectively, and for the benzofluoranthenes to ~ 10 h. These PAH relative rates are generally consistent with the Nielsen reactivity scale (Nielsen, 1984; Table 10.30). Note that McDow and co-workers (1995) reported that the water content of combustion particles of wood smoke and gasoline soot increased with increasing relative humidity and that PAH photodegradation rates are likely to increase with increasing water content (BaP was an exception) (see also a recent modeling study of humidity effects by Jang and Kamens (1998)). McDow, Kamens, and co-workers also conducted laboratory experiments on the effects of common organic constituents (e.g., methoxyphenols) on the rates, mechanisms, and products of the solution-phase photodegradation of PAHs associated with wood smoke and diesel soot (see, for example, Odum et al. (1994a), and McDow et al. (1994, 1995, 1996)). Figure 10.28, for example, shows the degradation of the reactive BaP (Class II reactivity) compared to BeP (Class V reactivity) in two solvents, hexadecane, taken as representative of aliphatics in diesel soot, or a mixture of 11 methoxyphenols found in particulate matter (McDow et al., 1994). As expected, BaP decays much more rapidly than BeP. In addition, the decay in the mixture of methoxyphenols is much faster than that in hexadecane.

512

10. AIRBORNE PAHs: CHEMISTRY A N D TOXICOLOGICAL IMPLICATIONS

,-- 100%

00%~*"

.m t'.m

_,

_,

_,

E d2

~E d2

50% I

50%

d2

-

BaP

13_

0

I

0

I

100

I

I

I

I

I

I

200 300 400 Irradiation time (minutes)

0

0

100 200 300 400 Irradiation time (minutes)

9 Hexadecane 9Methoxyphenols FIGURE 10.28 Rates of photodegradation of benzo[a]pyrene (Class II reactivity; see Table 10.30) and benzo[e]pyrene (Class V) in irradiated solutions of hexadecane or a mixture of methoxyphenols that are representative of important classes of organics present in particles of diesel soot and wood smoke, respectively (adapted from McDow et al., 1994).

OH OCH3

organic constituents (e.g., methoxyphenols) and on properties such as oxygen solubility and 10 2 lifetimes in the liquid-like layer of organic combustion aerosols. b. Photooxidations on Inorganic Solid-Air Surfaces

C

/\\ H

O Vanillin

Similarly, Jang and McDow (1995, 1997a) investigated the rates, products, and mechanisms of the photooxidation of BaA carried out in three solvents (toluene, benzene, and benzene-d 6) that contained three common constituents of ambient aerosols that are known accelerators of the photodegradation process: 9,10-anthraquinone, 9-xanthone, and the substituted methoxyphenol, vanillin. On the basis of this and previous studies, the researchers proposed "that there were at least two mechanisms" (see also Jang and McDow, 1995). One is singlet oxygen mediated and observed with BaA dissolved in benzene and benzene-d 6 solvents; it involves 10 2 addition to an aromatic ring of BaA followed by ring opening and products such as phthalic acid. The second mechanism is most important in toluene; it features H-atom abstraction from the solvent by aromatic carbonyl compounds, leading to complex free radical initiated oxidations. Jang and McDow (1995) have also suggested as other possible mechanisms for the photooxidation of PAH triplettriplet energy transfer to the PAH by a photosensitizer (e.g., polycyclic aromatic quinones or ketones and aromatic aldehydes) or free radical chain oxidations initiated by electronically excited PAH. In short, efficiencies of PAH photooxidation processes may depend in part on the H-atom-donating properties of some of the

Just as with organic combustion aerosols, the chemical and physical nature of inorganic solid substrates can have a dramatic impact on the photoreactivity of adsorbed PAH. In 1980, Korfmacher and co-workers reported that BaP, pyrene, and anthracene all photolyzed efficiently in liquid solution but were resistant to photodegradation when adsorbed on coal fly ash. Subsequent studies confirmed this observation and revealed that the carbon content of the ash (and the associated darkening of color) is a key factor in establishing the photostability of these PAHs. Indeed, they were stabilized at relatively small percentages of carbon, e.g., 5% or less (Behymer and Hites, 1985, 1988; Yokley et al., 1986; Dunstan et al., 1989; Miller et al., 1990). PAHs adsorbed on particles of carbon black were also photostabilized (Behymer and Hites, 1988). However, Barofsky and Baum (1976) demonstrated that BaP, anthracene, BaA, and pyrene deposited on carbon microneedle field desorption emitters and exposed to UV radiation were all photooxidized to carbonyl compounds. Similarly, PAHs can photodegrade efficiently in air when adsorbed to substrates of silica gel, alumina, or glass plates (e.g., see Lane and Katz, 1977; Kormacher et al., 1980; Behymer and Hites, 1985; Yokely et al., 1986). The effect of the physical state of BaP and perylene adsorbed on fused-silica plates on their reaction rates with ozone in air was studied by Wu and co-workers (1984), who measured the fluorescence of the two PAHs

E. A T M O S P H E R I C FATES OF PARTICLE-ASSOCIATED PAHs: HETEROGENEOUS R E A C T I O N S

as functions of concentration, reaction time, and 0 3 concentration. Interestingly, their reaction rates in aggregated states (inferred from eximer formation)were much slower than those in the dispersed state (only monomer fluorescence). They concluded that the reaction of 0 3 with aggregated PAH molecules was much slower than with monomers on the surface. AlebidJuretid and co-workers (1990) reported a similar observation for the PAH-silica gel-O 3 reaction. Such a "surface/bulk" reactivity phenomenon may in part be responsible for the low (or zero) reactivity reported for BaP deposited on, or present in, a variety of substrates and exposed to ambient levels of 0 3 (e.g., 100 ppb) in air (see Grosjean et al., 1983; and Coutant et al., 1988). These observations can be rationalized by assuming that, while BaP in fact does react rapidly with 0 3 in ambient particles, not all of it is at (or close enough to) the surface to be available for reaction (Atkinson et al., 1988a; Arey, 1998a). In the early 1990s, a series of articles appeared that described the rates, mechanisms, and products of the photooxidations of several simple gas-phase PAHs adsorbed on the surfaces of SiO 2 and A120 3 particles. These are models for some of the inorganic particles found in air (see Chapter 9). Studies of naphthalene and 1-methylnaphthalene (Barbas et al., 1993), acenaphthylene (Barbas et al., 1994), anthracene (Dabestani et al., 1995), phenanthrene (Barbas et al., 1996), and 1-methoxynaphthalene on SiO 2 (Sigman et al., 1996; see also below and the reviews by Dabestini (1997) and Sigman (1997)) in the presence and absence of air and using different photolytic wavelengths were carried out. Sigman and co-workers (1997) and Dabestini (1997) reported studies of the photochemistry of 2- and 3-ring PAHs adsorbed on SiO 2. Photooxidations were shown to proceed by two distinct mechanisms. Type I involves electron transfer from the photoexcited singlet state (S 1) of the PAH to atmospheric oxygen, forming superoxide, 0 2- (or transfer to "surface-active sites"). Indeed, Barbas and co-workers (1993) reported the room temperature EPR spectrum of 0 2- when naphthalene adsorbed on the surface of SiO 2 is irradiated. The type II mechanism is a singlet oxygen mediated, photooxidation process such as that described earlier for BaP in liquids--that is, triplet-triplet energy transfer from the T1 state of an electronically excited PAH molecule to the triplet ground state of an 0 2 molecule (3s producing O2(~zX), which then adds to the ground-state PAH, ultimately forming oxidized photoproducts. PAHs photooxidized exclusively by the type I electron transfer-superoxide mechanism include naphthalene and 1-methylnaphthalene (Barbas et al., 1993), fluorene (Barbas et al., 1997), and acenaphthene (Re-

513

yes et al., 1998). Typical products formed in the type I photooxidation of naphthalene and 1-methylnaphthalene are shown in Fig. 10.29A. PAHs photooxidized by the type II singlet oxygen mediated mechanism include acenaphthylene, whose oxidized products and yields are shown in Fig. 10.29B (Barbas et al., 1994), phenanthrene (Barbas et al., 1996), anthracene (Dabestini et al., 1995), and tetracene (Dabestini et al., 1996). An additional photochemical process, the formation of photodimers, is also observed for acenaphthylene, anthracene, and tetracene. Interestingly, Barbas and co-workers (1996) noted that, once sorbed to the SiO 2 surface, phenanthrene did not desorb into the gas phase even under high vacuum. As in previous studies of the PAHs cited above, no dark reactions were observed at the SiO2-phenanthrene interface. The authors suggested that the photostability of phenanthrene adsorbed on carbon black (half-life > 1000 h) and fly ash (half-life ~49 h) reported by Behymer and Hites (1985) may be due to a number of factors, including competitive absorption of incident radiation, energy transfer from excited phenanthrene to other PAHs in carbon black and "more facile reactions of singlet oxygen with the substrate."

5. G a s - Particle Reactions a. Ozonolysis

Lane and Katz reported in 1977 that the dark reaction of BaP deposited on the surface of glass Petri dishes with air containing 200 ppb of ozone was fast, with a half-life of ~38 min. Katz and co-workers (1979) exposed nine PAHs on thin-layer chromatography plates of cellulose in the dark to 200 ppb of 0 3 in air and found pronounced differences in their reactivities, e.g., half-lives of 36 rain for BaP, 2.9 h for BaA, 7.6 h for BeP, and 53 h for benzo[b]fluoranthene. Subsequently, in good agreement with Lane and Katz, a half-life of ~ 1 h was determined for BaP deposited on glass fiber filters and exposed (passively in a controlled atmosphere) to 200 ppb of 0 3 in the dark (Pitts et al., 1980). Pitts et al. (1986) exposed five individual PAHs, pyrene, fluoranthene, benz[a]anthracene, BeP, and BaP, deposited on glass fiber and Teflon-impregnated glass fiber filter (TIGF) substrates "passively" for 3 h in the dark in a 360-L Teflon environmental chamber to 50-300 ppb of O 3 in air at several relative humidities. These experimental conditions more nearly resemble the actual exposure of ambient particles to O 3 (in the dark) during transport than do exposures in Hi-Vol flow systems. Consistent with earlier studies, BaP, BaA,

514

10. AIRBORNE PAHs: CHEMISTRY A N D TOXICOLOGICAL IMPLICATIONS

hv > I ~

CO2H CO2H

SiO2

CH3 I

COCH 3

CO2H

SiO2

A Type I: Electron transfer

O--H, C

, H--O

/o\

C

O=C

[ hv SiO2

~ ~

O dicarb~

C=O

@ 3"8% 1;8~aphth~..~onhydride

Acenaphthylene

4.6%

2-Hydroxy-1-Acenaphthenone

1.7%

1,2-Acenaphthenedione

B Type I1: Singlet oxygen mediated addition

FIGURE 10.29 Typical photoproducts observed (A) in the irradiation (a = 300 nm) in air of naphthalene and 1-methylnaphthalene adsorbed on silica and formed by a Type I electron transfer (superoxide) mechanism (Barbas et al., 1993) and (B) in the irradiation in air of acenaphthylene (a = 350 nm) adsorbed on silica formed by a Type II singlet oxygen mechanism (Barbas et al., 1994) (adapted from Dabestani, 1997).

and pyrene were the most reactive ( ~ 6 0 - 8 0 % degraded at 1% RH) and, except for BaP, degradation was much less at 50% R H than at 1% RH. Fresh ambient particulate P O M sampled near a freeway was also exposed in this passive system for 3 h in the dark to 200 ppb of 0 3 in air at 1% RH. Concentrations of specific PAHs determined in the ambient particles and their percent reacted were similar for samples collected on both kinds of filters (glass fiber and TIGF). Again, under passive exposure conditions to approximately ambient levels of 0 3, BaP and BaA were found to be significantly more reactive than BeP. Results of a study of the ozonolysis of primary combustion aerosols by Van Vaeck and Van Cauwenberghe (1984a) are illustrated in Fig. 10.30. Shown are percent conversion profiles as a function of time for the decay of several 5- and 6-ring PAHs in diesel exhaust

particulate matter (Dp < 0.5 /xm) collected on glass fiber filters and exposed in the dark to 1.5 ppm 0 3 in air under Hi-Vol sampling conditions. Half-lives for degradation range from ~ 1 5 - 2 0 min for BaP to 2-3 h for low-reactivity benzo[k]fluoranthene. The authors noted that experiments at ambient levels of 0 3 of 100 ppb also "include significant conversion." Similarly, Lindskog and co-workers (1985) reported several PAHs in soot particles that were generated in a smoke gas generator and collected on a glass fiber filter were "transformed" when exposed to ozone. Overall, such results suggest that in ambient aerosols, the highly reactive PAHs, e.g., BaP and cyclopenta[cd]pyrene, could be susceptible to ozone degradation under certain atmospheric conditions, as well as during Hi-Vol collection of samples. Van Cauwenberghe (1985) proposed two different mechanisms for the degradation of unsubstituted PAHs

E. ATMOSPHERIC FATES OF PARTICLE-ASSOCIATED PAHs: HETEROGENEOUS REACTIONS

by ozone; one is a one-step electrophilic-nucleophilic attack on olefin bonds with high electron density. This forms an unstable primary ozonide that decomposes, with ring opening giving products with aldehyde a n d / o r carboxylic acid groups. Peroxides or a-keto hydroperoxides may also be formed. The second mechanism is a two-step electrophilic attack involving initial formation of a cr complex and ultimately quinones. Complex mixtures of oxidized products are produced in the ozonolysis of BaP on solid substrates. For example, BaP deposited on a glass fiber filter and exposed to 1 ppm 0 3 in air primarily formed products characteristic of the ring-opening mechanism, e.g., dialdehydes and dicarboxylic and ketocarboxylic acids. However, some quinones and phenols, were also formed. For details on the mechanisms and O-PAH products of these BaP-O3, gas-solid substrate reactions, see Van Cauwenberghe et al. (1979) and Van Cauwenberghe and Van Vaeck (1983 and references therein). Mixtures of products formed in the reactions of BaP deposited on glass fiber filters with 100-200 ppb of 0 3 in air are not only complex chemically, but their extracts are also directly mutagenic in the Ames assay (strain TA98, - $ 9 ) . Using the method of bioassay-directed fractionation and chemical analysis, as illustrated in Figs. 10.31a to 10.31d, Pitts and co-workers (1980) isolated, identified, and obtained the specific activity, 1600 rev /xg -1 (TA98, - $ 9 ) , of the major stable contributor to the activity of such mixtures, benzo[a]pyrene-4,5-oxide (Fig. 10.31d). It is a wellknown BaP metabolite, a powerful direct-acting frameshift mutagen, and a weak carcinogen on mouse skin (Levin et al., 1976). The value of 1600 rev/xg -1 is in good agreement with those of other researchers

100 [

• BaP

~

I 50

0

.

.

.

.

.

.

.

.

.

.

.

15 30

.

.

.

.

A gghiP O BeP

~ .

.

.

.

.

.

.

60

.

.

.

.

.

.

.

O,no .

.

.

.

90

.

.

.

.

.

.

.

.

.

.

120

180

240

Time (min)

FIGURE 10.30 Percent conversion-time profiles for the decay of 5 PAHs in diesel exhaust particulate matter (Dp = 0.5/xm) collected on glass fiber filters and exposed to 1.5 ppm of ozone in air under Hi-Vol sampling conditions. Half-lives (dotted line) decrease in order of the Nielsen (1984) electrophilic reactivity scale (Table 10.30): BaP, benzo[a]pyrene; BghiP, benzo[ghi]perylene; BeP, benzo[e]pyrene; IndP, indeno[1,2,3-cd]pyrene; BkF1, benzo[k]fluoranthene (adapted from Van Vaeck and Van Cauwenberghe, 1984).

515

(e.g., Wislocki et al., 1976, ~ 1370 rev/xg-1; Chang et al., 1979). Extrapolation of results from the ozonolysis or photooxidation of BaP deposited on solid substrates (e.g., filters) to possible reaction rates and products for BaP sorbed on or in ambient aerosols in real-world polluted atmospheres should be done with caution. For example, the yield of BaP epoxides is low (0.5%) for exposures in the dark and even less when the BaP-coated filter is exposed concurrently to actinic light (0.07%). However, this reaction, forming a BaP metabolite, is interesting per se and may have significant atmospheric and toxicological relevance if it were to take place on the surfaces of fine respirable ambient aerosols (Pitts et al., 1980; Pitts, 1983). Van Cauwenberghe and Van Vaeck (1983) noted that because of its increased solubility in lung tissues, its pulmonary absorption may be enhanced over nonpolar BaP. Similar considerations might apply to as yet unidentified products of the heterogeneous reactions of ozonolysis and photooxidation of, for example, the highly reactive and powerful human cell mutagen cyclopenta[cd]pyrene (XXVIII). b. Nitration

Certain 4- to 6-ring PAHs, deposited on a variety of substrates and exposed to a range of concentrations of gaseous NO 2 in the presence of trace amounts of H N O 3 near ambient levels, react to form mono- and dinitro-PAHs (Pitts et al., 1978; Pitts, 1979; J~iger and Hanus, 1980; Tokiwa et al., 1981). In the first product study in a laboratory-simulated atmosphere, ~ 2 0 % of the BaP deposited on a GF filter and exposed for 8 h to air containing 0.25 ppm of NO 2 and traces of gaseous H N O 3 was converted to 6-NO2-BaP; smaller amounts of the 1- and 3-isomers were also formed (Pitts et al., 1978; Pitts, 1979). Interestingly, in this heterogeneous process, a promutagen was converted to a stronger promutagen, 6-NO2-BaP, and two powerful direct mutagens, the 1- and 3-NO 2BaP isomers (TA98, - $ 9 ; see Table 10.16), were formed. The generality of this reaction of NO 2 with particleassociated PAHs was demonstrated when, under similar laboratory conditions, perylene (a weak promutagen) was converted to 3-nitroperylene, and pyrene (a nonmutagen), at a slower rate, to 1-nitropyrene; both nitro-PAHs are direct-acting frameshift mutagens. In a similar experiment, chrysene was not nitrated; this is consistent with its low reactivity in the Nielsen reactivity scale (Table 10.30). Tokiwa et al. (1981) reported that direct-acting mutagenic derivatives were formed when pyrene, phenanthrene, fluorene, fluoranthene, and chrysene (and carbazole) were deposited on filter paper and exposed to

516


80

a ~1~

BaP quinones m ~ Fraction13

[

70

BaP ~

60 ca. 50 c-

40

30 z

20 10

0

4

8

12

16

20 24 Time (min)

28

32

36

3.3092.5 ", 2 1 2 9n. .~. 0. 9 1 4 09 6 198 290 193 0 96 0 49 2 .90 9 BaF 1 2 3 4 5 6 7 8 9 1011 1 2 1 3 1 4 1 5 1 6 1 7 HPLC fraction number

13-4 Strongly mutagenic (direct)

200

m Q..

j

Q. if) r-

13-1

~Jl

13-2

] N.A.

Ill

~N.A.

/

z

~

I

0

I

I

4

I

I

8

I

I

12

100

I[

.5A.

I

I

16

I

I

20

Time (min)

/I

l[

/

0.01

I

0.03

I

0.05

I

0.08

I

0.10

9

gg sample per plate

F I G U R E 10.31 (a) Preparative HPLC separation (acetonitrile-water solvent) of products of the dark reaction of BaP deposited on a glass fiber filter with ~ 200 ppb 0 3 in air (asterisk marks scale change). (b) Specific direct-acting mutagenic activity of HPLC fractions (rev/xg - ] ) on TA98, - $ 9 . (c) Preparative HPLC separation of mutagenic fraction 13. (d) Dose-response curve for fraction 13-4 (TA98, -$9), BaP-4,5-oxide, whose specific activity is 1600 rev/xg-] (adapted from Pitts et al., 1980). N.A., no activity.

10 ppm NO 2. The 24-h nitration yields for pyrene rose from 0.02% with 1 ppm NO 2 in air to 2.85% when traces of H N O 3 were present. J~iger and Hanus (1980) found the order of reactivity for the reactions of PAHs adsorbed on several substrates with 1.3 ppm of NO 2 in air to be silica gel > fly ash > deactivated aluminum oxide > carbon. The qualitative composition of nitro-PAH products, however, was independent of the substrate. Subsequently, Ramdahl and co-workers (1984b) adsorbed six PAHs on several substrates and exposed them to 0.5 ppm of NO 2 in air containing water vapor and traces of H N O 3. The highest reactivity was observed for PAHs adsorbed on silica. The yields of three nitro-PAHs detected on alumina were only 14-24% of

those on silica. The relative PAH reactivity order was perylene > benzo[a]pyrene > pyrene > chrysene > fluoranthene = phenanthrene, similar to that found by other researchers in their solution-phase studies (e.g., see Dewar et al., 1956; and Nielsen, 1984). A similar order for relative reactivity was observed by Butler and Crossley (1981) for the loss of PAH deposited on soot particles and exposed to NO 2 in air. Guo and Kamens (1991) describe a system for studying gas-particle reactions on the surfaces of combustion aerosols in which they report a half-life of ~ 80 h for "high loadings" of particle-bound BaP in wood smoke particles reacting with ~200 ppb of NO 2 in air. Relatively small amounts of gaseous nitric acid in the NO2-air mixtures used to nitrate particle-phase

517

E. A T M O S P H E R I C F A T E S O F P A R T I C L E - A S S O C I A T E D P A H s : H E T E R O G E N E O U S R E A C T I O N S

PAHs on various substrates in simulated ambient air or combustion systems appear to play an important role in the reactivity (e.g., see Pitts et al., 1978; Pitts, 1979; Hughes et al., 1980; Lindskog et al., 1985; and Yokley et al., 1985). Grosjean and co-workers (1983) found no reaction of nitric acid free NO 2 in air with BaP, perylene, and 1-nitropyrene deposited on several substrates. Although Wu and Niki (1985) did not report an important role for HNO 3 in their spectroscopic study of the reaction of NO 2 with pyrene deposited on a silicon surface, acidity could have been provided to some extent by the silica surface. Thus, once again, substrate complexities may be involved. The nitration of PAHs by N O 2 / H N O 3 also occurs under laboratory conditions approximating plume gases, that is, higher concentrations of gases and deposition on coal fly ash as a substrate. Thus, Hughes and co-workers (1980) reported that BaP and pyrene reacted with 100 ppm of NO2; the presence of nitric acid (possibly on the surface of the fly ash)enhanced the rate of reaction. Reactions proceeded more rapidly on silica gel than fly ash substrates, and for pyrene, both mono and dinitro isomers were formed. At the 100 ppm plume gas level, neither NO nor SO 2 reacted with BaP or pyrene on the substrates studied; both PAHs reacted with SO3, but products were not characterized. In a study simulating stack gas sampling, Brorstram-Lund6n and Lindskog (1985) found that addition of NO 2 caused substantial degradation of the reactive PAHs present in soot generated from a propane flame. A strong enhancing effect was observed when gaseous HC1 was added to the laboratory stack gases, for example, 90% loss of BaP in 1 h with added HC1 compared to 20% in its absence. They suggested that, in addition to other processes that might occur in the hot effluent stream, under these acidic conditions, sampling artifacts may be a major problem in sampling stack gases, often in times as short as 15-30 min or less. The possibility of artifactual formation of nitroPAHs during the sampling of diesel exhaust was addressed soon after their discovery in diesel particles (e.g., Lee et al., 1980; Lee and Schuetzle, 1983). Schuetzle (1983) concluded that artifactual formation of nitro-PAHs "is a minor problem" (between 10 and 20% of the measured 1-nitropyrene) at short sampling times [e.g., 23 min, which is one federal test procedure (FTP) driving cycle], at low sampling temperatures (42~ and in diluted exhaust containing NO 2. The question of formation of nitroarenes during Hi-Vol sampling of ambient POM was considered in early studies (Pitts et al., 1978; Pitts, 1979) and addressed in several of the studies of PAH nitrations discussed above. In a definitive evaluation, Arey and co-workers (1988a) coated several perdeuterated PAHs

(e.g., perdeuterofluoranthene, perdeuteropyrene, and perdeutero-BaP) onto Hi-Vol filters loaded with previously collected ambient POM and exposed them for 7-10 h to ambient air during a high-NO X episode in southern California. Less than 3% of the total 1-nitropyrene collected during the episode was formed in the sampling process, and no formation of nitrofluoranthene was observed. Hence the authors concluded artifactual formation of nitroarenes during Hi-Vol sampling (e.g., the nitrofluorenes and nitropyrenes) is not significant (see Arey et al., 1988a, for references to other studies). Overall, while the combinations of substrate effects, ambient NO 2 levels, and other gas-particle phenomena preclude a definitive answer, the formation of significant amounts of nitroarenes in heterogeneous panicle-phase NO2-PAH, atmospheric reactions seems unlikely, e.g., much slower than photooxidation or ozonolysis. This conclusion also applies to heterogeneous reactions of N20 5 with particle-bound PAHs on diesel and wood soot (Kamens and co-workers, 1990; see also Pitts et al., 1985c, 1985d, 1985e). 10

1

1

2

6

5

2

8 7

7 6

5

XXVIII Cyclopenta[cd ]pyrene

Acenaphthylene

An exception may be particle-associated cyclopenta[cd]pyrene (XXVlII), which, with its exocyclic double bond (analogous to acenaphthalene), may react at a significant rate with the gas-phase NO 3 radical (or N20 5) and even faster in heterogeneous photooxidations or ozonolysis. Certainly, by analogy with acenaphthylene, whose atmospheric lifetimes under typical atmospheric conditions are 6 min for NO 3 radicals, 43 min for 03, and 1.6 h for OH radicals (see Atkinson and Arey, 1994), gas-phase cyclopenta[cd]pyrene would also be expected to decay rapidly by these reactions in nighttime and daytime polluted air environments. Recall that Fraser and co-workers (1998) found that over 50% of the cyclopenta[cd]pyrene was in the gas phase (see Figs. 10.2 and 10.27) during hot weather. As noted earlier, the rates of formation, mechanisms, chemical structures, and yields of products, as well as the mutagenic activities, bacterial and human cell, of such cyclopenta[cd]pyrene reactions are currently unknown but constitute an interesting and relevant research challenge to atmospheric chemists and toxicologists alike.

518


6. A t m o s p h e r i c Fates of Particle-Associated Nitroarenes As we have seen, key nitroarenes found in extracts of ambient particulate matter are 1-nitropyrene (1NO2-Py), predominant in primary combustion emissions, and 2-nitrofluoranthene and 2-nitropyrene, major products of gas-phase atmospheric reactions. Here we focus simply on their atmospheric fates as particlebound species participating in heterogeneous decay processes. Formation of such nitro-PAHs in gas-phase reactions is addressed in Section F.

Subsequently, Ioki (1977) used ESR spectroscopy to confirm the production of the benzo[a]pyrene-6-oxyl radical in the nitro-nitrite photorearrangement of 6NO2-BaP irradiated in benzene solution:

h~

O a. Gas-Particle Reactions

Heterogeneous reactions in sunlight of particlebound nitro-PAHs with ambient 0 3 or NO 2 are generally believed to be minor relative to photodegradation. For example, based on their experimental results and modeling study of the formation and decay of nitroPAHs in diesel exhaust emissions aged in the light and dark in an outdoor smog chamber, Fan and co-workers (1995) concluded that "photodecomposition was the main loss pathway for nitro-PAH in the atmosphere." Subsequently, Fan and co-workers (1996a) conducted an experimental and modeling study of the reactions of 0 3 and NO 2 with nitro-PAHs on heterogeneous soot particles. They concluded that while photodegradation is the major daytime loss process for the nitro-PAH, during the night, particle oxidation by 0 3 may be the most important decay pathway.

O

2 peri hydrogens 6-NO2-BaP

+ NO ~ Products

C( O

(

O

h~

slow reaction

>

Products

1 peri hydrogen 1-NO2-BaP

b. Photochemical Reactions

In 1966, Chapman and co-workers proposed a nitro-nitrite photorearrangement as an efficient primary photochemical process for nitroarenes in which the nitro group is out of the plane of the aromatic rings. This is followed by dissociation into NO and a phenoxy-type radical; ultimately quinones and other oxy products are formed (Chapman et al., 1966).

/ O

NO 2

COX)

>

fast reaction

C
310 nm) 1-nitropyrene, 1,8-dinitropyrene, and 3-nitrofluoranthene coated onto silica or in a dimethyl sulfoxide solution. Half-lives for photodecomposition of 1-nitropyrene in solution compared to those on silica were 1.2 and 6 days, respectively; for 3-nitrofluoranthene the half-lives were 12.5 days in solution and > 20 days on silica. Interestingly, 1,8-dinitropyrene photodecomposed with half-lives of 0.7 day in dimethyl sulfoxide compared to 5.7 days on silica; a major photodecomposition product was 1nitropyren-8-ol. Van den Braken-van Leersum and co-workers (1987) reported that on irradiation in methanol solutions (A > 300 nm), 1-nitropyrene undergoes a rapid conversion via the nitro-nitrite rearrangement, forming 1-hydroxypyrene (88%) and 1-hydroxy-2-nitropyrene (7%). Under the same conditions, the 4-nitro isomer is more stable than the 1-nitro isomer and 2-nitropyrene is very stable; it does not react either with or without oxygen present. In a study of substrate effects on the photodecomposition of several key nitroarenes, Fan and co-workers (1996b) added several key nitro-PAHs and their deuterated analogs, along with particles of diesel exhaust or wood smoke, to a 190-m 3 outdoor smog chamber. Rapid degradation was observed when they were aged in natural sunlight at temperatures from - 1 9 to +38~ For example, the half-lives on diesel soot particles at noon on June 15 were 0.8 h for both 1-nitropyrene and 2-nitropyrene and 1.2 h for 2nitrofluoranthene. The half-life was 0.5 h for 1-nitropyrene-d 9 on particles of wood soot. Overall, the authors concluded that the photodecomposition rates of these nitro-PAHs are strongly influenced by the physical and chemical nature of the substrates as seen in the photooxidation of PAHs discussed earlier, and on the solar radiation. This is consistent with a study by Feilberg and Nielsen (1999b), who investigated the influence of other aerosol components on the photodegradation rates of representative particle-associated nitro-PAHs in a model system consisting of the nitro-PAH dissolved in cyclohexane along with various known constituents of diesel exhaust and wood smoke particles. These "cosolutes" included PAHs, substituted phenols, hydroxyPAHs, oxy-PAHs, and substituted benzaldehydes.

519

In the absence of cosolutes, the photodegradation rates depended on the orientation of the nitro group. Thus 1-nitropyrene decayed relatively fast by the nitro-nitrite primary intramolecular photorearrangement process, followed by secondary radical reactions. However, 2-nitropyrene and 2-nitrofluoranthene were stable toward photolysis, consistent with the NO 2 group being in the same plane as the aromatic rings. However, when H-atom-donating cosolutes, e.g., certain phenols, were added, the photodegradation rates of both 1-nitropyrene and 3-nitrofluoranthene increased. In this case, the reaction occurred via H-atom abstraction from the phenol by the electronically excited nitro-PAHs. Feilberg and Nielsen concluded that the photodegradation of nitro-PAHs on both diesel particles and wood smoke proceeds primarily by radical formation. However, H-atom abstraction by the excited triplet states of 1-nitropyrene and 2-nitrofluoranthene may also contribute.

F. REACTIONS OF GAS.PHASE PAHs: ATMOSPHERIC FORMATION OF M U T A G E N I C NITROARENES From the late 1970s to the mid-1980s, the major source of nitro-PAHs in ambient air was thought to be combustion emissions of particulate POM. However, field and laboratory evidence subsequently pointed to a second major source of NO2-PAHs, their in situ atmospheric formation through gas-phase reactions of 2- to 4-ring volatile and semivolatile PAHs. Although both are potential sources of nitro-PAHs in air, as we shall see, the distributions of various isomers and compounds provide compelling evidence for a significant contribution of atmospheric reactions to the direct mutagenicities of polluted airsheds throughout the world. We discuss first their direct emissions and then their formation in atmospheric reactions of PAHs.

1. Combustion-Generated Primary Emissions of Nitroarenes Initially, major attention was focused on emissions of diesel soot from light- and heavy-duty motor vehicles, because the particle extracts contained such direct-acting bacterial mutagens and animal and possible human carcinogens (IARC Group 2B, 1989) as 1-nitropyrene (1-NO2-PY) and the 1,6- and 1,8-dinitropyrenes (DNP), as well as a third isomer, 1,3-dinitropyrene, IARC Group III. For example, Paputa-Peck and coworkers (Paputa-Peck et al., 1983; IARC, 1989) reported ~100 nitro-PAHs (17 positively identified) in

520


an extract of a light-duty diesel particulate sample. Among these, as seen in Table 10.33, 1-nitropyrene was the major nitroarene at levels of 7 5 / x g / g compared to 0.30, 0.40, and 0.53 /xg/g for 1,3-, 1,6-, and 1,8-DNP, respectively. For a comparison as to how the absolute and relative emission strengths of 1-nitropyrene and the three dinitropyrenes can vary, see the study of Japanese diesel-powered vehicles by Hayakawa and co-workers (1994). Also, for comparison with the 1983 Paputa-Peck et al. study, see the report of Feilberg and co-workers (1999a), who recently identified and quantified the semivolatile 1-nitronaphthalene in diesel exhaust (vide infra). Generally, 1-nitropyrene and 2-nitrofluorene are the most abundant nitro-PAHs in diesel exhaust (e.g., see also Schuetzle and Perez, 1983; and Beije and M611er, 1988). However, emission rates of these and other nitro-PAHs vary significantly with engine type, fuel, operating conditions, etc. (e.g., see Schuetzle and Frazier, 1986; IARC, 1989; and WHO, 1996). Thus generalizations should be viewed with caution (e.g., 2nitrofluoranthene is greater than 2-nitrofluorene in Table 10.33). Furthermore, the distribution of diesel compared to gasoline engine powered vehicles varies markedly from country to country (indeed, state to state), with associated variations in exhaust emission impacts on ambient air quality. For example, in 1994, diesel-powered vehicles constituted 18% of all vehicles in Japan, with the proportion continuing to increase. Murahashi and co-workers (1995) reported total dinitropyrene concentrations [E;(1,3-DNP + 1,6-DNP + 1,8-DNP)] of 3.6 pg m -3 in samples taken adjacent to a busy intersection in downtown Kanazawa, Japan. In ambient air, Berlin, Germany, concentrations of 5.2 ng m -3 of 2-nitrofluorene (also emitted in diesel exhaust) were reported by Beije and Maller (1988).

T A B L E 10.33 Concentrations of Selected Nitroarenes in Extracts of Diesel Particulate P O M a'b Nitroarene 1-Nitronaphthalene 2-Nitronaphthalene 2-Nitrofluorene 1-Nitropyrene 3-Nitrofluoranthene 8-Nitrofluoranthene 6-Nitrobenzo[ a ]pyrene 1,3-Dinitropyrene 1,6-Dinitropyrene 1,8-Dinitropyrene

Concentration (Ixg/g) 0.95 0.35 1.2 75 3.5 1.3 4.2 0.30 0.40 0.53

a Adapted from Paputa-Peck et al. (1983) and IARC (1989). b See also Table 10.35.

As seen in Table 10.34, the isomeric distribution of the nitro-PAHs in diesel exhaust particles is consistent with electrophilic nitrations of their parent PAHs. The dominance of 1-nitropyrene and the isomer distributions of the nitropyrenes and nitrofluoranthenes observed in diesel exhaust are generally consistent with the higher reactivity of pyrene on the electrophilic reactivity scale (Nielsen, 1984; see Table 10.30) and with the Ruehle et al. (1985) assignment of major electrophilic nitration products. [For evaluations of diesel exhaust, see IARC (1989), WHO (1996); and CARB (1998); reviews include Schuetzle (1983), White (1985), Tokiwa and Ohnishi (1986), and Schuetzle and Daisey (1990); articles include, e.g., Xu et al. (1982), Schuetzle et al. (1982), Gibson (1983), Schuetzle and Perez (1983), Paputa-Peck et al. (1983), Schuetzle and Frazier (1986), Robbat et al. (1986), Beije and M611er (1988), MacCrehan et al. (1988), Schneider et al. (1990), Westerholm et al. (1991), Rogge et al. (1993a), Hammerle et al. (1994), Hayakawa et al. (1994), Johnson et al. (1994), Lowenthal et al. (1994), Westerholm and Egeb~ick (1994), Murahashi et al. (1995), and Nielsen (1995, 1996).] Other combustion sources of nitro-PAHs and nitroPACs include, for example, the following: gasoline exhaust, Wang et al., 1978; Alsberg et al., 1985; IARC, 1989; coal fly ash, Fisher et al., 1979; Fisher, 1983; White, 1985; w o o d smoke, Alfheim et al., 1984b; indoor air particulate matter, e.g., emissions from a kerosene heater, Kinouchi et al., 1988; see also the review by van Houdt (1990) and references therein.

2. Atmospheric Formation of Nitro-PAHs and N i t r o - P A C s In the mid-1980s, the discovery of significant amounts of 2-nitropyrene in ambient particulate matter collected in a rural region of Denmark by Nielsen and co-workers (1984) and 2-nitrofluoranthene in southern California by Pitts and colleagues (1985b) provided unique initial evidence for the possible formation of nitroarenes by reactions of precursor PAHs in ambient air. Thus, these compounds are not electrophilic nitration products of their parent PAHs and are generally not observed in combustion sources such as diesel soot; see, for example, Table 10.34 and the report of Ciccioli and co-workers on the detection of emissions of 2nitrofluoranthene and 2-nitropyrene solely from a "very minor" Italian industrial source (see Ciccioli et al., 1993, 1995, 1996, and references therein). Subsequently, researchers confirmed the presence of 2-nitropyrene and 2-nitrofluoranthene in extracts of samples of ambient POM collected at sites throughout the world and reported 2-nitrofluoranthene levels that

T A B L E 10.34 Distribution and (Yields) of Nitro-PAH Isomers (a) Formed from the Gas-Phase Reactions of Selected Volatile and Semivolatile P A H s with OH Radicals and NO, Radicals (Each in the Presence of Oxides of Nitrogen), (b) Formed by Electrophilic Nitration, and (c) Present in Diesel Exhaust"9b Parent PAH

Gas-phase OH radical initiated reaction (yield)h

Gas-phase NO, radical initiated reaction (yield)h

Electrophilic nitrationC

NO,-PAH in diesel exhaust':

I-Nitronaphthalene > 2-Nitronaphthalene

I-Nitronaphthalene 2-Nitronaphthalene

1M4NN > 1M2NN > 1MSNN"

None reported

Naphthalene

I-Nitronaphthalene (0.3%) 2-Nitronaphthalene (0.3%)

1-Methylnaphthalene

1MSNN" > 1M4NN 2 1M6NN 1M3NN > 1MSNN 2 1M4NN (total yield -0.4'70)~ (total yield -30%)''

2-Methylnaphthalene

2M5NN > 2M6NN 2M7NN (total yield -0.2'70)~

2M4NN > 2MlNN (total yield -30%)"

Fluoieiie

3-Nitiofluoieiie (-1.4%) I-Nitrofluorene (-0.6%)

No data

2-Niti ofluoi eiie

2-Nitiofluoieiie

I-Nitroanthracene (low yield) 2-Nitroanthracene (low yield)

1-Nitroanthracene (low yield) 2-Nitroanthracene (low yield)

9-Nitroanthracene

9-Nitroanthracene

2-Nitrofluoranthene ( -3%)

2-Nitrofluoranthene ( -24%)

3-Nitrofluoranthene > 8-Nitrofluoranthene > 7-Nitrofluoranthene

3-Nitrofluoranthene

Anthracene

'

Fluoranthene

-

I-Nitronaphthalene (17%) 2-Nitronaphthalene (7%)

- 2M5NN

2MlNN

B

2M8NN > 2M4NN"

2-Methyl-1-nitronaphthalene

> 8-Nitrofluoranthene

7-Nitrofluoranthene ( - 1%) (-24%) 8-Nitrofluoranthene (-0.3%) Pyrene

2-Nitropyrene (-0.5%) 4-Nitropyrene ( -0.06%)

4-Nitropyrene (0.06%)

1-n1tropyrene

1-Nitropyrene

Biphenyl

3-Nitrobiphenyl ( 5 % )

No reaction observed

2-Nitrobiphenyl 4-Nitrobiphenyl"

2-Nitrobiphenyl

Adapted from Atkinson and Arey (1997) and Arey (1998a). Data from Atkinson and Arey (1994). ' From Ruehle et al. (1985). ' I From Eaborn et al. (1968). Nomenclature: e.g., 1M5NN = 1-methyl-5-nitronaphthalene. ' From Zielinska et al. (1990). 'Although 9-nitroanthracene was observed i n both OH and NO, reactions, it may not be a product as From Paputa-Peck et al. (1983); see Table 10.33. From Zielinska et al. (1989b). 'I

it

is also found

in

exposure to NO,/HNO, systems (Arey, 1998a).

522


generally exceeded those of 1-nitropyrene (e.g., Nielsen and Ramdahl, 1986; Ramdahl et al., 1986; Sweetman et al., 1986; Arey et al., 1987, 1988b; Atkinson et al., 1987a, 1988a; Nishioka et al., 1988; Zielinska et al., 1989a; Ciccioli et al., 1993, 1995, 1996; Legzdins et al., 1994; Atkinson and Arey, 1994; Wilson et al., 1995; Arey, 1998a). For example, Legzdins and co-workers (1994) used the bioassay-directed fractionation and chemical analysis technique to isolate, identify, and quantify 2nitrofluoranthene in extracts of ambient particles collected in Hamilton, Ontario, Canada. They found it accounted for ~ 70% of the total nonpolar direct bacterial mutagenicity (strain YG1021, standard reversion assay, Maron and Ames, 1983). Additionally, in two different monitoring campaigns conducted in the center of Milan, Italy, Ciccioli and co-workers (1993) reported 2-nitrofluoranthene, 2nitropyrene, and 1-nitropyrene were the only nitroarenes detected. Subsequently, in a comprehensive study of the atmospheric formation and transport of 2-nitrofluoranthene and 2-nitropyrene, they established their presence and levels in ambient particles collected at sites located in urban, suburban, forest, and remote areas in Europe, Asia, America, and Antarctica (Ciccioli et al., 1996, and references therein; see also Ciccioli et al., 1995). As an example of typical experimental data, Fig. 10.32 is a G C - M S selected ion monitoring (SIM) profile ( m / z 247) for the nitrofluoranthenes and nitropyrenes in an extract of ambient particles collected in southern California (Arey et al., 1988b). The 1-nitropyrene (1-NP) and 3-nitrofluoranthene (3-NF) presumably are from diesel emissions (Tables 10.33 and 10.34), but the dominance of 2-nitrofluoranthene and 2-nitropyrene reflects a second major source. Shortly after discovery of the 2-nitro isomers of pyrene and fluoranthene in ambient air, an OH radical initiated mechanism analogous to that of Nielsen and co-workers (1984) for 2-nitropyrene was proposed and experimental laboratory evidence reported for the in situ formation of 2-nitrofluoranthene (Pitts et al., 1985b; Ramdahl et al., 1986; Nielsen and Ramdahl, 1986; Sweetman et al., 1986; Arey et al., 1986; for background, see the discussion by Atkinson et al., 1987a). As seen in Fig. 10.33, the major pathway involves (a) OH radical attack at the sites of highest electron densities (the 3-position is favored for fluoranthene and the 1-position for pyrene, (b) the addition of NO 2 in the ortho (2-)position to the hydroxycyclohexadienyl-type radical, and (c) loss of water and formation of 2-nitrofluoranthene (and by analogy 2-nitropyrene). Note that analogous to the reactions of OH with simple aromatics, 0 2 can compete with NO 2 in adding to the

100 -

2-NF

m/z 247

o t-t~ r

-'

50

".~

2-NP 4-NP

7-NF 01 21

I

I

22

I

A

P I

23

I

I

24

Retention time (min)

F I G U R E 10.32 G C - M S selected ion monitoring (SIM) scan of the molecular ions (rn/z 247) of nitrofluoranthene (NF) and nitropyrene (NP) isomers in extracts of ambient particles collected in Torrance, California, January 1986 (adapted from Arey et al., 1988b, and Atkinson and Arey, 1997).

OH adduct radical (e.g., see Atkinson, 1994, and Chapter 6.G). However, in the presence of sufficient NO2, 2-nitrofluoranthene and 2-nitropyrene are formed and subsequently condense out on particle surfaces (see Pitts, 1987, Atkinson and Arey, 1994, Arey, 1998a, and references therein). For this mechanism to be effective, several requirements must be met: Under ambient conditions, a substantial portion of the fluoranthene and pyrene must be in the gas phase, which is indeed the case (see Figs. 10.2 and 10.3). Attack by OH radicals on the gas-phase fluoranthene and pyrene must be fast. As seen in Table 10.35, again this is true. In Table 10.36, the calculated atmospheric lifetimes of selected gas-phase PAHs due to reaction with OH are shown, e.g., lifetimes of ~2.9 h for fluoranthene and pyrene. Products (and their mutagenicities) of the gas-phase reactions of these and other 2- to 4-ring PAHs (fluorene, naphthalene, etc.) carried out under simulated atmospheric conditions should be consistent with the nitroarenes and nitro-PACs that have been identified in ambient air. This criterion has been well established over the years in environmental chamber studies conducted in several laboratories (e.g., see reviews by Atkinson and Arey, 1994; and Arey, 1998a; articles by Kamens et al., 1994; Fan et al., 1995; Feilberg et al., 1999a; and references therein).

F. REACTIONS OF GAS-PHASE PAHs: ATMOSPHERIC FORMATION OF MUTAGENIC NITROARENES

02N 9

H

523

NO2

I

OH

NO2

+ H20 2-NF

(3O/o)

/ OH+

Ho:r

NO2

+ H20 H

NO2 7-NF (1%)

F I G U R E 10.33 Arey, 1998a).

Mechanism of oxidation of fluoranthene by OH in air (adapted from

For example, Fig. 10.34 shows the mutagram (TA98,

-$9) of an extract of ambient particles collected in Claremont, California, August 1987 (Harger et al., 1992) together with mutagrams of extracts of environmental chamber reaction products for the simulated OH radical initiated reactions of phenanthrene, fluoranthene, and pyrene, respectively (Arey et al., 1992; Sasaki et al., 1995; Atkinson and Arey, 1994; Arey, 1998a). It can be seen that, in combination, the mutagrams of the photooxidation reaction products of phenanthrene, pyrene, and fluoranthene are very similar to the mutagen profile of the ambient air sample. 2-Nitrofluoranthene, found in fraction 4, dominated the fluoranthene reaction products, consistent with extracts of ambient POM. The more polar fraction 6 of the OH radical initiated phenanthrene reaction products contained the two nitrophenanthrene lactones previously identified in ambient air (Helmig et al., 1992a, 1992b), the potent direct mutagen 2-nitro-6Hdibenzo[b,d]pyran-6-one (XI) and its less mutagenic 4-NO 2 isomer (see Table 10.20). Recall that this 2nitro-PAC XI has been estimated to account for up to 20% of the direct mutagenicity of an ambient aerosol extract (Helmig et al., 1992b). Two nitropyrene lactone isomers were tentatively identified in fraction 6 of the pyrene reaction products. E1-Bayoumy and Hecht (1986) reported earlier that two nitropyrene lactones were

strong mutagens in the standard Ames plate incorporation assay (Maron and Ames, 1983). Figure 10.35 is the GC-MS m / z 247 profile of the nitrofluoranthenes and nitropyrenes in an extract of an ambient particle sample collected at night (Arey et al., 1988b). The high ratio of 2-nitrofluoranthene to 2-nitropyrene observed in this nighttime sample is indicative of nighttime gas-phase NO 3 radical reactions (for a review, see Kwok et al., 1994b). An NO 3 radical initiated mechanism for atmospheric formation of 2nitrofluoranthene is shown in Fig. 10.36 (Atkinson and Arey, 1997; Arey, 1998a, and references therein). Analogous to the OH reaction, NO 3 is proposed to add to the ring to form a fluoranthene-NO 3 adduct, followed by ortho addition of NO 2 and subsequent loss of HNO 3. This reaction is noteworthy because of its selectivity; i.e., only 2-nitrofluoranthene is formed, and in high yield (24%)compared to the OH-initiated daytime reaction (3%). Another example is the atmospheric formation levels, fates, and mutagenicities of 1- and 2-nitronaphthalenes (1-NN and 2-NN) and certain methylnitronaphthalene isomers (MNN). Naphthalene is the most abundant of the PAHs and its role, and that of the abundant MNNs, in the atmospheric chemistry of PAHs are being increasingly recognized as being important; e.g., see Pitts et al., 1985c; Atkinson et al., 1987b,

524


T A B L E 10.35

R o o m T e m p e r a t u r e Rate C o n s t a n t s , k, for the G a s - P h a s e Reactions of Selected P A H s and N i t r o - P A H s with the H y d r o x y l Radical, the N i t r a t e Radical, and Ozone (from A r e y 1998a) k (cm 3 m o l e c u l e - 1 s - l ) for r e a c t i o n w i t h

P A H or n i t r o - P A H

OH"

Naphthalene 1-Methylnaphthalene 2-Methylnaphthalene Acenaphthylene Acenaphthene Biphenyl Fluorene f Phenanthrene g Anthracene Fluoranthene Pyrene 1-Nitronaphthalene 2-Nitronaphthalene

2.16 5.3 5.2 1.1 1.0 7.2 1.6 1.3 1.7 ~ 5 ~ 5 5.4 5.6

• • • • • • • • • • • • •

O3 c

NO3 b

10-11 10-11 10-11 10-10 10 - l ~ 10-12 10-11 10-11 10-11 f 10 -11i 10-11i 10-12 10-12

3.6 7.7 1.08 5.5 4.6 .m m

80 "O 11.)

>

0 o

>, > ,,i-, o

O O

1-NP

60

A

40

2-NP ^

Relative retention time

20

o c

N E

rr

Phenanthrene

1

2

3

4

5

6

7

8

9

Fluoranthene

80

F I G U R E 10.35 GC-MS selected ion monitoring (SIM) of the molecular ions ( r n / z 247) of nitrofluoranthene (NF) and nitropyrene (NP) isomers in an extract of particles collected at night in Claremont, California, 1800 to 2400 hours on September 14, 1985 (adapted from Arey et al., 1988b).

60 40 20 1

2

3

4

5

6

7

8

9

Pyrene

80 60 40 20 1

2

3

I I 4

5

6

7

8

9

Fraction

F I G U R E 10.34 Mutagrams (Ames assay, microsuspension modifications, TA98, - $ 9 ) of extracts of ambient particulate matter collected on a filter in Claremont, California, August 28, 1987 (Arey et al., 1992; Harger et al., 1992) and the percentage of the total mutagenic activity from environmental chamber reaction products vs HPLC fraction number of the gas-phase atmospheric reactions of the semivolatile phenanthrene, fluoranthene, and pyrene, respectively (Sasaki et al., 1995). The mutagenicity of fraction 6 from the phenanthrene reaction is due to the product 2-nitro-6H-dibenzo[b,d]pyran6-one (the nitrophenanthrene lactone XI); 2-nitrofluoranthene (XXVII) is the mutagenic compound in fraction 4 from the fluoranthene reaction. Two nitropyrene lactones were tentatively identified in fraction 6 from the pyrene chamber reaction (adapted from Atkinson and Arey, 1994, and Arey, 1998a).

1990a,b, 1994; Zielinska et al., 1989b; Arey et al., 1992; Lane and Tang, 1994; Bunce and Zhu, 1994; Bunce et al., 1997; Sasaki et al., 1995, 1997a; and Feilberg et al., 1999a; see also reviews by Zielinska et al., 1990; Atkinson and Arey, 1994; and Arey, 1998a.

For example, Fig. 10.37 shows a mutagram of an extract of ambient gas-phase POM collected on polyurethane foam plugs in Claremont, California, along with mutagrams of the extracts of PUF samples collected from environmental chamber studies of the OH radical initiated reactions of fluorene and naphthalene, respectively (Arey et al., 1992; Harger et al., 1992; Atkinson and Arey, 1994). Clearly, the mutagrams from the OH-initiated reactions are very similar to that of ambient air. Thus, most of the activity of the products of the photooxidation of these two PAHs is in fraction 4. In the naphthalene reaction, key products in fraction 4 are 1- and 2-nitronaphthalene (Arey et al., 1992), with

~

ONO2

NO3 +

O2N

H

NO2 ONO2 H

~_ + HNO3 2-NF (24%)

F I G U R E 10.36 Mechanism of oxidation of fluoranthene by the nitrate radical (adapted from Atkinson and Arey, 1997).

526

10. AIRBORNEPAHs: CHEMISTRYAND TOXICOLOGICALIMPLICATIONS

H OH

Ambient PUF

80

o.+ @

60

[ ~ "

NO=//

40 20

H OH

"(3 (1)

1

2

3

4

~9

5

6

7

8

9

80

E

O"

-HONO/

60

/

40

OH

t--

.1-., 2~

H OH NOa

Fluorene + OH

02

.m ..i-, O

(andother isomer)

CHOH [~~/NO2

20

[ ~

CH=CHCHO

m O .1-.,

1

O

2

3

4

5

6

7

8

9

GHO

HO2+ [ ~

Naphthalene + OH

80

CH=CHCHO

FIGURE 10.38 Mechanism of oxidation of naphthalene in air by OH (from Sasaki et al., 1997a).

60 40 20 1

2

3

4

5

6

7

8

9

Fraction FIGURE 10.3 7 Mutagrams of extracts of an ambient sample of gas-phase direct mutagens collected on a PUF plug in Claremont, California, and, for comparison, extracts of the vapor-phase products from the OH radical initiated gas-phase reactions of naphthalene and fluorene in an environmental chamber. For naphthalene, the mutagenicity of fraction 4 is primarily due to the presence of 1- and 2-nitronaphthalene. Fraction 4 from the fluorene reaction contains four nitrofluorene isomers, with 3-nitrofluorene being dominant (adapted from Arey et al., 1992; Harger et al., 1992).

As discussed earlier (see Fig. 10.26), the contributions of 1- and 2-nitronaphthalenes and the methylnitronaphthalene isomers to the vapor-phase mutagenicity of extracts of daytime and nighttime samples of ambient air in Redlands, California, were determined by Gupta and co-workers (1996) and shown to be major contributors to the overall vapor-phase mutagenicity of ambient air. Furthermore, the researchers demonstrated that both daytime O H radical initiated and nighttime NO 3 radical initiated reactions of naphthalene and the methylnaphthalenes are important in pol-

H

the latter accounting for ~ 9 0 % of the activity of fraction 4. In the fluorene reaction, fraction 4 contained four nitrofluorene isomers; however, 3-nitrofluorene was the dominant isomer, accounting for 75% of the activity of this fraction. Figure 10.38 shows the mechanism of Sasaki and co-workers (1997a) for the daytime photooxidation of naphthalene in ambient air (i.e., O H attack), whereas Fig. 10.39 shows their nighttime mechanism for NO 3 radical attack. In contrast to the variety of ring-opened products and low total yields of 1- and 2-nitronaphthalene from O H radical attack (less than 1% for both compounds), no ring-opened products are formed, and the yields of the nitronaphthalenes are much larger for the NO 3 radical reaction, with the ratio of the 1- to 2-isomers being 2.4:1 (Sasaki et al., 1997a).

NO3+

~-

~

H

~

ONO 2

(andotherisomer)

ONO2

(Products)

H

H

NO2 ~~[ ~ ~ , ~ "

/NO2 + HNO3

OH ~

A A

+ NO 2

OH ,.---- [ ~ / N O 2

FIGURE 10.39 Mechanism of NO 3 radical reaction with naphthalene in air (from Sasaki et al., 1997a).

aEFEaENCES

luted ambient air atmospheres. Thus the MNN isomer pattern found in the daytime ambient PUF samples was significantly different from the nighttime profile. Furthermore, there was excellent agreement between the laboratory MNN isomer product patterns observed for OH attack and the profile from field measurements during the day and between the laboratory NO 3 studies and the nighttime isomer profiles. Recently, to test their model for the formation, decay, and gas-particle partitioning of nitronaphthalenes, Feilberg and co-workers (1999a) conducted outdoor environmental chamber studies of the reactions of naphthalene in the presence of diesel exhaust, propylene (to enhance OH radical formation), and additional NO. Experimentally, they determined that the gas-phase photolysis of 1-nitronaphthalene was its major degradation pathway; however, 2-nitronaphthalene photolyzed at a significantly slower rate so that other degradation processes may be important. These results differed from those of Atkinson et al. (1989), who found their photolysis rates comparable (see Table 10.36). Additionally, Feilberg et al. propose a somewhat modified mechanism of Atkinson et al. (1989) and Sasaki et al. (1997a) for the OH-initiated gas-phase reaction of naphthalene to form 1-nitronaphthalene. For the kinetics and mechanisms of the gas-phase reactions of other 2- to 4-ring PAHs with OH radicals, NO 3 radicals, and 03, see, e.g., phenanthrene, Kwok et al. (1994a), and indan, indene, fluorene, and 9,10-dihydroanthracene, Kwok et al. (1997); for dibenzothiophene, a volatile S-PAC that is ubiquitous in polluted ambient air environments and forms mutagenic photooxidation products, see Kwok et al. (1999); see also Kwok et al. (1994b) for a review of gas-phase NO 3 radical reactions with aromatics. In summary, as Arey (1998a) points out, one can assign much of the direct-acting mutagenicity (TA98, - $ 9 , microsuspension modification) of fractions 4 and 6 of particle-phase ambient samples to nitroarenes (fraction 4) and nitroarene lactones (fraction 6). Certain gas-phase nitroarenes are also major contributors to fraction 4. If one sums up the particle- and gas-phase contributions of these gas-phase reaction products of semivolatile and volatile PAHs, they could account for over 50% of the total vapor- and particle-phase direct mutagenicity of ambient air. Clearly this phenomenon has toxicological implications that should be addressed in risk assessments of the exposure and health effects of combustion-generated POM, e.g., diesel exhaust. References Aceves, M., and J. 0. Grimalt, "Seasonally Dependent Size Distributions of Aliphatic and Polycyclic Aromatic Hydrocarbons in Ur-

527

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Harrison, R. M., D. J. T. Smith, and L. Luhana, "Source Apportionment of Atmospheric Polycyclic Aromatic Hydrocarbons Collected from an Urban Location in Birmingham, U.K.," Environ. Sci. Technol., 30, 825-832 (1996). Hart, K. M., and J. F. Pankow, "High-Volume Air Sampler for Particle and Gas Sampling. 2. Use of Backup Filters to Correct for the Adsorption of Gas-Phase Polycyclic Aromatic Hydrocarbons to the Front Filter," Environ. Sci. Technol., 28, 655-661 (1994). Harvey, R. G., Polycyclic Aromatic Hydrocarbons, Wiley-VCH, New York, 1997. Hawthorne, S. B., M. S. Krieger, D. J. Miller, and M. B. Mathiason, "Collection and Quantitation of Methoxylated Phenol Tracers for Atmospheric Pollution from Residential Wood Stoves," Environ. Sci. Technol., 22, 470-475 (1989). Hayakawa, K., M. Butoh, Y. Hirabayashi, and M. Miyazaki, "Determination of 1,3-, 1,6-, 1,8-Dinitropyrenes, and 1-Nitropyrene in Vehicle Exhaust Particulates," Jpn. J. Toxicol. Environ. Health, 40, 20-25 (1994). Haynes, B. S., "Soot and Hydrocarbons in Combustion," in Fossil Fuel Combustion: A Source Book (W. Bartock and A. F. Sarofim, Eds.), pp. 261-326, Wiley, New York, 1991. Hecht, S. S., and K. E1-Bayoumy, "The Possible Role of Nitroarenes in Human Cancer," in Nitroarenes (P. C. Howard et al., Eds.), pp. 309-316, Plenum, New York, 1990. Heinrich, G., and H. Glisten, "Fluorescence Spectroscopic Properties of Carcinogenic and Airborne Polynuclear Aromatic Hydrocarbons," in Polynuclear Aromatic Hydrocarbons: Chemistry and Biological Effects (A. Bjorseth and A. J. Dennis, Eds.), pp. 983-1003, Battelle Press, Columbus, OH, 1980. Heinrich, U., M. Roller, and F. Pott, "Estimation of a Lifetime Unit Lung Cancer Risk for Benzo[a]pyrene Based on Tumour Rates in Rats Exposed to Coal Tar/Pitch Condensation Aerosol," Toxicol. Lett., 72, 155-161 (1994). Helmig, D., J. Arey, W. P. Harger, R. Atkinson, and J. L6pez-Cancio, "Formation of Mutagenic Nitrodibenzopyranones and Their Occurrence in Ambient Air," Environ. Sci. Technol., 26, 622-624 (1992a). Helmig, D., J. L6pez-Cancio, J. Arey, W. P. Harger, and R. Atkinson, "Quantification of Ambient Nitrodibenzopyranones: Further Evidence for Atmospheric Mutagen Formation," Environ. Sci. Technol., 26, 2207-2213 (1992b). Helmig, D., J. Arey, R. Atkinson, W. P. Harger, and P. A. McElroy, "Products of the OH Radical-Initiated Gas-Phase Reaction of Fluorene in the Presence of NOx," Atmos. Environ., 26A, 1735-1745 (1992c). Hering, S. V., R. C. Flagan, and S. K. Friedlander, "Design and Evaluation of New Low-Pressure Impactor. 1," Environ. Sci. Technol., 12, 667-673 (1978). Hering, S. V., S. K. Friedlander, J. J. Collins, and L. W. Richards, "Design and Evaluation of a New Low-Pressure Impactor. 2," Environ. Sci. Technol., 13, 184-188 (1979). Herreno-Saenz, D., F. E. Evans, T. Heinze, J. Lewtas, and P. P. Fu, "In Vitro Metabolism and DNA Adduct Formation from the Mutagenic Environmental Contaminant 2-Nitrofluoranthene," Chem. Res. Toxicol., 5, 863-869 (1992). Hieger, I., "The Spectra of Cancer-Producing Tars and Oils and of Related Substances," Biochemistry, 24, 505-511 (1930). Hinckley, D. A., T. F. Bidleman, and W. T. Foreman, "Determination of Vapor Pressures for Nonpolar and Semipolar Organic Compounds from Gas Chromatographic Retention Data," J. Chem. Eng. Data, 35, 232-237 (1990). Hirayama, T., H. Kusakabe, T. Watanabe, S. Ozasa, Y. Fujioka, and S. Fukui, "Relationship between Mutagenic Potency in Salmonella

typhimurium Strains and the Chemical Structure of Nitro Biphenyls," Mutat. Res., 163, 101-107 (1986). Hof, R. M., K. A. Brice, and C. J. Halsall, "Nonlinearity in the Slopes of Clausius-Clapeyron Plots for SVOCs," Environ. Sci Technol., 32, 1793-1798 (1998). Hoff, R. M., and K.-W. Chan, "Measurement of Polycyclic Aromatic Hydrocarbons in the Air along the Niagara River," Environ. Sci Technol., 21, 556-561 (1987). Hoffman, D., and E. L. Wynder, "Organic Particulate Pollutants--Chemical Analysis and Bioassays for Carcinogenicity, in Air Pollution (A. C. Stern, Ed.), Vol. II, 3rd ed., pp. 361-455, Academic Press, New York, 1977. Holder, P. S., E. L. Wehry, and G. Mamantov, "Photochemical Transformation of 1-Nitropyrene Sorbed on Coal Fly Ash Fractions," Polycyclic Aromat. Compd., 4, 135-139 (1994). Holloway, M. P., M. C. Biaglow, E. C. McCoy, M. Anders, H. S. Rosenkranz, and P. C. Howard, "Photochemical Instability of 1-Nitropyrene, 3-Nitrofluoranthene, 1,8-Dinitropyrene and Their Parent Polycyclic Aromatic Hydrocarbons," Mutat. Res., 187, 199-207 (1987). Hollstein, M., J. McCann, F. A. Angelosanto, and W. W. Nichols, "Short-Term Tests for Carcinogens and Mutagens," Mutat. Res., 65, 133-226 (1979). Holmberg, B., and U. Ahlborg, Eds., Symposium on Biological Tests in the Evaluation of Mutagenicity and Carcinogenicity of Air Pollutants with Special Reference to Motor Exhausts and Coal Combustion Products, Stockholm, Sweden, February 8-12, in Environ. Health Perspect., 47, 1-345 (1983). Howard, P. C., S. S. Hecht, and F. A. Beland, Eds., Nitroarenes: Occurrence, Metabolism, and Biological Impact, Environmental Science Research, Vol. 40, Plenum, New York, 1990. Howsam, M., and K. C. Jones, "Sources of PAHs in the Environment," in The Handbook of Environmental Chemistry, PAHs and Related Compounds (A. H. Neilson, Ed.), Vol. 3, Part I, SpringerVerlag, Berlin (1998). Hsieh, P., N. Y. Kado, J. N. Seiber, and T. Shibamoto, "Methods Development for Assessment of Vapor-Phase Mutagens and Carcinogens," Final Report to the California Air Resources Board, Contract No. A6--174-32, June 1990. Hsieh, D. P. H., N. Y. Kado, and R. Okamoto, "Measurement and Chemical Characterization of Vapor-Phase Mutagens in Diesel Exhaust," Final Report to the California Air Resources Board, Contract No. A032-095, March 1993. Hudgins, D. M., and S. A. Sandford, "Infrared Spectroscopy of Matrix Isolated Polycyclic Aromatic Hydrocarbons. 1. PAHs Containing Two to Four Rings," J. Phys. Chem. A, 102, 329-343 (1998a). Hudgins, D. M., and S. A. Sandford, "Infrared Spectroscopy of Matrix Isolated Polycyclic Aromatic Hydrocarbons. 2. PAHs Containing Five or More Rings," J. Phys. Chem. A, 102, 344-352 (1998b). Hudgins, D. M., and S. A. Sandford, "Infrared Spectroscopy of Matrix Isolated Polycyclic Aromatic Hydrocarbons. 3. Fluoranthene and the Benzofluoranthenes," J. Phys. Chem. A, 102, 353-360 (1998c). Hueper, W. C., P. Kotin, E. C. Tabor, W. W. Payne, H. L. Falk, and E. Sawicld, "Carcinogenic Bioassays on Air Pollutants," Arch. Pathol. Lab. Med., 74, 89-116 (1962). Hughes, M. M., D. F. S. Natusch, D. R. Taylor, and M. V. Zeller, "Chemical Transformations of Particulate Polycyclic Organic Matter," in Polynuclear Aromatic Hydrocarbons: Chemistry and Biological Effects (A. Bj~rseth and A. J. Dennis, Eds.), pp. 1-8, Battelle Press, Columbus, OH, 1980.

REFERENCES Hughes, T. J., D. M. Simmons, L. G. Monteith, and L. D. Claxton, "Vaporization Technique to Measure Mutagenic Activity of Volatile Organic Chemicals in the Ames/Salmonella Assay," Environ. Mutagen., 9, 421-441 (1987). Huisingh, J., R. Bradow, R. Jungers, L. Claxton, R. Zweidinger, S. Tejada, J. Bumgarner, F. M. Waters, V. F. Simmon, C. Hare, C. Rodgriguez, and L. Snow, "Application of Bioassay to Characterization of Diesel Particle Emissions," in Applications of Short-Term Bioassay in the Fractionation and Analysis of Complex Environmental Mixtures (M. D. Waters, S. Nesnow, J. L. Huisingh, S. Sandhu, and L. Claxton, Eds.), pp. 383-418, Plenum, New York, 1979. Imaida, K., L. K. Tay, M.-S. Lee, C. Y. Wang, N. Ito, and C. M. King, "Tumor Induction by Nitropyrenes in the Female CD Rat," in Carcinogenic and Mutagenic Responses to Aromatic Amines and Nitroarenes," Proceedings of the Third International Conference on Carcinogenic and Mutagenic N-Substituted Aryl Compounds, April 25-28, 1987, Dearborn, Michigan (C. M. King, L. J. Romano, and D. Schuetzle, Eds.), pp. 187-197, Elsevier, New York, 1988. International Agency for Research on Cancer (IARC), Monographs in the series Evaluation of Carcinogenic Risk of Chemicals to Humans: Vol. 32, PolynuclearAromatic Compounds, Part 1, 1984a; Vol. 33, Polynuclear Aromatic Compounds, Part 2, Carbon Blacks, Mineral Oils, and Some Nitroarenes, 1984b; Vol. 34, Polynuclear Aromatic Compounds, Part 3, Industrial Exposures in Aluminum Production, Coal Gasification, Coke Production, and Iron and Steel Foundry, 1984c; Vol. 35, PolynuclearAromatic Compounds, Part 4, Bitumens, Coal-Tars and Derived Products, Shale-Oils, and Soots, 1985; Supplement No. 7, Overall Evaluations of Carcinogenicity: An Updating of IARC Monographs, Vols. 1-42, 1987; Vol. 46, Diesel and Gasoline Exhausts and Some Nitroarenes, 1989. International Union of Pure and Applied Chemistry, Nomenclature of Organic Chemistry, Sections A, B, C, D, E, F, and H, Pergamon, Oxford, 1979. Ioki, Y., "Aryloxyl Radicals by Photorearrangement of NitroCompounds," J. Chem. Soc., Perkin Trans. II, 1240-1242 (1977). Ismail, Y., A. L. Lafleur, and R. W. Giese, "Polar Forms of Benzo[a]pyrene May Be Ubiquitous in the Environment," Environ. Sci. Technol., 32, 2494-2497 (1998). Izumi, K., and T. Fukuyama, "Photochemical Aerosol Formation from Aromatic Hydrocarbons in the Presence of NOx," Atmos. Environ., 24A, 1433-1441 (1990). J~iger, J., "Detection and Characterization of Nitro Derivatives of Some Polycyclic Aromatic Hydrocarbons by Fluorescence Quenching after Thin-Layer Chromatography: Application to Air Pollution Analysis," J. Chromatogr., 152, 575-578 (1978). J~iger, J., and V. Hanus, "Reaction of Solid Carrier-Adsorbed Polycyclic Aromatic Hydrocarbons with Gaseous Low-Concentrated Nitrogen Dioxide," J. Hyg. Epidemiol. Microbiol. Immunol., 24, 1-12 (1980). Jang, M., and S. R. McDow, "Benz[a]anthracene Photodegradation in the Presence of Known Organic Constituents of Atmospheric Aerosols," Environ. Sci. Technol., 29, 2654-2660 (1995). Jang, M., and S. R. McDow, "Products of Benz[a]anthracene Photodegradation in the Presence of Known Organic Constituents of Atmospheric Aerosols," Environ. Sci. Technol., 3I, 1046-1053 (1997a). Jang, M., R. M. Kamens, K. B. Leach, and M. R. Strommen, "A Thermodynamic Approach Using Group Contribution Methods to Model the Partitioning of Semi-Volatile Organic Compounds on Atmospheric Particulate Matter," Environ. Sci. Technol., 3I, 2805-2811 (1997b). Jang, M., and R. M. Kamens, "A Thermodynamic Approach for Modeling Partitioning of Semi-Volatile Organic Compounds on

535

Atmospheric Particulate Matter: Humidity Effects," Environ. Sci. Technol., 32, 1237-1243 (1998). Jenkins, B. M., A. D. Jones, S. Q. Turn, and R. B. Williams, "Particle Concentrations, Gas-Particle Partitioning, and Species Intercorrelations for Polycyclic Aromatic Hydrocarbons (PAH) Emitted during Biomass Burning," Atmos. Environ., 30, 3825-3835 (1996). Johnson, D. E., and H. H. Cornish, "Metabolic Conversion of 1- and 2-Nitronaphthalene to 1- and 2-Naphthylamine in the Rat," Toxicol. Appl. Pharmacol., 46, 549-553 (1978). Johnson, J. H., S. T. Bagley, L. D. Gratz, and D. G. Leddy, "A Review of Diesel Particulate Control Technology and Emissions Effects--1992 Horning Memorial Award Lecture," SAE, Spec. Publ. SP-1020, 1-35 (1994). Jung, H., A. U. Shaikh, R. H. Heflich, and P. P. Fu, "Nitro Group Orientation, Reduction Potential, and Direct-Acting Mutagenicity of Nitro-Polycyclic Aromatic Hydrocarbons," Environ. Mol. Mutagen., 17, 169-180 (1991). Junge, C. E., "Basic Considerations about Trace Constituents in the Atmosphere as Related to the Fate of Global Pollutants," in Fate of Pollutants in the Air and Water Environments (I. H. Suffet, Ed.), pp. 7-26, Wiley, New York, 1977. Kaden, D. A., R. A. Hites, and W. G. Thilly, "Mutagenicity of Soot and Associated Polycyclic Aromatic Hydrocarbons to Salmonella typhimurium," Cancer Res., 39, 4152-4159 (1979). Kado, N. Y., D. Langley, and E. Eisenstadt, "A Simple Modification of the Salmonella Liquid Incubation Assay," Mutat. Res., 121, 25-32 (1983). Kado, N. Y., G. N. Guirguis, C. P. Flessel, R. C. Chan, K.-I. Chang, and J. J. Wesolowski, "Mutagenicity of Fine ( [4 Entrance

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Exit

F I G U R E 11.1 (a) Schematic diagram of a multipass White cell, (b) sequence of images on filled mirror for White cell design, and (c) sequence of images on field mirror for Horn and Pimentel design (1971). (Adapted from Finlayson-Pitts and Pitts, 1986; and Hanst and Hanst, 1994.)

550

11. ANALYTICAL METHODS FOR GASES AND PARTICLES IN THE ATMOSPHERE

number of spots on the field mirror) has been made in this case. The advantage of such a White cell is that the source is reimaged on the field mirror M2 after each double traversal of the cell. This keeps the energy that enters the cell within the mirror system so that energy losses occur mainly through light absorption by the mirrors and, of course, by the gases in the cell. In practice, the loss of light energy through absorption by the mirrors imposes a major limitation on the number of passes that can be used. The fraction of the energy lost after n reflections from a mirror whose reflectivity is R is given by (1 - R'). Thus, if a mirror reflects 98% of the incident light and absorbs 2%, only 36% of the incident intensity will remain after 50 reflections from the mirror. After 100 reflections, only 13% of the incident intensity is left. While the path length and hence absorbance have increased, the energy loss may be so severe that such a large number of reflections becomes impractical. The number of reflections is also limited by the size of the image striking the entrance and the size of the mirror M2. As seen in Fig. l l . l b , the images that are refocused from M1 and M3 onto the field mirror M2 are "stacked" beside each other. The width of M2 therefore determines how many of these images can be accommodated (i.e., how many reflections are possible).

A practical problem arising when the images are too closely spaced (i.e., at long path length) is one of adjustment; temperature changes, for example, can cause very small changes in the mirror adjustments which result in moving the exit beam away from the exit aperture. Variations of the White cell are also in use. For example, Horn and Pimentel (1971) added a corner mirror assembly to redirect the beam that would normally exit the cell back into it. This doubles the number of passes, giving four rows of spots on the field mirror. The image pattern for such a design is shown in Fig. 11.1c (Hanst and Hanst, 1994). More complex multiple-reflection systems that give a much greater number of traversals have also been developed. For example, Tuazon et M. (1980) describe a system using four collecting mirrors that focus the light onto four field mirrors. The advantages and disadvantages of such multiple-mirror cells are discussed by Hanst (1971) and Hanst and Hanst (1994). An alternate design for folded optics was described by White in 1976. In this design, the light beam is folded back on itself, giving larger path lengths and greater optical stability. The effects of vibration, thermal expansion, and astigmatism are reduced and alignment errors are minimized with this design.

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F I G U R E 11.2 (a) Schematic diagram of multipass cell for infrared spectroscopy using astigmatic Herriott configuration (adapted from McManus et al., 1995), (b) spot configurations for normal Herriott multipass cell, and (c) spot configurations for astigmatic configuration (adapted from Zahniser et al., 1997).

A. GASES

A second multipass cell configuration is the Herriott cell (Herriott et al., 1964; Herriott and Schulte, 1965). This is particularly useful for coherent light sources such as lasers used in tunable diode laser spectroscopy but has also been used with incoherent light sources using optical fibers cemented to a ball lens at the entrance to the cell (Zahniser, personal communication). Two spherical mirrors are separated by a distance close to their radius of curvature, and the light beam enters through a hole in one of them, directed in an off-axis direction. After multiple reflections between the two mirrors, the light beam exits through the same hole as it entered, but at a different angle (Fig. l l.2a). The beam remains collimated throughout, in contrast to the White cell system, and gives the spot pattern shown in Fig. l l.2b. The path length is changed by changing the distance between the mirrors; in practice this means that this design is most useful for fixed path length systems. An astigmatic variant of the Herriott cell designed for use in ambient air studies is shown in Fig. l l.2a and described by McManus et al. (1995) and Zahniser et al. (1997). In this design, the two mirrors have different radii of curvature, giving the spot patterns shown in Fig. 11.2c. The spots more evenly fill the mirror, so that for a given number of passes, the spots are more widely spaced, or conversely, more passes can be obtained without problems of beam overlap (McManus et al., 1995). Major advantages of such cells are that they are relatively easy to align and folded optical paths can be obtained in small volumes. This is important when small amounts of sample are available, for example, in laboratory studies or when a fast response is needed; cells of smaller volume can be pumped out faster, giving shorter residence times in the cell. (2) F T I R Fourier transform infrared spectroscopy has been used for many years to measure atmospheric gases. Because F T I R has become such a common analytical method, we do not describe the technique itself here but rather refer the reader to several excellent books and articles on the subject (e.g., see Griffiths and de Haseth, 1986; Wayne, 1987). For reviews of some atmospheric applications, see Tuazon et al. (1978, 1980), Marshall et al. (1994), and Hanst and Hanst (1994). A problem in the application of F T I R to ambient air is that water vapor, CO2, and CH 4 are all present in significant concentrations and absorb strongly in certain regions of the spectrum. As a result, the spectral regions that are useful for ambient air measurements are 760- to 1300-cm -1, 2000- to 2230-cm -1, and 2390to 3000-cm- 1.

5 51

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(6)

and if the water is stripped from the cluster, the ion HT § remains and can be separated and detected by MS. That is, an (M + 1) peak results, where M is the molecular weight of the trace gas. Carbonyl compounds such as acetone have been detected with high sensitivity using this approach. For example, Fig. 11.17 is a schematic diagram of a mass spectrometer flown in an aircraft that was used to detect acetone in the lower stratosphere (Arnold and Hauck, 1985). The air is drawn in and the acetone reacts with H 3 0 + ( H 2 0 ) n formed using a high-frequency glow discharge ion source. The (M + 1) peak corresponding to [H. CH3C(O)CH3] § is sampled downstream through an inlet orifice into the quadrupole mass spectrometer. Near the earth's surface, there is sufficient ammonia in air that it undergoes a reaction with H 3 0 + ( H 2 0 ) , to form NH~-(H20),, and this can also act as an

Nitrate clusters with H 2 0 a n d / o r HNO 3 such as NO~(HNO3) n are common in the atmosphere (Perkins and Eisele, 1984). Proton transfer to such clusters can occur, but clearly, the trace gas must be more acidic than HNO 3. This limits the number of trace gases that can be ionized through this mechanism but includes the important atmospheric species H2SO4 and methanesulfonic acid, CH3SO3H (Tanner and Eisele, 1991; Viggiano, 1993). In short, positive and negative ions in air containing the trace gases of interest can be formed through discharge techniques and ions of the trace gases of interest generated via ion-molecule reactions. As discussed in more detail later, this approach has been used quite successfully to measure a number of species in air, including formic acid, acetic acid, dimethyl sulfide, and C12 (Spicer et al., 1994a, 1998). An alternate method is to add another compound to the mass spectrometer inlet, ionize this added species, and use its ion-molecule reactions to form ions a n d / o r ion adducts of the species of interest. This has been used to measure HONO, for example, in air where a chloride ion adduct of H O N O is formed when CHC13 is added in the corona discharge region (Spicer et al., 1993a). Other examples include the measurement of HNO 3. For example, as described in Section A.4a(5), radioactive ionization of added SF6 generates daughter ions that react with SiF4 to give SiF[. The SiF[ forms an adduct with HNO 3 and this adduct can be used to measure HNO 3 in air (Huey et al., 1998).

(2) Laser photoionization Another ionization method with great potential for ambient air applications is

A. GASES

563

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F I G U R E 11.17 Schematic of mass spectrometer used for stratospheric measurements (IG = ion getter pump, PS = pressure sensor) (adapted from Arnold and Hauck, 1985).

laser photoionization (see Letokhov (1987) and Pfab (1995) for reviews). Trace gases can be ionized if sufficient energy in the form of light is pumped in; for example, polycyclic aromatic hydrocarbons (PAH; see Chapter 10) in combustion mixtures have been measured by two-photon ionization at 248 nm (e.g., Castaldi and Senkan, 1998). In practice, for application to ambient air, efficient photoionization requires the use of pulsed lasers and multiphoton absorption methods. The terms "multiphoton ionization," or MPI, and "resonance-enhanced multiphoton ionization," or REMPI, are used to describe these processes. Figure 11.18 illustrates the principles of application of R E M P I to NO (discussed in more detail later). The electronically excited states of NO are shown in Fig. ll.18a and some potential ionization schemes in Fig. ll.18b (Pfab, 1995). Pulsed tunable lasers with wavelengths from ~190 to 1000 nm and spectral resolutions of 0.1 cm -1 are readily available. To ionize NO, the absorption of two, three, or four photons is needed. The first photon excites the NO into an intermediate state from which it is ionized using a second or, in some cases, two more photons. The transitions are described as an (n + m) transition, where n is the

number of photons that need to be absorbed simultaneously to reach the intermediate state and m is the number of photons to ionize the molecule from that state. The wavelength/energies of the photons involved in the various steps may be the same, which is referred to as a "one-color" process, or different, a "two-color" process. In the two-color case, the second photon is primed to indicate it is a different wavelength than the first photon. For example, in Fig. ll.18b, ionization via the A state can occur either by a (1 + 1) process using 226 nm or by a (1 + 1') process, where the A state is reached using 226 nm and ionization from this state occurs using 308 nm (Pfab, 1995; Hippler and Pfab, 1995). (The dashed arrows show transitions used for detecting NO by laser-induced fluorescence; see Section A.4a(1).) The high spectral resolution of laser radiation provides selectivity. For example, Figure 11.19 shows the R E M P I spectrum of the NO X (0,0) ~ A band using a (1 + 1') process with 226- and 308-nm light to photoionize NO (e.g., see Pfab, 1995; and Lee et al., 1997). As the laser is tuned into resonance with a particular rotational transition in this band, ions are generated and detected using a conventional electron multiplier. Clearly, high selectivity is possible by tuning on and off

564

11. A N A L Y T I C A L

METHODS

FOR GASES

AND

PARTICLES

(1+1)

IN THE

(1+1')

ATMOSPHERE

(2+2)

(2+1)

(2+1)

+

1s

Ionization limit

--V

60,000

70,000 II

v

D2'T-"+

E 60,000 I l l c.)

>, c~ r uJ

/ C2n

/~1A

9

A2s

!_... I~,,_//

0.10

o >, 131 r uJ

c

40,000

I I

v

/Predissociation /7" iim;

~.\

50,000

E

--

J

t,D 04 04

20,000

E r 04 04

X2H

Et-

04 04

%/, r~/

J~

J

E r E

o4

04 to

03

hv=

1187-220 nm

v,,v

E t-

03

I I

I

0.15 "c(nm)

FIGURE 11.18 (a) Potential energy diagram and (b) REMPI schemes for the excitation and ionization of NO (adapted from Pfab, 1995). The ~ indicates a virtual state. The dashed arrows show other transitions used to detect NO in the atmosphere by laser-induced fluorescence (see Section A.4.a(1)).

the rotational transitions. (Note, however, as discussed later, interferences may result if NO is generated by photolysis of other species such as NO 2 in the laser beam.) Pumping low-lying Rydberg states as the resonant intermediate is generally preferred over excitation of higher valence states because the latter often predissociate, undergo rearrangement, or intersystem cross before the subsequent absorption of photons can form

c.9 .

.

.

.

.

.

-~ . . . . . . . . . >,

. . . . . . . . . .

~ R21 2, ,~R11 + Q21 2 l r i " ~ l Q l l + P21 2

(~

E

89 I

a

I

I

II

'

Pll

N

ions. The Rydberg region can be reached by the simultaneous absorption of two or three photons, but with a loss of sensitivity (the absorption cross sections for two-photon absorption are of the order of 10 .5o cm 4 compared to 10 -is cm 2 for one-photon absorption). On the other hand, pulse energies from conventional excimer or N d : Y A G p u m p e d tunable laser systems are one to two orders of magnitude higher, since no second harmonic generation stages are required. Other interferences from the use of two- or threephoton resonant excitation compared to (1 + n) processes are photolysis of the analyte at the energy of the first (or second) photon and broadening of the absorption spectra due to the higher photon fluxes employed. Table 11.5 summarizes some potential R E M P I processes for the m e a s u r e m e n t of species of atmospheric interest (Pfab, 1995). This ionization technique clearly has a great deal of potential, although to date it has not been applied extensively to measurements in ambient air. c. M a s s Filters

I

225.8

225.9

226.0

226.1

226.2

226.:3

226.4

226.5

Wavelength (nm) FIGURE 11.19 Two-color (1 + 1') REMPI spectrum of the NO (X2II, v' = 0 --+ A2s +, v' = 0) band at 300 K (adapted from Pfab, 1995).

Two types of mass analyzers have been used extensively in atmospheric applications: quadrupole mass filters and time-of-flight (TOF) instruments. The use of ion traps is also being increasingly explored for this application. For the fundamental principles of mass

A. GASES TABLE

11.5

565

S o m e P o t e n t i a l R E M P I T r a n s i t i o n s for I o n i z a t i o n of M o l e c u l e s of A t m o s p h e r i c I n t e r e s t a

Type of process + + + +

1 1' 1' 1

Resonant transition

Excitation wavelength (nm)

X --+ A X --+ A X --+ B, C, D X --+ C

226 226 + 308 b 191 + 355 or 226 380

NO

1 1 1 2

CO

1 + 1' 2 + 1 2 + 1

X --+ B X --+ B X -+ E

115 + 345 229-231 215

NO 2

1 + 1' + 1 1 + 1' + 1"

3p Rydberg + 3p Rydberg +

484 + 248 + 484 c e.g., 482 + 275 + 460 d

N20

2 + 1

3p Rydberg +

230-250

NH 3

2 + 1

J~ --+ B

340-355

--+ b

286-289

CH3CHO

2 + 1

n --+ 3s

346-365

CH3COCH 3

2 + 2 2 + 1

n ~ 3s n ~ 3p

372-392 322-342 e 259 f

C6H 6

1 + 1

S O --+ S 1

Toluene

1 + 1

S O --+ S 1

267

Naphthalene

1 + 1

S O --+ S 1

270-310

d Campos et al. (1990). e McDiarmid and Sabljid (1988). f Boesl et al. (1980).

a A d a p t e d from Pfab (1995). b Hippler and Pfab (1995). c Benter et al. (1995).

separation using these techniques, see standard analytical chemistry texts such as Skoog et al. (1998). A potential limitation in the application of MS to near-surface measurements is the tremendous number of compounds in the atmosphere, particularly organics, and hence the increased complexity of interpretation of the single mass spectrum. In the MS ion source, the use of particular ion-molecule reactions to form the ions of interest or the ionization of one selected com-

Corona discharge needle . . . . ~ - -

N2 gas interface ,I, _~_ H/L~~,

,

-
2.5 /xm), that is, those in the coarse particle range. They include collection by gravitational sedimentation (e.g., dustfall jars) as well as by centrifugal

612

11. ANALYTICALMETHODSFOR GASESAND PARTICLESIN THE ATMOSPHERE

sedimentation, which allows collection in the submicrometer range (e.g., centrifuges and cyclone collectors). Gravitational sedimentation only collects the large particles that settle out of the atmosphere fairly quickly. This dustfall generally consists of particles that are relatively large and, as such, are not particularly relevant to the focus of this book. Thus dustfall collectors will not be discussed further. The principle of centrifugal collection is, of course, well known. Collection of particles using centrifugation involves passing the aerosol at a controlled rate through a rapidly spinning air mass. Collection of particles in ranges as small as ~ 0 . 1 - 1 /xm has been reported using this technique. The cyclone collector, a modification of the centrifuge technique, is based on bringing the air samples into a stationary cylindrical vessel at high velocity; a vortex is formed by the entry of the air tangential to the length of the vessel and particles in this vortex are subjected to a centrifugal force that depends on their size (Fig. 11.59). As a result, particles of different sizes are deposited at different locations along the length of the cyclone separator. Although cyclone collectors have been applied to size distribution measurements by using a series of cyclones in parallel, each having a different cut size, they are most commonly used as precollectors to remove larger particles ( ~ 3- to 30-/xm diameter) before the air sample enters a device such as an impactor designed for the measurement of particles in smaller size ranges.

Air out

Air in

2. M e a s u r e m e n t M a s s a n d Size

of P h y s i c a l C h a r a c t e r i s t i c s :

a. M a s s

The total mass of particles per unit volume of air is one of the major parameters used to characterize particles in air and, along with size, is the basis of air quality standards for particulate matter (see Chapter 2). Methods of mass measurement include gravimetric methods, /3-ray attenuation, piezoelectric devices, and the oscillating microbalance.

(1) Gravimetric methods The most straightforward method of determining the particle loading of the atmosphere is to weigh a collection substrate such as a filter before and after sampling. However, care must be taken to be sure that both temperature and relative humidity are carefully controlled when weighing both the loaded and clean substrate. As discussed earlier, some filters a n d / o r the collected particles are hygroscopic and unless care is taken to equilibrate them at a fixed temperature and relative humidity, the change in water content may completely mask the change in mass due to the particles. In addition, problems such as forces due to static electricity on the filter that interfere with accurate weight measurements must be controlled. Finally, particulate loading can change the sampling air flow rate and lead to large errors in determining the actual volume of air sampled. (2) fl-Ray attenuation /3-Particle beams (electrons) emitted from a radioactive source are attenuated when they pass through a filter on which particulate matter has been collected. (/3-particle beams rather than aparticle beams or y-rays are used because a particles do not penetrate typical thicknesses of filter well and y-rays are too penetrating and hence would require large sample thicknesses.) Figure 11.60 shows a

13-ray detector .....

FIGURE 11.59 Schematic diagram of one type of cyclone collector (adapted from Ayer and Hochstrasser, 1979).

FIGURE 11.60 Schematic diagram of a typical /3-ray attenuation device for measuring particulate mass.

B. PARTICLES schematic of a /3-ray attenuation device, which consists essentially of a / 3 source such as 14C, a /3 detector, and a means of positioning the filter paper containing the particulate matter between the two. The ratio of the transmission of /3-rays through a clean and loaded portion of the filter, respectively, is related to the particle loading via a B e e r - L a m b e r t type of relationship:

I / I o = e -~x.

(S)

I 0 and I are the intensities of the /3-rays that have passed through the clean and loaded portions of the filter, respectively, X is the thickness of the deposit, and /z is an attenuation constant that is approximately proportional to the density ( p ) of the material deposited. The mass per unit area deposited on the filter, given by pX, is the parameter desired in this measurement. Rearranging Eq. (S), one obtains

l n ( I o / I ) = ( t z / p ) pX.

(T)

The parameter tz/p is a constant known as the mass absorption coefficient; with the assumption that this is independent of the type of absorbing particles (an assumption that generally holds well enough to cause < 10% uncertainty), the value of ln(Io/I) is directly related to the parameter of interest, p X = mass per unit area. Such measurements can be carried out on filters with different cutoff sizes to obtain size resolution as well (e.g., see Spagnolo and Paoletti, 1994).

(3) Piezoelectric microbalance The piezoelectric microbalance is a resonant frequency device. The piezoelectric effect is the development of a charge on some crystals such as quartz when a stress is applied; the stress may be mechanical (e.g., added weight) or electrical. Such crystals may be used as part of a resonance circuit to provide very stable, narrow-band frequencies; the quartz crystal is plated on two sides with a thin conducting layer and leads are connected to the resonance circuit so the crystal replaces an L C network. The obtained frequency of vibration (v o) depends on a number of parameters of the crystal but is usually ~ 5 - 1 0 MHz. However, if a mass (Am) becomes attached to one side of the crystal, it changes the resonant frequency by an amount Av o such that Au0/u 0 = A m / m

(U)

as long as the increase in mass Am is much smaller than the mass (m) of the active part of the crystal. Particulate matter from ambient air can be deposited on the crystal in various ways, for example, by

613

using it as an impaction device. The mass of the collected particles can then be determined by following the change in the frequency. Alternatively, a reference crystal held at the same temperature and pressure as the crystal on which the particles are collected can be used, and the difference in frequencies between the two crystals can be determined. The piezoelectric microbalance is very sensitive, capable of detecting ~ 1 0 - s - 1 0 -9 g. The particles collected on the crystal surface can be chemically analyzed after collection using surface-sensitive techniques. One limitation is possible overloading of the crystal; thus when the collected mass reaches ~ 0 . 5 - 1 % of the mass per unit of the crystal, the surface must be cleaned.

(4) Oscillating microbalance The tapered-element oscillating microbalance is based on a similar principle to the piezoelectric microbalance. A hollow glass piece is mounted with the wider end fixed and a filter attached to the narrower end. The tip oscillates at a particular frequency in an applied magnetic field. As particles are collected on the filter, the resonant frequency of the glass tends to decrease. A feedback is used to maintain the oscillation frequency and provides a measurement of the collected mass (e.g., see Patashnick and Rupprecht, 1991). Good agreement has generally been observed between measurements made using the tapered-element oscillating microbalance and HiVol filter methods (e.g., Eldering and Glasgow, 1998). b. Size There are several different approaches that are commonly used to determine particle size distributions in air. One of them, impaction, has been discussed earlier. Multistage impactors with different cut points are used extensively to obtain both mass and chemical composition data as a function of size for particles with diameters > 0.2 /~m. Others, including methods based on optical properties, electrical or aerodynamic mobility, and diffusion speeds, are described briefly in the following section. The condensation particle counter (CPC) is used as a detector in combination with some of these size-sorting methods. The reader is cautioned to keep in mind that atmospheric particles are not all spherical nor even necessarily simple in shape. Thus, as discussed in Chapter 9.A, the term size cannot be uniquely defined for atmospheric particles. As a result, a measurement of the distribution of sizes using an impactor that is based on inertial characteristics, for example, may not give the same results as a size measurement based on optical techniques that use light scattering. With this caveat in mind, let us examine the most commonly used

614

11. A N A L Y T I C A L M E T H O D S F O R G A S E S A N D P A R T I C L E S IN T H E A T M O S P H E R E

methods of determining the size distribution of atmospheric particles.

(1) Optical methods Optical counters, optical microscopy, and electron microscopy fall under this heading. A review of optical methods is given by Baron et al. (1993). Single-particle optical counters. These instruments are used to measure particles in the ~0.1- to 10-/zm range by measuring the amount of light scattered by a single particle (Martens and Keller, 1968). As discussed in Chapter 9.A.4, the amount of this Mie scattering depends not only on the refractive index but also on the radius of the particle; hence the intensity of scattered light is a measure of the particle size. Assuming that the particles are spherical, smooth, and of known refractive index, one can calculate, using Mie theory, the intensity of scattered light of wavelength /~ at various angles (0) to the incident beam for a particle of a given size. Integrating over all scattering angles and wavelengths (since "white" incandescent sources are normally used in these instruments), one obtains the theoretical response of the single-particle counter, that is, the curve of scattered light intensity as a function of the particle diameter. Typical theoretical response curves are shown in Fig. 11.61 (Cooke and Kerker, 1975). Calibration of these single-particle counters is usually carried out using monodisperse polystyrene latex or polyvinyl latex spheres, which are available in sizes from ~0.1 to 3 / z m and have a refractive index of 1.6; alternatively, aerosols with lower refractive indices may be generated from liquids such as dioctyl phthalate (m = 1.49). Whitby and Willeke (1979) discuss the

importance of instrument calibration using standardized aerosols with an index of refraction as close as possible to the sample being measured; since the refractive index of atmospheric particles varies from 1.33 for water to 1.7 for minerals, they recommend using a calibration aerosol with m ~ 1.5. Because light scattering is very dependent on the particle shape, when measuring irregularly shaped particles such as coal dust, one should calibrate the instrument with aerosols generated from the same material. Figure 11.62, for example, shows the instrument response as a function of particle diameter for an ideal calibration aerosol of dioctyl phthalate and for coal dust particles. Potential problems with using single-particle counters in ambient measurements and ways to minimize these are discussed in detail by Whitby and Willeke (1979). Optical counters allow relatively rapid measurements of the size distribution and, unlike some of the other methods of size fractionation, include volatile particles in the measurement. However, some care must be taken in interpreting the detailed shape of the size distribution spectrum because of some anomalies that have been observed; for example, around the 1-/zm region, interference from light that is reflected or refracted from the front and back of the particle gives a "knee" in many calibration curves of number of particles versus their diameter (LBL, 1979). Electron and optical rnicroscopies. Counting the particles and measuring their sizes can be done by optical or electron microscopy, the former for particles with diameters from ~0.4 /zm to several hundred microns,

/

10_ 10 (1) co E

L

0

o_ co

/"

t~ t-0"1

v/

1--

DOP (m = 1.49)

/

L

133-

(1) E O o ,

1

,

./

_

tt

Coal dust

; 9149176

o

1 54-"//j 9 .~~/

o

(1) E~ E Or)

9

" ..... -:" Refractive ....: index 9 "...J -'x_ - - > ~ .... m = 1 5 4 005J

0.1-

1.70

_

0.1 0.1

, ,I. . tll

,

1 Particle d i a m e t e r

10

20

(gm)

FIGURE 11.61 Theoretical response of a typical single-particle counter (adapted from Whitby and Willeke, 1979; data from Cooke and Kerker, 1975).

0.01 0.1

i

i illllll

i

i illllll

1 Particle d i a m e t e r

i 10

i i 50

(gm)

FIGURE 11.62 Experimental calibration curves for a commercial single-particle counter and two types of calibration aerosols: dioctyl phthalate (DOP) and coal dust (adapted from Whitby and Willeke, 1979).

B. PARTICLES and the latter for particles from ~0.001 /xm and larger. The particles must be on the surface of the substrate and form less than a monolayer to minimize overlap of the particles. For electron microscopy, the sample and substrate must also be able to be subjected to high vacuum, heat, and electron bombardment without degradation over a period of time sufficient to make the measurement. Because of the instrumental requirements, these are usually not routine monitoring techniques. However, unlike other methods, they give detailed information on particle shapes. In addition, chemical composition information can be obtained using transmission electron microscopy (TEM) or scanning electron microscopy (SEM) combined with energy-dispersive spectrometry (EDS). The electron beam causes the sample to emit fluorescent X-rays that have energies characteristic of the elements in the sample. Thus a map showing the distribution of elements in the sample can be produced as the electron beam scans the sample. For example, Fig. 11.63 shows the T E M image of an NaC1 crystal (Fig. 11.63A) and of the same crystal after exposure to gaseous H N O 3 (Fig. 11.63B) and then small amounts of water vapor (Fig. 11.63(2) (Allen et al., 1996). After the crystal is dried, the formation of new microcrystallites attached to the NaC1 is observed. These can be shown using EDS to be crystals of NaNO 3. Thus Fig. 11.64b shows the EDS spectra obtained from the larger, original NaC1 crystal (but after exposure to H N O 3 and water vapor) and of small microcrystallites attached to it (Fig. 11.64a). Only Na and C1 are seen in the first case, but Na, N, and O in the second, and in the correct ratio for NaNO 3. The H N O 3 has reacted with the NaC1 surface to generate metastable surface nitrate ions. Exposure to water vapor generates a mobile quasi-liquid layer on the surface that contains Na+, CI-, and NO~- ions. On drying, the ions segregate to form separate microcrystallites of NaNO 3 and a flesh NaC1 crystal (Allen et al., 1996). The combination of SEM with EDS has also been applied to atmospheric particles (e.g., Pdsfai et al., 1995; Anderson et al., 1996; McMurry et al., 1996; Ganor et al., 1998). For example, individual sea salt particles were analyzed using T E M combined with EDS as well as selected-area electron diffraction (SAED) by Pdsfai et al. (1995) and Anderson et al. (1996). The crystal shapes correlated well with the chemical composition determined using EDS and SAED. For example, cubic crystals of NaC1 were observed. Sulfate occurred in either rod-shaped crystals, which had significant concentrations of (Mg + K + C a ) c o m p a r e d to Na, or tubular crystals, with much smaller concentrations of these three metals. In the latter case, the EDS showed

615

A

1.0pro

C

FIGURE 11.63 TEM images of an NaC1 crystal (A) before reaction and (B) after reaction with gaseous HNO 3 (1.2 • 1015 cm-3 for 15 min) and then (C) exposure to water vapor ( 1 /xm at altitudes of 17-19 km from 1976 to 1984 (adapted from Zolensky et al., 1989).

C. CHLOROFLUOROCARBONS

classified by their major chemical composition. The largest increase was in aluminum-rich particles, thought to be from the exhaust of solid rockets using fuels containing an aluminum additive and perhaps from ablation of the spacecraft paint. Silicate particles also increased, which may be due to ablation of the solid rocket motors and the spacecraft paint. Iron is found in exhaust and in alloys used to make the solid rocket motors. Noteworthy in Fig. 12.10 is the lack of significant increase in chondritic particles from extraterrestrial sources. The fact that they did not also increase over this period of time suggests that the measured increase in the other particles is a real effect rather than a sampling artifact, which has been of concern (Zolensky et al., 1989). In short, the launch of space vehicles is an additional source of emissions to the stratosphere, along with the proposed fleet of HSCTs.

C. C H L O R O F L U O R O C A R B O N S 1. T y p e s , N o m e n c l a t u r e ,

and Uses

Chlorofluorocarbons, that is, compounds containing chlorine, fluorine, carbon, and possibly hydrogen, have been used extensively in the industrialized nations in the past decades primarily as propellants in aerosol spray cans, as refrigerants, and as blowing agents, for example, for producing polyurethane foam. Their chemical characteristics have made them ideally suited for such uses in that they are generally nontoxic and chemically inert. Thus, they can be used around open flames, and leaks in refrigeration units do not present a health hazard as older units operated on coolants such as SO 2 once did. The principal chlorofluorocarbons which have been used are CC13F, CC12F2, and CHC1F 2. These are often referred to as CFC-11, CFC-12, and CFC-22, respectively; alternatively, the abbreviations F-11, F-12,

TABLE 12.2

669

and F-22 are used, after the Dupont trade name Freon. The numbers are based on a system developed by chemists at DuPont. The first number is the number of hydrogen atoms, plus one, the second gives the number of fluorine atoms, and the remainder are chlorine atoms. For chlorofluorocarbons containing two or more carbon atoms, a three-digit numbering system is used; the first digit gives the number of carbons minus one, and as for the one-carbon compounds, the second is the number of hydrogens plus one and the third is the number of fluorine atoms. Thus, CC12FCC1F2 is F-113, or CFC-113 (whose use for cleaning electronic components increased in the mid to late 1970s), and CCF2CC1F2 is F-114, or CFC-114. When more than one isomer is possible, the most symmetrical one has only a number; the letters a and b are added to distinguish the less symmetrical isomers. To provide some historical perspective, Table 12.2 gives the distribution of sales by use for CFC-11 and CFC-12, broken down into aerosol propellants, blowing agents, refrigerants, and other uses, from 1976 to 1992. On a global basis, the use as aerosol propellants accounted for more than half of the sales of CFC-11 and CFC-12 in 1976. Essentially all of this as well as most of that used in other applications has been released into the atmosphere. Figure 12.11 shows the history of the estimated annual global release rates of CFC-11 and CFC-12 from 1952 to 1980. Over this period of about three decades, the release rate into the atmosphere increased dramatically until about 1974, as these compounds found increasing use in our industrialized society. Figure 12.12 shows similar data from 1972 to 1992. The decrease in emissions as the Montreal Protocol took effect can be clearly seen (see Chapter 13.A). Data from 1986 to 1996 from the European Union show similar trends (McCulloch and Midgley, 1998). Figure 12.13 shows the trend in surface CFC-11 and CFC-12 concentrations in air measured at latitudes

Total CFC Sales by Use from 1976 to 1992 (in Thousands of Metric Tons) a Use

Compound

Year

Aerosol propellant

Blowing agent

CFC-11

1976 1984 1992

195.0 97.5 8.2

97.0 173.9 155.5

CFC-12

1976 1984 1992

237.3 121.3 14.7

15.0 49.2 14.3

a From AFEAS (1993).

Refrigeration

Other

Total

26.9 23.9 17.8

21.0 17.0 5.0

339.9 312.3 186.5

127.8 187.5 177.4

30.7 24.1 9.8

410.8 382.1 216.2

670

12. CHEMISTRY OF THE S T R A T O S P H E R E

'7,

CFC-12

280

C)3

240 -

~z~

400 -

Q.

200 160 C,f

'7 O 1.1_

o,,

/ /

00-

-"

11

,

I

I

I

I

I

I

I

I

I

I

I

I

I

I

I

I

500

E

o 0

200

u_ o

o

350 _

-

300

/ /

E O

~9

1 O0

eE

FS,"

"

I I 78

I I 80

I I 82

"1

1952

I

1960

I I 84

I I 86

I I 88

I I 90

I I 92

Year

/

0

0.1e

> 0.1e

Reaction

C1ONO 2 + HC1 --+ C12 + HNO 3

0.2

N205 + HC1 --+ C1NO2 + HNO 3

0.03

HOC1 + HC1

0.3

--+ C12 + H 2 0

C1ONO 2 + H20 -+ HOC1 + HNO 3 N205 + H20 --+ HOC1 + HNO 3 HO2NO 2 +

>0.1 0.02

3•

Liquid H2SO4-HNO3-H20

-3 0.1

,-~10-4-10 -2i

,-~10-2-10 -3e

3 • 10 -4

,-~10- 2

-~1 X 10-4-0.1 f 0.1g HOC1 + HNO3,

HC1 + HOC1 aerosol> C12 -+- U 2 0 , BrONO2 + H 2 0 HC1 + HOBr

(46) (39) (41) (42)

aerosol> HOBr + HNO3,

(47)

aerosol> BrC1 + H 2 0 .

(48)

Laboratory studies of the uptake of C10 into sulfuric acid (Martin et al., 1979, 1980), taken in light of a deficit in the inorganic chlorine budget at ~17 km after the Mount Pinatubo eruption, led Jaegl6 et al. (1996) to propose that a heterogeneous reaction of C10 on sulfuric acid aerosols to form perchloric acid, HC104, may also occur. The kinetics of these reactions in liquid solutions characteristic of the stratosphere, such as concentrated H 2 S O 4 - H 2 0 or ternary solutions with HNO3, depend on temperature as expected and in some cases at least, on acidity as well. For example, Donaldson et al. (1997) have shown that the second-order rate constant for the

692

12. CHEMISTRY OF THE STRATOSPHERE

a

1000

100 t~ O

Pre-Pinatubo

100

Mt. Pinatubo

O O

10

%

"-1

.ca t~ e~

0.1

z

E O

E "-

I

0.01

g

c(((c(((c((()

-

~

O O

~.

9 0

n

Ruiz

St. Helens, ~11~.Alaid

O .m .4.-,

N3 E

r

rr

1

10

I i I i I i

1979

9

[~ []

O

O

m

El Chichon

-

O .i.-, O r

Main aerosol layer, August 20, 1991

100

m

,!--

I

b

1000

10 m

O0

tf)

~,

%

mO

1

1981

1983

1985

1987

. i I i 1989

1991

1993

1 Time (years)

0.1

F I G U R E 1 2 . 2 9 Ratio of particle surface areas to those in the winter of 1978-1979, from 1979 to 1995 at 45~ and 20-km altitude. Data are not available for 1982-1984. (Adapted from Solomon et al., 1996.)

I

0.01 1000

z

100

z

May 7, 1993

10

0.1 0.01

0.

P

01

I

0.1

1

5

Particle radius (gm)

F I G U R E 1 2 . 2 8 Particle surface area distributions in the stratosphere (a) before Mount Pinatubo eruption, (b) August 20, 1991, over California, and (c) May 7, 1993, over California (adapted from Russell et al. (1996) and Goodman et al. (1994)).

reaction (42) of HOC1 with HC1 increases by more than an order of magnitude from 2.8 • 105 to 4.9 • 106 L mol-1 s-1 at 250 K as the weight percent sulfuric acid in a mixture with water increased from 49 to 67%. They propose that this is due to increasing protonation of the HOC1 to H2OC1 § in highly acidic solutions, followed by reaction of the protonated form with CI-. Of these heterogeneous reactions, the hydrolysis of N205 is particularly important in midlatitudes. For example, Fig. 12.30 shows the measured NO~ (NO~ = N O + N O 2) to NOy ( N O y = N O ~ + H N O 3 +N205 + ...) ratio at different latitudes compared to the predicted ratio using a gas-phase model as well as to a model that incorporates the N205 hydrolysis on aerosol particles (Fahey et al., 1993). Clearly, the inclusion of this reaction is necessary to bring the measurements

and models into better agreement (e.g., Rodriguez et al., 1991; Fahey et al., 1993; Kondo et al., 1997; Sen et al., 1998). This reaction (46) on aerosol particles of course occurs not only in midlatitudes but at the poles as well (Keys et al., 1993). As discussed by Morris et al. (1997), the photolysis of H N O 3 to OH + NO 2 and its reaction with OH to generate NO 3 are the major sources of NO X during the day. With large aerosol concentrations, the loss of NOx at night is largely through the formation of N205 (NO 3 + NO 2 ~ N205), followed by the hydrolysis of N 2 0 5. These reactions shift the N O J N O y ratio to smaller values in the presence of high particle concentrations.

... .... "'v..'"'. m~

o z

o

.., ...... ... ....... . .... .........

0.2

0z

gas phase only m IIo

........

.. . . . . . . . . . . . . . " ' . . v ...... .'.""'v-":;"

X

0.1 0 36

with N 2 0 5 + H20 I

I 38

I

I 40

I

I 42

I

I 44

I

I 46

I

I 48

I

50

Latitude (~

1 2 . 3 0 Comparison of predicted ratio NOx/NOy as a function of latitude at 19.5-20.5 km using only gas-phase chemistry (dotted line) or with the N205 hydrolysis on aerosol particles (solid line) compared to measured values shown as unfilled circles (from Fahey et al., 1993).

FIGURE

C. CHLOROFLUOROCARBONS

There are a number of measurements documenting changes in NO X and NOy in the stratosphere after the Mount Pinatubo eruption and which have been attributed to the removal of oxides of nitrogen due to reactions on aerosol particles. For example, a decrease in stratospheric NO 2 after the eruption followed by a return to normal levels has been reported (e.g., see Van Roozendael et al., 1997; and De Mazibre et al., 1998). Similarly, NOx decreases of up to 70% were reported, as well as increases in gaseous H N O 3 (much of that produced on the sulfate particles is released to the gas phase) (e.g., see Coffey and Mankin, 1993; Koike et al., 1993, 1994; David et al., 1994; Webster et al., 1994; and Rinsland et al., 1994). Not only does this heterogeneous hydrolysis alter the NO~ reactions that can lead to ozone destruction or formation, but it also changes the halogen cycles because less NO 2 is available to trap C10 as the nitrate. In addition, HO~ levels are increased. Thus there is relatively more C10 and the C10~ cycles leading to ozone destruction become more effective (e.g., McElroy et al., 1992; Avallone et al., 1993a, 1993b; Schoeberl et al., 1993b). Heterogeneous chemistry would be expected to shift the partitioning of chlorine away from HC1 toward more active forms. While increasing ratios of HC1/Cly subsequent to the Mount Pinatubo eruption have been reported in some studies, suggesting decreasing contributions from heterogeneous chemistry after the eruption (e.g., an increase in HC1/Cly from 0.40 in late 1991 to 0.70 in 1996 based on in situ measurements; Webster et al., 1998), satellite data suggest smaller changes in HC1/Cly (16 _+ 9% from 1992/1993 to 1995/1996) (e.g., Dessler et al., 1997). This effect can be seen in the midlatitude stratospheric measurements of Keim et al. (1996) shown in Fig. 12.31. In the tropopause region (shown by the

693

dotted line), the aerosol surface area increases. A significant increase in C10 and decrease in NO is seen at the same time, while NOy increases. This was attributed to the heterogeneous reaction of C1ONO 2 with HC1 to form H N O 3 on the aerosol particles. The C12 product generates C1 atoms, which react with 0 3 to give enhanced C10. Both C1ONO 2 and particle H N O 3 are measured as NOy so that conversion of one to the other should not lead to enhanced NOy. The latter was attributed to highly efficient sampling of sulfate aerosol particles containing nitrate. Similar observations of enhanced C10 and suppressed NO have been reported in other studies as well (e.g., Fahey et al., 1993; Toohey et al., 1993). Dessler et al. (1993) invoked heterogeneous reaction to explain measured C10 concentrations that were larger than expected for gas-phase chlorine chemistry. The hydrolysis reaction ( 4 6 ) o f N20 5 under many conditions in the atmosphere becomes limited by the rate of N20 5 formation, which only occurs at a significant rate at night (because of the rapid photolysis of the NO 3 precursor during the day). Hence under these conditions, reactions (39), (41), and (42) followed by photolysis of the chlorine-containing products become primarily responsible for the removal of gas-phase NOy and increase in C10 (Keim et al., 1996). Figure 12.32 shows the results of model calculations of the effects of the increased aerosols for October 1986 at 43.5~ (Solomon et al., 1996). The calculated change in the odd-oxygen loss rate when the measured aerosol particle surface area is incorporated into the

32

27

';i..

ciS;..;...

g

4ta

decomposition

, H C H O + Br.

(55)

The reaction of the alkoxy radical with 0 2 to give formyl bromide, HC(O)Br + HO2, was shown to be negligible compared to its decomposition. In short, the oxidation of brominated organics by OH and other species also ultimately generates atomic bromine. Oxidation of the organic bromine compounds to generate inorganic bromine is sufficiently fast that after a few years, inorganic forms comprise almost all of the stratospheric bromine. For example, Wamsley et al. (1998) showed that 97.5% of stratospheric bromine was in the form of inorganic species in an air parcel that had been in the stratosphere for about 6 years. The atmospheric chemistry of other bromochloromethanes that are present in smaller amounts in the troposphere (Fig. 12.44) is discussed by Orkin et al. (1997) and Bilde et al. (1998). Bromine is more effective, per atom, in destroying 0 3 than is chlorine for several reasons. First, unlike chlorine atoms, atomic bromine does not react rapidly with organics. For example, the reaction of Br with CH 4 is too slow to represent a significant loss process, unlike the analogous chlorine atom reaction. Atomic bromine does react with H C H O at a reasonable rate ( k 298K = 1.1 X 10-12 cm 3 molecule- 1 s - l ) as well as with higher aldehydes (whose concentrations in the stratosphere are not large enough for this to be important, however). As a result, Br is converted into its reservoir species, HBr, primarily through its reactions with HO 2 and HCHO:

(het)

Br + HO 2 ~ HBr + 02, F I G U R E 12.45 Schematic of gas-phase and heterogeneous bromine chemistry in the stratosphere. The heavier dark lines show the heterogeneous (het) chemistry.

k36 = 1.5 • 10-11e -6~176 cm 3 molecule -~ s -~ (DeMore et al., 1997),

(36)

704


Br + H C H O + HBr + CHO,

(56)

k56 = 1.7 • 10-11e -8~176 cm 3 molecule -1 s -1 (DeMore et al., 1997), These sinks for Br compete with its reaction (32)with

with O(3p). Despite these rapid sink processes for BrO, it still represents the major form of bromine in the lower stratosphere (Lary, 1996). H O B r is formed in the stratosphere by the reaction of BrO with HO2:

03:

BrO + H O 2 --+ HOBr + 0 2 , Br + 0 3 --+ BrO + 03,

(32)

k59 =

k32 = 1.7 • 1 0 - ~ e -8~176 cm 3 molecule -~ s -1

3.4 • 10-12e 54~

cm 3

(59)

molecule -~ s -~

(DeMore et al., 1997).


Measurements of the solubility of HBr in sulfuric acid at 220 K gave Henry's law constants from 8.5 • 103 M atm -1 for 72 wt% H2SO 4 to 1.5 • 107 M atm -1 for 54 wt% U2SO 4 (Williams et al., 1995; Abbatt, 1995). Application of these values to stratospheric aerosol particles typical of midlatitude conditions gives very small equilibrium concentrations of dissolved HBr; i.e., most of the HBr will remain in the gas phase. Another difference between bromine and chlorine compounds is their photochemistry. Thus, absorption cross sections of bromine compounds tend to be larger than those of the corresponding chlorine compounds, especially at longer wavelengths where the intensity of solar radiation available for photolysis is larger. As a result, conversion of bromine compounds into active forms is generally faster. For example, BrO has a peak absorption cross section of 1.5 • 10- ~7 cm 2 moleculeat 298 K, with the absorption occurring out beyond 360 nm; C10, on the other hand, has a peak absorption cross section of 6 • 10- ~8 cm 2 molecule- ~ and absorption approaches zero at 310 nm. As a result, conversion of BrO to Br + O through photolysis is important for BrO but the analogous reaction for C10 is not. Indeed, the lifetime of BrO with respect to photolysis in the stratosphere is of the order of seconds (Lary, 1996). BrO can also be converted back to atomic bromine via its reaction ( 3 1 ) w i t h C10 discussed earlier, its reaction with NO,

One issue of interest with respect to reaction (59) is whether there is a second channel that produces the reservoir species HBr, rather than the reactive HOBr, which would have a major effect on the partitioning of reactive bromine in the stratosphere (Lary, 1996). Larichev et al. (1995) place an upper limit of 1.5% on this potential path over the temperature range 233-298 K, which is consistent with indirect estimates of an upper limit of 0.01% (Mellouki et al., 1994) as well as with measurements of HBr and upper limits for the HOBr concentration in the stratosphere (D. G. Johnson et al., 1995). Cronkhite et al. (1998)observe no dependence of the overall rate constant on pressure between 12 and 25 Torr, indicating that a third possible channel producing an HO 2 9BrO adduct is not important. The best match of observed HBr concentrations at midlatitudes to modeled values is obtained with a branching ratio of ~ 1% (Chartrand and McConnell, 1998). This assumes that there is no HBr production from a minor channel in the OH + BrO reaction, analogous to the O H + C10 reaction discussed earlier. (The major path is OH + BrO --+ Br + HO2.) The atmospheric observations are also consistent with no production of HBr in reaction (59), but rather, a branching ratio for the production of HBr + 0 2 in the OH + BrO reaction in the range of 1 - 3 % (Chipperfield et al., 1997; Chartrand and McConnell, 1998). In this case, the O H + BrO reaction would be the major source of HBr in the 20- to 35-km region (Chipperfield et al., 1997). Another possible source of HOBr is the heterogeneous hydrolysis of BrONO 2. Once formed, HOBr photolyzes rapidly, with a lifetime of ~ 15 min at noon at the equinox at midlatitudes in the lower stratosphere (Lary, 1996). Above about 25 km, the reaction with O(3p) becomes important:

BrO + NO --+ Br + NO2,

H O B r + O(3p) ~ OH + BrO,

Like HC1, HBr is relatively slowly converted back to Br via reaction with OH: HBr + OH + Br + H 2 0 , k57 = 1.1 • 10-11e - ~

(57)

cm 3 molecule -1 s -1


k58 = 8.8 • 10-12e +26~

(58)

cm 3 molecule-~ s-

k6o =

1.4 • lO-l~ -430/T

cm 3

(60)

molecule -1 s -1

(DeMore et al., 1997),

(Nesbitt et al., 1995).

and, in the upper stratosphere, where there are higher ground-state oxygen atom concentrations, its reaction

Similarly, the absorption cross sections for the reservoir species BrONO 2 are much larger than those of

705

D. CONTRIBUTION OF BROMINATED ORGANICS

C1ONO 2 toward 400 nm, so that its lifetime with respect to photolysis at noon at the equinox in midlatitudes is only a few minutes (Lary, 1996). As discussed in Chapter 4.N, the photolysis products are not well established but appear to include Br + NO 3 as well as BrO + NO2, with possibly a small contribution from channels producing oxygen atoms. If the products are BrO + NO2, then the formation of BrONO 2 becomes a "do-nothing" reaction and will not contribute to ozone depletion. The same is true of the channel producing Br + NO3, since NO 3 photolyzes primarily t o O + N O 2 (see Chapter 4.G). As discussed in Section C.3, C1Ox chemistry and BrO~ chemistry are closely interrelated via the gasphase reactions of C10 with BrO: C10 + BrO --+ Br + OC10,

(31a)

--+ Br + C1OO,

(31b)

--+ BrC1 + 02.

(31c)

Indeed, these reactions play an important role in the Antarctic ozone hole and they have important implications for control strategies, particularly of the brominated compounds. For example, Danilin et al. (1996) examined the effects of C1Ox-BrO X coupling on the cumulative loss of 0 3 in the Antarctic ozone hole from August 1 until the time of maximum ozone depletion. Increased bromine increased the rate of ozone loss under the denitrified conditions assumed in the calculations by converting C10 to C1, primarily via reactions (31b) and (31c) (followed by photolysis of BrC1). Danilin et al. (1996) estimate that the efficiency of ozone destruction per bromine atom ( a ) is 33-55 times that per chlorine atom (the "bromine enhancement factor") under these conditions in the center of the Antarctic polar vortex, a ~ 60 calculated as a "global average" over all latitudes, seasons, and altitudes (WMO, 1999). However, although as much as 50% of the loss of 0 3 could be attributed to bromine chemistry at a Bry concentration of 25 ppt, a reduction in bromine did not give a proportional change in the total destruction of 0 3. Figure 12.46 shows the predicted cumulative 0 3 loss as a function of the Bry concentration; as bromine decreases, the contribution of the B r O - C 1 0 cycle decreases as well. However, the net effect on total ozone loss is quite small because near-total ozone destruction occurs even without a significant contribution from BrO. It should be noted that in the Arctic, the integrated concentrations of C10 are generally smaller than in the Antarctic; thus, the BrO + C10 cycle is relatively more important than the C10 + C10 cycle (Salawitch et al., 1993; WMO, 1995). In addition, only ~ 30-50% of the

~,

g

2 1.6

O

1.2

.=>

0.8

= E

0.4

O

0

CIO + CIO

5

I

10

I

15

I

20

25

Total Br (pptv)

F I G U R E 1 2 . 4 6 Model-calculated cumulative loss of ozone from August 1 to the day of maximum ozone depletion as a function of stratospheric bromine ([Cly] = 12.5 ppbv, [NOy] = 2 ppb, 70~ at an altitude corresponding to 50 mbar total pressure in these calculations) (adapted from Danilin et al., 1996).

0 3 in a given air mass is destroyed over the Arctic, compared to near-total ozone destruction over Antarctica. As a result, bromine does make a significant contribution to the total 0 3 destruction in the Arctic, and as a result, control of brominated organics is expected to have a greater effect on minimizing ozone destruction in this region. The BrO + BrO self-reaction occurs in a manner analogous to the C10 + C10 self-reaction: BrO + BrO - , 2Br + 0 2,

(61a)

--+ Br 2 + 02,

(61b)

M --+ B r 2 0 2 .

(61c)

The rate constant for paths a and b have been determined to be k61a = 5.3 • 10-12e -211/~ and k61b = 1.1 X 10-14e983/T c m 3 molecule-1 s - l , respectively (Harwood et al., 1998). At temperatures below 250 K, the termolecular reaction forming the dimer becomes relatively more important (Harwood et al., 1998). However, this self-reaction is not important because of the relatively small concentrations of BrO in the stratosphere. Reactions with C10 and HO 2 are much more important. Although there are fewer studies of the heterogeneous chemisty of BrONO 2 and H O B r than of the corresponding chlorine compounds, it is clear from the laboratory studies that have been done that analogous chemistry occurs, and at least as fast as for the chlorine compounds. Table 12.8 shows some of the most important reactions and typical values of the reaction probabilities. On ice, the hydrolyses of C1ONO 2 and BrONO 2 proceed at comparable rates (Tables 12.5 and 12.8). However, toward midlatitudes the particles are largely concentrated sulfuric acid-water mixtures, and on this surface the C1ONO 2 hydrolysis reaction probability

706


TABLE 12.8 SomeTypical Reaction Probabilities under Stratospheric Conditions for Some Heterogeneous Reactions Involving Bromine Reaction probability ( ~ ) Reaction BrONO2 + H20

het> H O B r + H N O 3

B r O N O 2 + HC1 H O B r + HC1 HOBr + HBr

het her her

H B r + C1ONO 2

BrC1 + H N O 3

BrC1 + H 2 0

~ Br 2 + H 2 0 her > BrC1 + H N O 3

Ice

Liquid H 2SO4-H 2 0

>0.3 b

=0.8 a

>0.3 b

>0.9 d

0.25 c

-~0.2 d

0.1 c

0.25 c

>0.3 e

a F r o m H a n s o n et al. (1996). b F r o m H a n s o n and Ravishankara (1993) and Allanic et al. (1997). c F r o m Abbatt (1994); 3' for H O B r + H B r ~ Br 2 + H 2 0 on 69 wt% H 2 S O 4 - H 2 0 at 228 K corresponds to a liquid-phase rate constant of > 5 • 104 L m o l - 1 s - 1 (Allanic et al., 1997). d F r o m H a n s o n and Ravishankara (1995); bimolecular rate constant in 70 wt% H z S O 4 is 1.4 • 105 L m o l - 1 s - 1 (Abbatt, 1995). e F r o m H a n s o n and Ravishankara (1992).

decreases. As a result, chlorine activation becomes less important at midlatitudes while bromine activation continues. Because of this, HOBr becomes an important intermediate in bromine chemistry. For example, model calculations by Randeniya et al. (1997) suggest that the increased OH produced upon photolysis of HOBr increases the loss of 0 3 in the 12- to 20-km range under conditions of large aerosol particle concentrations such as those after the Mount Pinatubo eruption. Another factor may be the strong interaction of ice and HOBr, which increases its residence time on the particles and hence the opportunity to carry out heterogeneous chemistry (Allanic et al., 1997). HOBr also serves to couple bromine and chlorine chemistry in an indirect manner. Thus, photolysis of HOBr generates increased OH concentrations, which then cause a faster recycling of HC1 back into chlorine atoms (Lary et al., 1996; Randeniya et al., 1996a,b; Tie and Brasseur, 1996). Lary et al. (1996) estimated that the lifetime of HC1 can be reduced by as much as a factor of three through this effect and suggest that the unexplained rapid rise in OH reported by Salawitch et al. (1994) at dawn may be due to the photolysis of HOBr formed overnight rather than of a nitrogen species such as HONO. Another coupling of these halogen chemistries takes place when HOBr reacts with HC1 in the condensed phase (Abbatt, 1994; Lary et al., 1996), forming BrC1, which is then released to the gas phase. Like C12 and Br2, BrC1 rapidly photolyzes so that active chlorine is again generated from what was HC1 in the particle. Reactions of BrONO 2 with HC1 and of HBr with HOC1 or C1ONO 2 give the same net result.

Thus, the effect of heterogeneous bromine chemistry is primarily to amplify the chlorine-catalyzed destruction of ozone through the more rapid conversion of the reservoir species HC1 back into active forms of chlorine (Lary et al., 1996; Tie and Brasseur, 1996). This becomes particularly important under conditions of enhanced aerosol particles, e.g., after major volcanic eruptions. In summary, there are a variety of paths by which bromine can contribute to stratospheric chemistry. Excellent reviews of the gas-phase chemistry and of the heterogeneous chemistry of bromine relevant to the stratosphere are found in Lary (1996) and Lary et al. (1996), and of the thermochemistry of bromine oxides in Chase (1996).

E. C O N T R I B U T I O N OF IODINE.CONTAINING ORGANICS In addition to chlorinated and brominated organics, iodine-containing organics are also emitted into the troposphere, primarily by biological processes in the oceans. Methyl iodide is believed to be the major species emitted, but others such as C1CH: I and CH: IBr may also be generated (e.g., see Cicerone, 1981; Klick and Abrahamsson, 1992; Moore and Tokarczyk, 1992; Schall and Heumann, 1993; Gribble, 1994; Happell and Wallace, 1996; and Carpenter et al., 1999), and ethyl iodide has also been measured recently (Yokouchi et al., 1997). There is a substantial difference in their tropospheric chemistry from that of the chlorine and bromine

G. PROBLEMS

compounds, however (e.g., see Huie and Laszlo, 1995). The carbon-halogen bond is very weak, 57 kcal mo1-1 in CH3-I compared to 70 kcal mol -~ for CH3-Br, 85 kcal mol-1 for CH 3-C1, and 108 kcal mol-~ for CH 3-F. In addition, the absorption spectra are red-shifted for the iodine compounds, so that their absorption spectra better overlap with increasing solar intensity. As a result, organoiodine compounds photolyze readily in the troposphere to generate iodine atoms (Calvert and Pitts, 1966). Combined with other fates such as reaction with OH and NO3, their tropospheric lifetimes are sufficiently short that they are not expected to reach the stratosphere in sufficient quantities to contribute to ozone destruction. Because of these rapid removal processes in the troposphere, the contribution of iodine to stratospheric photochemistry has not received much attention. However, Solomon et al. (1994)suggested that rapid transport from the lower troposphere into the upper troposphere and lower stratosphere via convective clouds could provide a mechanism for injecting such compounds into the stratosphere. While the relevant chemistry of iodine is not well known, it would be expected to interact with the C10 X cycles in much the same way as BrO, e.g., I + 0 3 ~ IO + 02

(62)

C10 + IO ~ I + OC10 I + C1OO IC1 + O 2 I +CI+

M

O2

(63a) (63b) (63c) (63e) (63f)

IO2C1

(63g)

The overall rate constant for the C10 + IO reaction has been measured to be k63 = 5.1 • 10-12e 28~ cm 3 molecule -~ s -~, with a branching ratio of 0.14 _+ 0.04 for all channels not producing I atoms at 298 K. (Turnipseed et al., 1997). This is in agreement with branching ratios for (63a) of 0.55 _+ 0.03, (63c) of 0.20 _+ 0.02, and (63d) of 0.25 _+ 0.02 reported by Bedjanian et al. (1997a). In addition, B r O - I O cross interactions would be expected; the major channel in this reaction appears to generate Br + OIO, with a branching ratio of ~1 within an uncertainty of ~ 35% (Bedjanian et al., 1997b, 1998; Laszlo et al., 1997; Gilles et al., 1997). Reaction of IO with HO2, O, and NO and photolysis will also occur (DeMore et al., 1997):

IO + O

~

+ 02,

0 2 -nt-

I,

IO + NO ~ I + NO2, I0 +hv~I

+O.

(66) (67)

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719

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C H A P T E R

13 Scientific Basis for Contro 1 of Halogenated Organics Despite basic understanding of the fundamental gas-phase chemistry discussed in Chapter 12 and the expectation based on this that CFCs would cause stratospheric ozone destruction, there was no clear and incontrovertible evidence that this was occurring until about the mid-1980s. The reason was the difficulty of reliably extracting small changes in stratospheric ozone superimposed on top of natural variations arising from a number of sources. However, with the discovery of the Antarctic "ozone hole" and ruling out natural causes for this dramatic ozone loss, it became firmly established that CFCs were indeed leading to stratospheric ozone destruction. Subsequent studies in the Arctic and at midlatitudes (discussed herein) have shown that ozone destruction has occurred globally. This area illustrates an interesting policy and regulatory dilemma that was faced in the late 1970s, because of the long delay times between release of the CFCs at the earth's surface and their impact on stratospheric ozone. Thus, although the science indicated the urgent need for controls starting with the seminal Molina and Rowland paper in 1974, there was at that time no concrete evidence for stratospheric ozone destruction from CFCs. However, had no policy decisions or regulatory actions been taken until such evidence was available and universally accepted, the magnitude and time horizon for ozone destruction would have been much more severe. In this chapter, we discuss the international agreements for phasing out halogenated organics and the resulting trends in the atmospheric concentrations of CFCs. The assessment of the relative importance of various compounds in terms of their ozone depletion potential (ODP) is summarized. Trends in stratospheric ozone and associated changes in surface UV radiation are treated, and finally, the chemistry of alternatives to CFCs that are finding increasing use as the phase-outs of CFCs take place.

A. INTERNATIONAL AGREEMENTS ON PHASEOUT OF HALOGENATED ORGANICS Because of the recognition in the mid-1970s that CFCs were a potential threat to stratospheric ozone, their use in aerosol sprays was banned in the United States in 1978. On an international level, the Vienna Convention for the Protection of the Ozone Layer agreed in March 1985 on the need to control emissions of CFCs and other chlorine-containing substances. This was followed by the Montreal Protocol in September 1987, which limited the production and import internationally of CFCs as well as halons. Specifically, the consumption of CFC-11,-12, -113, -114, and -115 was to be reduced to their 1986 levels by 1990, to 80% of these levels by 1994, and to 50% by 1999. The Halons1211,-1301, and -2402 were specified to return to the 1986 levels by 1992. For developing countries, the target for phaseout was the year 2010 (see International Legal Materials, 1987a,b, for details). However, about this time, a variety of research indicated that even with full implementation of the Montreal Protocol, the atmospheric abundance of chlorine could reach as much as 6-9 ppb between the years 2050 and 2075. This delay is due to the relatively long time between emission of these compounds into the troposphere and when they reach the stratosphere and photolyze to produce an active chlorine atom. Figure 13.1, for example, compares the estimated equivalent effective stratospheric chlorine from 1960 until the year 2100 with no controls and a 3% increase per year in CFC and methylchloroform emissions to those with the controls agreed to in the Montreal Protocol. Equivalent effective stratospheric chlorine loading depends on emissions as well as removal processes, which determine what fraction of the CFCs emitted at the earth's

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surface reach the stratosphere. In addition, the effectiveness of compounds in destroying stratospheric ozone depends on how readily they dissociate to form chlorine or bromine. As discussed in Chapter 12.D, bromine contributes significantly to ozone destruction as well and, indeed, is more efficient than chlorine under the current atmospheric conditions by about a factor of 60 (World Meteorological Organization, 1999). The contribution of bromine can be translated into equivalent chlorine by taking this increased efficiency into account. The net effect of chlorine and bromine is then expressed as "equivalent effective stratospheric chlorine" and it is this quantity that is shown in Fig. 13.1. While the growth in stratospheric chlorine should clearly be slowed by the Montreal Protocol agreements, it was still substantial and expected to lead to quite large losses of ozone. This recognition, bolstered by the dramatic appearance of the Antarctic ozone hole, led to further major amendments to the Montreal Protocol. The Montreal Protocol was first modified in June 1990 in what is known as the London Amendments. These amendments speeded up the phaseout of CFCs and halons to 50% by 1995, 85% by 1997, and 100% by 2000, or 2010 for the developing countries (e.g., see Chatterjee, 1995; and Parson and Greene, 1995). Carbon tetrachloride was to be fully phased out by 2000 and methylchloroform (CHsCC1 s) by 2005. Included in the controls were a variety of additional CFCs, including CFC-13,-111,-112,-211,-212,-213,-214,-215,-216, and -217, not covered in the original protocol. A non-

binding portion of the agreement included phasing out the interim replacement compounds for CFCs, the HCFCs (hydrochlorofluorocarbons), with a target date of 2020, but not later than 2040. (See International Legal Materials, 1991, for details.) A further amendment was subsequently signed in November 1992 in Copenhagen. This speeded up the complete phaseout of CFCs, CC14, and CHsCCls to 1996 and the halons from 2000 to 1994. The phaseout target for developing nations continued to be the year 2010. The HCFCs were to be reduced by 35% by 2004, 65% by 2010, 90% by 2015, and 99.5% by 2020, with full phaseout by 2030. Starting in 1995, the production of methyl bromide in developed countries was held at 1991 levels. In mid-July 1994, most countries had agreed to the restrictions; approximately 40 countries, with ~6% of the world's population, had not agreed to the controls (Holmes and Ellis, 1996). Figure 13.1 also shows the estimated equivalent effective stratospheric chlorine content as a function of year under the Copenhagen Amendments. Chlorine is expected to peak around the year 2000 and then decrease. As discussed by Holmes and Ellis (1996), noncompliance and allowed exemptions could lead to a much slower decline than projected in Fig. 13.1. In November 1995, it was agreed that CHsBr usage would be eliminated by 2010 in developed countries; in developing countries, its use is frozen in the year 2002 at the average 1995-1998 levels. Furthermore, HCFCs would be phased out in developed countries by 2020, with some production allowed until 2030 for application in existing refrigeration and air conditioning units. In developing countries, the use of HCFCs will be frozen at 2015 levels in the year 2016, with complete elimination by 2040 (Zurer, 1995). Figure 13.2 summarizes these international agreements for developed countries through ~ 1995 (Parson and Greene, 1995). In September 1997, it was agreed that the use of methyl bromide by developed countries would end earlier, in 2005, and in developing countries by 2015 (Spurgeon, 1997a). In addition, steps to try to reduce smuggling of CFCs, which had become a problem (Spurgeon, 1997b), were taken. As discussed in more detail in this chapter, detecting trends in stratospheric ozone and deconvoluting the causes are complex, particularly outside the polar regions. However, it is estimated that for the Antarctic, where the most dramatic loss of ozone has been observed, recovery may be experimentally observable by the year 2008 if the Montreal Protocol and associated amendments are followed (e.g., Hofmann et al., 1994). When first developed as refrigerants in 1928 by the DuPont chemists Midgley and Henne (1930), CFCs appeared to have ideal properties for this application

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2004 t-

2000 -

"C= N

Cut 50%

"C= N

E N

Cut 70%

2000 2000 2000 2000

199s

.

1987

~

1994

1996 o Z v

1992

Montreal protocol (September, 1987)

E

8

8

g9

-g e'-

,-

by 1995 1995 1995 Freeze

1990 -

8

o ~0

I Cut Cut Cut 50% 85% 30%

Cut _~0%

reez 199(

.

--,!

Cut 85~176 1997

1995 -

"E"

..~,o

"~ -rE N

75% 1994

Freeze

)9( 1996 1996 Freeze Cut 1995

,

Freeze

1995

50%

Freeze

1992 1992

London a m e n d m e n t (June, 1990)

85%

Copenhagen a m e n d m e n t (November, 1992)

Vienna a m e n d m e n t (November, 1995)

FIGURE 13.2 Summary of Montreal Protocol and subsequent amendments (adapted from Parson and Greene, 1995).

and, subsequently, for use as aerosol propellants and blowing agents. They were nontoxic, nonflammable, and highly stable. Of course, as described in the previous chapter, it is now known that it is these very properties that render them virtually nonreactive in the troposphere and hence give them sufficiently long lifetimes in the lower atmosphere to reach the stratosphere. Table 13.1 shows the most common CFC substitutes

currently in use. The physical and chemical properties that are needed for substitute compounds are described in an article by Manzer (1990) and the book edited by Miziolek and Tsang (1995). Overall, the approach in the short term has been to make these compounds more reactive in the troposphere by adding one or more abstractable hydrogens. The term " H C F C " is used for hydrochlorofluorocarbons, i.e., compounds containing hydrogen, chlorine, and fluorine, and " H F C "

730


TABLE 13.1 SomeCFC Alternatesa Name b

Compound

Use

HFC-134a

CH2FCF 3

HFC-125 HFC-152a HCFC-141b H CFC- 142b HCFC-22

CHF2CF 3 CH3CHF 2 CH3CC12F CH 3CF2 C1 CHC1F2

HCFC-123

CHC12CF 3

Automobile air conditioners, commercial chillers, residential refrigerators Industrial and commercial refrigeration Industrial and commercial refrigeration, aerosol sprays Polyurethane foams, solvent and cleaning applications Polyurethane foams Industrial refrigerant, manufacturing of polymers, polystyrene foams, aerosol sprays Commercial chillers

a From Kirschner (1994); for a discussion of halon replacements, see Freemantle (1995). b HFC = hydrofluorocarbon; HCFC = hydrochlorofluorocarbon.

for hydrofluorocarbons, i.e., compounds containing hydrogen and fluorine but no chlorine. Because the HCFCs contain chlorine, they will continue to contribute to ozone depletion, and their ultimate phaseout is included in the Montreal Protocol and its amendments. However, the fraction that reaches the stratosphere is much smaller than for the CFCs because of their partial removal in the troposphere by reaction with OH. For a discussion of the history of controls on ozonedepleting chemicals and factors that must be considered in finding transitional substitutes and long-term replacements, see Ko et al. (1994a), Parson and Greene (1995), the World Meteorological Organization Report (1995 and the 1999 update), and the volume edited by Miziolek and Tsang (1995).

B. OZONE DEPLETION POTENTIALS (ODP) The contribution of individual compounds to ozone depletion is characterized by the ozone depletion potential (ODP). The ODP of a compound as normally defined is the ratio of the global loss of ozone (i.e., integrated over latitude, altitude, and time) from that compound at steady state per unit mass emitted relative to the loss of ozone due to emission of unit mass of a reference compound, usually taken as CFC-11 (CFC13) (Wuebbles, 1983; Fisher et al., 1990; Solomon et al., 1992). The ODP thus provides a relative measure of the overall impact of a compound on ozone destruction over the long term. Values for ODPs have been derived using a variety of models (e.g., see World Meteorological Organization, 1995, 1999). However, the time scale for contribution to ozone depletion is not the same for all compounds (Fisher et al., 1990). For example, CFC-11 does not react in the

troposphere but photolyzes in the stratosphere, with an estimated stratospheric lifetime of ~ 5 0 - 5 5 years (World Meteorological Organization, 1995; Solomon and Albritton, 1992). The transition HCFC-123 (CF3CHC12) reacts with OH and hence has a finite lifetime in the troposphere (about 2 years; Nimitz and Skaggs, 1992). In the stratosphere, the rate of chlorine release from HCFC-123 is faster than for CFC-11, with an estimated stratospheric lifetime of ~47 years (Solomon and Albritton, 1992). Hence its contribution to ozone depletion is high initially, about 50% that of CFC-11, but falls off with time. In contrast, CFC-113 (CFC12CF2C1) reacts more slowly in the stratosphere (estimated lifetime of ~110 years; Solomon and Albritton, 1992), and hence it continues to give chlorine after the contribution from the reference compound, CFC-11, has started to decay. As a result, its effective ODP actually increases with time. Clearly, these different time scales for ozone depletion are important for designing both shorter term transitional alternatives for CFCs and longer term permanent substitutes. A semiempirical approach has been developed by Solomon et al. (1992) and applied to the estimation of ODPs on both short and long time scales. The ODP of compound X is calculated from Eq. (A) (Solomon and Albritton, 1992):

ODP(t) =

Fx

McFc-11

FCFC_ll

Mx

HX

~c~ 3

f te - (t-t~)/~x dt ts

:e-( t- ts)/,cvc_11 dt

fts

(A) The ratio Fx/Fcvc_11 is the fraction of the compound X that has been dissociated in the stratosphere compared to that of CFC-11, determined from atmospheric measurements. This term is most important at short times, i.e., right after release of the compound, in terms

B. OZONEDEPLETIONPOTENTIALS(ODP) 100

of determining the ODP. M denotes the molecular weight of each species, to convert the O D P to mass emissions, n x is the n u m b e r of chlorine, bromine, or iodine atoms in X (with the 3 in the denominator representing the three chlorine atoms in CFC-11), and a is an e n h a n c e m e n t factor that reflects the relative efficiency of ozone destruction by bromine (or iodine, should that prove ultimately to be important). Values of a for bromine are ~ 40-100 (World Meteorological Organization, 1999; Ko et al., 1998), and are larger for iodine, up to ~ 92,000 (Solomon et al., 1994a,b). (Note, however, that as discussed in Chapter 12.E, current evidence suggests that there is insufficient iodine in the stratosphere to contribute significantly to ozone depletion.) The time t s is the time for a molecule to be transported from the surface to the region of the stratosphere being examined, and t is the total time; i.e., the ( t - t s) term represents the total time in the stratosphere available for conversion into active forms of chlorine. The total atmospheric lifetimes of X and CFC-11 are denoted by r x and rcvc_11 , respectively. The total atmospheric lifetime of a compound is the time required for a pulse emitted into the atmosphere to decay to 1 / e of its initial value (see Chapters 5.A.lc and 12.C). It can be calculated from

10

,

ELoss rate

13_

CI

_

- ~ CH3Br . . . . . _~'_:. . . .

---'-'-'-= --- ---- ~

-

"'""

C F C - 113 9

.

HCFC-141b

~ ~

-

"'--... . . . . .

.

.

H C F C - 123

.

.01

.001

1

10

100

FIGURE 13.3 Time dependence of ODPs (adapted from World Meteorological Organization, 1995).

The integrals in Eq. (A) are of the form e -~t, where k x = 1 / r x and t is time in that region of the stratosphere. The ratio of these integrals in the last term in Eq. (A) can be thought of as the ratio of the timedependent concentrations of X and CFC-11 as they decay. As t--+ % the O D P calculated from Eq. (A) approaches the overall, time-integrated O D P calculated from models. This last term is the part of Eq. (A) that determines the time dependence of the ODP, and it generally dominates the O D P at longer times. The semiempirical nature of this calculation arises from basing the ratio F x / F c f c _ ~ on measurements of the compound of interest and CFC-11, combined with calculated atmospheric lifetimes. Figure 13.3 and Table 13.2 demonstrate the effects of different time scales for ozone depletion. The data

(B)

Time Dependence of Ozone Depletion Potentials for Some Compounds Using the Semiempirical Approach a Time Horizon (years)

Compound

CFC-113 CH 3CC13

HCFC-142b HCFC-22 HCFC-14 lb HCFC-123

CH3 Brb H-1121 H-1301 H-2402

1000

Time since reaching stratosphere (years)

where EBurdenat m represents the total amount of X in the atmosphere and s rate) the globally integrated loss rate due to all processes, including reactions, deposition, etc. r x can also be calculated from the global source strengths and the atmospheric concentrations.

TABLE 13.2

CF3Br

: -

O

~Burdenat m rx =

731

5

15

25

40

100

500

0.55 1.03 0.17 0.19 0.54 0.51 15.3 11.3 10.3 12.8

0.58 0.57 0.15 0.15 0.38 0.11 3.1 9.7 10.5 11.6

0.60 0.38 0.13 0.13 0.30 0.07 1.8 8.5 10.6 10.6

0.64 0.26 0.12 0.10 0.22 0.04 1.2 7.1 10.8 9.4

0.78 0.15 0.08 0.07 0.13 0.03 0.69 4.9 11.5 7.0

1.09 0.12 0.07 0.05 0.11 0.02 0.57b 4.1 12.5 5.9

a Using Eq. (A); from Solomon and Albritton (1992), for conditions of the Arctic lower stratosphere. b Ko et al. (1998) estimated an ODP of 0.39 based on kinetic and photochemical data from DeMore et al. (1997).

732

13. S C I E N T I F I C BASIS F O R C O N T R O L OF H A L O G E N A T E D O R G A N I C S

in Table 13.2 were derived using measurements of

et al. (1998) have summarized changes in the estimated

F x / F c F c _ 1 1 characteristic of polar regions. However, as

efficiency of bromine in stratospheric ozone destruction from c ~ = 3 6 in 1990 to c ~ = 5 8 in 1997 due to improved understanding of the kinetics and photochemistry of bromine compounds that are important in the stratosphere. They also show changes in the estimated atmospheric lifetimes of CH 3Br over this period of time, from 2.1 to 0.7 years. The net result of these changes is a variation in the estimated ODP from 0.73 in 1990 up to 0.99 using 1992 data and then to their estimate of 0.39 based on 1997 kinetic and photochemical data. Table 13.3 summarizes some of the values of globally and time-averaged ODPs (corresponding to the long times in Fig. 13.2) obtained using either models or the semiempirical method based on Eq. (A) at long times (Solomon et al., 1992). Also shown are the total atmospheric lifetimes for these gases. Note that the HFCs in Table 13.3 have such small ODPs that only upper limits can be placed on them. That is, their ODPs are essentially zero for all practical purposes. The reason is the obvious one, that they do not contain chlorine and, as discussed in Chapter 12.C, fluorine does not participate in ozone destruction cycles (Molina and Rowland, 1974; Rowland and Molina, 1975; Stolarkski and Rundel, 1975; Cicerone, 1979; Ravishankara et al., 1993; Sehested et al., 1994; Li et al., 1995). In summary, a combination of the time-dependent

discussed by Solomon and Albritton (1992), these are not expected to differ by large amounts from the globally averaged values of the ODPs. As expected from the earlier discussion, the ODP for HCFC-123 is initially high and then falls off rapidly, approaching a long-time, steady-state value of ~0.02. This rapid decline reflects its short atmospheric lifetime of ~1.7 years (Solomon and Albritton, 1992) due largely to its removal by reaction with OH in the troposphere (e.g., see Nimitz and Skaggs (1992) for a discussion of tropospheric lifetimes and ODPs). However, the ODP of CFC-113, with no known chemistry in the troposphere and a longer stratospheric lifetime than CFC-11, increases with time since it continues to generate atomic chlorine while CFC-11 decays. These data also demonstrate the impact of bromine chemistry on the stratosphere (see Chapter 12.D). The initial ODP for methyl bromide is 15, due primarily to the large c~ factor associated with bromine chemistry. However, since it is removed by reaction with OH in the troposphere as well as by other processes such as hydrolysis in the oceans and uptake by soils and foliage (see Chapter 12.D), it has a short atmospheric lifetime of ~1.3 years and hence the ODP decreases rapidly with time, toward a long-term steady-state value. The ODP for CH3Br also illustrates the sensitivity of these estimates to the input data. For example, Ko

TABLE 13.3 Atmospheric Lifetimes and Steady-State Ozone Depletion Potentials (ODP) Predicted Using Either a Two-Dimensional Model or a Semiempirical Method a'b Potential

Atmospheric

trace gas

lifetime (years)

Model

Semiempirical a

50 102 85 5.4 13.3 1.4 9.4 19.5 14 36 1.3 (0.7)d

1.0 0.82 0.90 0.12 0.04 0.014 0.10 0.05 CF3C(O)F + HO 2 .

(6)

The absolute kinetics of the decomposition, reaction (5), compared to reaction (6)with 0 2 have been somewhat uncertain (e.g., Wallington et al., 1992; Tuazon and Atkinson, 1993a; Rattigan et al., 1994). Hasson et al. (1998) studied the oxidation of HFC-134a by FTIR from 298 to 357 K at various concentrations of 0 2 and report that the ratio of rate constants is given by k 6 / k 5 = (3.8 _+ 1.6) • 10 -24 e(2400+500)/~ cm 3 molecule -1. Because of the strong temperature dependence, the relative importance of reaction (5) decreases significantly as one moves from the earth's surface to the much colder tropopause. Thus, about 20-25% of the loss of CF 3CHFO is estimated to proceed via (6) at the surface, increasing to ~ 8 0 % at the tropopause (Tuazon and Atkinson, 1993a; Rattigan et al., 1994). As discussed shortly, CF3C(O)F generated in reaction (6) ultimately hydrolyzes to trifluoroacetic acid (CF3COOH) in the atmosphere. The CF 3 radical adds O 2 to form CF302, which then forms the alkoxy radical CF30: M CF 3 + O 2 ---> CF302,

(7)

C F 3 0 2 + NO + CF30 + N O 2.

(8)

The values of k 0 and k~ for reaction (7) are 3.0 • 10 -29 cm 6 molecule 2 s -1 and 4.0 • 10 -12 cm 3 molecule -1 s-1, respectively, at 300 K (DeMore et al., 1997), giving an effective bimolecular rate constant of 3.7 • 10 -12 cm 3 molecule -1 s -1 at 1 atm pressure at 300 K. The lifetime of CF 3 with respect to reaction with O 2 is thus only ~50 ns under these conditions, and other reactions such as that with NO 2 (e.g., see Sehested et al., 1996; and Vakhtin, 1997)cannot compete. Reaction (8), reaction of CF30 2 with NO, is also sufficiently fast that this is a major fate of CF30 2 in the troposphere. Thus the rate constant at 298 K for reaction (8) is 1.6 • 10 -11 cm 3 molecule -1 s -1 (e.g., see Sehested and Nielsen, 1993; Bevilacqua et al., 1993; Turnipseed et al., 1994; and Bhatnagar and Carr, 1994), which gives a lifetime for CF30 2 of 4 min at 10 ppt NO. At small NO concentrations, reaction with HO 2 (or RO 2) can also occur: CF3OO

+ H O 2 --+ CF3OOH + O 2.

(9)

The rate constant for this reaction is k 9 = (4.0 -+- 2.0) x 10 -12 cm 3 molecule -1 s -1 (Sehested et al., 1997).

747

At an HO 2 concentration of i • 109 cm -3, the lifetime of CF30 2 with respect to reaction (9)with HO 2 is also 4 min. Thus, under these assumed concentrations of NO and HO2, the reactions are equal in magnitude. CF30 2 (and similar halogenated peroxy radicals) can also react with NO 2 to form trifluoromethyl peroxynitrate, CF3OONO2, but the relatively fast thermal decomposition of peroxynitrates makes this only a temporary reservoir under the conditions at the earth's surface, as is the case for unsubstituted alkylperoxy radicals discussed in Chapter 6 (Chen et al., 1993a; Niki et al., 1994; MOgelberg et al., 1994b). For example, the lifetime of CF3OONO 2 is estimated to be only ~ 1 min at room temperature and 1 atm total pressure; however, under typical conditions at the tropopause, the lifetime with respect to thermal decomposition is much longer, about 1 year (Mayer-Figge et al., 1996). Similar considerations apply to halogenated PAN analogs. For example, the lifetime of CF3C(O)OONO 2 is calculated to be only ~ 3 h at room temperature but to increase by about seven orders of magnitude under the conditions at 12 km! (Zabel et al., 1994; Wallington et al., 1994e). It should be noted that while the formation of hydroperoxides as shown in reaction (9) is common for simple hydrocarbon alkylperoxy radicals (see Chapter 6), recent work on chlorine-containing RO 2 radicals indicates that this is not always the case. Catoire et al. (1996) studied the reactions of C H C l 2 0 2 and CC130 2 with HO 2. In the latter reaction, the yield of phosgene, COC12, was ~100%. The reaction was interpreted to proceed via formation of the tetraoxide intermediate CC13OOOOH, which decomposes via a cyclic transition state to COCl 2 + HOC1 + 0 2. The tetraoxide intermediate from the C H C l 2 0 2 reaction was proposed to have two analogous decomposition paths available, one giving COC12 -Jr- H 2 0 -Jr- O2 (29% at 296 K and 700 Torr total pressure) and the other HC(O)C1 + HOC1 + 0 2 (71% at 296 K and 700 Torr). In neither case was the formation of a stable hydroperoxide observed, in contrast to the reaction of C H 3 0 2 with HO2, for example (see Chapter 6). In the case of C H 2 F O 2 radicals, both pathways occur, with 29% of the reaction with HO 2 giving the C H 2 F O O H and 71% HC(O)F as products (Wallington et al., 1994d). Other reactions of CF30 2 such as that with 0 3 are slow and cannot compete (e.g., see Nielsen and Sehested, 1993; Turnipseed et al., 1994; and Meller and Moortgat, 1995). There are no abstractable hydrogen atoms on CF30, so that reaction with 0 2 is negligible (k < 4 • 10 -17 cm 3 molecule- 1 s- 1 at 373 K; Turnipseed et al., 1994; Niki et al., 1994). In addition, the C - F bond is too strong for decomposition to give a fluorine atom and CF20 to take place. However, CF30 reacts with NO,

748


possibly with H 2 0 and with organics (e.g., Niki et al., 1994): CF30 + NO ~ CF20 + FNO,

(10)

CF30 + H 2 0 ~ CF3OH + OH,

(11)

CF30 + RH ~ CF3OH + R.

(12)

The reaction (10)with NO is fast, kl0 = (5-6) • 10 -11 cm 3 molecule-1 s-1 at 298 K (e.g., see Bevilacqua et al., 1993; Sehested and Nielsen, 1993; Bhatnagar and Carr, 1994; Jensen et al., 1994; Turnipseed et al., 1994; and Dibble et al., 1995), and has been shown to give CF20 and FNO as products (e.g., Chen et al., 1992a; Niki et al., 1994; Dibble et al., 1995). FNO reacts only slowly with atmospheric reactive species such as 0 3, O(3p), HO2, and HC1 so that photolysis appears to be its major atmospheric fate (Wallington et al., 1995). CF30 has also been observed to react with NO 2 to form CF3ONO 2 (Niki et al., 1994), but the kinetics were not reported. The importance of the reaction of CF 30 with H 2 0 is less certain; an upper limit to the rate constant kll at 298 K is 1 • 10 -16 cm 3 molecule -1 s -1 (Wallington et al., 1993b; Turnipseed et al., 1995). However, because of the relatively large amounts of water vapor in the troposphere, application of this upper limit gives a lifetime for CF30 with respect to reaction with water vapor of > 30 ms at 50% RH. This can be compared to a lifetime with respect to reaction (10) with 100 ppt NO of 7s. CF30 abstracts a hydrogen atom from alkanes in much the same way as OH, forming CF3OH (Chen et al., 1992b; Saathoff and Zellner, 1993; Bevilacqua et al., 1993; Sehested and Wallington, 1993; Kelly et al., 1993; Niki et al., 1994; Barone et al., 1994; Jensen et al., 1994; Bednarek et al., 1995; Wallington and Ball, 1995). It also reacts rapidly with alkenes and aromatics, in a manner suggesting addition as the initial reaction (e.g., see Chen et al., 1993b; Niki et al., 1994; Kelly et al., 1994; and Wallington and Ball, 1995). Table 13.6, for example, gives room temperature rate constants for the reaction of CF30 with some organics. Even the reaction with CH 4 is sufficiently fast that the lifetime of CF30 with respect to 1.7 ppm CU 4 is only 1 s. Hence reaction with organics is expected to be a major fate of CF30 in the troposphere. In the case of reaction with alkanes, the final product, CF3OH, is the same as the reaction with H20. CF30 has also been observed to react with CO with kinetics that are consistent with two channels, a pressure-independent channel and a pressure-dependent one. The existence of the pressure-dependent channel suggests CF30 adds to CO to form an adduct, analogous to the OH + CO reaction; this reaction may play

TABLE 13.6 Rate Constants at Room Temperature for the Reaction of CF 3 0 with Some Organics k(298 K)

Organic CH 4 C 2H 6

C3H s i-C4H 8 C2H 4 i-C4H10 C6H 6 CH3OH

(cm 3 molecule- 1 2.2 X 10 -14 1.4 ;K 10-12 4.7 x 10-12 6.1 • 10-12 3.0 • 10-11 7.2 ;K 10-12 3.6 • 10-11 2.5 X 1 0 1 2

Reference

s - 1)

Wallington Wallington Kelly et al., Wallington Kelly et al., Kelly et al., Kelly et al., Wallington

and Ball, 1995 and Ball, 1995 1993 and Ball, 1995 1993 1993 1993, 1994 and Ball, 1995

a minor role in the tropospheric removal of CF30 as well (Turnipseed et al., 1995; Wallington and Ball, 1995; Meller and Moortgat, 1997). While OH removes a significant amount of the CFC alternatives in the troposphere, some fraction does reach the stratosphere. One possible chain that has been considered in the stratosphere involves reactions of CF30 and CF30 2 with 0 3 (e.g., see Ko et al., 1994b). However, the rate constants for these reactions were measured subsequently and shown to be too slow relative to other removal processes such as CF30 + NO to be significant (Nielsen and Sehested, 1993; Maricq and Szente, 1993; Wallington et al., 1993a, 1994c; Maricq et al., 1994b; Ravishankara et al., 1994; Fockenberg et al., 1994; Meller and Moortgat, 1995; M6rs et al., 1995). The atmospheric fate of CF3OH appears to be primarily uptake into clouds in the troposphere and possibly into sulfate aerosols in the stratosphere. Lovejoy et al. (1995), for example, measured uptake coefficients for CF3OH on sulfuric acid-water mixtures from ~10 -3 to ~10 -1, the larger values occurring at the highest water concentrations. From these data, they suggest an uptake coefficient of about 0.1 for clouds, which gives a lifetime with respect to such uptake of the order of seconds under typical cloud conditions. As a result, uptake into tropospheric clouds would be controlled by the frequency of cloud formation. On the other hand, in the stratosphere, where the particle concentration is relatively low and the acid concentration high, the lifetime could be of the order of several years (Lovejoy et al., 1995). Other fates such as decomposition, photolysis, or reaction with species such as OH and C1 are believed to be negligible (Wallington and Schneider, 1994; Schneider et al., 1995; Huey et al., 1995). Figure 13.21 summarizes the chemistry of HFC-134a. b. Chemistry of H C F C - 1 2 5 (CHF2CF 3) With sufficient NO present, HCFC-125 (CHF2CF 3) reacts in an analogous manner to generate an alkoxy

749

D. TROPOSPHERIC CHEMISTRY OF ALTERNATE CFCs

02

CF3CH2F + OH

~--

H I C F 3 C - O - O + H20 I F

( C F 3 C H F O 0 - NO)*

1

CF302 _.4-_~,

02

NO2 CF30

F-C-F II

CF30H + OH

H I CF3-C-O

CF 3 + H - C - F + NO 2 II 0 CF 3 + H - C - F -~ II 0

CF30H + R

o

+

NO 2

Decomposition

02

CF 3 - C-F + HO 2 II 0

+ FNO

F I G U R E 13.21

Summary of major reaction paths in the tropospheric chemistry of HFC-134a.

radical:

M C F 3C F 2 + O2 ---+ CF3 CF2 O O ,

(14)

However, recall that the C-C1 bond is much weaker than a C - F bond. As a result, one additional potential route of decomposition in this case is elimination of a chlorine atom:

C F 3 C F 2 O O + NO --+ C F 3 C F 2 0 + N O 2.

(~5)

CF3CC120 --+ C1 + CF3C(O)C1.

CF 3CHF 2 + OH --+ H 20

-H

CF 3CF2,

(13)

With no abstractable hydrogen atoms on the alkoxy radical, reaction with 0 2 cannot occur and it decomposes by scission of the C - C bond (Edney and Driscoll, 1992; Tuazon and Atkinson, 1993b; Sehested et al., 1993): CF3CF20 + COF 2 + CF 3.

(16)

The CF 3 radical then reacts as described earlier, forming COF2, FNO, and CF3OH. Figure 13.22 summarizes the chemistry of HCFC-125.

c. Chemistry of HCFC-123 (CHC12CF3) The initial OH chemistry of HCFC-123 (CHC12CF 3) produces a fully halogenated alkoxy radical where further hydrogen abstraction by 0 2 is not an option: CF3CHC12 + OH -~ H 2 0 + CF3CC12,

(17)

M CF3CC12 + 0 2 ---+ C F 3 C C 1 2 O O ,

(18)

C F 3 C C I 2 O O + N O --+ C F 3 C C I 2 0 + N O 2.

(19)

(20)

In fact, this predominates over the competing C - C bond scission, CF 3CC12 O ---+ COC12 -~- CF3,

(21)

with yields of CF3C(O)C1 approaching 100% (Edney et al., 1991; Tuazon and Atkinson, 1993b; Hayman et al., 1994). Similar chlorine atom elimination has been observed for other chlorine-containing alkoxy radicals (e.g., see Wu and Carr, 1992; and Bhatnagar and Carr, 1995). The CF3C(O)C1 hydrolyzes in clouds and fogs to form CF3COOH, trifluoroacetic acid. It should be noted that, although it is not important in this particular case, there is an additional path available for some alkoxy radicals formed in the oxidation of certain halogenated organics, that is, the intramolecular elimination of HC1 from c~-monochloroalkoxy radicals. For example, the alkoxy radical CH2C1CHC10 both reacts with 0 2 and eliminates HC1,

750


02

CF3CHF2 + OH

CF3CF2OO + H20

,5:o

CH3CFCI2 + OH

02

NOe I

H-C-C-CI + HO2 < II I OF

CF3 + F-C-F

FIGURE 13.24 HCFC-141b.

"

Oa

NO2

y F-C-F

II o

CF30

~~ ~

H20

O

O

CF3OO

CF3OH + R

FIGURE 13.22 HCFC-125.

Summary of major tropospheric reactions of

with the latter favored at 1 atm in air (Wallington et al., 1996b). Figure 13.23 summarizes the chemistry of HCFC-123. d. Chemistry of NCFC-141b (CN3CFC12 )

Hydrogen atom abstraction from HCFC-141b (CH 3CFC12) produces the alkoxy radical OCH 2CFC12: CH3CFC12 + OH --+ H 2 0 + CH2CFC12, M

CH 2CFC12 + O 2 --+ OOCH 2CFC12,

(22) (23)

OOCH2CFC12 + NO --+ NO 2 + OCH2CFC12. (24) In this case, reaction of the alkoxy radical with 02, OCH2CFC12 + 0 2 --+ HO 2 + HC(O)CFC12,

CF3CHCI2 + OH

02

OCH2CFCI2

Summary of major tropospheric reactions of

OCH2CFC12 --+ HCHO + CFC12.

OH

+ FNO

02

dominates over decomposition by C - C bond scission (Edney et al., 1991; Tuazon and Atkinson, 1993b, 1994):

~',,3~

CF3OH +

OOCH2CFCI2

NOe

CI

CF3CF20

"~

(25)

CFaCCI2OO + H20

NOe CF3CCI20

CFa - C - e l + CI II O FIGURE 13.23 Summary of major tropospheric reactions of HCFC-123.

(26)

Further reaction of CFC12 with 0 2 and subsequently with NO generates C1C(O)F. Figure 13.24 summarizes the chemistry of HCFC141b. In short, the tropospheric chemistry of the CFC replacement compounds containing abstractable hydrogen atoms is very similar to the VOC chemistry discussed in Chapter 6. The major differences are found in the relative importance of the various potential fates of the alkoxy radicals formed in these reactions, which depends on the structure of the parent compound. Table 13.7 summarizes the major and minor products formed by the tropospheric oxidation of the four particular typical CFC replacement compounds treated here. For the chemistry of other CFC replacements, see papers by Chen et al. (1997) and a series of papers by Wallington, Nielsen, and co-workers (e.g., Wallington et al., 1994a; Wallington and Nielsen, 1995; MOgelberg et al., 1995a, 1995b, 1995c, 1996, 1997; Giessing et al., 1996). The atmospheric chemistry of a series of hydrofluoroethers (HFE), such as C4FgOCH 3 ( H F E - 7 1 0 0 ) , C 4 F 9 0 4 C 2 H 5 ( H F E - 7 2 0 0 ) , and CP 3cH2OCH2CF3, used as CFC alternates is discussed by Wallington et al. (1997, 1998), Christensen et al. (1998), and Cavalli et al. (1998)

e. Tropospheric Fates of Halogenated Products of H C F C Oxidation As seen in Table 13.7, oxidation of HCFCs by OH generates a variety of halogenated aldehydes and ketones as well as phosgene (COC12) , its fluorine analog (COF 2), C1C(O)F and HC(O)F, and the alcohol CF 3OH. The ultimate atmospheric fate of these products depends on their structures, of course, which determines their absorption cross sections as well as reactivity with OH, and their solubility in aqueous solutions such as clouds, rainwater, and the oceans.

751

D. TROPOSPHERIC CHEMISTRY OF ALTERNATE CFCs

TABLE 13.7 Typical Organic Products of the Tropospheric Oxidation of Some of the CFC Replacement Compounds Major fates of alkoxy radicals

Replacement

Compound

Major products

HFC-134a

CH2FCF 3

Decomposition and reaction with

HFC-125

CHF2CF3

Decomposition by C-C bond scission

COF 2, CF 3OH

HCFC-123

CHC12CF 3

Cleavage of C-C1 bond, small fraction by C-C bond scission

CF3C(O)C1

COF 2, CF3OH, COC12

HCFC-141b

CH3CC12F

Reaction with O2 primarily

HC(O)CFC12

C1C(O)F

0 2

The fate of CF3OH described earlier is believed to be primarily uptake into clouds. The same is true of COC12, COF2, HC(O)F, and CF3C(O)F. For all of these species, photolysis at the wavelengths found in the troposphere is negligible, as is reaction with OH (Nalle et al., 1992; Rattigan et al., 1993; Wallington et al., 1994a; Zachariah et al., 1995; World Meteorological Organization, 1995). Table 13.8 summarizes estimates of the lifetimes of these halogenated product species with respect to uptake by the oceans, clouds, and rainwater. Uptake into clouds followed by hydrolysis is the major removal mechanism from the atmosphere for these compounds. Hydrolysis forms HC1 and HF in the case of COC12, COF 2, and HC(O)F as well as formic acid in the latter case. CF3C(O)F hydrolyzes to HF and trifluoroacetic acid, CF3C(O)OH. The effects on human health of trifluoroacetic acid (TFA) from the oxidation of HFC-134a, halothane (1,1,1-trifluoro-2-bromo-2-chloroethane, used as an anesthetic), and some of the other CFC replacements, such as HCFC-123 and HCFC-124, have been of some

TABLE 13.8 Estimated Lifetimes of Halogenated Carbonyl Compounds in the Aqueous Phasea Compound HC(O)F COF2b C1C(O)F CF3C(O)F c CF3C(O)C1a

Ocean

Clouds

Rainout

3 months 4 months 2 years 2 years 3 years

4 days 4 days 6 days 6 days 6 days

180 years 72 years 265 years 675 years 900 years

a From Kanakidou et al. (1995). b Tropospheric lifetime estimated to be -0.5-3 days by De Bruyn et al. (1992, 1995) based on measurements of uptake by water surfaces; see also George et al. (1994a,b). c Tropospheric lifetime estimated to be -~0.5-3 days by De Bruyn et al. (1992, 1995) based on measurements of uptake by water surfaces; see also George et al. (1994a,b). d Tropospheric lifetime estimated to be -~2-10 days by De Bruyn et al. (1992, 1995) based on measurements of uptake by water surfaces; see also George et al. (1994a,b).

Minor products

HC(O)F, COF2, CF3OH , CF3C(O)F

concern. The toxicology of this compound is reviewed by Ball and Wallington (1993) and its effects on plants are discussed by Tromp et al. (1995). The reaction of TFA in the gas phase with OH is relatively slow (k ~ 1.7 • 10 -13 cm 3 molecule -1 s -1 at 296 K ) a n d is estimated to account for only ~ 10-20% of the loss of TFA, with the remainder being removed by rainout (Mc~gelberg et al., 1994a). Calculated maximum concentrations of TFA in rainwater in the future have been suggested to be in the range of 1-80 nmol L -1 (Ball and Wallington, 1993; Tromp et al., 1995), but a 3-D modeling study suggests that global annually averaged rainwater concentrations would be ~ 120 ng L-1, and in northern midlatitudes, monthly averaged concentrations could be as large as 450 ng L -1 in the summer (Kotamarthi et al., 1998). These higher estimated values are consistent with measurements of 40-1200 ng L-1 of TFA in rainwater in California and Nevada (Wujcik et al., 1997). Frank and co-workers (1996) report that in 1995, concentrations of TFA in Europe in rainwater were 0.26-2.1 nmol L -1, with up to 5.5 nmol L -1 in rivers; in the Dead Sea, where the TFA may have been concentrated through evaporation processes, the concentration was 56 nmol L-1. CF3C(O)C1 from the oxidation of HCFC-123 is rapidly taken up into cloudwater (Table 13.8). However, it also photolyzes (Rattigan et al., 1993; Wallington, 1994a; WMO, 1995), with an estimated tropospheric lifetime of ~ 33 days assuming a quantum yield for dissociation of unity (Rattigan et al., 1993): CF3C(O)C1 + h u ~ CF 3 + CO + C1.

(27)

The subsequent reactions of CF 3 are as discussed earlier. In the troposphere, the most likely fate of C1 is reaction with organics (see Chapter 6). The aldehyde HC(O)CFC12 formed in the oxidation of HCFC-141b is expected to photolyze in the troposphere, forming in part CFHC12 (Wallington et al., 1994a), which is itself oxidized in the troposphere by reaction with OH. By analogy to the photolysis of

752

13. SCIENTIFICBASISFOR CONTROLOF HALOGENATEDORGANICS

nonhalogenated aldehydes, formation of HCO + CFC12 is also expected. Because HC(O)CFC12 has an abstractable hydrogen atom, it reacts with OH in the troposphere: H C ( O ) C F C I : + OH ~ C(O)CFCI: + H : O .

(28)

With a rate constant at 298 K of 1.2 • 10 -12 cm 3 molecule- 1 s- 1, the lifetime of this aldehyde with respect to OH at 1 • 10 6 radicals cm -3 is about 10 days (Scollard et al., 1993). Either decomposition of C(O)CFC12 to CO + CFC12 or its reaction with O 2 can occur (Tuazon and Atkinson, 1994): C(O)CFC12 + CO + CFC12,

(29)

M C(O)CFC12 + 0 2 + OOC(O)CFC12.

(30)

Tuazon and Atkinson (1994) showed that 79 _+ 7% of the reaction proceeds via decomposition (29) and 21 _+ 5% by reaction ( 3 0 ) w i t h 0 2. However, the relative importance of these two possible fates depends on the particular radical. For example, for CF3CO, more than 99% proceeds by reaction with 0 2 (Wallington et al., 1994b), whereas for C(O)CF2C1, 39% decomposes and 61% reacts with 0 2 (Tuazon and Atkinson, 1994). Reaction of the alkylperoxy radical generated in (30) with NO followed by decomposition will generate the CFC12 radical: OOC(O)CFC12 + NO + OC(O)CFC12 + NO2, OC(O)CFC12 ~ CO 2 + CFC12.

(31) (32)

Reaction with NO 2 to form the peroxyacyl nitrate CFC12C(O)OONO 2 can also occur. Kanakidou et al. (1995) have carried out threedimensional modeling studies of the global tropospheric fates of HFC-134a and other HCFCs using projected emissions and the chemistry described earlier. Table 13.9 shows their calculated lifetimes for some of the CFC replacements with respect to oxida-

TABLE 13.9 Estimated Tropospheric, Stratospheric, and Overall Lifetimes of Some CFC Replacements a

Compound

HCFC-22 HFC-134a HCFC-123 HCFC-124 HCFC-141b HCFC-142b

Lifetime for oxidation by OH (years)

Lifetime for stratospheric loss (years)

Overall lifetime (years)

13.7 13.9 1.2 6.0 9.5 20.3

214 357 62 87 90 389

12.9 13.4 1.2 5.6 8.6 19.3

a From Kanakidou et al. (1995).

tion by OH, loss in the stratosphere, and the overall lifetime from both processes. Calculated lifetimes are in the range of 1-20 years. The model predicts that the oxidation products shown in Table 13.7 will not accumulate significantly in the troposphere, with concentrations typically 1% or less of their parent compound. The reason that the ODPs of these CFC replacements are much smaller than those of the original CFCs is the presence of an abstractable hydrogen with which OH can react. However, this also means that they can also contribute to ozone formation in the troposphere. Hayman and Derwent (1997) have used their photochemical trajectory model to calculate tropospheric ozone-forming potentials of some of these CFC replacements. Table 13.10 summarizes these relative ozone-forming potentials, expressed taking that for ethene as 100. Clearly, although they react in the troposphere, their contribution to tropospheric ozone formation is expected to be very small. While we have focused here on CFC replacements, similar chemistry applies to replacements for the bromine-containing halons. For example, CF 2BrH is a potential halon substitute that will react with OH in the troposphere (DeMore et al., 1997). Through the subsequent reaction with 0 2 and then NO, the alkoxy radical CF2BrO is formed. This decomposes via scission of the weak C - B r bond to form COF 2 (Bilde et al., 1996). Similarly, CF3I is a potential halon substitute. However, as for CH3I, it photolyzes rapidly to generate an iodine atom with an estimated lifetime of less than 2 days (Solomon et al., 1994a). While iodine in the stratosphere is expected to be very effective in ozone de-

TABLE 13.10 Calculated Ozone Formation Potentials for Some CFC Replacements a Formula

Ozone formation potential b

HCFC-22 HCFC-123 HCFC-124 HCFC-141b HCFC-142b HCFC-225ca

CHF2C1 CF3CHC12 CF3CHFC1 CH 3CFC12 CH3CF2C1 CF3CF2CHC12

0.1 0.3 0.1 0.1 0.1 0.2

HFC-23 HFC-32 HFC-125 HFC-134a HFC-143a HFC-152a HFC-227ea

CHF3 CH 2F2 CF3CHF2 CF3CH2F CH3CF3 CH 3CHF2 CF3CHFCF3

0.0 0.2 0.0 0.1 0.0 1.0 0.0

Compound

a From Hayman and Derwent (1997). b Relative to 100 for ethene, C 2 H4.

REFERENCES struction if present in sufficient quantities (Solomon et al., 1994a, 1994b), as discussed in Chapter 12.E, it appears unlikely that sufficient amounts reach the stratosphere to contribute significantly to ozone destruction (e.g., Wennberg et al., 1997; Pundt et a/.,1998). The ODP is estimated to be 290 nm) and hence photolyze in the troposphere. Use the data for CH3I in Table 4.42, Chapter 4, to estimate the lifetime of this compound with respect to photolysis at the earth's surface at a solar zenith angle of 50 ~ on May 1. How does this compare to its lifetime with respect to reaction with O H at 1 • 10 6 radicals cm -3 if the rate constant is 7.2 • 10-14 c m 3 molecules-~? 4. Using the heats of formation given in Appendix I, calculate zXH(298 K) for the reactions of CH 4 with C1 and Br, respectively. As discussed in Chapter 12, the C1 + CH 4 reaction plays an important role in the stratosphere whereas the analogous reaction of bromine atoms does not. C o m m e n t on whether this difference is due to enthalpy. }

}

}

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0. J. Nielsen, "Atmospheric Chemistry of CF30 Radicals: Reaction with H20," J. Phys. Chem., 97, 7606-7611 (1993b). Wallington, T. J., and W. F. Schneider, "The Stratospheric Fate of CF3OH," Environ. Sci. Technol., 28, 1198-1200 (1994). Wallington, T. J., D. R. Worsnop, O. J. Nielsen, J. Sehested, W. J. Debruyn, and J. A. Shorter, "The Environmental Impact of CFC Replacements--HFCs and HCFCs," Environ. Sci. Technol., 28, 320A-326A (1994a). Wallington, T. J., M. D. Hurley, O. J. Nielsen, and J. Sehested, "Atmospheric Chemistry of C3CO X Radicals: Fate of CF3CO Radicals, the UV Absorption Spectrum of CF3C(O)O 2 Radicals, and Kinetics of the Reaction CF3C(O)O 2 + NO --+ CF3C(O)O + NO2," J. Phys. Chem., 98, 5686-5694 (1994b). Wallington, T. J., T. Ellermann, O. J. Nielsen, and J. Sehested, "Atmospheric Chemistry of FCOx Radicals: UV Spectra and Self-Reaction Kinetics of FCO and FC(O)O 2 and Kinetics of Some Reactions of FCO~ with O2, 03, and NO at 296 K," J. Phys. Chem., 98, 2346-2356 (1994c). Wallington, T. J., M. D. Hurley, W. F. Schneider, J. Sehested, and O. J. Nielsen, "Mechanistic Study of the Gas-Phase Reaction of CHzFO 2 Radicals with HO2," Chem. Phys. Lett., 218, 34-42 (1994d). Wallington, T. J., J. Sehested, and O. J. Nielsen, "Atmospheric Chemistry of CF3C(O)O 2 Radicals. Kinetics of Their Reaction with NO 2 and Kinetics of the Thermal Decomposition of the Product CF3C(O)OzNO2," Chem. Phys. Lett., 226, 563-569 (1994e). Wallington, T. J., and J. C. Ball, "Atmospheric Chemistry of CF30 Radicals: Reaction with CH 4, CD4, CH 3F, CF 3H, 13CO, C 2H5 F, C2D6, C2H6, CH3OH, i-C4H8, and CzH 2 " J. Phys. Chem., 99, 3201-3205 (1995). Wallington, T. J., and O. J. Nielsen, "Atmospheric Chemistry of Hydrofluorocarbons," in Progress and Problems in Atmospheric Chemistry, Advanced Series in Physical Chemistry (J. R. Barker, Ed.), Vol. 3, pp. 616-685, World Scientific, Singapore, 1995. Wallington, T. J., W. F. Schneider, J. J. Szente, M. M. Maricq, O. J. Nielsen, and J. Sehested, "Atmospheric Chemistry of FNO and FNO2: Reactions of FNO with 03, O(3P), HO2, and HC1 and the Reaction of FNO 2 with 03,"J. Phys. Chem., 99, 984-989 (1995). Wallington, T. J., M. D. Hurley, J. M. Fracheboud, J. J. Orlando, G. S. Tyndall, J. Sehested, T. E. M0gelberg, and O. J. Nielsen, "Role of Excited CF3CFHO Radicals in the Atmospheric Chemistry of HFC-134a," J. Phys. Chem., 100, 18116-18122 (1996a). Wallington, T. J., M. Bilde, T. E. M0gelberg, J. Sehested, and O. J. Nielsen, "Atmospheric Chemistry of 1,2-Dichloroethane: UV Spectra of CHzC1CHC1 and CHzC1CHC10 2 Radicals, Kinetics of the Reactions of CHzC1CHC1 Radicals with O 2 and CHzC1CHC10 2 Radicals with NO and NO2, and Fate of the Alkoxy Radical CHzC1CHC10," J. Phys. Chem., 100, 5751-5760 (1996b). Wallington, T. J., W. F. Schneider, J. Sehested, M. Bilde, J. Platz, O. J. Nielsen, L. K. Christensen, M. J. Molina, L. T. Molina, and P. W. Wooldridge, "Atmospheric Chemistry of HFE-7100 (C4F9OCH3): Reaction with OH Radicals, UV Spectra and Kinetic Data for C4F9OCH ) and C4F9OCHzO ) Radicals, and the Atmospheric Fate of C4F9OCHzO" Radicals," J. Phys. Chem. A, 101, 8264-8274 (1997). Wallington, T. J., A. Guschin, T. N. N. Stein, J. Platz, J. Sehested, L. K. Christensen, and O. J. Nielsen, "Atmospheric Chemistry of CF3CHzOCHzCF3: UV Spectra and Kinetic Data for CF3CH(.)OCHzCF 3 and CF3CH(OO. )OCHzCF 3 Radicals and Atmospheric Fate of CF3CH(O 9)OCHzCF 3 Radicals," J. Phys. Chem. A, 102, 1152-1161 (1998).

REFERENCES Weatherhead, E. C., G. C. Tiao, G. C. Reinsel, J. E. Frederick, J. J. DeLuisi, D. Choi, and W.-K. Tam, "Analysis of Long-Term Behavior of Ultraviolet Radiation Measured by Robertson-Berger Meters at 14 Sites in the United States," J. Geophys. Res. 102, 8737-8754 (1997). Wennberg, P. O., J. W. Brault, T. F. Hanisco, R. J. Salawitch, and G. H. Mount, "The Atmospheric Column Abundance of IO: Implications for Stratospheric Ozone," J. Geophys. Res., 102, 8887-8898 (1997). Wenny, B. N., J. S. Schafer, J. J. DeLuisi, V. K. Saxena, W. F. Barnard, I. V. Petropavlovskikh, and A. J. Vergamini, "A Study of Regional Aerosol Radiative Properties and Effects on Ultraviolet-B Radiation," J. Geophys. Res., 103, 17083-17097 (1998). Wilkniss, P. E., J. W. Swinnerton, D. J. Bressan, R. A. Lamontagne, and R. E. Larson, "CO, CC14, Freon-11, CH 4 and Rn-222 Concentrations at Low Altitude over the Arctic Ocean in January, 1974," J. Atmos. Sci., 32, 158-162 (1975). World Meteorological Organization (WMO), "Scientific Assessment of Ozone Depletion: 1994," Global Ozone Research and Monitoring Project, Report No. 37, published February 1995; update, Report No. 44, February 1999. Wu, F., and R. W. Carr, "Time-Resolved Observation of the Formation of CF20 and CFC10 in the CFzC1 + O2 and CFC12 + O 2 Reactions. The Unimolecular Elimination of C1 Atoms from CFzC10 and CFCI20 Radicals," J. Phys. Chem., 96, 1743-1748 (1992). Wuebbles, D. J., "Chlorocarbon Emission Scenarios: Potential Impact on Stratospheric Ozone," J. Geophys. Res., 88, 1433-1443 (1983). Wuebbles, D. J., A. K. Jain, K. O. Patten, and P. S. Connell, "Evaluation of Ozone Depletion Potentials for Chlorobromomethane (CH2C1Br) and 1-Bromopropane (CH2BrCH2CH3)," Atmos. Environ., 32, 107-113 (1998). Wuebbles, D. J., R. Kotamarthi, and K. O. Patten, "Updated Evaluation of Ozone Depletion Potentials for Chlorobromomethane

7 61

(CH2C1Br) and 1-Bromopropane (CH2BrCH2CH3)," Atmos. Environ., 33, 1641-1643 (1999). Wujcik, C. E., D. Zehavi, and J. N. Seiber, "Trifluoroacetate Levels in 1995-1996 Fog, Rain, Snow, and Surface Waters from California and Nevada," National Meeting of the American Chemical Society, April 1997, Division of Environmental Chemistry Preprints of Extended Abstracts, Vol. 37, pp. 31-33, 1997. Zabel, F., F. Kirchner, and K. H. Becker, "Thermal Decomposition of CF3C(O)OzNO2, CC1F2C(O)OzNO2, CClzFC(O)OzNO2, and CC13C(O)O2NO2," Int. J. Chem. Kinet., 26, 827-845 (1994). Zachariah, M. R., W. Tsang, P. R. Westmoreland, and D. R. F. Burgess, Jr., "Theoretical Prediction of the Thermochemistry and Kinetics of Reactions of CF 20 with Hydrogen Atom and Water," J. Phys. Chem., 99, 12512-12519 (1995). Zerefos, C. S., A. F. Bais, I. C. Ziomas, and R. D. Bojkov, "On the Relative Importance of Quasi-Biennial Oscillation and E1 Nino/Southern Oscillation in the Revised Dobson Total Ozone Records," J. Geophys. Res., 97, 10135-10144 (1992). Zerefos, C. S., A. F. Bais, C. Meleti, and I. C. Ziomas, "A Note on the Recent Increase of Solar UV-B Radiation over Northern Middle Latitudes," Geophys. Res. Lett., 22, 1245-1247 (1995). Zerefos, C. S., K. Tourpali, B. R. Bojkov, D. S. Balis, B. Rognerund, and I. S. A. Isaksen, "Solar Activity-Total Column Ozone Relationships: Observations and Model Studies with Heterogeneous Chemistry," J. Geophys. Res., 102, 1561-1569 (1997). Ziemke, J. R., and S. Chandra, "Comment on 'Tropospheric Ozone Derived from TOMS/SBUV Measurements during TRACE A,'" by J. Fishman et al., J. Geophys. Res., 103, 13903-13906 (1998). Ziemke, J. R., J. R. Herman, J. L. Stanford, and P. K. Bhartia, "Total Ozone/UVB Monitoring and Forecasting: Impact of Clouds and the Horizontal Resolution of Satellite Retrievals," J. Geophys. Res., 103, 3865-3871 (1998). Zurer, P., "Controls Tightened on Methyl Bromide, HCFCs," Chem. Eng. News, December 18, 1995, p. 8.

C H A P T E R

14 Global Tropospheric Chemis try and Climate Change Over the past several decades, there has been increasing recognition in a number of areas of the environmental impacts, both realized and potential, of human activities not only on local and regional scales but also globally. This is particularly true of changes to the composition and chemistry of the atmosphere caused by such anthropogenic activities. One example, for which there is irrefutable evidence, is stratospheric ozone depletion by chlorofluorocarbons, discussed in detail in Chapters 12 and 13. Another area of global-scale dimensions that is commanding increased attention is the potential impact of atmospheric trace gases and aerosol particles on climate, the subject of this chapter. Climate is the longterm statistical characterization of parameters describing what we commonly term "weather," such as surface temperature. For example, the mean surface temperature with its associated variability over some time period, typically taken as 30 years, is one measure of climate. Thus, climate is distinguished from short-term, e.g., day-to-day, variations, which are typically referred to as "weather." The recognition that atmospheric gases play a central role in determining the earth's climate goes back more than a century to Joseph Fourier, who proposed in 1827 that heat is trapped by the atmosphere. In 1861, John Tyndall showed that 02, N2, and H 2 do not absorb infrared radiation but that CO 2 and N20 (as well as a number of organic compounds) do. Subsequently, Arrhenius (1896)considered the role of changes in atmospheric CO 2 on the earth's temperature due to the absorption of infrared radiation. He estimated, for example, that the temperature in the Arctic would rise by 8-9~ from an increase in atmospheric CO 2 by a factor of 2.5-3 over that present in 1896. Beginning in the late 1800s, Chamberlin explored the relationship between climate and atmospheric composition. A brief history of this area is given by Weart (1997) and by Fleming (1992).

As we shall see, the interrelationships between atmospheric composition, chemistry, and climate are very complex. For example, as discussed in more detail herein, it is clear that CO 2 emissions, primarily from fossil fuel combustion, have increased dramatically over the past century, leading to substantial increases in its atmospheric concentrations. The concentrations of a number of other greenhouse gases have been increasing as well (Ramanathan et al., 1985). In the simplest approach, these increases are expected to lead to a significant increase in the surface temperature, and indeed, there is general agreement that an increase of about 0.3-0.6~ over the past century has occurred (IPCC, 1996). Thus, there is a sound scientific basis for anticipating that chemical changes in the atmosphere will impact climate. However, the interplay between all of the contributing factors and hence the ultimate quantitative impacts are, at present, not well understood and the subject of intense research activity. For example, the increase in surface temperatures over the past century has not been continuous, occurring primarily from approximately 1910 to 1940 and from 1975 to the present, with recent years being some of the warmest since extensive record keeping began about 1860. Ice core studies have documented changes over the past approximately 100,000 years in the earth's climate, prior to extensive fossil fuel use, with some of the changes occurring quite rapidly (over time scales of a decade or less). Accompanying these have been changes in the concentrations of a number of atmospheric trace gases such as CH 4, indicative of complex feedbacks occurring between the atmosphere, land masses, and oceans. Such natural variability complicates assessment of the anthropogenic influences on climate, particularly when the time scales for the effects of emissions due to human activity to be manifested can be a century or more for some gases.

762

763

A. RADIATION BALANCE OF THE ATMOSPHERE: THE GREENHOUSE EFFECT

Given the breadth and complexity of the scientific issues involved in global climate, we shall focus in this chapter primarily on the current state of understanding of the role of atmospheric composition and chemistry in determining the radiation balance of the atmosphere. However, some of the variables affecting global climate and potential feedback mechanisms with associated implications for climate change as well as the global climate record are briefly discussed as well; the reader should consult the references cited in those sections for further details. An excellent summary of research in this expanding and dynamic area through 1995 is found in Climate Change 1995: The Science of Climate Change by the Intergovernmental Panel on Climate Change (IPCC, 1996). Radiative transfer in the atmosphere is not treated here in detail. The reader should consult books by Liou (1980), Goody and Yung (1989), and Lenoble (1993) for excellent detailed treatments of this subject.

also determines the temperature of the earth's surface through its absorption and reradiation. Although, as chemists, we tend to think of the absorption and emission of radiation in molecular terms, the greenhouse effect is best thought of in terms of the energy balance of the earth-atmosphere system taken as a whole (e.g., see Ramanathan et al., 1987; Ramanathan, 1976, 1988a,b; and Wang et al., 1995). Figure 14.1 shows the solar flux outside the atmosphere, which is approximated by blackbody emission at 6000 K, and at sea level, respectively. Absorption of incoming solar radiation by O3, O2, H 20, and CO2 as the light passes through the atmosphere to the earth's surface is evident. Recall in Chapter 3 that the average total incoming light intensity per unit area normal to the direction of propagation outside the earth's atmosphere, i.e., the solar constant, is 1368 W m -2. As seen in Fig. 14.2a, this is the energy density that would strike a planar disk of area ~r r 2 (where r is the radius of the earth) centered along the earth's axis. However, this incoming solar energy is spread over the entire 4 ~ r 2 surface area of the earth. The effective incoming solar radiation per unit area of the earth's surface is therefore 1368/4 = 342 W m -2. [It should be noted that, as discussed in Chapter 4.A, collision complexes of 0 2 with a second 0 2 molecule or with N 2 also may contribute an additional small amount to the absorption of incoming solar radiation; for example, 02"02 and O 2 . N 2 may contribute an additional 0.57-3 W m -2 to this total (Pfeilsticker et al., 1997; Solomon et al., 1998; Mlawer et al., 1998).]

A. RADIATION BALANCE OF THE ATMOSPHERE: THE GREENHOUSE EFFECT 1. G l o b a l A b s o r p t i o n a n d E m i s s i o n of R a d i a t i o n In Chapter 3 we examined the interaction of incoming solar radiation in the UV and visible regions of the spectrum with atmospheric gases, which drives atmospheric photochemistry. This incoming solar radiation

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Wavelength (nm) F I G U R E 14.1 Solar flux outside the atmosphere and at sea level, respectively. The emission of a blackbody at 6000 K is also shown for comparison. The species responsible for light absorption in the various regions (03, H20, etc.) are also shown (adapted from Howard et al., 1960).

764

14. GLOBAL TROPOSPHERIC CHEMISTRY AND CLIMATE CHANGE

! 1368W m2

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F I G U R E 14. 40 Fraction ( F ) of aerosol particles that are activated to form cloud droplets as a function of the total number of particles ( N tot). The horizontal line represents the 50th percentile for 10 sets of aircraft measurements. The 5th, 25th, 75th, and 95th percentiles are also shown (adapted from Gillani et al., 1995).

or as much as an order of magnitude smaller, depending on a number of factors, the most important of which are the particle sizes, the total particle concentration, and the local cooling rate in the cloud (Gillani et al., 1995). An increase in aerosol particles that can act as CCN can increase the number of cloud droplets and their size distribution, both of which can affect the light scattering properties of clouds and hence climate. We first briefly discuss the effects of clouds on climate and then the potential impacts of anthropogenic aerosols on the formation and properties of clouds. Clouds and global climate. Clouds in the troposphere interact with both solar and terrestrial radiation in complex ways, and either warming or cooling can result (e.g., see Ramanathan et al., 1989; Ramaswamy and Ramanathan, 1989; Fouquart et al., 1990; Harrison et al., 1990; Liou, 1992; King, 1993; Hartmann, 1993; Rossow and Zhang, 1995; Ramanathan, 1995; Crutzen and Ramanathan, 1996; and Baker, 1997). Thus, marine stratiform clouds found in the boundary layer backscatter solar radiation, leading to negative radiative forcing and a net cooling. Although such clouds also absorb terrestrial infrared radiation, they emit at about the same temperature as the earth's surface. As a result, as discussed for the greenhouse gases, there is little net effect (see Section A.2). However, cirrus clouds and deep convective cloud anvils found near the tropopause emit long-wavelength infrared out to space (Fig. 14.2c) at the colder temperatures characteristic of this region of the atmosphere. Because this energy emission is at lower temperatures, the net infrared emission out to space when they are present is smaller, leading to a positive radiative forcing, i.e., to warming (Twomey, 1991). As a result, the net radiative forcing

due to low clouds over the oceans is generally negative, whereas it is positive over some continental regions with high clouds. On a global basis, the mean net effect is negative, ~ - 2 0 W m -2 (Baker, 1997). For reviews of the relationship between clouds and climate, and anthropogenic effects on them, see Hobbs (1993a, 1993b), King (1993), Hartmann (1993), Andreae (1995), Schwartz and Slingo (1996), Schwartz (1996), and Baker (1997). Since some aerosol particles, which may be solids or liquids, e.g., H 2 8 0 4 , c a n serve as cloud condensation nuclei (CCN), increased particle emissions from anthropogenic processes have the potential for increasing the number of CCN. The concentration of droplets in a cloud is therefore expected to increase, although not necessarily in a linear fashion, with the increased concentrations of aerosol particles. The formation of a larger number of droplets for a given liquid water content will lead to each droplet being smaller, i.e., shift the size distribution to smaller droplets. This effect increases the cloud albedo and hence can contribute to global climate change (Twomey, 1974, 1977a,b, 1991; Twomey et al., 1984). Evidence for these effects is discussed shortly. Finally, the shift to smaller drop sizes may decrease the precipitation rate from clouds, increasing their lifetimes and hence the average amount of cloud cover (Albrecht, 1989; Lohmann and Feichter, 1997). This is also expected to have a significant effect on global climate. It should be noted that while aerosol particles affect clouds by serving as CCN for cloud formation, the reverse is also true, i.e., clouds also affect the formation and size distribution of aerosol particles. For example, the oxidation of SO 2 to sulfate in clouds generates larger particles whose light scattering cross sections are larger than for smaller particles formed by gas-phase oxidation processes (e.g., see Lelieveld and Heintzenberg, 1992). Aerosol particles incorporated into a cloud droplet will reappear as particles when the cloud evaporates. However, if new aerosol constituents are formed by in-cloud oxidation (e.g., of SO 2 to sulfate), the size of the resulting particle will be larger than the original particle. As discussed earlier, this is a potentially important process for growing aerosol particles, which are too small to serve as CCN at low values of Smax (e.g., for marine stratus), into sufficiently large sizes (i.e., to the right of the peak in Fig. 14.38) that they can act as CCN under the appropriate conditions of supersaturation (e.g., see Hegg, 1990; and Kaufman and Tanr6, 1994). For example, as seen in Fig. 14.41, aerosol particle number size distributions in the clean marine boundary layer outside of clouds are often observed to have a bimodal distribution. The larger mode above 0.1 /xm

C. AEROSOL PARTICLES, ATMOSPHERIC RADIATION, AND CLIMATE CHANGE 350

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has been attributed to aerosol particles that have been "cloud processed" (Hoppel et al., 1986; Hoppel and Frick, 1990; Anderson et al., 1994). That is, these particles served as CCN upon which clouds formed, followed by aqueous-phase reactions and evaporation of the cloud droplets to leave larger particles. Indeed, in some studies (e.g., Van Dingenen et al., 1995), it is assumed that particles in this mode can be taken as a measure of the CCN that were available for cloud formation in the prior cloud event. As discussed in Chapter 9.A.2, such aqueous-phase processes in the atmosphere are also believed to lead to two peaks in accumulation mode particles in urban areas. The albedo (R) of a thick, boundary layer cloud that does not absorb solar radiation over a surface with zero albedo can be approximated (Twomey, 1991; Schwartz and Slingo, 1996; Baker, 1997) by Albedo = R

~

(JJ)

1-g The value of the factor a is usually taken as 1 or 2 and r is the optical depth of the cloud, defined by I / I o = e - ' , where I and I 0 refer to the transmittance of direct solar radiation in the presence and absence of the cloud, respectively. The optical depth can be approximated by r ~ 2~rr2f Nh, where reff is an effective average droplet radius for scattering of solar radiation, h is the thickness of the cloud, N is the number of cloud droplets per unit volume, and g is the asymmetry parameter for single scattering, defined as the average of the cosine of the scattering angle. A typical value of g is ~0.85 for cloud droplets. (Note that the cloud height h and N are not independent, since the number of cloud droplets affects the precipitation rate, and this alters the cloud height; e.g., see Pincus and Baker,

807

1994.) Typical cloud albedos for thick clouds in the boundary layer are ~0.5 over the ocean in midlatitudes; i.e., half of the incoming solar radiation is scattered back out to space (Baker, 1997). This approximation, Eq. (JJ), illustrates why a change in the number of cloud droplets and their size affects the cloud albedo and hence the radiative forcing (see Problem 9). An important aspect of quantifying the indirect effects of anthropogenic emissions is the recognition that the changes generated in the cloud albedo are not constant for all clouds but rather depend on the particular cloud properties. For example, the effects of an absolute increase in CCN on a cloud with a low droplet number concentration will be larger than for one with a larger droplet number concentration, even if the two have the same albedo (e.g., see Platnick and Twomey, 1994; and Taylor and McHaffie, 1994). This "susceptibility" has been expressed in different ways, for example as d R ~ d O n N ) . Platnick and Twomey (1994) have derived an expression applicable to nonabsorbing clouds for the sensitivity ( d R / d N ) of cloud albedo, R, to changes in the number of cloud droplets, N, which they term cloud susceptibility: dR 6R 3 4WOw d N - r 6---~rv 9 W "

(KK)

In Eq. (KK), r - - 2 r r r ~ f N h is the optical thickness defined above, r v is the volume-weighted moment of the cloud droplet size distribution, which can be approximated by roll, Pw is the density of liquid water, and W is the liquid water content of the cloud. Using Eq. (JJ), the term r( 6 R / 6 r ) can be shown to be equal to R(1 - R) (see Problem 10). Platnick and Twomey (1994) have applied Eq. (KK) to marine clouds off the coast of California and southern Africa, to fogs in central California, and to ship tracks. Figure 14.42 shows a typical range of susceptibilities as a function of cloud droplet size. The measured susceptibilities in these studies covered three orders of magnitude, from 5 • 10 -5 cm 3 for fogs to 0.8 • 10 .3 cm 3 for marine clouds off south Africa and 2 • 10 .2 cm 3 for thin stratus clouds off the California coast. Similarly, Taylor and McHaffie (1994) report cloud susceptibilities in the range from 10 .4 to > 8 x 10 .3 at various locations around the world. The highest susceptibilities were those with the smallest aerosol particle concentrations below the cloud base. As the particle concentration increased beyond ~500 cm 3, the susceptibility was relatively constant at ~ 5 • 10 .4 cm 3. This means that the addition of new particles to a relatively clean air mass is far more effective than for a polluted one in terms of the effect on clouds. In short,

808

14. GLOBAL TROPOSPHERIC CHEMISTRY AND CLIMATE CHANGE

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1 0

,,r!,

I ~ ~I~

,,.

20

40

0"~ (.1

,,r2

II~ ~ ~ I ~ ~I~

60

I ~I~ ~ I ~ ~ ~

80

100

120

hnpl :~ Downwelling

e

o

Z

E=

0

c

~

xO

10

0

o

~>

o o

II

z

I

1o

, ,,,,,,,I

10 2

, ,,,,,,,I

10 3

, ,,,,,,,I

10 4

I

I I I l llll

105

Non-sea salt sulfate (ng m 3)

o~ [

:=L

H O N O + HNO3,

(1)

although the detailed mechanism is not known; gaseous H N O 3 is not generated in equivalent amounts, which has been attributed to its remaining adsorbed on the surface. This overall reaction occurs on a variety of surfaces in the laboratory and hence might be expected to also occur on surfaces in other environments, such as homes. This, indeed, is the case. Pitts et al. (1985) first used differential optical absorption spectrometry (DOAS) to establish unequivocally that NO 2 injected into a mobile home forms HONO. Interestingly, the dependence of the rate of H O N O generation on the NO 2 concentration was similar to that measured in laboratory systems, consistent with production in, or on, a thin film of water adsorbed on surfaces. A number of studies have confirmed that the behavior is similar to that in laboratory systems; i.e., the rate of production of H O N O increases with NO 2 and with relative humidity. Indoor levels of H O N O as high as 8 ppb as a 24-h average and 40 ppb as a 6-h

848

15. I N D O O R

average have been reported in normal, in-use buildings and homes (Febo and Perrino, 1991; Spengler et al., 1993; Weschler et al., 1994). The ratio of H O N O to NO 2 indoors can be quite large, up to ~0.15 (e.g., Febo and Perrino, 1991; Brauer et al., 1990, 1993; Spengler et al., 1993). This can be compared to typical values of a few percent outdoors (see Chapter 11). High levels of H O N O (up to ~ 30 ppb) have also been measurd in automobiles in use in polluted urban areas, and again, the ratio of H O N O to NO 2 was quite large, ~0.4, compared to 0.02-0.03 measured outdoors in the same study (Febo and Perrino, 1995). Figure 15.5, for example, shows one set of measurements of NO 2, HONO, and NO after injection of NO 2 into a home used for research purposes (Spicer et al., 1993). As NO 2 decays, H O N O is formed. Small amounts of NO are also generated, as has been observed in laboratory studies (e.g., Spicer et al., 1989). Delayed release of H O N O was observed, suggesting that H O N O was adsorbed in part on the surfaces and was subsequently released to the gas phase. For example, when the house was purged with outside air, the H O N O levels dropped; however, on closing up the house again, the levels increased immediately to a few ppb. Degassing of H O N O from the interior surfaces of automobiles has also been reported (Febo and Perrino, 1995). Interestingly, in measurements made in a commercial office building, indoor H O N O concentrations were observed not to be enhanced under conditions of high indoor 03, suggesting that dissolved H O N O in a surface film of water is readily oxidized to nitrate by 0 3 (Weschler et al., 1994). The generation of NO in Fig. 15.5 was attributed by Spicer and co-workers to a reaction of gaseous NO 2

350

AIR POLLUTION

with adsorbed HONO" NO2(g ) q- HONO(ad) --+ H + + NO;- + NO(g).

The same process was hypothesized to explain some time periods in a commercial office building when indoor NO actually exceeded outdoor NO (Weschler et al., 1994). As is the case in laboratory systems, equivalent amounts of H N O 3 are not observed as might be expected from the stoichiometry of reaction (1), likely due to H N O 3 remaining on the surface after formation a n d / o r being taken up by surfaces. For example, Spicer et al. (1993) used atmospheric pressure ionization mass spectrometry (see Chapter l l.A.2), which could measure H N O 3 with a sensitivity of 50 ppt, during the NO 2 decay and H O N O formation in the research home. H N O 3 was only observed, and in small concentrations ( < 1 ppb), when an unvented space heater was operating, and it decayed rapidly in the absence of sources. Similarly, Salmon et al. (1990) measured H N O 3 indoors and outdoors at five museums in the Los Angeles area and found that indoor concentrations were less than 40% of those outdoors, and typically about 10%. The accumulation of nitrate on indoor surfaces in a commercial building has been reported by Weschler and Shields (1996a) and attributed to the formation and uptake of H N O 3 via reactions of NO 3 (see later) a n d / o r oxidation of nitrite (i.e., adsorbed H O N O ) in an aqueous surface film. Subsequently, it was shown that H O N O is also directly emitted by gas stoves (Pitts et al., 1989). For example, Fig. 15.6 shows the concentrations of NO 2 and H O N O measured using DOAS when two top burners of a new, residential gas kitchen stove were turned on in a mobile home with both the central ventilation and air conditioning running. NO 2 from the gas stove emissions reached almost 300 ppb, and H O N O about

20

300

400

a~

m

Q. v

250

o

200

Z

150

o

100

Z

..Q v

o o

"1-

200

60

100 "

09:36 09:50

' 10:04

"~-7'

N~ '

I0:19

'

10:33 10:48

' 11:02

0 11:16

Time

FIGURE 15.5 Concentrations of NO2, HONO, and NO after injection of N O 2 into a research house (adapted from Spicer et al., 1993).

120 90

~

z

50

0

0

Exhaust hood on

300

z

Z

150

~

15

10

"{D t'cl:l

(2)

0

30

0

60

120

180

240

z~

2

0

300

Time (minutes)

FIGURE 15.6 N O 2 and HONO measured in a mobile home with two burners of a kitchen stove on; the above-range exhaust was turned on at the time shown (adapted from Pitts et al., 1989).

849

C. CO AND SO2 45 40 u.s. National Ambient Air 1-h CO Standard

35 E v

o o

30 2520-

... :.:.:.

Occupation 8-h CO Standard

// Hourly average

15-

~k

1050

and OSHA

_ _ 1 hr peak

:::

Heater off

] ]i~ ~ 1

i

~ L

2

~ L

3

4

~ 5

~ 6

Ambient 8-hU'S'NatioAir nalcostandard

[ 7

L

8

Home number

FIGURE 15.7 Measured concentrations of CO in eight homes using unvented kerosene heaters: dark shading, hourly average concentrations with heater off; no shading, hourly average concentrations with heater on; light shading, peak 1-h concentration (adapted from Mumford et al., 1991).

45 ppb. When the exhaust hood (which was externally vented) above the range was turned on, the NO 2 and H O N O decreased substantially. Similar production of H O N O from kerosene and propane space heaters has been observed (e.g., see Pitts et al., 1989; Brauer et al., 1990; Febo and Perrino, 1991; and Vecera and Dasgupta, 1994). While the mechanism generating such substantial concentrations is not known, it may involve the recombination of OH with NO as the combustion gases cool. In a house used for investigating indoor air pollution that had natural gas fueled applicances (a convective heater, a radiant heater, and a range with four burners), both the surface reaction of NO 2 and the direct combustion emissions contributed significantly to the measured indoor HONO. When an appliance was operational, the contribution of direct emissions was the more important source (Spicer et al., 1993). In short, the "dark reaction" of NO 2 with water on surfaces is ubiquitous and occurs not only in laboratory systems but also indoors. The combination of this heterogeneous reaction with combustion sources of H O N O can produce significant concentrations of H O N O indoors. As a result, there is a concern regarding the health impacts of nitrous acid, not only because it is an inhalable nitrite but also because it is likely the airborne acid present in the highest concentrations indoors.

CO concentrations in eight mobile homes with unvented kerosene heaters either off or on (Mumford et al., 1991). Both the hourly average and peak 1-h concentrations are shown for the situation with the heater on, whereas only the 1-h average for the heater off is shown. Also shown are the United States 1- and 8-h standards and the Occupational Safety and Health Adminstration (OSHA) 8-h standard. In three of the homes, the average CO concentrations exceeded the 8-h standard and in one home, the 1-h ambient air and OSHA standard was exceeded. The ratio of indoor to outdoor concentrations of CO in homes using gas stoves has been measured to be 1.2-3.8 (Wade et al., 1975), with the highest ratios found close to the source. Similarly, higher CO levels indoors compared to outdoors have been reported for restaurants in Korea, with those using charcoal burners as well as gas giving much higher concentrations (Baek et al., 1997). Figure 15.8, for example, shows the mean indoor-to-outdoor concentrations of CO and some other air pollutants measured in restaurants in Korea using either gas only or a combination of gas and

o :~

~' "o

C. CO A N D S O 2 As for NOx, combustion sources such as gas stoves and kerosene heaters can be significant sources of indoor CO. Figure 15.7, for example, shows measured

~

5 [ I CO 4k I /

I

3 I- I

~

2

~oo

1

Respirable

particles

co~

~--I

I Gas and charcoal ] [--I Gas only I I

I I

NO2

CdH6

FIGURE 15.8 Median ratios of indoor-to-outdoor concentrations of CO and some other air pollutants measured in restaurants in Korea where either gas alone ( [] ) or gas with charcoal ( 9 ) was used for cooking (adapted from Baek et al., 1997).

85 0

15. INDOOR AIR POLLUTION

charcoal. The use of charcoal increases this ratio to more than four; in one such restaurant, an indoor concentration of 90 ppm CO was measured. In buildings where motor vehicle exhaust can be entrained from outdoors or attached parking garages, elevated indoor CO levels may also result (e.g., Hodgson et al., 1991). On the other hand, in homes and offices where there was no direct indoor source of CO, the indoorto-outdoor ratio was about one, and sometimes less. For example, in Riyadh, Saudi Arabia, CO concentrations were measured indoors and outdoors; the indoorto-outdoor ratio varied from 0 to 2, but was typically below one (Rowe et al., 1989). There have been a number of measurements of CO in the "indoor environment" of automobiles. Given that cars are major CO sources in urban areas, one might expect higher concentrations of CO during commutes and this is indeed the case. Typical CO concentrations of ~ 9 - 5 6 ppm have been measured inside automobiles during commutes in major urban areas (Flachsbart et al., 1987; Koushki et al., 1992; Ott et al., 1994; Dor et al., 1995; Fernandez-Bremauntz and Ashmore, 1995). This can be compared to peak outdoor levels of ~ 10 ppm in highly polluted urban areas (see Chapter l l.A.4c). Thus, a significant enhancement of CO inside automobiles during commutes is common. For example, Chan et al. (1991b) report a ratio of the in-vehicle CO concentration to that outdoors of ~ 4.5 in Raleigh, North Carolina. As is the case for CO, SO 2 levels indoors and outdoors tend to be similar if there are no combustion sources indoors. For example, Hisham and Grosjean (1991b) report that the ratio of indoor-to-outdoor SO 2 concentrations averaged 0.89 for three museums in southern California, with a range from 0.36 to 1.92. On the other hand, quite high SO 2 concentrations can result when there are indoor combustion sources. For example, in China, where unvented stoves using coal are used extensively for cooking and heating, average SO 2 levels indoors are typically 250 /xg m -3 (96 ppb) during the summer and 750 /xg m -3 (287 ppb) during the winter, compared to average annual outdoor concentrations of 72-94/xg m -3 (28-36 ppb) (Florig, 1997; Ando et al., 1996).

D. VOLATILE ORGANIC COMPOUNDS Volatile organic compounds (VOC) are ubiquitous components not only of ambient air but also of indoor

air environments, including offices, commercial and retail buildings, and homes (Shah and Singh, 1988). There are three sources/categories for VOC: (1) entrainment of air from outside the building, (2) emissions from building materials, and (3) human activities inside buildings. As might be expected given the nature of the sources, a very large variety of organic compounds have been identified and measured indoors (e.g., Brown et al., 1994; Crump, 1995; Kostiainen, 1995). These number in the hundreds of different compounds, with the particular species and their concentrations depending on the particular sources present as well as the air exchange rates. Table 15.2 summarizes some of the types of organics that have been measured in indoor air and typical sources (Tichenor and Mason, 1988; Crump, 1995). Because of the VOC sources present indoors, the indoor-to-outdoor concentration ratios are quite large for many compounds. For example, Table 15.3 shows some typical ratios of indoor-to-outdoor concentrations for specific compounds found in each of the classes shown in Table 15.2, which are frequently present indoors (Brown et al., 1994). These data are based on a review of the literature and include data from a number of different countries. The ratio is for all but one compound substantially greater than one. Also shown in Table 15.3 is a typical range of concentrations expressed as the overall weighted average of the geometric mean, where the weighting was done using the number of available measurements. Some of the compounds associated with the three sources--entrainment from outdoors, emissions from building materials, and anthropogenic activi t i e s - a r e now briefly reviewed. E n t r a i n m e n t of air f r o m outdoor sources. Entrainment of outdoor air through ventilation systems brings with it the species found in ambient air, which have been discussed throughout this book. Some of them, such as H N O 3 and to a lesser extent 03, can be removed on surfaces such as those in air conditioning systems, and hence the indoor concentrations tend to be lower than those outdoors. Others such as NO tend to have similar concentrations indoors and outdoors if there are no significant combustion sources indoors (e.g., Weschler et al., 1994). In the case of hydrocarbons, the concentrations of compounds that do not have significant indoor sources tend to be about the same as the outdoor concentrations. For example, Lewis and Zweidinger (1992) measured VOC in 10 homes in winter and showed that the concentrations of ethene, benzene, 2-methylpentane, methylcyclopentane, 2,2,4-trimethylpentane, and 2,3-dimethylbutane indoors were within experimental error of those outdoors.

D. VOLATILEORGANIC COMPOUNDS TABLE 15.2 Class of Compounds

Hydrocarbons

Oxygenated organics

Other organics

Chlorinated organics

S51

Some VOCs Measured Indoors and Their Sources a Species

Typical Sources

Aliphatic hydrocarbons

Paints, adhesives, gasoline, combustion products, floor waxes

Aromatic hydrocarbons (toluene, xylenes, ethyl benzene, trimethylbenzenes, styrene, benzene)

Insulation, textiles, disinfectants, plastics, paints, smoking

Terpenes (limonene, c~-pinene)

Scented deodorizers, polishes, fabrics, fabric softeners, cigarettes, food, beverages

PAHs

Combustion products (smoking, wood burning, kerosene heaters)

Acrylic acid esters, epichlorohydrin

Monomers may escape from polymers

Alcohols

Aerosols, window cleaners, paints, paint thinning, cosmetics, adhesives

Ketones

Lacquers, varnishes, polish removers, adhesives

Ethers

Resins, paints, varnishes, lacquers, dyes, soaps, cosmetics

Esters

Plastics, resins, plasticizers, lacquer solvents, flavors, perfumes

Ethylene oxide

Sterilizers (hospitals)

Toluene diisocyanate

Polyurethane foam aerosols

Phthalic acid anhydride

Epoxy resins

Sodium dodecyl sulfate

Carpet shampoo

Benzyl chloride

Vinyl tiles

Tetrachloroethylene

Dry-cleaned clothes

Chloroform

Chlorinated water

1,1,1-Trichloroethane

Dry-cleaned clothes, aerosol sprays, fabric protectors

Carbon tetrachloride

Industrial-strength cleaners

p-Dichlorobenzene

Moth crystals, room deodorants

a Adapted from Crump (1995) and Tichenor and Mason (1988).

There are, however, some specific outdoor sources that can lead to higher concentrations of certain VOCs indoors than in the general outdoor air environment. For example, gases generated in landfills or from petroleum contamination can migrate through the soil and groundwater to adjacent buildings and homes to give larger indoor concentrations, particularly in basements and crawl spaces, than otherwise expected (e.g., Moseley and Meyer, 1992; Hodgson et al., 1992; Fischer et al., 1996). In one such case, the total hydrocarbon concentration was measured to be 120 ppm in a crawl space beneath the floor of a school where petroleum contamination was present from adjacent sources, compared to < 80 ppb outdoors (Moseley and Meyer, 1992). Although concentrations in various rooms

were lower, they were still elevated compared to outdoors, ranging from 0.13 to 3.4 ppm. The use of pesticides outside buildings can also lead to enhanced concentrations of these compounds indoors. For example, Anderson and Hites (1988) measured the concentrations of chlorinated pesticides indoors and found elevated levels inside, e.g., a factor of 7 times higher for y-chlordane compared to outdoor levels. One home that had the highest indoor concentrations had been treated with chlordane about a decade earlier, presumably by subsurface injection from which the pesticide migrated into the house through cracks in the basement walls. Enhanced levels of chlorpyrifos were observed indoors in homes where soil surrounding the home had been treated on a regular

85 2

15. INDOOR AIR POLLUTION TABLE 15.3

Compound n-Alkane n-Pentane n-Hexane n-Heptane n-Octane n-Nonane n-Decane n-Undecane n-Dodecane n-Tridecane n-Tetradecane n-Pentadecane Branched cycloalkanes 2-Methylpentane 2-Methylhexane 3-Methylhexane Cyclohexane Halogenated compounds Trichlorofluoromethane 1,2-Dichloroethane Dichloromethane Chloroform Carbon tetrachloride 1,1,1-Trichloroethane 1,1-Dichloroethene Trichloroethylene Tetrachloroethylene p-Dichlorobenzene m-Dichlorobenzene

Ratio of Indoor-to-Outdoor Concentrations for Some VOCs and Typical Ranges of Concentrations a Indoor / outdoor ratio (range of concentration (Ixg m -3))

Compound

Indoor / outdoor ratio (range of concentration (Ixg m-3))

Alcohols 2-Propanol n-Butanol

3 9 4 (1-5) 7 (1-5) 14 (1-5) 19 (5-10) 20 (1-5) 20 (1-5) >6 16 (1-5) >5 (1-5)

Aldehydes Acetaldehyde Butanal Hexanal Nonanal

2 2 3 4 (1-5)

12 (20-50) 4 (1-5)

Esters Ethyl acetate

15(5-10)

Terpenes Camphene c~-Pinene Limonene Total VOC

5 2 (1-5) >5 5(5-10)

Ketones Acetone Methyl ethyl ketone

Aromatic hydrocarbons Styrene Benzene Toluene Ethylbenzene rn- and p-xylene o-Xylene n-Propylbenzene 1,3,5-Trimethylbenzene 1,2,4-Trimethylbenzene (1-Methylethenyl)benzene Naphthalene

10 12 ( 375 nm, whereas the intensity of solar radiation is increasing significantly in this region. Such differences can significantly alter the photochemistry of important species such as 0 3, NO 3, and HCHO, even if the lamp output is normalized to give the same NO 2 photolysis rate constant (k~) as the sun. (3) Sunlamps A sunlamp is similar to a black lamp, except that a different type of phosphor is used and the lamp envelope transmits UV. Figure 16.5 shows a typical spectral distribution from a commercial sunlamp. The wavelength corresponding to maximum power is shifted to lower wavelengths (~310 nm), compared to black lamps, and there is significant intensity down to ~270 nm. However, the intensity decreases rapidly above ~ 330 nm. The mercury lines can again be seen superimposed on the phosphor fluorescence.

877

A. TROPOSPHERIC OZONE AND ASSOCIATED PHOTOCHEMICAL OXIDANTS

100

relatively low-energy radiation does not cause significant photochemistry in the troposphere. However, if desired, the intensity of these peaks can be decreased with the use of appropriate filters.

80-

.i-, ffl t-

60-

t>

40-

rr

200 200

I

250

300

350

400

AI

450

500

Wavelength (nm)

FIGURE 16.5 Typical spectral distribution from a sunlamp (reprinted with permission of North American Philips Lighting Corporation).

(4) Xenon lamps High-pressure xenon lamps provide the most faithful artificial reproduction of the solar energy distribution at the earth's surface in the wavelength region 290-700 nm. Figure 16.6 compares the output of an unfiltered xenon lamp to the zero air mass solar spectral irradiance (Winer et al., 1979). Unlike black lamps, xenon lamps have substantial intensity in the critical region around 300 nm; the region < 290 nm can be filtered out to match the solar energy distribution at the earth's surface using Pyrex of varying thickness (Fig. 16.10). In contrast to black lamps, the xenon lamp has maximum intensity at wavelengths above 400 nm. The Xe lamp also has a series of peaks in the 800- to 1000-nm region that do not appear in sunlight. This

(5) Measurement of light intensity In chamber studies, the spectral distribution of the irradiation source must be measured periodically (e.g., using a calibrated monochromator-photomultiplier combination) because the lamp and the windows in the chamber "age" (i.e., change with time). In addition to the spectral energy distribution of the lamp, the total absolute light intensity also must be measured. In particular, the intensity of the region below 430 nm where the most important photochemistry (e.g., of NO2, 03, and H C H O ) occurs is of greatest interest. Both of these calibrations are tedious and must be carried out with care. However, they are sufficiently critical to data interpretation that they are carried out frequently. For example, measurement of the total absolute light intensity is typically carried out after every four or five runs, and in some cases where knowledge of the light intensity is essential, after every run. The photolysis rate for NO 2 might be expected to be a good (although nonspecific) indicator of the intensity in the region < 430 nm since it absorbs strongly (see Chapter 4.C) and is also one of the major photochemically active species in V O C - N O X systems. Thus a standard procedure in smog chamber studies is to measure the rate of photolysis of NO 2 (k 1) as a relative measure of the total light intensity: NO 2 + hu

~'E

0.24 I-

Solar / spectral ~ irradiance II /~,

0.20

,i'..i~'l

0.16

i i:::: ~},

co 0.12

i:: x!)x

kl

, O(3P) + NO.

(1)

Determining k~ is not as simple as measuring the loss of NO 2, however, since secondary reactions of the O and NO produced in (1) lead to nonexponential decays of NO 2. Thus plots of ln[NO 2] against irradiation time are observed to be curved. In early smog chamber studies, a parameter known as k d was reported as a measure of the light intensity, where k d was defined by ~

Solarsimulator ( - d ln[NO2] )

0.08

kd =

0.04 250 400 600 800 1000 1200 1400 1600 1800 2000 2200 2400 Wavelength (nm) F I G U R E 1 6 . 6 Spectral irradiation of unfiltered solar simulator compared to zero air mass solar spectral irradiance (adapted from Winer et al., 1979).

dt

limit t ~ 0

k d was thus obtained experimentally by extrapolating the NO 2 concentration-time profile back to the beginning of the irradiation, t = 0. Since k 1 is the fundamental parameter of interest, however, there has been emphasis on measuring and reporting kl rather than k d in smog chamber studies. A procedure for determining k~ from measured rates

8 78

16. APPLICATIONSOF ATMOSPHERICCHEMISTRY

of photolysis of N O 2 is described in detail by Holmes and co-workers (1973). In this procedure, N O 2 is photolyzed in the smog chamber. W h e n 0 2 is absent (i.e., in 1 atm of N2) , the reactions of interest are (1) and (2)-(7): k2

O + NO 2 O + NO 2 + M

, NO + 02, k3 k4

O+NO+M

, N O 3 + M,

(3)

> N O 2 + M,

M N O 3 + N O 2 ~ N205,

(6,- 6)

N O 3 + N O 2 --+ N O + N O 2 + 0 2.

(7)

With appropriate steady-state assumptions for O, N O 3, and N205, the kinetic expressions for these reactions can be solved to yield the following equation for k~"

N0210 (N0210)] [NO2-------~

+R 2

[NO2 ] - 1

,

where k3[M] R1-

k2

k4[M] and

R2 -

E c~ o.

g O

o

k2

Thus, knowing the total pressure of N 2 and the rate constants k2, k3, and k4, one can determine k 1 from the initial concentration of N O 2 and its loss with time. The photolysis of N O 2 can also be carried out in the presence of 0 2. However, in this case additional reactions (e.g., O + 0 2 + M --+ 0 3 + M) must be considered and the kinetic expression for kl is more complex. Of course, k~ is not an absolute light intensity m e a s u r e m e n t per se but merely an indication of the intensity in one of the wavelength regions of interest for atmospheric chemistry. It has the advantage of being simple, convenient, and inexpensive, since only monitoring instruments for N O 2 are needed and these are generally standard components of the analytical apparatus. The disadvantage is that a n u m b e r of photochemically active species (e.g., 0 3 and H C H O ) have absorption coefficients and wavelength dependencies different from N O 2 and, for these, k~ will not necessarily be a good m e a s u r e m e n t of the rates of their photochemical reactions, depending on the spectral distribution of the light source. Direct measurements of light intensity using radiometers as described in Chapter 3.2 can also be made.

0.36

k

\//

0.27

J 2,--~ / - - - ~ - - -

/

f

~ ~ ~

"'... ~ ~.~ .......... / / / ~ , . . ~ j.// . . x L T _ --,,. CHacH~O

-II1\~,/ /"".

0.18

0.09

(4)

(5)

03

0.45 -:....C3H 6

(2)

N O 3 + N O + 2NO2,

kl=~ 7 (I+R1-R2)ln

0.54

V" 9

1/~, /

0

/""

\\

........

\\\

........

60

PAN

:.._~

ji:~:...."7 ""-.

NO 120

180

~

..........

240

. . . . . . . . .

NO2 300

360

Photolysis time (min)

FIGURE 16.7 Typical primary and secondary pollutant profiles in a propene-NO~ irradiation in a smog chamber (adapted from Pitts et al., 1975).

d. Typical Time-Concentration Profiles of Irradiated V O C - N O x - A i r Mixtures Figure 16.7 shows some typical c o n c e n t r a t i o n - t i m e profiles for irradiation of a p r o p e n e - N O mixture in the evacuable chamber of Fig. 16.3. The loss of the reactants, and the formation of the most commonly monitored secondary pollutants 03, PAN, and the oxygenates H C H O and C H 3 C H O are shown (Pitts et al., 1975). With in situ spectroscopic techniques, critical data on the formation of such species as H O N O , HNO3, and NO3, which are essential to understanding the chemistry of these systems, can also be obtained. Figure 16.8, for example, shows one portion of an F T I R spectrum obtained in a chamber run for a p r o p e n e - N O x

t 0.4

FHOONO2 HNOa l /

..

I/ I 0.2

N205

CHaONO2

I

HNOa

Oa -, 1HCOOH [

,

HCHOl CH3CH = CH2

,0, k7 0.0

IPANI

, , , , I , , , , I , , , , I , , , , O0 700 800 900 1000 11 Wavenumber (cm1)

FIGURE 16.8 Infrared spectrum in the 700- to 1100-cm-] region of a hydrocarbon-NO X mixture irradiated for 139 min. Initial conditions were 10 ppm propene, 1 ppm n-butane, 1 ppm neopentane, and 5.4 ppm NOx at 48~ Path length, 85 m; resolution, 0.125 cm -1 (adapted from E. C. Tuazon, unpublished data, and Pitts et al., 1977).

879

A. TROPOSPHERIC OZONE AND ASSOCIATED PHOTOCHEMICAL OXIDANTS

mixture carried out at relatively high reactant concentrations and 48~ (Pitts et al., 1977). A variety of species that are difficult to monitor with other techniques, such as H N O 3, N20 5, and peroxynitric acid (HO2NO2), are easily identified and measured by FTIR if their concentrations are sufficiently high. Clearly, such experiments can be used to examine the relationship between primary emissions and the formation of a host of secondary pollutants. For example, runs can be carried out at varying initial concentrations of hydrocarbon and NOx, and the effects on the formation of secondary pollutants such as 0 3 studied. The reactivity of various hydrocarbons can be examined by studying them singly or in combination. In addition, such parameters as temperature, relative humidity and total pressure, presence of copollutants, and spectral distribution of the light source can be systematically varied. One example of the use of chambers to study the effects of addition of copollutants is seen in Fig. 16.9 (Pitts et al., 1976). As discussed earlier, one source of HO 2 free radicals in ambient air is the photolysis of formaldehyde: H C H O + h u ~ H + HCO, M

H + 0 2 ~ HO2, HCO + 0 2 ~ H O 2 + C O . One might anticipate that the addition of H C H O prior to irradiation would increase the rate of conversion of the reactants to secondary pollutants such as 0 3 by providing an immediate source of the free radicals needed for the chain oxidations. As seen in Fig. 16.9, this is precisely what is observed.

Chambers also allow such parameters as the spectral distribution of the light source to be varied systematically. While this is impossible in ambient air studies, it is an important variable in order to simulate atmospheric chemistry at various altitudes. Since Pyrex glass absorbs radiation with A < 350 nm (Calvert and Pitts, 1966), different thicknesses of Pyrex can be used to provide different amounts of filtering in this portion of the actinic UV. Figure 16.10, for example, shows the percentage transmission of two different thicknesses of glass (3/32 and 1 / 4 in.) in the 250- to 400-nm region (Winer et al., 1979); as expected, the thicker the glass, the less transmitting it is below 350 nm. Since much of the important photochemistry (e.g., that of 0 3 and HCHO) occurs in this region, one might expect the rates of free radical formation and hence the overall rates of product formation to be less for the most filtered case. Figure 16.11 shows the formation of 0 3 in a chamber as a function of time for a p r o p e n e - n - b u t a n e - N O X mixture irradiated using a xenon lamp (vide infra)with two glass filters of different thicknesses (Winer et al., 1979). As expected, 0 3 is formed more slowly and the peak concentration is lower for the most highly filtered light source. In summary, chamber studies are a highly valuable experimental technique for studying atmospheric chemistry and the effects of varying parameters under controlled conditions.

100 -

0.5

80-

[ H C H O ~

-

E 03

_

J

. / / i 8 5 PPL

E

._o (/) E (/) -~

,,i-,

0

0.1

m 0

"-"i," opt,// .......... -" ............ . .... : 2

4

6

. 8

Irradiation time (hours)

FIGURE 16.9 Effect of added HCHO on ozone formation in irradiations of a hydrocarbon-NOXmixture. Average VOC was 2450 ppbC; average NOx was 0.33 ppm (adapted from Pitts et al., 1976).

//

60 -

(/)

3/32"/

(-

0.3

III /// /

/

40-

20-

/I I 0 250

,,"Jl

I 300

350

I~ 400

Z(nm) F I G U R E 1 6 . 1 0 Transmission spectra of two Pyrex filters of thicknesses ~ and z1 in., respectively (adapted from Winer et al., 1979).

~0

16. APPLICATIONS OF ATMOSPHERIC CHEMISTRY

0.84 0.72 /

Q.. c'~

0

0.48 0.36

0.24

\

I

0.60 E

/

\ \

3/32"/

"-,,,

-

/

_

i

/ 10 -12 >8 • 10 -11

>100 days 10-100 days 1-10 days 2-24 h [diand trialkyl aromatics, terminal alkenes] > ethylene > [monoalkyl aromatics] > [C 5 and larger alkanes] > [C2-Csalkanes]. The reactivities derived using two different scales are generally in agreement with each other (and with subsequent scales; see below), but there are some notable exceptions. For example, non-methane alkanes are predicted to be relatively more important on the basis of O H reactivity than the other scales would imply. This arises because the O H scale ignores important mechanistic aspects of reactions in irradiated V O C - N O X mixtures. Thus the initial rate of O H attack does not reflect whether that reaction ultimately leads to the generation of free radicals and hence to continued photooxidation; if the organic does not do so, that is, if it acts as an inhibitor of the chain photooxidations by removing OH, the initial rate of O H attack may assign too high a reactivity. In addition, the O H reactivity scale also does not take into account the nature of the products formed. For example, organics producing highly photolabile species that photolyze to form free radicals may be more reactive overall than indicated by the initial rate of O H attack. In addition, if products such as plant phytotoxicants or lachrymators are formed, the reactivity on the O H scale may be too slow. However, despite these potential deficiencies, the reactivity of organic mixtures in ambient air assessed using the O H reactivity scale has been found to give results that are generally consistent with the reactivity based on ozone formation (e.g., see Uno et al., 1985). The concept of O H reactivity has been applied to give a "first-cut" assessment of the contribution of various individual organics and sources to photochemical oxidant formation in a number of situations. For example, Chameides et al. (1992) scaled the contribution of various V O C concentrations in a variety of atmospheres from remote to polluted urban areas using O H reactivity. They concluded that while NOx concentrations decreased from polluted urban areas to rural to remote regions, the total V O C reactivity assessed in this manner was comparable at all continental areas from remote to polluted. In short, while the O H reactivity scale has a number of caveats associated with its use, it has proven useful in providing at least an initial assessment of relative contributions of organics to photochemical smog formation. 2. A p p l i c a t i o n t o C o n t r o l of M o b i l e Source Emissions There has been extensive work on quantitatively assigning reactivities to individual organics in auto ex-

909

haust based on their ozone-forming potential. This has in part been catalyzed (no pun intended!) by the development in California of reactivity-based emission standards for organics from passenger cars, light-duty trucks, and medium-duty vehicles (Table 16.7) (Croes et al., 1992). Four classes of "low-emission" vehicles were established: Transition (TLEV), Low (LEV), Ultra Low (ULEV), and electricity-powered, zero-emission vehicles (ZEV). The TLEV, LEV, and U L E V standards incorporate the concept of reactivity-weighted mass emissions of VOC, concurrent with increasingly strict NO~ control. The intent is to regulate based on equal ozone-forming potentials of the V O C emissions rather than simply on their total mass. That is, the emission standards for organics are set in terms of the amount of ozone formed in the atmosphere per mile traveled by a given vehicle/fuel combination rather than in terms of the simple total mass of V O C emitted per mile. The central element to this new strategy is a quantity called the reactivity adjustment factor (RAF). The R A F for a given "new generation" fuel is defined as the ratio of the ozone formed from 1 g of V O C exhaust emissions from a vehicle operating on that test fuel to the ozone formed by 1 g of V O C exhaust emissions from that vehicle operating on current, conventional, industry-average, base gasoline. The number of grams of ozone per gram of V O C exhaust emissions is defined

TABLE 16.7 California's Low-Emission Vehicle Standards for Passenger Cars and Light-Duty Trucks: Introduction of Ozone Reactivity Adjustment Factors a Vehicle category

1993 TLEV (transition) LEV (low) ULEV (ultralow) ZEVd (zero)

G r a m s / m i l e by pollutant NMOG b

NO x

CO

HCHO

0.250 0.125 0.075 0.040 0.000

0.4 0.4 0.2 0.2 0.0

3.4 3.4 3.4 1.7 0.0

0.015c 0.015 0.015 0.008 0.000

a These are 50,000-mi exhaust emission standards; fleet average NMOG requirements begin at 0.25 g/mi in 1994 and are progressively reduced in subsequent years to a level of 0.062 g/mi. Any combination of TLEV, LEV, ULEV, ZEV, and 1993 conventional vehicles can be used. b NMHC (non-methane hydrocarbons) for current and 1993 standards, NMOG (non-methane organic gases) with reactivity adjustments for others. c Methanol-fueled vehicles only. d Does not include power generation emissions, which are equivalent to 0.004, 0.02, and 0.017 g/tool for NMOG, NOx, and CO, respectively, based on South Coast air basin power plant emissions associated with electric vehicles.

910


as specific reactivity. Hence the R A F is given by Specific reactivity of emissions with test fuel RAF =

Specific reactivity of emissions with base fuel

(A) Determination of the specific reactivity of the exhaust emissions requires accurate knowledge of both the types and amounts of compounds emitted as well as how each contributes to O 3 formation. The latter factor, the ozone-forming potential, is treated in terms of its incremental reactivity (IR), which is defined as the number of molecules of ozone formed per VOC carbon atom added to an initial "surrogate" atmospheric reaction mixture of VOC and NOx: IR =

A[O 3]

A[C atoms of V O C added] "

(B)

Table 16.8 shows incremental reactivities calculated for CO and some typical VOCs found in exhaust emissions (Carter and Atkinson, 1987, 1989; Carter, 1994). An important feature is that the IR of a VOC varies with the V O C / N O ~ ratio of the air mass into which it is introduced. The IR values peak at a ratio of ~ 6 and drop off significantly at higher values, actually becoming negative for benzene and toluene. The reason for the latter is that during the oxidation of these organics, they act not only as free radical generators to convert NO to NO 2 but also as NO~ sinks. The IR of benzaldehyde is negative at all V O C / N O ~ since its photooxida-

T A B L E 16.8

tion does not result in radical generation, yet does remove NO X. The incremental reactivity of a V O C is the product of two fundamental factors, its kinetic reactiuity and its mechanistic reactiuity. The former reflects its rate of reaction, particularly with the OH radical, which, as we have seen, with some important exceptions (ozonolysis and photolysis of certain VOCs) initiates most atmospheric oxidations. Table 16.8, for example, also shows the rate constants for reaction of CO and the individual V O C with OH at 298 K. For many compounds, e.g., propene vs ethane, the faster the initial attack of OH on the VOC, the greater the IR. However, the second factor, reflecting the oxidation mechanism, can be determining in some cases as, for example, discussed earlier for benzaldehyde. For a detailed discussion of the factors affecting kinetic and mechanistic reactivities, based on environmental chamber measurements combined with modeling, see Carter et al. (1995) and Carter (1995). The peak IR value of a VOC is known as its maximum incremental reactiuity (MIR). The MIR of some VOCs are given in Table 16.9 and shown schematically in Fig. 16.35. (Note that the units of MIR used are grams of O 3 per gram of VOC added, rather than on a molecule per C atom basis as for the IRs in Table 16.8.) Note the very low reactivity for methane, as discussed earlier. These reactivities are in generally good agreement with experimental values measured

Typical C a l c u l a t e d I n c r e m e n t a l Reactivities and M a x i m u m Ozone as a F u n c t i o n of the V O C [ N O x Ratioa

V O C / N O x ratio Base case max 0 3 (ppb) b

4

6

8

10

72

160

214

215

Molecule (k on )c

CO (2.4 • 10-13)c Ethane (2.5 • 10 -13) n-Butane (2.4 • 10 -12) Ethene (8.5 • 10 -12) Propene (2.6 • 10-11) trans-2-Butene (6.4 • 10-11) Benzene (1.2 • 10-12) Toluene (6.0 • 10-12) Formaldehyde (9.2 • 10-12) Acetaldehyde (1.6 • 10-11) Benzaldehyde (1.3 • 10 -11) Methanol (9.3 • 10 -13) Ethanol (3.2 • 10-12)

12 209

16 194

20 180

40 139

Incremental reactivity (Molecules of O 3 / C atoms of VOC added)

0.011 0.024 0.10 0.85 1.28 1.42 0.038 0.26 2.42 1.34 -0.11 0.12 0.18

0.025 0.054 0.22 1.65 2.04 2.02 0.082 0.52 3.28 1.83 -0.15 0.27 0.37

0.022 0.041 0.16 0.90 1.03 0.97 0.033 0.16 1.20 0.83 -0.27 0.17 0.22

0.018 0.031 0.12 0.64 0.61 0.62 0.011 0.04 0.77 0.55 -0.34 0.12 0.14

0.016 0.026 0.098 0.50 0.51 0.48 0.003 -0.021 0.48 0.42 -0.37 0.091 0.10

0.012 0.018 0.069 0.33 0.39 0.31 -0.002 -0.036 0.32 0.29 -0.41 0.066 0.065

0.010 0.015 0.052 0.30 0.25 0.23 -0.004 -0.058 0.24 0.24 -0.41 0.055 0.038

0.005 0.007 0.019 0.14 0.14 0.054 -0.002 -0.051 0.051 0.098 -0.40 0.029 0.006

a From Carter and Atkinson (1989) calculated using the EKMA model with low dilution and an eight-component surrogate mixture of organics chosen to be representative of emissions into California's South Coast air basin. Note that the absolute values used may have changed for some components since the original publication due to further refinement of the model. b Peak ozone predicted from photolysis of initial base case mixture with the specified V O C / N O X ratio. c Rate constants for reaction with OH at 298 K in units of cm 3 molecule -1 s-1 taken from Atkinson (1989, 1994) and Atkinson et al. (1997a, 1997b).

B. REACTIVITYOF V 9 T A B L E 16.9

Maximum Incremental Reactivities (MIR) for Some V O C s MIR a (grams of 0 3 formed per gram of VOC added)

VOC

Carbon monoxide

0.054

Methane Ethane Propane n-Butane

0.015 0.25 0.48 1.02

Ethene Propene 1-Butene 2-Methylpropene (isobutene)

7.4 9.4 8.9 5.3

1,3-Butadiene 2-Methyl- 1,3-butadiene (isoprene)

10.9 9.1

c~-Pinene /~-Pinene

3.3 4.4

Ethyne (acetylene)

0.50

Benzene Toluene m-Xylene 1,3,5-Trimethylbenzene Methanol Ethanol

0.42 2.7 8.2 10.1 0.56 1.34

Formaldehyde Acetaldehyde Benzaldehyde

7.2 5.5 - 0.57

Methyl ten-butyl ether Ethyl tert-butyl ether Acetone C4 ketones

0.62 2.0 0.56 1.18

Methyl nitrite

9.5

a From Carter (1994).

using environmental chambers (e.g., Kelly and Chang, 1999). The R A F defined by Eq. (A) is calculated from the amounts of the individual compounds emitted and their MIR values as follows (Croes et al., 1992; California Air Resources Board, 1992): R A F = ~_,FT,i(MIRi) / ~_,FB,i(MIRi).

(C)

In Eq. (C) FT, i is the mass fraction of compound i in the exhaust from the test fuel, FB, i is the mass fraction of compound i in the exhaust from the base fuel, and MIR i is the maximum incremental reactivity of V O C i. Thus, a specific vehicle/test fuel combination with an R A F of 1.0 is expected to have the same contribution to ozone formation (in terms of specific reactivity, grams of 0 3 per gram of V O C exhaust emissions) as that vehicle operating on base gasoline. An R A F < 1.0 means that 1 g of the V O C exhaust emissions will form

91 1

less ozone than when the vehicle is run on the base fuel. As might be expected, different fuel compositions have a significant effect on the RAF; the amount of aromatics and oxygenates such as methyl tert-butyl ether (MTBE) in the fuel, sulfur content, and distillation temperature are important determinants of the R A F (e.g., see Ho and Winer, 1998). Table 16.10 shows the reactivity adjustment factors determined for light-duty motor vehicles operated on reformulated gasoline, M85 (85% methanol, 15% gasoline), compressed natural gas, and liquefied petroleum gas. There is no universally accepted definition of "reformulated gasoline"; that shown in Table 16.10 is the "Phase 2" gasoline mandated in California starting in 1996. It has a lower vapor pressure (which will lower evaporative emissions) as well as a lower sulfur content, the latter improving the efficiency of catalytic converters used for hydrocarbons, CO, and NO X. Other changes include lower concentrations of toxics such as benzene and added oxygenated organics such as alcohols or ethers, which may reduce CO. For a detailed description of reformulated gasolines, see Calvert et al. (1993) and Ho and Winer (1998). The allowed organic (NMOG) mass emissions (grams per mile) for TLEV, LEV, and U L E V are calculated as the appropriate standard given in Table 16.7 divided by the RAF. For example, for an R A F of 0.41, typical of T L E V fueled with M85, the allowed emissions are (0.125/0.41) = 0.30 g per mile (California Air Resources Board, 1993). Table 16.11 summarizes the allowable emissions for various potential fuels. The standards allow a larger total emission of organics as long as the ozone-forming potential of the emissions does not exceed those from conventional gasoline-fueled vehicles. For example, vehicles fueled on compressed natural gas can emit more than double the mass of those on reformulated gasoline because of the relatively low reactivity of the organics emitted. Indeed, the allowed emissions would be even larger in this case if only unburned methane was emitted; however, there are small concentrations of larger organics in the exhaust (and in the fuel itself) that increase the net reactivity of this fuel. It should be noted that while the RAFs give the relative contribution of a test fuel to 0 3 formation, they do not reflect changes in total mass emissions that may also occur with the change in fuel. The net impact of a fuel-vehicle combination can only be assessed in terms of the amount of 0 3 formed per vehicle mile traveled. For example, in California in 1996 the only gasoline allowed was the reformulated Phase 2 gasoline. The R A F for T L E V and LEV operating on Phase 2 gasoline is unity. However, the total organic mass emissions

912

16. APPLICATIONS OF ATMOSPHERIC CHEMISTRY ETBE MTBE / Acetaldehyde Formaldehyde Ethanol Methanol / m-Xylene Toluene Benzene /

Propene Ethene n-Butane Propane / Ethane I Methane CO

0

I

2

I

I

4

6

I

8

10

MIR (g 0 3 per g VOC)

F I G U R E 16.35 Maximum incremental reactivities of some organics (grams of 0 3 produced per gram of VOC) (data graciously provided by B. Croes, personal communication).

are significantly lower, so t h a t t h e c o n t r i b u t i o n to n e t 0 3 f o r m a t i o n is e x p e c t e d to b e s u b s t a n t i a l l y r e d u c e d . Similarly, B l a c k et al. (1998) m e a s u r e d e m i s s i o n s f r o m a n u m b e r of v e h i c l e - f u e l c o m b i n a t i o n s , i n c l u d i n g a v e h i c l e f u e l e d on c o m p r e s s e d n a t u r a l gas. T h e y calculate a v a l u e for t h e R A F of 0.87, a b o u t twice t h a t r e p o r t e d in T a b l e 16.10. H o w e v e r , t h e R A F v a l u e s a p p e a r to d e c r e a s e as t h e e m i s s i o n r a t e s i n c r e a s e , d u e

TLEV c

LEV c

ULEV c

to i n c r e a s e d c o n t r i b u t i o n s f r o m t h e u n r e a c t i v e unb u r n e d fuel. T h e l o w e r v a l u e s are c o n s i s t e n t with hyd r o c a r b o n e m i s s i o n r a t e s of ~ 0 . 1 2 5 g p e r mile. O t h e r ways of e x p r e s s i n g t h e reactivity of o r g a n i c s in t e r m s of o z o n e f o r m a t i o n a n d t h e i r i n t e r r e l a t i o n s h i p s are d i s c u s s e d in d e t a i l by C a r t e r (1994), a n d e x t e n s i o n of t h e p r i n c i p l e s of this a p p r o a c h to o t h e r c o m p o u n d s such as nitric acid, P A N , a n d a e r o s o l p a r t i c l e s is disc u s s e d by B o w m a n a n d S e i n f e l d (1994), B o w m a n et al. (1995), a n d D e r w e n t et al. (1998). B e c a u s e t h e v a l u e s of I R a n d M I R s h o w n in T a b l e s 16.8 a n d 16.9 w e r e d e v e l o p e d using t h e E K M A b o x m o d e l t e s t e d a g a i n s t e n v i r o n m e n t a l c h a m b e r d a t a (vide

0.98 0.41 1.00 1.00

0.94 0.41 0.43 0.51

0.94 0.41 0.43 0.50

TABLE 16.11 Allowable N M O G Emissions (g [ mi) for Various Potential Fuels for Light-Duty Motor Vehicles a

TABLE 16.10 Reactivity Adjustment Factors for Light-Duty V e h i c l e - Fuel Combinations a Fuel b Reformulated gasoline M85 CNG LPG

a RAFs are adopted in Section 13 of the "California Exhaust Emission Standards and Test Procedures for 1988 and Subsequent Model Passenger Cars, Light-Duty Trucks, and Medium-Duty Vehicles," and the process for establishing RAFs is in Appendix VIII of that document. The document is incorporated by reference in Title 13, California Code of Regulations, Section 1960.1(k); see also California Air Resources Board (1992a, 1993). The last amended date is March 19, 1998. b M85 = mixture of 85% methanol, 15% gasoline; CNG = compressed natural gas; LPG = liquefied petroleum gas. c TLEV = transition low-emission vehicle; LEV = lowemission vehicle; ULEV = ultralow-emission vehicle.

Fuel b

TLEV

LEV

ULEV

RF-A RFG M85 CNG LPG

0.125 0.128 0.305 0.125 0.125

0.075 0.080 0.183 0.174 0.150

0.040 0.043 0.098 0.093 0.080

a See Title 13, California Code of Regulations, Section 1960.1(g)(1). b RF-A is industry-average gasoline; RFG is California reformulated gasoline; M85 is a mixture of 85% methanol, 15% gasoline, CNG = compressed natural gas; LPG = liquefied petroleum gas.

913

C. FIELD OBSERVATIONS OF VOC, NO x, AND 0 3

supra), it is important to establish whether the same values for the various reactivity measures would be obtained using an airshed model. Russell and coworkers have applied the CIT airshed model to this problem and concluded that there is some difference, but it is small, < _+15%, and is primarily due to temporal and spatial differences in the emission patterns (McNair et al., 1994; Russell et al., 1995). Chang and Rudy (1990a) have also shown that this approach should be valid even if the total VOC emissions are substantially altered. Other uncertainties in the reactivities due to uncertainties in the kinetics and mechanisms are discussed by Yang et al. (1995, 1996). This approach has also been applied to sources other than automobiles to assess the relative importance of various organics and sources. For example, Blake and Rowland (1995) used the concept of maximum incremental reactivity to assess the relative importance of various organics in Mexico City. They concluded that liquefied petroleum gas was a major contributor to ozone formation and that relatively small fractions of highly reactive alkenes in the gas contributed disproportionately to ozone formation. The application of reactivity approaches to consumer products is discussed in detail by Dimitriades (1996). Analogous approaches to assessing VOC reactivities have been developed by Derwent and co-workers (e.g., see Hough and Derwent, 1987; Derwent and Jenkin, 1991; and Derwent et al., 1 9 9 6 ) w h e r e a trajectory model is used to calculate the additional ozone production due to the addition of a particular VOC under conditions typical of air masses advected across northwest Europe toward the British Isles. The photochemical ozone creation potential (POCP) index thus calculated is a measure of the reactivity of the particular VOC in terms of 0 3 formation. Andersson-Sk~51d and co-workers (1992) followed a similar approach for the summer conditions in southern Sweden. Table 16.12 compares the POCP values derived by Derwent et al. (1996, 1998) and Andersson-Sk~51d et al. (1992) to the MIR approach of Carter (1994). While the general trends in reactivities predicted by each approach are qualitatively similar, there are quantitative differences. For example, the POPC values for the simple alkanes relative to ethene are larger than the MIR values. This reflects in part the details of the mechanisms used in the calculations and the time scale over which the reactions are followed as well as differences in the assumed pollutant mix into which the VOC is injected, such as the V O C / N O ~ ratio. In short, the application of the principles of varying reactivities of organics as one component of the development of cost-effective control strategies is increas-

TABLE 16.12 Some Measures of Reactivity for Various Organics a HC o r V O C Methane Ethane Propane n-Butane Isobutane n-Pentane n-Hexane Cyclohexane Ethylene Isoprene Benzene Toluene o-Xylene rn-Xylene p-Xylene Methanol Ethanol 2-Propanol n-Butanol Formaldehyde Acetaldehyde Acetone Butanone Formic acid Acetic acid Dimethyl ether Methyl ten-butyl ether

MIR b

0.2 3 6 14 16 14 13 17 100 123 6 36 88 111 89 8 18 7 36 97 74 8 16

10 8

POCP c 0.6 12 18 35 31 40 48 29 100 109 22 64 105 111 101 13 39 14 61 52 64 9.4 37 3 10 17 15

POCP d

13 50 47 41 30 45 100 40 47 47 47 21 23 20 21 26 19 12 18

a Adapted from Derwent et al. (1996, 1998). Based on taking ethene as 100 for each scale. b M I R = maximum incremental reactivity scale; from Carter (1994). c POCP = photochemical ozone creation potential; from Derwent et al. (1998). d POCP = photochemical ozone creation potential under typical NO X conditions for Sweden; from Andersson-Sk~51d et al. (1992).

ingly being accepted as a sound approach to control of ozone and other photochemical oxidants. As discussed briefly below, this is also expected to impact the control of acids and particles as well.

C. FIELD O B S E R V A T I O N S OF VOC,

NOx, AND

0 3

The most direct approach to assessing the effectiveness of a control strategy or strategies would at first glance appear to be an examination of the relationship between reductions in emissions and measured changes in air quality. As discussed in detail in "Rethinking the Ozone Problem in Urban and Regional Air Pollution" (National Research Council, 1991), this unfortunately has not been possible for a variety of reasons. For example, as discussed earlier, pollutant concentrations

914

16. APPLICATIONSOF ATMOSPHERICCHEMISTRY

are very sensitive to the particular meteorological conditions. In comparing air quality statistics, it is then important to remove variations due to meteorological changes, or at least to reduce them to the same set of conditions. Unfortunately, there is no accepted methodology for doing so; for some relevant treatments, see Rao et al. (1992, 1994, 1995, 1996), Lefohn et al. (1993), Davidson (1993), and Zurbenko et al. (1995). In addition, there are relatively few comprehensive data sets of concentrations of the primary pollutants that can be used to assess whether, in fact, their concentrations have indeed decreased as expected from the mandated reductions in emissions. It is important to note that there are also some sources of 0 3 other than emissions in populated areas. For example, episodes of increased 0 3 due to stratospheric intrusion do occur (e.g., see Logan, 1989). In addition, there has been a well-documented (see Chapter 14) increase in global tropospheric ozone levels on top of which the local and regional sources are superimposed. In short, while a number of field studies and routine monitoring over the years have been directed to examining the relationship between emissions and air quality, with obvious implications for control strategy development, interpretation of the data is not straightforward or unambiguous. With these caveats in mind, we briefly discuss a few of the many studies that have been carried out. One of the urban areas in which long-term monitoring data for the criteria pollutants are available, in addition to very detailed meteorological and chemical data collected during some intensive field campaigns, is the Los Angeles area. This area has also had stringent controls on both mobile and stationary source emissions for a number of years. Hence it provides a comprehensive case for examining the effects of control strategies. Figure 16.36 shows the measured number of days that 0 3 exceeded 0.20 ppm for 1 h for the

-13 O .13

E z

120 100

~

80

//

years 1976 to 1991, as well as the number adjusted for differences in meteorology using an empirical model (Davidson, 1993). An overall decrease of 64% in the number of days can be deduced from the weatheradjusted trends, in spite of a large increase in the number of vehicle miles traveled due to population increases during this time. The number of days (meteorology adjusted) above the federal air quality standard of 0.12 ppm for 1 h decreased 22% over the same period. Figure 16.37 shows trends in NO X and particle nitrate as well as SO 2 and particle sulfate in the Los Angeles region over a similar time period (Dolislager and Croes, 1996). These are shown as the mean of the top 30 annual concentrations of each species. While they have not been adjusted for changes in meteorology, effects of such changes are deemphasized by using a 3-year moving mean. It is seen that the trend in NOx and the associated particle nitrate formed from it (see Chapter 7) is clearly downward, with a decrease of about 30% from 1977 to 1993, again despite the large increase in vehicle usage in this area (see Table 10.27). It is interesting that over the same period in many other regions of the United States and in Europe and Japan, ozone levels did not appear to change as dramatically (National Research Council, 1991; Lindsay et al., 1989; Rao et al., 1992, 1994, 1995, 1996; Zurbenko et al., 1995; Fiore et al., 1998; Oltmans et al., 1998). The major difference in control strategies in California compared to the U.S. federal approach has been an emphasis on both NO~ and V O C control, rather than primarily on VOC as has been the case at the federal level. For example, Table 16.3 shows the more stringent control of both NO~ and VOC from motor vehicles in California beginning in the mid-1970s. Since 1980, however, V O C emission standards in California have been comparable to the federal standards while the allowed NO~ emissions have been smaller by a factor of two or more.

II Actual F-I Weather-adjusted

04O 00 0

J ] 976

3978

] 980

] 982

] 984

1986

] 988

] 990

Year FIGURE 16.36 Trends in weather-adjusted Stage I episode days for 03 in California's South Coast air basin during the months from May through October, 1976-1991. Stage I episodes are > 0.20 ppm 03, 1-h average (adapted from Davidson, 1993).

915

C. FIELD OBSERVATIONS OF VOC, NO x, AND 0 3

0.35

NOx

0.30 0.25 v

O

0.20

,'-

0.15 m

9 PNO3

SO2

50

9 PSO41

40

m, E

30

O

""'0-.-0. 9~ . . . .

. O . 9o . ~

_

w

_..O

....

O3

X

oZ

m

0.10 _ l

"o t--t~

20

,r O

10

=o

Z

.... 9 .... 9 .... I I - - , I .... 9 .... 9 ....

0.05 0

r

SOe

I

I

I

I

I

I

I

9....

9....

9....

9....

9....

9....

9....

9 ....

9 ....

I

I

I

I

I

I

I

I

I

i-

m

t~ 13_

I 0

78 79 80 81 82 83 84 85 86 87 88 89 90 91 92 93 94 Middle of 3-year period

F I G U R E 16.37 Trends in NOx, particle nitrate (PNO3), 8 0 2 , and particle sulfate (PSO 4) as a function of year. The 3-year means of the annual top 30 daily means are shown. NO~ and SO 2 were measured at Long Beach (western end) and particle nitrate and sulfate at Riverside (the eastern end) (kindly provided by Leon Dolislager and B. Croes, California Air Resources Board).

Chameides and co-workers (1992) examined the observed concentrations of ozone and its precursors, NO X and VOC, in a variety of tropospheric locations, from remote marine areas to polluted urban regions. Figure 16.38 shows ranges of observed NOx and OH-reactivity adjusted VOC (expressed relative to propene) in four

different types of atmospheres: urban-suburban, rural typical of the eastern United States, remote tropical Brazilian forests, and remote marine boundary layer. Also shown are isopleths for 0 3 production during midday calculated using a box model. As discussed by Chameides and co-workers, the data

/Z// / I

100

10

_

k,J

_

z

-

ura,

~

+3 -

+1

-

~

0.11

0.01 0.1

I

I

Remote marine

Remote tropical

boundary layer

forest

I IIIII

I 1

I I IIIII 10

I

I

(Net 0 3 loss) IX

I I IIII 100

1 ~1--3 I IIII 1,000

OH reactivity adjusted VOC (ppbC) F I G U R E 1 6 . 3 8 Observed NO X and OH-reactivity-adjusted VOC (expressed as propene) in various regions of the troposphere. Isopleths shown are midday rates of 0 3 production (ppb h -1) calculated using a box model (adapted from Chameides et al., 1992).

916


in Fig. 16.38 indicate a strong relationship between 0 3 and NO~, but little with the reactivity-adjusted VOC concentrations in continental areas. For example, while the reactivity-adjusted VOC increases by almost two orders of magnitude from the remote marine region to the tropical forest sites due to biogenic organic emissions, 0 3 (and NO~) remain approximately constant. These trends are similar to the isopleths predicted using a simple photochemical box model and suggest that in remote and rural regions, 0 3 is most sensitive to NO~ control, whereas in polluted urban-suburban areas, it can be sensitive to either NOx or VOC control depending on the particular conditions, i.e., which side of the ridge line one is on in the top right corner of Fig. 16.14. Given our knowledge of the detailed chemistry leading to 0 3 formation from VOC and NO~ (see Chapters 6 and 7), one might expect that any NO~ that is oxidized will form either 0 3 or other oxidized nitrogen compounds such as PAN, HNO3, and N20 5. Trainer and co-workers (1993) showed that there was indeed a linear relationship between the amount of reacted oxides of nitrogen, expressed as NOz = ( N O y - NO~) (see Chapter 7 for definitions of NOz, NOy, and NO~), and 03, with the slope giving 8.5 ppb of 0 3 per ppb of NOx oxidized. This slope, i.e., the number of molecules of 0 3 generated per molecule of NO~ oxidized, is known as the ozone production efficiency. Figure 16.39 shows a similar relationship measured at a site downwind of metropolitan Toronto, Canada (Roussel et al., 1996). The slope of this line gives 13.8 ppb of 0 3 formed per ppb of NO~ oxidized. The

80 70 60 ~,

50

m

m

0

m

0

m

40 O

30

_

,,,,p,

9

20 10 I

0

0.5

1

,

I

NOy-

FIGURE 16.39

,

1.5

I

2

,

I

2.5

,

I

3

,

3.5

NOx (ppb)

Relationship between observed concentrations of 0 3 and reacted oxides of nitrogen expressed as (NOy - NO X) at a site downwind of Toronto, Canada (adapted from Roussel et al., 1996).

intercept in Fig. 16.39 is 26 ppb, which should represent the "background" 0 3. (Note that as discussed in Chapter 14, this is not the level of 0 3 in an atmosphere unperturbed by anthropogenic emissions but rather the current global level, which is a factor of 2-3 times higher than prior to the industrial revolution due to anthropogenic influences.) Both the slope and intercept are in good agreement with values of 6-12 reported for the slope and 24-42 reported for the intercept in other studies carried out in the eastern and southeastern United States (e.g., see Trainer et al., 1993; Kleinman et al., 1994; Olszyna et al., 1994; Hastie et al., 1996; and Ridley et al., 1998). These intercept values are also consistent with measurements of current boundary layer levels of 0 3 measured independently in a number of studies (e.g., Altshuller and Lefohn, 1996). Similar values for the ozone production efficiency have been reported in other areas around the world, including Athens (e.g., Peleg et al., 1997) and L1. Valby, Denmark (Skov et al., 1997). A slightly modified approach has also been used in a number of studies in which the sum of ( 0 3 + NO 2) is plotted against NO z (e.g., St. John et al., 1998). This minimizes the effects of short-term variations in 0 3 caused by its rapid reaction with NO. Thus, when 0 3 is titrated by the NO reaction, the measured 0 3 concentrations will be small; however, the NO 2 generated is a source of 0 3 through its subsequent photolysis. Hence the sum of ( 0 3 + NO 2) is sometimes chosen as a measure of the ultimate formation of ozone. The ozone production efficiency determined from slopes of plots of ( 0 3 + NO 2) against NOz in the Nashville, Tennessee, area was measured to be typically 5-6 if it was assumed that NOy is not removed by other processes. The production efficiency appeared to be about the same for the general urban plume and for an air mass in which a plume from a power plant was also embedded. Including other losses for NOy such as deposition lowers the estimated production efficiency by about a factor of two (St. John et al., 1998; Nunnermacker et al., 1998). This relationship between 0 3 and NOz = ( N O y NO X) in areas downwind of urban centers can be anticipated, based on the oxidation of NO to NO 2 by HO 2 and RO 2 radicals and the subsequent photolysis of NO 2 to form 0 3. As discussed in Chapter 6.J, at lower NOx concentrations, reactions of HO 2 and RO 2 with themselves and each other compete with their reactions with NO. However, the oxidation of NO to NO 2 leads to 0 3 formation since photolysis of NO 2, generating 03, is a major fate for NO 2. This then gives rise to the observed relationship between NOz and 0 3. Under these low-NO~ conditions, the formation of H 2 0 2 and other peroxides is important and deter-

C. FIELD O B S E R V A T I O N S OF V O C , N O x, A N D 0 3

mined largely by the rate of formation of the precursor free radicals (e.g., Kleinman, 1991, 1994). However, when the rate of NO~ emissions is larger than the rate of radical production, this relationship would not be expected to be as clear. In this region, reaction of NO 2 with OH to form HNO 3 becomes important, removing both the free radical OH and NO 2 without forming 03, weakening the O 3 - N O z relationship. Under these high-NO~ conditions, the HO 2 + HO 2 or RO 2 reactions are also less important, leading to a decreased formation of peroxides. Kleinman and co-workers (Kleinman, 1991, 1994; Kleinman et al., 1997)examined the utility of treating VOC-NO~ chemistry in terms of these two regimes defined in terms of the relative rates of free radical production compared to emissions of NOx. Given that the rates of free radical formation vary rapidly and from location to location, transitions from one regime to another can occur diurnally and seasonally, as well as geographically. For example, such a transition can occur during the fall when photolysis of 0 3 to form O(1D) and subsequently OH by its reaction with water decreases due to seasonal decreases in UV and relative humidity. Depending on the relative strength of the NO~ emissions, this can result in a transition from the low-NO~ to the high-NO~ regime, accompanied by decreases in H 2 0 2 production and a weakening of the correlation between 0 3 and NOz. This behavior is consistent with the observations of a number of field studies. For example, Jacob et al. (1995) report that in Shenandoah National Park in Virginia (U.S.) in early September, there was a good correlation between 0 3 and NOz, with a slope of 18 compared to the range of 8.5-14 observed in other studies. In the latter part of September, the correlation was weaker (r 2 = 0.23 vs 0.49 earlier) and the slope was only 7. This weakening of the relationship between 0 3 and NOz was accompanied by a decrease in concentrations of H 2 0 2 from an average of 0.86 ppb to 0.13 ppb, as expected for a transition from the NO~-limited to the VOC-limited regime. A number of modeling studies, combined with field measurements, suggest that VOC control may be more effective than NOx in controlling 0 3 at some locations, primarily urban. This is consistent with both box and airshed model predictions in that if one is in effect in the low V O C / N O ~ regime on the top of the ridge line in Fig. 16.14a, a decrease in NO~ could actually lead to an initial increase in 0 3 before it decreases. It should be noted, however, that these highly polluted locations are generally not those at which the ozone peaks occur. As illustrated in Fig. 16.14, as the air parcel moves downwind from these low V O C / h i g h NO~ regions, it generally moves into the NO~-limited regime and it is

917

in these downwind areas that 0 3 generally peaks. This type of behavior, i.e., transition from the VOC- to the NOx-limited regime, has been observed in many regions, including the Los Angeles area, Toronto (Fuentes and Dann, 1994), Munich (Fabian et al., 1994), and the eastern U.S. (McKeen et al., 1991b; Mathur et al., 1994). Modeling studies on the generation and transport of 0 3 in urban areas are also consistent with this increased sensitivity to NOx as the air mass moves downwind (e.g., Duncan and Chameides, 1998). Based on the chemistry discussed above, the use of indicator species has been proposed to differentiate air masses in which ozone formation is more sensitive to NO~ than to VOC and vice versa (e.g., Milford et al., 1994; Sillman, 1995; Sillman et al., 1997; Lu and Chang, 1998; Sillman, 1999). For example, model calculations suggest that high values of the ratios O3/{NOy - NO~}, H C H O / N O y , and H 2 0 2 / H N O 3 reflect air masses in the NO~-sensitive regime whereas low values reflect the VOC-sensitive regime (e.g., Sillman et al., 1997; Lu and Chang, 1998). The reasons for this are found in the complex chemistry discussed above, but some generalizations can be made. For example, high 0 3 and low NOy suggest relatively large free radical sources and smaller radical sinks such as the NO 2 + OH reaction. H C H O is a measure of the oxidation products of both anthropogenic and biogenic VOC; thus, the higher H C H O relative to NOy, the further to the right of the isopleths (Fig. 16.14a) is the air mass, i.e., toward the NO~-sensitive regime. Similarly, high concentrations of H 2 0 2 relative to HNO 3 indicate that the air mass is in the high-VOC/NOx regime where ozone is most sensitive to NO~ control. Some of the uncertainties in the application of such indicators in field studies are discussed by Sillman (1995) and Lu and Chang (1998). Related to the use of indicator species is the use of "species age" in photochemical modeling studies (Venkatram et al., 1998). In this approach, the VOC and NO~ are calculated at a particular point of interest in an air basin assuming no chemical reactions, i.e., only transport occurs. The age of the VOC and NO~ since the time of emission is also calculated. The amount of 0 3 formed is then estimated using the VOC-NO~ chemistry for that time period. This approach separates transport and chemistry in an explicit manner and allows the calculation of the effectiveness of various VOC and NOx reductions at a particular location. In short, a combined VOC-NO~ reduction strategy seems to be the optimum approach to controlling ozone and other secondary air pollutants, and there is evidence from the experience in southern California that this approach is effective. For differing viewpoints, however, see articles by scientists from General Motors

918


TABLE 16.13 Specificationsof Reformulated "Phase 2" Gasolinea

Research Laboratories (Chock et al., 1981, 1983; Klimisch and Heuss, 1983; Kumar and Chock, 1984; Wolff, 1993). Measure

D. A L T E R N A T E FUELS Motor vehicles are major sources of air pollutants worldwide, and the number of vehicles is anticipated to continue growing (e.g., see Walsh, 1990). A major focus of control strategy development for mobile sources in recent years has been on the development of alternate fuels. These range from relatively minor changes in the traditional composition of gasoline, such as reformulated gasolines, to compressed natural gas (CNG), liquefied petroleum gas (LPG), alcohol fuels and their blends with gasoline, or hydrogen. There have also been significant developments in electric vehicles fueled either by batteries or fuel cells. One might expect that vehicle emissions would be related to the composition of the fuel used, and a number of studies have confirmed this (e.g., see Schuetzle et al., 1994; Siegl et al., 1992; and Kaiser et al., 1991, 1992, 1993). For example, emissions coming from a single-cylinder engine have shown that the mass emissions increase as the molecular weight of a single-component fuel increases and that benzene emissions decrease as the aromatic content of the fuel decreases (Schuetzle et al., 1994). Hence, a feasible control strategy should be the use of fuels with smaller mass emissions, reduced reactivity of the emissions, or both. We discuss briefly some of the chemical implications of the use of some of these alternate fuels. For a more comprehensive treatment of the advantages and disadvantages of alternate fuels and technologies, see the National Research Council report (1991), and for a discussion of a variety of issues associated with motor vehicle emissions, see Cadle et al. (1996, 1997a, 1997b) and Chang et al. (1991).

1. Reformulated Gasolines As discussed by Calvert et al. (1993), there is no fixed, agreed-upon, definition of reformulated gasoline, although decreased VOC emissions and toxics are common goals. Table 16.13 shows the composition of what is known as Phase 2 reformulated gasoline compared to an average for gasolines sold in California in 1990. The vapor pressure at 100~ (known as the Reid vapor pressure, RVP) and the sulfur and benzene contents are lower, as are the aromatic and olefinic contents. The reduced vapor pressure reduces evaporative emissions (which can be as much as half of the total VOC emissions), while reducing the sulfur content improves

RVP (psi) c Sulfur (ppmw) Benzene (vol%) Aromatic HCs (vol%) Olefin (vol%) Oxygen (wt%) T9o (~ Tso (~

Flat limits b

1990 California average

7.0 40 1.0 25.0 6.0 1.8-2.2 300 210

8.5 150 2.0 32 9.9 0 330 212

a From California Air Resources Board, R. Pasek and Bart Croes, personal communication. b Flat limits met by each gallon; see Title 13, California Code of Regulations, Section 2262. c Applies to summertime only; definition of summertime varies by location.

catalyst operation for the reduction of VOC, CO, and NO X. The addition of oven-containing compounds such as methyl tert-butyl ether (MTBE), ethyl tert-butyl ether (ETBE), methanol, or ethanol is intended to counter the octane reduction due to reduced aromatic and olefinic contents and to decrease emissions of CO. Tunnel studies of vehicle emissions carried out before and after there was a significant increase in the oxygen content of fuel sold in the San Francisco area (from 0.3 to 2.0%) showed CO emissions decreased by 21 _+ 7%, and VOC by 18 _+ 10%, respectively, with no change in NOx (Kirchstetter et al., 1996). Formaldehyde, however, which is an oxidation product of MTBE added to the fuel, increased by 13 _+ 6% as did MTBE (see Problem 1). A similar decrease of 18 + 11% in CO emissions attributed to the use of oxygenated fuels was reported by Johnson et al. (1997). They used remote-sensing techniques to compare CO from automobiles leaving Las Vegas, where oxygenated fuels were in use at that time in winter (1991/92), to that from automobiles entering from California, where these fuels were not then in use. It is not entirely clear from measurements of CO in ambient air, however, whether such substantial decreases in CO are a general phenomenon. For example, Mannino and Etzel (1996) have analyzed ambient CO concentrations measured before and after oxygenated fuels were used in a number of western states in the United States and compared these data to analogous measurements made in states where oxygenated fuels were not used. They report an only slightly greater decrease in CO in the areas using oxygenated fuels compared to those not using oxygenated fuels.

D. ALTERNATE FUELS T A B L E 16. 14

Organic Alkanes Alkenes Formaldehyde Aldehydes Ethene Toluene Aromatics Methyl ethyl ketone Methanol Ethanol

9 19

Organic C o m p o s i t i o n of E x h a u s t and Evaporative Emissions from Gasoline and Some A l t e r n a t e Fuels a'b

Methanol M100

GasolineC

M85

0.632 0.040 0.021 0.004 0.031 0.199 0.059 0.015 0 0

0.224 0.007 0.067 0.004 0.005 0.032 0.023 0.005 0.633 0

0.023 0.001 0.050 0.001 0.001 0.009 0.005 0.001 0.911 0

Ethanol

Liquefied petroleum gas

Compressed natural gas

0.077 0.002 0.010 0.050 0.034 0.023 0.010 0.002 0 0.791

0.797 0.062 0.041 0.005 0.082 0.007 0.003 0.003 0 0

0.170 0.031 0.023 0.005 0.017 0.007 0.014 0.009 0 0

a From California Air Resources Board (1989). b Compositions given as fractions; the tests relied on a small number of vehicles, and the tests are not likely to be representative of the actual fleet of vehicles operating on each fuel. c Indolene used as the reference gasoline.

One concern with the use of MTBE in fuels is its high solubility in water and hence contamination of groundwater and surface water from fuel leakage and spillage (e.g., Squillace et al., 1997; Reuter et al., 1998); because of this, a phase-out of MTBE is occurring in some places, e.g., California. Increased concentrations of this toxic, highly volatile compound in air are also of concern. In one study in Finland, for example, where the gas contained 11% MTBE, concentrations around the perimeter of the gas stations were measured to average 0.1-0.4 ppb; however, concentrations at the pumping islands were as high as 37 ppb (Vainiotalo et al., 1998). It is interesting that the calculated reactivity of the speciated exhaust of oxygenated fuels containing MTBE did not change significantly because the lower reactivity of MTBE (Japar et al., 1991; Carter, 1994) is offset by the increased emissions of the highly reactive HCHO and isobutene. (It should be noted that these estimates of MTBE reactivity are for gas-phase reactions; heterogeneous oxidation on particles may also be important (e.g., see Idriss and Seebauer, 1996).) This is in agreement with the reactivity adjustment factors for reformulated gasoline in Table 16.10, which are all within 10% of that for the base gasoline. Gains in terms of ozone formation are then expected to be due primarily to changes in the mass of VOC emitted (although this is somewhat controversial--see Hoekman, 1992; Venturini, 1993; and Kirchstetter et al., 1996). This suggests that the major benefit of such reformulated gasolines and increased oxygen content may be reduced CO rather than reduced ozone formation, depending on the V O C / N O x regime into which the exhaust is emit-

ted. On the other hand, modeling studies suggest that the reduction in the vapor pressure and the sulfur and olefinic contents associated with the use of reformulated gasoline should contribute to reduction in ozone production (Dunker et al., 1996). 2. C o m p r e s s e d N a t u r a l G a s ( C N G ) Because CNG is primarily methane, it is expected to have relatively low reactivity, with the small amounts of reactive "impurities" such as small olefins and alkanes being responsible for most of its reactivity (see Table 16.14). Emissions of CO are smaller than from gasoline-powered vehicles, while the effect on NO X emissions is not clear (National Research Council, 1991). As seen in Tables 16.10 and 16.11, CNG shows the highest promise for low-reactivity exhaust emissions, and this appears to be the case for its use in "real" vehicles (Gabele, 1995). Figure 16.40, for example, shows the estimated ozone production per mile traveled for a vehicle fueled on CNG compared to vehicles fueled on reformulated gasoline (RFG) or the alcohol fuels M85 or E85 (vide infra). These measurements and estimates based on them include the contributions from both exhaust (including CO) and evaporative emissions (Black et al., 1998). Clearly, the reactivity of the CNGpowered vehicle emissions was substantially smaller than for the other vehicle-fuel combinations. Additional advantages of CNG are that because of the nature of the fuel, evaporative emissions are of significantly lower reactivity than conventional gasoline and since the emissions coming directly from the engine are lower, the likelihood of such vehicles being "super-emitters" is much less.

920

16. APPLICAT I O N S OF A T M O S P H E R I C CHEMISTRY

2000 ~"

1500 1000 500 0

FIGURE 16.40 Calculated ozone production per vehicle mile traveled for various car-fuel combinations. RFG = reformulated gasoline; M85 = 85% methanol, 15% gasoline; E85 = 85% ethanol, 15% gasoline; CNG = compressed natural gas (adapted from Black et al., 1998).

3. L i q u e f i e d P e t r o l e u m G a s ( L P G ) Liquefied petroleum gas is primarily propane but generally also contains significant amounts of olefins, which increase its reactivity substantially (Table 16.14). For example, LPG in the Los Angeles area contains about 2 mol% of alkenes whereas that in Mexico City contains almost 5% (Blake and Rowland, 1995). Hence while it is a significant improvement over gasoline in terms of reactivity (Tables 16.10 and 16.11), the exact amount of improvement is highly dependent on the nature and concentrations of these reactive impurities (Gabele, 1995). 4. A l c o h o l F u e l s a n d B l e n d s w i t h G a s o l i n e Methanol and ethanol fuels are used both "neat" and as a blend with gasoline. The terminology "M85" signifies a blend of 85% methanol, 15% gasoline (by volume), and similarly for "E85," used to denote ethanol-gasoline blends. The advantages of using an alcohol-gasoline blend are that the flame is visible (a flame from pure methanol is not), the vapor pressure is higher, which aids in ignition, and there is flexibility in fueling, depending on the availability of gasoline or alcohol fuels. Methanol is expected to oxidize to formaldehyde, both during combustion and after emission to the atmosphere. As discussed in Chapter 6.H, OH reacts with methanol primarily at the methyl group: CH3OH + OH ~ CH2OH + H 2 0 CH30 + H20

(85%), (15%).

(17a) (17b)

Both CH2OH and C H 3 0 then react with 0 2 to give HCHO (and HO2). However, as discussed in detail in

the National Research Council report (1991) and references cited therein, the results of modeling studies suggest that the amount of HCHO formed by oxidation of methanol in the atmosphere in many locations will not be significant compared to that formed by other processes. (Direct emissions of formaldehyde may be important, however, in some circumstances with low dilution rates such as underground parking garages or tunnels (e.g., see Chang and Rudy, 1990b).) Because methanol itself has low reactivity (Tables 16.8 and 16.9), some modest improvements in ozone may result from the use of methanol (e.g., see Dunker, 1990; National Research Council, 1991; Lloyd et al., 1989). It is interesting that the reactivity of the evaporative emissions from M85 is higher than that of the exhaust (Black et al., 1998). There has been an emphasis on the addition of ethanol to gasoline, due to the availability from grain sources. Ethanol is somewhat more reactive than methanol (Tables 16.8 and 16.9) and forms acetaldehyde upon oxidation (see Problem 2). Further oxidation of acetaldehyde produces PAN (see Chapters 6 and 7). There is evidence for such an oxidation sequence in Brazil, where extensive use is made of ethanol as a fuel. Thus relatively high concentrations of PAN have been observed in this area compared to those found elsewhere (Tanner et al., 1988; Grosjean et al., 1990). Increases in PAN have also been reported in urban areas such as Albuquerque, New Mexico, using ethanol-gasoline blends for CO reduction (e.g., Gaffney et al., 1997, 1998; Whitten, 1998); concentrations of aldehydes, particularly CH3CHO, were also higher. Alcohol fuels have also been reported to have higher emission rates of VOC as well as toxics such as HCHO, CH3CHO, 1,3-butadiene, and benzene (Gabele, 1995; Black et al., 1998). However, Stump et al. (1996) report decreased emissions of VOC, benzene, and 1,3butadiene with a gasoline fuel containing 8.8% ethanol but, in agreement with the other studies, increased emissions of HCHO and CH3CHO. Knapp et al. (1998) tested emissions from 11 vehicles at temperatures from - 2 0 to +75~ fueled on either gasoline or a blend with 10% ethanol and found the ethanol blend resulted in higher emissions of CH3CHO, in some cases almost by almost an order of magnitude. The changes in the emissions of benzene, HCHO, and 1,3-butadiene were variable with respect to both amount and sign. An additional problem with alcohol-gasoline blends is the increase in vapor pressure of gasoline in the mixture (e.g., see National Research Council, 1991; Calvert et al., 1993; and Timpe and Wu, 1995). This can contribute to much higher Reid vapor pressures, increasing the relative importance of evaporative emissions.

E. CONTROL OF ACIDS

Table 16.14 shows the VOC composition of the combination of exhaust and evaporative emissions measured on a limited number of vehicles. Similar data have been reported by Gabele (1995). The increased aldehydes associated with the use of alcohol fuels is evident. Overall, the use of ethanol blends is believed not to be effective in reducing ozone, but may actually increase it (Calvert et al., 1993; Dunker et al., 1996).

921

nature of the emissions, i.e., replacing many individual sources that are dispersed in time and space by a single source, may also be very important in determining air quality at various locations in an air basin. Fuel cells, in which a fuel such as methane, gasoline, or methanol is converted to electricity, to power vehicles also appear promising and indeed are currently being used on some buses. For a detailed discussion of this important area, see Lloyd (1992) and Lloyd et al. (1994).

5. Hydrogen Since the oxidation of H 2 produces water, it should be one of the cleanest possible fuels, and this indeed appears to be the case, with emissions consisting primarily of NO~ and very small amounts of VOC and CO from the combustion of some oil in the system. While it is being used in some buses, for example, it has not yet found widespread use due to technical and safety problems encountered on the passenger vehicle scale.

6. Electric Vehicles Electric passenger vehicles powered by batteries have been reintroduced in the United States. Although there are no tailpipe emissions associated with electric vehicles, charging the batteries requires power with which there are associated emissions. However, emissions from large sources are relatively easier to control than from millions of individual automobiles. In addition, the emissions of VOCs and CO from power plants are generally low, while those of NO~, SO 2, and particles depend on the type of fuel used. Natural gas-fired power plants generally have relatively low SO 2 and particle emissions, whereas coal-fired power plants have higher emissions of these pollutants and of NO~. Studies by Austin described by Cadle et al. (1996) suggest that the use of electric vehicles would result in lower total emissions of NO~ in the western United States, where gas is used at power plants, but higher emissions in the east, where coal is used. Larger PM10 and SO 2 emissions were projected regardless of the power plant fuel source. Another version of the electric vehicle that has been introduced into Japan and is scheduled to be introduced into the United States is a hybrid vehicle with both electric and gasoline engines. Such vehicles have the promise of lower emissions and much longer ranges than can currently be achieved with electric vehicles. For a discussion of various aspects of electric vehicles, see Wilkinson (1997) and the May 1998 issue of Enuironrnental Manager. It should be noted that in addition to the change in total emissions, the change in the spatial and temporal

E. CONTROL OF ACIDS As discussed in Chapter 8, the major contributors to acid deposition are sulfuric and nitric acids, with a significant contribution being increasingly recognized from organic acids. The chemistry of formation of nitric and organic acids has been discussed in Chapters 7 and 8 and has been shown to be part of the complex V O C - N O X chemistry that also leads to ozone formation. Hence control strategies applied for ozone will also impact the formation of these acids as well, although not necessary proportionately, as can be seen from the detailed chemistry. For example, Meng et al. (1997) applied an Eulerian grid model that included both gas and aerosol chemistry to the Los Angeles area for conditions representative of a smog episode on August 27 and 28, 1987. Model predictions showed that the changes in the gas-phase HNO 3 due to reductions in VOC and NOx were not proportionate, which is not surprising given the complex chemistry involved. For example, a 50% reduction in NO~ emissions alone gave a predicted 17% reduction in the maximum 1-h average concentration of HNO 3 at Riverside, in the eastern end of the air basin; a concurrent reduction of 50% in VOC gave a predicted reduction in peak HNO 3 of 39%. Interestingly, the peak HNO 3 was predicted to actually increase by 17% for VOC reductions from 35 to 50% without concurrent reductions in NO~ emissions. This increase in HNO 3 was associated with predicted simultaneous decreases in PAN of 58-76%. The HNO 3 increase was attributed to lowered concentrations of RCO 3 radicals, which resulted in less NO 2 being tied up in PAN and hence being available to react with OH to form HNO 3. In short, the development of control strategies for HNO 3 is intimately tied with that of 0 3. Although the control of organic acids has not been examined in detail, similar considerations are expected to apply there as well. As discussed in Chapter 8, the formation of H2SO 4 from SO 2 occurs largely in the aqueous phase. The

922


major oxidants in fogs and clouds are H202 and 03, so that control of photochemical oxidants is again expected to impact the rate of SO 2 oxidation and formation of H 2SO4. Given that the source of oxidants for SO 2 in both the gas and liquid phases is the VOC-NOx chemistry discussed earlier and that a major contributor to acid deposition is nitric acid, it is clear that one cannot treat acid deposition and photochemical oxidant formation as separate phenomena. Rather, they are very closely intertwined and should be considered as a whole in developing cost-effective control strategies for both. For a representative description of this interaction, see the modeling study of Gao et al. (1996). One of the key issues in developing effective control strategies for acid deposition has been what is known as "linearity." This term has been subject to a variety of interpretations and meanings and applied on microscopic, i.e., molecular, to macroscopic scales. A detailed treatment and discussion of linearity encompassing these scales is given by Hales and Renne (1992). In its simplest form of interest for policy and regulatory purposes, linearity is often treated in terms of source-receptor relationships. That is, if the emissions of the precursor SO 2 are lowered by 50%, will the deposition of sulfate also decrease by 50% at all receptor sites? A major factor involved in determining the relationship between SO 2 emissions and sulfate deposition is the chemistry. As discussed above, the oxidation of SO 2 by OH in the gas phase generates HO 2 and hence OH in the presence of NO. The regeneration of OH means that the oxidation will not be oxidant limited in the gas phase, and hence a reduction in SO 2 might be expected to be accompanied by a corresponding decrease in the formation of H 2SO4. However, the situation is not as clear-cut for the liquid-phase oxidation, which, we have seen, predominates in many (perhaps most) situations. In this case, a less than 1:1 relationship between the reduction in H2SO 4 formed and SO 2 emitted may result for a number of reasons operating on the microscopic scale. For example, less H 2 0 2 is available in many clouds than is needed to oxidize all of the S(IV) that is present, and hence the oxidation can be limited by the availability of oxidant (e.g., see Dutkiewicz et al., 1995). Another important factor that comes into play is the interplay between the acidity of the aqueous phase, the reaction kinetics, and the solubility of S(IV). Thus, as seen in Chapter 8, the solubility of S(IV) decreases as the aqueous phase becomes more acidic, limiting the total sulfur available for oxidation. In addition, all oxidations in the aqueous phase except that by H 2 0 2 are pH dependent; as a result, the contribution of

various oxidants can vary significantly with time in an air mass as the oxidations take place and removal by rainout and washout occurs. Examples of these complications have been reported in a number of field and modeling studies. For example, the amount of dissolved SO 2 in rain that has passed through a power plant plume has been found in some cases to be much less than might be expected, due to the acidification of the rain by other plume components that decreased the solubility of SO 2 (e.g., Dana et al., 1975). Similarly, studies using the PLUVIUS model developed for application to gas-phase species and their interaction with clouds and precipitation (Hales, 1989) suggest that the deposition rate of sulfate from an urban source mixed in with contributions from the "background" may initially be nonlinear (Hales, 1991). However, at longer times this deposition rate becomes linear, and since most of the removal occurs at these longer times, the overall integrated deposition does indeed appear to be linear. In short, as is the case for acids formed in V O C - N O X chemistry, the chemical and physical processes associated with the formation and deposition of sulfuric acid are also quite complex. There are a number of field measurements that have addressed this relationship between the mandated reductions in SO 2 emissions in the United States and the subsequent changes in sulfate deposition downwind. For example, one analysis of the trends in the atmospheric concentrations of sulfate in the northeastern United States suggests that from 1977 to 1989, the sulfate concentration decreased by about 22-28% during which the emissions of SO 2 were estimated to have decreased by 25% (Shreffler and Barnes, 1996). Similarly, Husain et al. (1998) have reported trends in sulfate at two sites in New York state, from 1979 to 1996 at Whiteface Mountain in a remote area in the eastern part of the state and from 1983 to 1996 at Mayville, in the western part. The trends at both sites were highly correlated. Figure 16.41 shows the relationship between sulfate or total sulfur (defined as sulfate plus gas-phase SO 2) at Whiteface Mountain as a function of the estimated anthropogenic emissions of SO 2 upwind in the Midwest. The relationship is well described as linear, with a correlation coefficient of r 2 = 0.81. (However, a negative intercept suggests that this cannot be directly extrapolated down to very small SO 2 emissions.) The results of a global 3-dimensional model simulation also suggest that in the boundary layer in the United States, a 50% reduction in anthropogenic SO 2 emissions in the United States will result in a similar (53%) annual reduction in the total (wet plus dry) deposition of sulfur (Chin and Jacob, 1996).

F. CONTROL OF PARTICLES

a E

1979

4-

9 ~

30 03

E

21 ~" 4

1995~~ 9 9 2

1996 b

9

,F" ~9 9 4 ~ ~

/

03 1996

m

10

I

15

I

20

I

25

I

30

I

35

SO 2 emissions in midwest (103 metric tons per day)

F I G U R E 16.41 (a) Sulfate and (b) total sulfur (sulfate + SO2(g)) at Whiteface Mountain, New York, as a function of estimated anthropogenic SO 2 emissions upwind in the Midwest from 1979 to 1996 (adapted from Husain et al., 1998).

about 30 times greater. The implications of such changes for the ultimate health impacts of primary emissions of particles from diesel engines are not clear. The use of alternate fuels in heavy-duty vehicles has also been the subject of many studies, and it appears that a significant reduction in emissions of particles can be obtained using such fuels. For example, Wang et al. (1997) measured the emissions of particles, NOx, CO, and hydrocarbons from more than 300 buses and heavy-duty trucks operating on either diesel fuel, natural gas, or alcohol fuels in 17 different states in the United States. Figure 16.42 shows their measurements of emissions of particulate matter and NOx. Particle emissions using compressed natural gas were reduced by more than an order of magnitude compared to those using diesel fuels; improvements, although not as dramatic, were also seen with the alcohol fuels. On the other hand, only the alcohol fuels gave significant reductions in NO~ emissions. As discussed in Chapter 9, in many locations secondary particles are more important contributors to atmospheric levels than are the primary emissions. The

In short, it appears that despite the potential complexities in the chemistry and meteorology involved in these long-range transport processes, at the SO 2 emission levels over the past decade, the deposition of sulfate has decreased approximately proportionately to the decrease in anthropogenic SO 2 emissions.

~)

1.8 1.6 1.4 1.2

n

0.8 0.6 0.4

E F. C O N T R O L

OF PARTICLES

As discussed in Chapter 2, epidemiological studies have suggested that small particles may cause significant health effects, including increased mortality rates. Depending on location and time, such particles may be either primary in nature, i.e., directly emitted, or secondary, i.e., formed from reactions of gases in air. Diesel engines are significant sources of particles, indeed the major source in some urban cores, and their use in many regions of the world is increasing (e.g., see Walsh (1997) and Chapter 10). As a result, there has been an emphasis on reducing these emissions, for example through engine redesigns. One concern has been the size distribution of the emissions. Thus, emission standards are written in terms of total mass, which reflects primarily the larger particles. In one laboratory study (Baumgard and Johnson, 1996), the total particulate mass emissions from a 1991 heavy-duty engine burning a 0.01 wt% S (low sulfur) fuel were much smaller than from a 1988 engine using the same fuel. However, in terms of the n u m b e r of ultrafine particles (0.0075-0.046 /xm), the newer engine emissions were

923

I-I

1

0.2 0

CNG

Diesel #1

35

(1)

E

v

o') x

O Z

25

Diesel #2

E93/E95

M100

FI

30 . m

I-I

m

20

I-I

15 10 5 0

CNG

Diesel #1

Diesel #2

E93/E95

M100

Fuel type F I G U R E 16.42 Measured emissions of (a) particles (PM) and (b) NO X from more than 300 buses and heavy-duty trucks running on compressed natural gas (CNG), two diesel fuels, or alcohol fuels (E93, E95, and M100). The bars are the means and the boxes encompass the 95% confidence intervals. E93 = 93% ethanol, 5% methanol, 2% K-1 kerosene; E95 = 95% ethanol, 5% gasoline; M100 = 100% methanol (adapted from Wang et al., 1997).

924

16. A P P L I C A T I O N S

OF ATMOSPHERIC

chemical components of the secondary particles consist largely of sulfate, nitrate, and complex polar organics, all of which are formed by oxidation processes in the VOC-NO~ system. This again suggests that control strategies applied to ozone should also impact the formation of secondary particles. For example, less NO\ might be expected to give less particle nitrate in general, and less VOC should reduce the formation of secondary organics, including those in particles. The actual relationship will be more complicated than a one-to-one correlation, of course, in that the concentrations of reactive species such as OH needed to carry out the oxidations are also affected, but not in a linear manner, by changes in VOC and NO~. A model for examining such relationships between individual organics and the production of secondary aerosol particles is discussed by Bowman et al. (1995). Trends in particulate matter with diameters less than 10 /,m, PMm, in the Los Angeles air basin from 1985 to 1994 are shown in Fig. 16.43 (Prasad, personal communication, 1997). A downward trend in both the mean and the maximum levels is evident, occurring along with similar trends in 03, NO~, and particulate nitrate and sulfate in this region (see Figs. 16.36 and 16.37). This is qualitatively consistent with expectations

100 \

90--

\

\

\ //

/

/

/--...-

"\

\

/

Maximum

\ \

80--

\

70--

E

~D3

o it.

60-50--

based on the chemistry described in earlier chapters and again illustrates the strong interrelationships between various species in this complex atmospheric system. Similar downward trends in PM10 have been observed at surface measuring sites in the United States in urban, suburban, and rural areas. Figure 16.44, for example, shows the trends in the annual average PM m levels in these three types of air environments (Darlington et al., 1997). Reductions of ~ 3 - 4 % per year have been observed. Simultaneously, the annual average gas-phase concentrations of SO 2 and NO\, precursors to sulfate and nitrate in particles, decreased. Downward trends of ~ 1.6-1.8% per year in the optically active aerosol over the United States has also been reported (e.g., Hofmann, 1993). Until relatively recently, air quality models focused on gas-phase reactions. However, with increases in our understanding of the formation, interactions, and fates of suspended particles, as well as increase in computing power, significant progress is being made in incorporating aerosol chemistry and physics into local and regional models (e.g., see Wexler et al., 1994; Eldering and Cass, 1996; and Meng et al., 1997). For example, Meng et al. (1997) used an Eulerian grid model for the Los Angeles area to predict the relationships between the precursors VOC and NO\ and their secondary oxidation products, including gaseous HNO 3 (discussed earlier), PM2.5, and nitrate in PM2. 5. Because HNO 3 is predicted to increase with reductions in VOC when there are no simultaneous reductions in NO~, both nitrate in PM2. 5 and total PM2. 5 also increase under this scenario. Reductions in NO~ without concurrent reductions in VOC do lead to reductions in PM2. 5 and nitrate in PM2.5, but not proportionately. For example, a 50% reduction in NO~ is predicted to decrease both PM2. 5 and its nitrate component by ~20%.

40-cO

30--

.4...,

cO 0 cO 0

20--

o

v--

10-0 84

CHEMISTRY

1.05 1.00

0.95 0.90 0.85

~ ~'~

..,,, "'~...,,

/ .... Urban

~. ,~

Suburban

0.80 I 86

I 88

I 90

I 92

94

Year

FIGURE 16.43 Trends in PM10 in the South Coast air basin in California (Los Angeles area) from 1985 to 1994. The averages for six sites having complete data throughout this time period are shown as well as the maximum concentrations. The 1994 data are for January to October (kindly provided by Dr. Shankar Prasad, California Air Resources Board).

N

E 0 Z

0.75 -

/ Rural

0.70 0.65 1988

I 1989

I 1990

I 1991

I 1992

I 1993

I 1994

1995

Year

FIGURE 16.44 Concentrations of PM10 normalized to 1988 as the base year in the United States for urban, suburban, and rural areas (adapted from Darlington et al., 1997).

G. ATMOSPHERIC CHEMISTRY A N D RISK ASSESSMENTS OF H A Z A R D O U S AIR P O L L U T A N T S

However, as expected from the chemistry discussed in Chapter 7.E, simultaneous control of NH 3 has a significant effect on particle nitrate formation, since the formation of ammonium nitrate is a major mechanism for conversion of gaseous HNO 3 to particulate nitrate. Thus, Meng et al. (1997) predict that a 50% reduction in both NO X and NH 3 would give about the same reduction in particulate nitrate. Not only does the chemistry generating 0 3 affect the formation of particles, but the reverse is true; i.e., the presence of particles is thought to affect ozone formation. The major reason for this is the effect of aerosol particles on light scattering and hence the actinic flux and photolysis rates. For example, scattering of light by particles in effect increases the path length through the atmosphere (see Chapters 3.C.2f and 14.C), leading to increased rates of photolysis. On the other hand, absorption of light by particles decreases the effective actinic flux. For example, Dickerson et al. (1997) used the Urban Airshed Model to examine the effects of increased light scattering by aerosol particles in the boundary layer in the northeastern United States. Using a refractive index typical of particles in that region; they predict that the amount of ozone generated increased by about 30% due to the increased rates of photolysis, particularly NO 2. The opposite effect was predicted for absorbing aerosol particles. Similarly, Jacobson (1998) predict decreased 0 3 formation (by 5-8%) in the Los Angeles area due to reduction in photolysis rates from UV-absorbing aerosol particles.

G. A T M O S P H E R I C CHEMISTRY A N D RISK ASSESSMENTS OF H A Z A R D O U S AIR P O L L U T A N T S As discussed in Chapter 2.G, risk assessment and risk m a n a g e m e n t are two separate elements of developing cost-effective control strategies for hazardous air pollutants, HAPs. The risk assessment portion ideally provides a complete understanding of the chemical and physical processes that apply to the HAP once it is emitted into the atmosphere. Thus, as we shall see in the examples that follow, some HAPs react in the atmosphere to form less toxic substances (sometimes referred to as "environmental deactivation"), whereas others form more toxic compounds ("environmental activation"). Understanding the atmospheric reactions of HAPs is clearly essential to developing appropriate health risk assessments upon which cost-effective control strategies can be based in the risk management phase.

925

In addition to the role of atmospheric reactions in the fate of airborne HAPs, atmospheric chemistry also plays a role in the formation of some of them, most notably formaldehyde and acetaldehyde. Thus, the potential formation of such compounds from the oxidation of precursors in the atmosphere must also be taken into account in their risk assessments. In the United States, the Clean Air Act Amendments of 1990 defined a list of 189 compounds or mixtures of compounds as "hazardous air pollutants," shown in Table 16.15 (Kao, 1994; Kelly et al., 1994). These include a range of chemicals, such as hydrocarbons, halogenated, oxygenated, and nitrogen- and sulfur-containing organics, pesticides, and inorganics. Although most are individual compounds, some such as "polycyclic organic matter" represent complex mixtures (see Chapter 10). A detailed review of the sources, atmospheric chemistry, and fates of HAPs is beyond the scope of this book. Instead, we concentrate here on several organic HAPs to illustrate the role of atmospheric chemistry in risk assessments. For other species, reviews and summaries are available in the literature (e.g., see California Air Resources Board (1997) for summaries of 243 species and Lin and Pehkonen (1999) for a review of the atmospheric chemistry of mercury). The California Air Resources Board has prepared risk assessments for a number of toxic airborne compounds and mixtures, designated as "toxic air contaminants," TACs (Table 16.15). For example, risk assessments for individual compounds such as benzene, benzo[a]pyrene (see Chapter 10), formaldehyde, and vinyl chloride have been carried out, in addition to complex mixtures such as diesel exhaust (California Air Resources Board, 1997a)and environmental tobacco smoke (California Environmental Protection Agency, 1997). These risk assessment documents form the basis for controls imposed as part of the risk management process (e.g., see Seiber, 1996). Kelly et al. (1994) and Kao (1994) have reviewed what is known about ambient concentrations of the HAPs and their reactions in air. Not surprisingly, given the number of compounds involved, the concentrations and atmospheric reactions and fates of many of these species are not known, precluding the development of accurate risk assessments. Also complicating such risk assessments are the wide variety of sources and their temporal and spatial variations, especially in urban areas (e.g., see Spicer et al., 1996; and Mukund et al., 1996). Some of the atmospheric fates of selected VOCs are treated in a review by Pitts (1993), in a series of papers by Grosjean (Grosjean, 1990a-c; 1991a-d) and the 1997 California Air Resources Board report. We examine here selected HAPs and related compounds,

926

16. APPLICATIONSOF ATMOSPHERIC CHEMISTRY TABLE 16.15

Compounds or Mixtures Designated as Hazardous Air Pollutants (HAPs) in the United States a and Toxic Air Contaminants (TACs) b in the State of California b'c I. Aliphatic and cyclic hydrocarbons

Methylchloroform (1,1,1-trichloroethane) Methyl iodide (iodomethane) *Methylene chloride (dichloromethane) Propylene dichloride (1,2-dichloropropane) 1,1,2,2-Tetrachloroethane 1,1,2-Trichloroethane

Saturated hydrocarbons Hexane 2,2,4-Trimethylpentane Unsaturated hydrocarbons *1,3-Butadiene Saturated halogenated hydrocarbons Bromoform *Carbon tetrachloride *Chloroform 1,2-Dibromo-3-chloropropane Ethyl chloride (chloroethane) *Ethylene dibromide (1,2-dibromoethane) *Ethylene dichloride (1,2-dichloroethane) Ethylidene dichloride (1,1-dichloroethane) Hexachloroethane Methyl bromide (bromomethane) Methyl chloride (chloromethane)

Unsaturated halogenated hydrocarbons Allyl chloride Chloroprene 1,3-Dichloropropene Hexachlorobutadiene Hexachlorocyclopentadiene *Tetrachloroethylene (perchloroethylene) *Trichloroethylene Vinyl bromide *Vinyl chloride Vinylidene chloride (1,1-dichloroethylene) II. Aromatic compounds

Aromatic hydrocarbons *Benzene Biphenyl Catechol Coke oven emissions Cumene Ethylbenzene Naphthalene Polycyclic organic matter Styrene Toluene Xylenes (isomers and mixture) o-Xylene m-Xylene p-Xylene Halogenated aromatic hydrocarbons Benzotrichloride Benzyl chloride Chlorobenzene 1,4-Dichlorobenzene

Hexachlorobenzene Polychlorinated biphenyls (aroclors) *2,3,7,8-Tetrachlorodibenzo-p-dioxin 1,2,4-Trichlorobenzene Phenolic compounds Cresols/cresylic acid (isomers and mixtures) o-Cresol m-Cresol p-Cresol Pentachlorophenol Phenol 2,4,5-Trichlorophenol 2,4,6-Trichlorophenol Phthalates Bis(2-ethylhexyl) phthalate (DEHP) Dibutyl phthalate Dimethyl phthalate Phthalic anhydride III. Nitrogenated organic compounds

Acetamide Acetonitrile 2-(Acetylamino)fluorene Acrylamide Acrylonitrile 4-Aminobiphenyl Aniline o-Anisidine Benzidine Diazomethane 3,3-Dichlorobenzidene Diethanolamine N,N-Dimethylaniline 3,3-Dimethoxybenzidine (Dimethylamino)azobenzene 3,3 '-Dimethylbenzidine Dimethylcarbamoyl chloride Dimethylformamide 1,1-Dimethylhydrazine 4,6-Dinitro-o-cresol and salts 2,4-Dinitrophenol 2,4-Dinitrotoluene 1,2-Diphenylhydrazine Ethyl carbamate (urethane) Ethylene imine (aziridine)

Ethylenethiourea Hexamethylene 1,6-diisocyanate Hexamethylphosphoramide Hydrazine Methyl hydrazine Methyl isocyanate 4,4 '-Methylenebis(2-chloroaniline) Methylene(diphenyl diisocyanate) (MDI) 4,4'-Methylenedianiline Nitrobenzene 4-Nitrobiphenyl 4-Nitrophenol 2-Nitropropane N-Nitroso-N-methylurea N-Nitrosodimethylamine N-Nitrosomorpholine Pentachloronitrobenzene (quintobenzene) p-Phenylenediamine 1,2-Propylene imine (2-methylaziridine) Quinoline 2,4-Toluene diamine 2,4-Toluene diisocyanate o-Toluidine Triethylamine (continues)

G. ATMOSPHERIC CHEMISTRY AND RISK ASSESSMENTS OF HAZARDOUS AIR POLLUTANTS TABLE 16.15

927

(continued)

IV. Oxygenated organic compounds

Glycol ethers Methyl ten-butyl ether (MTBE)

Alcohols Methanol Aldehydes *Acetaldehyde *Formaldehyde Propionaldehyde

Ketones Methyl ethyl ketone (2-butanone) Methyl isobutyl ketone (hexone) Oxides

Epichlorohydrin (1-chloro-2,3-epoxypropane) 1,2-Epoxybutane *Ethylene oxide Propylene oxide Styrene oxide

~, ~-Unsaturated carbonyls Acrolein Carboxylic acids Acrylic acid Chloroacetic acid

Esters Ethyl acrylate Methyl methacrylate Vinyl acetate

Other carbonyls and oxygenates Acetophenone Caprolactam 2-Chloroacetophenone Ethylene glycol Hydroquinone Isophorone Maleic anhydride Phosgene 1,3-Propane sultone /3-Propiolactone Quinone

Ethers Bis(chloromethyl) ether Chloromethyl methyl ether Dibenzofurans Dichloroethyl ether (bis(2-chloroethyl) ether) 1,4-Dioxane (1,4-diethylene oxide)

V. Pesticides and herbicides

Captan Carbaryl Chloramben Chlordane Chlorobenzilate 2,4-Dichlorophenoxyacetic acid Dichlorodiphenyldichloroethylene (DDE) Dichlorvos

Heptachlor Lindane (all isomers) Methoxychlor Parathion d Propoxur (baygon) Toxaphene (chlorinated camphene) Trifluralin VI. Inorganic compounds

Antimony compounds *Arsenic compounds (including arsine) *Asbestos Beryllium compounds *Cadmium compounds Calcium cyanamide Carbon disulfide Carbonyl sulfide Chlorine *Chromium compounds Cobalt compounds Cyanide compounds

Hydrochloric acid Hydrofluoric acid *Lead compounds Manganese compounds Mercury compounds Mineral fibers (fine) *Nickel compounds Phosphine Phosphorus Radionuclides (including radon) Selenium compounds Titanium tetrachloride VII. Sulfates

Diethyl sulfate

Dimethyl sulfate

a From Kao (1994) and Kelly et al. (1994). b Asterisks indicate HAPs for which the State of California has prepared detailed risk assessments and identified them as Toxic Air Contaminants (TACs). c For excellent summaries of general exposure and health effects data for 243 substances, see the California Air Resources Board report (1997b). Each summary describes the physical properties, sources, and concentrations both outdoors and indoors, atmospheric persistence, health effects, and other risk assessment information. d For a risk assessment, see "Evaluation of Ethyl Parathion as a Toxic Air Contaminant," California Department of Food and Agriculture (1988).

928

16. APPLICAT I O N S OF A T M O S P H E R I C CHEMISTRY

and the role of atmospheric chemistry in their risk assessments. Formaldehyde is a designated H A P / T A C . As we have seen on many occasions throughout this book, it not only is directly emitted by a number of sources, including by motor vehicles outdoors and by building materials indoors (see Chapter 15.D), but is also formed in air from the oxidation of both anthropogenic and biogenic organics. These secondary sources of H C H O vary geographically, seasonally, and diurnally and include the oxidation of anthropogenically emitted VOC as well as biogenic organics such as isoprene. In California, for example, oxidation of VOC is estimated to generate most of the H C H O (about 150,000 tons per year, with a large uncertainty of 50%), whereas direct emissions account for only about 10% of the total, i.e., for 18,000 tons per year (Cal EPA, 1992b; Harley and Cass, 1994). The importance of secondary formation of H C H O (and CH 3CHO) in the atmosphere is not unique to California (Kao, 1994). However, the relative amounts of direct emissions are undoubtedly much larger in urban areas where vehicles are responsible for 80% of the direct emissions. Formaldehyde reacts rapidly in air through photolysis (see Chapter 4.M) and through attack by OH, NOB, and, in coastal areas, likely chlorine atoms as well: H C H O + h v ~ H + CHO, H 2 + CO,

(18a) (18b)

H C H O + X(X = OH, NO3, C1) ~ XH + HCO. (19) Both H and HCO then react with 0 2 t o generate HO 2. Formaldehyde is therefore an example of a HAP that has both primary and secondary sources and that is relatively rapidly "deactivated" in the atmosphere through photolysis and atmospheric reactions (see Problem 3). A number of pesticides are listed as HAPs (see Table 16.16). These can be transported significant distances from their point of application, and during that time they undergo chemical transformations as well as deposition (Kurtz, 1990). Table 16.17 shows some pesticides and their transformation products in air (Seiber and Woodrow, 1995). One example of a reactive HAP that is a pesticide is 1,3-dichloropropene, used as a soil fumigant for nematodes. Both the cis and trans forms react with OH and O 3 (Tuazon et al., 1984), although the OH reaction is sufficiently fast that this reaction is expected to be the major atmospheric fate (see Problem 5). Thus, the rate constants for OH reaction with the cis and trans forms are 0.774 • 10 -11 and 1.31 • 10 -11 cm 3 molecule -1 s-1, respectively, giving lifetimes with respect to OH at

2 • 106 radicals cm -3 of 18 and 11 h. The products of the OH reactions are formyl chloride (HC(O)C1) and chloroacetaldehyde (C1CH2CHO), with one molecule of each formed for each molecule of 1,3-dichloropropene reacted (Tuazon et al., 1984)(see Problem 6). A number of the pesticide HAPs are organophosphorus compounds. Seiber and co-workers have measured some organophosphorus pesticides, including diazinon, parathion, chlorpyrifos, and methidathion, in air and fogs and wet deposition in a variety of locations in California (e.g., Glotfelty et al., 1990; Schomburg et al., 1991; Zabik and Seiber, 1993; Seiber et al., 1993; Baker et al., 1996). Both the parent pesticides, all containing reactive P = S groups, and their oxidation products, the corresponding oxon (P=O), have been identified and quantified. The ratio of the oxons to the corresponding parent compounds increased with distance from the site of release and also increased after dawn, suggesting photochemical reactions were involved in the conversion. These photochemical reactions likely involve OH chemistry. For example, Atkinson et al. (1989) reported that O H reacted with (CHBO)BP=S to form ( C H 3 0 ) 3 P = O and, similarly, that (CH30)2P(S)CH 3 formed (CH30)zP(O)CH3, in yields of 0.28 and 0.13, respectively. The reaction is believed to involve addition to the P = S bond, followed by secondary chemistry to generate the oxon (Goodman et al., 1988). Another example of this conversion of P = S found in pesticides to P = O is the oxidation of malathion in the atmosphere. Malathion itself is not a HAP and has relatively low acute mammalian toxicity because it is degraded by mammalian carboxylesterases. It is effective as a pesticide because in insects, it is activated to malaoxon, an acetylcholinesterase inhibitor. However, malathion itself typically contains impurities such as isomalathion whose mammalian toxicities are greater TABLE 16.16 Some Pesticides and Pesticide Transformation Products Listed as Hazardous Air Pollutants a 1,3-Dichloropropene 2,3,7,7-Tetrachlorodibenzodioxin 2,4,5-Trichlorophenol 2,4-D salts and esters 4,6-Dinitro-o-cresol and salts Acrolein Captan Carbaryl Chloroamben Chlordane DDE Dichlorvos Ethylene dibromide Ethylene oxide

Ethylene thiourea Heptachlor Lindane Methoxychlor Methyl bromide Parathion Pentachloronitrobenzene Pentachlorophenol Phosgene Phosphine Propoxur Toxaphene Trifluralin

a From Seiber and Woodrow (1995); see Table 16.15.

929

G. ATMOSPHERICCHEMISTRYAND RISKASSESSMENTSOF HAZARDOUSAIR POLLUTANTS TABLE 16.17

Some Pesticides and Their Transformation Products a

Pesticide

Structure

Reaction product

Merphos

(RS)3P

(RS)3PO Def

S

II

Metam sodium

CH3NCS MITC

H3CN"C~s - Na + S

II

Ethylenebis(dithiocarbamate) (EBDC)

H H2C~N

H2CmNH~C~S-

II HaC~NH~CmS-

\ C=S /

M++

HzC~N H Ethylenethiourea

H7 C~-,.N jC3H7 O2N~

H.

''J/NO2

/C3H 7 XN"

02 N ~ 7

NO2

Trifluralin CF3

CF3

S Ethyl parathion

C2HsOmP~O

I

O NO 2

C2HsOm

OC2H 5

PmO

I

NO 2

OC2H 5 Ethyl paraoxon

a Adapted from Sieber and Woodrow (1995).

than that of malathion itself (e.g., see Aldridge et al., 1979; and Ryan and Fukuto, 1985), as well as malaoxon. Malathion in air readily undergoes oxidation to malaoxon" O

~O ?CCHz~HC~, C2H50-OC2H5 atmospheric S oxidation pros

/\ /

O

O

CH 3 Malathion

\ CH 3

O?ccH2?Hc

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