Cyanide in Water and Soil: Chemistry, Risk, and Management

CYANIDE in WATER and SOIL Chemistry, Risk, and Management CYWS “L1666_C000” — 2005/10/22 — 00:42 — page ii — #2 CY...

Author: David A. Dzombak | Rajat S. Ghosh | George M. Wong-Chong

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CYANIDE in WATER and SOIL

Chemistry, Risk, and Management

CYWS “L1666_C000” — 2005/10/22 — 00:42 — page ii — #2

CYANIDE in WATER and SOIL

Chemistry, Risk, and Management David A. Dzombak Rajat S. Ghosh George M. Wong-Chong

Boca Raton London New York

A CRC title, part of the Taylor & Francis imprint, a member of the Taylor & Francis Group, the academic division of T&F Informa plc.

L1666_Discl.fm Page 1 Friday, October 21, 2005 4:20 PM

Published in 2006 by CRC Press Taylor & Francis Group 6000 Broken Sound Parkway NW, Suite 300 Boca Raton, FL 33487-2742 © 2006 by Taylor & Francis Group, LLC CRC Press is an imprint of Taylor & Francis Group No claim to original U.S. Government works Printed in the United States of America on acid-free paper 10 9 8 7 6 5 4 3 2 1 International Standard Book Number-10: 1-56670-666-1 (Hardcover) International Standard Book Number-13: 978-1-56670-666-7 (Hardcover) This book contains information obtained from authentic and highly regarded sources. Reprinted material is quoted with permission, and sources are indicated. A wide variety of references are listed. Reasonable efforts have been made to publish reliable data and information, but the author and the publisher cannot assume responsibility for the validity of all materials or for the consequences of their use. No part of this book may be reprinted, reproduced, transmitted, or utilized in any form by any electronic, mechanical, or other means, now known or hereafter invented, including photocopying, microfilming, and recording, or in any information storage or retrieval system, without written permission from the publishers. For permission to photocopy or use material electronically from this work, please access www.copyright.com (http://www.copyright.com/) or contact the Copyright Clearance Center, Inc. (CCC) 222 Rosewood Drive, Danvers, MA 01923, 978-750-8400. CCC is a not-for-profit organization that provides licenses and registration for a variety of users. For organizations that have been granted a photocopy license by the CCC, a separate system of payment has been arranged. Trademark Notice: Product or corporate names may be trademarks or registered trademarks, and are used only for identification and explanation without intent to infringe. Library of Congress Cataloging-in-Publication Data Catalog record is available from the Library of Congress

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Preface “Cyanide” is a chemical with a long and fascinating history of respectful and productive use by mankind. The fundamental cyanide species, the cyanide ion CN− , is a highly versatile and strong binder of metals in aqueous solution, a property that has been exploited in ingenious ways for commercial processes that have benefited society. The best known and largest volume uses of cyanide are in the gold mining and electroplating industries. In hydrometallurgical gold mining, aqueous solutions of CN− are used to extract and concentrate gold from ores containing very small amounts of gold. In electroplating, solutions of metal–CN species are used as the baths into which solid metals are dipped and coated with the metal from solution. The deposition of the metal from solution onto the solid metal is governed by the electrochemical gradient induced in the system, and by the metal–cyanide solution chemistry. Cyanide is also produced incidentally in significant quantities in a number of industrial processes, including coal coking and gasification, iron and steel manufacturing, aluminum manufacturing, and petroleum refining. This results in the need for control of cyanide releases in the form of gases, solids, and liquids. The substantial use of cyanide compounds in commerce coupled with the substantial incidental production of cyanide compounds means that significant amounts of cyanide are introduced into the environment on a continuous basis. Cyanide species are frequently occurring contaminants in water and soil. There are also natural sources of cyanide, such as black cherry and cassava plants. Indeed, there is a natural cycle of cyanide. However, anthropogenic inputs of cyanide to the environment are far greater in amount than natural inputs. Of course, “cyanide” is also widely known, and perhaps best known, as a potent human toxin. The most toxic form of cyanide is hydrogen cyanide, HCN, which is as toxic, and often even more so, to wildlife, especially aquatic life. There is great fear of “cyanide” in society, but some chemical forms of cyanide are nontoxic and in fact used regularly in food and cosmetic products. An example is the solid Prussian Blue, or ferric ferrocyanide, which is used as a blue pigment for use in inks, dyes, cosmetics, and other products. The chemistry of cyanide is both complex and diverse, and there are many different chemical forms of cyanide, including solid, gaseous, and aqueous species, and both inorganic and organic species. The particular chemical forms of cyanide that exist in a system, referred to as the speciation of the chemical, are all important in determining the environmental fate, transport, and toxicity of the cyanide. In our careers in environmental engineering and science, we have encountered many different problems involving cyanide in water and soil. Cyanide has been a focus in engineering and research projects that we have performed related to industrial and municipal wastewater treatment, groundwater treatment, industrial waste management, site remediation and restoration, and water quality assessment. These projects have been sponsored by a wide range of companies, industrial research organizations, and regional and federal government agencies. There is widespread interest in cyanide management for environmental and human health protection. We have learned much about cyanide use, management, emissions, and behavior in the environment in the course of these projects. Our education has been aided by useful knowledge and information acquired from many different sources and people. We undertook the preparation of this book to bring together in one place some of the current knowledge and information about cyanide release to, and behavior in, the environment, and means v

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of controlling or remediating these releases. No other broad-based examination of this topic exists. While there has been much good research and engineering development performed in the gold mining industry on cyanide management and control of environmental releases, most notably the work of Dr Terry Mudder and colleagues, this work has been focused on the industry with an orientation toward advancement of hydrometallurgical gold mining. There is much to be learned from the extensive knowledge about cyanide that has been gained in the gold mining industry, but there is a broader range of cyanide challenges in environmental engineering and science. Our book takes on this broader scope. This book tries to address the full range of issues pertaining to cyanide fate, transport, treatment, and toxicity in water and soil. We examine the sources of cyanide released to the environment, both anthropogenic and natural. We have tried to develop an appropriate balance of depth and scope of coverage. There have been compromises made on depth of coverage in some topical areas, but in all areas we have endeavored to provide good and current references to enable the reader to learn more about topics of particular interest. We developed this book to serve as a useful reference tool for engineers and scientists, including both practitioners and researchers, in academia, industrial organizations, government, and engineering and science consulting firms. We hope we have succeeded in our goal. Effective management and remediation approaches for cyanide in the environment require consideration of issues spanning many different fields. In this context, we have collaborated with a wide range of individuals possessing a wide range of expertise in our cyanide-related projects. To address the range of topics that we wanted to examine in this book, we engaged a number of our former and current collaborators to help us with the book. We are most grateful to the contributing authors, listed following this preface and in the header for each chapter. We are also grateful to Alcoa, Inc. and Niagara Mohawk Power Corporation for financial support that helped make this book project possible; and USFilter Corporation, the RETEC Group, Inc. and the Carnegie Mellon University Department of Civil and Environmental Engineering for providing assistance with preparation of graphics and the manuscript. We owe special thanks to Jacqueline Ziemianski, Donna Silverman, and Kacey Ebbitt of the RETEC Group, Inc. for their good work with preparation of graphics and securing permissions for use of copyrighted material, and to Nichole Dwyer of Carnegie Mellon University for her careful work in helping us with revising and formatting the text, with completing and formatting references, and with permissions. Finally, we thank our families for their understanding as we used many hours of family time to work on this book. David A. Dzombak Rajat S. Ghosh George M. Wong-Chong

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Editors

David Dzombak, Ph.D., P.E., DEE, is a professor in the Department of Civil and Environmental Engineering at Carnegie Mellon. Dr Dzombak’s research and professional interests include aquatic chemistry; fate and transport of chemicals in surface and subsurface waters; water and wastewater treatment; in situ and ex situ soil treatment; hazardous waste site remediation; abandoned mine drainage remediation; and river and watershed restoration. He has over 70 peer-reviewed publications and is the joint holder of three patents related to water and soil treatment. He has extensive research and consulting experience with cyanide management and treatment in soils, wastewaters, and process residuals. He has served as a member of the U.S. Environmental Protection Agency Science Advisory Board and is involved with numerous other professional service activities. Dr Dzombak received his Ph.D. in Civil-Environmental Engineering from the Massachusetts Institute of Technology in 1986. He also holds an M.S. in Civil-Environmental Engineering and a B.S. in Civil Engineering from Carnegie Mellon University, and a B.A. in Mathematics from Saint Vincent College. He is a registered Professional Engineer in Pennsylvania, and a Diplomate of the American Academy of Environmental Engineers. Dr Dzombak was elected a Fellow of the American Society of Civil Engineers in 2002. Other awards include the Professional Research Award from the Pennsylvania Water Environment Association (2002); Jack Edward McKee Medal from the Water Environment Federation (2000); Aldo Leopold Leadership Program Fellowship from the Ecological Society of America (2000); Distinguished Service Award from the Association of Environmental Engineering and Science Professors (1999); Walter L. Huber Civil Engineering Research Prize from the American Society of Civil Engineers (1997); Harrison Prescott Eddy Medal from the Water Environment Federation (1993); and National Science Foundation Presidential Young Investigator Award (1991). Rajat S. Ghosh, Ph.D., P.E., is a Program Manager with the EHS Science and Technology Group of Alcoa, Inc., the world’s largest producer of aluminum. He formerly was a Senior Technical Consultant in the Pittsburgh office of The RETEC Group, Inc., a U.S. environmental engineering and consulting company. Dr Ghosh’s research and professional interests are in geochemistry, transport and treatment of inorganic compounds (especially cyanide and heavy metals) in the subsurface; analytical method development for various inorganic and organic compounds; and subsurface multiphase flow and chemistry of organic compounds including coal tar, DNAPLs, and petroleum hydrocarbons. Dr Ghosh has extensive research and consulting experience with the electric power, natural gas, and aluminum industries in the United States in relation to cyanide management and treatment issues in soil and groundwater. In addition, Dr Ghosh serves as a senior technical reviewer for the U.S. Department of Defense basic environmental science and technology development program for site remediation under the auspices of the Strategic Environmental Research and Development Program (SERDP) and Environmental Security and Technology Certification Program (ESTCP). Dr Ghosh received his Ph.D. in Civil-Environmental Engineering from the Carnegie Mellon University in 1998. He also holds an M.S. in Chemical Engineering from University of Wyoming and a B.S. in Chemical Engineering from Jadavpur University, India. He is a registered Professional Engineer in Pennsylvania. He has over 20 professional publications in the open literature and is a joint holder of a U.S. patent on cyanide treatment technology. Dr Ghosh serves as a member of ASTM’s D-19 Committee on Water. Dr Ghosh was elected as a member of the Sigma Xi Honor Society. Other vii

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awards include the Jack Edward McKee Medal from the Water Environment Federation (2000) and the Graduate Student Award from American Chemical Society (1998). George M. Wong-Chong, Ph.D., P.E., DEE, retired director of process wastewater research at USFilter Corporation (Engineering and Construction), has over 35 years of experience in technology development, design, construction, operation, research and teaching of the management and treatment of contaminated groundwater, wastewaters, and solid hazardous waste. Dr Wong-Chong’s experience spans a range of industries including iron and steel, coal tar refining, organic chemicals, petroleum refining, munitions, aluminum manufacturing, coal gasification, live stock agriculture, and municipal wastewater. His experience in the iron and steel industry, where cyanide is a major concern, is internationally recognized; for coke plant wastewaters he developed a patented process, NITE/DENITE™, for the direct biological treatment of flushing liquor, which can contain very high concentrations of ammonia, cyanide, phenols, and thiocyanate. He also holds a patent for the physical/chemical treatment of municipal and industrial wastewaters. Dr Wong-Chong received his Ph.D. in Agricultural Engineering from Cornell University in 1974. He also holds an M.S. in Environmental Engineering from the University of Western Ontario, Canada, and a B.S. in Chemical Engineering from McGill University, Canada. He is a registered Professional Engineer in 10 states, and a Diplomate of the American Academy of Environmental Engineers. In 1999, Dr Wong-Chong received the Pennsylvania Water Environment Association Professional Research Award and the American Institute of Chemical Engineers Pittsburgh Section Award for Outstanding Professional Accomplishments in the Field of Consulting Engineering. Dr Wong-Chong has over 50 publications and presentations to his credit and remains very interested in waste water treatment technology development.

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Contributors

Todd L. Anderson, P.E. Malcolm Pirnie, Inc. Emeryville, CA

Sharon M. Drop, M.S. Alcoa, Inc. Pittsburgh, PA

Barbara D. Beck, Ph.D., DABT DABT, Gradient Corp. Cambridge, MA

David A. Dzombak, Ph.D., P.E., DEE Carnegie Mellon University Pittsburgh, PA

Brice S. Bond, M.S. Southern Illinois University Carbondale, IL

Stephen D. Ebbs, Ph.D. Southern Illinois University Carbondale, IL

Joseph L. Borowitz, Ph.D. Purdue University West Lafayette, IN

Robert W. Gensemer, Ph.D. Parametrix, Inc. Albany, OR

Joseph T. Bushey, Ph.D. Syracuse University Syracuse, NY

Rajat S. Ghosh, Ph.D., P.E. Alcoa, Inc. Pittsburgh, PA

Rick D. Cardwell, Ph.D. Parametrix, Inc. Albany, OR

Cortney J. Higgins, M.S. Carnegie Mellon University Pittsburgh, PA

Jeremy M. Clark Parametrix, Inc. Albany, OR

Gary E. Isom, Ph.D. Purdue University West Lafayette, IN

Rula A. Deeb, Ph.D. Malcolm Pirnie, Inc. Emeryville, CA

Michael C. Kavanaugh, Ph.D., PE, DEE Malcolm Pirnie, Inc. Emeryville, CA

David K. DeForest Parametrix, Inc. Bellevue, WA

Roman P. Lanno, Ph.D. Ohio State University Columbus, OH

Peter J. Drivas, Ph.D. Gradient Corp. Cambridge, MA

Richard G. Luthy, Ph.D., P.E., DEE Stanford University Stanford, CA ix

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Contributors

Johannes C.L. Meeussen, Ph.D. Energy Research Centre of the Netherlands Petten, The Netherlands Charles A. Menzie, Ph.D. Menzie-Cura and Associates Winchester, MA David V. Nakles, Ph.D., P.E. The RETEC Group Pittsburgh, PA Edward F. Neuhauser, Ph.D. Niagara Mohawk Power Co. Syracuse, NY Sujoy B. Roy, Ph.D. Tetra Tech, Inc. Lafayette, CA Mara Seeley, Ph.D., DABT DABT, Gradient Corp. Cambridge, MA Neil S. Shifrin, Ph.D. Gradient Corp. Cambridge, MA

John R. Smith, Ph.D., P.E. Alcoa, Inc. Pittsburgh, PA Angela J. Stenhouse, M.S. Parametrix, Inc. Bellevue, WA Thomas L. Theis, Ph.D., P.E., DEE Univ. of Illinois at Chicago Chicago, IL Jeanne M. VanBriesen, Ph.D. Carnegie Mellon University Pittsburgh, PA George M. Wong-Chong, Ph.D., P.E., DEE USFilter Corporation Pittsburgh, PA Thomas C. Young, Ph.D. Clarkson University Potsdam, NY Anping Zheng, Ph.D. URS Corp. Wayne, NJ Xiuying Zhao, Ph.D. Clarkson University Potsdam, NY

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Contents Chapter 1

Introduction George M. Wong-Chong, David A. Dzombak, and Rajat S. Ghosh

1

Chapter 2

Physical and Chemical Forms of Cyanide Rajat S. Ghosh, David A. Dzombak, and George M. Wong-Chong

15

Chapter 3

Natural Sources of Cyanide George M. Wong-Chong, Rajat S. Ghosh, Joseph T. Bushey, Stephen D. Ebbs, and Edward F. Neuhauser

25

Chapter 4

Manufacture and the Use of Cyanide George M. Wong-Chong, David V. Nakles, and Richard G. Luthy

41

Chapter 5

Physical–Chemical Properties and Reactivity of Cyanide in Water and Soil David A. Dzombak, Rajat S. Ghosh, and Thomas C. Young

57

Chapter 6

Biological Transformation of Cyanide in Water and Soil Stephen D. Ebbs, George M. Wong-Chong, Brice S. Bond, Joseph T. Bushey, and Edward F. Neuhauser

93

Chapter 7

Analysis of Cyanide in Water Rajat S. Ghosh, David A. Dzombak, Sharon M. Drop, and Anping Zheng

123

Chapter 8

Analysis of Cyanide in Solids and Semi-Solids David A. Dzombak, Joseph T. Bushey, Sharon M. Drop, and Rajat S. Ghosh

155

Chapter 9

Fate and Transport of Anthropogenic Cyanide in Surface Water Thomas C. Young, Xiuying Zhao, and Thomas L. Theis

171

Chapter 10

Fate and Transport of Anthropogenic Cyanide in Soil and Groundwater Rajat S. Ghosh, Johannes C.L. Meeussen, David A. Dzombak, and David V. Nakles

191

Chapter 11

Anthropogenic Cyanide in the Marine Environment David A. Dzombak, Sujoy B. Roy, Todd L. Anderson, Michael C. Kavanaugh, and Rula A. Deeb

209

Chapter 12

Cyanide Cycle in Nature Rajat S. Ghosh, Stephen D. Ebbs, Joseph T. Bushey, Edward F. Neuhauser, and George M. Wong-Chong

225

Chapter 13

Human Toxicology of Cyanide Joseph L. Borowitz, Gary E. Isom, and David V. Nakles

237

xi

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Contents

Chapter 14

Aquatic Toxicity of Cyanide Robert W. Gensemer, David K. DeForest, Angela J. Stenhouse, Cortney J. Higgins, and Rick D. Cardwell

251

Chapter 15

Toxicity of Cyanide to Aquatic-Dependent Wildlife Jeremy M. Clark, Rick D. Cardwell, and Robert W. Gensemer

285

Chapter 16

Human Health Risk Assessment of Cyanide in Water and Soil Barbara D. Beck, Mara Seeley, Rajat S. Ghosh, Peter J. Drivas, and Neil S. Shifrin

309

Chapter 17

Ecological Risk Assessment of Cyanide in Water and Soil Roman P. Lanno and Charles A. Menzie

331

Chapter 18

Regulation of Cyanide in Water and Soil David V. Nakles, David A. Dzombak, Rajat S. Ghosh, George M. Wong-Chong, and Thomas L. Theis

351

Chapter 19

Cyanide Treatment Technology: Overview George M. Wong-Chong, Rajat S. Ghosh, and David A. Dzombak

387

Chapter 20

Ambient Temperature Oxidation Technologies for Treatment of Cyanide Rajat S. Ghosh, Thomas L. Theis, John R. Smith, and George M. Wong-Chong

393

Chapter 21

Separation Technologies for Treatment of Cyanide David A. Dzombak, Rajat S. Ghosh, George M. Wong-Chong, and John R. Smith

413

Chapter 22

Thermal and High Temperature Oxidation Technologies for Treatment of Cyanide Rajat S. Ghosh, John R. Smith, and George M. Wong-Chong

439

Chapter 23

Microbiological Technologies for Treatment of Cyanide George M. Wong-Chong and Jeanne M. VanBriesen

459

Chapter 24

Cyanide Phytoremediation Stephen D. Ebbs, Joseph T. Bushey, Brice S. Bond, Rajat S. Ghosh, and David A. Dzombak

479

Chapter 25

Management of Cyanide in Municipal Wastewaters David A. Dzombak, Anping Zheng, Michael C. Kavanaugh, Todd L. Anderson, Rula A. Deeb, and George M. Wong-Chong

501

Chapter 26

Management of Cyanide in Industrial Process Wastewaters George M. Wong-Chong, David V. Nakles, and David A. Dzombak

517

Chapter 27

Cyanide Management in Groundwater and Soil Rajat S. Ghosh, David V. Nakles, David A. Dzombak, and George M. Wong-Chong

571

Index

591

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1 Introduction George M. Wong-Chong, David A. Dzombak, and Rajat S. Ghosh CONTENTS 1.1 1.2 1.3

Cyanide in History . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Cyanide Chemical Structure . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Cyanide and the Origin of Life . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.3.1 Role of Hydrogen Cyanide in the Production of Amino Acids . . . . . . . . . . . . . . . . . 1.3.2 Stanley Miller’s Experiment . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.4 Ubiquity of Cyanide Compounds in Nature . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.4.1 Cyanide in Outer Space . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.4.2 Hydrogen Cyanide in Earth’s Atmosphere. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.5 Cyanide in Industry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.6 Cyanide Releases to Water and Soil . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.7 Cyanide: Chemistry, Risk, and Management . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.8 Cyanide Regulations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.9 Cyanide Treatment Technology . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.10 Summary and Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

2 2 2 2 3 5 5 5 6 6 10 11 11 12 12

Cyanide compounds are produced and used in commerce in large quantities. In the United States, for example, approximately 200 million pounds of sodium cyanide are used annually just in heap leaching extraction of gold from ore [1], with much of this use taking place in one state, Nevada, which accounts for about 70% of U.S. gold production [2]. Large amounts of sodium cyanide are also used in electroplating [3]. Cyanide compounds are also produced incidentally in many processes, such as in aluminum and steel production, and are associated with wastewaters, solid wastes, and air emissions from these processes. In addition, cyanide compounds are present in legacy wastes disposed onsite at numerous manufactured gas plant sites in the United States and Europe. As a result, cyanide is a commonly encountered contaminant in water and soil. Because of the high degree of toxicity in certain forms of cyanide, primarily hydrogen cyanide (HCN), acceptable levels of cyanide compounds in water and soil are generally very low. For example, the U.S. drinking water maximum contaminant level for free cyanide (HCN and CN− ) is 0.2 mg/l, while the U.S. ambient water quality criterion for acute exposures in freshwater systems is 22 µg/l. As this thousandfold difference indicates, some aquatic organisms are significantly more sensitive to cyanide than are humans. Addressing problems of cyanide contamination in water and soil can be very challenging. Complicating factors include the complex chemistry and speciation of cyanide; the analytical challenges of measuring cyanide species in water and soil; the differential toxicity, reactivity, and treatability of the various cyanide species; overlapping and sometimes inconsistent regulations pertaining to cyanide; and the widespread public fear of cyanide, regardless of its form and location. Knowledge in all these areas is needed to develop effective strategies to remedy or manage cyanide 1

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Cyanide in Water and Soil

2

contamination in water and soil. This book presents current scientific understanding and engineering approaches for managing water and soil contamination with cyanide.

1.1 CYANIDE IN HISTORY Cyanide is a chemical well known to the public as a highly toxic agent [4]. For many, the word “cyanide” evokes emotions of death. This perception is prevalent in the history of cyanide dating back to antiquity, long before any understanding of the chemistry of this family of compounds was known. Traitorous Egyptian priests of Memphis and Thebes were poisoned using the pits of peaches [5]. In the 20th century, HCN gas was used in gas chambers in the World War II Holocaust, in prisons for execution of criminals with death sentences, and also as a chemical warfare agent. In 1782, the Swedish chemist Carl Wilhelm Scheele discovered a flammable, water-soluble acidic gas, later identified as HCN, when he heated the cyanide-bearing solid Prussian Blue in an aqueous sulfuric acid solution [6–8]. The name given to the evolved gas was Prussian Blue Acid, also referred to as prussic acid or blue acid [7]. This same gas caused Scheele’s death four years later [8]. The words “cyanine” and “cyanide,” derived from the Greek word “kyanos” for blue, soon came into use to describe the gas [7]. In 1811, Guy Lussac determined the composition of the gas as consisting of one molecule each of carbon, hydrogen, and nitrogen [6]. He referred to the HCN gas as hydrocyanic acid, or hydrogen cyanide.

1.2 CYANIDE CHEMICAL STRUCTURE Cyanide compounds contain the cyano-moiety, which consists of the carbon atom triply bonded to the nitrogen atom (−C≡N). The most basic, and most toxic, of these compounds is hydrogen cyanide (H−C≡N), hydrocyanic acid. HCN is a gas at ambient temperature, and is freely soluble in water. In water, HCN dissociates at high pH (pK a = 9.24 at 25◦ C) to form the cyanide anion, CN− . There are many different inorganic and organic cyanide compounds. Inorganic compounds include simple salts of cyanide with various metals such as sodium cyanide, NaCN(s), potassium cyanide, KCN(s), and more complex solids such as ferric ferrocyanide, Fe4 (Fe(CN)6 )3 (s), also known as Prussian Blue. The simple salts are highly soluble in water. The aqueous solubility of Prussian Blue and other similar complex cyanide solids are functions of pH and redox potential. There are also many organocyanide compounds, such as acetonitrile (CH3 CN), acrylonitrile (CH2 CHCN), and cyanogenic glycosides.

1.3 CYANIDE AND THE ORIGIN OF LIFE 1.3.1 ROLE OF HYDROGEN CYANIDE IN THE PRODUCTION OF AMINO ACIDS In the Precambrian or prebiotic period, about 4.6 billion years ago, primary components of the earth’s atmosphere were carbon monoxide, methane, hydrogen, nitrogen, ammonia, and water [9]. The German biologist E. Pfluger hypothesized that as the earth’s surface slowly cooled from an incandescent mass, HCN was formed by the chemical union of carbon and nitrogen, and that this compound had time to transform and polymerize to form proteins which constitute living matter [10]. Figure 1.1 and Figure 1.2 illustrate the polymerization of HCN, to form adenine, and the reaction of HCN and formaldehyde (another compound formed from the reaction of the constituents in the primitive earth’s atmosphere) to form glycine. These two compounds allow the synthesis of many amino compounds. Oro and Kimball [11,12] demonstrated the synthesis of adenine, a nucleic acid, and other purine intermediates from HCN under possible primitive earth conditions. These abiotically synthesized proteins were important stepping-stones to life, as we know it today.

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Introduction

3

N N

N

N +HCN

+HCN

N Hydrogen cyanide

N (HCN)3

(HCN)2 N N +HCN

N

N

Diaminomaleonitrile (HCN)4

N N N +HCN

N

N

Adenine (HCN)5 Hydrogen

Carbon

N

Nitrogen

FIGURE 1.1 Polymerization of hydrogen cyanide to form adenine. (Source: Barbieri, M., The Organic Codes: An Introduction of Semantic Biology, Cambridge University Press, Cambridge, MA, 2002. With permission.)

1.3.2 STANLEY MILLER’S EXPERIMENT In 1953, Stanley Miller demonstrated that HCN and certain organic compounds, including aldehydes and amino acids, can be formed from the constituents of the prebiotic earth atmosphere, that is, methane, ammonia, hydrogen, and water [9]. The experiments, which earned Miller a Nobel Prize, were performed in a spark-discharge reaction apparatus as shown in Figure 1.3. The apparatus, which was claimed to be a crude model of the primitive earth’s atmosphere, was charged with water and air was evacuated; then, a mixture of ammonia, methane and hydrogen was added. The water in the small flask was boiled to initiate a circulation of gases and water vapor into the reaction flask, in which an electric spark was generated. The spark initiated the reaction of the ammonia, hydrogen,

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4

+

N

+

N

N N

Formaldehyde

Hydrogen cyanide

Ammonia

Ammonitrile

N N

Ammonitrile

N

+

Water Hydrogen

Carbon

Water

N

+

Ammonia N

+

Glycine Oxygen

Nitrogen

FIGURE 1.2 Reaction of hydrogen cyanide and formaldehyde to form glycine. (Source: Barbieri, M., The Organic Codes: An Introduction of Semantic Biology, Cambridge University Press, Cambridge, MA, 2002. With permission.)

Electrodes

Gases

To a vacuum pump

spark discharge

Water out Condenser Water in Water droplets

Boiling water

Liquid water in trap conaining organic compounds

FIGURE 1.3 Apparatus for experiment by Stanley Miller that demonstrated formation of hydrogen cyanide from constituents of the prebiotic Earth atmosphere. (Source: Miller, S.L. and Orgel, L.E., The Origins of Life on Earth, Prentice-Hall, Englewood Cliffs, NJ, 1974. With permission.)

methane, and water to form HCN and aldehydes. A typical experiment entailed operating of the spark continuously for about 1 week with regular analysis of samples from the system. Figure 1.4 shows the reaction profiles for ammonia (charged material), and amino acids, HCN and aldehydes (reaction products). Miller’s data clearly demonstrated a mechanism for abiotic production of HCN in the atmosphere, one that also exists today during electrical discharges associated with thunderstorms [9].

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Introduction

5

8

NH3(×10)

Molar concentration

7 Amino acids(×103)

6 5

HCN (×102)

4 3 2

Aldehydes (×103)

1

25

50

75 100 Time (h)

125

150

FIGURE 1.4 Reactant and product concentrations in the experiment by Stanley Miller. (Source: Miller, S.L. and Orgel, L.E., The Origins of Life on the Earth, Prentice-Hall, Englewood Cliffs, NJ, 1974. With permission.)

The findings of Miller are further substantiated by the discovery of the presence of CO, HCN, OH• , formaldehyde and methanol in outer space [13].

1.4 UBIQUITY OF CYANIDE COMPOUNDS IN NATURE Cyanide compounds occur commonly in nature. HCN is present in outer space, in the earth’s atmosphere, in plants, animals, microbes, and fungi. Cyanide can be produced by certain plants, bacteria, fungi, and algae. Chapter 3, which examines natural sources of cyanide, discusses in detail the occurrence, role, and environmental impact of cyanide in plants, animals, microbes, and fungi. The natural cycle of cyanide in the environment is the focus of Chapter 12.

1.4.1 CYANIDE IN OUTER SPACE Hydrogen cyanide has been detected at a number of locations in outer space. For example, it is a trace constituent in the nitrogenous atmosphere of Titan, the largest moon of Saturn [14], and in the coma of the Hale–Bopp comet [15]. Polymerization products of HCN are the dominant components of dust grains sampled from the tail of Comet 81P/Wild2 in 2004 [16]. This presence of HCN in space is now being used to study the birth of massive stars [17]. The detection of large amounts of HCN toward the center of a protostar (an evolving star) means that it has already started to warm up; from this information it is possible to determine the degree of evolution and the age of the star [17].

1.4.2 HYDROGEN CYANIDE IN EARTH’S ATMOSPHERE Hydrogen cyanide is detectable in the troposphere and stratosphere of the earth. Its concentration in the nonurban troposphere of the northern hemisphere has been reported as approximately 160 pptv [18]. In the tropical upper troposphere, a range of HCN concentrations from 200 to 900 pptv

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has been reported [19]. From field measurements and modeling it has been established that biomass burning is a major global source of HCN emissions [19,20]. Estimates of the total release of HCN to the atmosphere from biomass burning range from 1.4 to 2.9 × 1012 g (as N) per year [19]. The residence time of HCN in the atmosphere is approximately two to four months [19]. The oceans of the world provide a sink for the atmospheric releases of HCN and other compounds from biomass burning [19], as discussed in Chapter 11.

1.5 CYANIDE IN INDUSTRY Substantial quantities of cyanide compounds are used and produced in commerce (Chapter 4). Today most cyanide compounds are manufactured starting with HCN, which is synthesized by the platinumcatalyzed reaction of ammonia and methane [3]. HCN is a basic chemical feed stock used in the manufacture of sodium cyanide for gold mining and electroplating; adiponitrile for nylon; methyl methacrylate for clear plastic; triazines for agricultural herbicides; methionine for animal food supplement; and chelating agents (e.g., nitrilotriacetate) for water and wastewater treatment [3]. Worldwide annual production and manufacturing capacity of HCN in 1992 were estimated to be 0.95 million tons and 1.32 million tons, respectively [3]. A 2001 estimate of worldwide cyanide production was 2.60 million tons [7]. In 2001, 0.75 million tons of HCN were produced in the U.S. (Table 1.1). A significant fraction, estimated to range from 8 to 20%, of HCN is used to produce sodium cyanide [3,21,22], much of which is used in hydrometallurgical gold mining. The production and use of cyanide is growing, as indicated by the chronological tabulation of HCN production in the United States in Table 1.1. In addition to use of cyanide compounds in gold mining, electroplating, and chemical production, cyanide compounds are also used in some applications that involve direct distribution to the environment. Sodium ferrocyanide, Na4 (Fe(CN)6 ) and ferric ferrocyanide, Fe4 (Fe(CN)6 )3 (s) are used as an anticaking agent in road salt [23]. It is the presence of ferric ferrocyanide that gives a blue color to salt in which it is used. These compounds can dissolve in water after placement on road surfaces. Sodium ferrocyanide is also used in some forest fire retardants [24].

1.6 CYANIDE RELEASES TO WATER AND SOIL Most cyanide that occurs in water and soil is anthropogenic, derived from industrial processes, but there are natural sources of cyanide as noted above. The combination of widespread industrial sources and natural sources leads to detectable concentrations of cyanide in many natural waters, though concentrations are usually low. In a 1981 evaluation of monitoring data in the USEPA STORET database, it was determined that the mean concentration of total cyanide in surface waters of the United States did not exceed 3.5 µg/l, but in 37 of 50 states there were sampling locations where total cyanide concentrations in excess of this level were reported [25]. Sample results from a number of industrialized areas had total cyanide concentrations greater than 200 µg/l. Total cyanide concentrations in U.S. drinking water intake supplies are usually very low (1000 3120 3000 2000 100–3120 2500 1059–1807 882 1924 1394–2317 1100 1134 1237 118–236 234 2513 0–181 39–113

Reference [50] [51] [52] [52] [53] [51] [52] [52] [52]

[54] [52]

[50] [50] [52] [52] [52]

[52] [52] [52] [52] [55] [55]

defense. For example, the larvae of the southern armyworm (Spodoptera eridania Cramer) showed a preference for the foliage of the cyanogenic lima bean [5]; army ant larvae feed on the cyanogenic grass (Cynodon plectostachyus) [8] and fungal leaf blight on the cyanogenic rubber tree (Hevea brasiliensis) [9]. Another example involves the cyanogenic plant birdsfoot trefoil (Lotus corniculatus), larvae of the five spotted burned moth (Zygaena trifolii), snails (Helix aspersa), and an ichneumonid wasp (Apantales zygaebarun) [10]. Production of cyanide by birdsfoot trefoil deters herbivory by the snail species, with the snail feeding only on tissues with low cyanide concentrations. In contrast, the moth larvae preferentially consume highly cyanogenic material and sequester the cyanide, to be used as their own defense compound. This deters predation, except by A. zygaebarum, which can detoxify the cyanide present in the larvae internally.

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Natural Sources of Cyanide

CH2OH

29

CH2OH

CN

O O

C CH3

HO

+

CN

O OH

CH3 H2O

b-glucosidase

OH

+ HO

HO

OH

C

CH3

CH3 Acetone cyanohydrin

OH

OH b-D-glucopyranose

Linamarin

CN hydroxynitrile HO

C

CH3

CH3

HCN

+

O

lyase

Acetone cyanohydrin

C

CH3

CH3 Hydrogen cyanide

Acetone

FIGURE 3.1 Decomposition of linamarin by plant enzymes. (Source: Adapted from Conn, E.E., Toxicants Occurring Naturally in Foods, National Academy of Science, Washington, DC, Chapter 14, 1973. With permission.)

The impacts of cyanogenesis in plants are discussed in relation to the cultivation of cassava and forage crops, sorghum, and sudan grass.

3.1.2.1 Impact of Cyanogenesis in Cassava (Manihot esculenta) The cassava plant is a staple food crop for over 500 million people in Asia, Africa, South America, and the Caribbean [11]. It is also the primary source from which plant starches, such as tapioca and farina, are derived. From 1985 to 1996, the world production of cassava increased from 134 to 164 million tons per year, an increase of about 2.0% per year [12]. With projected increases in world population, cassava production and usage will likely continue to increase. The HCN content of cassava can reach levels toxic to humans (Table 3.2). However, the HCN content of the bitter species of cassava protects the plant from pests and disease, making it more desirable for cultivation [13]. Consumption of bitter cassava by humans and livestock requires thorough processing to produce safe food products. Safe processing involves the following steps [13]: • Peeling of the root tubers; for animal food the peeling step is omitted • Milling of the peeled tubers (which initiates the HCN release) • Heating and drying to remove the HCN by volatilization In the traditional processing for human consumption, the milled tubers may be washed with water, or sun-dried to remove HCN, or fermented [13]. Traditional processing methods allow the volatilization and release of HCN to the atmosphere, a process that takes days to complete. It is conceivable that as processing moves towards more rapid methods (e.g., extraction in water), the extraction water may pose environmental problems. Similarly, the volume/quantity of peelings, if not reused as animal feed after processing, may also pose disposal problems. Human ingestion of improperly processed bitter cassava has severe consequences, and can result in epidemic spastic paraparesis, more commonly referred to as Konzo disease. The clinical features of the disease are characterized by an abrupt onset of a permanent, symmetrical, but nonprogressive

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TABLE 3.3 Examples of Plant Cyanogenic Glycosides, their Sources, and the Products from their Hydrolysis Glycoside

Plant sources

Amygdalin

Vicianin Dhurrin

Members of the Rosaceae, including almond, apple, apricot, cherry, peach, pear, plum, and quince Members of the Rosaceae, including cherry laurel; Eucalyptus cladocalyx; Linaria striata Dc. Sambucus niagra L. (elderberry), Acacia sp. (Australian acacias) Vicia sp. (common vetch) Sorghum sp. (sorghums, Kaffir corns)

Taxiphyllin

Taxus sp.

Linamarin

Lotaustralin Acacipetalin

Phaseolus lunatis L. (lima bean, many varieties); Linum usitatis-simum L. (linen flax); Manihot sp. (cassava or manioc); Trifolium repens L. (white clover); Lotus sp. (trefoils); Dimorphotheca sp. Occurs with linamarin Acacia sp. (South African acacias)

Triglochinin

Triglochin martimum L. (arrow grass)

Prunasin

Sambunigrin

Hydrolysis products Gentiobiose + HCN + benzaldehyde

D-Glucose + HCN + benzaldehyde

D-Glucose + HCN + benzaldehyde Vicianose + HCN + benzaldehyde D-Glucose + HCN + p-hydroxybenzaldehyde D-Glucose + HCN + p-hydroxybenzaldehyde D-Glucose + HCN + acetone

D-Glucose + HCN + 2-butanone D-Glucose + deimethylketone cyanohydrin D-Glucose + HCN + triglochinic acid

Sources: Data from Conn, E.E., Toxicants Occurring Naturally in Foods, National Academy of Science, Washington, DC, 1973, Chapter 14; Howe, R.H., Proceedings of the Conference on Cyanide and the Environment, Tucson, AZ, 1984, p. 331.

paralysis of both legs in previously healthy persons, and in severe cases, damage to arms and cranial nerves also may occur [14]. These outbreaks of Konzo epidemics have occurred in several African nations, including Mozambique, which experienced over 1000 cases in 1981 in Nampula province, 171 cases in 1988 in Namapa and Erati provinces, and large outbreaks in Mogincual provinces. In 1992 and 1993, outbreaks were also observed in the savanna zone of Bandundu region in Zaire, where 78 cases of Konzo were found in a population of 1936 inhabitants [15]. Mammals, including humans, can ingest sublethal quantities of HCN without adverse effects, as the HCN is enzymatically detoxified to thiocyanate and excreted in the urine [16,17]. Konzo symptoms arise with the rapid ingestion of large quantities of HCN in improperly processed cassava; ingestion rate exceeds the rate at which the body can safely process the ingested HCN. 3.1.2.2 Impact of Cyanogenesis in Forage Sorghum, sudan grasses, arrow grass, velvet grass, white clover, and alfalfa are all important forage crops used in livestock agriculture. These plants are all generally cyanogenic in nature, and can potentially be a hazard for grazing livestock. However, different animal species react differently after

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(a)

(b)

FIGURE 3.2 Biosynthetic pathway for cyanogenic glycoside. (a) Generalized biosynthesis of cyanogenic glycosides. (Source: Lechtenberg, M. and Nahrstedt, A., Cyanogenic Glycosides, John Wiley & Sons, 1999. Reproduced with permission.) (b) Biosynthetic pathway for dhurrin from L-tyrosine. (Source: Halkier, B.A. et al. Cyanide Compounds in Biology, John Wiley & Sons, 1988. Reproduced with permission.)

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ingesting plant material containing cyanogenic glycosides (whole intact plant material) or HCN (plant material which has been stressed, chopped, etc.). These differences are mainly due to differences in anatomical structure and different detoxifying abilities in different animals. Monogastric animals (e.g., horses and swine) have strong hydrochloric acid in their stomachs; this hydrochloric acid tends to react with HCN, liberated by the degradation of cyanogenic glycosides, to yield formic acid and ammonium chloride, substances that are essentially nontoxic. On the other hand, multistomached ruminant animals (e.g., cattle, sheep, and goats) have stomachs operating in the neutral pH range, potentially ideal for absorption of HCN released from the breakdown of cyanoglycosides. It must be recognized that toxic episodes will depend on the rates of ingestion, the quantity of cyanoglycosides or HCN ingested, and the capacity of the animal to oxidize the HCN to thiocyanate. If the rates of ingestion and cyanide concentration are greater than the detoxification rate, then a fatal event will occur [18]. It is estimated that a 450-kg cow should be able to detoxify HCN at a rate of 0.5 g h−1 , and rapid ingestion of 1 g would be fatal within 15 to 20 min [19]. One gram of HCN can be contained in about 2 kg of forage, with a cyanide concentration of 500 mg kg−1 . The cyanide content of forage depends on the following factors [19]: • Species and variety of forage crops: Sorghum generally has much higher cyanogenic content than sudan grass, and the sorghum–sudan grass hybrids contain higher cyanogenic material than sudan grass. • Stage of crop growth: Young growing tips, or shoots or leaves, tend to have the highest cyanogenic contents; mature plant material (e.g., leaves, stalks, etc.) tends to have the lowest content. Sudan grass should not be grazed or green chopped until it reaches a height of at least 45 to 50 cm, and sorghum–sudan grass hybrids should not be grazed or green chopped until they have reached at least 60 to 75 cm in height. Sorghum is generally unsafe for pasture or green chopping until full maturity. Greater details on forage management can be obtained in bulletins from any of the U.S. state universities that provide agricultural extension service.

3.1.3 ETHYLENE PRODUCTION IN PLANTS In addition to the production of cyanogenic glycosides, plants that synthesize ethylene via the enzyme 1-aminocyclopropane-1-carboxylic acid (ACC) oxidase also produce HCN as a by-product [20]. Ethylene is produced by almost all parts of higher plants, and is essential for seed germination, root and shoot growth, flower development, senescence and abscission of flowers and leaves, and ripening of fruit. Ethylene is also produced during periods of stress, including drought, flooding, chilling, exposure to ozone, and mechanical wounding. Figure 3.3 depicts the methionine cycle and ethylene biosynthesis. As shown in this figure, ethylene in higher plants is synthesized from S-adenosyl-L-methionine (SAM) via the intermediate ACC. The last step in the pathway, involving the conversion of ACC to ethylene, requires oxygen, and is catalyzed by the enzyme ACC oxidase. HCN is produced as a final by-product of the ACC oxidase reaction in a stoichiometric 1:1 ratio with ethylene.

3.2 CYANIDE IN MICROORGANISMS Cyanide is produced by, and contained in, various microorganisms. Cyanogenic microorganisms have been identified in the general classes of algae, bacteria, and fungi.

3.2.1 CYANOGENIC ALGAE In a review by Vennesland et al. [21], three species of algae are identified as having the capability to produce HCN. These are Chorella vulgaris and cyanobacteria (blue-green algae) Anacystis nidulans,

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Methi (Yang) cycleonine SAM synthetase

MTA nucleosidase

ACC synthetase

O H N

H2C

C

CoASH CH2

Malonyl-CoA

C

C COO–

H2C

+ NH3

H2C

COO– ACC N-malonyltransferase

N-malonyl-ACC

H2C

COO–

1-Aminocyclopropane1-carboxylic acid(ACC)

ACC oxidase H

H C

C

H ½ O2

CO2 + HCN + H2O

H Ethylene

FIGURE 3.3 Production of hydrogen cyanide during ethylene synthesis. (Source: Adapted from Buchanan, B.B., Biochemistry and Molecular Biology of Plants, American Society of Plant Biologists, Rockville, MD, 2002.)

and Nostroc muscorum. These organisms produce HCN from an aromatic amino acid preursor (e.g., histidine, tryptophane, phenylalanine, and tyrosine) via an amino acid oxidase–peroxidase enzyme system in an oxic and illuminated environment. Histidine produces the greatest amount of HCN. The reaction occurs in a sequence of stages, as shown in Figure 3.4. In the first stage, oxidation of histidine is catalyzed by amino acid oxidase to an imino acid; oxygen is reduced to hydrogen peroxide. In the second stage, imino acid oxidation occurs, and is catalyzed by peroxidase to yield HCN and imidazole aldehyde [21]. It must be noted that in the absence of peroxidase, the imino acid is rapidly hydrolyzed, nonenzymatically, to keto acid and ammonia [21].

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Histidine

FIGURE 3.4 Production of hydrogen cyanide from histidine in amino acid oxidase–peroxidase system. (Source: Reprinted from Cyanide in Biology, Vennesland, B. et al., 349, 1981. With permission from Elsevier.)

2H+ H2N - CH2 - COOH

HN

2H+ CH - COOH

N

≡ C - COOH

HCN + CO2

FIGURE 3.5 Production of hydrogen cyanide from glycine. (Source: Adapted from Knowles, C.J., Bacteriol. Rev., 40, 652, 1976.)

3.2.2 CYANOGENIC BACTERIA Cyanogenesis in bacteria is limited to Chromobacterium violaceum and certain pseudomonad species (i.e., P. chlorophis, P. aureofaciens, P. aeruginosa, and P. fluorescens) [22–32]. These organisms require glycine for the production of HCN, which occurs only in the transition stage of growth, from log phase to stationary phase, under aerobic conditions [22,23,25,26,28,30]. A two-step oxidative reaction model involving two flavoproteins has been proposed for the production of HCN from glycine. This reaction sequence is shown in Figure 3.5, which entails the oxidation of glycine to iminoacetic acid and subsequent oxidation to HCN and carbon dioxide. The HCN in this figure originates from the methylene group, and carbon dioxide from the carboxyl group of glycine [26].

3.2.3 CYANOGENIC FUNGI Formation of HCN was first observed in the fungus Marasmius oreacles in 1871 [25]. Numerous species of fungi have now been identified as cyanogenic. These organisms include species of the genera Actinomycetes, Basidiomycetes, Clitocybe, Marasmius, Pholiota, Polyporus, and Tricholoma [25]. The production of cyanide by fungi varies, depending on the growth stage and species

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TABLE 3.4 Listing of Plant Pathogenic Fungi Organism Basidiomycetes species Marasmium oreacles Stemphylium loti

Disease Snow mold disease Fairy ring disease Copperspot disease

Plants affected Alfalfa and other forage plants Grasses (lawns and pasture) Birdsfoot trefoil

Sources: Data from Castric, P.A., Castric, K.F., and Meganathan, R., Cyanide in Biology, Vennesland, B., Conn, E.E., Knowles, C.J., Westley, J., and Wissing, F., Eds., Academic Press, London, 1981, p. 236; and from Knowles, C.J., Bacteriol. Rev., 40, 652, 1976.

of the organism. In Pholiota aurea, HCN was formed in young, fresh, fruiting bodies; old fruiting bodies and old mycelia produce HCN only when damaged or stressed [33]. In snow mold basidiomycetes grown in liquid culture, HCN was produced during active growth [34]; in a B-type isolate of the same snow mold basidiomycetes, HCN was produced only during the autolytic growth phase [35]. Fungi are also capable of producing cyanogenic compounds that in turn produce free cyanide. The B-type isolate of the snow mold basidiomycetes produced an unstable cyanogenic compound in shake cultures, with increased rate of production observed during active growth. The cyanogenic compound released free cyanide by a nonenzymatic reaction [35], and was identified predominantly as a glyoxylic acid, cyanohydrin [35,36]. These snow mold fungi also produced a β-glucosidase enzyme [36], which could be responsible for HCN release from plant cyanogenic glycosides. (e.g., linamarin, lotasistrain, and amygaldin). Many cyanogenic fungi are plant pathogens. Table 3.4 presents a listing of these pathogenic fungi, the disease caused by each of them, and the plants affected by these pathogens. It is generally believed that the pathogenic effects are partly because of the ability of the organism to release HCN and partly because of the release of HCN from the host plant’s cyanogenic glycosides, when acted upon by the fungal β-glucosidase enzyme. This may be especially true in the case of the snow mold basidiomycetes [25].

3.2.4 ROLE OF CYANOGENIC COMPOUNDS IN MICROORGANISMS The role of cyanogenic compounds in microorganisms is unclear, except in the case of fungi. It is possible to conclude that in algae and bacteria, the cyanogenesis is simply an end result of the metabolic process under opportunistic circumstances, because of the strict environmental and substrate specificity required for HCN production. In addition, the small quantities of HCN produced, measured as nano-moles in reported tests [22,24,27], likely will not be environmentally significant, and may not be an effective antibiotic deterrent in the natural environment. For example, the bacterial species Pseudomonas aeroginosa has the capacity for cyanogenesis, but also has the capacity to degrade cyanide compounds [37]. Thus, there may not be a net release of HCN in the natural environment. Fungi, on the other hand, are recognized plant pathogens, and their use of HCN, either self-excreted or enzymatically induced from the host plant, is a mechanism to provide growth nutrients. In the field of plant pathology, it is generally accepted that certain microorganisms provide some benefit to “host” plants. Species of pseudomonas (including species known to be cyanogenic) have been recognized to be involved in the suppression of plant root pathogens [38]. This control on pathogens may involve the bacterial secretion of secondary metabolites, including antibiotics, siderophores, and HCN. Cyanogenic processes in microorganisms thus have a range of effects, but generally do not seem to pose any threats to ecosystems and the natural environment.

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3.3 CYANIDE IN ANIMALS 3.3.1 CYANOGENIC ANIMALS The production of HCN by animals is restricted almost exclusively to one phylum, arthropods; even within this phylum, composed of 11 extant classes of animals, the phenomenon is restricted to certain members of Chilopoda (centipedes), Diplopoda (millipedes), and Insecta (insects) [39]. Based on the literature, the following arthropods have been identified as being cyanogenic: • Chilopods (centipedes): 7 out of the 3,000 species which make up the class of chilopoda [39] • Diplopods (millipedes): 46 out of the 7,500 species which make up the class of diplopoda [39] • Insecta (insects): 68 out of the 750,000 species which make up the class of insecta [40] Specific class, families, and species name listings for these animals can be found in the reviews by Duffey [39] and Davis and Nahrstedt [40].

3.3.2 BIOSYNTHESIS OF CYANIDE There has been limited study of the biosynthesis of cyanide in arthropods [39], and the studies performed indicate that to a large degree, cyanide biosynthesis processes in two species of millipedes mirror the processes known to occur in plants [39], where a biogenetic precursor (e.g., L-phenylalanine) is processed to form a cyanogenic glycoside (e.g., mandelonitrile), as illustrated in Figure 3.6. Figure 3.6 shows the cyanogenic release of HCN and an aldehyde (e.g., benzaldehyde). Biosynthesis of cyanogenic glycosides in insects is slightly different from biosynthesis in plants, in that these animals appear to have the capacity to sequester the cyanogenic compounds from their food source, and also synthesize these compounds from amino acids such as valine and isoleucine, the precursors to linamarin and lotaustralin [40]. Also, some insects have the biosynthetic ability to assimilate ingested cyanogenic glycosides into other cyanogenic compounds, as shown in Figure 3.7.

Phenylacetonitrile

FIGURE 3.6 Biosynthetic pathway for hydrogen cyanide and benzaldehyde in the polydesmoid millipede. (Source: Duffey, S.S., Cyanide in Biology, Vennesland, B., Conn, E.E., Knowles, C.J., Westerley, J., and Wissing, F., Eds., Academic Press, London, 1981, p. 385. With permission.)

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FIGURE 3.7 Transformation of plant cyanoglucoside prunasin to (R)-mandelonitrile by the beetle Paropsis atmoria. (Source: Nahrstedt, A. et al., Cyanide Compounds in Biology, Evered, D. and Harnett, S., Eds., John Wiley & Sons, 1988. Reproduced with permission.)

3.3.3 ROLE OF CYANOGENIC COMPOUNDS IN ARTHROPODS The role of cyanogenesis in arthropods is generally accepted as a defensive or antipredator device [21,39,40]. These defensive chemicals are actively secreted when a predator attacks [21]. However, as in plants, there are apparent exceptions to this defensive function. Certain beetle larvae are voracious predators of both cyanogenic and benzoquinone-producing millipedes, and starlings and toads consistently eat millipedes [39]. These phenomena reflect opportunistic traits of nature. Another example is the opportunistic trait of the larvae of some insects to sequester the cyanogenic materials from plants and in turn, use them for their own defense. Further detail on this behavior is provided in Chapter 12.

3.4 CYANIDE IN FOREST FIRES Annually, there are over 42,000 forest fires (wildfires) in the United States [41]. In the summer of 2000, over 7 million acres of land was burned in the western United States. Similar events occur worldwide [42]. These fires occur as a result of natural events (i.e., electrical storms) and man’s activities (e.g., slash and burn, and accidents). Cyanide is produced in forest fires, and is released to the atmosphere [43]. Emissions of cyanide may be exacerbated in some forest fires, where cyanideamended (as an anticorrosion agent) fire retardant is used to assist in the fighting of the fire [44,45]. Usage of these cyanide-treated fire retardants can result in adverse environmental consequences. A forest fire is an uncontrolled thermal event where organic matter is thermally consumed, at times in the presence of abundant oxygen, and at times with limited oxygen. The organic matter naturally contains some nitrogen along with abundant carbon, and the circulating air contains abundant nitrogen. In those regions of a fire where oxygen is limited, there will be opportunity for HCN to form. In fact, this type of thermal process, where organic matter is fired in a furnace, was one of the early processes for producing cyanide salts [46]. Laboratory simulation of the smoldering remains of a fire, which can persist for weeks after the passage of the flame front and underground fires, has demonstrated the production of HCN methane, ethane, ethylene, acetone, acetonitrile, acetylene, propene, formaldehyde, methanol, acetic acid, formic acid, glycolalehyde, phenol, furan, and ammonia [47,48]. Forest fires release significant quantities of cyanide to the atmosphere. Estimates of the amount of HCN and methyl cyanide released to the atmosphere from biomass burning are in the range of

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(1.4 to 2.9) × 1012 g (as N) per year and about 0.5 × 1012 g (as N) per year, respectively [43,49]. It is believed that the oceans of the world provide the primary receiving sinks for this cyanide and in the absence of any reports of adverse effects, one might assume that these ocean sinks [43] are capable of handling this quantity of cyanide at the rates currently produced. At this time it is difficult to estimate the quantity of cyanide released from fire retardant materials used in fighting forest fires. However, given the large number of outbreaks each year, these releases may be significant.

3.5 SUMMARY AND CONCLUSIONS • HCN and cyanogenic compounds are constituents of the earth’s biosphere. Cyanide contains carbon and nitrogen, key components in the evolution of life on this planet. • Cyanogenesis occurs in plants, microorganisms, and animals. It is believed that all plants produce low levels of cyanogenic compounds, and some 2650 species produce significant levels of free cyanide. Many cyanide-producting plant species bear social and economic importance in our daily life, in areas such as food, forage, and horticulture. • In the animal world, arthropods (centipedes, millipedes, moths, and butterflies) use cyanogenesis as a defense device against predators (e.g., venom in the centipedes and millipedes). • The role of cyanogenesis in algae and bacteria in the natural environment is varied and uncertain. Fungi are plant pathogens, and their use of HCN, either self-excreted or enzymatically induced from the host plant, is a mechanism to provide growth nutrients. • Plant pathogenic fungi use HCN, either self-produced or induced from the host plant, to infest pasture and lawn grasses and other forage plants, as a means of securing growth nutrients. Pathogenic fungi cause plant diseases such as snow mold, fairy ring, and copper spot diseases. • The level of naturally occurring cyanide from the plant, microbial, and animal kingdoms does not appear to pose a major environmental threat to ecosystems. However, with ever increasing demands for food, and increased cultivation of cyanide-producing plants, there could be environmental implications in the processing of food plants like cassava. • Forest fires release significant quantities of cyanide to the atmosphere. The oceans of the world appear to provide the primary receiving sinks for this cyanide.

REFERENCES 1. Oparin, A.I., The Origin of Life, Dover Publications, Mineola, NY, 1953. 2. Allen, J. and Strobel, G.A., The assimilation of H14 CN by a variety of fungi, Can. J. Microbiol., 12, 414, 1966. 3. Thatcher, R.C. and Weaver, T.L., Carbon–nitrogen cycling though microbial formamide metabolism, Science, 192, 1234, 1976. 4. Lechtenberg, M. and Nahrstedt, A., Naturally occurring glycosides, in Cyanogenic Glycosides, Ikan, R., Ed., John Wiley & Sons, Chichester, U.K., 1999, Chapter 5. 5. Jones, D.A., Cyanogenesis in animal–plant interactions, in Cyanide Compounds in Biology, Evered, D. and Garnety, S.F., Eds., John Wiley & Sons, 1988, p. 151. 6. Bolhuis, G.G., The toxicity of the cassava root, Netherlands J. Agr. Sci., 2, 176, 1954. 7. Seigler, D.S., Cyanogenic glycosides and lipids: structural types and distribution, in Cyanide in Biology, Vennesland, B., Conn, E.E., Knowles, C.J., Westley, J., and Wissing, F., Eds., Academic Press, London, 1981, p. 133. 8. Georgiadis, N. and McNaughton, S., Interactions between grazers and a cyanogenic grass, Cynodon plectostachyus, Oikos, 51, 343, 1988.

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9. Lieberei, R., Relationship of cyanogenic capacity (HCN-c) of the rubber tree (Hevea brailiensis) to the susceptibility to Microcyclus ulei, the agent causing South American leaf blight, J. Phytopath. (Berl), 122, 54, 1988. 10. Jones, D.A., Selective eating of the acyanogenic form of the plant Lotus corniculatus by various animals, Nature, 193, 1109, 1962. 11. Padmaja, G., The culprit in cassava toxicity: cyanogens or low protein? in Consultative Group on International Agricultural Research News Letter, Volume 3, www.worldbank.org/html/cgiar/newsletter/Oct96/ 6cgnews.html, 1996. 12. Gottret, M.V., Cassava, Section 2.1.B, Annual Report 97, Centro International de Agricultura Tropical (CIAT), Cali, Colombia, www.ciat.org/impact/iannual97/lanual97c.htm, 1997. 13. Cereda, M., Processing of cassaca roots in Brazil, Proceedings of the International Workshop on Cassava Safety, ISHS Acta Horticultura, 375, 21, 1994. 14. Howlett, W.P., Konzo: a new human disease entity, Proceedings of the International Workshop on Cassava Safety, ISHS Acta Horticulturae, 375, 32, 1994. 15. Cliff, J.L., Cassava safety in times of war and drought in Mozambique, Proceedings of the International Workshop on Cassava Safety, ISHS Acta Horticulture, 375, 37, 1994. 16. Jones, D.A., Why are so many food plants cyanogenic? Phytochemistry, 47, 155, 1998. 17. Westley, J., Mammalian cyanide detoxification with sulphane sulfur, in Cyanide Compunds in Biology, Evered, D. and Garnety, S.F., Eds., John Wiley & Sons, Chichester, U.K., 1988, p. 201. 18. USAT, Cherry tree leaves killed foals, scientists conclude, USA Today, July 13, 2001. 19. Vough, L.R. and Cassel, E.K., Prussic acid poisoning of livestock, University of Maryland, College of Agriculture, Maryland Cooperative Extension Bulletin, www.inform.um.edu/prussic_acid_poisoning_ of_ livestock.htm, 1987. 20. Yip, W.K. and Yang, S.F., Ethylene biosynthesis in relation to cyanide metabolism, Botanical Bull. Acad. Sin. Taipei, 39, 1, 1998. 21. Vennesland, B., Pistorius, E.K., and Gewitz, H.S., HCN production by microalgae, in Cyanide in Biology, Vennesland, B., Conn, E.E., Knowles, C.J., Westley, J., and Wissing, F., Eds., Academic Press, London, 1981, p. 349. 22. Castric, P.A., Hydrogen cyanide, a secondary metabolite of Pseudomonas aeruginosa, Can. J. Microbiol., 21, 613, 1975. 23. Castric, P.A., Influence of oxygen on Pseudomonas aeroginosa hydrogen cyanide synthase, Curr. Microbiol., 29, 19, 1994. 24. Castric, P.A., Castric, K.F., and Meganathan, R., Factors influencing the termination of cynogenesis in Pseudomonas aeruginosa, in Cyanide in Biology, Vennesland, B., Conn, E.E., Knowles, C.J., Westley, J., and Wissing, F., Eds., Academic Press, London, 1981, p. 236. 25. Knowles, C.J., Microorganisms and cyanide, Bacteriol. Rev., 40, 652, 1976. 26. Michaels, R. and Corpe, W.A., Cyanide formation by Chromobacterium violaceum, J. Bacteriol., 89, 106, 1965. 27. Nazly, N., Collins, P.A., and Knowles, C.J., Cyanide production by harvested Chromobacterium violaceum, in Cyanide in Biology, Vennesland, B., Conn, E.E., Knowles, C.J., Westley, J., and Wissing, F., Eds., Academic Press, London, 1981, p. 289. 28. Niven, D.F., Collins, P.A., and Knowles, C.J., The respiratory system of Chromobacterium violaceum grown under conditions of high and low cyanide evolution, J. Gen. Microbiol., 90, 271, 1975. 29. Wissing, F., Cyanide formation from oxidation of glycine by Pseudomonas species, J. Bacteriol., 117, 1289, 1974. 30. Wissing, F., Growth curves and pH optima for cyanide producing bacteria, Physiol. Plant., 21, 589, 1968. 31. Wissing, F. and Andersen, K.S., The enzymology of cyanide protection from glycine by Pseudomonas species. Solubilization of the enzyme, in Cyanide in Biology, Vennesland, B., Conn, E.E., Knowles, C.J., Westley, J., and Wissing, F., Eds., Academic Press, London, 1981, p. 275. 32. Castric, P.A., personal communication, 2003. 33. Bach, E., The agaric Pholiota aurea: physiology and ecology, Dan. Bot. Ark., 16, 1, 1956. 34. Lebeau, J.B. and Dickson, J.G., Preliminary report on production of hydrogen cyanide by a snow mold pathogen, Phytopathology, 43, 581, 1953.

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35. Ward, E.W.B. and Lebeau, J.B., Autocatalylic production of hydrogen cyanide by a certain snow mold fungi, Can. J. Bot., 40, 85, 1962. 36. Stevens, D.L. and Strobel, G.A., Origin of cyanide in culture of psychrophilic basidiomycetes, J. Bacteriol., 95, 1094, 1968. 37. Dhillon, J.K. and Shivaraman, N., Biodegradation of cyanide compounds by Pseudomonas species, Can. J. Microbiol., 45, 201, 1999. 38. O’Sullivan, D.B. and O’Gara, F., Traits of fluorescent Pseudomonas species involved in the suppressions of plant root pathogens., Microbiol. Rev., 56, 662, 1992. 39. Duffey, S.S., Cyanide and arthropods, in Cyanide in Biology, Vennesland, B., Conn, E.E., Knowles, C.J., Westerley, J., and Wissing, F., Eds., Academic Press, London, 1981, p. 385. 40. Davis, R.H. and Nahrstedt, A., Cyanogenesis in insects (Chapter 15), in Comprehensive Insect Physiology Biochemistry and Pharmacology, Vol. II, Kerkut, G.A. and Gilbert, L.I., Eds., Pergamon Press, Oxford, 1985, p. 635. 41. Johnson, R., Wildfires, www.geo-outdoors.info/wildfires.htm, accessed: August 18, 2004. 42. Levine, J.S., Global Biomass Burning: Atmospheric, Climatic, and Biospheric Implications, MIT Press, Cambridge, MA, 1991. 43. Li, Q., Jacob, D.J., Bey, I., Yantosca, R.M., Zhao, Y., Kondo, Y., and Notholt, J., Atmospheric hydrogen cyanide (HCN): biomass burning source, ocean sink? Geophys. Res. Lett., 27, 357, 2000. 44. Milstein, M., Forest service tells fire retardant maker to remove cyanide, The Oregonian, September 20, 2000. 45. Marshall, P., Red rain — effective? Yes. Toxic? Probably, Forest Magazine, http://www.fseee.org/ forestmag/0303redrain.shtml, accessed: April 23, 2005. 46. Robine, R. and Lenglen, M., The Cyanide Industry, John Wiley & Sons, New York, 1906. 47. Bertschi, I., Yokelsom, R.J., Ward, D.E., Babbitt, R.E., Susott, R.A., Goode, J.G., and Hao, W.M., Trace gas and particle emission from forest fires in large diameter and below ground biomass fuels, J. Geophys. Res., 108, 8472, 2003. 48. Holzinger, R., Warneke, C., Jordan, A., Hansel, A., Lindinger, W., Sharffe, D.H., Schade, G., and Crutzen, P.J., Biomass burning as a source of formaldehyde, acetaldehyde, methanol, acetone, acrylonitrile, and hydrogen cyanide, Geophys. Res. Lett., 26, 1161, 1999. 49. Mauresberger, A., Methyl cyanide (CH3 CN) and hydrogen cyanide (HCN): tracers for biomass burning, Max-Planck Institute, Germany, http://mpi-hd.mpg.de/mauersberger/arnold/biomass.htm, accessed: April 23, 2005. 50. Shibamoto, T. and Bjeldanes, L.F., Introduction to Food Toxicology, Academic Press, San Diego, CA, 1993. 51. Tewe, O.O., Detoxification of cassava products and effects of residual toxins on consuming animals, in Proceedings of the FAO Expert Consultation on Roots, Tubers, Plantains and Bananas in Animal Feeding, Machin, D. and Nyvold, S., Centro International de Agricultura Tropical, Cali, Colombia, http://www.fao.org/docrep/003/T0554E/T0554E06.html, January 1991. 52. Nartey, F., Cyanogenesis in tropical foods, in Cyanide in Biology, Vennesland, B., Conn, E.E., Knowles, C.J., Westley, J., and Wissing, F., Eds., Academic Press, London, 1984, p. 115. 53. Ravindran, V., Preparation of Cassava leaf products and their use as animal feeds, in Proceedings of the FAO Expert Consultation on Roots, Tubers, Plantains and Bananas in Animal Feeding, Centro International de Agricultura Tropical (CIAT), http://www.fao.org/DOCREP/003/T0554E/T0554E00.htm, January 1991. 54. Wheeler, J.L., Implications for domestic animals of cyanogenesis in sorghum forage and hay, in Proceedings of the International Workshop on Cassava Safety, ISHS Acta Horticulturae, 375, 25, 1994. 55. Goodger, J., Characterization of cyanogenesis in Australian eucalyptus, in Proceedings of the Plant Biology, http://www.rycomusa.com/aspp2001/public/P39/0478.html, 2001. 56. Conn, E.E., Cyanogenic glycosides, in Toxicants Occurring Naturally in Foods, National Academy of Science, Washington, DC, 1973, Chapter 14. 57. Howe, R.H., The presence of cyanides in nature, in Proceedings of the Conference on Cyanide and the Environment, Tucson, AZ, 1984, p. 331.

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and the Use 4 Manufacture of Cyanide George M. Wong-Chong, David V. Nakles, and Richard G. Luthy CONTENTS 4.1

Production of Cyanide Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.1.1 Hydrogen Cyanide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.1.2 Production of Sodium Cyanide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.1.2.1 Global and U.S. Production . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.1.2.2 Production Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.1.3 Production of Ferrocyanides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.1.4 Production of Acrylonitrile . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.1.4.1 Global and U.S. Production . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.1.4.2 Production Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.2 Incidental Industrial Production of Cyanide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.2.1 Coking and Gasification of Coal . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.2.2 Blast Furnace Operations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.2.3 Aluminum Production . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.2.4 Municipal Waste and Sludge Incineration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.3 Summary and Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

42 42 42 42 44 48 49 49 50 51 52 52 52 53 53 54

Cyanide, a natural compound found in plants and animals, is believed to be a key component in the origin of life (see Chapter 1) and plays a pivotal role in today’s commerce. It is a basic component in the manufacture of a number of products including synthetic fibers and plastic, gold, agricultural herbicides, fumigants and insecticides, dyes and pigments, animal feed supplements, chelating agents for water treatment, and specialty chemicals and pharmaceuticals [1,2]. Table 4.1 presents a breakdown of the overall industrial use of hydrogen cyanide, including as a feedstock chemical for production of other cyanide compounds, as of 1991. Table 4.2 presents a list of some industries that use cyanide compounds in the manufacturing process, along with the cyanide compounds employed. The cyanide industry traces its history to about 1710 with the discovery of Prussian Blue (or ferric ferrocyanide), an iron cyanide compound, which at that time was used almost exclusively in dyeing [3,4]. However, it was not until about 1885 that substantial commercialization of cyanide, specifically potassium cyanide, occurred with the development of the McArthur-Forest process, known today as the cyanidation process, for the extraction of gold from low-grade ores [3]. This discovery represents a major sustaining factor in today’s cyanide commerce, with about 20%, or an estimated 0.6 million tons, of the worldwide production of cyanide used in mining [5,6]. This chapter discusses the manufacture of cyanide compounds, especially hydrogen cyanide, sodium cyanide, ferrocyanide, and acrylonitrile, as well as the uses of these compounds and their 41

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TABLE 4.1 Use of Hydrogen Cyanide in Manufacturing in the United States (1991 Estimate) Product

HCN used (%)

Adiponitrile for nylon Acetone cyanohydrin for plastics Sodium cyanide Cyanuric chloride for pesticides and agricultural products Chelating agents (e.g., EDTA) Methionine, animal feed Misc.: specialty chemicals and pharmaceuticals

41 28 13 9 4 2 3

Source: Data from Pesce, L.D., Kirk–Othmer Encyclopedia of Chemical Technology, Vol. 7, John Wiley & Sons, New York, 1993.

production rates. The chapter also discusses those industries where cyanide production is an incidental occurrence, such as in coking and gasification of coal, metal ore reduction in blast furnaces, the reduction of alumina, and municipal waste and sludge incineration.

4.1 PRODUCTION OF CYANIDE COMPOUNDS 4.1.1 HYDROGEN CYANIDE In 2001, the worldwide production of hydrogen cyanide was approximately 2.6 million tons [6]. The U.S. production in the period 1983 through 2001 was 0.33 to 0.75 million tons per year, as shown in Table 4.3. There are four commercial processes for the production of hydrogen cyanide. Two of these are synthesis processes involving the reaction of ammonia, methane (natural gas), and air over a platinum catalyst: (1) the Andrussow process and (2) the Blausaure–Methan–Ammoniak (BMA) process. A third process, the Shawinigan process, uses a carbon fluid bed in an electrical fluohmic furnace to react ammonia and propane. The fourth process is the acrylonitrile production process where hydrogen cyanide is produced as a by-product and which accounts for about 30% of worldwide supply [2]. Table 4.4 presents summary information about the synthesis processes for hydrogen cyanide. The Andrussow process, which is by far the dominant manufacturing process, produces hydrogen cyanide via the following reaction [2]: CH4 + NH3 + 1.5O2 → HCN + 3H2 O

(4.1)

Figure 4.1 presents a schematic flow diagram of the Andrussow process. This diagram shows the recovery/recycle of ammonia and waste heat-design features that improve the efficiency and economy of the process. Details of the process are available in the Kirk–Othmer Encyclopedia of Chemical Technology [2].

4.1.2 PRODUCTION OF SODIUM CYANIDE 4.1.2.1 Global and U.S. Production The McArthur-Forest patent for gold extraction from ore with cyanide was issued in 1887 and the cyanidation process was first used in the Crown Mine in New Zealand and then elsewhere in the

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Manufacture and the Use of Cyanide

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TABLE 4.2 Use of Cyanide Compounds in Manufacturing Industries Industry Adhesives Cement stabilizer Electroplating

Fire retardant Herbicides Fumigant, poison gas, pesticides, insecticides, parasiticide

Mining

Petroleum Photography

Pharmaceuticals (includes antibiotics, steroids, prescription and nonprescription drugs)

Primary cyanide compounds used in the process Ammonium thiocyanate Calcium cyanide Potassium- or sodium-cyanide (degreasing) Propionitrile (solvent, dielectric fluid) Nickel cyanide Silver cyanide Barium cyanide Zinc cyanide Copper cyanide Hydrogen cyanide Cyanogen chloride (metal cleaner) Mercuric potassium cyanide (mirror manufacturing) Potassium ferrocyanide Ammonium thiocyanate Cyanogen Cyanogen chloride Cyanogen bromide Zinc cyanide Copper cyanide Calcium cyanide Hydrogen cyanide Ammonium thiocyanate (pesticides) Sodium cyanide Malononitrile Cyanogen bromide Barium cyanide Calcium cyanide Ferrocyanide (used as a flotation agent for copper and lead/zinc separation) Malononitrile (lubricating oil additive) Propionitrile (solvent) Ferricyanide bleach Mercuric cyanide Hydrogen cyanide Ferricyanide Ferrocyanide Propionitrile Ammonium thiocyanate (ingredient in antibiotic preparations)

References [14] [15] [14–18]

[19,20] [14,21] [14,15]

[14–17]

[15] [17,22–24]

[14,15,22,24]

(continued)

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TABLE 4.2 Continued Industry Pigments, paints, dyes, ink, personal care products

Road salt

Rocket and missile propellant Synthetic fiber, acrylic fiber, nylon, synthetic rubber

Wine

Primary cyanide compounds used in the process Ferricyanide Ferrocyanide Ferric ferrocyanide (Prussian blue, Fe4 (Fe(CN)6 )3 ) Malononitrile Mercuric cyanide (germicidal soap) Copper cyanide (marine paint) Sodium ferrocyanide Ferric ferrocyanide (Prussian blue, Fe4 (Fe(CN)6 )3 ) Potassium ferrocyanide Cyanogen Ammonium thiocyanate Malononitrile Adiponitrile (intermediate in the manufacture of nylon) Cyanogen bromide Cyanogen chloride Hydrogen cyanide (production of nylon and other synthetic fibers and resins) Ammonium thiocyanate (improve the strength of silks) Potassium ferrocyanide

References [15,25–27]

[17,28–30]

[14,15] [14–16,31]

[32]

Source: Data from MPI, Final Technical Memorandum: Summary of cyanide investiation at SRWTP and preliminary conclusions and recommendations, report by Malcolm Pirnie, Inc., Emeryville, CA to the Sacramento Regional County Sanitation District, Sacramento Regional Wastewater Treatment Plant, Regulatory Compliance Group, Sacramento, CA, 2004.

1890s. This process started the new field of hydrometallurgy. With the advent of this process, world production of potassium cyanide rose from 5,900 tons per year in 1899 to 21,000 tons per year in 1915 [2,3]. Sodium cyanide eventually replaced the potassium salt for economic reasons, and has been the cyanide salt used in hydrometallurgical gold extraction solutions for many years. Production and use of sodium cyanide has been growing. Global annual usage of sodium cyanide in 1989 was about 340,000 tons. In the early 1990s, the total world production of sodium cyanide was estimated to be in excess of 450,000 tons. In 2001, the global production rate was about 600,000 tons per year [2,6]. 4.1.2.2 Production Methods In 1906, Robine and Lenglen [3] cited 79 processes for the production of potassium cyanide: 10 processes involving extraction from ferrocyanide; 13 processes involving extraction from thiocyanate; 28 processes involving synthesis from atmospheric nitrogen; 24 processes involving synthesis from ammonia; and four other processes. In 1891 through 1899, the Beilby process — involving synthesis from ammonia, sodium and potassium carbonate, and powdered charcoal — accounted for about 50% of the total European production of alkali cyanide [2]. In 1900, the Castner

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TABLE 4.3 Production of Hydrogen Cyanide in the United States, 1983–2001 Year

Productiona , 103 tons/yr

2001 2000 1999 1998 1997 1996 1995 1994 1993 1992 1991 1990 1989 1988 1987 1986 1985 1984 1983

750 765 745 725 710 695 675 645 600 570 565 585 565 500 470 430 365 365 330

a

Production estimates for 1983–1988; Source: Data from Pesce, L.D., Kirk–Othmer Encyclopedia of Chemical Technology, Vol. 7, John Wiley & Sons, New York, 1993. Production estimates for 1989–2001; Source: Data from Myers, E., American Chemistry Council, Washington, DC, personal communication, 2002.

TABLE 4.4 Synthesis Processes for Hydrogen Cyanide Process Andrussow Blausaure–Methan–Ammoniak Shawinigan Acrylonitrile process

Catalysts

Temperature, ◦ C

Feed

Platinum/rubidium Platinum Carbon fluid bed in a fluohmic furnace By-product

1100 1100 1350–1650 400–510

NH3 , air, and CH4 NH3 and CH4 NH3 and C3 H8 NH3 , air, and C3 H6

Source: Data from Pesce, L.D., Kirk-Othmer Encyclopedia of Chemical Technology, Vol. 7, John Wiley & Sons, New York, 1993.

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Wastewater

NH3 Recycle

NH3 Fractionator

Waste-Heat Boiler

NH3 / Water

NH3 Feed Air Feed

Reactor

Steam NH3 Absorber Diammonium Phosphate Solution

Natural Gas Feed

NH3 Stripper

Off-Gas Minus NH3

Monoammonium Phosphate Solution

Acid Waste Gases to Boiler or Flare

HCN Absorber

HCN/Water Coolers

Waste Water

HCN Stripper

Steam

HCN/Water

HCN Fractionator

SO2

HCN with SO 2 Inhibitor

FIGURE 4.1 Schematic flow diagram of the Andrussow hydrogen cyanide production process. (Source: Pesce, L.D., Kirk–Othmer Encyclopedia of Chemical Technology, Vol. 7, John Wiley & Sons, New York, 1993. Reprinted with Permission of John Wiley & Sons, Inc.)

process replaced the Beilby process and dominated production through 1960 for both potassium and sodium cyanide. For the production of sodium cyanide, the Castner process employs elemental sodium and a reaction with ammonia and carbon as follows: 2Na + 2NH3 + 2C → 2NaCN + 3H2

(4.2)

Low yields and elevated costs led to the obsolescence of the Castner process. This process was replaced by the neutralization or wet processes that react hydrogen cyanide from the Andrussow or BMA processes with a sodium hydroxide solution: HCN + NaOH → NaCN + H2 O

(4.3)

Most modern, high tonnage production plants use essentially purified anhydrous liquid hydrogen cyanide to react with sodium hydroxide to produce a product consisting of 99% sodium cyanide. The manufacturing process includes the evaporation of water and crystallization of the sodium cyanide. Control of the process is critical to maximize the average crystal size; to avoid hydrogen

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Dehumidifier Scrubber

Vacuum system

Dust Scrubber

Condenser

Waste Cyclone Separator Air Heater

50% Caustic

Hydrogen cyanide

Crystallizer System

Filter

Briquetter

Mixing Conveyor Steam Screens

Product to packaging and storage

FIGURE 4.2 Production process flow diagram for sodium cyanide. (Source: Pesce, L.D., Kirk–Othmer Encyclopedia of Chemical Technology, Vol. 7, John Wiley & Sons, New York, 1993. Reprinted with permission of John Wiley & Sons, Inc.)

cyanide polymer formation, which produces an off-white product; and to minimize sodium formate formation, which reduces product purity. Figure 4.2 presents a process flow diagram for a typical sodium cyanide production plant. An occasionally used, alternative process is the direct absorption of crude hydrogen cyanide gas from the manufacturing operation into a sodium hydroxide solution. However, the purity of the sodium cyanide product is lower, that is, approximately 96 to 97% [2]. The primary impurities are sodium carbonate and sodium formate. The formation of larger crystals facilitates the dewatering in the filtration step. In many plants, the moist salt from the filter is passed through a mixing conveyor to destroy the lumps. Often, heated air (450◦ C) is passed through the cake on the filter and through the mixing conveyor. Drying is completed in a hot-air conveyor-dryer. This approach to drying avoids the overheating of the sodium cyanide crystals, thus minimizing the formation of sodium formate in the dried product. A slight excess of sodium hydroxide must be maintained at all stages of processing to maintain an elevated pH, which

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prevents the formation of a black or brown hydrogen cyanide polymer. The product, as shipped, must also contain a slight excess of sodium hydroxide, to ensure that the product yields clear solutions following arrival at its destination. Inherently, the sodium cyanide forms a 50-µm diameter crystal, yielding a dusty solid of low bulk density that must be compacted or fused into larger particles for safer handling. Due to the expense associated with melting the product for casting it in molds, most processes employ mechanical compacting devices that produce either briquettes or granular products. The compaction process occurs using heat and pressure. Most sodium cyanide is sold in dry form to minimize transportation costs although appreciable tonnage is also sold as a 30% aqueous solution [2]. About 90% of today’s sodium cyanide is used in gold extraction [5,6]. Plants for the production of sodium cyanide, using these processes, are operating in the United States, Italy, Japan, the United Kingdom, Australia, Germany, and China.

4.1.3 PRODUCTION OF FERROCYANIDES Ferric ferrocyanide, also known as Prussian Blue (Fe4 [Fe(CN)6 ]3 ), was the first cyanide compound put to commercial use. The compound was discovered by a Berlin color maker in 1704 [3]. This led to a long history of ferrocyanide chemistry, which has resulted in the use of these compounds in a wide variety of industrially significant applications. A treatise on the chemistry of ferrocyanides [7] describes some 22 applications, and these are listed in Table 4.5. In the late 1700s through the early 1900s, ferrocyanide salts were produced by (1) the synthetic fusion of nitrogenous organic residues (e.g., animal blood, hides, hornes, waste/scrap leather, etc.), potash, and iron, and (2) the direct extraction from illuminating-gas and from the by-product

TABLE 4.5 Uses of Ferrocyanides and their Derivatives in Industry Analytical chemistry Anticaking agent Blueprints Case hardening and heat treatment of steel Chemical synthesis: catalysts, reaction intermediates, and reagents Chemotherapy Corrosion inhibitors Desulfurization of coke oven gas Detergents Dying of textiles Electrical equipment treatment: corrosion resistance; arc stabilization and lowering of grounding resistance Electroplating Minerals dressing, beneficiation, and mining Pesticides Petroleum refining: trace metals removal Photography Pigments and dyes Pickling of steel Rubber: peptizing agent, stabilization agent, and accelerator Separation and identification of organic bases Trace metals removal in fermentation Source: Data from ACC, The Chemistry of the Ferrocyanides, American Cyanamid Co., New York, NY, 1953.

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of illuminating-gas clean-up (e.g., spent iron oxide boxes for gas purification) [3]. It was estimated that about 1.8% of the nitrogen in coal reacted to form hydrogen cyanide during coal gasification. In the direct gas extraction processes, the illuminating gas was scrubbed with an alkaline iron salt solution. Robine and Lenglen [3] discussed in detail nine processes for the direct extraction of cyanide from illuminating gas, three processes for extraction from ammoniacal liquor, and 11 processes for recovering ferrocyanide from spent iron oxide. In the synthesis from nitrogenous organic matter, the process chemistry for making potassium ferrocyanide was thought to be: K2 CO3 + Nitrogenous Matter + Energy → KCN + · · ·

(4.4)

6KCN + Fe2+ → K4 Fe(CN)6 + 2K +

(4.5)

In the first reaction, hydrogen cyanide is produced by the thermal breakdown of the organic matter in an oxygen controlled environment (Equation [4.4]). Subsequently, the hydrogen cyanide reacts with potassium to form potassium cyanide. The potassium cyanide, in turn, reacts with the iron to form potassium ferrocyanide as shown in Equation (4.5). Today, ferrocyanide production utilizes the crude sodium cyanide, produced as described in Section 4.1.2, and ferrous sulfate to form sodium ferrocyanide: 6NaCN + FeSO4 + Heat → Na4 Fe(CN)6 + Na2 SO4

(4.6)

The sodium ferrocyanide is recovered by crystallization as the decahydrate salt. The potassium salt is produced by reacting sodium ferrocyanide with calcium hydroxide and potassium chloride and carbonate according to the following reactions: Na4 Fe(CN)6 + 2Ca(OH)2 → Ca2 Fe(CN)6(s) + 4Na(OH)

(4.7)

Ca2 Fe(CN)6 + 2K2 CO3 → K4 Fe(CN)6 + CaCO3(s)

(4.8)

In earlier times, ca. 1900, Prussian Blue was produced in a two stage process. The first stage reacted potassium ferrocyanide and ferrous sulfate to form a grayish-white precipitate of potassium ferric–ferrocyanide. In the second stage, the potassium ferric–ferrocyanide is oxidized to the tetrairon(III) tris(hexakiscyanoferrate), Fe4 [Fe(CN)6 ]3 [3]. Today, the production of Prussian Blue is more direct, where ferrocyanide is reacted with excess iron(III) to produce the intense blue precipitate [2].

4.1.4 PRODUCTION OF ACRYLONITRILE Acrylonitrile [C3 H3 N], also called vinyl cyanide, is among the top 50 chemicals produced in the United States as a result of the tremendous growth in its use as a starting material for a wide range of chemical and polymer products. Acrylic fibers remain the largest use of acrylonitrile. Other significant uses are resins and nitrile elastomers and as an intermediate in the production of adiponitrile and acrylamide. 4.1.4.1 Global and U.S. Production Worldwide production of acrylonitrile was approximately 3.2 million tons in 1988 [8]. As shown in Table 4.6, more than one-half of that production was located in Western Europe and the United States. In the United States, BP Chemicals dominated production, supplying more than one-third of domestic production. Nearly one-half of the United States production was exported in 1988, with most going to Japan and the Far East [8]. This export market grew steadily from the mid-1970s

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TABLE 4.6 Worldwide Acrylonitrile Production, 1988 Production, 103 tons

Region Western Europe United States Japan Far East Mexico

1200 1170 600 200 60

Total

3230

Source: From Brazdil, F., Kirk-Othmer Encylopedia of Chemical Technology, Vol. 1, John Wiley & Sons, New York, 1993.

TABLE 4.7 Worldwide Acrylonitrile Demand, 103 Tons per Year Region Western Europe Japan United States Far East Mexico/South America Total

1976

1980

1985

1988

880 570 590 200 81

880 510 660 270 130

1140 635 640 385 200

1200 680 660 560 250

2321

2450

3000

3350

Source: Data from Brazdil, F., Kirk–Othmer Encyclopedia of Chemical Technology, Vol. 1, John Wiley & Sons, New York, 1993.

to 1988. During this period, it increased from 10% in the mid-1970s to 53% and 43% in 1987 and 1988, respectively. The large exports to the Far East were the result of higher raw material costs (i.e., propylene costs) relative to the United States. A more detailed breakdown of the world demand for acrylonitrile for the period between 1976 and 1988 is provided in Table 4.7. Growth in demand during this period averaged about 3.6% per year between 1984 and 1988. Projections beyond 1988 were 3% per year through 1993.

4.1.4.2 Production Methods Prior to 1960, processes based on either ethylene oxide and hydrogen cyanide or acetylene and hydrogen cyanide were used to produce acrylonitrile. Growth in the demand for acrylic fibers around 1950 spurred improvements in process technology and resulted in the discovery of a heterogeneous vapor-phase catalytic process. This process, which produced acrylonitrile using selective oxidation of propylene and ammonia, is commonly referred to as the propylene ammoxidation process. This process was introduced in 1960 and eventually displaced all other acrlyonitrile manufacturing processes. As of 1988, over 90% of the approximately 3.2 million metric tons of acrylonitrile produced worldwide each year was manufactured using the propylene ammoxidation process [8].

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Fluid-Bed Catalytic Reactor 400° - 510° C 49-196 kPa

Absorber

51

Acrylonitrile recovery column

Acetonitrile recovery column

Lights column

Product column

Crude acrylonitrile Crude acetonitrile

Off-Gas

H.P. Steam

Crude HCN

Product acrylonitrile

H2O

BFW

Air Ammonia START

H2O

Proplene

Heavy impurities

FIGURE 4.3 Process flow diagram for ammoxidation process. (Source: Brazdil, F., Kirk–Othmer Encyclopedia of Chemical Technology, Vol. 1, John Wiley & Sons, New York, 1993. Reprinted with permission of John Wiley & Sons, Inc.)

The primary chemical reaction of the propylene ammoxidation process is as follows: C3 H6 + NH3 + 1.5O2

CATALYST

−→

C3 H3 N + 3H2 O

(4.9)

A process diagram of the commercial process is shown in Figure 4.3. This process uses a fluidized bed reactor, in which propylene, ammonia, and air contact a solid catalyst at 400 to 510◦ C and 49 to 196 kPa gauge. It is a single pass process that achieves about 98% conversion of the propylene, and uses about 1.1 kg of propylene per kilogram of acrylonitrile produced. As shown in Figure 4.3, hydrogen cyanide is a by-product of the acrylonitrile production process. This hydrogen cyanide can be processed as a salable product or used in the manufacture of methyl methacrylate and acetonitrile, common industrial solvents [8].

4.2 INCIDENTAL INDUSTRIAL PRODUCTION OF CYANIDE Many industrial operations that employ thermal processing of carbonaceous materials produce small quantities of cyanide. Included among these operations are • • • •

Coking and gasification of coal Blast furnace processing for iron and nonferrous metal oxide reduction Alumina reduction Municipal waste and sludge incineration

Brief discussions of these operations follow. More details and discussions of these operations are presented in Chapter 26.

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4.2.1 COKING AND GASIFICATION OF COAL The coking operation involves distillation of coal by indirectly heating the coal in the absence of air to temperatures in the range of 900 to 1100◦ C to vaporize all volatile constituents in the coal [9]. These volatile constituents include a range of hydrocarbons (e.g., benzene, toluene, methane, naphthalene, phenols, xylenes, and polynuclear aromatic hydrocarbons), nitrogenous compounds including ammonia, and sulfurous compounds including hydrogen sulfide. The heating period can range from 18 to 36 h [9]. Hydrogen cyanide is formed, albeit in relatively small amounts, due to the high temperature, reducing atmosphere, and the presence of nitrogen and carbon. The quantity of cyanide produced in the coking of coal has been reported to be about 1.5 to 2.0% of the nitrogen content of the coal [3]. A portion of the cyanide remains in the coke oven gas while the remainder leaves the coking system in the waste ammonia liquor wastewater [3]. In the early days of the cyanide industry, ca. 1900, cyanide was recovered from illuminating-gas production, which was essentially a coking process [3,10]. The cyanide was recovered by absorbing the hydrogen cyanide from the gas stream into an alkaline iron salt solution to form an alkali ferrocyanide product. As previously noted, Robine and Lenglen [3] provide detailed descriptions of nine processes for the direct extraction of hydrogen cyanide from illuminating gas and three processes for the extraction of cyanide compounds from the ammoniacal liquors. The coal gasification process is similar to coking, in that the coal is heated to similar temperatures. However, air is introduced into the coal gasifier to combust a portion of the coal, which produces the heat for the steam/coal reactions that take place. The quantity of air that is injected is controlled to maintain the gasification temperature and the quality of product gas that is produced (high or low BTU content). Again, similar to the coking process, some hydrogen cyanide is formed given the conditions of high temperature, abundance of carbon (as low molecular weight hydrocarbons, carbon dioxide or carbon monoxide), and nitrogen from the injected air or coal. Just as in the coking operation, some of this cyanide remains in the product gas and the remaining portion exits the plant in the gas cleaning wastewater.

4.2.2 BLAST FURNACE OPERATIONS In blast furnace operations, the furnace is charged with coke, metal oxide, and limestone flux. The furnace is heated to maintain a temperature profile of about 900◦ C in the upper section of the furnace and greater than 1770◦ C in the bottom section around the tuyeres. Periodic blasts of air and supplemental fuel (oil) are fired into the reaction mixture [9]. This environment of high temperature, available carbon as carbon monoxide and carbon dioxide, and nitrogen from the air blast provide conditions for hydrogen cyanide formation. Flue gas emitted from the top of the blast furnace, about 2.5 to 3.5 tons of gas per ton of iron produced, has a heating value of 80 to 90 BTU/SCF [9] and fuels auxiliary stoves or powers boilers and blowers. The blast furnace gas is laden with dust, as much as 0.05 tons per ton of iron produced, that must be removed to prevent clogging of combustion equipment. About 70% of the dust is removed by bag-houses and dry scrubbing and the remaining 30% by wet scrubbing. The hydrogen cyanide that is present in the gas is removed during the wet scrubbing process and reports to the blast furnace gas scrubber water.

4.2.3 ALUMINUM PRODUCTION Aluminum is manufactured via the electrometallurgical reduction of alumina (Al2 O3 (s)) in the Hall– Heroult process [2,11]. The alumina is placed in a molten cryolite (Na 3 AlF6 ) bath in a carbon-lined cell, or “pot.” An electrical current is passed through the reaction mixture using carbon anodes placed in the molten mixture. The carbon pot liner, which typically is 15 in. thick, serves as the cathode. The molten mass attains a temperature of about 950 to 1,000◦ C [2,11]. The carbon anode is rapidly

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consumed while the carbon pot liner cathode is re-used, usually for several years until it becomes unacceptably contaminated [11]. In this reactor, the conditions of high temperature, available carbon, and available nitrogen provide the opportunity for hydrogen cyanide formation. Air entering at reactor seals is the primary source of nitrogen. The hydrogen cyanide produced becomes absorbed into the carbon of the pot liner and at the end of a pot liner processing life, the cyanide concentration can be as high as 0.9% by weight [11]. Cyanide levels vary within a pot, with highest concentrations observed in the potlining at the side walls. Spent potlining containing cyanide and other contaminants is removed for treatment and disposal. Until the early 1970s, spent potlining was managed as an inert residue and was often used as onsite fill material. Today, spent potlining is a listed hazardous waste (K088) in the United States. Based on data from 1991 to 1997, the annual production of K088 waste is approximately 80,000 to 100,000 tons [12]. Past practices for the disposal of spent potlining have resulted in adverse environmental impacts such as contamination of soil and groundwater, as discussed in more detail in Chapter 27.

4.2.4 MUNICIPAL WASTE AND SLUDGE INCINERATION In areas where land-disposal of municipal wastewater treatment sludge is not practical, sludge incineration is an accepted disposal alternative. In some instances, the economics are improved by combined incineration of municipal refuse and sludge coupled with co-generation of electricity. In the northeastern states of Massachusetts and Rhode Island, municipal wastewater sludge incineration is widely practiced, for example, at the Cranston, Rhode Island and Fitchburg, Massachusetts municipal wastewater treatment plants. These sludges contain about 5 to 6% nitrogen on a dry weight basis and are essentially organic in composition. They are charged to the incinerator with about 70 to 80% water. During the course of sludge incineration, there tends to be pockets of reduced conditions, especially in the early stages of incineration. These localized conditions, coupled with the high temperature and availability of nitrogen and carbon, provide the opportunity for hydrogen cyanide formation. This cyanide leaves the incinerator in the exhaust gases and is transferred to the off-gas scrubbing water. At the Cranston, Rhode Island POTW, an average of 2.08 g cyanide per kilogram of dry sludge incinerated is removed in the scrubber water [13].

4.3 SUMMARY AND CONCLUSIONS • Hydrogen cyanide and other cyanide compounds are used extensively in manufacturing, including the production of synthetic fibers and plastics, agricultural herbicides, fumigants and insecticides, dyes and pigments, animal feed supplements, chelating agents for water treatment, and specialty chemicals and pharmaceuticals. • Sodium cyanide is used extensively for the extraction of gold from ore in hydrometallurgical gold mining. • In 2001, the worldwide production of hydrogen cyanide was approximately 2.6 million tons. • Hydrogen cyanide is manufactured primarily by the Andrussow process, which involves the reaction of methane, ammonia, and oxygen. Hydrogen cyanide from this process is reacted with sodium hydroxide solution to form sodium cyanide, the most common solid form of cyanide that is used in commerce. • Potassium ferrocyanide (K4 [Fe(CN)6 ]), and related solid-phase precipitates, especially ferric ferrocyanide (Fe4 [Fe(CN)6 ]3 ), or Prussian Blue, is produced in large quantity for a variety of specialty uses. • Acrylonitrile, also called vinyl cyanide, is produced in large quantities because of its use as a starting material for a wide range of chemical and polymer products.

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• Various cyanide compounds are produced incidentally during the manufacture of coke, steel, and aluminum, and during the incineration of municipal waste and wastewater sludges.

REFERENCES 1. Homan, E.R., Reactions processes and materials with potential for cyanide exposure, in Clinical and Experimental Toxicology of Cyanides, Ballantyne, B. and Marrs, T.C., Eds., John Wright, Bristol, UK, 1987, p. 1. 2. Pesce, L.D., Cyanides, in Kirk–Othmer Encyclopedia of Chemical Technology, Vol. 7, John Wiley & Sons, New York, 1993. 3. Robine, R. and Lenglen, M., The Cyanide Industry, John Wiley & Sons, New York, 1906. 4. Ritter, S.K., Prussian Blue still a hot topic, Chem. & Eng. News, May 2, 2005, p. 32. 5. Logsdon, M.J., Hagelstein, K., and Mudder, T.I., The management of cyanide in gold extraction, International Center for Metals and the Environment, Ottawa, Canada, 1999. 6. Young, C.A., Cyanide: just the facts, in Cyanide: Social, Industrial and Economic Aspects, Young, C.A., Twidwell, L.G., and Anderson, C.G., Eds., The Minerals, Metals & Materials Society, Warrendale, PA, 2001, p. 97. 7. ACC, The Chemistry of the Ferrocyanides, American Cyanamid Co., New York, NY, 1953. 8. Brazdil, F., Acrylonitrile, in Kirk-Othmer Encyclopedia of Chemical Technology, Vol. 1, John Wiley & Sons, Inc., New York, 1993, p. 352. 9. USS, The Making, Shaping and Treating of Steel, McGannon, H.E., Ed., United States Steel Corp., Pittsburgh, PA 1971. 10. Hayes, T.D., Linz, D.G., Nakles, D.V., and Leuschner, A.P., Eds., Management of Manufactured Gas Plant Sites, Vol. 1, Amherst Scientific Publishers, Amherst, MA, 1996, Chapter 2, p. 5. 11. USEPA, Best demonstrated available technology (BDAT) background document for spent aluminum potliners — K088, U.S. Environmental Protection Agency, Office of Solid Waste, http:// www.epa.gov/epaoswer/hazwaste/ldr/k088/k088back.pdf, 2000. 12. USEPA, Land disposal restrictions — background document to establish the effective date for amended treatment standards for spent aluminum potliners (proposed rule), U.S. Environmental Protection Agency, Washington, D.C., http://www.epa.gov/epaoswer/hazwaste/ldr/k088/landdisp.pdf, 2000. 13. Bratina, C., Cranston, Rhode Island wastewater treatment plant, personal communication, 2004. 14. ATSDR, Toxicological profile for cyanide (update), U.S. Department of Health and Human Services, Public Health Service, Agency for Toxic Substances and Disease Registry, Atlanta, GA, 1997. 15. USNLM, Toxicology and Environmental Health Information Program. Toxicology Data Network (TOXNET), U.S. National Library of Medicine, http://toxnet.nlm.nih.gov, accessed: February 18, 2005. 16. Boucabeille, C., Bories, A., Olliver, P., and Michel, G., Microbial degradation of metal complexed cyanides and thiocyanate from mining wastewaters, Environ. Pollut., 84, 59, 1994. 17. Kjeldsen, P., Behaviour of cyanides in soil and groundwater: a review, Water, Air, Soil Pollut., 115, 279, 1999. 18. Patterson, J.W., Cyanide, in Industrial Wastewater Treatment Technology, Butterworth, Boston, 1985, p. 115. 19. Hopkins, S.J., Special water quality survey of the Pecos and Gallinas Rivers below the Viveash and Manuelitas fires of 2000, Surveillance and Standards Section, New Mexico Environment Department, 2001, Available at http://www.nmenv.state.nm.us/swqb/viveash_fire_report_02-2001.html, accessed: February 25, 2005. 20. Little, E. and Calfee, R., The effects of UVB radiation on the toxicity of firefighting chemicals, U.S. Department of Agriculture, Forest Service and Wildland Fire Chemical Systems, 2002. 21. Rapean, J.C., Johnson, R.A., and Hanson, T.P., Biodegradation of cyanide: Nitrite interference in the standard test for total cyanide, in Proceedings of the 35th Purdue Industrial Waste Conference, West Lafayette, IN, 1980, p. 430. 22. USEPA, Sustainable industry project phase I report, Chapter 3, The photoimaging industry, U.S. Environmental Protection Agency, Office of Policy Development, Washington, D.C., 2002.

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23. Meeussen, J.L., Keizer, M.G., and de Haan, F.A.M., Chemical stability and decomposition rate of iron cyanide complexes in soil solutions, Environ. Sci. Technol., 26, 511, 1992. 24. Owerbach, D., Analysis and sample stability of cyanide in industrial effluents, J. Water Pollut. Control Fed., 52, 11, 1980. 25. USEPA, Seminar publication — National conference on urban runoff management: Enhancing urban watershed at the local, county and state levels, EPA-625/R-95-003, U.S. Environmental Protection Agency, Washington, D.C., 1995. 26. USEPA, Profile of the wood furniture and fixtures industry, EPA-310/R-95-003, U.S. Environmental Protection Agency, Office of Enforcement and Compliance Assurance, Washington, D.C., 1995. 27. USEPA, Pharmaceutical manufacturing category effluent limitations guidelines, Pretreatment standards and new source performance standards, Final Rule, U.S. Environmental Protection Agency, 40 CFR, Parts 136 and 439, 1998. 28. USEPA, Seminar publication — Wellhead protection: a guide for small communities, EPA-625/R-93002, U.S. Environmental Protection Agency, Washington, D.C., 1993. 29. ICF, Construction and demolition waste landfills, Report by ICF, Inc. Fairfax, VA, Contract No. 68-W30008, U.S. Environmental Protection Agency, Office of Solid Waste, Washington, D.C., 1995. 30. Paschka, M.G., Ghosh, R.S., and Dzombak, D.A., Potential water-quality effects from iron cyanide anticaking agents in road salt, Water Environ. Res., 71, 1235, 1999. 31. USEPA, Profile of the motor vehicle assembly industry, EPA-310/R-95-009, U.S. Environmental Protection Agency, Washington, D.C., 1995. 32. USEPA, Consumer fact sheet on cyanide, U.S. Environmental Protection Agency, Office of Ground Water and Drinking Water, http://www.epa.gov/ogwdw/dwh/c-ioc/cyanide.html. Accessed: February 25, 2005. 33. MPI, Final Technical Memorandum: Summary of cyanide investigation at SRWTP and preliminary conclusions and recommendations, report by Malcolm Pirnie, Inc., Emeryville, CA to the Sacramento Regional County Sanitation District, Sacramento Regional Wastewater Treatment Plant, Regulatory Compliance Group, Sacramento, CA, 2004. 34. Myers, E., American Chemistry Council, Washington, DC, personal communication, 2002.

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Properties 5 Physical–Chemical and Reactivity of Cyanide in Water and Soil

David A. Dzombak, Rajat S. Ghosh, and Thomas C. Young CONTENTS 5.1

Free Cyanide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.1.1 Cyanide Ion Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.1.2 HCN Formation and Dissociation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.1.3 HCN Volatilization . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.1.4 Free Cyanide Adsorption to Soil and Sediment . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.1.5 Free Cyanide Oxidation. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.1.6 Free Cyanide Hydrolysis. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.2 Metal Cyanides: Aqueous Species . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.2.1 Weak Metal–Cyanide Complexes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.2.1.1 Formation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.2.1.2 Dissociation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.2.1.3 Adsorption on Soil and Sediment . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.2.1.4 Oxidation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.2.2 Strong Metal–Cyanide Complexes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.2.2.1 Formation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.2.2.2 Dissociation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.2.2.3 Adsorption on Soil and Sediment . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.2.2.4 Oxidation–Reduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.3 Metal–Cyanides: Solid Phase Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.3.1 Simple Cyanide Solids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.3.2 Alkali or Alkaline Earth Metal–Metal Cyanide Complex Solids . . . . . . . . . . . . . . . 5.3.3 Other Metal–Metal Cyanide Complex Solids. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.4 Cyanate . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.5 Thiocyanate. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.6 Organocyanides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.7 Summary and Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

58 58 58 60 61 62 64 65 65 65 67 68 71 73 73 75 76 78 79 80 80 80 82 84 86 88 88

The reactivity, fate, and toxicity of cyanide in water and soil is highly dependent on the chemical speciation of the cyanide. As outlined in Chapter 2, many different soluble and solid forms of cyanide exist. The simplest form of soluble cyanide is the negatively charged cyanide ion, CN− , which is composed of a carbon atom triple bonded to a nitrogen atom (–C≡N). The nature of this triple bond controls the reactivity of the cyanide anion, including complexation with other metal cations, 57

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formation of molecular hydrogen cyanide (HCN), oxidation of cyanide to cyanate, and adsorption onto clays and other soil components. In environmental systems, wastewaters, and wastes, cyanide usually is found in free and complexed forms, as HCN and as metal–cyanide complexes. Because of a reactive electronic arrangement, cyanide anions can readily form metal–cyanide complexes with most metal cations. Most of these complexes exist as soluble species, but many, particularly iron-cyanide complexes, can react further with metal cations to form stable cyanide solids. The soluble and solid phase cyanide species that occur most often in water and soil are outlined in Chapter 2 and examined in more detail here. In this chapter, the specific physical–chemical properties and reactivity characteristics of the different chemical forms of cyanide are presented. Included are examinations of the nature of bonding in and with the cyano group and free cyanide speciation; the properties and reactivities of soluble metal–cyanide complexes; the properties and reactivities of metal–cyanide complex solids; and the properties and reactivities of cyanate, thiocyanate, and organocyanide compounds.

5.1 FREE CYANIDE 5.1.1 CYANIDE ION BONDING Free cyanide consists of the cyanide anion, CN− , and molecular hydrogen cyanide, HCN, both existing as water soluble entities. The cyanide ion acts as a monodentate ligand with the carbon acting as the donor atom, and also as an ambidentate ligand acting as a donor at both ends of the ion [1]. Several structural factors govern the reactivity of free cyanide. The triple bonded structure of a cyanide anion is comprised of a sigma bond, two π bonds, and two empty bonding orbitals [2]. The “s” and the “p” orbitals are filled with maximum number of electrons, while the “d” and “f” orbitals are empty. This configuration allows for a number of bonding arrangements. Since halogens also have filled “s” and “p” orbitals, the behavior of the cyanide anion is similar to that of halogens [3]. The cyanide ion is considered a pseudo-halide in that it can form π -acceptor covalent bonds with transition metals [3]. It may also share electrons at the triple bond with the Group VI elements oxygen and sulfur, forming cyanate, CNO− , or thiocyanate, SCN− [3], or may act as a strong nucleophile in reactions with organic molecules, for example, nucleophilic addition reactions with aldehydes and ketones to form cyanohydrins [4]. The cyanide ion readily forms neutral compounds or anionic complexes with most major metal cations. The partially or wholly filled “d” orbitals of transition series metals can form covalent bonds with the empty anti-bonding orbitals of the cyanide ion. This involves acceptance of electron density into π orbitals of the carbon atom. The cyanide ion is a strong σ donor, which is responsible for the high stability of some of the metal–cyanide complexes [3].

5.1.2 HCN FORMATION AND DISSOCIATION The cyanide anion protonates in water to form hydrocyanic acid, HCN, the most toxic form of cyanide (see Chapters 13 and 14). The pKa for HCN dissociation reaction is 9.24 at 25◦ C [5]. Thus, at pH greater than 9.24, cyanide anion dominates free cyanide speciation, while soluble HCN is the dominant species under acidic to neutral pH conditions (pH < 9.24). The free cyanide dissociation reaction is as follows: HCN = H+ + CN− ,

pK a = 9.24 at 25◦ C,

I=0

(5.1)

Figure 5.1 shows the distribution of HCN and CN− species as a function of pH for a simple aqueous solution at 25◦ C. The temperature dependence of the equilibrium constant governing the species

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Physical–Chemical Properties and Reactivity

59

1.0 CN–

Ionization fraction ([HCN]/CNT, [CN-]/ CNT)

HCN 0.8

0.6

0.4

0.2

0

5

6

7

8 pH

9

10

11

FIGURE 5.1 Free cyanide species distribution as a function of pH at 25◦ C(pKa = 9.24 for HCN dissociation at T = 25◦ C, I = 0).

distribution of free cyanide can be calculated via the van’t Hoff equation: ln(K2 /K1 ) = (Hr,25C /R)[1/T1 − 1/T2 ]

(5.2)

where Hr,25C is the standard enthalpy change of reaction at 25◦ C(298 K), R is the molar gas constant (8.314 × 10−3 kJ mol−1 K −1 ), T1 is the reference temperature (298 K), and T2 is the temperature of interest in K. The standard enthalpy change for the reaction given in Equation (5.1) is 146 kJ mol−1 , as calculated using the thermodynamic data compiled in Stumm and Morgan [6]. Substitution of this value, and assuming it is approximately constant for the temperature range 5 to 30◦ C, enables calculation of the temperature dependence of the acidity constant in Equation (5.1): KT = exp[1.756 × 104 K(3.356 × 10−3 K −1 − T−1 ) − 21.28]

(5.3)

where KT is the equilibrium constant for HCN dissociation at the temperature T (K) of interest. Combining Equation (5.3) with the mass action equation for the reaction in Equation (5.1), and the mass balance equation for free cyanide (molar concentrations in [ ]), TOTCN = [HCN] + [CN− ]

(5.4)

yields the following expression for the species distribution fractions for HCN and CN− : αHCN = [HCN]/TOTCN = {H+ }/[{H+ } + exp[1.756 × 104 K(3.356 × 10−3 K −1 − T−1 ) − 21.28]] αCN = [CN− ]/TOTCN = 1 − αHCN

(5.5) (5.6)

where {H+ } is the hydrogen ion activity, 10−pH . The species distribution fraction for HCN, αHCN , is presented in Figure 5.2 for temperatures between 5 and 30◦ C (278 and 303 K), and zero ionic

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60

1 0.9 0.8

T = 5°C T = 25°C

[HCN]/CNT

0.7

T = 15°C

0.6 0.5

T = 10°C

T = 20°C

0.4

T = 30°C

0.3 0.2 0.1 0 6

7

8

9

10

11

pH

FIGURE 5.2 Ionization fraction for HCN as a function of pH and temperature (I = 0).

TABLE 5.1 Literature Values of Henry’s Law Constant (KH,HCN ) for HCN Temp. (◦ C) 25 Not given Not given 25

KH (atm L mol−1 )

−1 KH (mg L−1 a /mg Lw )

0.122 0.073 0.104 to 0.114 0.115

0.005 0.003 0.0043 to 0.0047 0.0047

Reference Bodek et al. [7] Doudoroff [80] Smith and Mudder [2] Avedesian [81]

strength (I). As is evident in Figure 5.2, temperature has a significant effect on free cyanide species distribution. As temperature decreases, dissociation of HCN decreases, extending the species dominance of HCN to higher pH values.

5.1.3 HCN VOLATILIZATION Hydrogen cyanide has a very low boiling point (25.7◦ C) and thus is volatile in water under environmental conditions. The equilibrium air–water partitioning of HCN can be described by Henry’s Law: PHCN = KH,HCN [HCN]

(5.7)

where PHCN is the partial pressure of HCN gas, atm, KH,HCN the Henry’s Law constant, atm L mol−1 , and [HCN] the equilibrium aqueous phase concentration of HCN, mol L−1 . Table 5.1 lists reported values of Henry’s Law constant for HCN. Henry’s Law constants with units relevant to Equation (5.1) are provided, along with dimensionless analogs corresponding to an equilibrium partitioning expression in which both the aqueous and gas phase concentrations are expressed in the same mass concentration units. Note that the Henry’s Law constant is a function of temperature. There are various empirical relationships that express Henry’s Law constant as a function

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of temperature. One such relationship, reported by Bodek et al. [7], is as follows: log KH,HCN = −1272.9/T + 6.238

(5.8)

where KH,HCN is the Henry’s Law constant, mm Hg/M and T the Temperature, K. Equation (5.8) is reported to be valid for HCN concentrations ranging from 0.01 to 0.5 M and temperatures from 20 to 95◦ C.

5.1.4 FREE CYANIDE ADSORPTION TO SOIL AND SEDIMENT Free cyanide (CN− , HCN) adsorbs weakly on soils and sediment. The cyanide anion can be retained by soils with anion exchange capacity, but in the pH range 4 to 9 of interest for most soils, HCN is the dominant form of cyanide and CN− concentrations are very low. HCN adsorbs weakly or not all to inorganic soil components such as iron oxide [8], aluminum oxide, clay, and sand [9]. However, HCN has been shown to adsorb significantly to soils with appreciable organic carbon content. The magnitude of cyanide adsorption onto soils tested by Chatwin et al. [10] showed excellent correlation with organic carbon content. Higgins and Dzombak [9] further demonstrated the interaction of HCN with organic carbon in experiments with activated carbon and freshwater sediment. They developed an expression relating sorbed HCN concentration, CS , to aqueous phase concentration, Cw , through an organic carbon normalized distribution coefficient Koc (=Kd /foc ). CS = Koc Cw foc = (6.5 L/gs )Cw foc

(5.9)

where CS is in µg/gs , CW is in µg/L, and foc is the fraction of organic carbon in the adsorbent. The experiments upon which this linear relationship is based all involved low concentrations of free cyanide in water ( 2, so electrostatic repulsion of the negative Ni(CN)2− 4 species is in part responsible for the absence of surface binding. For the most part, however, it is the presence of cyanide that inhibits adsorption of nickel. Figure 5.8 shows data for similar sets of batch adsorption experiments performed with FeOOH(s). Once again, the adsorption of nickel observed in the system with CNT = 10−5 M is very similar to the adsorption of nickel on FeOOH(s) in the absence of cyanide. In the systems with CNT = 10−4 and 10−3 M, adsorption of nickel is enhanced at lower pH values, and inhibited at higher pH values. The FeOOH(s) surface is positively charged up to about pH 6, or even higher, so electrostatic attraction of Ni(CN)2− 4 explains in part its adsorption at lower pH values. Electrostatic attraction alone is not sufficient to explain the extent of removal observed, however. Through surface interaction modeling, Theis and Richter [34] demonstrated that Ni(CN)2− 4 must bond at specific surface sites on goethite, in surface complexation reactions that involve high free energies of interaction. They proposed the formation of a goethite–cyanide–metal surface complex via a surface complexation reaction: 2− + ≡Fe2 (OH)2+ 2 + Ni(CN)4 + 2H = ≡Fe2 –(CN)2 –Ni–(CN)2 + 2H2 O

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(5.27)


69

(a) 100 Ni2+

Ni(CN)42–

Percent total Ni

80

60

CNT = 10–4 M

40

20

0 3

4

5

6

7

8

9

10

pH (b) 100 Ni(CN)42–

Ni2+

CNT = 10–3 M

Percent total Ni

80

60

40

20

0 3

4

5

6

7

8

9

10

pH

FIGURE 5.6 Theoretical distribution of nickel in the presence of (a) 10−4 M cyanide (NiT = 10−4.77 M, I = 0.01 M), and (b) 10−3 M cyanide (NiT = 10−4.77 M, I = 0.01 M). (Source: Reprinted with permission from Theis, T.L. and Richer, R.O., Particulates in water, 189, 73, 1980. Copyright 1980 American Chemical Society.)

where ≡Fe2 (OH)2+ 2 is a surface hydroxyl site on the surface of goethite in aqueous suspension, and ≡Fe2 –(CN)2 –Ni–(CN)2 is the surface species formed by adsorption of Ni(CN)2− 4 on the goethite. The uptake of H+ shown in the reaction occurs commonly in adsorption of inorganic anions on oxides, and is related to the commonly observed pH dependence for anion adsorption: maximum adsorption at lower pH and decrease in adsorption with increasing pH [13]. Formation of a metal– ligand–metal ternary surface complex as shown in the reaction of Equation (5.27) has been proposed for other metal–ligand systems [35–38]. Theis and West [8] studied the adsorption of cadmium, copper, and zinc divalent cations and their metal–cyanide complexes on goethite in aqueous suspensions. Some typical results for adsorption of Cd2+ , Cu2+ , and Zn2+ in the absence of cyanide are presented in Figure 5.9. All of the experiments were conducted with total metal concentration of approximately 2 × 10−5 M and with 0.6 g/L

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70

100

90

80

Percent nickel removed

70

60

50

40

30

20

10

0 3

4

5

6

7 pH

8

9

10

11

FIGURE 5.7 Nickel adsorption as a function of pH in the presence of silicon dioxide and Cyanide. NiT = 10−4.77 M, I = 0.01 M, SiO2 = 29.41 g/L. () CNT = 10−5 M, () CNT = 10−4 M, () CNT = 10−3 M, ( ) CNT = 0. (Source: Reprinted with permission from Theis, T.L. and Richer, R.O., Particulates in water, 189, 73, 1980. Copyright 1980 American Chemical Society.)

•

FeOOH(s). The pH adsorption edge plots shown in Figure 5.9 exhibit the typical characteristics for cation adsorption on metal oxides: an increase from 0 to 100% adsorbed with increasing pH. Batch adsorption experiments conducted with free cyanide showed no adsorption of the free cyanide on goethite for any pH from 3 to 11 (data not shown). Experiments with free cyanide added in excess of the metal concentrations were also performed. Results for the three metals are given in Figure 5.10. At lower pH values, adsorption of the cadmium, copper, and zinc was unaffected by the free cyanide as may be seen by comparison with Figure 5.9. Above pH 6.5 to 7.0, however, adsorption of the metals was inhibited by the presence of the cyanide. At the higher pH values, metal–cyanide complexes dominate the speciation of the metals (e.g., see the aqueous phase speciation diagram for zinc in Figure 5.5). The data in Figure 5.10 indicate that the cadmium–, copper–, and zinc–cyanide species have no affinity for the goethite surface at neutral to alkaline pH values. The examples presented in this section demonstrate that some weak metal–cyanide complexes can adsorb on soils and soil components under some conditions, but the extent of adsorption depends strongly on the particular metal–cyanide species, mineral adsorbent, and solution conditions. Solution pH is an especially important governing parameter, as is the case for adsorption of all ions on oxidic minerals [13]. The data presented also demonstrate that the presence of free cyanide in

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100

90

80

Percent nickel removed

70

60

50

40

30

20

10

0 3

4

5

6

7

8

9

10

11

pH

FIGURE 5.8 Nickel adsorption as a function of pH in the presence of iron oxide (goethite) and cyanide. NiT = 10−4.77 M, I = 0.01 M, α-FeOOH = 0.59 g/L. () CNT = 10−5 M, () CNT = 10−4 M, () CNT = 10−3 M, ( ) CNT = 0. (Source: Reprinted with permission from Theis, T.L. and Richer, R.O., Particulates in water, 189, 73, 1980. Copyright 1980 American Chemical Society.)

•

a systems with metals, leading to the formation of metal–cyanide complexes, can result in enhanced or reduced adsorption of the metals. The metal–cyanide complexes may interact with the surface to a greater or lesser extent than the metals alone. An interrelated, complex group of factors governs metal–cyanide species adsorption, and it is diffcult to form generalizations. 5.2.1.4 Oxidation Weak metal–cyanide complexes generally are readily oxidized by oxidizing agents such as chlorine or ozone. The more strongly bonded complexes in the WAD category, such as nickel, silver, and mercury cyanide complexes, oxidize more slowly [15]. The more weakly-bonded complexes, including those of cadmium, copper, and zinc, decompose rapidly in the presence of oxidizing agents. As discussed in Section 5.1.5, and in more detail in Chapter 20, alkaline chlorination is the most common approach used to treat waters bearing free cyanide. A number of weak metal–cyanide complexes are also readily oxidized in this process. In order to identify the fraction of measured total cyanide, which includes metal–cyanide complexes and free cyanide, that is treatable by alkaline

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72

100 90 80 Percent removed

70 60 50

Cd+2 = 2.08 10–5 M Cu+2 = 1.92 10–5 M

40

Zn+2 = 2.41 10–5 M

30

α −FeOOH = 0.6 g/L I = 0.1 M

20 10 0 3

4

5

6

7

8

9

10

11

pH

FIGURE 5.9 Adsorption of Cd, Cu, and Zn on goethite. I = 0.01 M, α-FeOOH = 0.6 g/L. () CuT = 10−4.72 M, () ZnT = 10−4.62 M, () CdT = 10−4.68 M. (Source: Theis, T.L. and West, M.L., Environ. Technol. Lett., 7, 309, 1986.) 100 ZnT = 1.73 10–5 M/ CNT = 2.00 10–3 M CuT = 1.95 10–5 M/ CNT = 1.00 10–4 M CdT = 1.88 10–5 M/ CNT = 2.00 10–3 M

90 80 Percent removed

70 60 a−FeOOH = 0.6 g/L I = 0.1M

50 40 30 20 10 0 3

4

5

6

7

8

9

10

11

pH

FIGURE 5.10 Effect of cyanide on adsorption of Cd, Cu, and Zn on goethite. I = 0.01 M, α-FeOOH = 0.6 g/L. () CdT = 10−4.73 M, CNT = 10−2.70 M; () ZnT = 10−4.76 M, CNT = 10−2.70 M; () CuT = 10−4.71 M, CNT = 10−4.00 M. (Source: Theis, T.L. and West, M.L., Environ. Technol. Lett., 7, 309, 1986.)

chlorination, an analytical measurement known as “cyanide amenable to chlorination” [15] has long been employed. The method involves measurement of total cyanide on samples with and without treatment by chlorination, with the difference giving the amount of cyanide in the sample amenable to chlorination (Chapter 7). While the CATC method has limitations, as discussed in Chapter 7, the existence of the method speaks to the facile oxidation of a number of weak metal–cyanide complexes. In some cases, the presence of weak metal–cyanide complexes can enhance the rate of free cyanide decomposition through catalysis by the metal. This has been demonstrated for copper cyanide complexes [20]. Gurol and Holden [20] studied the effect of copper(I) on the removal of free cyanide by

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ozone in alkaline solution. They performed experiments in solutions at pH 11.5 in systems with an excess of free cyanide over copper(I), giving Cu(CN)3− 4 as the dominant copper species. They found that the presence of copper increased the rate of free cyanide oxidation significantly. Comparison of initial rates of cyanide disappearance for systems with and without copper indicated a fivefold higher rate for the system with copper cyanide species. Further investigations revealed that the observed enhancement was likely due a very fast, independent oxidation–reduction reaction between Cu(I) and free cyanide. The following reaction sequence was proposed: 2− − 2Cu(CN)3− 4 + 2O3 → 2Cu(CN)3 + 2CNO + 2O2

(5.28)

− + 2Cu(CN)2− 3 + O3 + 2H → 2Cu(CN)3 + O2 + H2 O

(5.29)

− 2Cu(CN)− 3 → 2Cu(CN)2 + (CN)2

(5.30)

2Cu(CN)− 2

(5.31)

+ 4CN

−

= 2Cu(CN)3− 4 − −

(CN)2 + 2OH− → CN + CNO + H2 O

(5.32)

The reaction in Equation (5.28) represents the direct oxidation of cyanide to cyanate by ozone. In Equation (5.29), Cu(I) is oxidized to Cu(II) in a fast reaction. The Cu(II) species subsequently oxidizes cyanide to cyanogen (C2 N2 ), being reduced back to Cu(I) in the process (Equation [5.30]). An equilibrium between the copper(I) cyanogen species and Cu(CN)3− 4 is rapidly established (Equation [5.31]). In the last step, Equation (5.32), cyanogen goes through a disproportionation reaction to yield free cyanide and cyanate. The net reaction from the above sequence is thus as follows: 3− − − 2Cu(CN)3− 4 + 3CN + 3O3 → 2Cu(CN)4 + 3CNO + 3O2

(5.33)

Thus, the oxidation of 3 mol free cyanide requires 2 mol of ozone and produces 3 mol of cyanate, as would be expected, but the rate of the reaction is much accelerated due to the presence of the Cu(I).

5.2.2 STRONG METAL–CYANIDE COMPLEXES 5.2.2.1 Formation The cyanide anion can form strong complexes with a number of transition heavy metals, the most notable among them are cobalt, platinum, gold, palladium, and iron. Most of these metals fall in Groups IB, IIB, and VIII of the periodic table. As iron is by far the most abundant of these elements in the environment and in process waters, iron–cyanide complexes are the strong metal–cyanide complexes of greatest interest. Gold–cyanide complexes are of great commercial interest, as the strength of the gold–cyanide bond is exploited in hydrometallurgical gold mining for aqueous extraction of gold from ores (see Chapters 4 and 26). The metal–cyanide bonds in these complexes are mostly arranged in tetrahedral or octahedral forms with strong bonding energy existing between the heavy metal atom and the cyanide ligand [1]. Because they can only dissociate under strongly acidic pH conditions (pH < 2), they are referred to as strong acid dissociable complexes, or simply as strongly-complexed cyanide [15]. As some of these species are formed very slowly [1,33], it is difficult to determine the equilibrium formation constants. Formation data determined by direct thermodynamic methods are available only for complexes of gold(I), and palladium(II) [1]. For other metals, like iron, electron transfer between complex ions of the element in the +2 and +3 oxidation states is rapid enough such that the ratio of the formation constants can be determined from measurement of redox potentials. This ratio can then be combined with standard enthalpy and entropy change measurements for the formation reaction of interest.

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TABLE 5.4 Equilibrium Constants for Formation of Selected Strong Metal–Cyanide Complexes Reaction Ba2+ + Fe2+ + 6CN− = BaFe(CN)2− 6

Ba2+ + Fe3+ + 6CN− = BaFe(CN)− 6 Ca2+ + Fe2+ + 6CN− = CaFe(CN)2− 6 Ca2+ + Fe3+ + 6CN− = CaFe(CN)− 6 2Ca2+ + Fe2+ + 6CN− = Ca2 Fe(CN)06 Ca2+ + H+ + Fe2+ + 6CN− + e− = CaHFe(CN)2− 6 Fe2+ + 6CN− = Fe(CN)4− 6 Fe2+ + H+ + 6CN− = HFe(CN)3− 6 Fe2+ + 2H+ + 6CN− = H2 Fe(CN)2− 6 Fe3+ + 6CN− = Fe(CN)3− 6 2Fe2+ + 6CN− = Fe2 (CN)06 2K+ + Fe2+ + 2H+ + 6CN− = K2 H2 Fe(CN)06 3K+ + Fe2+ + H+ + 6CN− = K3 HFe(CN)06 K+ + Fe2+ + 6CN− = KFe(CN)3− 6 2K+ + Fe2+ + 6CN− = K2 Fe(CN)2− 6 K+ + Fe2+ + H+ + 6CN− = KHFe(CN)2− 6 Li+ + Fe2+ + 6CN− = LiFe(CN)3− 6 2Li+ + Fe2+ + 6CN− = Li2 Fe(CN)2− 6 Li+ + Fe2+ + H+ + 6CN− = LiHFe(CN)2− 6 Mg2+ + Fe3+ + 6CN− = MgFe(CN)− 6 Mg2+ + Fe2+ + 6CN− = MgFe(CN)2− 6 2+ + 6CN− = NH Fe(CN)3− NH+ 4 4 + Fe 6 2+ + 6CN− = (NH ) Fe(CN)2− + Fe 2NH+ 4 2 4 6 + 2+ + 6CN− = NH Fe(CN)2− NH+ 5 4 + H + Fe 6 Na+ + Fe2+ + 6CN− = NaFe(CN)3− 6 2Na+ + Fe2+ + 6CN− = Na2 Fe(CN)2− 6 Na+ + Fe2+ + H+ + 6CN− = NaHFe(CN)2− 6 Sr 2+ + Fe3+ + 6CN− = SrFe(CN)− 6 Tl+ + Fe2+ + 6CN− = TlFe(CN)3− 6

log K (at 25◦ C, I = 0) 49.40 55.44 49.69 55.47 51.00 52.71 45.61 50.00 52.45 52.63 56.98 52.31 50.22 48.12 48.98 51.47 47.69 48.53 51.22 55.39 49.43 48.07 48.87 51.40 47.99 48.74 51.43 55.62 48.75

Source: Data from Sehmel, G.A., Cyanide and antimony thermodynamic database for the aqueous species and solids for the EPA-MINTEQ geochemical code, PNL-6835, Pacific Northwest Laboratory, Richland, WA, 1989, (Table 5).

Table 5.4 lists the equilibrium constants for the reversible formation of iron–cyanide complexes, which are of primary interest with respect to cyanide in the environment. The constants reported are from the compilation by Sehmel [5], which was selected for the reasons discussed in Section 5.2.1.1. Among all the iron–cyanide complexes, the most commonly occurring are ferrocyanide, Fe(CN)4− 6 , 3− where iron is the +2 oxidation state, and ferricyanide, Fe(CN)6 , where iron is in the +3 oxidation state. Another iron–cyanide complex only recently identified and not presented in Table 5.4, with a chemical formula, Fe(CN)5 NHCH4− 3 , has been found to dominate groundwater at many former manufactured gas plant sites [39].

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75

25 20

Water oxidized

15 10 pE

HCN

Fe(CN)6 (tot)

5 0 –5

CN–

–10

Water reduced

–15 0

2

4

6

8

10

pH

FIGURE 5.11 Predominance diagram for dissolved cyanide species in equilibrium with hydrous ferric oxide at T = 25◦ C, as calculated with MINEQL+ (Schecher et al., 1998) using the reactions in Equation (5.1) and Table 5.4. TOTCN = 0.6 mM, TOTFe = 0.5 mM, TOTK = 0.4 mM, TOTNa = 0.06 mM, TOTCl = 0.06 mM, and I = 0.06 M NaCl. (Source: Ghosh, R.S. et al., Water Environ. Res., 71, 1205, 1999.)

In soils and aquifer systems where iron is ubiquitous, the aqueous speciation of cyanide is influenced significantly by reactions with iron dissolved from iron oxides [40,41]. Equilibrium with hydrous ferric oxide, the common amorphous iron oxide, typically is important because Fe(OH)3 (s) serves as the source of iron that becomes dissolved, which in turn regulates the cyanide speciation. Figure 5.11 presents a species predominance diagram for dissolved cyanide species in a system in equilibrium with hydrous ferric oxide. The diagram was calculated with MINEQL+ [32] using the reactions and equilibrium constants in Equation (5.1) and Table 5.4, and in the MINEQL+ thermodynamic database for the iron dissolution, hydrolysis, and redox reactions. In the area denoted “Fe(CN)6 (tot),” cyanide is predicted to exist at equilibrium primarily as the iron cyanide species 4− Fe(CN)3− 6 (oxic conditions) or Fe(CN)6 (anoxic conditions). In the remaining area HCN is the predominant form of dissolved cyanide, except for a small region at pH > 9.2, the pK a for HCN, above which CN− dominates free cyanide speciation. 5.2.2.2 Dissociation As indicated in the equilibrium species predominance diagram of Figure 5.11, iron–cyanide complexes require acidic conditions to dissociate and form free cyanide. It is important to remember, however, that the species distribution shown in Figure 5.11 reflects equilibrium conditions. Actual species distributions for systems with iron present are strongly governed by kinetics. Dissociation of iron–cyanide complexes in the dark is very slow [42]. Like weak metal–cyanide complexes, the dissociation properties of iron–cyanide complexes in aqueous solutions are functions of their stability constants, pH, temperature, and redox potential of the solution. Meeussen et al. [42] ◦ studied the dissociation of ferrocyanide, Fe(CN)4− 6 , in 1 mM solutions in the dark at 15 C. Based on the results, Meeussen et al. [42] projected half-lives ranging from 1 year under reducing conditions (pE ≈ 5) at pH 4 to 1000 years at the same pH under oxidizing conditions (pE ≈ 10). In these laboratory experiments, maximum decompostion rates were observed at low pH and pE. Actual decomposition rates in the environment could be quite different, for example, through enhancement by catalysts. Nevertheless, the results of Meeussen et al. demonstrate the high degree of stability of

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(a)

(b)

k 6

[Fe(CN)6]4–

Ic

k -6

k r

[Fe(CN)5H2O]3– + CN–

kf

[Fe(CN)6]4*

Kd HCN + OH–

approx. 0.02 [Fe(CN)5H2O]2–+ CN–

h [Fe(CN)6]3– hn [Fe(CN)6]3–* approx. 0.98 [Fe2(CN)10]5– ? [Fe2(CN)10]5–

OH– [Fe(CN)5(OH)]3– h hn … [Fe2(CN)10]6–

Fe(OH)3

FIGURE 5.12 Ferro- (a) and ferricyanide (b) photodissociation reaction pathways. (Sources: Information from: Gaspar, V. and Beck, M.T., Polyhedron, 2, 387, 1983; Fuller, M.W., Aust. J. Chem., 39, 1411, 1986.)

iron–cyanide complexes and the importance of considering kinetics when evaluating dissociation of these complexes. Since strong acid conditions are required to dissociate these complexes, the total cyanide analytical method, which is designed to recover strong metal–cyanide complexes in addition to free cyanide and weak metal–cyanide complexes, employs strong acid pH conditions (pH = 1.5) and heat (125◦ C, for 2 h) to achieve dissociation of all metal–cyanide complexes present ([15,43]; Chapter 7). While ferro- and ferricyanide complexes are quite stable in the dark, they can dissociate rapidly when exposed to light [42,44–46]. For example, in experiments with 1.28 µM ferrocyanide solutions at pH 12 exposed to diffuse daylight, Meeussen et al. [42] observed an initial decomposition rate of approximately 8% per hour. Light in the ultraviolet (UV) range (wavelengths less than 420 nm) is responsible for the photolysis of ferro- and ferricyanide [45]. Some proposed photodissociation pathways for ferro- and ferricyanide are shown in Figure 5.12. The photoactivated dissociation rate depends on light intensity, light wavelength, temperature, presence of catalysts, and other parameters [45,47–49], but the results of Meeussen et al. [42] demonstrate clearly that photodissociation of iron–cyanide complexes upon exposure to natural light can be very rapid. Based on photolysis experiments with ferro- and ferricyanide, Broderius and Smith [45] estimated mid-day half-lives (for mid-summer at the latitude and climatic conditions of St. Paul, MN) for 25 to 100 µg/L concentrations of these species to be 18 and 64 min, respectively. Photolytic degradation of ferro- and ferricyanide follows approximately first-order kinetics [45,49], at least initially, but the rate slows as free cyanide accumulates in solution [49]. While some differences in the rates of photolysis of hexacyanoferrate have been observed at different pH values [45,49], Kuhn and Young [49] found no consistent pattern of initial rate coefficient dependence on pH in studies on solutions at pH 4 to 12. The presence of natural organic matter, or other photoreactive substances in water, can significantly decrease the rate of hexacyanoferrate photolysis [49]. The rate of photochemical dissociation in natural waters is dependent on various environmental factors, including free cyanide content of the solution, sunlight intensity, temperature, turbidity, and depth of the water column [45,50]. In many surface waters, significant photolysis will occur only in the top 50 to 100 cm of the water column where sunlight intensity is sufficient, providing opportunity for dilution of any free cyanide produced [45]. Free cyanide could possibly be undetectable or shortlived [51]. There are scenarios, however, where sunlight intensity may be uniform across the entire water column, such as in shallow ponds and in surface runoff. Hexacyanoferrate contamination of the latter can occur, for example, through the spreading of road salt containing iron cyanide as an anticaking agent [52], that is, the commonly used “blue salt.” More details on the photolysis of iron cyanide species and the role of photolysis in fate and transport of hexacyanoferrate in surface waters are provided in Chapter 9. 5.2.2.3 Adsorption on Soil and Sediment Like the weak metal–cyanide complexes (Section 5.2.1.3), strong complexes such as ferrocyanide, 3− Fe(CN)4− 6 , and ferricyanide, Fe(CN)6 , can adsorb on common soil and sediment components

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Percent Fe(CN)–4 6 adsorbed

100

75

50

3.3 mg g–1 1.7 mg g–1

25

0.83 mg g–1 0.50 mg g–1 0 5

6

7

8

9

10

pH

FIGURE 5.13

Ferrocyanide adsorption on alumina at different adsorbate/adsorbent ratios (mg Fe(CN)4− 6 as

CN per g γ-Al2 O3 (s)). 10−5.19 M Fe(CN)4− 6 and various γ-Al2 O3(s) solid doses in 100 mL 0.01 M NaCl solution. (Source: Bushey, J.T. and Dzombak, D.A., J. Coll. Int. Sci., 272, 46, 2004, John Wiley & Sons. Reproduced with permission.)

such as iron, aluminum, and manganese oxides, and clays [8,12,53]. Adsorption of metal–cyanide complexes occurs through a combination of electrostatic attraction and surface complexation [8]. As most strong metal–cyanide complexes are anionic, they can be substantially adsorbed onto soils with high anion exchange capacity. Solution conditions, especially pH, also affect the extent of adsorption of metal–cyanide complexes in aqueous systems. Alessi and Fuller [12] conducted laboratory column mobility tests in which ferricyanide solution was passed through five different soils of varying physical and chemical characteristics. Based on these tests, it was concluded that soil properties, such as low pH (pH < 5), free iron oxide content, and kaolin, chlorite, and gibbsite type clay (high anion exchange capacity) material increased adsorption of iron–cyanide complexes to soil material. Conversely, soils, sediments, and aquifer materials dominated by sand or other components with high cation exchange capacities tend to be weaker adsorbents for iron–cyanide complexes [40]. Mobility tests performed in fixed-bed columns packed with sand-dominated aquifer material and ferrocyanide-contaminated site groundwater as the mobile phase revealed minimal interaction between the dissolved ferrocyanide complexes and site sand [40]. In a mobility test performed by Ghosh et al. [40], ferrocyanide was observed to break through the column in one pore volume, similar to transport of a conservative tracer. It has been demonstrated in a number of studies [8,53,54] that aluminum and iron oxides, two very common and surface-reactive components of soils and sediments [6,55], can adsorb iron–cyanide species significantly, especially at lower pH values (1.5 mg l−1 , followed by gradual decreases for the following 21 days. This increase in free cyanide concentration was accompanied in both systems by a 2–3 unit decrease in pH, reaching values of 5.0–5.5 in three to four days. An investigation of the ability of proprietary aerobic bacterial strains to degrade iron–cyanide complexes was performed with laboratory batch reactors and solutions of potassium ferrocyanide, spent potliner leachate, and groundwater affected by spent potliner leachate [47]. The bacterial strains were added to the batch reactors, in which environmental conditions (oxygen and nutrient levels, pH, temperature) were carefully controlled. For test periods up to 200 h, there was no indication of degradation of potassium ferrocyanide in the synthetic test solution, or of total cyanide in the spent potliner leachate or groundwater. Monitoring of the bacterial populations by enumeration showed that the bacteria were tolerant of the cyanide species present at concentrations from 10 to 1000 mg l−1 , but there was no evidence that the strains could grow on these substrates as a sole source of carbon and energy. While there were clear losses of cyanide from the systems utilized in these studies, it was not possible to demonstrate conclusively that the micro-organisms were directly responsible for the degradation of the hexacyanoferrates. In these studies, parallel efforts with sterile systems generally showed that neither time nor pH contributed substantially to the increase in free cyanide concentration. Nevertheless, free cyanide appears in the media in the presence of bacteria. The mechanism by which this occurs remains a mystery. Unambiguous evidence for direct degradation of hexacyanoferrates, or other metal–cyanide complexes for that matter, is lacking. An alternate explanation may be that bacterial exudates promote dissociation of the complexes and that these bacterial strains degrade the liberated free cyanide. While clarifying this specific question will require additional study, this does not necessarily preclude the use of bacteria in biotreatment systems for metal–cyanide complexes. For example, the treatment system at the Homestake Lead, ND consistently remove 91 to 95% total cyanide (which includes iron and other metal complexes); significant portion of the iron complex was removed by adsorption onto the microbial mass [54]. The wastewater processed also contained significant concentration of thiocyanate [6]. 6.1.3.4 Bacterial Degradation of Thiocyanate Thiocyanate is typically considered a product of the biological transformation of cyanide (Section 6.6.2.2), rather than as a substrate for degradation. Nonetheless, bacterial degradation of thiocyanate has been reported and two possible pathways have been proposed (Table 6.2). Most evidence supports a cyanate pathway, in which thiocyanate is converted first to cyanate, with subsequent formation of ammonia and carbon dioxide. Evidence for this pathway has been obtained from studies of several bacterial species [9,51,55–59], including Thiobacillus thiocyanoxidans and Pseudomonas stutzeri. Most of these studies demonstrated that thiocyanate was used as source of carbon, nitrogen and, in some instances, sulfur. Thiocyanate degradation occurs under a range of physicochemical conditions, including highly alkaline soda lake sediments and soils [9]. A carbonyl pathway has also been reported. In this pathway, thiocyanate is converted directly to ammonia and carbonyl sulfide. This pathway was observed in the chemolithotroph Thiobacillus thioparus. Further efforts allowed for the identification of the enzyme responsible (thiocyanate hydrolase) and genes encoding this enzyme in other thiocyanate-degrading bacterial cultures [60]. Interestingly, this

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hydrolase shows a significant homology to nitrile hydratases, suggesting some similarity between nitrile and thiocyanate degradation.

6.2 FUNGAL DEGRADATION 6.2.1 FUNGI CAPABLE OF DEGRADING CYANOGENIC COMPOUNDS Much of the evidence for fungal degradation of cyanide was obtained from the study of pathenogenic fungi, particularly those that attack cyanogenic plants. The capacity to degrade cyanide evolved in these species, including Stemphylium loti, Gloeocercospora sorghi, Leptosphaeria maculans, and Helminthosporium trucicum, as a means of countering the cyanogenic chemical defenses of plants. The fact that fungal degradation of cyanide is more prevalent in pathogenic fungi than nonpathenogenic [61] supports this contention. Nevertheless, the capacity to degrade cyanogenic compounds has been observed in other fungi as well, including species from the genera Fusarium [62,63], Trichoderma [64], Cryptococcus [21,65], Scytalidium, and Penicillium [11]. There is also a report of cyanide degradation by white rot fungi [66], attributed to the activity of lignin-degrading enzymes. Since this mechanism differs greatly from those identified in the aforementioned fungi and there has been no further support for this pathway, it is not discussed here.

6.2.2 CONDITIONS REQUIRED FOR FUNGAL DEGRADATION OF CYANOGENIC COMPOUNDS Fungal degradation occurs under conditions generally similar to those for bacteria. Degradation is commonly observed at neutral pH, although there are reports of cyanide degradation at pH 4.0 [11] and under alkaline conditions up to pH 10.7 [6,62]. The extent of degradation is lower under alkaline conditions than at neutral pH, most likely due to changes in cyanide speciation. Like the hydrolytic pathways of bacteria, the hydrolytic fungal degradation pathway preferentially utilizes HCN over CN− as a substrate [39]. Cyanide speciation shifts from HCN to CN− as the pH increases, suggesting that the decrease in degradation may be due to substrate availability. Fungal degradation is enhanced by the presence of HCN. The presence of simple sugars and other organic nutrients in the medium enhances fungal degradation, while little to no growth is observed in media without these compounds. Apparently cyanide alone (as the sole source of carbon and nitrogen) cannot support fungal growth [62]. Unlike the bacterial oxygenase pathways, fungi require no addition cofactors as growth and degradation on basal complete media containing cyanide was possible.

6.3 PATHWAYS FOR FUNGAL DEGRADATION OF CYANOGENIC COMPOUNDS The fungal pathways for the degradation of cyanogenic pathways are more limited than those in bacteria and in some instances show mechanistic differences. While fungi show the capacity to degrade free cyanide, nitriles, and perhaps metal–cyanide complexes, this capacity is reportedly limited to a single hydrolytic pathway. Thiocyanate degradation has also been observed and involves the cyanate pathway described earlier for bacteria.

6.3.1 FUNGAL DEGRADATION OF FREE CYANIDE AND NITRILES Fungal degradation of free cyanide is a two-step hydrolytic process, in which cyanide hydratase (also called formamide hydrolase) converts HCN to formamide. An amidase such as formamidase is then responsible for the formation of formic acid and ammonia (Table 6.2). Both steps of the pathway are obligate for cyanide degradation as fungi such as Fusarium solani are unable to grow on either

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formamide or formic acid as the sole sources of carbon and nitrogen [62]. In fact, growth in media supplemented with cyanide and nutrients does not occur until the formamide has been converted to formic acid. Activity of the first enzyme, cyanide hydratase, is induced only in the presence of cyanide [63]. The reaction is irreversible with a pH optimum in the 7.0 to 9.0 range [5,62]. Cyanide hydratase purportedly requires no energetic cofactor and shows little dependence upon nutrients from the media. The cyanide hydratases from F. lateritium and G. sorghi show 65% homology at the gene sequence level and 75% homology at the amino acid level, indicating a high degree of similarity between fungal enzymes. The fungal enzyme shows only 35% homology to the cyanide hydratase from Alcaligenes faecalis [63]. The second reaction is dependent upon nutrients in the media as well as oxygen. When F. solani was grown in minimal media in the presence of cyanide, formamide is the sole product formed. However, when yeast extract was present the formamide was converted to formic acid after a lag period [62]. The rate of the amidase reaction is slower than that of cyanide hydratase. For example, in F. solani the formation of formic acid from formamide was only 25% of the rate of formamide formation from HCN. The identity of this amidase has not been conclusively demonstrated, although it has been suggested to be a formamidase similar to those from A. xylosoxidans and P. putida. Fungal degradation of nitriles has also been demonstrated. Two yeast strains, Cryptococcus sp. UFMG-Y28 and Candida guilliermondii CCT 7207, were able to use nitriles such as benzonitrile as the sole nitrogen source for growth [65]. Both free and immobilized cultures of Candida guilliermondii CCT 7207 degraded aliphatic and aromatic nitriles, although immobilized cultures showed a slightly slower growth and acetic acid formation rate [67]. However, immobilization allowed for degradation under a higher initial nitrile concentration.

6.3.2 FUNGAL DEGRADATION OF METAL–CYANIDE COMPLEXES Evidence for fungal degradation of metal–cyanide complexes is comparable to that for bacteria. That is, there are studies that demonstrate growth of fungal cultures in media containing metal cyanides as the nitrogen source concomitant with losses of cyanide from solution. A mixed fungal culture consisting of Fusarium solani, Trichodera polysporum, Fusarium oxysporum, Scytalidium thermophilum, and Penicillum miczynski isolated from a former gasworks site was capable of using tetracyanonickelate (TCN) and hexacyanoferrate as nitrogen sources at pH 4.0 and 7.0 but not at alkaline pH [10,11]. Similar results were obtained for Fusarium oxysporum N-10 [68] and Cryptococcous humicolus MCNZ [21] when grown on TCN. The results of these studies have suggested a role for cyanide hydratase in the fungal degradation of metal cyanides. In most of these studies, formic acid and ammonia were reported as terminal degradation products, with a formamide intermediate detected in some instances. For the transformation of TCN by F. solani and T. polysporum at pH 7, the rates of ammonia evolution and the rates of TCN dissociation in solution were almost identical (∼1.9 mmol min−1 l−1 ) [10]. Carbon dioxide, most likely arising from the metabolism of formic acid, has also been observed [10], a result consistent with the cyanide hydratase pathway. Degradation of both tetracyanonickelate and hexacyanoferrates was greater at pH 4.0 than at pH 7.0 when heterotrophic fungal cultures were used [11]. Although the metal–cyanide complexes supported fungal growth as nitrogen sources, the rate of growth was only approximately 20% of that observed when HCN was provided as the nitrogen source. Nevertheless, when a mixed culture of F. oxysporum, Sctyalidium thermophilum, and Penicilium miczynski isolated on hexacyanoferrate was grown in the presence of this metal–cyanide complex, >50% loss of cyanide was observed after 28 days of growth. Maximal degradation coincided with maximal growth. Growth of a culture containing F. solani and Trichoderma polysporum on TCN resulted in >90% loss of cyanide in the same time period. When F. solani was grown alone on hexacyanoferrate, there was a 90% loss of cyanide and a 95% loss of iron from the solution in 34 days.

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The conversion of formamide appears to be the rate-limiting step in this process [21]. However, some fungal studies have observed an accumulation of formamide without the accompanying conversion to formate. As the extent to which formate is synthesized by fungi depends upon the culture conditions, there may also be external factors that influence the conversion of formamide. Formation of formate was not observed in washed-cell suspension and cell-free extract systems, but was present in cultures undergoing active growth [21]. Studies showing formation of amidase products were obtained from actively growing cultures [10,68]. Finally, as for bacteria, it appears unlikely that fungi degrade the metal complex itself, but transform the cyanide liberated from the complex as it dissociates in solution. Several factors support this contention. First, the activity of the cyanide hydratase/amidase system is consistent with the fungal mechanism for free cyanide degradation. Since significant structural work has been done on this family of enzymes [39], the ability to accept metal–cyanide complexes substrates appears unlikely. Second, the propensity to degrade metal–cyanide complexes parallels their stability in solution, just as it does for bacteria. Fungal degradation of TCN was more rapid than for hexacyanoferrate. Similarly, measurements of cyanide hydratase and amidase activity were at least tenfold greater for fungi grown on TCN at pH 7.0 than for cultures grown on hexacyanoferrates at pH 4.0. The stability constant for the iron complex is threefold greater the constant for TCN. Furthermore, there was no fungal growth on hexacyanoferrates at pH 7.0, conditions where the dissociation of hexacyanoferrate in solution would be more limited compared to pH 4.0 [69]. Apparent degradation of hexacyanoferrates was also associated with a loss of blue color (Prussian blue) from the nutrient culture [10], which would also be coincident with dissociation of the complex. Nevertheless, little loss of cyanide was observed in sterile systems, indicating that some biological component contributes to the removal of cyanide by fungi in systems containing metal–cyanide complexes.

6.3.3 FUNGAL DEGRADATION OF THIOCYANATE There is limited evidence for the fungal degradation of thiocyanate. Candida tropicalis was identified in a rotating biological contactor used for thiocyanate and phenol treatment, but its specific contribution to degradation was not determined. In another study, Acremonium strictum, isolated from the activated sludge obtained from a wastewater treatment plant for coke-oven-gas condensate, showed >90% degradation of thiocyanate in three days, given an initial concentration of 1.2 g l−1 [55]. Degradation was pH dependent, with an optimum near 6.0, and was influenced by the initial thiocyanate concentration. The maximal rate of degradation was observed at a concentration of 2.1 g l−1 , with the rate decreasing at concentrations >4 g l−1 . Degradation of thiocyanate was inhibited to a limited extent by nitrate and ammonia, but was significantly decreased by elevated concentrations of free cyanide, nitrite, or phenol, again demonstrating the importance of nitrogen sources and cocontaminants in the waste stream on the degradation of cyanogenic compounds. Although the authors of a study with A. strictum concluded otherwise [55], the mechanism of thiocyanate degradation is consistent with the cyanate pathway described in Section 6.1.2.1 for bacteria. The products of thiocyanate degradation by A. strictum were ammonia and sulfate. Although the initial products of the cyanate pathway are cyanate and sulfide, subsequent oxidation of sulfide gives rise to sulfate. Although no cyanate was observed following thiocyanate degradation by A. strictum, there may be several possible reasons for this. Unlike most studies of cyanide degradation, this study was performed upon agar plates, rather than in solution culture. Perhaps these conditions promoted sulfide oxidation by the fungal culture. As there was nearly complete degradation of thiocyanate during experiment, subsequent oxidation to sulfate may have been required to prevent sulfide toxicity. Alternately, there may have been an increased demand for sulfate that may have induced the enzyme responsible for sulfur oxidation in A. strictum. Obviously additional study is required to resolve this question and more thoroughly characterize the mechanism responsible for the fungal degradation of thiocyanate.

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6.4 BACTERIAL ASSIMILATION 6.4.1 BACTERIA CAPABLE OF ASSIMILATING CYANOGENIC COMPOUNDS Bacterial assimilation of cyanide is spread over a wide range of genera, including both Gram-negative and Gram-positive bacteria (Table 6.1). Some of these bacterial species, such as Chromobacterium violaceum and some species of Pseudomonas, are cyanogenic, meaning that cyanide compounds are preferentially synthesized through secondary metabolism (Chapter 3). However, most other bacterial species capable of assimilating cyanide are not cyanogenic. While a primary role of assimilatory pathways is to provide protection from cyanide intoxication, the contribution of cyanide assimilation to primary metabolism is perhaps of greater importance. Several studies have demonstrated that bacteria can utilize cyanide as a source of nitrogen and carbon for growth [7]. Since a wide variety of bacterial, fungal, and plant species synthesize cyanogenic chemicals for release as defensive compounds and deterrents, cyanide in soil and aqueous systems represents a reduced form of carbon and nitrogen that can be more easily and efficiently assimilated that other more oxidized forms of the same elements. It is not surprising then that specific pathways have evolved to provide organisms with the ability to incorporate these compounds directly into primary metabolism. Cyanide assimilation has been studied most extensively in C. violaceum, E. coli, and B. megaterium. Chromobacterium violaceum has been the primary focus [70] principally because this species possess pathways for cyanide both degradation and assimilation [71]. While other species such as Citrobacter freundii and Enterobacter aerogenes may also assimilate cyanide, it has not been clearly established whether this is due to enzymes specific for cyanide assimilation or ancillary activity of other enzymes.

6.4.2 CONDITIONS REQUIRED FOR BACTERIAL ASSIMILATION OF CYANOGENIC COMPOUNDS Unlike the reactions involved in cyanide degradation, evidence of cyanide assimilation has been studied predominantly under in vitro conditions with isolated enzymes. Few studies [71] have attempted to estimate in vivo activity, but the results obtained have been comparable to the in vitro studies. In micro-organisms such as C. violaceum, cyanide assimilation can be detected in vivo under aerobic conditions at the end of the exponential growth phase during periods of active cyanogenesis [71,72]. The activity of enzymes involved in cyanide assimilation, β-cyanoalanine synthase and γ-cyano-α-aminobutyric acid synthase, as well as sulfur transferase enzymes (e.g., rhodanese, 6.6.6.2), tend to increase in cyanogenic bacterial species following periods of cyanogenesis, consistent with their role in preventing cyanide self intoxication [70,72]. β-Cyanoalanine synthase is believed to play the predominant role in the removal of endogenous cyanide produced during highly active periods of cyanogenesis in C. violaceum [71] and in nonmammalian organisms in general [1], primarily because thiocyanate, the product of the reaction catalyzed by rhodanese, is not detected in bacterial cultures [70,71]. The assimilatory reaction mediated by β-cyanoalanine synthase does not require O2 or NAD(P)H [1] and is therefore more energy-efficient than cyanide degradation. Studies of cyanide assimilation by C. violaceum often use growth conditions (addition of glutamate + glycine) that promote cyanogenesis. Growth of C. violaceum on glutamate and glycine induces cyanide production, which consequently results in an increase in the synthesis of enzymes involved in cyanide assimilation [72]. A glutamate + glycine + methionine treatment resulted in a greater removal of cyanide from the media (91%) than cells treated with glutamate alone (62%). Methionine alone has been shown to inhibit cyanide assimilation, as does chloramphenicol, as indicated by a 10% reduction in formation of cyanide derivatives by cells of C. violaceum exposed to chloramphenicol in the presence of glutamate + glycine + methionine.

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Likewise, stationary phase cells of this same species resuspended in solution containing cyanide in conjunction with either glycine or succinate + serine secrete cyanide assimilatory products to the media [73,74]. Evidence also suggests that cysteine plays an important role in cyanide detoxification and perhaps assimilation in bacterial species such as Citrobacter freundii and Enterobacter aerogenes [75]. As discussed in Section 6.4.3.1, cysteine is a cosubstrate for the assimilatory enzymes.

6.4.3 PATHWAYS FOR THE BACTERIAL ASSIMILATION OF CYANOGENIC COMPOUNDS Cyanide assimilation in bacteria is accomplished by substitution reactions carried out by one of three pyridoxal phosphate enzymes: cysteine synthase, β-cyanoalanine synthase, or γ-cyano-αaminobutyric acid synthase. These pathways require a three-carbon skeleton to accept the cyanide ion (CN− ), resulting in the corresponding release of a characteristic ion in exchange. The product of these substitution reactions is the nitrile derivative of an α-amino acid. The production of ninhydrinpositive material (e.g., amino acids) is used as a diagnostic tool to differentiate between cyanide degradation and assimilation [76]. The nitrile produced and the fate of this compound differs between pathways and between bacterial species.

6.4.3.1 Bacterial Assimilation via the β-cyanoalanine Synthase Pathway The β-cyanoalanine synthase pathway is believed to play a primary role in cyanide assimilation in cyanogenic bacterial species like C. violaceum and in acyanogenic species such as B. megaterium and E. coli. The first step in the pathway (Table 6.3) is a pyridoxal phosphate substitution reaction that can reportedly use several potential three-carbon substrates, including cysteine, cystine, serine, O-acetylserine [1,71,73]. This has been observed for enzymes isolated from B. megaterium [77] and Enterobacter strain 10-1 [78,79]. For most bacterial species examined, including E. coli, B. megaterium, and Salmonella typhimurium, cysteine synthase activity is responsible for the formation of β-cyanoalanine. Nevertheless, cyanoalanine synthase (CAS) from C. violaceum has been shown to be the enzyme specifically involved in cyanoalanine synthesis [70]. The subsequent fate of cyanoalanine formed from bacterial CAS and cysteine synthase enzymes has been a matter of conjecture as conflicting reports appear in the literature. In plants, cyanoalanine is converted to the amino acid asparagine (Table 6.3) and subsequently to aspartate with hydrolytic release of ammonia. However, in C. violaceum, amino acids were not detected following the formation of cyanoalanine [71]. Instead, cyanoalanine persisted, which was suggested as evidence that formation of this compound represented a detoxification step, rather than a step toward assimilation. In contrast, most studies have reported that assimilation by C. violaceum, E. coli, and B. megaterium, parallels plant assimilation with asparagine or aspartate detected subsequent to cyanoalanine formation [1,70,71,73,80–83]. The enzyme responsible for this reaction seems to differ by bacterial species. β-Cyanoalanine hydratases are responsible for this reaction in Pseudomonas sp. [84], and perhaps also in B. megaterium and C. violaceum. Chromobacterium violaceum also has an alternate enzymatic system for cyanide assimilation (Section 6.4.3.2). The enzyme from Pseudomonas sp. strain 13, which catalyzes the formation of asparagine and aspartate is a large 100 kDa protein specific for cyanoalanine, yielding the products asparagine and aspartate in a 2.2:1 ratio. A comparable cyanoalanine hydratase has not been characterized from any other bacterial source. In contrast, for Endobacter strain 10-1 and E. coli, this reaction reportedly occurs through ancillary activity of asparaginase, converting cyanoalanine to asparagine or aspartate [76,82].

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TABLE 6.3 The General Categories of Chemical Reactions Responsible for the Assimilation of Free Cyanide by Living Organisms Substitution reactions ␤-Cyanoalanine synthase HS9CH2CH(NH2)COOH + HCN → H2S + NC9CH2CH(NH2)COOH H3C9COO9CH2CH(NH2)COOH + HCN → CH3COO– + NC9CH2CH(NH2)COOH

H2NCO–CH2CH(NH2)COOH

␥-Cyano-␣-aminobutyric acid synthase – 9 S (CH2)2CH(NH2)COOH + CN– → HS9(CH2)2CH(NH2)COOH + NCS9(CH2)2CH(NH2) → NC9(CH2)2CH(NH2)COOH + SCN– Sulfur transferases S2O32– + CN– → SO32– + NCS–

Amino acid synthesis reactions α-Aminobutyric acid synthesis CH3 CH2 CHO + NH3 + HCN → CH3 CH2 CH(NH2 )–CN → CH3 CH2 CH(NH2 )COOH Alanine synthesis + CH3 CHO + NH+ 4 + HCN → CH3 CH(NH2 )–CN + 2 H2 O → CH3 CH(NH2 )COOH + NH4 Glutamate synthesis HOOC(CH2 )2 CHO + NH+ 4 + HCN → CH3 CH2 CH(NH2 )–CN + 2 H2 O → HOOC(CH2 )2 CH(NH2 )COOH + NH3

For the β-cyanoalanine pathway, there are two possible substrates, which can accept free cyanide. The synthesis of γ-cyanoα-aminobutyric acid from free cyanide uses the ionized form (− S–) of homocysteine. This reaction also forms thiocyanate (SCN− ), as do reactions mediated by sulfur transferase enzymes. As α-aminobutyric acid is a possible amino acid precursor, this reaction is included here.

6.4.3.2 Bacterial Assimilation via γ-Cyano-α-aminobutyric Acid Synthase An alternate pathway for cyanide assimilation is apparently unique to C. violaceum [85]. This pathway, which also requires pyridoxal phosphate, operates in parallel with the β-cyanoalanine synthase and is induced by similar growth conditions, such as growth on media containing glutamate + glycine [74]. As mentioned earlier, this pathway represents only a minor pathway for assimilation, with the γ-cyano-α-aminobutyric acid possibly serving as a precursor for the synthesis of amino acids such as glutamate. The synthesis of γ-cyano-α-aminobutyric acid from cyanide occurs via a two-step reaction (Table 6.3), the first of which is nonenzymatic and the second catalyzed by γ-cyano-α-aminobutyric acid synthase. The second reaction is fundamentally an enzymatic replacement reaction for thiocyanate (SCN). The available data demonstrates that γ-cyano-α-aminobutyric acid synthase is highly specific in its enzymatic activity. This enzyme shows no comparable activity as a homoserine sulfydrylase or OAS homoserine sulfhydrylase synthesizes cystathionine at a low rate (8%) in comparison to the formation of γ-cyano-α-aminobutyric acid. The importance of this amino acid derivative and this pathway have not been established.

6.5 FUNGAL ASSIMILATION 6.5.1 FUNGI CAPABLE OF DEGRADING CYANOGENIC COMPOUNDS Fungal assimilation of cyanide has a more limited distribution than is observed in bacteria and plants. Assimilation has been observed in as few as five genera, representing only eight species in total. Three species capable of assimilating cyanide appear in the genus Pholiota, and include P. adipose,

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P. aurivella, and P. praecox. Another species, identified as a psychrophilic basidiomycete and the pathogen responsible for winter crown rot in forage crops, is also capable of cyanide assimilation [86]. This species may be a snow mold fungus, perhaps Fusarium nivale. Previous studies have shown F. nivale to be cyanogenic [87], as well as a species capable of cyanide assimilation. Two other genera that reportedly assimilate cyanide, Rhisopus and Marasmius, are also clearly cyanogenic. Rhizopus oryzae, a fungus associated with cassava spoilage, reportedly detoxifies cyanide with extracellular rhodanese [88], but the contribution of this process to cyanide assimilation is unclear and may represent detoxification rather than assimilation.

6.5.2 CONDITIONS REQUIRED FOR THE FUNGAL ASSIMILATION OF CYANOGENIC COMPOUNDS Fungal assimilation of cyanide has been observed in cultures grown in standard media formulations at both room temperature and decreased temperature (13◦ C). Synthetic media have been used in most cases [86,89], with potato dextrose used in one case where synthetic media failed to support growth of specific fungal strain [90]. Similar to bacteria, assimilation is observed when specific substrates are present in the media when cyanide is present. For fungi, assimilation has been observed when ammonia and acetaldehyde are supplied simultaneously with cyanide [91], or in the presence of ammonia and succinic semialdehyde [92]. The additional substrates are required for the specific cyanide assimilatory pathways present in fungi, both of which lead to the synthesis of specific amino acids.

6.5.3 PATHWAYS FOR THE FUNGAL ASSIMILATION OF CYANOGENIC COMPOUNDS As with plants and bacteria, the fungal pathways for cyanide assimilation are associated with amino acid metabolism. One of these pathways, present in F. nivale, is comparable to the β-cyanoalanine pathway in that the product is asparagine. Little is known about this pathway other than the product formed, so this pathway will not be discussed in detail here. Previous studies have shown that β-cyanoalanine synthase is not the enzyme responsible for cyanide detoxification in fungal species such as Stemphylium loti [93]. A second pathway, similar to that found in bacteria, has been described in Rhizoctonia solani and involves an ammonia-dependent conversion of cyanide to α-aminobutyric acid [89]. This product of this pathway reportedly differs from the corresponding bacterial pathway in the carbon atom onto which the cyanide moiety is substituted and the substrates required. Most fungi capable of assimilating cyanide utilize a second pathway leading to the formation of alanine, including the aforementioned psychrophilic basidomycete. This species can also form glutamic acid via a similar pathway. Production of both amino acids in response to cyanide exposure has been reported in one study [86]. 6.5.3.1 Fungal Assimilation via α-Aminobutyric Acid Synthesis Rhizoctonia solani is a pathogenic fungus that causes root rot in crop plants. In the presence of cyanide and ammonia, this species forms the intermediate α-aminobutyronitrile from proprionaldehyde (Table 6.3). The proprionaldehyde is believed to arise from the nonoxidative decarboxylation of α-ketobutyrate [94]. Studies of cyanide assimilation in liquid cultures of R. solani at room temperature demonstrated rapid formation of α-aminobutyronitrile within 30 min, followed by the appearance of α-aminobutyric acid with the next 30 min. This observation was interpreted as indicative of a precursor–product relationship [89]. This second step requires water, resulting in an enzymatic hydrolysis to product within the regeneration of one molecule of ammonia. Presumably, the resulting α-aminobutyric acid can be used to synthesize amino acids. The α-aminobutyric acid formed differs from that produced by the corresponding bacterial pathway in that the cyanide is added to the aldehyde carbon atom along with the ammonia

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molecule. The subsequent hydrolysis reaction converts the cyanide group to a carboxyl group, resulting in the complete assimilation of cyanide in these two steps. In the bacterial pathway, the cyanide is incorporated into α-aminobutyric acid via substitution onto a sulfhydryl group at the terminal carbon on the opposite end of the molecule. The cyanide molecule persists on the product γ-cyano-α-aminobutyric acid. 6.5.3.2 Alanine and Glutamate Synthesis The pathways for alanine and glutamate synthesis in a previously unidentified physchrophilic basidiomycete are similar to the α-aminobutyric acid pathway above in that ammonia is required as a precursor and the second hydrolytic step regenerates ammonia. The pathways in this basidiomycete also lead directly to amino acids, rather than an amino acid precursor. Formation of alanine utilizes acetaldehyde as the precursor (Table 6.3). The ubiquity of acetaldehyde as an intermediate in cellular metabolism perhaps explains why the majority of fungal species capable of cyanide assimilation produce alanine following exposure to cyanide [91]. Little kinetic information has been reported for this proposed reaction. The synthesis of alanine by this pathway occurs at a somewhat slower rate than α-aminobutyric acid. Fungal cultures exposed to cyanide showed a rapid increase in α-aminopropionitrile during the 12 h following exposure, followed by a subsequent decrease over the next 12 h to initial levels. In contrast, alanine levels increase steadily over the same 24 h period. While there is a linear relationship between formation of αaminopropionitrile formation and enzyme (as total protein) concentration, α-aminopropionitrile also formed nonenzymatically when the protein extract was omitted from the reaction [91]. The enzyme responsible for the formation of alanine has not been described, but the reaction is reportedly similar to a plant nitrilase reaction. The initial studies of cyanide assimilation in this fungal species demonstrated initial formation of both alanine and lower levels of glutamate [86]. This observation led to the hypothesis that the glutamate was formed from alanine [91]. Subsequent study revealed, however, that succinic semialdehyde can be converted to 4-amino-4-cyanobutyric acid in the presence of ammonia and cyanide (Table 6.3). As for the formation of alanine, crude isolated protein extracts of the fungal mycelia showed evidence of nitrilase activity [92]. Nitrilase activity showed a linear increase with time and enzymes concentration at room temperature, with a pH optimum of 8.0. Similar protein extracts also possessed other enzymatic activity, including glutamic acid decarboxylase, succinic semialdehyde dehydrogenase, and γ-aminobutyric acid transaminase. Based upon these various activities, a pathway was proposed that relates formation of 4-amino-4-cyanobutyric acid to basic elements of fungal carbon (succinic semialdehyde) and nitrogen (glutamate) metabolism (Figure 6.1).

6.6 BOTANICAL TRANSFORMATION PROCESSES Cyanide is involved in a wide variety of plant pathways (Chapter 3). Free cyanide (CN− ) is produced concomitantly with the gaseous plant hormone ethylene and is the active constituent of the cyanogenic glycosides used to deter herbivory and store nitrogen, and may regulate aspects of nitrogen metabolism. Another cyanide compound, cyanate (CNO− ), is produced from the decomposition of urea and other compounds. Not surprisingly, there are corresponding pathways that control endogenous levels of these and other cyanide compounds in plant tissues, preventing self-intoxication. Despite periods of active cyanogenesis, concentrations of free cyanide are typically maintained at low levels in plants [95]. There are two primary pathways that contribute to the assimilation of cyanide compounds. Plants appear not to have pathways for cyanide degradation. Free cyanide (HCN, CN− ) is removed by either the cyanoalanine pathway (Section 6.4.3.1) or the sulfur transferase pathway (Section 6.6.2.2). The presence of a plant sulfur transferase pathway in plants similar to that of rhodanese in animals was long disputed but was recently confirmed when genes encoding proteins

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Succinic semialdehyde NH4+ Ketoglutaric acid HCN

gAminobutyric acid 4-Amino 4-cyanobutyric acid

2H2O

CO2

Glutamic acid

NH4+

FIGURE 6.1 A biochemical model illustrating the assimilation of free cyanide into 4-amino-4-cyanobutyric acid by a psychrophilic basidiomycete (fungi). The 4-amino-4-cyanobutyric acid is converted to glutamate by the action of a potentially unique nitrilase. (Source: Strobel, G.A., J. Biol. Chem., 242, 3265, 1967. With permission.)

with sulfur transferase activity were identified. The specific role of the sulfur transferase pathway and cyanase in plant metabolism is still largely a matter of conjecture. The relative importance of these pathways with respect to one another and to cyanide assimilation/detoxification in plants is also unknown. The interaction between these pathways has only been examined in insects, where the cyanoalanine pathway has been shown to play the primary role [96]. Other enzymes, such as cyanase and cyanamide hydratase, contribute to the metabolism of other natural and anthropogenic cyanide compounds. The role of these enzymes in plant metabolism is not clear, although roles for cyanase in primary carbon metabolism have been proposed [29].

6.6.1 CYANIDE-RESISTANT RESPIRATION Cyanide disrupts numerous metabolic processes, principally by binding to the metal cofactors in enzymes, although reactions with functional groups such as carbonyls or disulfide bonds also occur. The most detrimental inhibition is blockage of the mitochondrial electron transport chain, through complexation of cyanide with the iron (Fe3+ ) in the terminal cytochrome, cytochrome c oxidase. In animals, exposure to cyanide can quickly decrease respiration rates to 500 mg/l)

Yes

Yes

Perform standard extraction with iso-octane/ hexane/ methylene chloride

Add Ca(OH)2 raise pH between 12 and 12.5

Perform cyanide analysis

FIGURE 7.3 Cyanide spot testing and pretreatment decision diagram.

rate faster at higher pH values [8–11]. Prior to pH preservation, prolonged exposure to sunlight can result in volatilization losses of photo-liberated HCN, thus reducing the concentration of total cyanide and metal–cyanide complexes. In pH-preserved samples, on the other hand, photodecomposition can result in a positive bias in free cyanide concentration. Care should be taken to minimize any light exposure during all procedural steps from sample collection to distillation and analysis. Figure 7.3 provides a decision diagram regarding spot testing and pretreatment. This chart provides guidance for when and how to pretreat samples prior to analysis.

7.2 TOTAL CYANIDE MEASUREMENT TECHNIQUES Operationally, the term “total cyanide” refers to all cyanide-containing groups that can be collectively measured as the cyanide ion, CN− , after some treatment steps to liberate the cyanide ion [3].

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Analysis of Cyanide in Water

Cooling water

129

In Out

To vacuum Condenser

Inlet tube

Caustic absorber

Distilling flask

Heater

FIGURE 7.4 Manual distillation apparatus for total cyanide test.

Total cyanide measurement can either be performed manually or by using automated techniques. As indicated in Figure 7.2 and Table 7.2, the total cyanide measurement does not capture all species bearing the −C≡N group. Further detail about the various total cyanide measurement techniques is provided below.

7.2.1 APHA/AWWA/WEF (STANDARD METHODS), ASTM, USEPA, AND USGS METHODS The most common total cyanide measurement is via manual distillation followed by colorimetric, titrimetric, or electrochemical finish techniques to measure cyanide ion concentration. The various analytical finish techniques are described in the following subsection. Figure 7.4 shows the total cyanide apparatus employed in the manual distillation procedure. The manual distillation technique is approved by all U.S. regulatory, government, and consensus organizations (see Table 7.2 for method numbers) and measures the free cyanide and metal-complexed (including both weak acid and strong acid dissociable) forms of inorganic cyanide present in an aqueous sample. The total cyanide method does not measure the following cyanide-related compounds: cyanates, thiocyanates, most organic-cyanide compounds, and most cobalt and platinum cyanide complexes. The detection limit in reagent water matrix is in the range of 1 to 5 µg/l. The conventional total cyanide method involves prolonged distillation of the sample at 125◦ C under strongly acidic conditions (pH < 2), which breaks apart most strong and weak

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metal–cyanide complexes and liberates free cyanide. Most organocyanides and some extremely strong metal–cyanide complexes such as cobalt cyanide do not fully decompose under the total cyanide distillation conditions. Additional sample predistillation treatment by ultraviolet radiation is required to obtain recovery of cobalt-cyanide complex. If incomplete recovery of cobalt cyanide occurs even after UV pretreatment followed by distillation, a second distillation is recommended [3]. The free cyanide present or generated by decomposition of complexes is converted to HCN under the acidic conditions and removed from solution by distillation as HCN(g). The distillate is passed through an alkaline absorber solution, which traps the HCN and converts it to the form of CN− . Final determination of CN− in the absorber solution is then performed via colorimetric, titrimetric, or cyanide ion-selective electrode methods. Standard Methods 4500-CN-C, Total Cyanide After Distillation [3]. Hydrogen cyanide (HCN) is liberated from an acidified sample (with sulfuric acid, H2 SO4 ) via a 2-h distillation and purging with air. HCN in the distillate is collected by passing the distillate gas through an NaOH scrubbing solution and then analyzed by any of the three finishing procedures, titration, colorimetric, or selective ion. The colorimetric analysis employs pyridine–barbituric acid for color development. This total cyanide method has been approved by USEPA for drinking water and NPDES compliance testing. ASTM Method D 2036-98, Standard Test Methods for Cyanides in Water, Method A — Total Cyanides After Distillation. This ASTM method [12] is technically similar to the Standard Methods total cyanide method 4500-CN-C except that it employs a 1-h reflux distillation compared to 2-h distillation for Standard Methods 4500-CN-C. Cyanide concentration can be quantified by either the titration, colorimetric, or selective ion electrode procedure. The method is approved by the USEPA for total cyanide determination. USGS Method I-4302-85, Total Cyanide, Colorimetric with Pyridine–Barbituric Acid. This method [13] is technically similar to Standard Methods 4500-CN-C and ASTM Method D 2036-98. USGS Method I-3300-85, Total Cyanide,Colorimetric with Pyridine–Pyrazolone. This method [14] is technically similar to ASTM Method D 2036-98 except that it utilizes pyridine–pyrazolone reagent for color development in the colorimetric procedure compared to pyridine–barbituric acid reagent used for ASTM Method D2036-98 and Standard Methods 4500-CN-C. This is also a USEPAapproved total cyanide analytical method. The above four methods, all use manual distillation with acidic reflux for the decomposition of complex cyanides prior to manual finish techniques for quantification of liberated cyanide. There are several semi-automated and automated analytical methods for the determination of total cyanide that are used primarily in municipal and industrial wastewater treatment plants. Some of the automated methods are presented below. USEPA Method 335.4, Determination of Total Cyanide by Semi-Automated Colorimetry. This USEPA method [15] utilizes a manual reflux-distillation operation to release hydrocyanic acid (HCN) from cyanide complexes, which is then absorbed in a scrubber containing sodium hydroxide solution. The cyanide ion in the absorbing solution is determined colorimetrically by automated, continuousflow analysis equipment designed to deliver and mix sample and reagents in the required order and ratios. USEPA Method 335.3, Total Cyanide by Colorimetric and Automated UV. This USEPA method [16] utilizes an automated UV digestion and distillation unit to decompose cyanide complexes and release HCN. Cyanides are determined automatically by a colorimeter and a recorder. This method was withdrawn for use in drinking water analysis by USEPA in 1994 [17] due to the concern about incomplete UV digestion, but still remains approved for reporting cyanide concentrations as required by NPDES permits. Total Cyanide by Low-Power UV Digestion Method. This is an automated analysis method implemented using an Skalar SAN plus segmented flow analyzer (model SA2001) with an SA1050 random access sampler, an SA 5570 in-line distillation unit, and an SA 555 UV-B inline digester

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(Skalar Analytical B.V., The Netherlands). The method is technically equivalent to the USEPA Method 335.3. Cyanide is released as HCN from cyanide complexes by means of UV digestion and distillation. HCN is then determined by a colorimetric procedure as described previously. ASTM Method D 4374-00, Standard Test Methods for Cyanides in Water, Automated Methods for Total Cyanide, Dissociable Cyanide, and Thiocyanate. This method [18] utilizes alkaline UV irradiation, acidification, and thin film distillation for cyanide-containing samples in an automated system. The breakdown of the strong metal–cyanide complexes, prior to the thin film distillation, is achieved by UV irradiation. Absorption of the liberated HCN gas is carried out using a glass coil and NaOH solution. It also employs the standard colorimetric determination of the recovered cyanides with an automated colorimeter. The method was developed and is employed routinely by the Municipal Water Reclamation District of Greater Chicago [19].

7.2.2 ANALYTICAL FINISH TECHNIQUES The three common analytical finish techniques used for final quantitation of cyanide ion liberated from the samples are titrimetric, colorimetric, and cyanide ion-selective electrode methods. Both the titrimetric and colorimetric procedures are approved by USEPA, Standard Methods (APHA / AWWA/ WEF), and ASTM; the cyanide ion-selective electrode method is approved only by Standard Methods. This section presents brief descriptions of each of the quantitation methods. Also discussed is amperometric measurement of cyanide ion, which is employed in the available cyanide by ligand displacement method. Titrimetric Procedure (Standard Methods 4500-CN-D, ASTM D2036-98). In this procedure [3,12], cyanide ion from the alkaline absorber solution following distillation is titrated with silver nitrate standard solution to form a soluble cyanide complex, Ag(CN)− 2 . Any presence of excess Ag+ in the solution is detected by a rhodanine indicator, which immediately turns from yellow to a salmon hue color indicating the endpoint. The concentration of cyanide in the absorber solution is calculated from the titrant normality and the volume of titrant used to reach the color endpoint. Typical calibration range is between 0.1 and 10 mg/l. Colorimetric Procedure (Standard Methods 4500-CN-E, ASTM D2036-98). In this procedure [3,12], the most commonly employed, cyanide ion from the alkaline absorber solution following distillation is converted to CNCl by reaction with chloramine-T at pH < 8. Following the formation of CNCl, pyridine–barbituric acid is added to the solution, which converts CNCl to a red–blue complex. The colored complex exhibits a stable absorption maximum at 578 nm. The concentration of cyanide in the absorber solution is determined spectrophotometrically by comparison against a standard calibration curve of absorbance vs. concentration. Typical calibration range is between 0.02 and 0.2 mg/l. Because of its detection sensitivity and broad calibration range, the colorimetric procedure is the analytical finish technique of choice for most cyanide analysis methods. Cyanide Ion-Selective Electrode (Standard Methods 4500-CN-F, ASTM D2036-98). The ionselective electrode is one of the few methods [3,12] that measures individual cyanide species, in this case CN− at high pH. In this procedure, cyanide ion from the alkaline absorber solution following distillation is determined potentiometrically using a CN− ion-selective electrode and associated meter. The concentration of cyanide ion in the absorber solution is determined by comparison against a standard calibration plot of log concentration of cyanide versus potential (mV). The cyanide electrode method is less sensitive in terms of detection than the colorimetric method and subject to numerous interferences. For these reasons, it is the least used of all the analytical finish techniques. Typical calibration range is between 0.05 and 10 mg/l. Amperometry (USEPA Method OIA-1677). In the “available cyanide by ligand displacement” method for measurement of free and weakly complexed cyanide [20,21], discussed in Section 7.3.2, liberated cyanide is isolated by use of a membrane, and the amount of cyanide ion collected is

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measured by amperometry. The cyanide isolation is achieved by acidifying the sample and allowing HCN to pass through a gas diffusion membrane. HCN is captured in a sodium hydroxide acceptor solution and thus converted to CN− ion. The use of a membrane selective for HCN makes it possible to employ a nonselective electrical conductivity measurement for quantitation of the amount of cyanide ion in the NaOH acceptor solution.

7.3 WEAK METAL–CYANIDE COMPLEXES There are three well-known analytical methods that render cumulative measurements of weak metal– cyanide complexes. These are (i) the weak acid dissociable cyanide (WAD) method; (ii) the available cyanide by ligand displacement method; and (iii) the cyanide amenable to chlorination (CATC) method. All of these analytical methods measure free cyanide in addition to the weak metal–cyanide complexes.

7.3.1 WEAK ACID DISSOCIABLE (WAD) CYANIDE METHOD The weak acid dissociable method is approved by ASTM [12] and Standard Methods [3] and involves distillation of the sample under slightly acidified (pH 4.5 to 6.0) conditions. This method does not recover CN− from strong metal–cyanide complexes, as indicated in Table 7.2. The most commonly used weak acid dissociable cyanide methods are Standard Methods 4500-CN-I “WAD Cyanide by Distillation” and ASTM Method D 2036-98, “Standard Test Methods for Cyanides in Water, Method C — Weak Acid Dissociable Cyanides.” In addition, there is an automated WAD cyanide method, ASTM Method D 4374-00, “Automated Method for Dissociable Cyanide by Thin Film Distillation [18].” Standard Methods 4500-CN-I, Weak Acid Dissociable Cyanide. Hydrogen cyanide (HCN) is liberated from a slightly acidified sample (acetate buffer, pH = 4.5) via a 2-h distillation and purging with air. HCN in the distillate is collected by passing the distillate gas through an NaOH scrubbing solution and then analyzed by any of the three finishing procedures, titration, colorimetric, or selective ion electrode (see Section 7.2.2 for analytical finish techniques). Figure 7.4 shows the instrument set-up for this method. The WAD cyanide method [3] is not yet approved by USEPA for drinking water or NPDES compliance testing, but has gained acceptance in several states (e.g., Pennsylvania and Texas). The WAD cyanide method has been observed to be less prone to interferences than the CATC Method [22,23] , which is a USEPA-approved method. The detection limit for this method is usually 1 to 5 µg/l for the colorimetric finish procedure. Zinc acetate buffer is used, prior to distillation, in the method to precipitate iron-cyanide complexes and enhance the selectivity of the method. However, for samples dominated by iron-cyanide complexes (>50%), an intermediate filtration step (using 0.45 µm filter) following zinc acetate buffer addition and prior to distillation is desirable to remove the precipitated zinc–iron-cyanide complexes. Otherwise, they can redissolve and dissociate under the conditions of the distillation and create a positive bias [24]. ASTM Method D 2036-98, Standard Test Methods for Cyanides in Water, Method C — Weak Acid Dissociable Cyanides. This ASTM method is technically similar to the Standard Methods weak acid dissociable cyanide method 4500-CN-I, except that it employs a 1-h reflux distillation in contrast to 2-h distillation for Standard Methods 4500-CN-I. Cyanide determination can be conducted colorimetrically, titrimetrically, or by cyanide ion-selective electrode. ASTM Method D 4374-00, Standard Test Methods for Cyanides in Water — Automated Methods for Total Cyanide, Dissociable Cyanide, and Thiocyanate. This is an automated method capable of detecting total cyanide, WAD cyanide, and thiocyanate [18]. For determination of WAD cyanide, the sample is acidified to pH 4.5, and exposed to a continuous thin film distillation unit to liberate HCN from the sample. Absorption of the liberated HCN gas is carried out using a glass coil and an NaOH

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Acceptor

Gas diffusion cell

Carrier

Acid

FIGURE 7.5 1999.)

Detector

Injection valve Pump

Waste

Mixing coil

Waste

Schematic for available cyanide Method OIA-1677. (Source: USEPA, Method OIA-1677,

trap solution. Cyanide determination is conducted colorimetrically at 578 nm by pyridine–barbituric acid reagent, as described in Section 7.2.2.

7.3.2 AVAILABLE CYANIDE BY LIGAND DISPLACEMENT METHOD USEPA and ASTM approved the available cyanide by ligand exchange method developed by OI Analytical (Wilsonville, OR), Method OIA-1677 [20,21,25], which measures mercury-cyanide complexes in addition to all the conventional weak acid dissociable complexes (i.e., cyanide complexes with Cu, Ni, Zn, Cd, Ag) and free cyanide. The list of analytes captured in the method is provided in Table 7.2. Researchers at the University of Nevada, Reno [26], in association with ALPKEM (a division of OI Analytical), developed this rapid, distillationless, flow-injection ligand-exchange method to determine available cyanide. The method consists of two parts: sample pretreatment, followed by cyanide quantification using amperometric detection. In the sample pretreatment step, ligand-exchange reagents are added to displace the cyanide ions from weak and intermediate strength metal–cyanide complexes. In the cyanide quantification step, a portion of the aliquot of the pretreated sample is injected into the flow injection manifold. The addition of hydrochloric acid converts cyanide ion to HCN, which diffuses through a membrane into an alkaline trap where it is reconverted to CN− . The alkaline trap solution is then analyzed for CN− amperometrically using a silver working electrode, a silver/silver chloride reference electrode, and a platinum counter electrode. Figure 7.5 shows a diagram of the flow injection system employed in the available cyanide method. The method detection limit in reagent water is approximately 0.5 µg/l. This method is particularly sensitive to any amount of sulfide in the solution, which is a positive interferent [23,27]. It is essential to treat samples with PbCO3 to remove sulfide prior to preservation when available cyanide analysis is to be performed. Section 7.1 should be consulted for information about the interferences and pretreatment steps required for any weak metal–cyanide complex analysis.

7.3.3 CYANIDE AMENABLE TO CHLORINATION This analysis method is applicable to those cyanide complexes and species that are “amenable to chlorination,” that is, upon chlorination, the cyanide complexes are dissociated and the liberated free cyanide is destroyed. In general, all weak metal–cyanide complexes (i.e., cyanide complexes of Cu, Ni, Zn, Ag, and Cd), are amenable to chlorination. Cyanides amenable to chlorination (CATC) are measured by separating and analyzing total cyanide in two sample portions: one portion chlorinated, and the other left as is and not chlorinated. Both sample portions are then subjected to the total cyanide analysis procedure. The difference in the CN− concentrations between the chlorinated and unchlorinated portions is designated as the cyanide amenable to chlorination. The CATC method has serious limitations owing to the fact that some organic compounds, including compounds that do not contain the −CN group as well as organocyanide compounds,

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can react to form free cyanide during chlorination, giving higher results for cyanide after chlorination than before chlorination [2,5,28,29]. This may lead to a negative value for the calculated concentration of cyanide amenable to chlorination, a problem that has often been encountered with steel industry samples, petroleum refinery distillate, and paper and pulp industry samples [3]. Also, the cyanide amenable to chlorination method has exhibited only partial recovery of some weak metal–cyanide complexes (e.g., nickel, silver, and mercury-cyanide complexes) in some matrices [26]. For the CATC method employing colorimetric finish, the detection limit of this method is usually 5 µg/l in reagent water matrix.

7.4 METAL–CYANIDE COMPLEX ANALYSIS BY LIQUID CHROMATOGRAPHY Various liquid chromatography techniques have been developed for the separation, identification, and quantification of the metal–cyanide complexes in water samples. Ion chromatography, which employs ion exchange resins for separation of metal–cyanide complexes, is the most common technique employed. Reversed-phase ion-pair partition chromatography, involving a nonpolar adsorbent, has also been used. Otu et al. [30] provide a review of the liquid chromatography techniques that have been developed for analysis of metal cyanides in water. A summary of the techniques is provided here.

7.4.1 ION EXCHANGE CHROMATOGRAPHY An ion chromatography method was developed by Dionex Corporation to measure specific metalcyanide complexes (i.e., cyanide complexes with Ag, Cu, Au, Ni, Fe, and Co) at mg/l levels of detection. The method may also be applicable for determining additional metal–cyanide complexes, such as platinum, and palladium cyanide complexes. This method has been approved by the USEPA [31] and by ASTM [32]. The original Dionex method has been modified to include both µg/l and mg/l level detection capability for the six target metal–cyanide complexes in that method [31,32], listed in Table 7.2. In aqueous solution, cyanide forms relatively stable anionic coordination complexes with most transition metals of the form [M(CN)x ]n− (M = the transition metal, x = the number of cyanide groups, and n = the electronic charge of the complex). Due to the stable nature of these complexes, they can be separated using anion exchange [31–33]. Following separation, detection is typically accomplished via low-wavelength ultraviolet light absorption at 215 nm. Figure 7.6 shows the ion chromatography instrumentation set-up for measurement of mg/l concentration levels of metal-cyanide complexes (high-level method). Determination of µg/l concentrations may also be accomplished using an automated online sample preconcentration [31,32]. Figure 7.7 presents typical chromatograms obtained for waters spiked with the six pertinent complexes. Method detection limits and calibration ranges for each metal–cyanide complex species are presented in Table 7.3. Further details about this method are available in ASTM [32]. The major interferences to this method are similar to those for the total and WAD cyanide methods, but there are some different considerations in analysis of metal–cyanide complexes by ion chromatography. Primary interferences and the corresponding pretreatments are discussed below. Photodecomposition. Some metal-cyanide complexes can photodissociate, resulting in decreases in their concentrations. Light exposure during sample collection, preparation, and analysis should be prevented, as much as possible. This applies to analytical samples as well as standards. Following acquisition, samples should be treated for any other interference (e.g., chlorine) and stored in amber bottles with preservatives. Chemical interferences. As in all ion chromatographic methods, certain chemicals can interfere with the analysis of metal–cyanide complexes and reduce the performance of the method, in terms of column performance and peak resolution. This becomes especially important when performing

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Anion trap column

Gradient pump

Autosampler Guard column

Separator column

Data analysis

UV-VIS detector

FIGURE 7.6 Ion chromatography instrument set-up for determination of metal–cyanide complexes. (Source: ASTM, Designation D 6994-04, Annual Book of ASTM Standards, Vol. 11.02, 2004. Copyright ASTM INTERNATIONAL. Reprinted with permission.)

[Au(CN)2]–

[Ag(CN)2]–

0.0

4.0

8.0

[Co(CN)6]3– [Cu(CN)3]2–

[Fe(CN)6]4–

12.0 Min.

16.0

Absorbance units

Absorbance units

[Cu(CN)3]2– [Ni(CN)4]2–

20.0

[Ag(CN)2]–

[Au(CN)2]–

0.0

4.0

8.0

[Co(CN)6]3– [Ni(CN)4]2– [Fe(CN)6]4–

12.0 Min.

16.0

20.0

FIGURE 7.7 Ion chromatograms from analysis of metal–cyanide complexes in reagent water (left) and groundwater sample (right). Metal–cyanide species concentrations are in the mg/l range. (Source: ASTM, Designation D 6994-04, Annual Book of ASTM Standards, Vol. 11.02, 2004. Copyright ASTM INTERNATIONAL. Reprinted with permission.)

sample preconcentration. Carbonates, organic acids, and high total dissolved solids can contribute to unstable baselines and large front-end tailing. While the presence of such species may not be avoided in samples, the use of high purity water for reagent and eluent preparation is essential for ensuring the highest quality chromatography. As most samples for cyanide analysis are preserved at pH 12 or greater, optimum results are achieved when the matrix of the calibration standards is matched to those of the samples.

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TABLE 7.3 Method Detection Limits and Calibration Ranges for the Analysis of Individual Metal– Cyanide Complexes by Ion Chromatography Study level

Ag(CN)− 2

Au(CN)− 2

Co(CN)3− 6

Cu(CN)2− 3

Fe(CN)4− 6

Ni(CN)2− 4

High level (mg/l)

MDL Calibration range

0.77 1–100

0.64 1–50

0.43 1–100

0.09 0.1–2

0.09 0.1–20

0.83 1–200

Low level (µg/l)

MDL Calibration range

8.66 10–125

2.8 5–100

0.99 1–200

4.56 0.1–5

0.21 0.5–20

7.33 50–100

Source: From ASTM, Designation D 6994-04. Annual book of ASTM Standards, Vol. 11.02, ASTM International, West Conshohocken, PA, 2004.

Oxidizing agents. Oxidizing agents like chlorine can decompose certain weak metal-cyanide complexes, thereby causing a decrease in their concentration. Refer to Section 7.1 for pretreatment steps.

7.4.2 REVERSED-PHASE ION-PAIR PARTITION CHROMATOGRAPHY In ion pair (or ion interaction) chromatography, ion pairs are partitioned between a polar mobile phase and a hydrophobic stationary phase. This approach has been applied for the chromatographic separation and quantification of metal–cyanide species in water samples [34–37]. There is no standard method that has been adopted by governmental or consensus organizations, as the technique has been applied mostly in research contexts. However, Waters Corporation has published a method [38] based on the approach of Hilton and Haddad [36]. The ion-pair chromatography methods that have been developed are quite similar to one another. In the methods used by Hilton and Haddad [36] and Grigorova et al. [34], for example, a C-18 stationary phase is employed (C18 Novapak cartridge column, Waters Corporation), and the mobile phase consists of a solution of 2 to 5 mM tetrabutylammonium hydrogensulfate-methanol (approximately 70:30 by volume). Detection is by UV absorption. The mobile phase reagents cause minimal background interference in the low-UV range used in the UV detector, for example, 205 to 215 nm. Elution and separation of metal–cyanide complexes progress in order of decreasing ion-pair polarity, with the more polar ion pairs being eluted earlier. Experience with ion-pair chromatography indicates that the method is capable of measuring rapidly a range of metal–cyanide complexes at mg/l levels. Grigorova et al. [34], for example, − 4− 3− demonstrated the ability of the method to distinguish Cu(CN)− 2 , Ag(CN)2 , Fe(CN)6 , Co(CN)6 , 3− − Ni(CN)2− 4 , Fe(CN)6 , and Au(CN)2 when present together at 10 mg/l each in a synthetic solution. Separation and measurement were completed in 30 min. Experiments with hydrometallurgical gold mining effluent yielded similar results, with individual metal–cyanide species separated, identified, and measured rapidly at mg/l levels. The limited efforts made to apply ion-pair chromatography for measurement of metal–cyanide species at lower (µg/l) level concentrations, indicate the need for further method development. Haddad and Kalambaheti [39], using an analytical approach similar to that described above, studied recovery of various metal-cyanide complexes at low concentrations. A preconcentration step was implemented in the method. The studies were conducted with various electrolyte compositions. 3− It was found that recovery from spiked solutions was low for Fe(CN)4− 6 and Cu(CN)4 , and significant interferences were caused by the presence of other common anions such as Cl− and SO2− 4 .

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Overall, available information indicates that the ion-pair chromatography method is promising, but needs further development for measurement of metal–cyanide species in complex waters at µg/l concentrations. Haddad and Kalambaheti [39] identified some potential ways to address the problems they encountered. There apparently has been little effort to develop the method further, however, since the work in the late 1980s to early 1990s. Huang et al. [37] worked to optimize various aspects of the method, but their efforts were not focused on enhancing method performance at low concentration levels.

7.5 FREE CYANIDE ANALYSIS In analyzing the cyanide content of a water sample, free cyanide, that is, the sum of HCN and CN− , is usually of primary interest, as HCN is a highly toxic form of cyanide. The cyanide anion CN− is easily converted to HCN via a simple change in pH (pK a for HCN is approximately 9.2 at 25◦ C). Since the 1950s, there has been sustained interest and effort in development of reliable, rapid, lowcost techniques for measuring free cyanide at low (µg/l) concentrations in water samples of varying composition. These efforts have yielded a number of different methods, only a few of which are in widespread use today. The differentiating factors among these methods are primarily cost and ease of application. In this section, the leading approaches for analysis of free cyanide in water are described. The most widely used methods are identified, and the reasons for their widespread adoption are discussed.

7.5.1 GAS CHROMATOGRAPHY Since HCN is a volatile species, gas chromatography (GC) may be used for measurement of HCN content in water. GC techniques for HCN analysis were initially developed in the 1960s [40,41] and used in the performance of the first definitive studies of cyanide toxicity to aquatic organisms [42]. In developing the GC technique for HCN analysis, an objective of Schneider and Freund [41] was to avoid shifting the HCN/CN− equilibrium. The technique of Schneider and Freund [41], later modified by Claeys and Freund [40], involves first stripping HCN from an unaltered water sample by passing finely dispersed compressed air through the sample. A relatively large volume sample is used (e.g., 20 l) so that the stripped HCN comprises only a small portion of the HCN in the system. The gas exiting the sample is then passed through a heated drying tube, and into a cooled concentration column containing an adsorbent coated on a granular support. The HCN is thus cold trapped in the concentration column. Subsequently, the concentration column is heated and the contents of the column are injected via carrier gas into a gas chromatograph equipped with a thermal conductivity detector [41] or flame ionization detector [40]. Claeys and Freund [40] demonstrated that HCN concentrations as low as 1 µg/l can be detected with this technique. Modern gas chromatography equipment permits measurement of HCN at sub-ppb concentrations. While purge and trap GC techniques for measurement of HCN are well established, the approach is not widely used, primarily because of the expense of the analysis. Unlike the case for volatile organic compounds, for which GC is used extensively, the instrument configuration needed for HCN analysis often is not useful for the simultaneous measurement of other compounds of interest. Less expensive techniques with detection limits and accuracy similar to that of gas chromatography are available, and are used more frequently.

7.5.2 DIRECT COLORIMETRIC DEVELOPMENT The colorimetric procedure used to measure the amount of cyanide ion present in the NaOH scrubber solution for the total cyanide test (Section 7.2.2) is sometimes used directly on water samples, without distillation. For example, in many of the studies used as the basis for U.S. ambient water quality

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criteria for cyanide [43], free cyanide was determined using the pyridine–pyrazalone colorimetric method from the 13th edition of Standard Methods [44]. In that earlier edition, the introduction to the total cyanide method indicates that distillation can be omitted if “it is known that the sample contains only simple cyanides of the alkalis and is completely free of all interferences [44].” The same document has a list of interferences that includes sulfides, heavy metal ions, fatty acids, thiocyanate, cyanate, glycine, urea, and oxidizing agents. The possibility of omitting the distillation step in total cyanide analysis is not mentioned in more recent editions of Standard Methods [3]. In fact, this is now explicitly discouraged: “The importance of the distillation procedure cannot be overemphasized [3].” Direct colorimetric development to analyze free cyanide is not a standard method of any U.S. government agency or consensus organization. While it is still practiced by some, there are important calibration and interference issues that make the accuracy of the method doubtful, especially with higher ionic strength water samples. With respect to calibration, for example, the issue of matrix matching is significant. Calibration standards for colorimetric analysis of the distillate in the total cyanide test are prepared in NaOH solution of the same composition as the NaOH scrubber solution in which the HCN in the distillate gas is trapped. Using a calibration curve based on standards in NaOH solution to interpret color developed directly in a water sample is problematical. The chemical composition of the calibration standard solution should match that of the sample. Another source of error in direct colorimetric development is pH variation between sample and calibration standards. Variations in pH will affect the cyanide ion distribution in the sample. Calibration standards in NaOH solution are all at a common pH. Finally, for samples dominated by metal–cyanide complexes, it will be difficult to quantify the free cyanide content of the sample with a reasonable degree of accuracy, without first separating the complexes from the solutions to be analyzed for free cyanide. All of the problems listed are eliminated by inclusion of distillation prior to colorimetric development. The accuracy of direct colorimetric development for measurement of free cyanide was evaluated by Dzombak and Higgins [45]. To examine the performance of the direct colorimetric development technique for measurement of free cyanide, spiked samples of freshwater and seawater were tested using the 1971 Standard Methods direct colorimetric method [44], the current colorimetric procedure [3], and the free cyanide by microdiffusion method [46]. The 1971 direct colorimetric method [44] is also sometimes referred to as the “pyridine–pyrazolone method.” The current colorimetric procedure [3] uses pyridine–barbituric acid for color development. Samples of filtered water from a moderately hard, freshwater lake in Pennsylvania and filtered seawater from coastal Oregon were spiked with 20 to 150 ppb free cyanide (KCN) and then analyzed using the methods cited. Results are presented in Table 7.4. It was found that direct colorimetric analysis for measuring free cyanide yielded results comparable to, though less consistent than, those obtained by microdiffusion for free cyanide in freshwater, but that direct colorimetric analysis was grossly inaccurate for measuring free cyanide in seawater. The clear failure of direct colorimetric analysis with the seawater was most likely due to the high ionic strength of the seawater, a known interferent for the colorimetric procedure, and difficulty in buffering the analysis solution to the correct pH.

7.5.3 GAS–LIQUID DIFFUSION Various techniques have been developed that exploit the volatility of HCN, to measure free cyanide in aqueous samples by allowing the HCN to volatilize and then by capturing and measuring the trapped HCN. The basic approach involved with these methods is to provide a confined gas (air) volume above a water sample, and to also put in contact with the gas phase an NaOH solution that serves as a sink for the HCN that diffuses out of the water sample and into the gas phase. Kruse and Thibault [47] first proposed the use of a microdiffusion cell for measurement of free cyanide in water samples. A Conway microdiffusion cell (Figure 7.8) consists of two concentric

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TABLE 7.4 Resultsa of Free Cyanide Analyses in KCN-Spiked Freshwater and Seawater Samples using Direct Colorimetric Analysisb and Microdiffusionc [CN− ] using direct colorimetric analysisb (µg/l)

[CN− ] using free cyanide by microdiffusionc (µg/l)

Percent of free (diffusible) cyanide detected using colorimetric method

Freshwater 20 60 150

36.9 69.2 144

20.2 58.2 145

183 119 99.7

Seawater 20 60 150

11 1.9 4.8

20.4 58.5 155

53 3.3 3.1

Sample nominal value (µg/l)

a Source: Data from Dzombak, D.A. and Higgins, C.J., Quarterly Progress Report (May 27)

WERF Project 01-ECO-1, Water, Environment Research Federation, Alexandria, VA, 2004. b APHA Method 207, Standard Methods for the Examination of Water and Wastewater,

13th ed., American Public Health Assoc., American Water Works Assoc., and Water Environment Research Federation, Washington, DC, 1971. c ASTM D 4282-95. Annual Book of ASTM Standards, Vol. 11.02, ASTM International, West Conshohocken, PA, 1998.

NaOH absorber solution

Sample

FIGURE 7.8 Test apparatus for analysis of free cyanide by microdiffusion.

chambers and a plastic cover or lid. The water sample is introduced in the outer chamber, while NaOH absorber solution is introduced in the inner chamber. After addition of the sample to the outer chamber, a pH 6 buffer solution is added to the sample and the cover or lid is replaced. The entire cell is then placed in the dark for a specified period of time, to allow passive diffusion of HCN to occur. At the end of the diffusion period, the NaOH absorber solution is analyzed for free cyanide content by the colorimetric method outlined in Section 7.2.2. From tests with spiked solutions at different pH conditions, Kruse and Thibault [47] found that complete recovery of free cyanide could be obtained at pH 7 in diffusion periods of 5 h or less.

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The microdiffusion method subsequently was recommended by the American National Standards Institute for measuring free cyanide in photographic development process effluents [48]. The method gained more widespread use and was added as a standard method for free cyanide analysis by ASTM in 1983 [46], with a modification that involved the addition of CdCl2 as a precipitation agent for iron-cyanide complexes following the addition of pH buffer to the sample in the outer chamber of the microdiffusion cell. Broderius [49] developed a similar diffusion procedure for isolating and concentrating HCN in NaOH, but he employed a larger reaction vessel (flask) containing 1.5 l of test solution and about 1.3 l of gas phase volume. A glass dish containing 8 to 20 ml of 0.02 N NaOH was suspended above the test solution inside the flask. A 2 h diffusion period was employed in HCN testing with the apparatus. Broderius chose a larger-scale reactor design and relatively short diffusion period in order that only a small fraction of the total HCN would be removed during the test, thus not disturbing significantly the cyanide species equilibrium distribution in the sample. Also, the Broderius [49] test procedure involved no reagent additions to the water sample for pH adjustment and control. Broderius explicitly wanted to avoid sample acidification as he expected this to induce “conversion to HCN of simple cyanides and, most likely, some portion of certain metallocyanides.” While the Broderius procedure is advantageous for assessing the true HCN concentration under particular solution conditions, it is more often of interest to know the total free cyanide concentration (HCN + CN− ) in a sample. Free cyanide is defined as the sum of the two species because CN− is converted instantaneously to HCN upon a simple lowering of pH in a solution that contains it. Moreover, except for very high pH samples and samples with significant amounts of weak metal– cyanide complexes present, microdiffusion and the Broderius method will yield the same results as HCN is the dominant form of free cyanide for all systems with pH < 9.2. For these reasons, and because of the small scale and simplicity of the method, the microdiffusion method has become the standard and preferred gas–liquid diffusion method for measuring free cyanide in water. Free cyanide by microdiffusion is an ASTM approved method [46] and involves the measurement of HCN evolved passively from an acidified sample. It is a simple technique and easily applied to waters of complex composition. With respect to the microdiffusion method, free cyanide is defined as the cyanide that diffuses at room temperature from simple cyanides or weak metal–cyanide complexes as hydrogen cyanide gas, from a solution of pH 6 to 6.5. The test method does not measure metal– cyanide complexes or organocyanide compounds that resist dissociation, such as, iron cyanide, cyanohydrin, etc. The microdiffusion method does recover some fraction of weaker metal–cyanide complexes if present [23]. The ASTM microdiffusion method is performed in a Conway microdiffusion cell (Figure 7.8). The water sample is introduced in the outer chamber, while a specified amount of NaOH absorber solution is introduced in the inner chamber. After the addition of the sample to the outer chamber followed by the addition of cadmium chloride solution to precipitate the hexacyanoferrates, a pH 6 buffer solution is added to the sample and the cover or lid is replaced. The closed cell is then placed in the dark to allow 4 h of diffusion [46]. During this time, free cyanide diffuses as HCN gas and is absorbed into the sodium hydroxide absorber solution in the center chamber, where it is converted to CN− . At the end of the specified diffusion period, the CN− in the absorber solution is analyzed using the standard colorimetric procedure (Section 7.2.2). The typical calibration range is between 0 and 150µg/l. In addition to the potential for weak metal–cyanide complexes to decompose under the acidic conditions imposed for the microdiffusion test, the method is subject to some other interferences, primarily from other volatile species that may also be trapped in the NaOH absorber solution. Volatile species that can affect the results obtained in the microdiffusion test include ammonia, sulfide, and phenol [50]. These substances need to be present in fairly large quantities in the NaOH absorber solution, in order to influence the colorimetric procedure used to evaluate its free cyanide content, for example, ammonia and phenol greater than 50 mg/l, and sulfide greater than 1 mg/l. The microdiffusion method has undergone an inter-laboratory study and has been modified to eliminate the use of CdCl2 for iron-cyanide complex precipitation [51]. The revised method will

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be submitted to the USEPA Office of Solid Waste for potential inclusion in the USEPA analytical methods manual SW-846 [52].

7.5.4 ION-SELECTIVE ELECTRODE As discussed in Section 7.2.2, the cyanide ion-selective electrode can be used for measurement of CN− concentration in the NaOH absorber solution used to capture HCN from the distillate in conventional total cyanide analysis. The cyanide ion-selective electrode is also employed directly for analysis of free cyanide in water samples. For example, in the United States, the cyanide ion-selective electrode is allowed to be used directly for analysis of finished drinking water, to demonstrate compliance with the U.S. national primary drinking water regulations [53]. The cited procedure, Method 4500CN F from Standard Methods, is presented there as a finish procedure to analyze CN− ion in the NaOH absorber solution used to trap HCN in the distillation of a water sample for total cyanide analysis. Use of the cyanide ion-selective electrode for direct analysis of a water sample would require dilution with NaOH prior to analysis, to match the matrix of calibration standards and to convert all HCN to CN− . There is no discussion in Standard Methods [3] of the use of the ion-selective electrode directly on a water sample made alkaline by NaOH addition, but this approach has been employed and reported in the literature. Sekerka and Lechner [54] presented a method for analysis of free cyanide in a water sample using the cyanide ion-selective electrode without distillation. To a 20 ml sample of water, 10 M NaOH is added to achieve a pH of 11.5. Small quantities of some additional reagents (orthophosphoric acid and bismuth nitrate), related to a procedure for identification of interfering halides, are also added. Cyanide ion concentration is determined by the standard colorimetric method. In tests with analysis of synthetic solutions of KCN in double-distilled water, cyanide ion concentrations as low as 10−7 M(2.6 µg/l) were measured. Sekerka and Lechner [54] also studied the magnitude of known interferences with the cyanide ion-selective electrode. The cyanide electrode malfunctions if anions that form salts with silver are present, including sulfide (S2− ), iodide (I− ), bromide (Br − ), thiocyanate (SCN− ), and chloride (Cl− ). Measurements performed on samples with 5 × 10−7 M cyanide ion indicated that chloride and thiocyanate did not interfere until their concentrations exceeded the cyanide concentration by 104 , and bromide did not interfere up to a concentration ratio of 102 . On the other hand, iodide ion interfered at a 1:1 ratio, and sulfide species interference was significant at all concentrations studied. It is thus feasible to employ the cyanide ion-selective electrode for direct analysis of free cyanide in water, but only for relatively simple, low ionic strength water samples. Without distillation, there is a large potential for interference with the cyanide ion-selective electrode measurement from ions and organic substances present in the water. Because of this, and the availability of less complicated alternative methods such as the microdiffusion technique [46], the cyanide ion-selective electrode is infrequently used for direct measurement of free cyanide on water samples without distillation.

7.6 THIOCYANATE AND CYANATE MEASUREMENT TECHNIQUES 7.6.1 THIOCYANATE MEASUREMENT PROCEDURE Thiocyanate (SCN− ), which is not captured by conventional total cyanide tests, may be measured colorimetrically using Standard Methods 4500-CN M [3]. The procedure involves sample titration under acidic pH conditions (pH = 2) with ferric nitrate. Under the analysis conditions, the Fe3+ reacts with SCN− to form an intense red color complex that exhibits an absorption maximum near 460 nm. The concentration of thiocyanate is determined spectrophotometrically at the wavelength maximum by comparison against a standard calibration curve of absorbance versus concentration.The

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typical calibration range is between 0.1 and 2 mg/l. Further details about the method are available in Standard Methods [3]. Common interferences to the method and their corresponding remedies are discussed as follows: Hexavalent Chromium. Hexavalent chromium is a positive interference and is removed by addition of ferrous sulfate after acidifying the sample to between pH 1 and 2 using nitric acid. Reducing Agents. Reducing agents can reduce Fe3+ to Fe2+ , thereby preventing the formation of the color complex. Reducing agents can be destroyed by adding few drops of hydrogen peroxide. Sulfide. Sulfide can convert any free cyanide ion in the sample into SCN− during the test procedure and therefore could serve as a positive interferent. Sulfides are removed by adding lead salts prior to the iron addition. Refer to Section 7.1 for sulfide removal steps. Microbes. Thiocyanate is biodegradable. Samples are thus preserved with acid addition to pH < 2 and stored at 4◦ C to inhibit microbial activity. In addition to the titration method, an ion chromatography technique is also available to measure SCN− in solutions. Separation of thiocyanate is achieved using an AG5 guard column and an AS5 analytical column with AMMS-I and 25 mM H2 SO4 as column suppressor and regenerant, respectively [33]. Isocratic elution is maintained using the four different eluent preparations (2 mM NaOH, 4.5 mM Na2 CO3 ; 2 mM NaOH; 0.8 mM 4-cyanophenol and 2% acetonitrile) and final detection is performed using a conductivity detector. This method has not yet been approved by any consensus or governmental organization.

7.6.2 CYANATE MEASUREMENT PROCEDURE Cyanate is also not captured in conventional total cyanide tests and must be measured individually when its presence is suspected and is of interest. The Standard Methods 4500-CN-L [3] procedure for measuring cyanate (CNO− ) involves high temperature hydrolysis at low pH. At a pH of 2 to 2.5 and at a temperature between 90 and 95◦ C, cyanate hydrolyzes to ammonia according to the following reaction: 2NaCNO + H2 SO4 + 4H2 O → (NH4 )2 SO4 + 2NaHCO3

(7.1)

Ammonia concentration is measured before and after the treatment using either an ammonia ion-selective electrode, or by colorimetric development followed by direct nesslerization. This test is applicable for measurement of cyanate compounds in industrial waste and natural waters. The typical calibration range is between 1 and 200 mg/l with method detection limit around 1 to 2 mg/l. Common interferences to the method and corresponding pretreatments are as follows: Organic Compounds. Organic nitrogenous compounds can hydrolyze to ammonia following acidification. Thus, acidification and heating must be controlled carefully to minimize such interferences. Metals. Metallic compounds can form colored complexes with nessler reagents during ammonia analysis. Addition of Rochelle salt or EDTA during the ammonia analysis by nesslerization can overcome these interferences. Oxidizing agents. Certain powerful oxidizing agents can oxidize cyanate to CO2 and nitrogen. Refer to Section 7.1 for detection and pretreatments of oxidizing agents. Further details about the method are available in Standard Methods [3].

7.7 CYANOGEN HALIDE ANALYSIS Cyanogen halides, especially cyanogen chloride (CNCl), are highly toxic species and are not captured in the total cyanide test. Hence, their formation and presence is of interest, especially in situations in which water containing cyanide is chlorinated or brominated [4,29].

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Cyanogen halides (CNX, where X = Cl, Br, etc.) react with pyridine barbituric acid (see Section 7.2.2 under Colorimetric Procedure) to produce a red–blue color complex at pH < 8 that exhibits an absorption maximum at 578 nm. The concentration of cyanogen halides is determined spectrophotometrically by comparison against a standard calibration curve of absorbance versus concentration. The typical calibration range is between 5 and 150 µg/l. A colorimetric technique specifically for cyanogen chloride is given as Method 4500-CN J in Standard Methods [3]. The primary interference to this method is the instability of the analytes [3]. Cyanogen halides can hydrolyze to cyanate at a pH of 12 or more. Standard Methods [3] recommends that samples should be collected without any form of preservation, including NaOH addition, and analyzed as quickly as possible following acquisition to minimize any hydrolysis. If the presence of oxidizing agents or sulfide is detected or suspected, pretreatment steps outlined in Section 7.1 should be implemented. Gas chromatography can also be used for analysis of cyanogen halides [4,55]. This approach has been used primarily for analysis of drinking water samples. To prevent the hydrolysis degradation of cyanogen halides during sample storage, it is recommended that the pH of samples be lowered to 3.0–3.5 with sulfuric acid [4].

7.8 ORGANOCYANIDE MEASUREMENT TECHNIQUES Organocyanide compounds generally are not measured with the conventional analytical tests, which involve treatment to liberate free cyanide followed by measurement of the free cyanide. Most organocyanides are resistant to the release of cyanide ion in the total cyanide, WAD cyanide, and CATC test conditions. There can be some partial recovery of organocyanide compounds in the conventional test methods. Table 7.5 presents results obtained by Yi et al. [5] for recovery of cyanide from solutions spiked with four different organocyanide compounds in the standard total cyanide test. As shown there, for two of the compounds there was essentially no recovery, while for the other two compounds, mass recoveries (as CN) were 19 and 73%, respectively. The effect of chlorinating the solutions on recovery was also studied. There was no change in recovery for the two compounds not detected in the total cyanide test without chlorination (Table 7.5). For the other two compounds, recovery was reduced to zero for one and doubled for the other. The reduction in recovery probably resulted from oxidative destruction of the compound, including the −CN group. Overall, the results of Yi et al. [5] demonstrate that recovery of cyanide from organocyanide compounds in the standard total cyanide and CATC method is quite variable and substantially less than 100%. Methods other than the conventional cyanide analysis techniques need to be employed to analyze for organocyanide compounds in water. Specific organocyanide compounds can and have been measured with liquid chromatography techniques [56], but there is no available method for measuring total organic cyanide in a water sample. Theis et al. [57] developed and tested a scheme for assessment of total organocyanide content of a water sample that involved the following steps: • Separation of inorganic cyanides by passing the preserved sample (pH ≈ 11) through an anion exchange resin bed (100 to 200 mesh resin). • Measurement of concentrations of inorganic species remaining after ion exchange pretreatment (

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