Nitrosation Reactions and the Chemistry of Nitric Oxide

Nitrosation Reactions and the Chemistry of Nitric Oxide Elsevier, 2004 Author: D.L.H. Williams ISBN: 978-0-444-51721-0 ...

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Nitrosation Reactions and the Chemistry of Nitric Oxide Elsevier, 2004 Author: D.L.H. Williams ISBN: 978-0-444-51721-0

by kmno4

CONTENTS INTRODUCTION 1. Reagents effecting nitrosation 1.1. Nitrous acid 1.1.1.

Acid catalysed pathways

1.1.2.

Reaction via dinitrogen trioxide N203

1.1.3.

Nucleophile catalysed pathways

1.2. Nitrosyl halides HalNO 1.3. Nitrosonium salts NO'X1.4. Alkyl nitrites RON0 1.4.1.

Reactions in aqueous acid solution

1.4.2.

Reactions in aqueous basic solution

1.4.3.

Reactions in non-aqueous solvents

1.5. N-Nitrososulfonamides RS02N(NO)R1 1.6. Nitrogen oxides 1.6.1.

Nitric oxide NO

1.6.2.

Nitrogen dioxide NO2

1.7. Miscellaneous reagents 2. Nitrosation at nitrogen centres 2.1.

Primary aromatic amines

2.2.

Primary aliphatic amines

Contents

v1

2.3. Secondary amines 2.4. Tertiary amines 2.5. Amides and related compounds 2.6. Other nitrogen-containing compounds 3. The reactions of N-nitrosamines and related compounds 3.1. Rearrangement of aromatic N-nitrosamines Ar(R)NNO (the Fischer-Hepp rearrangement) 3.2. Denitrosation of nitrosamines R(R1)NNO 3.3. Denitrosation of nitrosamides RCON(N0)R' 3.4. Denitrosation of nitrososulfonamides ArS02N(NO)R 3.5. Carcinogenic properties of N-nitroso compounds 4. Aliphatic and alicyclic C-nitrosation 4.1. Nitrosation of ketones 4.2. Nitrosation of other carbonyl compounds 4.3. Nitrosation of nitroalkanes 4.4. Nitrosation of carbanions 4.5. Addition to alkenes 5. Aromatic C-nitrosation 5.1. Products of the reactions 5.2. Reaction mechanisms 5.3

Nitrous acid-catalysed nitration reactions

6. O-Nitrosation 6.1.

Nitrosation of alcohols

6.2.

Nitrosation of hydrogen peroxide

6.3.

Nitrosation of ascorbic acid

6.4.

Other O-nitrosation reactions

7. S-Nitrosation 7.1. Nitrosation of thiols 7.2.

Nitrosation of thiocarbonyls

7.3.

Nitrosation of organic sulfides

7.4. Nitrosation of sulfinic acids 7.5.

Nitrosation of thiocyanate ion

7.6. Nitrosation of sulfite/bisulfite 7.7.

Nitrosation of thiosulfate ion

7.8. Nitrosation of inorganic sulfide 8. Synthesis, properties and reactions of S-nitrosothiols 8.1.

Synthesis

8.2. Physical properties 8.3.

Thermal and photochemical decomposition

8.4.

Decomposition in aqueous solution

8.5. Reaction with nucleophiles 8.5.1.

Reaction with thiols (transnitrosations)

Contents

8.6.

8.5.2.

Reaction with ascorbic acidlascorbate

8.5.3.

Reaction with hydrogen peroxide

8.5.4.

Reaction with other nucleophiles

Detection and quantitative determination of S-nitrosothiols 8.6.1. Spectrophotometric determination of RSNOs 8.6.2.

Electrochemical determination of RSNOs

8.6.3.

Capillary zone electrophoresis

8.6.4.

Conversion to nitrite ion

8.6.5.

Conversion to nitric oxide

9. Nitrosation involving metal-nitrosyl complexes 9.1.

Sodium nitroprusside (pentacyanonitrosylferrate 11)

9.2.

Other metal nitrosyls

9.3.

Iron-sulfur cluster nitrosyls

9.4. Nitrosation of nucleophiles co-ordinated to metals 10. The biological chemistry of nitric oxide 10.1. Background 10.2. Reactions of nitric oxide with oxygen and superoxide 10.3. Reactions of nitric oxide with haem proteins 10.4. Reactions of nitric oxide with thiolate ion 10.5. Analytical methods for the determination of nitric oxide 10.5.1. Colorimetric methods 10.5.2. Electrochemical methods

Contents

ix

10.5.3. The oxyhaemoglobin method

182

10.5.4. Chemiluminescence method

182

10.5.5. Spin-trapping methods

183

10.5.6. Fluorescence spectroscopy

186

1 1. Nitric oxide in biological systems

187

11.1. Biological properties of nitric oxide

187

11.2. Biosynthesis of nitric oxide

192

11.3. NO synthase enzymes

194

11.4. Nitric oxide in the treatment of foods

197

12. Nitric oxide-releasing compounds (NO donors)

199

12.1. Organic nitrates

199

12.2. Organic nitrites

20 1

12.3. S-Nitrosothiols

202

12.4. Inorganic nitrosyl complexes

208

12.4.1. Sodium nitroprusside Na2[(CN)5FeNO](SNP)

208

12.4.2. Ruthenium nitrosyls

209

12.4.3. Dinitrosyl-iron-thiol complexes (DNICs)

209

12.4.4. Iron-sulfur cluster nitrosyls

210

12.5. NONOates (Diazeniumdiolates)

211

12.6. Sydnonimines

213

12.7. Furoxans (oxadiazoles)

214

12.8. Hydroxyurea

2 15

x

Contents

12.9. Other NO-donors 13. Nitroxyl (HNO) and the nitroxyl anion (NO-) 13.1. Generation of HNO/NOP 13.2. Physical and chemical properties of HNO/NO13.3. Determination of HNO/NO13.4. Biological implications Appendix References Index

XI

INTRODUCTION Nitrosation reactions have been known for over 150 years. The early work was concerned with the reactions of amines. Piria first appeared to have carried out the reactions with ahphatic primary amines in 1846 and isolated some deamination products. Later Hofmann worked with primary aromatic amines. Griess started the work on the chemistry of diazonium ions, and nitrosamines were characterised a little later from the reaction of secondary amines. The rearrangement of aromatic nitrosamines was discovered by Fischer and Hepp. Victor Meyer showed in 1873 that nitrosation at carbon was possible when he isolated nitrolic acids from aliphatic nitro compounds, and addition products of nitrosyl chloride with alkenes were then reported by Tilden. Alkyl nitrites were then generated from alcohols and some early work at the beginning of the last century generated some S-nitrosothiols from thiols. Other reactions continued to be discovered later. Many of these reactions are now standard laboratory procedures and a number have been adapted to the industrial scale. The most widely known of these is diazotisation and the subsequent coupling of arenediazonium ions with a variety of amines and phenols to generate azo dyes which have been widely used to dye garments. Hydroxylamine (an important intermediate) is made by the Raschig process involving nitrosation of bisulfite; nitrosation of cyclohexane derivatives leads to s-caprolactam and hence Nylon 6 on polymerisation; nitroso compounds are used in the rubber industry and alkyl nitrites and metal nitrosyl complexes (notably sodium nitroprusside) have played a part as vasodilators in medicine. Since the discovery in 1956 that nitrosamines were carcinogens in animal experiments, there was an explosion of interest in their chemistry and particularly in methods to destroy them in a variety of consumer products in which they occur, often as very minor by-products. In more recent times, since the discovery of the amazing biological properties of nitric oxide there has been a large interest in the chemistry of S-nitrosothiols, not only as possible vasodilators, anti-platelet aggregation agents etc., but also since there is a general belief (not fiilly authenticated) that they in some way act in vivo as storage and transport vehicles for nitric oxide. Apart from the large synthetic aspect there have been also a large number of mechanistic studies carried out on nitrosation reactions. It has proved a rich and rewarding field for physical organic chemists, initially concentrated on amine nitrosation, but widely extended in more recent times to nitrosation at centres other than nitrogen. A large number of features have stood out, including acid catalysis, base catalysis, nucleophilic catalysis, rate limiting proton transfer, difftision controlled reactions, intramolecular rearrangements, trans-nitrosation reactions etc., all of which have been established under various experiment conditions. Much is now known about the detailed reaction mechanisms of nitrosation reactions.

xii

Introduction

However in spite of the wide diversity and application of such reactions, they have not received as much coverage as have for example the corresponding nitration reactions, and until 1988 there was no single book devoted to their study, even though most standard organic texts and review articles covered various aspects. The short monograph 'Nitrosation', published in 1988 was an attempt to fill this rather obvious gap. Since that time there has been something of a burst of activity in a large number of areas, but notably resulting from the discovery of the biological properties of nitric oxide. Many aspects of nitric oxide chemistry are closely inter-linked with nitrosation reactions. It was thus felt to be an appropriate time to generate a larger textbook devoted to nitrosation reactions and the chemistry of some of the nitroso products of these reactions, and an account of the chemistry of nitric oxide. The early chapters are heavily influenced by the format of the 1988 monograph, which is now out of print, and attempts are made to bring together both the synthetic and mechanistic aspects of the reactions. Again there is a chapter which discusses the reagents which can effect nitrosation. The basic early ideas are summarised and there is a concentration on the more recent developments. Thereafter the book contains chapters relating to nitrosation at different element centres and there are additional chapters which give the chemistry of nitric oxide and of nitric oxide donors, in so far as they are related to the range of biological properties. It is not appropriate in a book written for chemists for the detailed biology to be discussed, but there are short chapters on the biological chemistry of nitric oxide and on the biology of nitric oxide. There are a large number of review articles which discuss the biological aspects more fiilly, to which the reader is referred. The book concludes with a short chapter on the known chemistry of nitroxyl HNO and of the nitroxyl anion NO" , which may be closely linked with some of the biological chemistry of NOcontaining compounds. The Appendix contains some physical data of important compounds relating to nitrosation and nitric oxide chemistry. I am indebted to a number of people with whom I have discussed this area over a number of years. In particularly I would like to thank in this regard. Professor John Ridd, Dr. Geoffrey Stedman, Dr. Tony Butler and Professor Ramon Leis and his colleagues from the University of Santiago de Compostela, Spain, from all of whom I have learnt much. I also wish to thank members of my research group who have worked in this area (too numerous to name individually) over the past 30 years or so, who have made major contributions, particularly to the understanding of the various reaction mechanisms involved. Any errors and omissions remain my responsibility. Durham, April 2004

Lyn Williams

Nitrosation Reactions and the Chemistry of Nitric Oxide D.L.H. Williams © 2004 Elsevier B.V. All rights reserved.

Chapter 1

Reagents effecting nitrosation l.l.

Nitrous acid The most commonly used reagent by far for bringing about electrophilic nitrosation is nitrous acid, generated in aqueous acid solution from a nitrite salt, usually sodium nitrite, and used in situ. Pure nitrous acid has never been isolated, since decomposition occurs giving various oxides of nitrogen as final products. This decomposition is usually represented by Eq. (1), but it does not 3HNO2 = 2N0 + HNO3 + H2O

(1)

give the full picture as nitric oxide will be further oxidised to nitrogen dioxide. Fortunately the decomposition is relatively slow at room temperature and at low [HNO2], around 1 x 10~^ - 1 x 10"^ M, and this allows both synthetic and mechanistic work to be carried out at 0°C and 25 °C, where much of the kinetic work has been done. In the absence of oxygen the decomposition pathway involves the two equilibria Eq. (2) and (3), whereas in the presence of oxygen 2N02 + 2H2O

^ =^ HNO2 + NO3- + H3O

(2)

2HN02 ^S=^ NO + NO2 + H2O

(3)

2N0 + O2

(4)

- ^ 2NO2

Eq. (4) comes into play, and decomposition is significantly faster [1,2]. Where decomposition is a problem e.g. for slow nitrosation reactions, better quantitative work can be achieved by carrying out the reactions anaerobically e.g. under nitrogen. The structure of nitrous acid is well-known [3]. The molecule exists in both cis and trans forms (structures 1 and 2) with the trans form dominant. / N—O // O

H N—Q // \ O H

The bond lengths and angles are known and there is some dipolar character which confers partial double-bond character to the N-OH bond.

2

Nitrosation Reactions and the Chemistry of Nitric Oxide

Nitrous acid is a weak acid, Eq. (5). Its ^K^ value has been measured many times by different methods. The most reliable value is probably 3.148 at HNO2 + H2O ^5=^ NO2" + H3O+

(5)

25°C and at zero ionic strength [4]. Measurements carried out on nitrous acid nitrosation reactions at pH > -2.5 require the correction due to N02~ formation to be made. Early attempts at mechanistic studies, particularly kinetic studies, in nitrosation presented a confused situation. This was resolved and put together in an excellent review article by Ridd [5], who rationalised the results and gave a comprehensive account of the various mechanistic pathways involved, which is the basis of current-day thinking in this area. Nitrous acid exists in aqueous solution in equilibrium with dinitrogen trioxide N2O3, Eq. (6), which is also an effective nitrosating species, which can 2HNO2 ^^^^

N2O3 + H2O

(6)

be regarded as nitrosonium nitrite. The bluish tinge seen in fairly concentrated solutions of nitrous acid is due to N2O3. Quantitative work on the reactivity of N2O3 generated in this way in situ requires an accurate value for the equilibrium constant for N2O3 formation (i.e. [N203]/[HN02]^). There is a wide spread of values in the literature, most of which were determined spectrophotometrically. However there is now general agreement that a value of 3.0 X 10~^ M~^ is probably the most reliable [6]. This is also based on a measured extinction coefficient for N2O3, which is in agreement with a value obtained by pulse radiolysis. Use of this value leads to sensible values for the rate constants for N2O3 attack on very reactive substrates, which are believed to occur at the diffusion-controlled limit. Two other equilibria are important in the context of nitrosation by nitrous acid. One involves the conversion to the nitrosonium cation NO"^ which becomes important at very high acidities, Eq. (7). In 60% perchloric HNO2 ^ ^^3^^ "^^

N0+ + 2H2O

(7)

acid and 60% sulfuric acid the conversion to NO"^ is virtually complete. This was demonstrated by observation of a new band at 260 nm in the UV spectrum and also by the appearance of the Raman line at 2300 cm~^ both of which are characteristic of NO"^, as measured for a number of nitrosonium salts NO'^X", which can be isolated and which are commercially available. There is however some uncertainty as to whether NO"^ is involved as the nitrosating species in quite dilute acid solution (see Section 1.1.1.).

Reagents Effecting Nitrosation

3

The other important equilibrium is that which is estabhshed in the presence of a range of non-basic nucleophiles X~ (some can be neutral), which generates a new family of potential nitrosating species XNO, Eq. (8), some of HNO2 + ^ 3 ^ ^ + ^~

^*=^ ^^Ô + 2H2O

(8)

which can be isolated in the pure state, usually by other routes. As to whether acidic solutions of nitrous acid react via N2O3, NO"^ or XNO will depend on a number of factors particularly the pH, [HNO2] and the nature and concentration of X~. Often there are spectacular catalytic effects for a number of X~ species, which will be discussed later in detail in Section 1.1.3. 1.1.1. Acid catalysed pathways In relatively dilute aqueous acid solution (typically 0.1 M) and at relatively low [HNO2] (typically 1 x 10"-^ - 1 x 10~^ M), nitrosation of a large number of nucleophilic substrates (S) follows the rate equation Eq. (9). Here Rate = ^[HN02][H30+][S]

(9)

reaction is acid catalysed and first order in each of the reagents. There has been some controversy regarding the exact nature of the nitrosating species. The rate equation is consistent with mechanisms involving H2N02'^ and also NO"^, Eq. (10, 11). There is a close analogy with the situation in electrophilic HNO2 + H3O+ ^5=^ H2N02^ + H2O H2N02^ + S

—•

NO"^ + S — •

(10)

^

(11)

Product

HNO2 + ^ 3 0 ^ ^5=^ H2N02^ + H2O H2N02^ —

1

NO+ + H2O Product

nitration where the reagent is N02'^. However, in nitrosation there is no spectroscopic evidence in favour of either reagent for reaction in dilute acid solution [7]. Part of the evidence in favour of NO2"*' as the reagent in nitration, derives from the observation of a rate law which is zero-order in S for very reactive substrates, when there is rate limiting formation of N02^. In nitrosation, although there have been claims of such a zero-order law (for the nitrosation of hydrogen peroxide [8] and of alcohols [9]), the evidence is not convincing, since very high substrate concentrations have to be used and the

4


curved plots of k^^^ vs [S] can be interpreted in terms of a medium effect. However in another solvent, (acetonitrile) there is a clear zero-order dependence on [S] for reaction with an alcohol and also a thiol (at low concentration), which points to rate limiting NO"^ formation [10]. For reactions in water there is a persuasive argument in favour of reaction via H2^^2^ from experiments involving ^Ô exchange, and by comparison of the rate constant for ^Ô exchange between nitrous acid and water, and those for reactions with a number of anionic species including azide ion at relatively high concentrations [11]. There is also some indirect evidence from ^^N NMR studies [12] that the nitrous acidium ion H2N02'^ does play a part in acidic solutions of nitrous acid, but the definitive physical evidence is still lacking. Two independent theoretical studies [13,14] conclude that the protonation of nitrous acid occurs at the hydroxylic oxygen atom (thus creating a good leaving group H2O for NO"*" transfer) and that reaction occurs via a nitrosonium ion-water complex (H20'^—NO) rather than via the free nitrosonium ion itself. Whatever the exact nature of the electrophile, it is of some interest to compare some of the values of the third-order rate constant defined by Eq. (9) for a range of different substrates. Some of the literature values are collected in Table 1. For detailed references see [15]-[19]. The values of k accord qualitatively with a process involving electrophilic attack (e.g. aniline > 4-nitroaniline > 2,4-dinitroaniline). There is also a clear tendency for the values to approach a limiting value at higher reactivity. This limit is ---7,000 M~^ s~^ for neutral substrates at 25°C, -12,000 M"^ s~^ for mono-negatively charged species at 25°C and --2,000 M~^ s~^ for mono-negatively charged species at 0°C. It has been argued that this represents the approach to a diffusion controlled process where the rate limiting step is the formation of an encounter pair, which precedes the actual chemical process [20]. If we assume that the rate constant for encounter pair formation is ~ 7 x 10^ M~^ s~^ then we get a value for the equilibrium constant for formation of the nitrosating species of ~1 X 10~^ M~^ This is within the extremely large range (3.0 x 10"^ 7.8 X 10"^ M~^) variously reported in the literature [21] for NO"^ formation at high acidities. Given such a large range of values however, this result cannot be taken as definitive evidence for the involvement of NO"^ at low acidities. For mono-negatively charged species the limiting value of k appears to be --12,000 M~2 s"^ (and maybe --18,000 M~^ s"^ for the one example, 8203^" of a doubly-negatively charged species) as expected from the extra electrostatic forces involved in the formation of an encounter pair with a positively charged reagent. Until recently the most reliable value of K for NO"^ formation, Eq. (7), has been taken as 3 x 10~^ M~^ [7],[20]. However, measurements of the solubility of nitrous acid in sulfuric acid have enabled a value of 1.2 x 10"^

Reagents Effecting Nitrosation Table 1 Values of A: in Rate = A:[HN02][H30+][S] Substrate S Urea Methylurea Hydrazoic acid Hydrazinium ion Methanol Sulfamate ion 2,4-Dimtroaniline 4-Nitroaniline Aniline 4-Methylaniline Cysteine Mercaptopropanoic acid Thiourea Chloride ion Bromide ion Iodide ion Thiocyanate ion Nitrite ion Ascorbate ion Acetate ion Azide ion Thiocyanate ion Benzenesulfinate ion Thiosulfate ion Enolate of Meldrum's acid

i^/M-2 s-1 0.89 10 160 620,611 700 1,167 2.5 2,600 4,600 6,000 456,443 4,764 6,960 975 1,170 1,370 1,460 1,893 2,000 2,200 2,340 11,700 11,800 18,000 12,700

Temperature/°C 25 25 25 25 25 25 25 25 25 25 25 25 25 0 0 0 0 0 0 0 0 25 25 25 25

M~^ to be deduced [22]. This agrees well with an earlier measurement by Seel and Winkler [23] of 1.4 x 10"^ M~^ so this is probably the most reUable figure currently available. This highlights one of the problems in this area, when subsequent calculations depend so much on the magnitude of K here, and reinforces the point regarding the difficulty of obtaining a reliable value for such a small equilibrium constant. 1.1.2. Reactions via dinitrogen trioxide 1^2^3 As mentioned earlier, solutions of nitrous acid in water are in equilibrium with dinitrogen trioxide Eq. (6), which is itself a powerful nitrosating species, Eq. (12). Pure solid and liquid forms (which are blue) of

6


N2O3 + S — • +S-NO + NO2"

(12)

N2O3 can be readily isolated at low temperatures. There is however a strong tendency to dissociate to NO and NO2 as the temperature rises. This can readily be demonstrated in a sealed system. In aqueous solution at low acidities and reasonably high [nitrous acid] however, the decomposition pathway is sufficiently slow to allow kinetic work to be undertaken. The rate equation, Eq. (13) has been established many times for a range of substrates. This is Rate = yi:[HN02]^[S]

(13)

easily distinguished experimentally from Eq. (9) by the absence of acid catalysis and by a second-order dependence upon [HNO2]. This is the rate equation expected for rate limiting attack by an equilibrium concentration of N2O3 on S, and the third-order rate constant k in Eq. (13) is given by ^2^N70Q where ^2 is the second-order rate constant for N2O3 attack and ^ N ^ Q ^ ^^ ^^^ equilibrium constant for N2O3 formation. Values of k2 are thus crucially dependent upon the value taken for K^^Q , and the most reliable appears to be 3.0 X 10~^ M~^ [6]. Some values 01 K2 obtained in this way (mostly for amines) are given in Table 2. The detailed references are cited in [15] p.5. It is clear that kj increases with the reactivity of S as measured (for the amines) by the p^^ value of the conjugate acid and again the ^2 values approach that expected for a diffusion controlled process as the reactivity is increased. More recently obtained values include 1.2 x 10^ M~^ s~^ for the enol form of acetone [16], and 1.2 x 10^^ M~^ s"^ for the enolate form of indane-l,3-dione [24], both of which probably react at the diffusion limit. For some of the reactions, the activation energies were determined. Values in the range 10-20 kJ mol"! were found for those reactions with ^2 values ~1 x 10^ M~^ s~^ [25], supporting the suggestion that we are indeed at, or close to the diffusion limit here. For a number of very reactive substrates reaction is zero-order with respect to [S] and the observed rate equation is given by Eq. (14). This is Rate = klUNOj]^

(14)

readily interpreted as rate limiting N2O3 formation, and will occur when the reaction of N2O3 with S is much faster than its hydrolysis, resulting in a change of rate limiting step. This has been achieved experimentally for

Reagents Effecting Nitrosation Table 2 Values of ^^2 in Rate = ^2[N203][S] at 25°C Substrate (S) 2-Chloroaniline 3-Chloroaniline 4-Chloroaniline 2-Methylaniline Aniline 3-Methylaniline AT-Methylaniline 4-Methylaniline 4-Methoxyaniline (0°C) Piperazine Hydroxylamine Mononitrosopiperazine Morpholine Methylbenzylamine Ethylbenzylamine 7V-Methylglycine 2-Butylamine Propylamine Methylamine Diisobutylamine Dimethylamine Diethylamine Dipropylamine Methylcyclohexylamine Diisopropylamine Piperidine Dibutylamine 2-Methylbut-2-ene (0°C) 2,3-Dimethylbut-2-ene (0°C) 2-Phenylindole (3°C) 2-Phenylindolizine Azide ion 3 -Chloropheny Ihy drazine

P^. 2.63 3.46 3.92 4.39 4.60 4.69 4.85 5.07 5.29 5.55 5.90 6.80 8.70 9.54 9.68 10.20 10.56 10.67 10.70 10.82 10.87 10.98 11.00 11.04 11.20 11.20 11.25 -

k-ylM-^ s-1

1.4x107 9.6x107 2.8 X 108 4,2x108 7.5x108 8.2x108 4.0x108 1.9x109 1.8x108 1.3x108 2.0x108 7.5 X 107 2.2x108 1.8x108 3.1 x l 0 8 1.5x108 1.6x108 9.4 X 107 1.6x108 1.3x108 1.2x108 1.1x108 1.2x108 2.2 X 108 1.6x108 1.3x108 1.8x108 5.5 X 104 3.9x105 0.6x108 9.3 X 109 2.1 X 109 5.9x109

reactions with N-methylaniline, ascorbic acid, azide ion, thiosulfate ion (see ref. [15] p.3) and 4-mercaptopyridine [19]. Literature values of K vary somewhat in the range 4-30 M~^ s~^ for reactions at 25"^, possibly as a result of different ionic strength conditions and more probably from the possibility that

8


not all reactions were fully zero-order in S and there may be a small component of a reaction taking place via the acid catalysed pathway, not detectable in the second order-analysis. 7.7.3. Nucleophile catalysed pathways A range of non-basic nucleophiles added to aqueous acidic solutions of nitrous acid often generate major catalytic effects. The nucleophiles (X~) include chloride ion, bromide ion, iodide ion, thiocyanate ion, thiourea (and alkyl thioureas), thiosulfate ion, 4-mercaptopyridine (and other heterocyclic thiones) and dimethyl sulfide. This catalytic effect arises from the equilibrium formation of XNO which acts as the effective nitrosating species, Eq. (8). The rate equation expected from such a mechanism is given in Eq. (15), and has Rate = ^xN0^XN0[HN02][H30+][X-][S]

(15)

been established many times for a range of S species. Reaction is first-order in each of HNO2, H30"^, X~ and S; ^XNO ^^ ^^^ second-order rate constant for the reaction of XNO with S and ^XNO ^^ ^^ equilibrium constant for XNO formation. Reactions are often carried out with [S] » [HNO2]. The other components are regenerated during reaction and their concentration is thus constant in any one experiment. Plots of the measured first-order rate constant ôbs ^^ [^T ^^^ generally linear. In many cases ^xNO ^^^ ^QQW determined independently (usually by spectrophotometry in the UV spectra). Values of ^XNO ^^^ ^^^^ ^^ readily obtained from the slope of plots of k^^^ versus [X-] and the various concentration terms. The catalytic efficiency depends on the product ^XNO^XNO* Values of A^xNO ^^^ given in Table 3. The range of values is large, covering ten powers of ten. There is a reasonable correlation between the magnitude oiKy^^ and the Edwards nucleophilicity parameter of X~, and this has been used to interpolate values of ^xNO ^^^ those species where it has not proved possible to make a direct measurement [31]. Some values of A:x^Q are given in Table 4. The trends are much the same for a wide range of other nucleophiles. Two points emerge from the results:(a) that the overall catalytic effect of the nucleophiles is governed almost entirely by the magnitude of Ky^^Q, (values of k^^Q change in the opposite way and their range is quite small), and, (b) for the most reactive reagents, again reaction occurs at the diffusion limit. Even so, thiourea is a much more spectacular catalyst than is chloride ion since ^XNO here is 10^ larger for ONSC^(NH2)2, than it is for CINO formation. An example of the range of catalysis is given in Fig. 1 for the catalytic effects of Br~, SCN" and SC(NH2)2 in the nitrosation of morpholine. There are a

Reagents Ejfecting Nitrosation

Table 3 Values ofKyôlM-'^ at 25°C XNO CINO BrNO ONSCN ONSC+(NH2)2 S-Nitroso-2mercaptopyridine QNS^Q^-

^XNO 1.1 X 10-3 5.1 X 10-2 30 5000 --1 X 105

Reference

1.7x107

[30]

[26] [26] [27] [28] [29]

Table 4 Values of A:xNO^~' ^-^ for reaction with two aniline derivatives at 25°C [32, 33]

CINO BrNO ONSCN ONSC+(NH9)2

C,H,NH, 2.2x109 1.7x109 1.9x108 1.3 X 106

4-NH7C6H4COOH 1.1X109 4.3 X 108 1.4x106 1.8x104

400

Fig. 1. Catalysis by X , (a) bromide ion, (b) thiocyanate ion and (c) thiourea in the nitrosation ofmorpholine.

10


large number of more examples in the literature. Values of ^ ^ N O ^^^ ^BrNO are shown in Table 5 for a larger range of reactants. The detailed references are given in ref [15] p. 15. For (b) the approach of log ky^Q to the diffusion limit is seen clearly in Fig. 2 for both CINO and BrNO in the nitrosation (diazotisation) of aniUne derivatives. As expected the activation energies for the reactions of these two reagents here are much smaller than for the other nitrosating agents. There has been no report of catalysis of nitrosation by fluoride ion, presumably since the A^XNO value for nitrosyl fluoride (a well-characterised species) is much too small. Similarly, although iodide ion catalysis is known and is a significant feature, there has been no independent measurement of ^XNO ^^^ ^Ô, since the latter has a strong tendency to decompose to iodine and nitric oxide. Just as for N2O3 (but not for H2N02"^), it is possible to arrange the appropriate experimental conditions where the formation of XNO is the rate Table 5 Rate constants ^ciNO ^ ^ ^BrNO (^"^ ^~^) ^^^ ^^^ reactions of nitrosyl chloride and nitrosyl bromide in water at 25 °C (except where stated) for a range of reactants Reactant Cysteine Thioglycolic acid Benzenesulfinic acid Benzenesulfinate ion Dimethylamine Glycine 4-Nitroaniline 4-Chloroaniline Sulfanilic acid Aniline 4-Methylaniline 4-Methoxyaniline Hydroxylamine (0°) 1,2-Dimethylindole (3°C) Azide ion (0°C) Morpholine(31°C) N-Methylaniline

^riNO

^RrNO

1.2x106 1.4 X 107 4.6x107 2.4x108 3.1 X 107 1.7x107 2.1 X 108 1.9x109 1.8x109 1.4x109 2.6x109 2,5 X 109 3.0x109 3.4x109 5.1 X 109 3.5x107 9.9x108 1.8x106 -

5.8x104 1.1X106 9.5 X 106 1.2x107 3.6x107 4.3 X 107 2.5x109 9.9x108 3.2x109 1.7x109 2.5x109 2.8x109 3.7x107 4.7x107 5.0x109


11

Encounter limit

Fig. 2. Rate constants for the reaction of (a) nitrosyl chloride, (b) nitrosyl bromide, (c) nitrosyl thiocyanate and (d) S-nitrosothiouronium ion (all generated in situ) with aniline derivatives as a function of^K^,

limiting step. This has been achieved for some very reactive substrates such as some aniline derivatives, azide ion and some thiols. Experimentally this can be detected by the observation of dov^nwârd curved plots for A:obs vs [S]. The full zero-order dependence upon S is not usually achieved. The change in the rate limiting step will occur when k_^ becomes comparable with A:2[S], defined in Eq. (16). Analysis via a double reciprocal plot allows the determination of XNO + 2H2O

HNO2 + H3O+ + X-

(16)

k-i

XNO + S

S-NO + X-

both k\ and k^\lk2 and since k2 values can be obtained from reactions at low [S], all parameters can be deduced. Hence we have a kinetic method for the determination of A^XNO (^ k\lk_\). Results are shown in Table 6 for the reaction of three thiols [34], which form the corresponding S-nitrosothiols upon nitrosation (See Chapter 7). Values of ^1 are reasonably close together for all three nucleophiles and for both thiols, suggesting that they are close to the diffusion limit. A similar set of results for the determination of k^ values was

12


Table 6 Values of A:i/M-2 s-^ and k_ils-^ for the nitrosation of three thiols, together with the derived ^XNO values/M-2 at 25°C

ki{C\-) kiiBr-) ki (SCN-) k_x (CINO) k_i (BrNO) k_Y (ONSCN) i^QNO ^BrNO ^QNSCN

N-Acetylcysteine 2.3x103 4.6x103 6.0 X103 1.8 X 106 7.8 x 10^ 1.8 X 102 1.3x10-3 5.9x10-2 34

Thioglycolic acid 2.9x103 4.0x103 1.2 x 10^ 2.7 X 106 9.6 x 10^ 2.8 x 10^ 1.1x10-3 4.2x10-2 43

Mercaptosuccinic Acid

1.1 x 10^

5.1 x 10^

21

reported earlier [35] for reaction of anilines and azide ion at 0°C where the constancy of the values for X~ = CI", Br~ and T was interpreted in the same way. Now we have k__^ values for the hydrolysis of XNO (or the O-nitrosation of water) where the familiar trend CINO > BrNO > ONSCN is apparent. There is a remarkable agreement between the kinetically derived ^XNO ^^^^^ ^ ^ ^ study and the values measured directly (Table 3), when we consider the inherent errors in analysis via double reciprocal plots, particularly when some of the values depend on the measurement of a small intercept on the (ôbg)"^ axis. 1.2.

Nitrosyl halides HaINO Apart from the generation of nitrosyl halides in low concentration in aqueous acid solutions of nitrous acid containing halide ions, the nitrosyl halides (which are commercially available) can be used in solution in a variety of non-aqueous solvents to bring about nitrosation. At room temperature nitrosyl chloride is an orange-yellow gas (bp -6.4''C) which is highly irritating to the skin, eyes and mucous membranes and requires great care when working with it (fume cupboard). It dissolves in most organic solvents such as alcohols, ethers, chloroform etc. In water it is hydrolysed to give nitrous acid and other derived products such as nitric acid and nitric oxide. It can react violently with acetone. Solutions in non-aqueous solvents are best prepared by reaction of alkyl nitrites and hydrogen chloride, and used in situ. The nitrosyl halides can all be synthesised by direct reaction of the halogen with nitric oxide, Eq. (17), in a reaction which is reversible, so 2N0 + X^ - ^ = ^ 2XN0 that solutions have a tendency to decompose.

(17)


13

Most reactions involving nitrosation with nitrosyl halides have been carried out with nitrosyl chloride. It is a very powerful nitrosating agent and will generate diazonium ions from primary aromatic amines, nitrosamines from secondary amines, alkyl nitrites from alcohols, S-nitrosothiols from thiols, nitrosyl halide adducts from alkenes, nitroso compounds from ketones etc., Eq. (18)-(23). The reactions with alkenes were much used synthetically to characterise alkenes particularly in terpene chemistry [36]. The nitroso-chloro products generally exist as nitroso dimers or as oximes depending on the alkene structure. PhNH2 + CINO —*R2NH + CINO — • ROH + CINO —^

PhN2^

(18)

R2NNO

(19)

RONO

(20)

RSH + CINO — ^ RSNO \ / \ / C=d + CINO — C(NO)QCl)

(22)

R(Me)C=0 + CINO —

(23)

RCOCH=NOH

(21)

Solutions of nitrosyl halides in organic solvents have proved particularly useful in bringing about amide nitrosation (see Eq. (24), for example with a diaryl urea [37]) which are difficult to nitrosate using nitrous acid (see Chapter 2). ArNHCONHAr + CINO = ArN(NO)CONHAr + HCl

(24)

A more recently reported reaction which gives aromatic C-nitroso compounds [38] involves the reaction of organomercurials with CINO, Eq. (25), where X = CI, Br, OAc and R is an aromatic group. When R is aliphatic RHgX + CINO —*-

RNO

(25)

gem-chloronitroso compounds are formed e.g. 2-chloro-2-nitrosocyclohexanol when R is cyclohexyl. Similarly, aromatic C-nitroso compounds can be generated, Eq. (26), from organo tin derivatives in methylene chloride at -25 to O^C [39]. XC^H4SnMe3 + CINO —^

XC6H4NO

(26)

Liquid CINO has had some use as an ionising solvent, for example to generate some nitrosonium salts Eq. (27). Aqueous solutions in neutral or

14


CINO + FeCl3 ^^=^ NO^FeCl4-

(27)

alkaline solution have also proved to be effective reagents when acidic conditions need to be avoided [40]. All of the reactions described thus far are consistent with an electrophilic process, i.e. where the nitroso group is transferred in the NO"^ sense without necessarily forming the free cation. With some bicyclic systems (e.g. norbomene etc.) there is mechanistic evidence that addition to the alkene system occurs via a four-centre transition state, i.e. addition takes place in the syn sense [41]. There has been no major effort elsewhere to determine the stereochemistry of addition. Nitroso compounds can also be obtained from nitrosyl halides by free radical reactions often initiated photochemically. The best known example is the formation of nitroso compound from cyclohexane, Eq. (28), which then OH + CINO

•

•

J

(28)

forms the oxime by proton transfer. This reaction has a major industrial application since the oxime can rearrange in acid solution to give caprolactam which on ring-opening polymerisation yields nylon 6 [42]. Nitrosation by this pathway is often accompanied by the formation of other products, e.g. chloro compounds, as expected for free radical reactions. At higher temperatures than those generally used for nitrosation reactions, nitrosyl chloride will often act as a chlorinating agent, particularly of alkanes. No mechanistic study has been reported, but again free radical pathways are likely. 1.3.

Nitrosonium salts NO"''X" A large number of nitrosonium salts NO^X" have been synthesised and characterised; a number are readily available commercially. The most conmionly encountered examples are the tetrafluoroborate, the hexafluorophosphate and the hydrogen sulfate. All are reasonably stable crystalline materials in dry air and are readily prepared from dinitrogen tetroxide, dinitrogen trioxide or nitrosyl chloride together with a source of the anion, in acetonitrile or nitromethane solvents. In water, nitrosonium salts are rapidly hydrolysed to give nitrous acid and so must be used under anhydrous conditions. Suitable solvents include nitromethane, toluene and acetonitrile. Their structure has been documented using X-ray crystallography, vibrational spectroscopy and


15

conductivity measurements in solution. A complex of N2O4 with 18-crown-6 is believed to be [NO"^.crown.H(N03)2~] and has been used to generate Snitrosothiols in solution from thiols and N-nitrosamines from secondary antiines [43]. As expected, nitrosonium salts are very efficient nitrosating agents. There are many examples in the literature of the nitrosation of amines, alcohols etc. They are particularly useful for the nitrosation of the less reactive species such as amides, Eq. (29), and sulfonanfiides. Primary nitrosamides are readily hydrolysed ArCON(R)H + NO+BF4- = ArCON(R)NO + HF + BF3

(29)

to give the carboxyhc acid. N-Metallation of the amide followed by treatment with a nitrosonium salt in ether at 0°C is also an efficient synthetic procedure BuLi RCON(R')H

•

NO^ RCON(R')Li -^^^^^

RCON(R0NO

(30)

for the nitrosation of secondary nitrosamides, Eq. (30), [44]. Reactions of nitrosonium salts were much studied by Seel and co-workers [45] in a series of papers in the period 1950-7. These included reactions with some inorganic species, such as azide ion in liquid SO2 solvent, Eq. (31), where the products N2O and N2 are clearly indicative of an electrophilic nitrosation reaction. NO+X- + N3- = N2O + N2 + X-

(31)

Mechanistic studies of nitrosation of anilines with nitrous acid at high acidities (> ~6 M HCIO4) have identified NO"^ as the effective reagent [46]. At these acidities physical measurements have shown that nitrous acid is almost quantitatively converted to NO"^. The rate equation, Eq. (32), is very different Rate = )i:[ArNH3'^] [NO+]h^-2

(32)

from that found at low acidities. Reaction occurs at the protonated form of the amine and there is a negative dependence on the acidity (using the Hammett acidity function h^). In addition there is a large kinetic isotope effect ( % / ^ D -^ 10), which suggests that the rate limiting step is now the proton transfer to the solvent S, rather than the attack of the nitrosating species, Eq. (33). It is worth noting that nitrosonium salts can also act as one-electron oxidising agents as well as effecting two-electron nitrosation reactions. Two well-known examples of the former are the oxidation of iodide ion and sulfite ion where nitric oxide is generated, Eq. (34, 35). The salts are widely used in

16


fast ArNH3+ + NO+ ^*=^ ArNH2N0+ + H+ slow ArNH2N0+ + S ArNHNO + SH+

^

(33)

fast ArN2"^

J

2N0+ + 2 r = 2N0 + I2

(34)

2N0+ + SO32- = 2N0 + SO3

(35)

synthetic organic chemistry since the reaction conditions are fairly mild and reactions are quite selective. Among the many examples are the oxidative cleavage of ethers and oximes, initiation of polymerisation reactions, initiation of some condensation reactions and the generation of radical cations, some of which can be isolated as their salts [47]. A number of sulfur compounds including thioureas, thiones and thioketones, when treated with NC^BF^" in acetonitrile, immediately form the disulfide dication and nitric oxide, Eq. (36), 2 C=S

+ 2NO^ — •

C-S-S-d

+ 2NO

(36)

[48]. Some cyclic sulfides yield long-lived radical cations which show (by EPR analysis) evidence of S-S transannular interaction [49]. Occasionally the one-electron reaction occurs when a conventional electrophilic nitrosation reaction is expected. For example some diphenylamine derivatives give the radical cations (which then dimerise) when the nitrosamine product would be expected, Eq. (37), [50]. Ar2NH + NO^ = Ar2NH + NO 1.4.

Ar2NNAr2

(37)

Alkyl nitrites RONO Alkyl nitrites or nitrite esters have been known for a long time and are, under many conditions, effective nitrosating agents. They are mostly colourless-yellow volatile liquids derived in the main from aliphatic alcohols. They are conveniently prepared from alcohols and nitrous acid in the presence of an acid catalyst in an equilibrium reaction, Eq. (38). The alkyl nitrite is


ROH + HNO2

-*

RONO + H2O

17

(38)

generally distilled from the reaction mixture since it has a lower boiling point than the alcohol, resulting from H-bonding interactions. This drives the reaction to completion. They can in principle be prepared from the corresponding alcohol and any carrier of NO"^. There is a recent report [51] which claims that good yields can be obtained when the alcohol is treated with nitric oxide in the presence of air, in organic solvents, when the reagent is probably dinitrogen trioxide N2O3, as in Eq. (39). ROH + N2O3

^

RONO + HNO2

(39)

Alkyl nitrites tend to decompose slowly in air but can be kept indefinitely in sealed glass containers at 4°C. Apart from their ability to act as electrophilic nitrosating agents, they possess powerful vasodilatory and antianginal effects - facts which have been known for over a century. In earlier days alkyl nitrites were used to treat angina attacks. There is even a reference to the use of "nitrite of amyl" in the writings of Conan Doyle (The case of the resident patient). Nowadays the use of alkyl nitrates, particularly glyceryl trinitrate (nitroglycerine) has overtaken the use of alkyl nitrites in this field. However alkyl nitrites have been increasingly used as recreational drugs of abuse, and are constituents of "poppers". It is probable that the biological properties of alkyl nitrites result from release of nitric oxide. This can be achieved in vitro by NO group transfer to a thiol and decomposition of the S-nitrosothiol generated (see Chapter 8), but it is more likely that in vivo, NO release is achieved enzymatically. 1.4.1. Reactions in aqueous acid solution Alkyl nitrites undergo ready acid-catalysed hydrolysis in water. Reactions are much faster than are the corresponding reactions of carboxylic esters, and usually require fast reaction techniques to follow them kinetically. This is of course the reverse reaction of Eq. (38), and sets up the equilibrium position. The synthesis of RONO compounds is usually carried out with excess ROH to drive the reaction to the right. Allen first established the rate law, Eq. (40), for the hydrolysis of Rate = ^[RONO] [H3O+]

(40)

propyl and ^butyl nitrites in 72% dioxan-water at 0°C [52]. This has been confirmed for reactions in water at 25°C [53] when reaction was examined in both directions, starting from nitrous acid and the alcohols, and also later [54]

18


for reactions starting with the alkyl nitrite. It was shown, for a number of substrates, including sulfamic acid, hydrazoic acid, thioglycolic acid, cysteine, N-methylaniline and thioureas, that hydrolysis of RONO occurs rapidly and nitrosation occurs via the protonated form of nitrous acid (H2N02'^/NO"^). As expected the reaction rate was reduced by addition of ROH and at high [substrate] the hydrolysis of RONO becomes rate-limiting. A full kinetic analysis yields the equilibrium constants for RONO formation, Eq. (38), and the rate constants for nitrous acid nitrosation, which are in good agreement with the literature values obtained by direct measurements. For ^butyl nitrite the extent of hydrolysis is so great (and the reaction so rapid) that the kinetics are identical with those when nitrous acid itself is used as the reagent. Strangely neither the acid catalysed hydrolysis of RONO nor the nitrosation of ROH are catalysed by non-basic nucleophiles [54], in spite of the fact that nitrosation of amines, thiols etc. are catalysed in this way. The hydrolysis of RONO is subject to general acid catalysis and there is a solvent isotope effect, suggesting that protonation is involved in the rate-limiting step. This has led to the proposal that reaction is a concerted process involving proton transfer and breaking of the O-N bond in the transition state, Eq. (41).

RONO + HA

r

X

A---H--0—N=0

ROH + A" + NO^

(41)

R

As far as synthesis of nitroso compounds by alkyl nitrites is concerned in aqueous acid solution, there is absolutely no advantage in using alkyl nitrites over nitrous acid, since rapid hydrolysis always precedes the nitrosation step. 1.4,2. Reactions in aqueous basic solution Hydrolysis of alkyl nitrites also occurs by base catalysis, but is a much slower process than the corresponding acid hydrolysis pathway and is also essentially irreversible since the nitrous acid product is converted to nitrite RONO + 20H- - ^

RO" + NO2" + H2O

(42)

anion which is not an effective nitrosating agent, Eq. (42). There is a close analogy with the base catalysed reactions of carboxylic acid esters. Bond fission occurs in the O-N sense as shown by stereochemical and ^Ô experiments. Absence of ^Ô exchange between the ester and the solvent argues against the reversible formation of a tetrahedral intermediate, ruling out an addition-elimination mechanism and favouring a one-step process.


19

Alkyl nitrites have been widely used as nitrosating agents for a large range of substrates. Diazonium ions can be generated from primary aromatic amines, nitrosamines from secondary amines, alkyl nitrites from alcohols (a socalled transnitrosation or nitroso group exchange reaction), S-nitrosothiols from thiols and a number of C-nitroso and derived products. Two recent examples include the ready reaction with ascorbic acid [55], which generates NO (just as the reaction with nitrous acid) and the reaction with malononitrile [56], which occurs via the anion of malononitrile, Eq. (43). The main OHCH2(CN)2 - ^ = ^

RONO

CH(CN)2-

(43)

HON=C(CN)n

advantage in the use of a basic medium, is to avoid the rapid hydrolysis which occurs in acid solution. Synthetically there appears to be no major advantage in using alkyl nitrites rather than nitrous acid, at least in aqueous solution, although alkyl nitrites containing P-electron-withdrawing groups are particularly effective at bringing about nitrosations of amines, Eq. (44), in aqueous alkaline solution [57]. \ OW XCH9ONO + NH

(44)

XCH2OH + NNO / (X = CH2CI, CH2F, CF3, CH2OC2H5, etc.)

Kinetic studies show that reaction occurs via the unprotonated form of the amine [58] and there is generally a large negative entropy of activation, consistent with a highly ordered transition state, probably four-centred, Eq. (45). The nitroso group exchange with an alcohol occurs readily in alkaline

RONO +

/

NH

.01 '^

R- -Or-^ I

I

\ ROH +

NNO / (45)

solution, Eq. (46), and is believed to be an example of 0-nitrosation at the alkoxide ion [59].

RONO + RO"

V RO—N—OR'

RO" + R'ONO

(46)

20


1.4.3. Reactions in non-aqueous solvents There is a large literature describing nitrosation reactions of alkyl nitrites in non-aqueous solvents. Many of these reactions have been adapted to produce good synthetic routes for the formation of nitroso compounds but have not, in general, been much studied mechanistically. t-Butyl nitrite has proved to be an excellent reagent for the nitrosation of amines, alcohols and thiols in acetonitrile or chloroform solution [60] and the more reactive alkyl nitrites such as 2,2,2-trichloroethyl nitrite gave quantitative yields of nitrosamines from secondary amines in a range of solvents including cyclohexane, dichloromethane, 1,4-dioxane, chloroform and dimethyl sulfoxide [61]. Alkyl nitrites in dimethyl sulfoxide containing copper(I) cyanide give products of diazotisation [62] and butyl nitrite with tertiary aromatic amines give the dealkylated nitrosamines among other products [63]. Interestingly alkyl nitrites in mildly basic conditions (K2CO3-DMF) give mainly the substituted anisole derivative, Eq. (47), whereas in acid solution using nitrous

5,

OMe

Base

ÔNO

Acid ^ OMe

^^^^

11

II

^4y^

Y ÔMe NOH

"6L

^ ^ " Y ^

(48)

^^=^"OMe

acid the 4-substituted nitroso compound is the main product, from the phenol derivative, Eq. (48), [64]. Propyl nitrite in acid solution in 1-propanol is only an effective nitrosating agent towards aniline derivatives in the presence of halide ion or thiourea as catalysts [65], which implies the intermediacy of the corresponding nitrosyl halides etc. as the effective reagents. There is kinetic evidence that the protonated form of the alkyl nitrite in alcohol solution will nitrosate more reactive species such as thiols, without the necessity for nucleophilic catalysts [66]. In acetonitrile solution a number of alkyl nitrites react readily with amines, alcohols and thiols, again without the need for a nucleophilic catalyst. Acid catalysis occurs and for the reactions of the alcohols and thiols, the measured rate constant was independent of [substrate] and consistent with rate limiting NO"^ formation, Eq. (49)-(51), whereas for the less reactive amines, the reaction of NO"^, Eq. (51), is rate limiting [67].


RONO + H2SO4 ^S=^ RO(H)NO+ ^^=^

RO(H)NO^ + HSO4-

ROH + N0+

NO"^ + substrate — •

21

(49) (50)

product

(51)

A detailed kinetic study has been carried out [68] of the reaction of several alkyl nitrites (2,2-dichloroethyl nitrite etc.) with a number of secondary amines in a range of organic solvents such as cyclohexane, dichloromethane, chloroform, acetonitrile, dimethyl sulfoxide etc. In all cases the nitrosamine was formed quantitatively. The dependence of the measured rate constant on [amine] was sigmoid. The experimental results were consistent with a mechanism in which a hydrogen-bonded intermediate is first formed which generates a zwitterionic intermediate which can break up to give the RONO + R2NH ^^=^

RONO.R2NH

R'ONO.R2NH T~*" RON

•

+NR2

(52) ROH + R2NNO

(53)

k^

H R2NH products spontaneously and also in a catalytic pathway involving a second amine molecule, Eq. (52),(53). This stepwise mechanism is analogous to that involved in the aminolysis of carboxylic acids. As the polarity of the solvent is increased there is a change towards a concerted mechanism, Eq. (45). There is a good correlation with the Taft function for solvents. Changing the solvent (using acetonitrile-water or dioxane-water mixtures) can bring about a change of rate limiting step. For example in the nitrosation of ureas, in up to 70% acetonitrile-water mixture there is no halide ion etc. catalysis (just like for reactions in water), whereas as the % acetonitrile is increased beyond 70%, catalysis becomes apparent and the rate limiting step is reaction of the nitrosating species with the substrate [69]. In cyclohexane as solvent (non-polar), reactions of N-methylaniline with a number of reactive alkyl nitrites give the expected nitrosamines. Spectrophotometric scanning reveals the build up of absorbance of a reaction intermediate taken to be T^, since in such a non-polar solvent the zwitterionic intermediate T^- will be very

22


R'0^^

^

R'ONf

NHR2

(54) NR2

poorly solvated, Eq. (54), [70]. In recent years there has been much activity, principally from Iglesias and co-workers, on the reactions of alkyl nitrites in microorganised media notably in aqueous P-cyclodextrin solutions and in aqueous micellar solutions. Their work has concentrated on the hydrolysis of alkyl nitrites, both in acid and base solution, and on the nitrosation reactions of amines. There is a comprehensive review of this work [71] which also includes for comparison, a summary of the reactivity of alkyl nitrites in water and organic solvents. 1.5.

N-Nitrososulfonamides RS02N(N0)R' Sulfonamides undergo N-nitrosation to give products of deamination if they are primary, Eq. (55), and stable N-nitrososulfonamides if the structure is RSO2NH2 RSO2NHR'

*" RSO3H + N2

(55)

• RS02N(N0)R'

(56)

a secondary sulfonamide, Eq. (56). The latter, particularly 4MeC^H4S02N(NO)CH3 (MNTS), is well-known as the Diazald reagent for the generation of diazomethane as a methylating agent. In acid solution MNTS breaks down quantitatively to give nitrous and the sulfonamide, Eq. (57), [72]. RS02N(N0)Me

H O"*" —^—•

RS02NHMe + HNO2

(57)

The reaction rate is unaffected by both traps for nitrous acid (sulfamic acid) and nucleophilic catalysts. Reaction is subject to general acid catalysis and there is a kinetic solvent isotope effect k^^kj^ of 1.5. The facts are consistent with a mechanism in which the proton transfer from the solvent to the nitrogen

+ RS02N(N0)Me + HA ^5==^ RS02NH(NO)Me + A"

(58)

atom, Eq. (58), is rate limiting. This behaviour is similar to that encountered in the hydrolysis of nitrosamides [73] and in the diazotisation of some deactivated amines (see chapter 2). The consequence is that just as for alkyl


23

nitrites any nitrosation reactions proceed via hydrolysis and ttie active reagent is nitrous acid. In neutral or basic solution however, hydrolysis is much slower, again just as for the alkyl nitrites, and this allows the possibility of direct NO group transfer. Leis and co-workers [74-6], have shown that reaction does indeed occur with a range of nitrogen, carbon, oxygen and sulfur-centred nucleophiles. Typical examples include a number of primary, secondary and tertiary amines, azide ion, hydrazine, carbanions (or enolates) derived from dimedone, 2-4pentanedione, nitroethane, 2,2,2-trifluoroethoxide ion, hydroperoxide ion, sulfite ion, thiourea and thiosulfate ion. In many cases the products were as expected for a conventional electrophilic nitrosation reaction, but for the 'harder' nucleophiles (including 0H~), reaction occurred at the sulfur atom. Thus MNTS shows ambient behaviour in this regard [77]. Correlation of the rate constant data with various nucleophile parameters was examined, the best fit being with the Ritchie N^ scale, which suggests a frontier orbital-controlled reaction with a significant diradical element in the transition state. Reaction with cysteine in 25% EtOH-water solvent, occurs via the anion and the initial product is the S-nitroso species [78]. Substituent effects have been studied. The reagents containing electron-withdrawing groups (X = CI or NO2) are particularly reactive, Eq. (59), and a powerful case has been made ^

V-s02-N(N0)Me + S

•

X—/

y-S02NHMe

+ S-NO^

(59)

[79] for the use of these reagents as effective nitrosating agents due to their relative stability to hydrolysis compared with alkyl nitrites. They are particularly "good" reagents within the pH range 3-8 where there are difficulties with other conventional reagents. 1.6.

Nitrogen oxides

1.6.1. Nitric Oxide NO Nitric oxide or nitrogen monoxide is not a conventional two electron transfer electrophilic nitrosating agent and requires oxidation to nitrogen oxidation state (+3) before it can become so. However because of its ready oxidation in air to give NO2/N2O4 (which can effect nitrosation, see 1.6.2) there has been much confusion in the early literature, particularly the biological literature, where experiments indicating nitrosation have occurred readily because of the failure to eliminate completely oxygen from the system. Under

24


fully anaerobic condition however, nitric oxide itself will not bring about nitrosation. Nitric oxide is of course a stable free radical and under some circumstances, will react with, for example, a carbon centred radical to generate a C-nitroso compound. A classic example of this is the reversible (C^H5)3C* + NO ^ i = ^

(C6H5)3CNO

(60)

reaction in ether with the triphenylmethyl radical, Eq. (60), [80]. The nitroso product is blue and the reactants can be recovered by solvent evaporation. Nitric oxide has been widely used in gas phase reactions to trap out, and thus demonstrate the existence of, radical intermediates, particularly carbon-centred radicals. Radicals generated photochemically react with nitric oxide in situ to give e.g. the oxime from cyclohexane, Eq. (61), [81]; alkenes behave similarly, Eq. (62), [82]. ^^2^

r

1*

NO

CF2=CH ^ ^ ^ C F 2 = C F

^ ^

hv

a -a CF2=C(N0)F

(61)

(62)

Diethylamine can be converted to the nitrosamine by nitric oxide in the presence of Cu(II) salts, with the probable intermediacy of a copper nitrosyl complex, Eq. (63), [83]. Similarly, silver ion allows the formation of the CuCl2(aq) + NO ^^=^ Cu(II)NO complex Cu(II)NO complex + Et2NH

•

(63)

Cu(I) complex + Et2NNO

nitrosamine from nitric oxide, possibly via the intermediate formation of a cation radical, Eq. (64), [84]. Iodine will generate nitrosyl iodide (a good

R2NH

Ag"^ •

+• R2NH

MO i l i ^ R^NNO + H+

(64)

nitrosating agent) from nitric oxide, again allowing conventional electrophilic nitrosation to occur, Eq. (65), [85].

Reagents Ejfecting Nitrosation

2NO + I2 ^^=^

2IN0

R^NH ^

R2NNO

25

(65)

A reaction of nitric oxide with another free radical is believed to be important in the context of the in vivo biological properties of nitric oxide. It reacts with the radical anion the superoxide anion to give the peroxynitrite NO + 02* = ONOO-

(66)

anion (a 0-nitroso compound), Eq. (66), at a rate very close to that expected for a diffusion-controlled process [k = 4.3 - 6.7 x 10^ M~^ s"^), [86-8]. It has been suggested that this reaction might be responsible for removal of excess toxic nitric oxide in the body, but this view has been challenged and the matter is not resolved. The use of nitric oxide in solution in the presence of air has recently become an important nitrosating pathway. It has the advantage of potential use in a range of organic solvents as well as in water over a range of pH values. Thus, amines react to give the expected products [89], amides generate Nnitrosamides [90], thiols give S-nitrosothiols [91], alcohols yield alkyl nitrites [51] and hydrogen peroxide forms peroxynitrous acid [92]. The generally accepted mechanism here is that nitric oxide is rapidly oxidised to nitrogen 2NO + 02 = 2NO2 NO2 + NO = N2O3

y

(67)

N2O3 + S = +SNO + NO2" dioxide, which reacts with further nitric oxide to give the well-known nitrosating agent dinitrogen trioxide, Eq. (67). This can occur in most organic solvents and in water irrespective of the pH except that the final step must compete effectively with the hydrolysis of N2O3, which in basic media would give the non reactive nitrite ion. It appears, and there are rate constant data to back this up, that the reaction of NO2 with NO competes very favourably with its own hydrolysis to nitrite and nitrate ions. In some cases hydrolysis of N2O3 competes with the nitrosation, thus for the reaction of morpholine at pH 7.4 [89], a mixture of nitrite ion and N-nitrosomorpholine accounts for the "total nitrogen" yield. Interestingly for both the reactions of thiols (and one amine) [91] and hydrogen peroxide [92], at pH values >7.5, the rate equation, Eq. (68), is identical with that found for the autoxidation of nitric oxide in aerated Rate = k\HO]\0^

(68)

26


water, and there is also good agreement between the measured rate constants. The reaction rate is independent of the substrate concentration and the rate limiting step must be the same as that for the autoxidation of nitric oxide. There are many examples where metal nitrosyl complexes can act as nitrosating agents. These will be discussed in Chapter 9.

1.6.2. Nitrogen dioxide NO2 Nitrogen dioxide is also a free radical, but in contrast to nitric oxide, does readily dimerise, setting up the equilibrium, Eq. (69), which has been 2NO2 ^^^^

N2O4

(69)

much studied by physical chemists. Any future reference to nitrogen dioxide is taken to mean the mixture NO2/N2O4. Most of the reactions of dinitrogen dioxide involve the formation of intermediate free radicals. For example in many cases the addition to alkenes generates dinitro or nitro-nitrite products. Aromatic nitro compounds can also be synthesised from nitrogen dioxide, providing an alternative route to that involving the nitronium ion N02"^. There is a review of the literature up to 1964 by Schechter [93], in which some of the early puzzling results are rationalised. Under certain circumstances however, N2O4 can react in the NO"^03~ sense and is thus an electrophilic nitrosating species. In high acid concentration of sulfuric and perchloric acid, the conversion to the nitrosonium ion is essentially quantitative and the spectroscopic properties of NO"^, notably the 2300 cm~^ band in the Raman spectrum, have been observed. In relatively polar solvents such as chloroform or ether, heterolytic fission can occur leading to nitrosation products [94]. Dinitrogen tetroxide has been used to nitrosate amides, when more conventional reagents have proved unsuccessful [95]. It also has been used to effect deamination of primary amines, giving fewer by-products than is the case with nitrous acid deamination [96]. There is a review which also covers reactions with alcohols and thiols [97]. There are more recent references to nitrosation of both amines and thiols by NO2/N2O4 [98-9]. Frequently, dinitrogen tetroxide yields products of nitration and nitrosation simultaneously, which makes it less useful as a synthetic reagent. Secondary amines in dichloromethane or acetonitrile give both N-nitroamines and N-nitrosamines, Eq. (70), and alcohols yield alkyl nitrates as well as alkyl nitrites, Eq. (71), [100-102]. Alkenes can yield nitroso nitrates at low temperatures in liquid ethane-propane as solvent [103]. Recently Ridd and coworkers [104] have shown that the reaction with hexenes in hexane solvent


R2NH-

N2O4

27

-^R2NN02 (70)

-^R2NN0 ROH-iiÔi

-RONO2 (71) -RONO

yields the dinitro compound as the major product with smaller amounts of nitronitrite and nitroalcohol, whereas in chloroform the main product is by far the nitrosonitrate, Eq. (72). hexane \

/

^C(N02)q(N02) + Other products

i:J20i \ chloroform

/ ;C(N0)q(0N02) (72)

A number of N2O4 complexes e.g. with copper(II) nitrate [105] and also with silica-polyethyleneglycol [106], act as nitrosating agents, particularly with thiols, although the final isolatable product is the disulfide derived from firstformed S-nitrosothiol. Nitrogen dioxide encapsulated in a calixarene, produces a stable complex which is particularly suitable for bringing about the nitrosation of secondary amines [107]. There is one reference in the literature [108] to a kinetic study of the reaction of N2O4 with aniline in donor non-aqueous solvents such as acetonitrile and ethyl acetate. The reaction was claimed to be kinetically firstorder in N2O4 and zero-order in aniline, suggesting that the rate limiting step is the formation of some nitrosating species (NO"^ or a derivative of) from N2O4, prior to its reaction with the amine function. 1.7.

Miscellaneous reagents The chemical literature contains a large number of references to other nitrosating species which are effective with certain substrates under certain conditions. In general these have not been examined mechanistically so very little is known about their reaction mechanism. Nitro compounds Tetranitromethane, C(N02)4, has been used to bring about nitrosation and also nitration reactions. Morpholine is converted in good yield to Nnitrosomorpholine when heated in tetrahydrofuran for four hours at 70°C,

28


CJ O^

C(N02)4

r

>

(73)

THF, l(fC

Eq. (73), [109]. The tertiary aromatic amine N,N-dimethylaniline yields Nmethyl-N-nitrosoaniline by heating in pyridine [110]. Under these rather forcing conditions, it may be that homolysis occurs generating NO2/N2O4 which could nitrosate amines. Similarly the antibacterial and antifungal reagent 2-bromo-2nitropropane 1,3-diol converts diethanolamine to the N-nitroso compound at around pH 12, Eq. (74), [111]. (CH20H)2C(Br)N02 + (HOCH2CH2)2NH —

(HOCH2CH2)2NNO

(74)

There is an example of an inorganic nitro complex which acts as a nitrosating agent for aromatic amines [112]. Sodium hexanitrocobaltate(III) converts aniline derivatives to the 1,3-diaryltriazene, Eq. (75), almost certainly via the intermediacy of the diazonium ion and subsequent coupling with the ArNH2 + Na3Co(N02)6 —^

ArNH-N=N-Ar

(75)

reagent. This reagent also converts hydrazine derivatives to the azides, but will not react with aliphatic amines because of interfering complex formation. Inorganic nitrates There are a number of reports of nitroso compound formation from inorganic nitrates. The reagent with the best synthetic potential is probably the Clayfen reagent, which is K-10 clay-supported ferric nitrate. Alcohols are readily converted to alkyl nitrites at room temperature in hydrocarbon solvents [113], and thiols to the corresponding S-nitrosothiols [114], although the isolated products are again the disulfides formed by the breakdown of the Snitrosothiols. This reagent will also convert hydrazines to azides [115]. Nitrite ion Normally nitrite ion has no capacity for effecting nitrosation. However in the presence of some carbonyl group-containing catalysts, nitrosation can be achieved in neutral and basic solution [116]. Catalysts include formaldehyde, chloral, benzaldehyde derivatives and pyridoxal. Reaction is particularly effective for the conversion of secondary amines to nitrosamines, Eq. (76). Product formation was rationalised by suggesting the formation of an iminium


29

OHR2NH + R'CHO -* R2N=CHR' + H2O + R2N=CHR' + NO2" — ^ R2N-CHR'0N0 R2N-CHR'0N0

—^

(76)

R2NNO + R'CHO

ion intermediate, which reacts with nitrite ion giving a dialkylamino nitrite ester which breaks down rapidly and intramolecularly to give the nitrosamine and regenerating the catalyst. This pathway has also been suggested for the reaction of secondary amines with solid sodium nitrite in halogenated solvents [117]. Fanning and Keefer [118] have shown that bis(triphenylphosphine) nitrogen(+l) nitrite [N(PPh3)2]"^N02~ in dichloromethane generates an intermediate which converts secondary amines to nitrosamines in high yield. Nitrosyl carboxylates (Acyl nitrites) RCOONO It is generally believed that when sodium nitrite is dissolved in a carboxylic acid, it generates an equilibrium concentration of the nitrosyl carboxylate, which acts as the nitrosating agent [119]. Many nitrosating reactions have been carried out preparatively by the use of e.g. sodium nitrite dissolved in acetic or formic acids, Eq. (77). Subsequently, nitrosyl acetate HNO2 + ^ 3 ^ ^ + RCOO" ^ ^ = ^

RCOONO + 2H2O

(77)

CH3COOAg + CINO = CH3COONO + AgCl

(78)

(acetyl nitrite) has been synthesised from silver acetate and nitrosyl chloride [120], Eq. (78), at liquid nitrogen temperatures. It is a pale brown liquid at room temperature, a green liquid at -78°C and a green solid at -196°C. It is rapidly hydrolysed in water but is sufficiently stable in solvents such as pyridine or acetic acid to act as an effective nitrosating agent. It converts 1octanol to the nitrite ester in pyridine, Eq. (79), and gives a similar distribution of deamination products from 1-octylamine as does the use of the reagent sodium nitrite/acetic acid, Eq. (80). CH3(CH2)6CH20H + CH3COONO — ^ CH3(CH2)6CH20NO

CH3(CH2)6CH2NH2 — "^

octenes CH3(CH2)5CH(CH3)OH • CH3(CH2)5CH(CH3)OOCCH3 CH3(CH2)5CH20H CH3(CH2)6CH200CCH3

15% 6% 20% 3% 48%

(79)

(80)

30


Kinetic results for reaction in the presence of carboxylate buffers often show a pathway that takes place via the nitrosyl carboxylate. Stedman [121] showed that the reaction between nitrous acid and hydrazoic acid in acetate buffers has two pathways, involving both nitrosyl acetate and dinitrogen trioxide as the nitrosating species. Later [122], more detailed kinetic work on the nitrosation of N-methylaniline and piperazine under similar conditions, has revealed pathways via CH3COONO, N2O3 and H2N02"^/NO"^. By comparison with other substrates, it appears that CH3COONO might react at the diffusion limit, which enables an equilibrium constant for the formation of nitrosyl acetate, Eq. (77), to be estimated as -1.4 x 10~^ M~^ Such a small value would not allow the observation of nitrosyl acetate in these solutions, by any spectroscopic technique. Sodium nitrite in trifluoroacetic acid has been shown to be effective in the nitrosation of water-insoluble amides [123]. The same reagent has been used successfully to bring about diazotisation (and subsequent reaction with azide ion) of highly deactivated aniline derivatives such as perfluoroaniline and 2,6-difluoroaniline, Eq. (81). Similarly sodium nitrite in anhydrous propionic

NaNO, ^

.

„

_

^

,

„

(31)

CF3COOH acid generates 2-nitroso products from phenols regioselectively [124], maybe via the in situ formation of nitrosyl propionate. Fremy's salt K2[(SO^Ô] This was first prepared by Fremy as early as 1845 [125]. The stable yellow solid contains the dimeric anion, whilst in solution (which is violetblue) the monomeric ion exists as a relatively stable free radical. It has mainly been used in organic chemistry as a specific oxidising agent involving a oneelectron transfer, notably for the oxidation of phenols and anilines to quinones [126]. In addition, however, secondary and tertiary amines can be nitrosated either in pyridine solution or in aqueous sodium carbonate [127]. Nothing is known about the mechanism of the reaction but it is possible that the tertiary amine generates a radical cation which, after proton loss undergoes C-N bond fission to give the secondary amine (which then undergoes nitrosation), and the aldehyde product.


31

\ N'Nitrosamines N—NO / Nitrosamines are not so reactive as nitrosating agents as are their counterparts the alkyl nitrites, probably because of the strength of the N-N bond. However there are a number of cases which have been reported which show that nitrosation can occur using nitrosamines. The hydrolysis (or denitrosation) or nitrosamines has been well studied. Reaction occurs in an acid-catalysed process generating nitrous acid. This is of course the reverse f^NO

—-—^

/

/

NH + HNOo ^

^^^^

reaction to the synthesis of nitrosamines and the equilibrium position lies well over to the left hand side. Nevertheless the denitrosation reaction can be studied if steps are taken to remove the nitrous acid at such a rate that it is much faster than the reverse reaction of Eq. (82). This can be achieved by the use of so-called 'nitrous acid traps' such as urea, sulfamic acid, hydrazine, hydroxylamine, hydrazoic acid etc. So long as there is sufficient trap present to ensure that reversibility does not interfere, denitrosation occurs smoothly in a process which is independent of the nature and concentration of the nitrous acid trap. Other non-basic nucleophilic species X~ can now be added and the rate of formation of XNO can be measured. Apart from the halides and thiocyanate ion, this kinetic study revealed that sulfur nucleophiles were "good" nucleophiles, including cysteine, glutathione, methionine, thiourea and alkyl thioureas [128-9]. A number of aliphatic heterocyclic nitrosamines including piperazine derivatives react similarly. On the synthetic side, there are many examples of reactions whereby nitrosamines transfer the NO group to amines, alcohols, thiols etc. In many cases, particularly those taking place in aqueous solution it is likely that reaction takes place by prior hydrolysis to give nitrous acid which can effect nitrosation of the added substrate. An equilibrium situation can develop, Eq. (83), and the position of equilibrium will depend on the relative concentrations \ HgO^. R2NNO -*

NH RjNH + HNO2

*

/

NNO

(83)

of the nitrosamine and added amine and also on their relative N-N bond strengths. Such reactions have been observed for aliphatic [130], alicyclic [131] nitrosamines as well as for nitrosoureas [132]. The indirect nature of the transfer can be demonstrated by interception of the nitrous acid generated by denitrosation, by a substrate more reactive than the amine [133].

32


The denitrosation reaction has been used to generate 4-nitrosophenol which condenses with a second phenol molecule to give a quinone-imine which is coloured bright red in acid solution and blue in alkaline solution, Eq. (84). This is the chemistry behind the Liebermann qualitative test for a nitrosamine \

H30'^ f^NO

^

\ NH + HNO2

phenol ^

Of^

^^—\ V-OH

(84)

phenol

Blue

Red

[134], or an alkyl nitrite. A direct NO transfer can also occur i.e. without the intermediate formation of nitrous acid or any other nitrosating species. NNitrosodiphenylamine reacts in acid solution with N-methylaniline to give Ph2NNO + f V NV~N(Me)H ^ ( M e ) H ,c

^ "^ Ph2NH + f f >V-N(Me)NO - r (85)

N-methyl-N-nitrosoaniline, Eq. (85), in a reversible reaction which is not catalysed by halide ion or thiocyanate ion, under conditions where the denitrosation of Ph2NN0 is markedly catalysed by these anions [135]. Later, in an attempt to avoid the reversibility [136], the reaction with aniline derivatives was studied, when the diazonium ions are formed. The dependence of the rate constant on the acidity suggested that reaction occurred (unusually) with the protonated form of the aniline. S-Nitrosation of a cysteine residue in proteins by N-methyl-Nnitrosoaniline has also been reported [137], although it is not known if this is a direct reaction. Nitrosation by nitrosamines can also be brought about in non-aqueous solvents. Thus indoles undergo C-nitrosation, Eq. (86), when treated with Nnitrosodiphenylamine in a mixture of chloroform and trichloroacetic acid as solvent [138]. Aromatic N-nitrosamines also undergo the Fischer-Hepp rearrangement to give the 4-nitroso compound. This reaction will be discussed folly in Chapter 3.


VR

+ PhjNNO

f T V R

33

(8«)

Alkyl Sulfur-nitroso compounds Both thionyl chloronitrite SOCIONO and thionyl dinitrite S0(0N0)2 have been prepared [139], by reaction of thionyl chloride with silver nitrite in dry tetrahydrofuran. Both give excellent yields of alkyl nitrites from alcohols, and oximes from carbonyl compounds, Eq. (87), and are claimed to be the RCH^COR' + SOCIONO —

RC(=NOH)COR'

(87)

most efficient nitrosating agents yet synthesised. There are a number of sulfur compounds which act as very powerful catalysts in nitrosation reactions. These include thiourea and alkylthioureas, thiosulfate ion, dimethylsulfide and a range of thiones. In each case the \+ effective nitrosating species is almost certainly a species containing the S—NO entity. These are not reagents which have been isolated and used in synthesis reactions, but nevertheless constitute an important family of nitrosating agents generated and used in situ. Their chemistry will be discussed later in Chapter 7. S-Nitrosothiols (thionitrites) RSNO can be isolated in some cases and will effect nitrosation of a range of species They are usually however, also generated in situ, and again their chemistry will be discussed in detail in Chapter 8. S-Nitrothiols (thionitrates) RSNO2 can also be effective nitrosating agents in neutral aprotic solvents. Aromatic amines in the presence of copper(II) halides yield aryl halides, presumably via the intermediate formation of the corresponding diazonium ion, and secondary amines give nitrosamines [140]. It was suggested that reaction might occur by way of rearrangement of the nitrothiols to the isomeric nitrite form RSONO, although there is no definitive evidence on this point. Other reagents There has been an upsurge, in recent years, in the range of rather unusual nitrosating agents which have been developed. Some do not appear to have major synthetic advantages over the reagents already discussed, but others might be appropriate reagents under certain circumstances. These include sodium nitrite under phase-transfer conditions [141], oxyhyponitrite [142],

34


various complexes of HalNO and N2O4 [143 and references therein], sodium nitrite on silica and silica chloride bases [144] and N2O4 immobilised on polyvinylpyrrolidone [145].

Nitrosation Reactions and the Chemistry of Nitric Oxide D.L.H. Wilhams © 2004 Elsevier B.V. All rights reserved.

35

Chapter 2

Nitrosation at nitrogen centres By far the best-known and most widely studied nitrosation reactions are A^-nitrosation of amines and related compounds. Many of these reactions have been known for a long time, indeed one of the very first examples reported, involved the conversion of aspartic acid into malic acid in 1846 by Piria [146]; deamination of simple aliphatic and also of aromatic primary amines was achieved soon thereafter [147]. A^-Nitrosations occur widely in organic chemistry in many standard synthetic pathways, and a number (including diazotisation and azo dye formation) have large scale industrial applications. Because of their importance, such reactions have also been much studied mechanistically, and in general are now well understood. Many reactive intermediates in nitrosation have been identified in kinetic studies, using amine (or closely related) substrates. It is worth noting that a large range of mechanistic features can appear in nitrosation. These include acid catalysis, base catalysis, nucleophilic catalysis, reversibility, encounter-controlled reactions, intramolecular rearrangements, C-C bond cleavage reactions, and /pô-attack. Many of these have been demonstrated unambiguously in reactions involving attack at the nitrogen atom. The whole area has been a particularly fruitful one for mechanistic investigations [5,148]. Since the discovery that nitrosamines are powerful carcinogens in all animal species which have been tested, the nitrosation of secondary (and tertiary) amines has been very widely studied, particularly from the viewpoint of the possible in vivo formation of nitrosamines from naturally occurring secondary amines and sources of nitrous acid in foods and in water supplies. 2.1.

Primary aromatic amines Reaction of nitrous acid with primary aromatic amines has probably been the most widely studied nitrosation reaction, both from the synthetic and mechanistic point of view. It is known that N-nitrosation occurs which is followed by proton transfers and loss of a water molecule to give the relatively stable arenediazonium ion RN2"^, Eq. (88). The diazonium ion is a very HNO2 • RNH2^N0 ^ i = ^ RNHNO T'^ HsO^ _, + RN=N-OH ^ >• RN=N-OH2 • RN=N RNH2

RN=N-OH (88) + H2O

important intermediate in organic chemistry and reacts as an electrophile with a large variety of nucleophiles to generate usefiil products, e.g. haloaryl

36


derivatives, azo dyes, phenols, cyano compounds etc., and are widely discussed in standard organic chemistry texts. Some primary aromatic nitrosamines have been detected spectroscopically at low temperature [149] and a number of heterocyclic primary amines yield stable primary nitrosamines, Eq. (89). Their stability has N—N

N—N NHNO

Ar

(89)

Ar

been attributed to internal H-bonding in the diazohydroxy form of the nitrosamine [150]. Reactions are usually carried out in aqueous acidic solutions of sodium nitrite at ~5°C, and any subsequent reactions of the diazonium ions carried out in situ, sometimes after a pH change. For amines less soluble in water, reaction can be carried out in mixed alcohol-water solvents or in non-aqueous solvents using alkyl nitrites. Some salts of diazonium ions can be isolated and are reasonably stable in the solid form, whereas others decompose, often explosively, so care is needed when they are isolated. The tetrafluoroborate ArN2^BF4~ is particularly stable and can be obtained commercially. There are a number of excellent, comprehensive reviews of the diazotisation reaction, notably by Zollinger [151-4] who has made a major contribution to this area of chemistry. He has also produced a review of dediazonisation reactions of arene diazonium ions [155]. A number of mechanisms have been identified, including:(a) a unimolecular decomposition to give the aryl cation and nitrogen, Eq. (90), which may include both a heterolytic and homolytic pathway, ArN2^ —

Ar"^ + N2

(90)

(b) a bimolecular mechanism involving synchronous attack by the nucleophile and nitrogen loss, and, (c) an elimination-addition pathway in which the intermediate is a carbene. Arenediazonium ions can also act as a source of aryl radicals and the procedure has been used to effect arylation of aromatic and other compounds. Cyclisation reactions are also known, generating heterocyclic species. As with most nitrosation reactions, diazotisation of aniline derivatives is catalysed by the addition of non-basic nucleophiles. For many nucleophiles

Nitrosation at Nitrogen Centres

37

(X~), the extent of catalysis can be quite spectacular, e.g. for SCN", SC(NH2)2 and I~, resulting from the large equilibrium constant ^XNO ^^^ ^'^ formation of XNO for these nucleophiles [33]. The second-order rate constants for reaction of XNO with the free base form of the aniline correlate quite well with the p^^ values of the protonated anilines except for the most reactive BrNO and CINO, where the values tend to level off at the higher ^K^ values towards the limit expected for a diffusion controlled process, as shown in Figure 2 (p. 11). A similar pattern was observed for the reaction of nitrosyl chloride with amines in non-acidic aqueous solution using a competition method to get the rate constants for CINO attack relative to that of CINO hydrolysis (to N02~ and C r ) . The product ratios [RN2"^]/[N02~] give the relative rate constants, which tend to level off for amines where the pX^ values are > --3 [40], as shown in Figure 3. Another feature which emerges from the kinetic data is that it is possible, for very reactive amines at relatively high concentration of the free base form (i.e. at low acidity), to change the rate limiting step to that of the formation of XNO. This has been achieved for aniline itself for both bromide and iodide ion-catalysed reactions [156]. In addition, at high [X~] and in particular for anilines containing electron-withdrawing groups, X~ catalysis can disappear. For example, the reactions of 4-nitro, 2-nitro, 4-carboxy, 4-chloro and 3-methoxyanilines showed curved plots of observed rate constant vs [Br""], and 2,4-dinitroaniline showed no evidence at all of nucleophilic catalysis. This has been interpreted [32] in terms of a reversibility of the initial nitrosation reaction, Eq. (91). If ArNH2 + BrNO

,

^1

*

+ ArNH2N0 + Br"

(91)

several stages ArN2^ ^_2[Br~] and ^2 ^^^ comparable, then curved plots should follow, and if ^_|[Br~] » k2 then bromide ion catalysis should disappear as the proton transfer from Ar"4s[H2NO to the solvent becomes rate limiting. This changeover is more likely at the higher [Br~] levels and also when Ar contains electron-withdrawing substituents (expected for a S^2-type reaction at the nitroso nitrogen atom). The same situation exists for the nitrosation of amides which will be discussed later in this chapter. The same also applies to


38

O

-4

-2

4

6

O

10

12

Fig. 3. Rate constants for the reaction of nitrosyl chloride (relative to its hydrolysis reaction) with amines as a function of pATg.

diazotisation of 1-naphthylamines, where again at high [nucleophile], curved plots are found [157-8]. The tendency for reversibility of Ar~^NH2N0 formation is much more marked in methanol (and probably in other alcohol solvents) than it is in water. Curved plots of the observed rate constant vs [Br~] also occur for reaction using hydrogen bromide as the catalyst [159], and for some aniline derivatives for diazotisation in MeOH/HCl the rate limiting step is the deprotonation of RC^H4^TvJH2NO ^*^^^ * ^ reagent is CINO. Little is known about the diazotisation of diamines. One interesting example is the reaction with 2,3-diaminonaphthalene. The final product here is the triazole derivative, Eq. (92), which is a fluorescent material and this


39 (92)

HNO2 )

property is the basis of one analytical procedure for the determination of nitrite/nitrous acid [160]. A kinetic study of the diazotisation [161] shows that both the free base and the monoprotonated forms of the reactant are involved the latter predominating at the higher acidity. The reaction with the free base form occurs close to, or at, the diffusion limit and the presence of a 2-NH3'^ substituent reduces reactivity by about a factor of 800. Presumably the diazonium ion undergoes cyclisation in a rapid step to give the triazole product. At much higher acidities, it is possible to generate the doubly-diazotised ion from 2-aminoaniline, Eq. (93). NH2

NH2

NO^

N2' (93)

NO

NH2

N2^

N.

The most widely used quantitative test for nitrite/nitrous acid is undoubtedly the Griess test [162]. This is based on the diazotisation of an aniline derivative followed by a coupling reaction which generates an azo dye with a high extinction coefficient. The aniline derivative much used in the analysis is sulfanilamide and the coupling agent 1-naphthylamine which has been replaced in Shinn's modification by N-(l-naphthyl)-ethylenediamine [163]. Other modifications have included the diazotisation of 4-methylaniline and coupling with 2-naphthol-3,6-disulfonic acid. The first azo dyes to be produced commercially were derived from aniline derivatives, and colour changes were brought about by substituent group changes, both in the anilines and in the coupling reagents. The latter usually contained sulfonic acid groups in order to confer water solubility on the final dye. Later, the range was extended to include exocyclic heterocyclic primary amines, since many had better colour properties. Some commercially useful examples are shown below in (94). Industrially, diazotisation is usually carried

Xx

XX

JD

,NH2

n H (94)

40


out using nitrosyl sulfuric acid in moderately concentrated sulfuric acid. Generally, nucleophilic catalysis occurs so there is the scope for the use of milder conditions. The amine group here is much less basic than it is in the aniline derivatives, but in many cases ring-protonation occurs. Much of the work is described in the patent literature. 2.2.

Primary aliphatic amines Aliphatic primary amines also undergo ready nitrosation with the usual range of nitrosating reagents to generate the diazonium ion. Reactions in acid solution are usually much slower than are those of the corresponding aromatic amines, since the aliphatic amines are more basic and therefore a much smaller fraction of the reactant is present as the free base form. The diazonium ions are now much less stable, since the positive charge cannot be delocalised into an aromatic system. Loss of the nitrogen molecule occurs readily in solution, generating a source of carbocations, Eq. (95), which can undergo a variety of RNH2

•

RN2^

•

R^ + N2

(95)

further reactions. There are one or two cases where the diazonium ion has been detected spectroscopically, when there is a major stabilising factor such as the presence of strongly electron-withdrawing groups such as in the hexafluoropropane-2-diazonium ion [164]. There is no comparable chemistry to that of the arenediazonium ions and the reaction is usually referred to as the deamination reaction. The reactions of primary aliphatic amines have also been reviewed by Zollinger [165]. In general the carbocations can generate alcohols or ethers by attack of the solvent SOH, alkenes by proton loss, or rearrangement products resulting from either R or H migration, see (96); quite often there is a large spread of products so that in marked contrast to the reactions of primary aromatic amines, the reactions of the aliphatic amines have no significant synthetic potential. One of the problems concerning this simplified mechanistic scenario is that the product distribution is sometimes quite different from that observed when the same carbocations are generated by other means, e.g. by solvolysis of a halide or a tosylate. One explanation is that the branching point for the different product routes is not the free carbocation itself, but rather its immediate precursor the diazonium ion R N / [166]. In organic solvents reaction takes pmce with e.g. alkyl nitrites, usually generating a smaller range of products, but very rarely a single product, so again synthetically there is not much potential here. The position was reviewed [167], noting that in solvents such as nitromethane, chloroform and a range of


41

^ I —C—C—OS ;—C—NHo H

?u

HN02 ^

IT

'

H

H migration.

SOH

H -H+

'

I

X

R migration

4

H

SOH

H

rr H

-H^

SOH R

R >

SOH=hydroxylic solvent, H2O, ROH etc.

~3. The rate law, Eq. (100), has been identified, where h^ is (100)

Rate = ^[ArNH3+][HN02]/2o

the Hammett acidity function, for reactions in the acid region 0.1-6.5 M perchloric acid [169-70]. This was interpreted by Ridd as reaction taking place now via the anilinium form of the reactant in the series of steps outlined in

H2N02^ (or N0+) + fX

RNH HA "^

f^ ^ ^

X = N02,0H, Hal(R = COCH3), NRPh(R = H),etc.

(136)

The early experiments showed that:(a) the yields appear to be better with hydrogen chloride than with other acids, (b) a number of nitrosation products, e.g. addition products of CINO to alkenes, could be trapped out from the reaction mixture, and, (c) the reaction of the 3-nitro derivative in the presence of urea gave only the product of denitrosation i.e. the secondary amine. These observations were rationalised in 1913 by Houben [238] in terms of a mechanism set out below where denitrosation occurs, promoted by chloride ion, and the released CINO effects 4-substitution of the amine to give the final product, Eq. (137-8), in an electrophilic aromatic substitution. R>JNO

RNH + HCl

:i==^

r

I

+ CINO ( - • decomposition)

(137)

60


RNH )

RNH +C1NO

- ^

^

^YiC\

(138)

NO

(139) NO2 This outline mechanism had achieved general acceptance until the 1970s and was given in many textbooks, although some authors commented that the full experimental evidence was lacking [239]. Later it was noted by two independent groups [240-1], that rearrangement occurred even in the presence of quite high concentrations of nitrous acid traps such as urea or sulfamic acid. It was also shown that 3-nitro-N-methyl-N-nitrosoaniline gave no rearrangement product even in the absence of added urea, Eq. (139). Rate measurements showed that for the reaction of N-nitrosodiphenylamine in methanol [242], acid catalysis occurs but no chloride ion catalysis, in contrast to the situation for the Orton rearrangement of N-chloroanilides. As a result of these observations it was proposed in 1972 that rearrangement occurred intramolecularly in parallel with the reversible denitrosation reaction, Eq. (140-1), [243]. This outline mechanism explains RNNO

RN(H)NO + H30^ ^ T ^

I

RNH ^ —

I

H

+ YNO

(140)

f

Decomposition RN(H)NO

RNH intramolecular

|

r

|1

NO

(141)

The Reactions of N-nitrosamines and Related Compounds

61

equally well all of the early observations. The yield of rearrangement product then depends on the relative rates for the intramolecular pathvây and the decomposition of YNO. The denitrosation pathway can be examined in isolation by the addition of a nitrite trap (sulfamic acid etc.) in sufficient quantity to make denitrosation effectively irreversible and by the use of a sufficiently high concentration of a good nucleophile Y~ to make the rate of rearrangement effectively neglible. Similarly, the rearrangement reaction can be made dominant by the addition of the secondary amine PhN(H)R which ensures that the denitrosation pathway is fully reversed. Both of these situations have been achieved experimentally for reactions in aqueous solution using sulfuric acid or hydrochloric acid catalysts. The correct mechanism was established beyond doubt by experiments involving the progressive addition of nitrite traps (which remove YNO). With the intermolecular mechanism, addition of e.g. sulfamic acid should progressively decrease the rearrangement product yield eventually to zero. All nitrite traps should behave similarly. However the prediction from the intramolecular pathway mechanism is that the yield of rearrangement product should again decrease with increasing [nitrite trap], but this time to a constant limit, governed by the relative rates of the denitrosation and rearrangement pathways, Eq. (142). The rate of denitrosation can be varied by changing the concentration and nature (thiourea > thiocyanate > Br~ > Ct) of the nucleophile, whereas the rate of + PhN(R)HNO

Y"

•

Denitrosation

(142)

Rearrangement rearrangement for any one reactant is fixed only by the acidity of the medium, which affects the rate of denitrosation in the same way. The results obtained for reactions in sulfuric acid are shown in Tables 8, 9 and 10. It is quite clear from Table 8 that the rearrangement yield is constant at around 21% and the rate constant is also constant within the experimental error, over a very wide range of concentration for four nitrite traps of quite different structure [244]. Similarly for the 3-methoxy compound (Table 9), which gives a much higher % rearrangement, the yield is again constant at around 85% over again a wide range of trap concentrations [245]. These results are completely at odds with the intermolecular mechanism, Eq. (137-8). The effect of the nature and concentration of the nitrite traps is shown in Table 10. As the nucleophilicity of the nucleophile is increased the % rearrangement product

62


Table 8 Rearrangement yield and overall rate constant for the reaction of N-methyl-N-nitrosoaniline in sulfuric acid (2.75 M) in the presence of various nitrite traps. Nitrite trap HN3 (6.5 X 1 0 ^ M) HN3(16.3x 10-4 M) NH2S03H(3.1 X 10-3 M) NH2SO3H (7.8 X 10-3 M) CO(NH2)2(0.10M) N H 2 0 H ( 1 . 6 x 10-3 M) NH7QH(2.6x 10-3M)

% Rearrangement 21 21 21 22 21 20 20

104)t/s0.65 0.67 0.65 0.64 0.62 0.66 0.62

Table 9 Rearrangement yield and overall rate constant for reaction of 3-methoxy-N-methyl-Nnitrosoaniline in sulfuric acid (3.5 M) in the presence of various nitrite traps Nitrite trap None HN3 (1 X 10-3 M) HN3 ( 5 x 1 0 - 3 M) N H 2 N H 2 ( l x 10-3 M) NH2NH2(5X10-3M)

NH2S03H(1 X 10-3 M) NH7SO3H(5X10-3M)

% Rearrangement 95 84 85 85 85 84 84

103 ^/s 2.94 3.10 3.22 3.25 3.39 3.33 3.48

Table 10 Rearrangement yield and overall rate constant for the reaction of 3-methoxy-N-methyl-Nnitrosoaniline in sulfuric acid (3.5 M) containing hydrazoic acid (5 x 10-3 M) and added nucleophiles Nucleophile (Water) CI-(0.077 M) Br-(0.077 M) SCN- (0.077 M) (NH2)2CS (0.079 M) Bi-(0.004 Bi-(0.008 Br- (0.016 Br- (0.032 Br-(0.100

M) M) M) M) M)

% Rearrangement 80 73 16 0 0

103^s~^ 3.27 3.51

65 56 36 29 11

3.86 4.14 6.6 9.4 -


63

decreases, to - 0 % for both SCN~ and thiourea at these concentrations. The Table also shows the effect of increasing the concentration of one nucleophile, in this case bromide ion, where again the rearrangement yield is progressively decreased. The same effect has been found for reactions of the 3-methoxy derivative in ethanol containing hydrogen chloride using thiourea as the nucleophile. Under these circumstances, when denitrosation is effectively irreversible, the measured rate constant k^ will be the sum of the rate constants for rearrangement {k^ and denitrosation (A:j)), Eq. (143), thus predicting a

linear dependence of k^ upon [nucleophile] with a positive slope and positive intercept. Again this is borne out experimentally. Addition of N-methylaniline to the reaction of N-nitroso-Nmethylaniline in sulfuric acid increases the rearrangement yield from --20% to >80%. Under these conditions, the rate constant is unaffected by the addition of significant amounts of C r [246]. This is shown quite dramatically in Table 11 where the rate constants for the reaction of 7V-methyl-N-nitrosoaniline, (a) leading to rearrangement only, and (b) leading to denitrosation only, are compared when halide ion is added. In case (a), achieved by the addition of excess N-methylaniline, the rate constant is unchanged by the addition of both c r and Br~, whereas for (b), achieved by addition of excess sulfamic acid, there is marked catalysis by both halides, with Br~ > CX" as expected. Both rearrangement and denitrosation are subject to acid-catalysis with a solvent isotope effect ^H90^'^D90 ^^ ^^' ^'^' consistent with a rapid pre-equilibrium protonation i.e. specific hydrogen ion catalysis. Protonation at the nitrogen atom makes the most sense mechanistically, and there is evidence, by observation of changes in the UV spectrum at very high acidities which are consistent with a substantial degree of protonation. The ^K^ of A^-methyl-Nnitrosoaniline has been estimated, from these measurements to be — 2 [247]. The rearrangement is an example of an electrophilic substitution since 3methyl and 3-methoxy substituents increase the rate of reaction, whereas 3chloro and 3-nitro substituents much reduce the rate of reaction [246,248]. Further there is a kinetic isotope effect of k^^k^ -2.4 when the 4-ring hydrogen atom is replaced by deuterium [246]. This identifies a Wheland intermediate from which proton transfer to the solvent is in part rate-limiting, which is not often encountered in electrophilic aromatic substitution. These results identify the two intermediates given in Eq. (144-5). A 2-methyl substituent activates slightly, probably by affecting the basicity of the nitrosamine, whereas the 2,6-dimethyl derivative is very stable both to rearrangement and denitrosation, no doubt due to steric inhibition to solvation in the protonated nitrosamine [249].

64


Table 11 Rate constants for (a) the rearrangement and (b) the denitrosation of N-methyl-Nnitrosoaniline in the presence of added CI" and Br~ Hahde added

10%s-l (a) 1.75 1.79 1.77

None Sodium chloride (0.24 M) Sodium bromide (0.10 M)

RNNO

(b) 16.3 41.7 75.1

RNHNO

i

+H2O

(144)

(145) NO It is not possible from the kinetic data to deduce the actual nature of the rearrangement, so it remains a matter of conjecture. Various suggestions have been made for other intramolecular rearrangements of N-substituted anilines such as the formation of a 7C-complex intermediate, a polar cyclic transition state or a caged radical ion mechanism. An attractive possibility is a suggestion, first made to account for the rearrangement of sulfanilic acid [250] and later extended to other N-substituted anilines [251], which might provide the reaction pathway for the Fischer-Hepp rearrangement. This involves the formation of a bent-boat form intermediate which is ring protonated and which might allow the nitrogen atom of the nitroso group to approach the 4-ring position sufficiently to allow bonding to take place directly, Eq. (146). + RNHNO (146)


65

Whatever the nature of the intermediates and transition states, it is quite clear that the rearrangement takes place intramolecularly. A fundamental error made in early years makes the assumption that if an intermediate can be trapped out, then it is an intermediate on the pathway to the product. This is not, of course, necessarily the case as we have shown here where reversible denitrosation to give the secondary amine and a nitrosating agent occurs, in parallel with the rearrangement. It is very likely that the reaction of aromatic secondary amines with nitrosating agents will first of all attack the amino nitrogen atom, and not the 4-ring position, and this will be followed by the intramolecular rearrangement. There is spectral and kinetic evidence to back up this suggestion [252]. 3.2.

Denitrosation of nitrosamines R(R')NNO The formation of nitrosamines from secondary amines and a nitrosating agent XNO is reversible, with the equilibrium position lying generally well over towards nitrosamine formation. The reverse reaction can be brought about if, starting with the nitrosamine in acid solution, the nitrosating agent generated can be removed with a nitrous acid trap, at a rate which is much faster than the reverse reaction, the re-nitrosation of the secondary amine generated, Eq. (147). The solvent (in this case water) -promoted reaction is

RR'NNO + H2O RR'NNO + X"

• •

RR'NH + HNO2 ( — • removed) RR'NH + XNO ( — • removed)

(147) (148)

often quite slow, and the process can be speeded up by the addition of nucleophilic species X~, Eq. (148). A procedure which has proved useful synthetically is to reflux the nitrosamine in alcoholic sulfuric acid in the presence of urea [253]. The range of nucleophiles which have been examined is quite large and their relative reactivities have been established [128,129]. These are shown in Table 12. There is a reasonable correlation between reactivity and the Pearson nucleophilicity parameters n for two different nitrosamines, (a) N-methyl-N-nitrosoaniline and (b) N-nitrosodiphenylamine, as shown in Figure 4, but the range for which n values are known is rather small ( C r , Br~, I~, SCN", thiourea), so the correlation may not be indicative of the detailed mechanism. The reactions are all acid-catalysed and follow quantitatively the Hammett acidity function h^ quite well for the relatively high acid concentrations required to effect reaction.

66


Table 12 Relative reactivity of a range of nucleophiles towards N-methyl-N-nitrosoaniline in acid solution. Nucleophile Chloride ion Cysteine Glutathione S-Methylcysteine Bromide ion Methionine Thiocyanate ion Thiourea Methylthiourea Iodide ion

Relative reactivity (Cl~ = 1) 1 2 3 35 55 65 5500 13000 14250 15750

The most basic site in nitrosamines is likely to be the oxygen atom, just as it is in amides, as this allows delocalisation of the positive charge to the amino nitrogen atom, Eq. (149). Indeed there is physical evidence for the R2N-N

^

•

R2N=N

OH

(149) \ H

existence of this species from NMR studies of solutions of nitrosamines in strong acids [254]. Calculations also indicate that the 0-protonated form is more stable than is the N-protonated species [255]. However it is not easy to write a convincing mechanism via the 0-protonated species, and most authors write the protonation on the N atom. This allows the ready expulsion of a

R2NH-N

^

X-

•

R2NH + XNO

(150)

O good leaving group R2NH, Eq. (150). More recently, Keefer and co-workers [256] have produced evidence that N-protonation of dimethylnitrosamine occurs to a measurable extent. This was based on NMR measurements of the Z - ^ ^ E equilibrium following selective deuteriation of one of the methyl groups (syn to the oxgen atom of the nitroso group). An estimate for the pX^ for the N-protonated speices is —12 to -13. In the absence of added nucleophiles, denitrosation is quite slow unless high acid concentrations are employed. The reaction carried out in the presence of say bromide ion and azide ion is a procedure which has been used


n Fig. 4. The reactivity of nucleophiles in the denitrosation of nitrosamines in terms of the Pearson nucleophiUcity parameter, n, (a) for N-methyl-N-nitrosoaniline and (b) N-nitrosodiphenylamine.

67

68


to remove small quantities of nitrosamines present in a range of consumer products. The reaction, in a modified form in the Thermal Energy Analyser (TEA) is used to determine amounts of nitrosamines quantitatively. A typical procedure is to treat the sample with 5-10% solution of hydrogen bromide in acetic acid, which at high temperature (reflux) generates nitric oxide which is analysed using a chemiluminescence detector after reaction with ozone, generating nitrogen dioxide in an excited state, which produces the chemiluminescence [257]. Other solvents have been used successfiiUy as has a procedure involving heating the sample to ~300°C to generate nitric oxide directly. In order to separate the nitrosamine components, GC-TEA and HPLC-TEA modifications have been introduced to commercial instruments. Other methods for the destruction of nitrosamines include UV irradiation, [258], reaction with transition metal complexes, reaction with CuCl/HCl, oxidation with potassium ferrate K2Fe04 (a powerful oxidant), reaction with carboxylic acid halides, reduction with an Al-Ni alloy with aqueous alkali, use of exchange resins and incorporation into zeolites sometimes containing metal catalysts (e.g. copper). Much of the detail is within the extensive patent literature. At high nucleophile concentrations, for the more reactive species, catalysis is lost. This occurs for the reaction of N-methyl-N-nitrosoaniline in aqueous acid in the presence of thiocyanate ion or thiourea. Plots of the observed rate constant {k^ against [nucleophile] are initially linear, but downward curvature sets in at about 0.2 M [259], and both curves level off at the same limiting value. This will occur, Eq. (151), when A:2[X~] » k_^ and the rate limiting step will be the protonation of the nitrosamine. The limiting

RR'NNO

HsO^, kx " k-i

H30^ \+ RR'NN(OH)"^

+ XT, k2 RR'NHNO (^ ^x \ k_2 )

RR'NH + XNO Trap

Destroyed (151)

value of k^ should be independent of the nature and concentration of X as is found. Bromide ion is not sufficiently reactive for this to be observed, even at concentrations of -- 0.8 M. This scheme is supported by the observation of a kinetic isotope effect ^HOO^^DOO ^^ ^^^ range 1.3-3.8 when ^2[^~] ^^ ^-i ^^^


69

0.3 when ^2[^~] ^^ ^-i* This pattern of behaviour had been previously noted for the denitrosation of nitrosamides, which will be discussed in detail in the next section. Similarly there is no nucleophilic catalysis for the denitrosation of N-methyl-N-nitrosoaniline in ethanol solvent. All of these results are consistent with those obtained for the reverse reaction i.e. N-nitrosation, discussed in Chapter 2, where again for amines, nucleophilic catalysis can disappear at high [nucleophile] for the most reactive nucleophiles and is totally absent for the nitrosation of amides. There are many examples of the so-called trans-nitrosation reaction i.e. where one nitrosamine transfers the NO group to another amine. This is generally a reversible reaction, Eq. (152), and the final equilibrium position R2NNO + R'2NH -j—^

R2NH + R'2NN0

(152)

will be governed by the concentrations of the reactants and the nature of R and R. A few examples are given in Eq. (153-5). A direct reaction can be ruled .°

,0CH3 OH

.OCH3

p ^ O C H 3 k^CH=NOH

OH (174)

An important industrial process, the SNIA Viscose process, is based upon the nitrosation of cyclic ketones in which C-C bond fission occurs - a common event for systems where the nitroso group cannot form an oxime by tautomerism. The sequence of events is laid out in Eq. (175), for the reaction

?< ^ AT

NO+HSO4"

^Y^^"

NOH H^

(

mi

o

(175)

Aliphatic and Alicyclic C-nitrosation

81

of an aryl cyclohexyl ketone. A similar reaction occurs for other substituted cyclohexanes e.g. cyclohexane carboxylic acid. The reaction is usually carried out in 85-90% sulfuric acid and the oxime formed undergoes a Beckmann rearrangement to give a seven membered cyclic amide, 8-caprolactam which undergoes ring opening polymerisation to give nylon 6. There is some kinetic evidence that the nitrosation goes via the enol form of the ketone under these conditions [281] as outlined in Eq. (176).

(yr = o=' ,COAr

/—\

Âr

"ÔH

(176)

N0+ Ar

+ ArC02H Later a more comprehensive kinetic study [16] has shown that enolisation is a common feature for ketone nitrosation, just as it is for other electrophilic reactions of ketones e.g. halogenation. For a number of ketones, nitrosation in water by nitrous acid in the presence of added nucleophiles X~ ( C r , Br~, SCN~, SC(NH2)2) is first-order in ketone, is acid catalysed and independent of the concentration of nitrous acid and of the concentration and nature of X~. These enolisation rate constants for a number of ketones were shown to be the same, within experimental error, as those obtained from halogenation and other experiments. Further, by adjusting the concentrations of both X~ and HNO2 it is possible to change the rate-limiting step from that of enolisation to the step where the nitrosating species attacks the enol - again as was shown to be the case for halogenation at low [halogen]. Table 13 shows the second-order rate constants {k^ derived for acid catalysed enolisation, defined by Eq. (177), for 2-propanone, 2-butanone and 1,3-dichloroacetone, together with the average values in the literature derived Rate = ^JHN02][H+]

(177)

fi-om halogenation and hydrogen-isotope exchange reactions. The agreement is very good. For the first two ketones the values are for the acid-catalysed

82


Table 13 Values of the rate constant for enolisation Ketone 2-Propanone

RQ from nitrosation 3.8 X 10-5 M-1 s-l

KQfromhalogenation etc.

2-Butanone

4.9 X 10-5 M-1 s-l

4.8 X 10-5 M-l s-l

1,3-Dichloroacetone

3.2 X 10-^s-l

3.2x10-6 s-l

2.8 X 10-5 M-l s-l

Table 14 Values of ^ in, rate = /:[XNO][Enol]

Enol

^/M-1 s-l BrNO 7.4x108

QNSCN

CH3C(OH)=CH2

CINO 1.5x109

CH3C(OH)=C(H)CH3

4.6 xlO^

3.8x109

3.0x108

CH3C(OH)=C(H)COCH3

1.0x105

1.4x104

500

C1CH2C(0H)=C(H)C1

1.2x104

2.8x103

(NH2)2C SNO

38

reaction, whereas for 1,3-dichloroacetone the reaction is not acid-catalysed and so the rate constant is a first-order rate constant. Generally when the equilibrium constant for enol formation K^ is very small, the rate-limiting step is usually that of enolisation, whereas when K^ is quite large then reaction of the enol form is rate-limiting. However, it is possible (just as in the case of halogenation of e.g. 2-propanone), to adjust the experimental conditions such that nitrosation is rate-limiting, even for systems with very small K^ values. Some results of such studies are presented in Table 14. Two points emerge:(a) the introduction of electron withdrawing groups, -COCH3 and -CH2CI and CI, deactivate the enol to electrophilic attack as predicted, and, + (b) the reactivity sequence CINO > BrNO > ONSCN > (NH2)2CSNO which occurs in N-nitrosation is also the pattern in aliphatic C-nitrosation. For 2-propanone and 2-butanone it is possible to identify a pathway via N2O3 in the absence of added nucleophiles. The rate constants for N2O3 addition to the enols are both close to the diffusion controlled limit. The values of rate constants for reaction of the enol forms are crucially dependent on the values of the equilibrium constant for enolisation K^, This may not be known accurately when K^ values are very small, such as for 2propanone and 2-butanone.

Aliphatic andAlicyclic C-nitrosation

83

In a number of cases the enol is sufficiently acidic for there to be present enough of the enolate anion for reaction to occur via this form of the reactant. This has been detected, by a detailed analysis of the variation of the observed rate constant with acidity for the reactions of some cyclic P-diketones:Meldrum's acid [18], dimedone [282] and indane 1,3-dione [24], and also for trifluoro-substituted derivatives of 2,4-pentanedione [283]. In the latter case 2,4-pentanedione itself reacts exclusively via the enol form, whereas 1,1,1trifluoro-2,4-pentanedione reacts via both the enol and enolate forms and l,l,l,5,5,5-hexafluoro-2,4-pentanedione reacts exclusively via the enolate form. These substitutions show clearly the acid-strengthening effects of the -CF3 groups and also their deactivating effects to electrophilic addition in the enol/enolate. As expected the enolates are much more reactive than are the corresponding enols, and many react at the diffusion limit. The nitrosation of ethylcyclohexanone-2-carboxylate, which exists in solution almost exclusively in its enol form, Eq. (178), is inhibited by the

^

ii

^ V

addition of p-cyclodextrin, as a consequence of the formation of unreactive complexes between the enol and P-cyclodextrin [284]. Anionic micelles of hydrogen- and sodium dodecyl sulfate generate substantial catalysis of the nitrosation of benzoylacetone, under conditions where the nitrosation step is rate-limiting. The results are treated by the pseudo phase models and are explained in terms of reaction in both the micellar and aqueous phases and that catalysis arises from an increase in the concentrations of both the nitrosating species and the enol in the micellar phase [285]. Interestingly a method has been developed for the determination of K^ values for P-dicarbonyl compounds, based on the addition of surfactants to the reaction medium of their nitrosation reactions [286]. The method relies on the ability of the micelles to provide a medium (or phase) within the bulk solvent, and on the fact that K^ values are extremely sensitive to the nature of the solvent. Values have been reported for benzoylacetone, acetylacetone, ethyl acetoacetate and ethyl benzoylacetate, many of which (but not all) agree with published literature values. 4.2

Nitrosation of other carbonyl compounds Any reactant containing the CH3CO- or -CH2CO- groupings will in principle undergo nitrosation yielding oxime, or nitroso (monomers or dimers)

84


products. Some well-known examples are shown in Eq. (179-180), and others can be found in the Touster compilation [278]. As for simple ketone reactions, CH3COCH2CO2C2H5

^^^*

CH3COCCO2C2H5

(179)

NOH C6H5CH2CO2C2H5

"^^^>

C6H5CCO2C2H5

(180)

NOH C-C bond fission can occur in some cases. Two examples are shown in Eq. (181-2). The reaction of cyclohexane carboxylic acid is an alternative RC0CH(R')C02R"

^^^*

R'CC02R"

(181)

NOH RCH(C02H)2

XNO

„^^^ „ *- RCCO2H

(182)

NOH

a

C02H

/v^NOH

/—\

procedure for the industrial production of 8-caprolactam, Eq. (183). This process is activated by UV light which suggests a different reaction pathway involving radical intermediates and possibly the ketene C ^ H | Q = C = 0 , but this is not known with any certainty. Kinetic studies with ethyl cyanoacetate and diethyl malonate show that the results are consistent with reaction via the enol and enolate forms of both NCCH2C02Et ^ ; = ^

NCCH=CX^

:^:^

XNO X

NCCF^CQ XXÔ

NCCC02Et NOH

^^^^>*

Aliphatic and Alicyclic C-nitrosation

85

reactants [287]. The products are the expected oximes, Eq. (184). The ester behaves in the same way as do simple ketones towards nitrosation and also towards halogenation. Malonic acid and its 2-substituted derivatives also undergo nitrosation via their enol forms. Enol tautomers of carboxylic acids and esters are difficult to detect and to isolate, since K^ values are much smaller than they are generally for simple ketones. Some examples exist, notably when there are large bulky groups present which destabilise the keto form by inhibition of solvation of the carbonyl group [288]. It is possible, by changing various concentration terms, to make either enolisation of the malonic acids or the further reaction of the enols rate-limiting. The enolisation rate constants agree reasonably well with the literature values, where they exist. The existence of a second term in the rate equation which is second order in the malonic acid (MA), Eq. (185), is readily explained by another pathway Rate -A:[MA] + )t*[MA]2

(185)

to enolisation which is base catalysed, brought about by the malonate anion [289], which is the dominant pathway at pH > 3.2. Interestingly, enolisation of malonic acid is not catalysed by mineral acid, but rather by an intramolecular acid catalysed pathway involving proton transfer from one of the carboxylic acid groups to the carbonyl oxygen atom of the other acid group. This could precede proton transfer from the methylene group to the solvent, or take place concurrently with it, Eq. (186).

"

^-

R—^

•

.CHO R—< 0S020Et

(202) ^ '

[302], Eq. (201-2) respectively, leads to the formation of substituted ketones or aldehydes via preliminary nitrosation, isomerisation to the oxime and hydrolysis of the oxime to give the carbonyl derivative. 7i-Complexes of NO"*" with alkenes have been postulated as reaction intermediates, but now one has been synthesised and characterised, as the tetrachloroaluminate, by the reaction of nitrosonium tetrachloroaluminate with

+ N0ÂlCl4-

* Vt. •

II

-li—NO^ AICI4(203)

Aliphatic andAlicyclic C-nitrosation

91

l,2,3,3,4,5,6,6-octamethyl-l,4-cyclohexadiene, Eq. (203), in SO2/CD2CI2 solvent at-90°C [303]. Nitrogen dioxide (NO2/N2O4) generally adds to alkenes yielding dinitro compounds or nitro nitroso compounds, but in the absence of weakly basic solvents and at low temperature, reaction can occur by an ionic mechanism via the N0"^03"~ f^^^*"^- Thus 2-methylpropene gives the nitroso nitrate in its CH2=C(CH3)2

NO^NOs" - ^

[CH2(NO)C(ON02)(CH3)2]2

(204)

dimeric form, Eq. (204), which can be converted to the oxime on warming [103]. The reaction conditions are, liquid ethane-propane solvent in the temperature range -196 to -78°C. Nitric oxide in its pure form will not react with alkenes, but if traces of oxygen or nitrogen dioxide are present then nitroso nitro alkanes are readily formed. Nitrogen dioxide will be present in small concentrations if oxygen is not rigorously excluded from the system, but can also be present, particularly if nitric oxide is at high pressure, by the disproportionation reaction, Eq. (205). In one study [304], reaction was initiated by addition of a small quantity of 3N0 = N2O + NO2

(205)

nitrogen dioxide which produced P-nitroalkyl radicals, detected by EPR, which were then captured by nitric oxide to give the nitroso nitroalkane. Radicals were detected when air was introduced into the nitric oxide stream. Earlier [305], the product from cyclohexene and nitric oxide under seven atmospheres pressure was shown by X-ray analysis to be the trans-dimer, Eq. (206).

'NO' • •

•

•

(206)

There are thus two explanations for the reactions of nitric oxide in the presence of even traces of oxygen, (a) the radical explanation given above involving attack of NO2 and (b) a heterolytic mechanism given in Chapter 1, section 1.6.1., when considering nitric oxide as a potential nitrosating species (again in the presence of oxygen), where there are many examples given for the nitrosation of amines, amides, thiols and alcohols, where the products have been rationalised in terms of N2O3 formation, which then effects conventional electrophilic nitrosation. In these cases the products are usually the nitroso

92


products, which cannot stem from initial attack by nitrogen dioxide. Maybe both mechanisms operate under different sets of experimental conditions. Dinitrogen trioxide reacts readily with alkenes, generating nitrosonitroalkanes. For example, the reaction of 2-methylpropene [306] gives the trans dimer of l-nitro-2-methyl-2-nitrosopropane, Eq. (207). The orientation CH2=C(CH3)2 + N2O3

*- [CH2(N02)C(NO)(CH3)2]2

(207)

of addition here suggest a radical reaction, perhaps involving initial attack by nitrogen dioxide. Typical reaction conditions are, (a) a solution of dinitrogen trioxide in ether at low temperature, or (b) gas mixtures of nitric oxide and nitrogen dioxide. Reactions in water where dinitrogen trioxide is formed in situ from nitrous acid, are clearly different, in that electrophilic nitrosation has been established on many occasions including for the reaction with alkenes. Examples of reactions where nitric oxide reacts with a free radical (usually a carbon-centred radical) generating nitroso compounds have been discussed in Chapter 1, section 1.6.1. There is a recent review [307] of the reactions of alkenes with nitrogen oxides and other nitrosating (and nitrating) agents.

Nitrosation Reactions and the Chemistry of Nitric Oxide D.L.H. WiUiams © 2004 Elsevier B.V. All rights reserved.

93

Chapter 5

Aromatic C-nitrosation Nitration of aromatic systems is one of the classical reactions in organic chemistry and has been widely studied. The early work of the Ingold school established the mechanism for reaction via the nitronium ion N02'^, and subsequently, radical pathways, particularly involving nitrogen dioxide have been discovered. More recent developments have included the consequences of ipso nitration and the search for direct spectroscopic evidence for proposed intermediates. The topic has been well covered in the literature and the monograph by Schofield [308] has covered the important aspects of the area comprehensively. In addition, all textbooks on organic chemistry cover particularly the pathway involving electrophilic aromatic substitution. By contrast, aromatic nitrosation reactions have not received much attention, although in recent years there has been a resurgence of interest. Substitution of a hydrogen atom in the aromatic nucleus by the NO group has been mostly limited to reactants which contain -OH, -OR, -NH2 and -NHR groups i.e. groups which activate electrophilic aromatic substitution in the ortho and para position by their mesomeric effect. Aromatic nitrosation can in principle be accomplished using any of the reagents discussed in Chapter 1. 5.1.

Products of the reactions The most familiar reactions are those of phenols and naphthols. Both systems undergo ready nitrosation by nitrous acid in dilute aqueous mineral acid to give the products outlined in Eq. (208-9). Phenol itself gives mainly

H

, NO

J

900/0

N^^^

+ HNO2 + HsO^

(208) NO

^^k^N-OH

10%

94


(209)

+ HNO2 + HsO"^

OH

4-nitrosophenol (90%) and a little 2-nitrosophenol (10%). In solution all the C-nitroso aromatic products exist overwhelmingly in their tautomeric forms, the benzo (or naphtho) quinone monooximes, which accounts for the lack of blue colour in the final reaction solutions. The nitrosation of phenol itself has been the first stage of some syntheses of the analgesic paracetamol. Reduction of the 4-nitrosophenol to give the amine, followed by acetylation with acetic anhydride leads in a three stage synthesis to paracetamol. Phenyl ethers e.g. anisole also undergo ready nitrosation which has often been accompanied by dealkylation, so that the major isolated product is 4nitrosophenol, Eq. (210). The mechanism of the dealkylation reaction has not

(210)

+ HNO2 + H NO

OH

been established. Much later, 4-nitrosoanisole has been isolated (in 56% yield) using sodium nitrite in dichloromethane/trifluoroacetic acid solvent [309], and also quantitatively from nitrosonium tetrafluoroborate in acetonitrile under argon [310]. The latter procedure has been used successfully to obtain 4nitrosoanisole derivatives fi-om 2-methyl, 3-methyl, 2,6-dimethyl, 3,5dimethyl, 2-bromo, 3-bromo and 2-mesityl substituted anisoles [311]. A novel procedure has been developed for the nitrosation of anisole using acetic acid/sulfuric acid mixtures or in trifluoroacetic acid under a stream of nitric oxide containing injected oxygen [312-3]. In this way anisole and mxylene were nitrosated in good yield and toluene and o-xylene (for the first

Aromatic C-nitrosation

95

time) in modest yield. It is believed that the nitric oxide/oxygen mixture generates N2O3 in solution, maintaining it in fixed concentration. This has the effect of reducing the problems arising from the oxidation of both nitrous acid and the aromatic substrate, which previously resulted in low yields of the expected 4-nitroso products. This lias a major benefit in allowing a greater range of substituted arenes to undergo nitrosation in a reasonable yield. In addition, aromatic nitrosation is much more regioselective than is aromatic nitration, and given the ease with which these nitroso compounds undergo oxidation to the nitro compound, this translates into a method where a more regiospecific range of nitro compounds can be synthesised. It is reported that the reaction of nitrosonium ethyl sulfate reacts with alkyl aryl ethers (and also dialkyl aniline derivatives) to give the 4-nitroso arenes without any of the usual side products (dealkylated and nitro products) [314]. Reactions can also be brought about by alkyl nitrites under mildly basic conditions. For example m-cresol reacts with i-amyl nitrite in DMF containing

RONO CH3

K2C03/DMF^

cu.-\c^

^^^^^ OH

potassium carbonate to give the benzoquinone monooxime derivative, Eq. (211), in 71% yield [315]. The generality of this procedure was shown by the nitrosation of 3-methoxyphenol and 1-naphthol. In both cases the para quinone monooximes were isolated, whereas the reaction under acidic conditions (sodium nitrite in aqueous acetic acid) gave the ortho products [316]. Other alkyl nitrites and N-methyl-N-nitroso-p-toluenesulfonamide also generate aromatic C-nitroso products from phenol derivatives 317-8]. Phenol itself undergoes a rapid reaction with a nitrosothiol, S-nitroso-Nacetylpenicillamine in the pH range 2-9, Eq. (212). A radical mechanism was

+ RSNO

•

[1^ J 'OH

(212)

96


originally proposed [319], but the results are equally compatible with the formation of nitric oxide from RSNO in a copper catalysed reaction (see Chapter 8), subsequent oxidation and nitrosation via N2O3. Reaction is halted when EDTA is added to remove the copper catalyst, so it appears that the reaction via NO formation and oxidation of NO is the more likely [320]. Pires and co-workers have shown that the NO/O2 system generates an efficient nitrosating species which will convert phenol to 4-nitrosophenol [321]. Secondary aromatic amines are readily nitrosated at the nitrogen atom and the nitrosamine formed can rearrange intramolecularly in acid solution to give the 4-nitroso isomer, Eq. (213), by the Fischer-Hepp rearrangement discussed fully in Chapter 3, section 3.1. CH3l)JH

CH3NNO

CH3NH

NO Tertiary aromatic amines also readily form the 4-nitroso product. Dealkylation can occur here, giving the N-nitroso-4-nitro derivative as one of the isolated products. A likely series of steps involves C-nitrosation, oxidation of the nitroso group, dealkylation and further N-nitrosation as outlined in Eq. (214) [322]. There is no reason to discount the possibility that C-nitrosation occurs via N-nitrosation and Fischer-Hepp rearrangement - quite a likely scenario given the reactions of secondary amines. NR2 HN02^ H30^

r l

I}JR2

TjJR(NO)

HN02^ NO2 (214)

Nitrosation of phenol can also be accomplished photochemically using nitrite ion [323], as shown in Eq. (215). There have been no mechanistic studies, but it is possible that nitroso phenoxy radicals are involved as reactive


NO2"

97

hv

H ^

(215)

KJ OH

intermediates. In the same way, but using nitrosamines as the source of nitroso radicals, naphthols and anthrols are also converted via the nitroso product to the oximes [324]. There have been a number of cases reported where the nitroso group displaces a group (other than the hydrogen atom) from an aromatic system. The most well-known reaction is that of the nitrosative decarboxylation of benzoic acid derivatives as in Eq. (216). The reaction is essentially quantitative for the 3,5-dibromo-4-hydroxy derivative shown, and is also rapid, with immediate evolution of carbon dioxide, for reaction in water at room NO HNO2 »

+ CO2

(216)

H30^ temperature [325]. Other groups can also be displaced, e.g. -CH(OH)Ph, -CH2C^H4NMe2, -N2Ph, -COCH3 and -CHO, from the 4-position in 4substituted N,N-dimethylanilines. In these cases however the final product is the 4-nitro derivative and not the 4-nitroso, presumably formed by oxidation of the first formed nitroso product. A number of organometallic compounds also react in a similar way, by displacement of the organometallic group. This group includes the reactions of aryl mercury halides, aryl Grignard reagents, and tin, silicon and thallium compounds [326]. 5.2.

Reaction mechanisms Mechanistic studies have concentrated on the reactions of phenols naphthols and aryl ethers. Very little work has been undertaken with the dialkylanilines. A series of papers by Challis and co-workers in the 1970s set up most of what is known regarding aromatic C-nitrosation. The kinetic and other measurements are all consistent with the outline mechanism given in Eq. (217). Here we have reversible addition of NO"^ from a carrier of the

98


h B H

+ BH^

NO

nitrosonium ion, XNO (where X can be a range of groups, e.g. H2O-, C1-, SCN- etc.) to form the conventional Wheland intermediate, which can be stabilised by proton loss to give the dienone intermediate which in turn loses a proton to a base B (which can be the solvent) to give the product. This leads to the rate equation, Eq. (218). Most of the features predicted by this equation ^^^^ ^ /^l/^2[C6H50H][HN02][H^][X-][B]

^îg)

have been borne out experimentally :(1) there is no acid catalysis in the pH range 1.0-4.5 when ^_i[H'^][X~] » /:2[B]. It is more difficult to achieve fully the other limiting condition of ^_|[H^][X~] « A:2[B], but it has been partially achieved and acid catalysis appears [327-8], (2) similarly when ^_j[H"^][X~] » A:2[B] there is no nucleophilic catalysis by e.g. halide ion and thiocyanate. Again nucleophilic catalysis appears, for the 2naphthol reaction at low [H^] and low [X"], [329], (3) again when A:_j[H"^][X~] » A:2[B], the final proton loss to B is rate-limiting and this results in a kinetic isotope effect (typically %/A:j) --3.5) when the 4position in phenol is substituted by deuterium [327,330]. Base catalysis is also found under these conditions, which follows the Bronsted equation giving a P value for phenol itself of 0.37, [328], and,


99

(4) when nucleophilic catalysis does occur, the experimentally obtained second order rate constants for XNO reactions follow the now familiar order CINO > BrNO > ONSCN, but the actual values (for reaction with 2-naphthol [329]) are well below those obtained for the reactions of aromatic amines such as aniline itself. Similarly the third-order rate constant (for when X = H20'^-) for the reaction of 2-naphthol is 4700 M~^ s~^ and is close to the diffusion limit suggesting that H2N02"^/NO"^ is the most reactive of the nitrosating agents. There is no kinetic evidence for a reaction pathway involving dinitrogen trioxide. The reaction mechanism is thus the familiar A-S£2 pathway followed by many other aromatic systems. One major difference for nitrosation is, that under the "usual reaction conditions" the rate-limiting step is the final proton loss and not attack of the electrophile. The dieneone intermediate has not been observed spectroscopically as it has been in the bromination reaction of phenols. A kinetic study of the nitrosation of five 4-substituted phenols [331] showed the same characteristics as did the reaction of phenol itself even though the products are now the 2-nitro products. This suggests that nitrosation occurs at the 2-position, which is then followed by rapid oxidation to the nitro compound. At much higher acidities > -0.5 M, acid catalysis is a general feature of phenol nitrosation. It has been suggested that there is now an additional acidcatalysed pathway for the decomposition of the dienone intermediate which competes with the spontaneous breakdown, and which probably involves protonation of the nitroso oxygen atom. The kinetics of the reactions of aryl ethers are also consistent with the A-S£2 mechanism, and show many of the features exhibited by the phenol reactions [327, 332-3], even though the final products are the tautomers of the 4-nitroso phenols, which are believed to be formed after the rate-limiting stage by hydrolysis of the 4-nitroso aryl ethers. The nitrosation of substituted indoles yields the tautomeric oxime forms of the 3-nitroso product, Eq. (219), [334], even though there is a secondary

\ X^^îjj R H

™^2_ H30+

.

II

^

(219)

amine function present in the molecules. The mechanism appears to be the same as that for phenols and aryl ethers, but for the more basic compounds

100


(with pK^ > -3.5) formation of the Wheland intermediate is rate-Hmiting. Pathways via H2N02'^/NO"^, N2O3 and CINO were identified and each occurred at, or close to, the diffusion limit. For the less basic indoles (pK^ < -6.0), the second stage i.e. proton loss to the solvent, is rate-limiting as is normally the case with phenols etc. If however the 3-position is substituted, then the final product is the N-nitroso compound the nitrosamine. Detailed kinetic analysis reveals that even in these cases the initial attack occurs by attack at the 3-position and after proton loss from the nitrogen atom, an internal rearrangement of the NO group occurs from carbon to nitrogen [3356]. Studies of reactions of alkyl nitrites with a range of substituted phenols [318] in basic aqueous media, have led to the conclusion that reaction occurs at the oxygen atom of the phenolate anion generating unstable aryl nitrites which undergo either an intramolecular NO group rearrangement to the 4-position leading to the quinone oximate anion product, or simultaneously, homolysis to give the phenoxy radical and nitric oxide, Eq. (220). The latter, in the presence of oxygen is known to generate an efficient nitrosating species, probably N2O3, which can itself effect C-nitrosation in competition with its hydrolysis to give nitrite ion. The evidence in favour of these pathways comes from:(a) the observation of substantial concentrations of nitrite ion in the products, and lower yields of the quinone oximate product, (b) the good correlation between the reactivity and the pK^ of the substituted phenols, and with the Hammett a~ values, and, (c) methyl group substitution in the aromatic ring does not result in the expected decrease in reactivity.

X-b

11 + RONO

•

X-4-

I

+ RO

X-4H

II

+ NO

NO (220)


101

There is a claim [337] that the Wheland intermediate proposed in aromatic C-nitrosation has been identified by transient absorption spectroscopy, whereas a theoretical paper [338] failed to find a stable Wheland intermediate but identified a strongly bonded Electron Donor-Acceptor n complex between benzene and toluene with NO"^. Calculations involving the formation and stability of pre-equilibrium charge transfer complexes between donor aromatic systems and Br2, NO"^ and N02"^ [339], have produced an explanation as to why N02'^ is much more reactive than NO"^ (estimated to be a factor of 10^^ [332]) in these reactions. High level MO calculations and semiquantitative application of Marcus-Hush theory have resulted in the proposal of a charge transfer mechanism for both nitration and nitrosation. This may account for the bright yellow and red colours seen when aromatics and NO"*" salts are mixed. In nitration, two intermediates are proposed but only one in nitrosation. Here, pre-equilibrium complex formation is followed by a slow transformation to the radical cation Ar*"^ and NO which combine to form the stable C-nitroso products, Eq. (221). The relative unreactivity of NO"^ is

ArH + NO"^ :^=^

[ArH-NO^]

slow < ^

.+ ArH + NO

(221)

t ArNO + H^ believed to arise from the slow nature of the radical cation formation. As yet there is no convincing experimental evidence for such a mechanism. Some polymethylbenzenes generate donor-acceptor complexes with NO"^ derived from alkyl nitrites in acidified solutions of dichloromethane [340]. These complexes, which are formed almost quantitatively from hexamethylbenzene etc. have characteristic spectral properties and were identified by a combination of UV-VIS and IR spectroscopy. Remarkably stable complexes are formed between nitrosoaromatics and N0+ [310], which can be identified by NMR, and in one case [341] a somewhat unstable 1:1 complex has been isolated from 4-nitrosoanisole and a nitrosonium salt, Eq. (222). The equilibrium formation constants have been determined by NMR measurements and are typically >40000 M~^ for 4nitrosoanisole in acetonitrile solvent [313]. It is believed that the stability of such complexes prevents further ring-substitution and accounts for the slow final proton loss.


102

J

+N0^PF6-

^

.NO^

+ PF6"

(222)

NO 5.3.

Nitrous acid-catalysed nitration reactions It has long been known that nitrous acid, present in low concentration in nitric acid solutions, or added as sodium nitrite, can have a major catalytic effect on the nitration reaction. Initially it was believed that nitrosation first occurred followed by oxidation to give the nitro product. It is now known that although this process can occur, more fi'equently an alternative pathway is taken [342]. Nitrous-acid catalysed nitrosation can be recognised by the change in reaction rate and also in the regioselectivity which occurs if steps are taken to remove the nitrous acid. This was achieved in the early experiments by the addition of urea, but more recently, more effective nitrous acid scavengers, such as hydrazine and sulfamic acid have been used successfully. Most of the reported work refers to the reactions of phenols, anisoles, aniline derivatives and reactive alkyl benzenes, but the reaction also occurs in the O-nitrosation of alcohols. The well-known reaction of ethanol itself at room temperature with concentrated nitric acid leads to a violent autocatalytic reaction which generates clouds of nitrogen dioxide. However if the reaction is carried out in the presence of nitrous acid scavengers, ethanol and concentrated nitric acid can safely be refluxed to generate ethyl nitrate. The product ratios in aromatic nitration of e.g. phenol are quite different for the nitrous acid catalysed reaction. The former yields a mixture of 2- and 4-nitro products whilst the latter gives almost exclusively the 4-isomer. The nitration of N,N-dimethylaniline by nitric acid in the presence of nitrous acid is also much faster than the C-nitrosation reaction when there is no nitric acid present. In addition to the expected nitration product, 4-nitro-N,Ndimethylaniline, significant amounts of N,N,N',N'-tetramethylbenzidine were formed, in yields which increased with acidities. The presence of the latter product very much suggests that a one-electron transfer process is involved and a mechanism was proposed assuming a one-electron transfer in the first step, Eq. (223), [343]. This yields a radical cation-radical pair with nitric oxide, which in nitric acid releases the radical cation which can either dimerise to give the benzidine product, or which can exchange with the nitronium ion giving another radical cation - radical pair, this time with nitrogen dioxide, from which the nitro product is obtained. The key to the mechanism is the exchange

Aromatic C-nitrosation HN(CH3)2

103

(CH3)2 + NO""

;+«H-NO

+ H"

(223)

HNO3 \ N 0 2 N(CH3)2

N(CH3)2 (CH3)2N-

rx-TX

•N(CH3)2 NO2

reaction, NO + N02'^ ^ ^ NO"^ + NO2 which may represent an oversimplification. The mechanism also fits the results obtained from the reactions of phenols and anisoles [344]. Kinetic evidence also supports this mechanism. In particular, if the assumption is made that attack by NO"^ or by NO2 can be rate limiting, then a consequence (derived from a detailed steady state treatment) is that the kinetic order with respect to HNO2 should change from 1 to 0 as the concentration of HNO2 is increased. This has been observed experimentally [345-6]. Strong evidence for the intermediacy of radical cations comes from the observation of CIDNP effects when reactions are carried out with ^^N-labelled nitric acid [347] for the reaction of dimethylaniline, and also for a number of other reactants including mesitylene [347] and phenols [348]. As expected the CIDNP signals are not present when nitrous acid is removed from solution using hydrazoic acid. The actual polarisation effect is believed to arise from the partitioning of the radical cation-radical pair ArH*"^02, between dissociation to separated

104


radicals, and combination to give the Wheland a intermediate and thence the nitro product. For the nitrous acid-catalysed nitration of naphthalene in trifluoroacetic acid or in a mixture of methanesulfonic and acetic acids, the kinetics reveal a term second-order in naphthalene, which is dominant at higher [naphthalene]. This is interpreted in terms of the formation and reaction of dimer radical cations [349]. Radical cations can be generated from nitrosonium salts such as N0"^BF4~ and pyrene and anthracene in dichloromethane solvent [350]. There is a comprehensive review of the whole range of radical processes which have been identified in nitration, which includes a section on the nitrous acidcatalysed nitration reactions [351].


105

Chapter 6

O-Nitrosation 6.1.

Nitrosation of alcohols

Alcohols react readily with acidified aqueous nitrous acid solutions to generate alkyl nitrites, Eq. (224). The reaction sets up an equilibrium which can be

ROH + HNO2

=C 5

0-nitrosation

115

the N-nitroso product. Hydroxylamines, particularly at high acidities are also believed to undergo O-nitrosation as the first step. We can of course regard the hydrolysis of a large number of nitrosating species, such as the nitrosyl halides, dinitrogen trioxide, alkyl nitrites etc., as examples of the O-nitrosation of water generating nitrous acid in acid solution, Eq. (245). In line with many reactions the familiar reactivity sequence CINO > BrNO > ONSCN is again borne out, see Table 6 p. 12.

CINO +

/

H

H +/ ONO

H

+ cr

(245)

H H2O

HNO2 + H3O+ In basic solution the corresponding reactions of alkyl nitrites are examples of O-nitrosation of the hydroxide ion, generating nitrite anion.


117

Chapter 7

S-Nitrosation In recent times the most widely studied conventional nitrosation reactions have been those where reaction occurs at a sulfur atom. These reactions often generate unstable products, which may account in part for the lack of information in the early literature, but much interest has been generated in this field, particularly for the nitrosation of thiols, stemming from the discoveries surrounding the biological properties of nitric oxide. In particular the chemistry of the products of thiol nitrosation, S-nitrosothiols, has been much examined. There is now a large body of literature, and given the importance of these compounds, a separate chapter (Chapter 8) has been devoted to their chemistry. 7.1.

Nitrosation of thiols The simplest example of S-nitrosation in organic chemistry is probably the nitrosation of thiols, generating S-nitrosothiols (formerly called thionitrites) RSNO, as in Eq. (246). The reaction is written here for nitrous acid

RSH + HNO2

HsO"^ ^

RSNO + H2O

(246)

nitrosation, but in principle any of the nitrosating agents discussed in Chapter 1 will be effective. The reaction formally resembles the nitrosation of alcohols, but there are a number of important differences which will be discussed later. Even though the reaction was discovered early in the 20th century, it has not been much studied until recent times, mainly because the product is often unstable in its pure state and also the reactions are very rapid and have required fast reaction techniques to study the kinetics. Since the discovery of the important biological properties of nitric oxide (discussed in detail in Chapter 11), intense interest has been aroused in the chemistry of RSNO species since some are naturally occurring, they have powerftil biological properties and can release nitric oxide. It has been suggested that RSNO compounds act as storage and transport agents of nitric oxide in vivo. A few rather unstable RSNO compounds were isolated about a hundred years ago, notably the phenyl derivative 6 [388], and later the rather more stable PhSNO 6

(CH3)3CSNO 7

Ph3CSNO 8

118


t-butyl, 7, and triphenylmethyl, 8, derivatives [389-90]. A variety of nitrosating agents have been used successfully, including nitrosyl chloride, alkyl nitrites, dinitrogen tetroxide, dinitrogen trioxide and nitrous acid. To this list must now be added the use of NO/O2 as a reagent. The most well-known examples, both of which are stable in their pure solid states, are S-nitroso-Nacetylpenicillamine (SNAP), 9, and S-nitrosoglutathione (GSNO), 10, both of which can readily be prepared using aqueous acidified nitrous acid, often using SNO H3C^.SNO HOc H02C"^NHAc 9

N^^C02H NH2

H

O

10

methanol as a co-solvent [391-2]. Solutions of other RSNO species can readily be generated from the thiol and nitrous acid and the solutions are often sufficiently stable, particularly in acid solution, to be used in situ in synthetic and mechanistic studies. Stability of the pure materials and stability of their solutions have often been confiised. In contrast to the situation with alcohol nitrosations, thiol nitrosations are effectively (at least fi-om the synthetic viewpoint) quantitative. However an attempt has been made to determine the equilibrium constants for example of Eq. (246), [393]. There is no measurable intercept for the plots of the firstorder rate constants against excess thiol concentration, as there is for the alcohol reactions, so the equilibrium constants (AT) must be large. Analysis of the final reaction mixtures for thiol concentration using EUman's reagent reveals low concentrations of fi-ee thiol. Approximate values for K have been determined for a number of thiols, and are shown in Table 17. The values are necessarily approximate because of the error of measurement of a very low thiol concentration. It has been shown that solutions of SNAP and GSNO made up from the pure solid materials contain approximately 0.8% free thiol. These levels are not significant fi-om the synthesis angle, but are important in the generation of nitric oxide from RSNO solutions, which will be discussed in Chapter 8. A kinetic study of the reaction of nitrous acid with a number of thiols has shown that the well-known third-order rate equation, Eq. (9), is obeyed. There is also substantial nucleophilic catalysis [34]. Some results are shown in Table 18. There is not a large spread of k values which suggests that reactions are perhaps close to the diffiision limit for H2N02"^/NO^ attack, even though

119

S-nitrosation

Table 17 Values ofATforEq. (246) Thiol Penicillamine

Kiyr

Cysteine

6x10^

Thiomalic acid

3x10^

N-Acetylpenicillamine

7x10'^

3x10^

Table 18 Values of A:/M~^s~^ in, rate = A:[RSH][HN02][H30^] and kîyr^s-^ ^2[XN0][RSH] RSH Cysteine

in rate =

k 350

k2{cmo)

^2(BrN0)

1.2x106

5.4 X 104

;t2(0NSCN) 7.2x102

220

1.1x106

4.6x104

7.5 X 102

N-Acetylcysteine

1600

1.0x107

4.6x105

1.7x103

Penicillamine Glutathione

790 1080

1.2x107

5.6x105

3.9x103

Thioglycolic acid

2630

1.4x107

9.2 X 105

2.4x104

Mercaptosuccinic acid

1300

-

8.5 X 104

5.8x103

Cysteine methyl ester

values are a factor of 3 or so below that expected. For the nucleophile catalysed reactions there is a further demonstration of the reactivity sequence CINO > BrNO > ONSCN. The reaction mechanism is quite different from that proposed for the corresponding reactions of alcohols, where no nucleophilic catalysis occurs (Chapter 6). For thiols, it appears that reactions take place by direct rate-limiting attack of XNO at the sulfur atom, just as is the case for amines, there being no evidence for any intermediate or rate-limiting proton transfer. Reactions are of course much faster for thiols than for amines, since there is little protonation at the reactive site. Reactants containing both amine and thiol groups undergo preferential S-nitrosation, though there are reported cases where subsequent internal S to N rearrangement of the NO group can occur. Normally the intial product of nitrosation of cysteine is the S-nitroso derivative which often decomposes to the disulfide (see chapter 8). In one case, over a longer timescale the isolated product is the thiurancarboxylic acid, 11, which can be rationalised in terms of initial S-nitrosation (a strong yellow colour develops), NO group rearrangement to N, development of the diazonium ion, 12, and ring closure by internal attack of the thiol group and

120


COOH

-7S / COOH 11

HSCH2CH V

V 12

NSN

loss of nitrogen. This work has been extended to include the reactions of cysteine esters [394]. Nitrous acid nitrosation of cysteine has been subjected to a more detailed kinetic study [395], in which it was revealed that both HSCH2CH(^>ffl[3)COOand HSCH2CH("4NIH3)COOH forms react in acid solution, both with rate coefficients close to the encounter limit. A number of RSNO species have been detected in vivo including GSNO and some S-nitroso proteins. A S-nitroso protein derivative of serum albumin has been isolated and characterised [396-7], and it has been proposed, since nitric oxide itself has a very short half-life in vivo, that it circulates in mammalian plasma mainly as the S-nitroso derivative of serum albumin, although this has not been demonstrated unequivocally. The nitrosation of the free thiol group (Cys-34) in both human and bovine samples of serum albumin, by nitrous acid showed the almost immediate formation of the absorption peaks characteristic of RSNO structures at 345 and 545 nm. Nucleophilic catalysis also occurred and the rate equations and values of the appropriate rate constants were very similar to those obtained for cysteine itself These results argue strongly in favour of a conventional S-nitrosation effected by H2N02^/NO"^ or XNO [398]. A second, slower reaction occurred when nitrous acid was present in excess over the thiol concentration, which was interpreted in terms of a Nnitrosation at the indole nitrogen atom of the tryptophan residue in the protein. It has been reported that other nitrosating agents can also effect Snitrosation of-SH groups in proteins, notably a dinitrosyl iron complex [399] and also an aromatic N-nitrosamine [137]. Three reports [91,400-1] give details of a kinetic study of the nitrosation of a number of thiols, including cysteine and glutathione, using NO/O2 at physiological and other pH values. Reactions are zero-order kinetically in the thiol and the rate equation is the same as that reported earlier for the oxidation of nitric oxide in aerated water, i.e. Eq. (68) on p.25. There is also agreement between the measured third-order rate constants here and the published value for the autoxidation of nitric oxide. The results are interpreted in terms of rapid reaction of N2O3 with the thiols, Eq. (247), and that this reaction must compete very effectively with the hydrolysis of N2O3, Eq. (248), under these conditions.

S-nitrosation

121

RSH + N2O3 + H2O = RSNO + NO2"" + H3O+

(247)

N2O3 + 3H2O = 2N02~ + 2H3O+

(248)

The kinetics of the reactions of a number of alkyl nitrites with thiols have been published [402]. Reactions have been examined in water over a pH range 6-13. In all cases the S-nitrosothiols are formed. The pH dependence shows clearly that reaction occurs via the thiolate form RS", but there are complications in many cases e.g. for cysteine and glutathione, where there are many protonation equlibria. For N-acetylcysteine and thiolglycolic acid, where only one RS~ species is possible, there is a good correspondence between the calculated curves and the experimental curves i.e. there was good agreement with the published p^^ values. Table 19 shows some of the results. All the thiols studied show much the same reactivity except for thioglycolic acid which is somewhat more reactive. Within the alkyl nitrites studied, the dramatic activating effect of an electron-withdrawing substituent is clearly marked, along with the smaller steric effects for the primary, secondary and tertiary structures. Most of the reactions are sufficiently rapid to require stopped-flow techniques, whereas the reactions of the trichloro-substituted nitrite are too fast to measure, even by this method. All of the thiols showed the S-shaped curves when the rate constants were plotted against pH, and all levelled out at high pH to give k(\\m), indicating that reaction occurred via the thiolate anion and not the free thiol. An example is shown in Figure 5, for the reaction of cysteine with three alkyl nitrites. 7.2

Nitrosation of thiocarbonyls Thiocarbonyl compounds including thioketones, thioureas and thiones are particularly powerful sulfur nucleophiles, so it is no surprise that electrophilic nitrosation occurs readily at the sulfur atom in these compounds. In general, the range of nitrosating agents will react, generating initially the nitrososulfonium ion, Eq. (249). These species, usually coloured yellow in / 0 = S + XNO — ^

^C^S-NO ^ - ^

^C-S-NO

+ X"

(249)

solution, are rather unstable and often appear only as transient species before decomposition occurs. The most widely studied reactions have been those of thiourea and alkyl thioureas. Werner showed [403] that two further reactions occur, one dominant at low acidities which leads to nitrogen and thiocyanate

122


"o E £ CM

Fig. 5. Second-order rate constants for the reaction of cysteine with 1, (CH3)2CH(CH2)20NO, 2, (CH3)2CHONO and 3, (CH3)3CONO plotted against pH.

S-nitrosation

123

Table 19 Values of the limiting rate constant ^(lim)/M~"^ s~^ for the reactions of some alkyl nitrites with thiols RONO (CH3)3CONO (CH3)2CHONO CH3CH2ONO C2H50(CH2)20NO Cl(CH2)20NO CI3CCH2ONO

1 1.7 11 28 169 1045

2 1.5 12 25 165 1100

A:(lim)/M-1 s-1 4 3 1.8 1.8 11 12 28 31 159 169 1070 1010 1

H

OH

13

14

1 15

H 16

M"-^ and the rate constants for their formation suggest that they all occur at the diffiision limit and finally, all have considerable potential as catalysts in nitrosation reactions. 7.3.

Nitrosation of organic sulfides It is to be expected that the sulfur atom in simple sulfides R2S would be approximately as nucleophilic as the sulfur atom in thiols. However there is now no convenient leaving group, so there is no obvious product to the reaction once the S-nitroso ion has been formed, Eq. (263). There is however some R2S + XNO ^^=^ R2S"^0 + X-

(263)

evidence that these ions do form. Very early work reported that simple alkyl sulfides react both with nitrous acid and alkyl nitrites giving coloured solutions, characteristic o f - S - N O species, although no products were identified [414]. The presence of ions of the type R2S"^0 can be inferred however from the substantial catalysis of the nitrosation of N-methylaniline by added dimethyl (CH3)2S + PIN02 + H3O+ ^ ; = ^

(CH3)2SÔ + 2H2O

(CH3)2SÔ + C6H5N(H)CH3 —

C6H5N(NO)CH3 + (CH3)2S + H"^

(264)

(265) sulfide, readily interpreted via Eq. (264-5). Catalysis is quite substantial and is approximately the same as that generated by bromide ion in the same reaction [415]. Diazotisation of aniline is also catalysed by S-methyl cysteine. Similarly there is strong kinetic evidence that reactions of this type can also occur intramolecularly. The high reactivity towards nitrous acid nitrosation, of both methionine and S-methylcysteine, compared with alanine has been interpreted [416] in terms of initial nitrosation at the sulfur atom followed by an internal S- to N-rearrangement of the nitroso group. Later, a detailed kinetic investigation of the nitrosation of thioproline in acid solution confirmed that there is an additional reactivity here relative to the same reaction of proline, which must be associated with the presence of the

S-nitrosation

129

sulfur atom. The authors [417] proposed an initial S-nitrosation followed by an internal S- to N-rearrangement of the nitroso group, after a proton-loss, as outlined in Eq. (266). A similar pathway has been identified for the nitrosation of thiomorpholine [418]. Molecular orbital calculations using the frontier-orbital approach have concluded that the 'soft' electrophile NO"^ will react preferentially at sulfur rather than at nitrogen, as observed in the examples discussed.

K H

^C02H H

a N

(266)

!i\

^92^^

(^

K H

^C02H H

(3^

CO2H

N

NO

CO2H

H

Other nitrosating agents such as nitrosamines and a nitrososulfonamide appear also to be able to transfer the nitroso group to both S-methylcysteine and methionine. In the case of the nitrosamine reaction, kinetic measurements were carried out in the presence of hydrazine [129], and this drives the reaction to completion, which strongly suggests that the S-nitroso ion reacts rapidly with hydrazine, Eq. (267). In this instance S-methylcysteine and methionine R(R')S"^0 + NH2NH2

fast

^

RSR' + Decomposition products

(267)

are acting as catalysts in the denitrosation of the nitrosamine. Both sulfides have a reactivity comparable to that of bromide ion in the same reaction, whereas cysteine and glutathione are significantly less reactive. The S-methyl group does then appear to contribute to the nucleophilicity of the sulfiir atom. The reactions of a disulfide with dinitrogen tetroxide probably involves an initial S-nitrosation [419]. The isolated products however, under these rather forcing conditions are products of oxidation, Eq. (268), with evidence

130


N2O4 PhSSMe

•

PhSS(02)Ph + MeSS(02)Me

(268)

for the intermediacy of both C^H5SN0 and a sulfmyl nitrite C^H5S(02)NO. The photolysis of disulfides in the presence of nitric oxide yields a nitrosothiol [420], Eq. (269), presumably by reaction of the dimethyl sulfide radical generated, with nitric oxide. CH3SSCH3

hv

•

2CH3S*

NO •

2CH3SNO

(269)

7.4.

Nitrosation of sulfinic acids It has long been known that alkyl and aryl sulfinic acids react readily with nitrous acid to generate sulfonyl hydroxylamine derivatives, Eq. (270). The reaction has been used to synthesise some hydroxylamine derivatives, to 2RSO2H + HNO2 = (RS02)2NOH + H2O

(270)

characterise long-chain aliphatic sulfinic acids and for the quantitative determination of sulfinic acids by direct titration with an acidic solution of sodium nitrite [15], p. 187 and references therein. It appears from the stoichiometry that two consecutive nitrosation reactions are involved here. Kinetic measurements show that the rate of disappearance of reactant is the same as that for the appearance of the product which suggests that the initial nitrosation is rate-limiting [421]. The kinetic dependence upon the acidity reveals that there are two pathways, one via the undissociated acid and one via the sulfinate anion. The latter is, as expected, the more reactive species, reacting at or close to the diffusion limit. Over the acidity range studied (0.052.0 M) reaction is first-order in nitrous acid and there is the normal catalysis by C r , Br-, SCN- and SC(NH2)2. Thus the reagent is either H2N02"^/NO^ or XNO. All the experimental results accord with the outline mechanism in Eq. (271-2), where S-nitroso sulfinate (or sulfonyl nitrite) is first formed, which RSO2H + HNO2 = RSO2NO + H2O

(271)

Fast = (RS02)2NOH

(272)

RSO2NO + RSO2H

then act as a nitrosating agent itself, rapidly reacting with another sulfinic acid molecule to give the hydroxylamine derivative.

S-nitrosation

131

Alkyl nitrites also yield sulfonyl hydroxylamine derivatives from their reaction with sulfinic acids [422]. No doubt all nitrosating agents are capable of bringing about this reaction. Nitrososulfinates have been postulated as reactive intermediates on many occasions. Some have been isolated from the reaction of sulfinic acids with dinitrogen tetroxide in ether at - 2 0 T [423]. They form unstable brown crystals with the N=0 IR absorption band near 1840 cm~^ at somewhat shorter wavelength than in nitrosothiols (1490-1700 cm~^), no doubt due to the powerful electron-withdrawing effect of the sulfonyl group. They decompose readily on warming giving off nitric oxide, and in solution in non-aqueous solvents bring about nitrosation of alcohols, amines and thiols [424-5]. 7.5.

Nitrosation of thiocyanate ion Nitrosyl thiocyanate ONSCN is too unstable to be isolated in its pure state. Nevertheless it is known as a blood-red species stable in solution at low concentration [426]. Such solutions, which can be prepared in a variety of ways, are known to be potent electrophilic nitrosating agents. Synthetic procedures have included the reaction of nitrous acid with thiocyanic acid, nitrosyl chloride with silver thiocyanate and ethyl nitrite with thiocyanic acid [427]. In addition, solutions in aqueous acid can be generated from nitrous acid and thiocyanate ion as discussed in Chapter 1, section 1.1.3. It is generally believed that the nitrogen atom of the nitroso group is bound to the sulfur atom of the thiocyanate, so this is an example of S-nitrosation. Molecular orbital calculations support this view [428]. The equilibrium constant for ONSCN formation i.e. for Eq. (273), is HNO2 + H3O+ + SCN" ^ i = ^ ONSCN + 2H2O

(273)

30 M~^ at 25°C [27], and this relatively large value is primarily responsible for the major catalytic effect of thiocyanate ion when it is added to solutions of nitrous acid. The rate constant for its formation, Eq. (274), has not been Rate = ^[HN02][H30+][SCN-]

(274)

measured directly, but rather by the further reaction of ONSCN with a very reactive substrate, such that the rate of this reaction is much greater than that for the reverse reaction of ONSCN formation, i.e. its rate of hydrolysis, so that the rate of ONSCN formation is rate-limiting. Values of ^ (Eq. (274)) at O^C were obtained in this way, of 1500 and 1460 M~'^s~^ using aniline and azide ion as the reactive scavengers [35] and more recently values of 1.2 x 10"^ and 1.1 x 10^ M'-^s"^ were determined at

132


25°C using hydrazoic acid and thioglycolic acid as the reactive species [34,188]. The thiocyanate-catalysed nitrosation of two thiones also gave values in the same region [19]. All the results are consistent with a rate-limiting step where H2N02'^/NO^ reacts with SCN". Thiocyanate ion also undergoes a nitrosation reaction using nitrosamines [128], S-nitrosothiols [429] and a nitrososulfonamide [77] as the nitrosating species. The powerful catalytic effect of thiocyanate is utilised in a procedure for the denitrosation of carcinogenic nitrosamines, if reaction is carried out in the presence of a nitrite trap to drive the reaction to completion i.e. to the quantitative formation of the secondary amine (see Table 12, p.66). Similarly it can allow the ready transfer of the nitroso group from a nitrosamine to another amine (a trans-nitrosation reaction) [130-2], and in the transfer from an alkyl nitrite to a secondary (or primary) amine in an alcohol solvent [408], Eq. (275-6). The role of thiocyanate catalysis of nitrosamine formation in vivo, at the stomach pH, particularly of smokers, has been examined [430].

R2NNO + R2'NH

HsO^, SCN" •

H30^, SCN" RONO + R^'NH -^ • ^ ROH

R2NH + R2'NN0

ROH + R^'NNO ^

(275)

(276)

7.6.

Nitrosation of sulfite/bisulfite Nitrous acid in mildly acid solution reacts with sodium bisulfite very readily, generating a yellow-coloured solution believed to contain the nitrososulfonic acid anion, Eq. (277). This is unstable and reacts further with HNO2 + HSO3- = ONSO3- + H2O

(277)

any excess bisulfite to give hydroxylamine disulfonate, Eq. (278). This is the basis of the important industrial process, the Raschig synthesis, since ONSO3- + HSO3-

2H30^ ^—

(H03S)2NOH

(278)

hydroxylamine can readily be obtained from the disulfonate by hydrolysis. The final product of the nitrosation reaction is made up of two successive nitrosation steps, and is reminiscent of the corresponding reaction with sulfinic acids. The reaction is made more complicated by further sulfonation of the product and also by the hydrolysis of a proposed intermediate which results in

S-nitrosation

133

some nitrous oxide formation. Kinetic studies, which are complicated by these further reactions, nevertheless support the general idea of the rapid formation of 0NS03~ and a further slower nitrosation by this intermediate [431]. As expected, other nitrosating agents, including dissolved nitric oxide in aerated solution [432], some iron nitrosyl complexes [433], alkyl nitrites [77], nitrososulfonamides [77], and S-nitrosothiols [429] react similarly. Reactions in acid solution (with nitrous acid) take place via the bisulfite form, whereas at higher pH with some of the other nitrosating species, reaction occurs via the sulfite dianion. There appear to be no references to the use of sulfite/bisulfite as a catalyst for other nitrosation reactions. Strangely, sodium bisulfite and ascorbic acid are among the best nitrous acid scavengers in the prevention of nitrosamine formation from secondary amines in cosmetic products [434]. It may well be that the decomposition of 03SSN0~ is, under these circumstances much faster than the amine nitrosation reaction, leading to reduction of nitrous acid to nitric oxide. 7.7.

Nitrosation of thiosulfate ion Solutions of nitrous acid and thiosulfate ion also generate a yellow colour very rapidly. This is believed to be a S-nitrosation reaction, generating the nitrosothiosulfate anion, Eq. (279). This species is not very stable, even in HNO2 + H3O+ + S2O32- = ONSSO3- + 2H2O

(279)

2ONSSO3- = S40^^- + 2NO

(280)

solution and decomposes to give nitric oxide and the tetrathionate anion, Eq. (280). The rate law for the formation of the yellow species, Eq. (281), is made up of two terms, one representing reaction of H2N02"^/NO"^ with S203^~ and Rate = A:[HN02][S2032-][H30+] + )t'[HN02]^

(281)

the other the rate limiting formation of N2O3 [30]. The value of k here is 18,000 M~2s~^ which is somewhat larger than that believed to be the encounter-controlled limit (7,000 and 11,000 M'^s"^ for neutral and singly charged substrated respectively), but may have an additional electrostatic effect here for a doubly-charged reactant. So a value of -18,000 M'-^s"^ may represent the encounter-limit in this case. To date no other doubly negatively charged species have been examined to check this possibility. Analysis of the spectrum of the yellow species gave a value of 1.66 x 10' M~2 for the equilibrium constant for its formation from nitrous acid and thiosulfate ion. This is one of the largest values known for this type of

134


reaction, and if 0NSS03~ has any reactivity as a nitrosating species then it follows that thiosulfate ion should be a good catalyst for nitrous acid nitrosation reactions. This turns out to be the case in the nitrosation of Nmethylaniline, hydrazine and sulfamic acid [435]. However because of the very large value of ^XNO ^^^^' î^^^^s acid is quantitatively converted to 0NSS03~ even at quite low 8203^" concentration. This means that the measured rate constant will not increase with [8203^"]. In the reaction with Nmethylaniline, 0NSS03~ is approximately 10"^ less reactive (as measured by the bimolecular rate constants) than is ONSCN. A later study of the diazotisation of anilines in the presence of these nucleophile catalysts [436] showed that, in the case of the 8203^" reactions, the results are consistent with nitrosation via the ON8803~ species, and the following reactivity order was established:- CINO > BrNO > N2O3 > 0N8CN > [(NH2)2C8NO]^ > ON8803"^. It is not surprising that ON8803~ is such a poor electrophilic nitrosating agent given that it is a negatively charged species. As expected, other nitrosating agents such as alkyl nitrites, nitrososulfonamides and 8-nitrosothiols also react with thiosulfate, generating the 8-nitroso anion. 7.8.

Nitrosation of inorganic sulfide There appears to be no report of a reaction with sulfide ion or the hydrosulfide with nitrous acid in acid solution, since even though both are powerful 8-nucleophiles, protonation will occur in acid solution generating hydrogen sulfide. However there are some fragmented accounts of reactions, which generate the typical yellow colour of 8-nitroso species, with nitrosating agents which are effective in neutral or alkaline media. Thus, aerobic solutions of nitric oxide [437], alkyl nitrites [429], 8-nitrosothiols [429] and a nitrososulfonamide [77] all generate a yellow colour when treated with either Na28 or Na8H (and even Na282 in one case) at pH values greater than 7. The spectra look identical in each case and it is likely that the same product is formed. The species has been identified as the nitrosodisulfide (or perthionitrite) ion ON88~, rather than the expected ON8~ ion. It is not obvious how this ion is formed. A number of possibilities can be proposed, perhaps the most likely is that ON8~ is first formed and that this reacts fiarther with H8~ to generate the disulfide product, Eq. (282-3), but this is a matter for speculation. Another possibility is that H88~ is first formed and XNO + H 8 ' = ON8- + H^ + X-

(282)

ON8- + H8~ -^ ON88-

(283)

S-nitrosation

135

it is this species which undergoes S-nitrosation. The yellow colour with an absorption maximum 410 nm is sufficiently stable in solution to allow the possibility that this reaction might be used as a quantitative determination procedure for S-nitrosothiols or alkyl nitrites, or even sulfide ion.


137

Chapter 8

Synthesis, properties and reactions of S-nitrosothiols 8.1.

Synthesis As outlined in Chapter 7, S-nitrosothiols can be generated in solution by nitrosation of the corresponding thiols by, in principle, any of the conventional nitrosating agents. By far the most well-known stable (in the pure form) examples are SNAP, 9 and GSNO 10 described earlier (p. 118). There is currently an overwhelming interest in the chemistry and physiology of Snitrosothiols (henceforth described as RSNOs), given their probable involvement in the biological properties of nitric oxide. There has consequently been a considerable effort directed at the isolation and characterisation of a whole range of RSNO species. Here are some examples which have been synthesised and characterised in recent years:- the cysteine derivatives 17 [438] and 18 [439], S-nitrosocaptopril 19 [440], the sugar derivatives 20 [439], 21 [441] and 22 [442], the dinitroso derivative 23 [443], the highly hindered aromatic structure 24 [444] and the SNAP derivative containing the amidine group 25 [445]. In addition, S-nitrosoproteins and peptides have been isolated along with derivatives containing different fluorophores with the aim of using such compounds as intracellular probes [446]. Much of the chemistry of S-nitrosothiols however has been conducted using solutions of S-nitrosothiols generated, usually from nitrous acid and thiols. It appears that the stability of such solutions is sufficient to allow the solutions to be used in situ. The reasons for this stability will be discussed later in this Chapter. As a result of the major interest in the chemistry and physiology of S-nitrosothiols a huge literature has built up since -1990. There have been a number of review articles, many of them dealing with the biological properties associated with those of nitric oxide. This topic will be dealt with in Chapters 10, 11 and 12. There are in addition, a number of review articles which deal with the chemistry of S-nitrosothiols, to which the reader is referred, [424], [447], [448] and [449]. 8.2.

Physical properties Stable RSNOs in their pure form are either coloured green or red/pink. In general, tertiary structures such as SNAP 9 are green and primary structures such as GSNO 10 and the sugar derivative 20 are pink. The 3-mercapto-l,2,propanediol derivative forms as an unstable red gelatinous liquid before decomposition sets in at room temperature [439]. The UV-visible absorbances

138


CH3 ONS.

\^SNO

NHs^cr

HO2C

17

(_J-C02H

NHCOCH3

19

18

OAc Ao A.

OAc SNO

AcO-^^-^Y~^*

OAc

OAc

20

21

H

NHAr

SNO

ONS'

H' OH

NHAcH

o

23

22

R C(CH3)3 24 NH2

O N S ^ ^ ^ ^ N ^ ^ ^

Y

NHAcH

25

f \

Synthesis, Properties and Reactions of S-nitrosothiols

139

occur in the 330-350 nm region with extinction coefficients around 500 M~^ cm~^ (due to the n° ~> 7i* transition) and also in the 550-600 nm region with smaller extinction coefficients around 20 M""^ cm"^ (due to the n^,^ -^ 7C* transition). Both of these bands (but more usually the first) have been used to monitor the formation and reactions of RSNOs in solution. The IR spectra have been analysed and assigned [450]. The N-0 stretching (1480-1530 cm"^) and bending frequencies have been identified, as has the C-S bond stretching frequency (600-730 cm~^). Both ^H and ^-^C NMR spectra have been analysed [438]. There is a downfield shift of both a-carbon and a-proton resonances compared with the corresponding values of the relevant thiols. The crystal structure of SNAP 9 has been determined [391] and the only unexpected feature is a rather long C-S bond. There are linear relationships between ^ N NMR shifts and the reduction potentials of RSNOs and also with the ^K^ values of the corresponding parent thiols [451]. The partial double bond character of the S-NO bond allows the possibility of syn and anti isomers, Eq. (284), the preferred isomer depending on the nature of R.

S-N R

^ O

^S-N

(284) \,

Early ideas concerning the formation of nitric oxide from RSNOs concentrated on the homolysis of the S-NO bond. However, thermodynamic, kinetic and theoretical calculations [446] reveal that the bond dissociation energies are around 130 kJ mol'^, which is much too high for spontaneous dissociation at around room temperature. Alternative mechanisms will be discussed later. 8.3.

Thermal and photochemical decomposition On heating, either in the pure state or in a variety of solvents, decomposition occurs to give the disulfide and initially nitric oxide, Eq. (285), which is fiirther oxidised in air to give the characteristic brown ftimes of 2RSN0 = RSSR + 2N0

(285)

nitrogen dioxide. The penicillamine derivative SNAP 9, is quite stable in the solid form up to 148°C, when examined by differential scanning calorimetry and by thermogravimetric analysis [452], whereas refluxing in methanol solvent yields the disulfide readily [391]. However S-nitrosocysteine is too unstable in its pure form to be examined by these techniques. The difference is clearly due to a steric effect not on the unimolecular homolytic fission of the

140


S-N bond, Eq. (286), but rather on the subsequent dimerisation reaction, Eq. (287), which may involve a second RSNO molecule, Eq. (288). RSNO = RS* + NO

(286)

2RS* = RSSR

(287)

RS* + RSNO = RSSR + NO

(288)

Arguments have also been presented in favour of an autocatalytic chain process, which is sensitive to steric effects, in which the chain carrier might be N2O3 [453]. There is EPR evidence [454] that thiyl radicals are formed during thermal decomposition, and this has been reinforced more recently for the thermal or photochemical decomposition by the observation, again by EPR of nitroxide species, which have well-characterised spectra, when reactions were carried out in the presence of radical generators [455]. The photochemical reaction has been studied in more detail. Irradiation of solutions of GSNO 10 at either 340 nm or 540 nm results in decomposition and formation of nitric oxide [456]. Kinetically the reaction is approximately first-order in GSNO. All of the results are consistent with a mechanism based on Eq. (289)-(292) [457]. In the absence of oxygen, homolysis occurs, GSNO = GS* + NO

(289)

GS* + GSNO = GSSG + NO

(290)

GS* + 02 = GSOO*

(291)

GSOO* + GSNO = GSSG + NO + O2

(292)

followed by the reaction of GS* with the reactant, yielding GSSG and NO. Arguments are presented which suggest that the dimerisation of GS* is not a kinetically important step here. In the presence of oxygen, the peroxy radical is formed, which also reacts rapidly with GSNO generating GSSG and NO. Addition of photosensitizers such as Rose Bengal, promotes the reaction significantly [458]. The cytotoxic effect of GSNO on some leukemia cells is much enhanced by irradiation at 340 nm or 545 nm, which is diminished by the addition of oxyhaemoglobin (which reacts rapidly with NO), suggesting strongly that the cytotoxic agent is nitric oxide. This process may well find a clinical use in photochemotherapy.


8.4.

141

Decomposition in aqueous solution Solutions of RSNOs in water, particularly at pH around 7 tend to decompose generating the disulfide and nitric oxide. There is a wide range of reactivity with structure. Nitric oxide has been detected quantitatively using a NO-specific electrode when reaction is carried out anaerobically and also as nitrite ion if carried out aerobically. The products are thus the same as those obtained from the thermal and photochemical reaction, Eq. (285), but can occur in the absence of incident radiation and also at room temperature. The mechanism of this reaction was for some time a puzzle and kinetic measurements gave rates which varied very widely for the same reactant in different laboratories and sometimes in the same laboratory. The puzzle was solved in 1993, [459], when it was realised that this reaction is catalysed by copper ions, and that there is enough copper in distilled water to bring about the reaction. This explains the variation in rate results, given that the copper concentration will vary from laboratory to laboratory and even within the same laboratory from day to day. In the presence of a metal ion chelator such as EDTA, reaction is effectively brought to a standstill, even for the most reactive RSNOs, leaving usually a very small component of the reaction due to the thermal decomposition. This is shown graphically in Figure 7 which shows the absorbance of SNAP at 350 nm as a function of time. Trace (a) shows the effect of added EDTA and traces (b) to (e) the effect of added Cu^"^ [447]. Subsequently it was shown that the active reagent is in fact Cu"^ [460], since reaction was progressively reduced in rate and eventually stopped by addition of the specific Cu'^-chelator neocuproine, and the characteristic spectrum of the Cu"^-neocuproine complex was observed. This finding was confirmed independently by studies in another laboratory [461]. It was further shown that Cu"^ is generated in solutions of RSNOs containing Cu^"^, by thiolate reduction - a well-known reaction. It had been shown that solutions of RSNOs contain low equilibrium concentrations of free thiol [393]. When the thiol concentration is reduced by generating RSNOs with an increasing excess of nitrous acid, the reaction is again progressively halted [462]. This is shown graphically in Figure 8, which gives the absorbance-time plots for the decomposition of SNAP, (a) prepared in situ with excess [HNO2], (b) with equimolar [HNO2] and [thiol] and (c) with an excess [thiol]. There is a huge variation in rate which accounts (together with the possible different copper concentrations) for the large variation reported in the literature. Thus it is possible to stabilise RSNO solutions either by the addition of EDTA (or neocuproine) or by the generation of the RSNO solutions from a significant excess of HNO2 ^^^^ ^^^ ^^^^^' ^^^ example the half-life for the decomposition of S-nitrosocysteine in solution at 25°C in the presence of EDTA is -55 hours.


142

Absorbance 0.4

0.3 4

0.2 ^

0,H

120 Fig. 7. Effect of Cu2+ on the decomposition of SNAP with (a) EDTA, (b) no added Cu2+, (c) added 5 ^M Cu2+, (d) added 10 îM Cu2+ and (e) 50 îM Cu2+ added.

RSNO + H2O set up in acid The equilibrium RSH + HNO2 solution is frozen when the pH is raised to ~7 as the nitrous acid becomes fully deprotonated. For many RSNO compounds the rate equation, Eq. (293), is obeyed. 2+1 Rate = yt[RSNO][Cu^l

(293)

although in some cases there is a small autocatalytic component at the start, which may represent the formation of Cu"^. For some very reactive RSNO compounds there is a zero-order dependence on [RSNO], when the rate of Cu^"^ reduction by thiolate is fully rate-limiting.


143

1.20r

Absorbance

0-60

0.00 1100

Fig. 8. Effect of added N-acetylpenicillamine (NAP) on the decomposition of SNAP prepared in situ with (a) excess HNO2, (b) equimolar HNO2 and NAP and (c) excess NAP.

In principle, values of k in Eq. (293), should enable structure-reactivity patterns to be established, but this is complicated by the different amount of free thiolate present in the different RSNO solutions. There are many other complications which occur in the presence of added thiols. Initially this will promote reactions by increasing the rate of Cu"^ formation, but at higher concentrations there is evidence in some cases that Cu^"^ can become complexed with thiolate, rendering it less available for Cu"^ generation, with a consequent decrease in reactivity. In some cases this effect is very pronounced, for example in the reaction of S-nitrosopenicillamine. Penicillamine itself is used medically to treat Wilson's disease, which relies on its ability to complex free Cu^"^. This is the dominant effect when even low concentrations of penicillamine are added to S-nitrosopenicillamine [462]. The outline mechanism is given by Eq. (294) and (295). As both Cu^"^ and RS~ are regenerate, they can be present in truly catalytic quantities. 2Cu2+ + 2RS- = 2Cu+ + RSSR

(294)

144


2Cu"^ + 2RSN0 = 2Cu2+ + 2RS- + 2NO

(295)

Nothing has been estabUshed regarding the mechanism in Eq. (295), but it is likely that Cu"^ becomes bonded to the sulfur atom and also possibly (bidentately) to an amino group, if present, as in the S-nitroso derivatives of penicillamine, cysteamine and cysteine, all of which are particularly reactive. A full structure-reactivity correlation has not been attempted, but there are some startling differences in the A: values, Eq. (293), as shown in Table 20. These results are difficult to explain, given the complication due to the presence of low thiolate concentrations, which have different effects upon different RSNOs. Also the results are probably not reproducible, since it is likely that the thiolate concentration will vary from one sample preparation of the same RSNO to the next. It is also possible that various internally coordinated intermediates will play a part, for example as in structures 26, 27, 28 and 29. The large reduction in rate constant, for example on A^-acetylation of S-nitrosopenicillamine may arise from the much reduced extent of coordination of the copper to the amino group when it is acetylated. Similar effects arise on esterification of the carboxylic acid derivatives.

27

28

29

No metal ion catalysis was observed for added Zn^""", Ca^"*", Mg^"*", Ni-^"^, Co2+, Mn2+, Cr3+ or Fe^"*", although there was some indication of a small effect by Fe^"^. Another decomposition pathway, leading to thiol and nitrous


145

Table 20 Values oik, Eq. (293), for the copper-catalysed decomposition of some S-nitrosothiols. RSNO S-nitrosocysteine ethyl ester S-nitrosopenicillamine S-nitrosocysteamine S-nitrosocysteine S-nitrosothiolactic acid S-nitrosomercaptoacetic acid S-nitroso-N-acetylpenicillamine

klM-h~^ 270,000 67,000 65,000 24,500 1100 300 20

acid (not nitric oxide) formation is brought about by Hg^"^ and to a lesser extent by Ag^. This will be discussed later. A quantum mechanical study [463] has confirmed the high value for the S-N bond dissociation energy, and has further shown that in RSNO bonded to Cu"^, a stable intermediate is formed which results in a weakening (and lengthening) of the S-N bond and a strengthening of the N-0 bond, which means that binding with Cu"^ promotes NO release from RSNOs. Another complication has been observed and rationalised in the case of the decomposition of GSNO. At GSNO concentrations around 1 x 10~^ M (where it is convenient to follow its disappearance spectrophotometrically at 350 nm) in the presence of added Cu^^ (~1 x 10"^ M) at pH 7.4, very little decomposition occurs and very little nitric oxide can be detected. Reactions can be made to occur however by substantially increasing the copper concentrations or by the addition of glutathione GSH or any thiol e.g. cysteine [461,448]. However, the reaction occurs quantitatively at much lower GSNO concentrations - 1 x 10~^ M, without the addition of a thiol, or increasing the Cu^"^ concentration, where now reaction has to be followed by measurement of nitric oxide generated (anaerobically) using a commercial NO electrode system. All of these results are consistent with complexation of Cu^"^ by the disulfide product GSSG as shown in 30, [464]. Complexes of this type are well-known and have been identified spectroscopically [465]. The suggested structure is a 1:1 complex as shown in 30, but at much higher [Cu^"^] a 2:1 complex has been identified and characterised [466]. It was possible to identify a characteristic shoulder at 250 nm in the final spectrum from the GSNO reactions [464]. The % yield of nitric oxide increased steadily towards 100% as the initial concentration of GSNO was decreased from 3.7 x 10"^ to 3.0 x 10"^ M, when the added Cu^"^ concentration was 1 x 10"^ M, The mechanistic suggestion was confirmed by experiments in which GSSG was added initially. As the

146


30 added [GSSG] was increased, the % yield of nitric oxide decreased progressively from 90% to 10%. It appears that GSSG is acting as a metal ion (for Cu^"^ in this case) chelator. Further support for these ideas comes from experiments using the Snitroso derivatives of the two dipeptides (GluCys and CysGly) which occur in GSH, and cysteine itself, [467]. The S-nitroso derivative of GluCys behaves just like GSNO, in that at the higher concentrations very little decomposition occurred, since the glutamate residue in the disulfide product which complexes the free Cu^"^, is present. When the reactant concentration is decreased and Cu^"^ is in excess, reaction proceeds smoothly. Conversely, for the S-nitroso derivative of CysGly (and for S-nitrosocysteine), where there is no glutamate residue, the yield of nitric oxide is virtually quantitative at all reactant concentrations. The lower [GSNO] used in these experiments are more closely related to the situation in vivo, so that in principle, nitric oxide release should occur readily. It transpires that Cu"^ can also be generated by thiolate reduction of Cu^"^ when the latter is bound to proteins and peptides. Experiments with Cu^"*" bound to the tripeptide GlyGlyHis, to histidine and also to the protein human serum albumin, [468], were carried out. In each case, treatment with a thiol led to Cu"^ production, observed as the neocuproine complex, and RSNO decomposition was achieved quantitatively using all three sources of Cu2+, although at a somewhat reduced rate compared with free (hydrated) Cu^"^ itself. These results explain why decomposition of GSNO (at the higher concentrations) can be achieved, by the addition of thiol, since this will generate Cu"^ even from the complexed Cu form. As expected, no thiyl radicals were detected in the Cu^'^-catalysed reactions, [469], in contrast to the photochemically-induced reactions.


147

In principle any suitable reducing agent should be capable of reducing Cu^^ to Cu"^ and thereby effect decomposition of S-nitrosothiols. There is clear evidence that ascorbic acid can fulfil this role [385]. At quite low ascorbic acid concentrations (typically 1 x 10""^ M) decomposition of GSNO and SNAP occurs readily as measured by the disappearance of the absorbance due to RSNO and also by the appearance of nitric oxide, when reactions were carried out anaerobically. The reaction rate increased with [ascorbic acid] and was completely halted by the addition of EDTA. At higher ascorbate concentrations a different reaction occurs where ascorbate acts as a nucleophile - a reaction which will be discussed later. There has been much interest in the chemistry of S-nitrosoproteins derived by S-nitrosation of a cysteine residue within the protein. A number have been detected spectrophotometrically and an S-nitrosoprotein derivative of serum albumin has been isolated and characterised [396-7]. The S-nitroso derivatives of both bovine and human serum albumin are quite stable in solution at pH 7.4, at concentrations --1 x 10~^ M, even in the presence of added Cu^"^, but decompose more rapidly at concentrations around 1 x 10~" M yielding nitric oxide [398]. This behaviour, similar to that observed for GSNO, probably results from the complexation of Cu^"^ within various sites in the protein. Interestingly there appears to be a significant thermal reaction for the decomposition of some S-nitroso sugar derivatives. This is shown diagramatically in Figure 9, for structure 20, where there are clearly two reactions, one copper catalysed and one not. Figure 10 shows the complete elimination of the copper catalysed reaction, leaving a significant thermal decomposition pathway. Similar behaviour was observed for the decomposition of the S-nitroso derivative of 1-thio-P-D-glucose, which was generated and used in situ [439]. It has been long known that S-nitrosothiols decompose rapidly in the presence of mercuric ion Hg^"^; this has been incorporated, and much used, into an analytical procedure for the quantitative determination, originally for thiols, [470], but more recently also for S-nitrosothiols. There is a detailed account of the procedure for RSNO determination [471]. The probable mechanism is given in Eq. (296-7), where mercuric ion first forms a complex RSNO + Hg2+ : ; r : ^

RS(Hg)NO^+

RS(Hg)N02+ + H2O — •

RSHg+ + HNO2 + H^

(296) (297)

with RSNO, probably bonding to the sulfur atom; the complex then is attacked by water to yield nitrous acid and the thiol-mercury complex. In the analytical method, nitrous acid is then determined by the Griess method. Rate

148


0

100

200

300

400

500

600

700

time / min Fig. 9. Absorbance-time plots (at 343 nm) for the decomposition of the sugar derivative GPSNO 20 in the presence of added Cu2+, (a) no added Cu2+, (b) 0.1 x 10-5 M, (c) 0.5 X 10-5 M, (d) 1.0 X 10-5 M and (e) 5.0 x 10-5 M , all at pH 7.4.

50

100 time / min

150

200

Fig. 10. Absorbance-time plots (at 343 nm) for the decomposition of the sugar derivative GPSNO 20 with added Cu2+ (a) without EDTA and (b) with added EDTA, both at pH 7.4. measurements using either mercuric chloride or nitrate gave the expected rate equation, Eq. (298), [472]. The reaction was very much faster wîth mercuric


Rate = A:[RSN0][HgCl2] (or [Hg^^])

149

(298)

nitrate than with the chloride; it is known that mercuric chloride in solution exists primarily as the undissociated molecule. There was very little dependence of the rate on the structure of a range of cysteine-based reactants. Reaction also occurred with silver ion, but much less rapidly and there was a curious 2.5 order dependence upon the silver ion, suggesting perhaps that more than one silver ion is involved in complex formation. This reaction is reminiscent of the reaction involving the hydrolysis of RSNO species [409] shown in Eq. (299). This is of course the reverse of the H30^ RSNO + HÔ ^ ^

RSH + HNO2

(^^^^

reaction of nitrous acid with thiols and can only be studied in the presence of a nitrous acid 'trap', to ensure that the reaction is driven to the right. This can be achieved, using azide or sulfamic acid. Hydrolysis is quite slow and occurs at a suitable rate only at > IM H2SO4. Catalysis by CI", Br", SCN~ and SC(NH2)2 is a feature, as expected given the catalytic action of these nucleophiles in the reverse reaction. One further reaction of RSNOs in solution needs to be discussed. It has been found that decomposition in the presence of quite high thiol concentrations leads to a quite different spread of products, including ammonia and nitrous oxide but not nitric oxide. It is therefore not an important reaction in the context of S-nitrosothiols as nitric oxide-donors, but nevertheless deserves to be addressed. Three reports showed clearly that something quite different was taking place at high thiol concentration. Singh et al [473] showed that GSNO in the presence of GSH gives in addition to the disulfide product, ammonia, nitrite ion and nitrous oxide. Independently, Swift [474] established the rate equation, Eq. (300), and showed that the reaction rate was Rate = A:[GSNO][GSH]

(300)

unaffected by the addition of either EDTA or Cu^"^, and that it was quite general for a number of RSNOs. The same rate equation was established for the reaction of S-nitroscysteine and cysteine [475]. Rate measurements over a pH range show that the reactive species is the thiolate anion [476]. A complex mechanism has been proposed [473], based on the formation of an Nhydroxysulfenamide leading eventually to hydroxylamine formation which is then reduced by thiolate to give ammonia, the main nitrogen product. The minor products arise from various side reactions.

150


8.5.

Reactions with nucleophiles Apart from the homolytic S-N bond fission for the thermal and photochemical decompositions, the possibiUty arises that heterolytic bond fission could occur, when RSNO species would act as electrophilic nitrosating agents in the same way as do alkyl nitrites, nitrosamines and nitrososulfonamides. This has proved to be the case and S-nitrosothiols have been shown to react with a wide variety of nucleophiles in this way. The general reaction is set out in Eq. (301), for a negatively charged nucleophile, RSNO + X- = RS- + XNO

(301)

although this need not necessarily be the case. The fate of XNO depends on the nature of X. Some important reactions are discussed separately. 8.5.1. Reactions with thiols (transnitrosation) This is the exact analogy of the transnitrosation reaction which occurs between alkyl nitrites and alcohols, allowing in this case a new S-nitrosothiol to RSNO + R'SH ^ ^ = ^

RSH + R'SNO

(302)

be generated by reaction with a thiol, Eq. (302). Early experiments [424 and references therein], showed that these reactions occur readily. Additionally, decomposition products of both S-nitrosothiols, the disulfides, as well as the mixed disulfide RSSR' were among the final products. Rate and equilibrium constants for the exchange process have been determined by spectrophotometry [477-8]. The reaction takes place even in the presence of EDTA and is unaffected by the presence of Cu^"^, and is generally much faster than the copper-catalysed reaction. These reactions are generally rapid and require the stopped-flow technique in order to determine rate constants. By using a reasonable excess of R'SH over RSNO it is possible to drive the reaction to the right and obtain the rate constants for that process alone. Such kinetic studies [479] show clearly that the reactant nucleophile is the thiolate anion and this is supported by substituent effects when the reactivity order is:- ClC^H4CH2SNO > CH3OCOCH2CH2SNO > OHCH2CH2SNO > CH3CH2SNO. An unreactive RSNO, such as GSNO (at 4 x 10"^ M) will in the presence of cysteine (1 x 10"^ M) generate S-nitrosocysteine, which decomposes quite readily in buffer solutions containing "adventitious" Cu^"*", [480]. Equilibrium constants and rate constants have been obtained for transfer of the NO group from S-nitroso bovine serum albumin to a number of thiols.


151

The results are explained in terms of the electronic structure of the thiol and its pii:^ value [481]. As yet there is no experimental evidence for the presence of a discrete intermediate in these reactions, which are presumed to occur in a one-stage process (S^2-like), but a theoretical study [482] has suggested that an intermediate of structure R'SN(SR)0~ might be involved. This transnitrosation reaction has been widely discussed as a possible reaction in vivo with regard to the biological properties of NO and of RSNOs, particularly with regard to the storage and transport of NO around the body. This is discussed in more detail in Chapter 11. 8.5.2. Reaction with ascorbic acid/ascorbate The reaction where ascorbate brings about reduction of Cu^"^ which then leads to the formation of NO and disulfides from RSNOs has already been discussed in Section 8.4. This occurs at quite low concentrations of ascorbate, even as low as 1 x 10~^ M. It had been reported additionally that decomposition of GSNO occurs with ascorbate even in the presence of a metal ion chelator. Further, the biological properties are enhanced in the presence of ascorbate suggesting that this occurs as a result of increased NO production [483]. A detailed mechanistic study [385] has shown that in fact two reactions occur, one at low concentration of ascorbate, which is the copper-catalysed reaction, Eq. (303), and the other which is dominant at higher ascorbate concentrations, typically 1 x 10~^ M, Eq. (304), which is unaffected by copper.

RSNO + Ascorbate RSNO + Ascorbate

Cu^^

•

RSSR + NO

(303)

•

RSH + NO

(304)

Both reactions generate nitric oxide, but whereas one leads to the disulfide product, the other forms the corresponding thiol. For GSNO it is possible, by choice of ascorbate concentrations to separate the two reactions, but in other cases they can overlap. The second reaction is interpreted in terms of nucleophilic attack by ascorbate at the nitroso nitrogen atom, leading eventually to the formation of nitric oxide (as in the reaction of nitrous acid and alkyl nitrites with ascorbate), the thiol and dehydroascorbic acid. The reaction has been examined over the wide pH range 3.6-13.7, in the presence of EDTA and a detailed analysis of the variation of the rate constant (which increases with pH) with pH shows that both the mono-anion and the di-anion are involved as nucleophiles, with the dianion being the more reactive by a factor of -10^. Calculations show that at the physiological pH of 7.4, most of the reaction (> 90%) proceeds via the reaction of the dianion. There is no evidence of

152

Nitrosation Reactions and the Chemistry of Nitric Oxide 3.00 n

•

Data

- - - Calculated 2.00

yl

1.00 #» ^ i

^ o E

0.00

1E

-1.00

-1

1

4

5

r-

6

-I

8

1 ^

J

•

PH

I

1

1

1

1

10

11

12

13

14

TO-2.00 -3.00 -4.00

r

-5.00 J Fig. 11. The calculated curve and the experimental points for the plot of log (second-order rate constant) vs pH for the reaction of GSNO with ascorbic acid in the presence of EDTA.

reaction via the undissociated form of ascorbic acid even at the lowest pH values studied. Using the literature values for the pK^ of ascorbic acid and its mono-anion, there is excellent agreement between the calculated and observed plot of log (measured rate constant) vs pH, see Figure 11. 8.5.3. Reaction with hydrogen peroxide. The reaction of RSNOs with hydrogen peroxide is treated separately here because the product peroxynitrite, a powerful oxidising agent, is known to react with a variety of biological targets bringing about cell damage. The possibility of its formation in vivo is therefore a matter of some concern for biologists. Its formation from the reaction of nitric oxide with superoxide anion, Eq. (305) is also a reaction which has been much studied and debated in biological NO + O2*

ONOO"

(305)


153

Fig. 12. Repeat spectra for the reaction of S-nitrosocysteine with hydrogen peroxide in the presence of EDTA at pH 11.6 showing the decreasing absorbance due to S-nitrosocysteine at around 340 nm, and the build up, and later falling away of the absorbance at 301 nm due to the peroxynitrite anion.

circles. This reaction is believed to occur at the diffusion limit and has been thought of as a possible route of detoxification for excess NO production. Incidentally there are recent reports that peroxynitrite is formed in good yield and high purity from the reaction of NO with solid potassium superoxide [484] and also with tetramethylammonium superoxide in liquid ammonia [485]. S-Nitrosothiols generally, react with hydrogen peroxide to give initially the peroxynitrite anion, Eq. (306). At very high pH (13.1) the peroxynitrite is RSNO + HOO- = RS- + ONOO" + H^

(306)

relatively stable and is observed as an absorbance at 302 nm (8 1670 M"^ cm~^). At lower pH (11.6) one can observe the decreasing absorbance at 350 nm due to RSNO, concurrently with the build up, and subsequent slower decay, of the 302 nm peak (see Figure 12), whilst at even lower pH (7.4) the only observable spectrum is the decreasing absorbance at 350 rmi, since here, peroxynitrite is sufficiently protonated (p^^ of peroxynitrous acid 6.5) for isomerisation to nitrate to occur rapidly [486]. These results point strongly to a reaction where the nucleophile is the hydroperoxide anion. A more detailed kinetic analysis over a range of pH values shows that this is the case, and also yielded, by two procedures, for the reaction of S-nitrosocysteine, values of

154


11.7 and 11.6 for the ^K^ of hydrogen peroxide, in excellent agreement with the literature value of 11.6. These experiments were all carried out in the presence of EDTA to eliminate any possibility that reaction occurs by preliminary nitric oxide formation, arising from the copper-catalysed reaction discussed earlier (Section 8.4). The biological properties of peroxynitrite will be discussed further in Chapter 11. 8.5.4. Reaction with other nucleophiles A kinetic study has been carried out in which rate constants were obtained for reaction of a number of RSNOs with a large range of nucleophiles. These include the nitrogen nucleophiles [223], primary, secondary and tertiary aliphatic amines, hydrazine, hydroxylamine, azide ion, ammonia, semicarbazide, thiomorpholine and S-methylcysteine, together with the range of sulfur centred nucleophiles [429], sulfite ion, thiourea, thiocyanate ion, thiosulfate ion, thiomethoxide ion, and sulfide ion. There is a large range of reactivity within this group, as expected, and some of the reactions were measured using stopped-flow spectrophotometry. Most were studied over a pH range to establish the exact nature of the reactive form of the nucleophile. The amines all react for example via the free base form of the amine. This predicts an expression for the measured first-order rate constant k^ (when [amine]^ » [RSNO]^) given in Eq. (307), where k is the second-order rate constant for k^=k A:jamine]/(i^^ + [H+])

(307)

reaction of RSNO with the non-protonated form of the amine, [amine] is the total stoichiometric concentration of amine and K the acid dissociation constant of the protonated amine. Good double reciprocal plots of \lk^ vs [H ] were obtained, and there was also good agreement between the determined and literature values ofK^, Table 21 shows the values of A: (Eq. (307)) obtained for the reaction of S-nitrosopenicillamine with the range of nucleophiles. Generally the sulfiir-centred nucleophiles are more reactive than the nitrogen centred ones, as expected. There is some correlation with ^K^ values for the family of secondary amines, but this does not extend to other nitrogen compounds. Similarly there is not a correlation with the Pearson nucleophilicity parameter, the better correlation is with the Ritchie N^ parameter, for those values which have been determined. This is similar to the behaviour of the nitrosation of a wide variety of nucleophiles by a nitrososulfonamide, which is discussed in Section 3.4.


155

Table 21 Values of ^/M~^s~^ for the reaction of S-nitrosopenicillamine with a range of nucleophiles.. Nucleophile N-Methylaniline Semicarbazide Thiocyanate ion Trimethylamine Morpholine Ethylamine Proline Propylamine Hydroxylamine Thiourea Piperidine Pyrrolidine Azideion Thiomorpholine Hydrazine S-Methylcysteine Thiosulfate ion Thiomethoxide ion Sulfite ion

A: (Eg. (307)) Too slow to measure 4.4 x 10"^ 2.5 x 10^ 3.6 x 10-^ 3.8x10-4 4.9x10-4 5.8x10-4 7.8x10"^ 9.9 x 10-4 1.3 xlO-^ 1.7 x 10"^ 7.2 x 10"^ 1.1x10-2 1.5 x 10-^ 3.3 x 10-2 3.8 x 10-2 2.3 40 527

All of these kinetic experiments were carried out in the presence of EDTA to eliminate the possibility of the formation of NO and hence N2O3 via the copper-catalysed reaction. The much less basic secondary aromatic amine N-methylaniline is clearly too weak a nucleophile to undergo this reaction. The reported, rapid formation of the N-nitroso product from N-methylaniline must arise from the copper-catalysed reaction [320]. A feature of the kinetic results in Table 21 is the remarkable reactivity of the sulfite ion 803^", which in turn arises in large part from a relatively low value of the activation energy. Its reactivity is significantly greater than that predicted by its N^ value; it is not clear why this is so, particularly since reactivity towards the nitrososulfonamide is much as expected. There is one reference in the biological literature to the reaction of GSNO and of S-nitroso serum albumin with sulfite ion [487]. It was tentatively suggested that the toxicity of sulfite might be related to its high reactivity towards RSNOs or even to NO itself The hydrolysis of RSNOs in acid solution has already been discussed. In alkaline solution, reaction also occurs and the rate equation Eq. (308) established. The expected 'nitrogen' product from OH" attack at the nitroso

156


Rate = A:[RSNO][OH-]

(308)

nitrogen atom is nitrite anion. This was measured at -50% of the theoretical maximum, which suggest that the harder nucleophile OH~ reacts also at the sulfur atom, 31, which could lead to nitrous oxide formation, though this has

r

OH

N=0 R—S "OH 31 never been shown. This would account for the quite substantial positive intercepts at [amine] = 0 for the plots of the observed rate constants vs [amine] for a range of aliphatic amines given in Table 21, and also parallels the situation for the corresponding reactions of the nitrososulfonamide [75]. No attempt has been made to correlate RSNO structure with its reactivity as a nitrosating species. It is to be expected that electron-withdrawing groups within R would promote reaction as in the case of alkyl nitrites. One feature however does stand out from the kinetic study with sulfite ion [429] - the reactivity of the two S-nitroso sugar derivatives studied, S-nitroso-thio-P-Dglucose and particularly structure 20 is significantly greater than for the structures based around the cysteine molecule. For example the k value for 20 in reaction with sulfite ion is 1.2 x lO"^ M'^s"^ compared with 527 M~ s~ for S-nitrosopenicillamine. This may well be because of the electron-attracting properties of the ring oxygen atoms in the sugars. The biological properties of RSNOs will be discussed along with those of nitric oxide itself, in Chapter 11. 8.6.

Detection and quantitative determination of S-nitrosothiols Since the discovery of the biological properties of RSNOs, associated with those of nitric oxide, there has been intense activity in the analytical field, where a variety of procedures have been applied to this problem, much of it directed at the determination of rather low concentrations of RSNOs which are present in body fluids. Basically there have been two approaches, (a) directed at the direct determination of RSNOs, often preceded by a separation procedure, and (b) aimed at the determination of a derived product from RSNOs, usually nitric oxide itself. In this section only methods appropriate to (a) will be considered, whereas the more numerous methods in (b), aimed at


157

the quantitative determination of nitric oxide will be discussed in Chapter 10, Section 10.5. There have been numerous review articles in this area. More detailed accounts which relate to (a) are given in references [449], [471], [488]. 8.6.1. Spectrophotometric determination ofRSNOs In the absence of competing absorbances, the two maxima at 320-360 nm (e ~ 500-800 M~^cm~^) and around 540 nm (8 ~ 15 M"^cm~^) characteristic of RSNOs can and have been used to determine RSNO concentrations. Some of the detailed wavelengths and 8 values are given in Table 22 and are taken from reference [471]. The extinction coefficients are relatively modest, even for the 320-360 nm absorbances, which means that even with the best modem spectrophotometers the range of detection is usually only in the range 0.1-10 mM, so the method is somewhat limited when analysing RSNOs in body fluids. Nevertheless the method, which is straightforward, has had some application. Often there are several RSNO species present, so a separation procedure is necessary before the assay - this is frequently a high-performance liquid chromatography arrangement. S-Nitrosoglutathione (GSNO) and the S-nitroso derivatives of sulfur-containing proteins have been detected and quantified in blood plasma and in other bodily fluids. The total RSNO concentration in plasma is approximately 1 \\M. 8.6.2. Electrochemical determination ofRSNOs An electrochemical probe for the analysis of low concentrations of thiols has been adapted for the determination of RSNOs as well. Using a Ag/AgCl reference electrode, the working electrode Au/Hg is set in series at both oxidising and reducing potentials. Concentrations in the nanomolar range have been determined in this way. Again, if there is a mixture of compounds, separation by HPLC can precede the electrochemical analysis. The halfreactions given in Eq. (309), and in Eq. (310), are those taking place at the

2RSN0 + 2H+ + 2e — • 2RSH + 2N0

(309)

2RSH + Hg — •

(310)

Hg(SR)2 + 2H+ + 2e

oxidising and reducing electrodes respectively. 8.6.3. Capillary zone electrophoresis In principle this is a simple and reliable procedure in which RSNOs, thiols and disulfides can be determined simultaneously. Individual components become separated on the basis of relative molecular mass and charge. High

158


Table 22 Absorbance maxima (nm) and extinction coefficients (M~lcm~l) for some commonly encountered RSNO species RSNO S-Nitrosocysteine S-Nitroso-N-acetylcysteine S-Nitrosocysteamine S-Nitrosoglutathione S-Nitroso-N-acetylpenicillamine

(^max)l

£335 nm

(^max)2

£545 nm

335 334 333 335 335

503 507 536 586 519

544 545 545 544 591

14.9 16.3 16.1 17.2 7.0 at 591 nm

voltages in the range 10-30 kV are applied across narrow bore capillaries and separation occurs in the flow towards a detector, usually a UV detector set at -335 nm for the RSNO analyses. The Umits of detection are currently in the micromolar range. There is usually a direct relationship between concentration and peak heights, but the method does suffer from a lack of sensitivity, and has thus far, not been applied to analyses of body fluids. 8.6.4. Conversion to nitrite ion All RSNOs react rapidly with mercuric ion as described earlier in the Saville procedure, generating the corresponding thiol and nitrite anion. There are various methods for the reliable determination of nitrite anion. The most widely used has been the Griess reaction which generates, after acidification, nitrous acid, which in turn is converted to a diazonium ion which is coupled with a phenol or an aromatic amine to generate an azo dye with a large extinction coefficient. Unless a separation is effected beforehand, this will of course yield the total RSNO content. The method is highly sensitive with a detection limit in the range 0.1-0.5 îM. Other procedures for the determination of nitrite, include high performance capillary electrophoresis, where nitrite and nitrate can be analysed simultaneously, the use of GC-MS, which, although sensitive and specific, again for the determination of nitrite and nitrate, is not very amenable to routine analyses, and the diazotisation of an aromatic diamine and subsequent formation of a triazole which is highly fluorescent. Usually the latter method involves the diazotisation of 2,3-diaminonaphthalene, Eq. (311), which leads to

"^^2^

I

II

I

N

(311)


159

the formation of the naphthotriazole which can be quantified by fluorescence spectroscopy and a caUbration graph. The detection limits are in the micromolar range. 8.6.5. Conversion to nitric oxide Much more attention has been given to the analysis of nitric oxide itself and a variety of procedures will be discussed in Chapter 10. There are a number of ways in which RSNOs can be converted quantitatively to nitric oxide; consequently there are a variety of analytical procedures for RSNOs which depend on their conversion to NO. This can be achieved in the following ways:(a)

thermal degradation,

(b)

photochemical degradation,

(c)

decomposition in solution by the Cu^'^-catalysed route,

(d)

reaction with ascorbic acid.

The principal methods for nitric oxide determination, which will be discussed more fully later, include :(a) the chemiluminescence method, based on reaction of nitric oxide with ozone and the consequent release of radiation from nitrogen dioxide produced in an excited state, (b) electrochemical methods based either on a Clark-type modified oxygen electrode or an electrode where nitric oxide is oxidised electrochemically on a polymeric metalloporphyrin, (c)

methods based on electron paramagnetic resonance spectroscopy (EPR),

(d)

use of mass spectrometry,

(e)

the oxyhaemoglobin assay procedure.

Whilst it is relatively straightforward and reliable to analyse for both the RSNOs and NO at typically milli- or even micromolar concentration levels, it is an order of magnitude more difficult to do so at the nanomolar levels which occur frequently in vivo. This has been highlighted recently in a paper entitled "Measurement of physiological S-nitrosothiols: a problem child and a

160


challenge" [489]. It is claimed that many of the analyses of RSNOs carried out in plasma, serum and urine depend on conversion to NO or N02~ and in many cases will produce high values due to artifactual formation of RSNOs from RSH and N02~ on acidification due to incomplete removal of endogenous plus blank N02~. It is suggested that this extra contribution to physiological RSNOs may even exceed the naturally occurring levels.


161

Chapter 9

Nitrosation involving metal-nitrosyl complexes There is now a rich chemistry involving the formation and reactions of metal-nitrosyl complexes following the earlier explosion of interest in the related metal carbonyl complexes. Much of the chemistry has been reviewed at various times and an excellent book appeared in 1992 [490] which covers the subject comprehensively. Additionally, in a recent review of many aspects of nitric oxide chemistry, there are contributions on mechanistic aspects of the reactions of nitric oxide with transition-metal complexes, non-heme iron nitrosyls in biology and the coordination and organometallic chemistry of metal-NO complexes amongst others [491]. More recently there is a review article [492] which discusses the coordination chemistry of nitric oxide with reference to bio-inorganic systems. This chapter will deal only with, (a) metalnitrosyl compounds which have been used as electrophilic nitrosating agents and (b) the nitrosation of nucleophile species coordinated to metal centres. Another aspect relating to metal-nitrosyls as NO-donors will be discussed in Chapter 12. Examples of (a) are given in a compilation by Bottomley [493]. Metal nitrosyls have been synthesised by a large variety of procedures. One textbook [494] quotes thirteen different ways which have been used successfully. The more common methods have involved the reactions of nitric oxide, nitrous acid, nitrosyl halides, nitrosonium salts or alkyl nitrites. Some examples are shown in Eq. (312-316), which are typical reactions of these reagents. Reactions can also be brought about by NO-CO ligand exchange and there is one report of the formation of nitrosyl complexes of Ru, Rh and Co using S-nitrosocysteine as the nitrosyl donor. [Ru(NH3)^]^+ + NO + H+ = [Ru(NH3)3NO]^+ + N H /

(312)

Ir(CO)ClL2 + NO+ = [Ir(N0)(C0)ClL2]^

(313)

Ni(PPh3)4 + CINO = Ni(NO)Cl(PPh3)2

(314)

Fe(CO)3(PPh3)2 + RONO + H+ = Fe(CO)2(NO)(PPh3)2 + CO + ROH (315) [CpCr(CO)3]-

Diazald

•

CpCr(NO)(CO)2

(316)

Diazald = N-methyl-N-nitroso-p-toluenesulfonamide 9.1.

Sodium nitroprusside (pentacyanonitrosylferrate II) Sodium nitroprusside [Fe(CN)3NO]^~ 2Na"^ (SNP) has been known since 1850 [495]. It is of course neither a nitro compound nor a prusside but the

162


trivial name has been widely adopted. Its structure has been well established from UV, IR, Mossbauer and crystallographic studies. Much of the early work has been well reviewed by Swinehart [496]. Under appropriate conditions SNP can act as a donor of NO or of NO"*" and also can deliver cyanide ion. Reaction occurs readily with amines, where primary amines yield products of deamination, i.e. nitrogen and an alcohol, (and alkenes), Eq. (317) [497]. Secondary amines give nitrosamines, sometimes as the free species and [Fe(CN)5NO]2- + RNH.

[Fe(CN)5(H20)]-

+ N2 + ROH

(317)

sometimes co-ordinated to the metal in a complex. Any reaction carried out in basic solution is complicated by the reaction of hydroxide ion (which is reversed in acid solution), generating the nitro ligand within the complex, Eq. (318) [496]. The reaction with secondary amines has been examined [Fe(CN)5NO]2- + 20H- = [Fe(CN)5N02]'^ + H2O

(318)

kinetically and a rate equation established, which has both a first-order and second-order term in R2NH [498]. This has been interpreted in terms of a complex formation which can react either with water or with another secondary amine molecule, Eq. (319)-(321). Reaction with ammonia proceeds similarly

[Fe(CN)5NO]^~ + R2NH

,0" ./ Fe(CN)5N^+ NHR2

^

+ H2O

.0" ./ |Fe(CN)5N^+ NHR2

(319)

[Fe(CN)5H20]^~ + R2NNO + H^ (320)

. /,o" |Fe(CN)5N^+ NHR2J

+ R2NH

[Fe(CN)5NR2H]^

+ R2NNO + H^ (321)

leading to nitrogen as a product. These reactions are characterised by the formation of bright and often changing colours, attributed to complex formation. Such colours have been observed with a number of other nitrogen

Nitrosation Involving Metal-nitrosyl Complexes

163

nucleophiles such as indole, pyrrole, phenylhydrazine and hydrazones, although the products have not been identified. Reaction also occurs with hydrazine [Fe(CN)5NO]^- + N2H4 - ^ [Fe(CN)5(H20)]^- + N2O + NH3 + H^

(322)

[499] as outlined in Eq. (322). When the nitroso group in SNP is labelled with '^N, the label appears in the N2O product but not in the NH3 product. So the ammonia derives from the hydrazine reactant. Again it is proposed that a complex of the type [Fe(CN)3(NO)N2H4]^~ is first formed leading to [Fe(CN)5N20]^~ which is unstable with respect to loss of nitrous oxide. Ammonia, hydroxylamine and azide ion react similarly with SNP generating the expected nitrosation products. The reaction with thiolate ion has been much studied. A purple-red solution is formed from cysteine, which is the basis of a well-known qualitative test for cysteine and other thiols. The reaction occurs with the thiolate ion only and not the thiol, and is beheved to be involved in the mechanism of the wellknown hypotensive action of SNP. The overall reaction leads to the disulfide product, but no nitric oxide is formed, except in the case of the photochemical reaction. A reasonable explanation is that a reduced species (Fe^^) is formed, which is reoxidised in the air to give SNP, Eq. (323), (324), (325), so that the O RS + [Fe(CN)5N0]^

i3-

Fe(CN)5K

(323)

SRJ 3-

O

[Fe(CN)5N0]^ + RS'

Fe(CN)5H

(324)

SRJ 2RS*

(325)

RSSR

overall process is just one of SNP-catalysed oxidation of thiolate to disulfide [500]. However the EPR signal assigned to the reduced species has been queried [501] and was reassigned to [Fe(CN)4N0]^~. In addition, when reaction is carried out anaerobically, and allowed to stand for some time, nitric oxide but not cyanide ion can be detected in the products. A modified version of events is given in Eq. (326), (327) and (328), [502]. [Fe(CN)3NO]^-

-^

[Fe(CN)4N0]^- + CN"

(326)

[Fe(CN)4NO]^-

-*^

[Fe(CN)4]^- + NO

(327)

164

Nitrosation Reactions and the Chemistry of Nitric Oxide 45[Fe(CN)^]^+ Fe.2+

6[Fe(CN)4]2- + 6CN-

(328)

Although the transient red intermediate has been examined by electronic and EPR spectroscopy, its structure has not been established beyond doubt. Recently a FTIR examination [503] has been carried out. A band is observed at 1380 cm~^ which correlates with the loss of the red colour (520 nm), which suggests that the structure contains RSNO bound to the iron atom. Free RSNO species show the NO vibrational frequency at about 1500 cm~\ and this might be expected to decrease when RSNO is bound to an electron rich situation. This would identify the structure of the transient red intermediate as [Fe(CN)3(ri'-N-RSNO]^-. The discussion of SNP as an NO-donor and its place in medicine is described in Chapter 12. Other sulfur-centred nucleophiles also generate transient coloured solutions with SNP; these include thioureas, alkyl thioureas, sulfite ion and the hydrosulfide ion. No detailed structure analyses of the intermediates have been carried out, neither has a full product analysis. SNP also reacts with carbon nucleophiles, particularly ketones and other carbonyl compounds containing an acidic hydrogen atom. The reaction with acetone has been the most widely studied. Again it is postulated that an intermediate is formed where the carbanion (or possibly the enolate anion) is bound to the nitrogen atom in the complex, Eq. (329), which then undergoes hydrolysis releasing the oxime derivative, Eq. (330), [504]. Reactions have also been reported for the corresponding carbanions derived from malonic acid derivatives [505].

[Fe(CN)5NO]^

+ -CH2COCH3

,0 Fe(CN)5N( ^. CHCOCH3

4-

(329) 1 4-

.0 Fe(CN)5N( CHCOCH3 J

9.2.

H2O

i3-

[Fe(CN)5H20]^~ + CH3COCH NOH (330)

Other metal nitrosyls Examples of electrophilic nitrosation reactions are much more common for SNP than they are for any other metal nitrosyl. It seems that this is related to the question of whether the M-N-O bond is linear or bent, which is reflected in the value of the N-O stretching frequency V^Q. In nitric oxide itself V-^Q is


165

at 1870 cm~^ and when bound to a metal atom this can either increase or decrease depending on the other ligands co-ordinated to the metal and other factors. In general V^^Q values in the range 1950-1450 cm""^ are associated with a linear or nearly linear arrangement, whereas "bent" bonds show V^^Q values around 1720-1400 cm~^. There is considerable overlap and so there is no quantitative link between the linearity and V^Q. Nevertheless Bottomley proposed in 1973 [506], that this rough correlation could be used to predict whether the arrangement is Hnear and consequently, as a result of the maximum charge transfer from NO to the metal, such complexes would act as NO transfer agents to the usual range of nucleophilic species, amines, thiols etc. Conversely a bent MNO configuration will result in less charge transfer, reducing the capability to bring about electrophilic nitrosation. In the limiting case of a 120° angle, the polarity would become reversed and the complex would act as a source of NO", reacting for example with protons. In SNP the V^Q is at 1944 cm~^ and towards the other extreme is 1515 cm~^ for [Cr(CN)5N0]'^-. The latter shows no tendency to act as a nitrosating agent. Ruthenium nitrosyl complexes also participate in a number of reactions in which electrophilic nitrosation occurs. Here the RuNO group is almost linear in most complexes studied and the N-0 stretching frequency is at 1887 cm~^ in [RuCl^NO]^". Some examples of the nitrosation reactions of ruthenium nitrosyls are shown in Eq. (331), [507], (332), [508] and (333), [509]. [Ru(Cl)(das)2N0]^- + N3- -^ [Ru(hedta)NO]^- + OH" - •

Ru(Cl)(das)2N3 + N2 + N2O [Ru(hedta)N02]^

[Ru(bipy)2(NO)Cl]2+ + N3- — •

(331) (332)

[Ru(bipy)2(H20)Cl]+ + N2 + N2O (333)

Meyer and co-workers [509] showed that a ruthenium nitrosyl will bring about aromatic C-nitrosation, in acetonitrile, when the product becomess co-ordinated, Eq. (334). [Ru(bpy)2(NO)X]2+ + C6H5NR2 —

[Ru(bpy)2X(ONC6H4NR2)]^ (334)

Examples are also known for the metal nitrosyls of osmium and iridium, Eq. (335), [510], and (336), [511]. The iridium complex undergoes a rather [Os(das)2(NO)Cl]2+ + N2H4 —

[Os(das)2(Cl)N3] + H2O + 2H+

(335)

166


[Ir(PPh3)(Cl)3NO]+ + ROH - ^

[Ir(PPh3)(Cl)3RONO] + H"^

(336)

rare nitrosation of an alcohol to give the alkyl nitrite, bonded to the metal in the product [512]. A more recent example [513], reports the synthesis of an osmium nitrosyl [OsX(NO)Q2] where X is CI or Br and Q the anion of quinolin-8-ol. The nitrosyl has a high NO stretching frequency at 1910 cm~^ indicating a high degree of NO"^ in the complex. This is borne out chemically by the formation of oxime products with acetylacetone and propiophenone. When nitric oxide is passed into solutions of copper(II) halides in anhydrous acetonitrile, soluble copper nitrosyl complexes are formed, Eq. (337), [514], which have characteristic N-0 stretching frequencies typical of (CuX2)j, + nNO ^5=^ >/n(CuX2.NO)2

(337)

linearly coordinated nitrosyls. Earlier it had been shown [83] that amine and alcohols could be nitrosated by nitric oxide in the presence of Cu(II) salts. The reaction with amines appears to be quite complex but has been interpreted as a heterolytic fission of the dimer of nitric oxide (N2O2) into NO" and NO"*". The former ends up as nitrous oxide after protonation and loss of a water molecule, and the latter effects conventional electrophilic nitrosation with amines or alcohols. Deamination of primary aliphatic amines is believed to occur both for the free amine and also for the amine coordinated to the Cu(II) halide, both of which accompany some loss of nitric oxide in a displacement reaction. Nitrosyl haems have been synthesised from nitric oxide and a number of heam derivatives, and are examples of iron nitrosyl compounds. These will, under certain conditions, bring about nitrosation of secondary amines [515], although little is known about the reaction mechanism. 9.3.

Iron-sulfur cluster nitrosyls In 1858 the French chemist Roussin prepared salts of the two ions Fe4S3(NO)y~ and ^^2^2^^^)^''^ universally known as Roussin's black and red salt respectively [516]. Their synthesis is amazingly simple experimentally, but mechanistically must be very complex. The black salt precipitates as the ammonium salt simply by mixing sodium nitrite, with iron(II) sulfate and ammonium sulfide. Their structures are shown in 32 and 33, and their general chemistry has been reviewed [517]. There has been a renewed interest in one aspect of their chemistry since the discovery of the amazing biological properties of nitric oxide, since the salts can readily release nitric oxide upon oxidation. This aspect will be covered later in Chapter 12. This section will deal with the possibility of these species acting also as electrophilic nitrosating species. Structural analysis has revealed that the FeNO grouping is linear, although the N-0 vibrational frequency is less than 1850 cm~^ However there

167


NO 2-

ON^

..^S.. ^NO Fe' Fe ON^ ^ S ^ NO

33 32 is NMR evidence [518] in favour of the linear arrangement. There are also references in the literature to reactions with secondary amines such as morpholine where nitrosamines are generated [519]. It is claimed that these reactions are faster than the conventional nitrous acid nitrosation. Formation of carcinogenic nitrosamines in this way may well account for the high incidence of oesophageal cancer is the Linxian valley in Northern China. Here an item in the local diet is a pickled vegetable, prepared using the local water supply, which is high in nitrate and nitrite levels. Analysis of this material has shown that it contains high levels of the red ester [Fe2(SMe)2(NO)4] derived from Roussins' red salt. A fiill chemical study has not been carried out but it is equally possible that nitrosamines could arise from loss of nitric oxide from the red ester, oxidation, generation of N2O3 leading to nitrosamine formation or as a direct nitrosation process from the red ester in the electrophilic NO"^ transfer sense. The co-administration of the red ester and proline to rats results in the formation of nitrosoproline. As a result of these studies the eating habits have been changed in the Linxian Province, resulting in the reduction of the incidence of cancer in the area. 9.4.

Nitrosation of nucleophiles co-ordinated to metals There are a number of cases where conventional nucleophilic species which normally undergo nitrosation with a range of nitrosating species, will also undergo nitrosation when these nucleophiles are co-ordinated to metals as ligands. Usually the products are the same as those generated from the free nucleophiles, but in some cases they remain co-ordinated to the metal. The most studied reaction is probably that of the nitrosation of azide ion co-ordinated to cobalt(III) in an octahedral complex. This was examined in some detail by Haim and Taube, [520], who established the stoichiometric equation Eq. (338), and the rate equation Eq. (339), for the reaction with

168


[Co(NH3)3N3]2+ + HNO2 + H^ = [Co(NH3)5(H20)]^+ + N2O + N2

(338)

Rate = ^[[Co(NH3)5N3]2+][HN02][H+]

(339)

nitrous acid in water. The products N2O and N2 are the same as those obtained from the nitrosation of free azide, and the kinetics suggest that the nitrosating agent is H2N02"^/NO"^. Catalysis by halide ion etc. is also a feature, so reaction can also be initiated by XNO species. The probable mechanism involves the rate-limiting formation of the nitrosyl azide complex [Co(NH3)3N3NO]-^^, which breaks up to give N2O and N2 leaving an unstable pentacooordinate species [Co(NH3)5]^"^ which reacts readily with water giving the octahedral aquo complex. A similar reaction has been observed for a number of other metal azide complexes including [Rh(NH3)5N3]^'^, [Co(en)2(N3)2]^, [Co(en)2(N3)(H20)r^ and [Cr(H20)5N3]2^. Rate constants (Eq. (338)) have been obtained and are given in Table 23. The detailed references are given in [521]. Rate constants are not very different from each other over the range which suggest that the outside nitrogen atom is attacked which is to a large degree protected from the metal environment. There are many examples of reactions of other co-ordinated nucleophiles, notably ammonia [522], amines [523], hydroxylamine [524], water [525], methylene groups [526] etc. Examples of such reactions are shown in Eq. (340)-(344). [Os(NH3)5N2]^^ + HNO2 ~ ^ [Pt(en)2Cl2]^^ + KNO2 —

[Os(NH3)4(N2)2]^^ [Pt(en)[N(NO)CH2CH2NH2]Cl2]'^

[Co(en)2Cl(NH20H)]2++ HNO2 — [Co(NH3)5(H20)]^+ + N2O3 —

(340) (341)

[Co(en)2Cl(H20)]2+-f N2O + N2 (342) [Co(NH3)50NO]2^

[Co(CN)3CH2C3H4NH]2- + HN02 —

HNC5H4CH=NOH

(343) (344)


169

Table 23 Values of A:, Eq. (339), for the nitrosation of azide derivatives Reactants

10-3 klMr^s-"^

Temperature/°C

HN3

0.034

N3-

2.5

0

[Co(NH3)5N3]2+

1.55

25

0

[Rh(NH3)5N3]2+

0.40

25

[cis-Co(en)2(N3)2]+

2.86

25

[trans-Co(en)2(N3)2]+ [cis-Co(en)2(N3)(H20)]2+

0.85

25

0.44

25

[trans-Co(en)2(N3)(H20)]2+ [Cr(H20)5N3]2^

0.11 2.4

25 25

Nitrosation Reactions and the Chemistry of Nitric Oxide D.L.H. WilHams © 2004 Elsevier B.V. All rights reserved.

171

Chapter 10

The biological chemistry of nitric oxide 10.1. Background Until 1987 nitric oxide was very much regarded as an atmospheric pollutant generated at high temperatures from nitrogen and oxygen, Eq. (341). At 298 K the equilibrium constant is ~10~^^, whereas because this is an N2(g) + 02(g) = 2N0(g)

(341)

endothermic reaction (AH° - 90 kJ mol"^) the equilibrium constant increases sharply with temperature to ~10~^ at 2000°C. Thus significant amounts of nitric oxide are generated at these high temperatures which are encountered in industrial processes and car engines. Nitric oxide is readily oxidised in the air to give NO2/N2O4 which leads to respiratory problems, particularly for asthmatics, and which also react with hydrocarbons in the air to give eye irritants. There are also further problems associated with photochemical smogs in which one of the key reactions is the photochemical reaction of nitrogen dioxide with oxygen leading to ozone formation which itself is also a major atmospheric pollutant at ground level. The problems have been much reduced by the use of catalytic converters in car exhausts in which nitric oxide is reduced to nitrogen (and carbon monoxide is simultaneously oxidised to carbon dioxide), and also by lowering the temperature of operation of many industrial processes which are exposed to the atmosphere. It has been estimated that --10^^ tons of nitric oxide are produced naturally, annually, through lightning strikes, which of course becomes dissipated in a very large volume. It was a major surprise to find in 1987 and the following years through the work of Furchgott, Murad, Ignarro, Moncada and others, that nitric oxide is synthesised in vivo and is responsible for a range of physiological processes including vasodilation, inhibition of platelet aggregation, neurotransmission, anti-viral and -bacterial activity, penile erection etc. In addition nitric oxide either in excess or deficiency is believed to be responsible to some degree in a range of diseases in humans. The discovery stemmed from experiments in the 1970s on blood flow in animals, when it was observed that acetylcholine (known to bring about smooth muscle relaxation and hence increased blood flow) acted indirectly on the lining of the blood vessels, the endothelium layer. Furchgott and Zawadzki showed that endothelial cells played a role in the relaxation of smooth muscle by acetylcholine [527]. Some "messenger molecule" was being generated which brought about the relaxation. This was christened "the endotheliumderived relaxing factor" EDRF, which was not identified for some years. Two

172


groups suggested in 1987 that EDRF was indeed nitric oxide [528-9], and even though this was challenged on a number of occasions, there is now general acceptance of this fact. Later it was shown that nitric oxide is synthesised in vivo from L-arginine and its involvement in other physiological roles established. Consequently there has been a massive surge of research activity in this area and there is a very large literature, which continues to grow. The subject has been reviewed on many occasions, mostly covering the biological aspects [530], but some which have concentrated on the chemistry involved [531]. 10.2. Reactions of nitric oxide with oxygen and superoxide One of the most important reactions of nitric oxide is its oxidation by oxygen. In aqueous solution, when oxygen has not been rigorously excluded, the final product is nitrite ion which is formed quantitatively. This reaction has been mentioned earlier in Section 1.6.1. The generally accepted sequence of events is given in Eq. (342-345). Whilst the intermediacy of N2O3 has not 2N0 + 02 = 2NO2

(^42)

NO + NO2 =N203

(343)

N2O3 + H2O = 2NO2- + 2H^

(344)

N2O3 + S — •

(345)

+S-NO + NO2"

been unambiguously demonstrated, its presence is a reasonable inference, given that the addition of a nitrosable substrate S will compete with N2O3 hydrolysis to give nitrosation products. The rate equation has been established [532-4] as Eq. (346) and the value of the third-order rate constant is ~5 x 10^ M~^ s~^ at Rate = it[NO]2[02]

(346)

25°C. This value is constant over the pH range 1-13. Rate constants were obtained from spectrophotometric observation of N02~ or HNO2 formation (depending on the pH of the solution) and also by a conductance method. The rate equation is consistent with the rate-limiting formation of NO2. The reaction of NO with NO2 must compete very effectively with the hydrolysis of the latter (which would give an equal mixture of nitrate and nitrite ion). Data obtained from a numerical analysis of pulse radiolysis experiments predict that nitrite ion will be the only significant product [535]. Two mechanisms have been proposed for the oxidation of NO, one based on the dimerisation of NO, Eq. (347-9), and the other on the intermediacy of the peroxynitrite radical, Eq. (350-2). Both involve a pre-equilibrium

The Biological Chemistry of Nitric Oxide

173

2NO :=^=^ N2O2

(347)

N2O2 + O2 = N2O4

(348)

N2O4 ^5=^ 2NO2

(^"^9)

N O + O2 ^^^

(350)

ONOO*

ONOO* + NO — • ONOONO

(351)

O N O O N O — • 2NO2

(352)

formation of an intermediate and so can account for the very small value measured for the activation energy, (AH^ = 4.6 ± 2 . 1 kj mol~^) but it is not easy to distinguish between them. The solubility of NO in water is ^1.7 x 10"^ M at 25°C and P ^ Q of 1 atmosphere, i.e. is comparable to that of oxygen and nitrogen. In the complete absence of oxygen, a solution of NO in water is indefinitely stable, and undergoes no hydration reaction. The third-order nature of the rate equation for the oxidation of NO in water means that the reaction rate is particularly sensitive to the concentrations of the reactants. For example for an oxygen-saturated aqueous solution ([O2] -^ 1 X 10"^ M) the half-life for NO oxidation, when [NO] is 1 x 10"^ M (a typical value in biological systems), is -50 hours and is clearly of no significance. This values falls to ~5 hours when [NO] is 1x10"^ M, etc. The reaction of the nitrosating agent (probably N2O3, Eq. (343)) developed in aerated solutions of NO has already been discussed in Section 1.6.1. In addition, the oxidation of Fe(CN)^'^~ and 2,2'-azinobis(3ethylbenzthiazoline-6-sulfonic acid) (ABTS) by NO/O2 aqueous solutions showed the same rate law, Eq. (346), as did the oxidation of NO by O2 [533-4], which means that the same reagent is responsible for reaction in all cases, and the rate of its formation is the rate limiting step. It is generally assumed that this species generated as an intermediate in N2O3, although originally the authors [533] were not of this view but believed that some other 'as yet uncharacterised species' was involved. The same third-order rate law is obeyed for the autoxidation of NO in aprotic solvents such as carbon tetrachloride [536], where the final product is nitrogen dioxide, just as for reaction in the gas phase. These solutions have the ability to nitrosate, presumably via N2O4. There is a dramatic increase (~x300) in the rate constant for the aqueous autoxidation when it is within the

174


hydrophobic region of a phospholipid, a biological membrane or a detergent micelle [537] - a finding which may have a bearing on the reaction in vivo for the generation of N2O3 to bring about the formation of S-nitrosothiols from thiols. S-Nitrosothiols have been detected many times in biological systems and there is no clear-cut explanation as to how they are formed. It has been argued [91] that the reaction in water in the absence of biological membranes is too slow, for S-nitrosation by this route, to be realistic. Nitric oxide reacts with superoxide radical anion (produced from xanthine/xanthine oxidase) generating peroxynitrite anion, Eq. (353), [86]. NO + 02*

= ONOO"

(353)

This reaction is believed to be important biologically since peroxynitrite, a powerful oxidising agent, reacts with a large number of targets, effecting cell damage. There has built up a considerable biological literature centred on the formation and reactions of peroxynitrite. The implications of nitric oxide, superoxide and peroxynitrite formation in biological systems have been discussed [538]. The reaction of NO with superoxide has been examined kinetically in vitro, and the measured rate constant, obtained by flash photolysis, is close to the calculated diffusion limit, as expected for a radicalradical reaction [88] Superoxide also brings about the degradation of S-nitrosothiols [539]. The inorganic products are a nitrite/nitrate mixture which accounts for 90% of the total nitrogen; the nature of the organic product was not established. Kinetic measurements with S-nitrosoglutathione and S-nitrosocysteine give second-order rate constants of 1.3 x 10"^ and 7.7 x 10"^ M~^ s~^ respectively. These values are very much smaller than the value for the corresponding reaction of NO, and have been used to support the idea that S-nitrosothiols are generated from NO in biological systems to protect it from very rapid destruction by superoxide, and hence that RSNOs could serve as a reservoir and carrier of nitric oxide in biological systems. Peroxynitrite reacts with a range of enzymes, thiols, oxyhaemoglobin etc. It can effect the nitration of tyrosine, although the mechanism (heterolytic vs homolytic) has not been fully established, the rupture of DNA strands and bring about hydroxylation. The nitration of tyrosine seems to be associated with a number of diseases since nitration products occur in these cases, notably 3-nitrotyrosine. A number of possible candidates have been suggested as the source of the nitration products, but current opinion regards peroxynitrite as the most likely, which effects nitration via a free radical mechanism involving the intermediate formation of NO2. A detailed kinetic analysis of the reactions of NO with both oxygen and superoxide [540] has revealed that the formation of


175

peroxynitrite from NO and superoxide is at least in part responsible for the toxicity of NO, but if the peroxynitrite radical ONOO*, Eq. (350) is also responsible for biological damage, then the reaction of NO with oxygen is important in this context, Peroxynitrite ion also reacts rapidly with carbon dioxide to give the nitroso peroxycarbonate ion, Eq. (351). This in turn breaks down to give ONOO- + CO2 = 0N00C02~

(351)

0 N 0 0 C 0 2 ~ = NO3- + CO2

(352)

nitrate ion and carbon dioxide, Eq. (352). The breakdown is believed [541] to involve homolysis of the O-O bond, generating NO2 and C03* as observable intermediates. The reaction with carbon dioxide is likely to affect the biological properties of peroxynitrite given the in vivo levels of carbon dioxide. This is currently an area of chemistry which requires further study. 10.3. Reaction of nitric oxide with haem proteins The most studied class of iron nitrosyl complexes is that where nitric oxide is bonded to an iron atom in a haem complex. Such complexes with haemoglobin (Hb) and myoglobin (Mb) are well known. Haemoglobin contains four protein chains and four haem groups whereas myoglobin has a single protein chain and one haem group. A typical haem group is shown as structure 34. In these complexes the iron atom is formally in the Fe(II)

oxidation state. The structures and reactions of nitrosyl metalloporphyrins have been fully described [542]. Such species have been characterised by EPR spectroscopy although the interpretation of the spectra is not always clear-cut. Nitric oxide binds very strongly both to Hb and IVIb and the nitrosyls have been synthesised in the laboratory; they do not possess the biological properties of

176


free nitric oxide e.g. of vasodilation. Each iron atom binds nitric oxide, so haemoglobin can bind up to four nitric oxide molecules. Other structural tools such as resonance Raman spectroscopy, NMR and notably X-ray crystallography have been used successfully to characterise the species involved. The observed Fe-NO bond-angle for example has been shown to be 145° in the nitrosyl derivative of horse heart Hb, which is in agreement with the prediction made for (FeNO)^ species. It is a reasonable assumption to make, that nitric oxide is transported in the blood in the form of a nitrosyl haem, although this has never been established. There is here a major problem:- oxyhaemoglobin (OxyHb), which has an oxygen molecule coordinated to iron, reacts with nitric oxide at least in vitro, rapidly and irreversibly to give methaemoglobin (metHb) in which the iron is now in the Fe(III) oxidation state, and nitrate ion, Eq. (353). In addition, deoxyhaemoglobin scavenges nitric oxide to give, reversibly an iron-nitrosyl complex HbFe(II)NO, Eq. (354). Both reactions have relatively oxyHb + NO — • Hb + NO ^^=^

metHb + NO3-

HbFe(II)NO

(353) (354)

large rate constants - 1 x 1 0 ^ M~^ s~^ These reactions are believed to be those responsible for the removal of excess nitric oxide in vivo. One is also the basis of an analytical procedure for nitric oxide (see later). Why is it then that nitric oxide, generated in the endothelial cells of blood vessels, is not rapidly destroyed via Eq. (353-4), as soon as it enters the blood stream? A number of explanations have been advanced, but none is as yet compelling. Blood flow creates a differential gradient of red cells, lowest at the endothelium and increasing towards the centre of the blood vessel. This treatment has been looked at using mathematical models which quantify the situation [543]. Others have shown, although there is no explanation given, that the rate of scavenging of nitric oxide by red cells is - three orders of magnitude less than it is by cell-free oxyHb. A third explanation which has been much discussed and is not accepted by many workers in the field, is that nitric oxide is transported in the blood stream, not as an iron-nitrosyl complex but as a Snitrosothiol, where the NO group is bound to a thiol group of a cysteine fragment which is part of a protein chain at position 93 [544]. S-Nitrosohaemoglobin (HbSNO) can be generated in solution from simple RSNOs such as S-nitrosoglutathione (GSNO) and oxyHb. Solutions are reasonably stable and can bring about vasodilation and inhibition of platelet aggregation. HbSNO has also been detected in biological systems. There are however widely differing reports regarding the level of HbSNO both in arterial and venous blood, from 200 nM to 5 |i,M, obtained by different analytical


111

procedures, but mostly relying on the release of NO and its detection by the chemiluminescence method. The deoxyHb interestingly does not form a Snitroso compound, but rather the iron nitrosyl HbFe(II)NO. The original papers [544] postulated that S-nitrosation occurred via a transnitrosation reaction as in Eq. (355) involving essentially a direct transfer of NO"^. Others GSNO + oxyHb ^^=^

GSH + HbSNO

(355)

have suggested that NO reacts with a vacant haem site to give an iron nitrosyl followed by an intramolecular transfer of NO"^ to the thiol in the cysteine residue [545]. Subsequently it has been shown [546] that S-nitrosation of CysP93 in oxyhaemoglobin is virtually halted in the presence of the copper(I) chelator neocuproine. This indicates strongly that NO is generated from GSNO by the now well-known copper-catalysed process, Eq. (356). It is believed that GSNO + Cu(I) ^^P^

GS- + NO + Cu(II)

(356)

the autoxidation of NO at these low concentrations is too slow to allow N2O3 formation and subsequent nitrosation, so it has been suggested [546] that NO generates HbSNO in a Cu(II)-catalysed reaction i.e. the reverse of Eq. (356). This is not a well-known reaction, but there is one interesting report [547] which shows that the presence of Cu(II) brings about rapid S-nitrosation of both bovine and serum albumin. It is possible that this occurs via a copper-NO complex or even by the oxidation of NO to nitrogen oxidation state III by Cu(II). This is an area which would indeed profit from further mechanistic study. It has further been shown that the formation of the iron nitrosyl from deoxyhaemoglobin, which is not a vasodilator, by GSNO also requires the presence of copper [548]. Again, this argues strongly in favour of the intermediacy of NO, generated by the copper-catalysed decomposition of GSNO. Although the hypothesis regarding involvement of HbSNO is in many ways an attractive one, it is by no means widely accepted, and it is currently the subject of fierce debate. One group has measured the HbSNO levels in human arterial red cells at around 50 nM [549], which is considered to be much too low to act as a reservoir of NO. Further, HbSNO is unstable in the reductive environment of red cells. Another group concludes that NO is consumed by reaction with oxyhaemoglobin, rather than conserved [550]. It has been claimed that there is a large amount of experimental evidence, much of it based on comparative rates of reaction and on measured levels of HbSNO in biological systems, which argues against the involvement of HbSNO as a carrier of NO in vivo. Rather, evidence points to a finely balanced and

178


regulated position, between the generation of NO biosynthetically and its destruction (to nitrate ion) by reaction with oxyhaemoglobin [551]. There is a review article [552] which discusses the strengths and weaknesses of the various theories which have been advanced in recent times. Clearly much more work is needed in this extremely complicated area, before a definitive position is achieved. 10.4. Reaction of nitric oxide with thiolate ion Nitric oxide in the presence of oxygen generates a nitrosating agent which will react with thiols generating S-nitrosothiols. This reaction has already been discussed, not only for thiols but for a range of nucleophilic species. However if oxygen is rigorously excluded from the system another reaction occurs, which does not give S-nitrosothiols, but rather the disulfide derivative and nitrous oxide. This reaction occurs at relatively high pH and probably therefore involves the thiolate anion [553]. The sequence of reactions in Eq. (357-361) is a reasonable account of the probable mechanism. The rate limiting step is believed to be the first step involving the reaction of NO with RS" + NO — •

RSNO*

(357)

RSNO* + H3O+ ^^^=^ RSNOH* + H2O

(358)

2RSN0H* — ^

(359)

RSN(OH)-N(OH)SR

RSN(OH)-N(OH)SR HON=NOH — •

—•

RSSR + HON=NOH

N2O + H2O

(360) (361)

thiolate ion. An alternative pathway has been suggested [554] involving the formation of a sulfonic acid which reacts with a thiol to give the disulfide Eq. (362-3). RSNOH* + NO — • RSOH + RSH — •

RSOH + N2O RSSR + H2O

(362) (363)

The results of kinetic studies have been reported by three independent groups, one measured the rate of N2O appearance [555], another the rate of glutathione disappearance [556] and a third followed the disappearance of nitric oxide using a chemiluminescence method [557]. The first and third groups are in general agreement whilst the second reports a rate constant which is larger


179

by about 100. On balance the concensus of opinion is that the reaction of nitric oxide with thiolate ion, at least for simple low molecular weight thiols such as glutathione or cysteine, is too slow to be important in vivo, but the area awaits the results of more appropriate model thiols. 10.5. Analytical methods for the determination of nitric oxide 10.5.1. Colorimetric methods One of the most widely used assays for nitric oxide relies upon its oxidation and conversion to nitrite ion. When acidified (if necessary) the solution will nitrosate/diazotise amines; in the case of aromatic primary amines this generates the diazonium ion which when coupled with an activated aromatic compound gives an azo dye with a large extinction coefficient. This is the basis of the Griess test [162]. Nitric oxide in aerated water is known to give nitrite ion quantitatively. There are a number of variants which have been used as diazotisation and coupling reagents. Sulfanilamide, sulfanilic acid and 4-methylaniline have been used successfully as the aromatic primary amines and 1-naphthylamine, 2-naphthol, N-(l-naphthyl)-ethylenediamine and 2naphthol-3,6-di-sulfonic acid are well-known coupling agents. The reactions are shown in in Eq. (364-5), for the diazotisation of sulfanilamide and coupling with N-(l-naphthyl)-ethylenediamine. The X^^^^ for the azo dye is in NH2 (364) SO2NH2 NHCH2CH2NH2

NHCH2CH2NH2

SO2NH2

the region of 540 nm and extinction coefficients are --53,000 M""^ cm~^ The azo dye is quite stable in solution once formed, so that decomposition is not a problem here which would lead to an error. Nitric oxide at concentrations as low as -- 5 \xM can be analysed in this way. The method is of course based on

180


a sampling procedure, and is often not the preferred option, when direct measurement methods are available. Another colorimetric method is based on the oxidation of ferrocyanide by the intermediate from the autoxidation reaction of NO, probably N2O3, Eq. (366). This reaction results in an increase in the absorbance at 420 nm, and a [Fe(III)(CN)6] 3-

4[Fe(II)(CN)^]^+ NO/O2

(366)

calibration graph can readily be prepared. This sampling procedure has been used with nitric oxide concentrations as low as 25 fxM. The change in extinction coefficient is only -1000 M"^ cm~^ and so it is not a particularly sensitive method. A third colorimetric assay depends on the oxidation of ABTS 2,2'(azinobis(3-ethylbenzthiazoline-6-sulfonic acid), which is colourless, to the green ABTS^ radical cation, Eq. (367), which is stable in solution for 24

"Xcw^xr"'" CH2CH3

CH2CH3

NO/O2

•03S^^-^S

CH2CH3

ABTS

(367)

S^,,^x/S03-

r

CH2CH3

hours. The extinction coefficient of ABTS"^ at 600 nm is -12,000 M~^ cm"^ The method can be used to determine nitric oxide down to - 2 |LiM. All three methods can be useful in laboratory experiments within their limits of measurement, but are not feasible methods for the determination of nitric oxide in biological fluids, where the concentrations are much lower and a direct measurement procedure (rather than one relying on a sampling procedure) is called for. 10.5.2. Electrochemical methods There has been much effort directed at the electrochemical determination of nitric oxide, (a) since it will be a direct method not relying on sampling


181

procedures and (b) because the method has the potential to analyse rather low nitric oxide concentrations, such as those found in biological systems in the nanomolar range. There has also been much development work to enable very small electrode systems to be used in biological cells and other situations. There has been a large measure of success and commercial kits are available for work in the laboratory at relatively high nitric oxide concentrations and also in biological systems for very low concentrations. Basically there are two current procedures which have been used successfully, (a) the classical Clark probe (originally designed for the determination of oxygen in solution), where electrochemical oxidation of nitric oxide takes place on a platinum electrode (the anode), Eq. (368), with a silver NO + 2H2O — •

NO3- + 4H+ + 3e-

(368)

wire as the counter-electrode (the cathode), and the system is run in the amperometric mode. A constant potential of 0.9 V is applied and the current is measured, which for a limited range is linear with nitric oxide concentration. Calibration graphs are easily generated by the formation of nitric oxide in solution by a range of suitable reactions e.g. the reduction of nitrite/nitrous acid by iodide ion or ascorbate etc. The whole procedure is carried out in the complete absence of oxygen. The detection limit is of the order of 10~^ M and the response time a little slow at 1-3 s. The Clark electrode was first used by Shibuki and Okada [558] for the measurement of nitric oxide in brain slices. Modifications to the working electrode include the use of a lacquer-coated Pt/Ir electrode and membrane- and polymer-coated carbon fibre electrodes. Malinski and Taha [559] developed an alternative electrode based on the electrochemical oxidation of nitric oxide on a polymeric porphyrin. This is said to have significant advantages over the Clark electrode, in that response times are shorter, the detection limit is a little lower and a wider range of linearity exists between the current and the nitric oxide concentration. It does not appear at this stage to be available commercially and constructing the electrode is a skilled procedure. For the measurement of reasonably high nitric oxide concentrations (10~-^-10~^ M) in a laboratory situation, such as measurement of nitric oxide released from a nitric oxide donor, the Clark electrode is probably the better choice, whereas for the determination of very low nitric oxide concentrations particularly in small-scale biological situations, the use of the polymeric porphyrin electrode of Malinski is the better option. More detail of both electrode systems, including their relative strengths and deficiencies are given in references [560] and [561].

182


10.5.3. The oxyhaemoglobin method This method is based on the reaction already encountered between oxyhaemoglobin and nitric oxide yielding methaemoglobin and nitrate ion, Eq. (369). There is a change in the UV/vis spectrum from a A^j^^x ^t 415 nm with oxyHb + NO — •

metHb + NO3-

(369)

e = 131,000 M-^ cm"^ (oxyHb) to a X^^ at 406 nm with e = 162,000 M"^ cm~^ (metHb). Since the change is not large, better results can be obtained using the difference spectrum method rather than using the absolute spectra, i.e. by placing oxyHb at the same concentration in both the sample and reference cuvettes in the spectrophotometer. There is also a procedure using two wavelengths simultaneously and measuring the difference spectra between the two wavelengths. The method is discussed fully, including the difficulties and limitations of the method in reference [562]. The relatively fast reaction of nitric oxide with MetHb however (giving the iron nitrosyl) may lead to sources of error in this method. 10.5.4. Chemiluminescence method This is essentially the same method as has been applied for the determination of nitrosamines (mentioned in section 3.2, p.68), which are converted to nitric oxide by a number of possible methods. Reaction with ozone, Eq. (370), then leads to the rapid formation of nitrogen dioxide in an NO + O3 — • NO2* — ^

NO2* + O2

NO2 (ground state) + hv

(370) (371)

excited state, which returns to the ground state, Eq. (371), emitting light as a chemiluminescence. This occurs in the red and infrared part of the spectrum (-640-3000 nm) with the peak at about 1100 nm. The signal is amplified and recorded. The signal (mV) is usually linearly related to the nitric oxide concentration. Calibration is normally carried out with standard nitric oxide samples in nitrogen, which are available commercially. Initially the method was used for the determination of nitric oxide in the atmosphere with reference to the pollution aspect, but is now much used because of its sensitivity for the determination of nitric oxide associated with its production in biological systems. For analysis of nitric oxide released into the gas phase from, for example, endothelial cells or nitric oxide donors or by photochemical decomposition, this method is appropriate. It is also used in hospitals to monitor nitric oxide gas (at low concentration) given to new bom infants


183

suffering from pulmonary hypertension and to adults for the treatment of some chronic respiratory disorders. Nitric oxide dissolved in a liquid phase can be displaced into the head space by a stream of inert gas and this can be transferred directly to the ozoniser and analyser. Concentrations as low as 10~^^ M can be detected using this method. Practical details of the method, including possible error sources and how they are overcome are given in reference [560] pp.83-92 and in reference [562] pp.309-318. 10.5.5. Spin-trapping methods Although nitric oxide has an unpaired electron and is therefore, paramagnetic, it cannot in dilute aqueous solution be detected by the Electron Paramagnetic Resonance (EPR) spectroscopy method, as a result of the coupling of the spin and orbital angular momentum. However a number of species react with nitric oxide to give adducts/complexes which are EPR-active and have been widely used to detect and quantify nitric oxide. Complexes with haem proteins. Both deoxyhaemoglobin and deoxymyoglobin form stable complexes with nitric oxide, Eq. (372), and the EPR spectrum was HbFe(II) + NO

—•

HbFe(II)NO

(372)

reported as early as 1955 [563]. The characteristics of the broad EPR spectrum (g values, hyperfme splitting and linewidth) have been used to detect and quantify nitric oxide widely since that time. The spectrum is quite complex at room temperature, but is much simplified at 77 K in the frozen state, and is generally characterised by three absorptions. Nitrosylmyoglobin MbNO behaves in a similar fashion. Significant advantages of using haemoglobin in this way is that it has a very high affinity for nitric oxide (equilibrium constant ~3 X 10^^ M~^) and the spectrum is very characteristic. One disadvantage is that, particularly in an oxygen-rich environment there is a tendency for Hb to form oxyHb which readily oxidises nitric oxide to nitrate ion. Other nitrosylhaem complexes which have been studied by EPR in this regard include nitrosyl cytochrome P450, nitrosyl cytochrome P420, nitrosyl catalase and nitrosyl guanylate cyclase. All give spectra at 77 K which are readily interpreted. Iron-nitrosyl dithiol complexes. Nitric oxide will generally bind Fe^"^ to form an iron-nitrosyl complex. Some iron-thiol complexes have been used to determine nitric oxide by EPR. The most well-known is the complex with diethylthiocarbamate (DETC), Eq. (373-4). The nitrosyl complex is relatively


184 C2H5 \

/

C2H5

C2H5 \

s+ Fe-2+

N

/ C2H5

S

,/K \

1

^S

C2H5

DETC (373) NO C2H5 \

DETC + NO

K K

/ N—I

/ C2H5

^

,C2H5 (374)

> - < C2H5

stable even in the presence of oxygen, and gives a distinct three-line EPR spectrum both at room temperature and at 100 K. A more water soluble ligand is the derivative N-methyl-D-glucamine dithiocarbamate (MGD), which is

\

rPH3

S

CH2—(CHOH)4-CH20H MGD

easily synthesised. This forms an iron-nitrosyl complex, which is also quite stable and water soluble, Eq. (375), again with a characteristic EPR spectrum. 2MGD + Fe2+ + NO

(375)

Fe2+(NO)(MGD).

A related ligand is 2,3-dimercapto-l-propanesulfonic acid (DMPS) which also complexes Fe^"*" and then binds nitric oxide, Eq. (376). The NO SO3 'O3S

S

S

+ NO

S03" O3S

s s

(376)

nitrosyl complex has been characterised. The Fe^''"(DMPS)2 compound was first used in industrial situations as a scrubber of flue gases to remove nitric oxide (a pollutant), because of its strong affinity for nitric oxide [564]. The


185

iron complex is coloured red with absorbance maxima at 358 and 509 nm and is stable in solution under anaerobic conditions. The nitrosyl complex is colourless in solution so it is easy to follow the reaction in Eq. (376) readily by spectrophotometry. In general Fe^'^(NO)(dithiol)2 complexes are also easy to detect by EPR and have been used to measure nitric oxide release from NOdonors and also in vivo. Other spin traps. Two other organic spin-traps have been developed recently, both of which have resulted in EPR changes which are measurable when reaction with nitric oxide has occurred. The first is a biradical generated by photolytic decomposition of a 2-indanone during which carbon monoxide is

hv

r

T-

-

NO

N-0

Ô (377) eliminated , Eq. (377), [565]. The second is a nitronyl nitroxide which gives an imino nitroxide on reaction with nitric oxide, Eq. (378), [566]. There is a

V-R

JÔ^

^

3

V-R

+ NO2

(378)

O' dramatic change in the EPR spectrum on reaction, which can readily be analysed. A range of different nitronyl nitroxides have been synthesised and used in this way - including R = N"^(CH3)3, CH2C(CH3)2CH20H and C^H4C02Na. Nitronyl nitroxides have been used to monitor nitric oxide release from NO-donors and to probe the mode of action of vasodilators. Continuous monitoring is possible as is the possibility of use in biological membranes. A relatively new application for a spin-trap, orginally designed as a trap for C-centred radicals has involved the use 3,5-dibromo-4-nitrosobenzene sulfonate (DBNBS). One reported application [567] has been an investigation of NO participation in the human platelet system.

186


10.5.6. Fluorescence spectroscopy Most methods based on measurement of a fluorescence have used the reaction of 2,3-diaminonaphthalene (DAN) with nitrous acid and measurement of the fluorescence of the resulting 2,3-naphthotriazole (NAT). This procedure has already been outlined in section 8.6.4., p. 158, when dealing with the analysis of S-nitrosothiols. The method was originally designed for the analysis of nitrite ion and nitrate ion. The later has to be reduced quantitatively to nitrite ion, before acidification and diazotisation of DAN. Nitric oxide in the presence of even traces of oxygen will effect diazotisation quantitatively. DAN itself has a very low fluorescence whereas NAT is highly so. The excitation wavelength is usually 375 nm and the emission measured at either 415 or 450 nm. Fluorescence is linearly related to the triazole concentration and it is claimed that nitric oxide concentrations as low as -10 nM can be determined. Other techniques have been applied to the determination of nitric oxide concentrations, but are less frequently encountered. They include, resonance Raman spectroscopy, gas chromatography and mass spectrometry. For full details of all reported methods see references [568-9] and also [570] where a whole issue of the journal is devoted to this subject. A number of articles concentrate on the question of which method is appropriate for a given situation.

Nitrosation Reactions and the Chemistry of Nitric Oxide D.L.H. Wilhams © 2004 Elsevier B.V. All rights reserved.

187

Chapter 11

Nitric oxide in biological systems 11.1. Biological properties of nitric oxide Since the identification in 1987 of nitric oxide as the endotheliumderived relaxing factor (EDRF), and its role in vasodilation, there has been an enormous research effort directed at establishing the mode of action and its biosynthesis. The literature in this area has expanded at such a rate that it is extremely difficult to keep track of events. In the years following 1987 it was also found that nitric oxide is a key factor in a whole range of other physiological processes, including inhibition of platelet aggregation, neurotransmission, cytotoxicity and penile erection. It is also believed that nitric oxide is involved in the brain in the processes of memory and learning, as well as being involved in a whole range of human diseases relating either to a deficiency or overproduction of nitric oxide. Further discoveries continue to emerge. Such is the interest that has built up in this relatively short time, that nitric oxide was nominated as 'Molecule of the Year' by Science magazine in 1993 [571]. Previously, nitric oxide had principally been regarded as a major contributor to atmospheric pollution, deriving from its synthesis from nitrogen and oxygen at high temperatures, Eq. (379), and subsequent oxidation to N2(g) + O^Cg) = 2NO(g)

(379)

nitrogen dioxide. The chemistry here has been much studied, including that involved in the development of catalytic converters, where nitric oxide is reduced back to nitrogen by carbon monoxide, Eq. (380), in the presence of NO(g) + CO(g) = N2(g) + C02(g)

(380)

platinum, palladium and rhodium catalysts. The "bad" (i.e. polluting) properties and the "good" (i.e. biological) properties have recently been brought together in a monograph [572]. Whilst there is still much to be learnt regarding many aspects contributing to the biological properties of nitric oxide, quite a lot of facts have been established and there is some understanding of the mechanisms involved. Chemists perhaps do not fully appreciate some of the difficulties here, when dealing with nitric oxide at the nanomolar level and also with the extremely large and complex structures of the enzymes involved. It is now reasonably certain that nitric oxide "activates" (i.e. reacts with) the enzyme guanylate cyclase (sometimes called guanylyl cyclase) which results in the relaxation of the so-called smooth muscles found principally in blood

188


vessels. This occurs via a series of biochemical reactions resulting (after nitric oxide activation) in an increased concentration of cyclic guanosine monophosphate cGMP (35) which is derived from guanosine triphosphate GTP

HN H2N

N^"¥

^

cGMP

^

OH 35

1

GTP

•0-P-0-P-0-P-a-H20 O-

O'

O" HO' 36

(36). The increased cGMP level decreases the intracellular Ca^^ which is required for contraction of smooth muscle. This allows the muscle to relax and consequently the blood vessel to dilate, thus lowering the blood pressure. The soluble form of guanylate cyclase occurs in most cells, and although it has been isolated, a full structure has not yet been obtained. It is however a ferrohaem in which the iron atom is bound to a histidine grouping and also to the four nitrogen atoms of the porphyrin system in the haem in a pentacoordinate arrangement. When reaction with nitric oxide occurs there is evidence which suggests that the histidine grouping is lost, retaining a pentacoordinate arrangement for the iron as outlined in Eq. (381). Evidence

Nitric Oxide in Biological Systems

NO

Hist-105

189

(381)

Hist-105

for this "activation" process comes from a variety of spectral and kinetic measurements made on model compounds. The Fe-histidine stretching frequency in guanylate cyclase measured by resonance Raman spectroscopy, is very low, indicating that the Fe-N bond is very weak and so prone to displacement on reaction with nitric oxide. The iron atom is now above the plane of the porphyrin ring allowing bonding with GTP in the cavity created, bringing about reaction to give cGMP. The ability of nitric oxide to inhibit platelet aggregation is also thought to involve guanylate cyclase and the formation of cGMP. Nitric oxide is generated in the endothelial cells and can diffuse into the blood stream and into the platelets. Guanylate cyclase and cGMP (following activation by nitric oxide) are involved in a series of reactions which results in a reduction in the affinity of platelets for each other and also for the endothelial surface, thus reducing the clotting of blood. Some nitric oxide donors, notably sodium nitroprusside, have been used routinely for this purpose during medical operations, to reduce the clotting problem during surgery. Nitric oxide also plays a large part in the immune system. Macrophages are cells which are part of the immune system and protect against various microbial attacks. In the 1980s it was found that both nitrite and nitrate ions were found in the supernatant liquid (in laboratory experiments) and in the urine (in animal experiments), after bacterial infections had occurred. This led to the conclusion that nitric oxide is generated in the macrophages (by the same pathway as in the vascular system) which acts to kill bacterial and also tumour cells. Nitric oxide in relatively high concentrations is toxic and destroys various cellular structures. It is likely that the toxicity does not arise directly from nitric oxide but from some derived product, possibly peroxynitrite, ironsulfiir clusters or iron-haem compounds. Much remains obscure in this area, but it has been suggested that the toxicity of nitric oxide arises by its attack on iron atoms, which are involved in cell respiration and DNA synthesis. It is believed that nitrite and nitrate anions

190


arise by oxidation of nitric oxide and hydrolysis. This reaction in the laboratory generates exclusively nitrite, but in biological systems it is conceivable that some is oxidised further to give nitrate. What is known however is that nitric oxide generated in the immune system is synthesised by the same route as nitric oxide generation in the endothelium i.e. by the enzymatically catalysed oxidation of arginine. Nitric oxide is known to produce mutations in bacterial and mammalian cells. Experiments carried out under aerobic conditions at pH 7 clearly demonstrated mutagenicity. One explanation given depends on the oxidation of nitric oxide to give N2O3 which effects nitrosation/diazotisation of amine groups, resulting in base mis-pairing and hence causing mutation to occur. The detailed chemistry has not, however, been fully established. A detailed account of the role of nitric oxide in the immune system is given in reference [573] and in references therein. Perhaps the most surprising area where nitric oxide is known to play a pivotal role is in the nervous system where it acts as a neurotransmitter. One of the many chemical processes taking place in the brain results in cGMP formation and that this requires the presence of arginine. Soon after these discoveries the presence of nitric oxide in brain tissue was confirmed. It is now known that nitric oxide is one of a number of identified messengers in the central and peripheral nervous system. A detailed account of the biology here is beyond the scope of this book and the reader is referred to many review articles on the subject in the biological literature e.g. [574]. Less of a surprise is to find that nitric oxide plays an important part in the process of penile erection, given its role already described in vasodilation. Erection results from a relaxation of the smooth muscle and consequent dilation of the blood vessels in the corpus cavemosum in the penis. This can be activated by a number of stimuli activated in the brain and passed through the peripheral nervous system using neurotransmitters which include nitric oxide. Dilation of the blood vessels in the corpus cavemosum occurs by elevated cGMP levels generated as before, by activation of the enzyme guanylate cyclase by nitric oxide. Thus any problems leading to a reduction of nitric oxide production can be a cause of impotence or male erectile disfunction (MED). An apparently simple remedy would be the application of nitric oxide donors such as glyceryl trinitrate, GTN, (see later in chapter 12). Indeed local application of GTN does have the predicted effect. However there are problems arising from this approach and to date, in spite of much research activity there is no nitric oxide donor on the market aimed at treating impotence. However another approach has been much more successful. The process of dilation of the blood vessels is quickly reversed by the spontaneous


191

hydrolysis of cGMP to give guanosine monophosphate (GMP) as shown below:NO (from arginine)

Guanylate cyclase

GTP GMP

cGMP

Dilation of bloodvessels

Phosphodiesterase

The hydrolysis of cGMP to give GMP is catalysed by a phosphodiesterase, and if this reaction can be inhibited then loss of cGMP can be much reduced. This is the principle of the mode of action of the wellknown drug Viagra, prescribed for the treatment of MED. The structure of Viagra (sildenafil) is given as 37.

C2H5O HN

37 Curiously, sildenafil was originally designed as a drug for the treatment of high blood pressure, for which it is not very effective, and the effect on penile erection was discovered accidentally during clinical trials. More recently two other drugs have been introduced in the market which are more effective than Viagra, which work on the same principle i.e. they act as phosphodiesterase enzyme inhibitors. There are a number of reviews of the role of nitric oxide in penile erection in the biological literature, for example reference [575] which discusses the pharmacology involved.

192


Nitric oxide continues to be implicated in a large number of other areas, although in many cases there is little understanding to date of the detailed pathways involved. There is for example an abnormally high level of nitrite ion in the synovial fluid of joints of sufferers of rheumatoid arthritis, which suggests an overproduction here, and perhaps in other inflammatory conditions, of nitric oxide. Recent developments indicate that nitric oxide is involved in wound healing, and also in the formation of melanin in the skin following exposure to UV light. Overproduction of nitric oxide is also a possible contributory factor in patients suffering from Parkinson's disease, and may also contribute to brain damage in stroke patients. The serious condition of sceptic shock which results in a severe reduction of blood pressure, often fatal, could be the result of massive overproduction of nitric oxide by macrophages. This can in some cases be reversed by the rapid administration of enzyme inhibitors of the biosynthesis of nitric oxide. New results have shown that increased nitric oxide levels are found in patients with urinary disorders and multiple sclerosis and there seems to be a link between schizophrenia and a malfianction of the biosynthesis of nitric oxide. Other medical disorders continue to be reported which could be the result of deficiency or overproduction of this amazing molecule. 11.2. Biosynthesis of nitric oxide In the laboratory or in industrial processes, nitric oxide is very conveniently prepared by oxidation of ammonia (catalytically) or reduction of inorganic species in oxidation states of nitrogen IV and V, usually nitrite or nitrate salts, with a range of conventional reducing agents. Nature has chosen a totally different route by the enzymatic oxidation of the amino acid L-arginine in the presence of oxygen. Biosynthesis occurs under very mild conditions of neutral pH and body temperature. The final organic product is L-citrulline which returns to L-arginine as part of the urea cycle. Apart from the enzyme(s) nitric oxide synthase (NOS), co-factors are the reduced form of nicotinamide adenine dinucleotide phosphate (NADPH), tetrahydrobiopterin, calcium ions and haem. Reaction is a two-stage process, first forming Nhydroxyarginine as an intermediate in a two-electron transfer and then conversion to L-citrulline and nitric oxide in a three-electron transfer as laid out schematically in the scheme overleaf:-


H2N^NH2

H2N^NOH NOS NADPH

H3N

193

CO2" O2

L-arginine

NOS NADPH

.NH

H2O H3N

CO2

+ NO

H2O H3N

O2

X02~

L-citrulline

Urea cycle

The overall five-electron transfer oxidation has no simple equivalent in non-enzymatic chemistry. The history of the discovery of these reactions has been laid out elsewhere [576]. Labelling of the terminal guanidine nitrogen atoms shows that the nitrogen atom in nitric oxide derives from this position, and the oxygen atom derives from molecular oxygen, as shown by oxygen ^Ô labelling experiments. The intermediate has been isolated and identified, although it is believed to remain bound to the enzyme in the biosynthesis. It is relatively easily synthesised by a number of routes and is also now available commercially. N-Hydroxylation significantly decreases the pK^ value of the guanidine nitrogen (from 13.6 to 8.1). The N-hydroxyarginine is unstable in base, but is stable in acid solution as the protonated form. Oxidation can occur as a two-electron transfer resulting in HNO elimination which rapidly forms nitrous oxide N2O. A one-electron transfer oxidation has also been identified leading to radical species which have not been fully characterised. As in other NADPH oxidations, it acts as a source of electrons which generates an active oxidising agent from molecular oxygen which reacts according to Eq. (382). The second oxidation uses half a mole of NADPH, Eq. (383). H9N ,NH 2'^^^'^^

H2N..N-OH

T

T

. 02 + H* + NADPH —

+ NADP

+ H2O (382)

194


H2N^^N-0H T /NH

+ O2 + 0.5 H^ + 0.5 NADPH (383)

H2N^0 + NO + 0.5 NADP+ + H2O

As yet the role of tetrahydrobiopterin (BH^), 38, is not clear, although H

i'/x

HN H2N

N^ ^N" H

38 it is known that its presence is necessary in all of the NOS synthase pathways. BH4 is a well-known co-factor for a number of monooxygenases and its biochemistry here has been much studied; it seems to act as an activator of molecular oxygen by the formation of a 4-hydroperoxy derivative. The position regarding its involvement in nitric oxide biosynthesis is discussed fully in reference [576]. 11.3. NO synthase enzymes Three forms of NOS have been isolated and partially identified. These are called isoforms of NOS. They are, endothelial NOS (eNOS), inducible NOS (iNOS) and neuronal NOS (nNOS). All act on the substrate L-arginine to give L-citruUine and nitric oxide. Reactions are stereospecific and D-arginine is not a substrate for NO biosynthesis. The three isoforms are different in that they do not contain the same amino acid sequence in one of the active sites of the enzymes. eNOS as the name suggests is found in the endothelial cells of blood vessels and generates EDRF(NO) which results in relaxation of the smooth muscle thus increasing blood flow. iNOS develops in the macrophages in response to stimuli associated with infection. It is not normally present in healthy cells. nNOS is active in the neurons of the nervous system and also is present in the brain. It was the first NOS to be isolated and purified. All three forms catalyse both steps of the oxidation of L-arginine.


195

There is a haem function in all NOS forms which is believed to play a central role in the biosynthesis. A reasonable scenario is as follows:- Larginine becomes bound near the haem site, an electron from NADPH is transferred to the haem, allowing molecular oxygen to bind to the iron atom. A second electron transfer to the haem allows 0 - 0 bond breaking, one oxygen atom ending up as water and the other remaining bound, which becomes inserted into a N-H bond in L-arginine generating N-hydroxy-L-arginine bound to the enzyme and reforming the haem in the normal ferric state. The second step - the oxidation of the intermediate hydroxy compound - again is believed to involve binding of oxygen to a haem following an electron transfer followed by a series of somewhat speculative reactions involving a peroxideiron species and the N-hydroxy-L-arginine cation radical resulting finally in the formation of L-citrulline, nitric oxide, water and ferric haem [573]. Indirect evidence that a haem function plays a crucial part in the biosynthesis comes from the observation that if the iron site is coordinated with carbon monoxide then the ability to generate nitric oxide is lost. There has been much interest in the development of enzyme inhibitors for NOS. This stems principally from the quest to be able to stop overproduction of NO in sceptic shock and also in a range of diseases where it is believed that excess NO contributes to, or is responsible for, the condition e.g. as in rheumatoid arthritis. The first indication of enzyme inhibition occurred even before NO was identified as EDRF etc., when Hibbs and coworkers [577] showed that N-methyl-L-arginine, 39, reduced the cytotoxic effects of an L-arginine-promoted response in macrophages. Since then a large range of other A^-substituted L-arginine derivatives, including the N-nitro, 40, N-amino, 41, and N,N-dimethyl, 42, derivatives among others. In addition a C H 3 ( H ) N ^ . N H N02(H)N,^.NH NH2(H)N^^NH

(CH3)2Nv,^NH

^NH

^NH

^NH

/NH

39

40

41

42

large number of other related derivatives including L-thiocitruUine, 43, and its S-methyl derivative (one of the most potent inhibitors), aminoguanidine, 44, and methylguanidine, 45. N,N-Dimethylarginine, 42, is itself naturally occurring and so can adversely affect nitric oxide production. There is however an enzyme also present in vivo, which converts 42 to L-citrulline

196


NH NH2

CH3

H3N" ^C02" 43

44

45

and dimethylamine. It is possible however that 42 could act under conditions of NO overproduction as a self-regulating procedure. In some cases of patients suffering from mild coronary heart disease, administration of L-arginine over some weeks has proved beneficial, since it can now compete more favourably with the inhibitor. A number of these NOS enzyme inhibitors have been, and continue to be used medically, particularly in the treatment of sceptic shock. There are a number of naturally occurring materials in foods for which claims have been made of their efficacy in cardiovascular health generally, although the scientific evidence is not yet totally convincing. One such food is garlic, where there are claims that blood pressure is lowered, platelet aggregation inhibited and that it has antibacterial and antifungal properties. The components of garlic believed to activate NOS in this regard are allicin, 46, and agoene, 47. Both have shown to inhibit platelet aggregation and to mirror the effects of nitric oxide itself, although how this occurs is a matter for speculation at this stage.

46

47

Currently the main thrust of research in this area is the search for NOS enzyme inhibitors which are specific to only one of the isoforms. Generally speaking most of the inhibitors are effective against all forms with only a marginal specificity for one or another. A specific inhibitor would enable the treatment of particular diseases/infections, without affecting other functions. Some success has been reported [578]. The most selective nNOS inhibitor available is the structure 48, called ARL 17477. This resulted in a virtually


197

48 complete inhibition of NOS activity in the brain cortex without affecting the mean arterial blood pressure. 11.4. Nitric oxide in the treatment of foods Both nitrite and nitrate ion, usually added as the sodium or potassium salts, have been used in aspects of food preservation for a long time. There are two reasons for this, (a) to protect, particularly all cured meat products, from bacterial growths especially Clostridium boutulinum, and (b) to produce an attractive red colour which is imparted to cured meats. Originally nitrate salts were used, but it was found that nitrate undergoes a bacterial/enzymatic reduction to nitrite, which is the effective reagent, and so nitrate salts were gradually replaced by nitrite salts. Nitrite (via nitrous acid) is readily reduced to nitric oxide by a variety of standard reducing agents such as ascorbate, in the laboratory, and is also reduced to nitric oxide in biological systems. It is thus difficult to establish whether the antibacterial properties arise directly from nitrite ion in from nitric oxide. Nitric oxide can be trapped as the nitrosyl complex in vitro by the addition of myoglobin in the reduced Fe(II) state, Eq. (384), and the same reaction occurs in vivo. Myoglobin is a ferrohaem, which MbFe" + NO = MbFe^NO

(384)

has been traditionally regarded as a store for oxygen in the muscle tissue, since it can (like haemoglobin) reversibly bind oxygen. The nitrosyl is the material responsible for the red/pink colouration in cured meats. Haldane in 1901, [579], was the first to prepare nitrosylhaemoglobin by the addition of nitrite to haemoglobin. MbFe^TÔ has been much studied, since it is a good model for the nitrosyl derivative of guanylate cyclase and hence gives an insight into the mode of action of biological NO signalling. A major health concern arises concerning the addition of sodium nitrite to cured meats, following the discovery of the carcinogenicity of nitrosamines. Nitrite ion in the acid environment of the stomach exists in its protonated form and will readily generate nitrosamines from any secondary amines present in foods or added drugs. Later, it has been shown that a greater risk comes from other nitrite sources e.g. in vegetables and principally in the water supply containing increased levels of nitrate ion, which undergoes rapid bacterial

198


reduction to nitrite ion in the saliva. Nevertheless, permitted levels of sodium nitrite in cured meat products has been reduced, and in many cases ascorbic acid is also added which reduces nitrous acid rapidly to give nitric oxide, Eq. (385). This has the effect of reducing the nitrosamine levels, a Tocopherol 2HNO2 "^ ascorbate = 2NO + dehydroascorbate + H2O

(385)

(Vitamin E) has the same effect. Current food regulations in the U.S.A. require the reduction of the nitrite level added in bacon manufacture to 120 ppm (from a previous level of 200 ppm), and the incorporation of 500 ppm ascorbate in the brine solutions used in the curing process. Interestingly, previous commercial practice added ascorbic acid in the curing process for another purpose, to increase the rate of nitric oxide production, and so ensure a more rapid development of the red/pink colour. Experiments have been carried out using the direct application of nitric oxide gas to meats. There were no differences in the colour development when compared to the nitrite-treated meat. This procedure has not been approved for commercial use.

Nitrosation Reactions and the Chemistry of Nitric Oxide D.L.H. WiUiams © 2004 Elsevier B.V. All rights reserved.

199

Chapter 12

Nitric oxide releasing compounds (NO donors) There has been intense and wide ranging research interest in the quest for chemical compounds which will liberate nitric oxide under fairly mild conditions. This work has been aimed at (a) the development of drugs which could be applied medically to make up for an inefficient biosynthesis of nitric oxide in vivo, resulting from disease etc., (b) again in the medical field, to be able under certain circumstances, e.g. in operations, to apply NO-donors which will produce some of the properties of nitric oxide tailored to a specific need and (c) being able to generate nitric oxide in a controlled manner in laboratory and animal experiments. One of the major problems here in the understanding of the mechanism of NO release is that many, but not all of the donors can generate NO in vitro in non-enzymatic reactions, whereas in vivo there is the probability that most, if not all of these species will generate NO by an enzyme-catalysed process. So the fact that compound X will not generate NO in vitro is not necessarily a meaningful criterion for the possible use of X in vivo. This makes the search for suitable NO-donors more difficult. 12.1. Organic nitrates By far the most widely used donor of nitric oxide are the alkyl nitrates, particularly glyceryl trinitrate GTN, 49. On the face of it, GTN, a well-known high explosive is not an obvious choice for use as a NO-donor, particularly when it is not easy to see a plausible mechanistic pathway to NO generation. Nevertheless, before any understanding of mechanism and recognition of the role of nitric oxide, GTN was used clinically (since 1879) to treat the condition of angina pectoris. This continues to be one of the most effective and most prescribed drug for this purpose, producing almost immediate relief from the acute chest pains associated with the condition. It is currently within the range of the top 100 prescribed drugs for all purposes. A variety of delivery procedures have been used, including transdermal (patches attached to the skin), sublingual (tablets) and spray methods. GTN does have some drawbacks -ONO2 I

ONO2

'—ONO2 49

O2NO—1 rO2N0

9NO2

T

•—ONO2 50

0-,N

200

Nitrosation Reactions and the Chemistry of Nitric Oxide H(

ONO2

ÔN02 (CH2)2

52 in that it is a short-acting drug, and so has to be taken frequently, and more seriously, it does generate a tolerance to itself in a number of patients. It can also result in some severe headaches. Other alkyl nitrates have been tested and are on the market, including pentaerythrityl tetranitrate, 50, isosorbide dinitrate, 51, isosorbide mononitrate, 52, and the more complex nicorandile, 53. None has been used as widely as has GTN, even though 50 is claimed not to generate a tolerance problem. It is generally assumed that GTN produces an EDRF i.e. nitric oxide, in vivo which effects vasodilation and allows greater blood flow through the affected narrow arteries. It has now been demonstrated that indeed NO is generated from GTN in vivo. The situation in vitro however is not so clear cut. Some workers have claimed that NO is generated in the presence of thiols, using the oxyhaemoglobin method, whilst others, using the very sensitive NOspecific electrode system have failed to detect NO even in the presence of cysteine. The known chemistry of alkyl nitrates does not include references to breakdown to nitric oxide. Alkyl nitrates are readily synthesised by 0-nitration of the corresponding alcohols, although in the case of GTN, this has to be carried out at fairly low temperature, since explosions can occur on detonation at high temperatures. In general alkyl nitrates are stable in aqueous solution at physiological pH values, but do decompose in the presence of excess thiol to give the disulfide and nitrite ion [580], More studies have been carried out on the decomposition of GTN in vivo. Nitric oxide has been detected in the exhaled air, blood, the liver and in other organs of patients after the administration of GTN. It is certain that GTN breaks down enzymatically. Two enzyme systems have been proposed, although in neither case is the mechanism well established. One of the enzymes is an NADPH-dependent cytochrome P450 and the other, some ioszymes of the glutathione-S-transferase family. The site of the enzyme binding has not been established and both explanations have been criticised by other workers. Recently however Stamler and co-workers [581], have isolated an enzyme, mitochondrial aldehyde dehydrogenase, from the mitochondria of cells, which generates NO from GTN. It is claimed that this work represents a major advance in the understanding of the mode of action of a drug which has been so widely used for such a long time.

Nitric Oxide Releasing Compounds (NO donors)

201

Initially it was believed that the tolerance of patients to long-term use of GTN arose from a local depletion of thiol functions. This is now thought not to be the case and the latest results suggest that this tolerance arises on repeated dosage by the actual using up of the enzyme. GTN has also been widely used medically to treat patients after heart attacks and congestive heart failure. It also has had some use as an alternative to Viagra in the treatment of erectile disfunction, when it is applied directly, locally in the form of a spray. Recent developments have included the synthesis of the so-called hybrid drugs which include two active functional centres. A number of organic nitrates have been designed in which the nitrate function is built-in to other drugs such as aspirin. The basic idea is that incorporation of a NO-donor moiety might not only reduce or remove the toxicity of the drug but also provide an NO-donor which will contribute to the biological activity. One such hybrid nitrate-aspirin drug is given in structure 54, in which a second

ONO2 54 benzene ring has been brought into the molecule to which the nitrate group is attached. It was hoped that this approach would reduce or eliminate the gastric problems associated with the use of aspirin. Other nitrate (and nitrite)-hybrid drugs have been developed. At the present time all are undergoing trials and are not yet on the market. 12.2. Organic nitrites The antianginal effects of alkyl nitrites have also been known for over a hundred years, but are not currently prescribed for the purpose, since alkyl nitrates have proved to be more acceptable. In the early years however a number of alkyl nitrites, particularly iso-amyl nitrite, 55, were used clinically. ;CH(CH2)20NO H3C 55

202


Since alkyl nitrites are more volatile than alkyl nitrates, administration was usually in the vapour phase. However the use of alkyl nitrites produces some undesirable side-effects such as some severe headaches and dizziness (resulting from vasodilation). Iso-amyl nitrite however has been much used or abused as a recreational drug particularly in the gay community, and is a constituent of "poppers", again making use of the vasodilatory property, and increased blood flow to the brain. Alkyl nitrites are easily prepared from the corresponding alcohol and a nitrosating agent. The 0-NO bond dissociation energy is quite high (150-170 kJ mor^) and so they are thermally stable at room temperature. They are prone to nucleophilic attack at the nitrogen atom, e.g. by thiolate ion, Eq. (386), generating S-nitrosothiols which readily decompose to nitric oxide RONO + -SR' = RO- + R'SNO

(386)

RSNO + Cu+ = R'S- 4- NO + Cu^+

(387)

by the copper-catalysed route, Eq. (387). This provides a ready non-enzymatic route to nitric oxide from alkyl nitrites in the presence of thiols at physiological pH, since the pAT^ of most thiols is - 8. It has been shown that alkyl nitrites release nitric oxide in vivo and their antianginal effects are, as for alkyl nitrates, rapid and short lasting. Their actual vasorelaxing property is greater than that of alkyl nitrates, which may reflect the higher rate of NO formation, although there are no data on this point. There appears to be no tolerance problem with alkyl nitrites. Although there is this well understood non-enzymatic pathway to nitric oxide, it is probable that there is additionally an enzymatic pathway which takes over in vivo, but there is very little firm information on this point. 12.3. S-Nitrosothiols The chemistry of S-nitrosothiols RSNO has been thoroughly reviewed in Chapter 8. This includes a discussion of the formation of nitric oxide, thermally, photochemically and by the copper-catalysed pathway, for all of 2RSN0 = RSSR + 2N0

(388)

which the overall reaction is Eq. (388). At body temperature and in the absence of UV light the only realistic non-enzymatic route to yield nitric oxide is the copper-catalysed one. The active reagent is Cu"^, which is needed only in trace amounts, and this can readily be formed by reduction of Cu^"^ by thiolate and possibly other reducing agents such as ascorbate. It is not necessary for the


203

Cu^"^ to be free as the hydrated ion, since it can be accessed when it is bound to amino acids, peptides and proteins. A number of RSNO species have been detected in vivo, including Snitrosoglutathione, S-nitrosoalbumins and S-nitrosohaemoglobin. There is currently a belief that nitric oxide is transported around the body as RSNO species, principally the albumin derivatives, although the evidence for this is somewhat circumstantial, and derives principally from the observations that NO itself has a very short half-life in vivo, estimated variously in the region of 100 ms, whereas RSNOs are much more stable. In general, nitric oxide release from RSNOs is favoured by high copper concentrations (free or bound) and by the presence of reducing agents such as thiolate or ascorbate. Increasing the thiolate concentration can have a variety of effects however. In many cases this facilitates the formation of Cu"^, the active reagent, but in some cases increasing the thiolate concentration can complex Cu^"^ and thus make it more difficult to access Cu"^. This is very much a structure-specific effect, and the most well-known and pronounced example is the complexation of Cu^"^ by penicillamine. When the thiolate concentration is reduced to very low levels however and in the absence of any other reducing agent, nitric oxide release is much retarded, and can be reduced, in effect, to virtually nothing. Much use has been made of the use of Cu^"^ chelators, notably EDTA, to demonstrate that these reactions are Cu^~^promoted reactions. At low concentrations ascorbate behaves in the same way, acting as a reducing agent for Cu^"^, but at higher ascorbate concentrations, RSNOs act as electrophilic nitrosating agents, developing nitric oxide by a NO"^-transfer reaction which is totally independent of the presence or absence of Cu^"^. No other reducing agents have been examined in this way in vitro, but it is quite conceivable that others do exist in vivo and will promote nitric oxide release. There is much confusion in the literature regarding the relative "stability" of RSNOs. For example it is often quoted, even in the present day literature that "S-nitrosoglutathione is more stable than is S-nitrosocysteine". This is certainly true for the pure material since the former is indefinitely stable as the pure solid at room temperature, whereas it is extremely difficult to obtain a pure sample of the latter, except at very low temperatures. The statement is also true for solutions of those two RSNOs at concentrations in the millimolar range; this is due to complexation of the Cu^"^ by the disulfide product from S-nitrosoglutathione which does not occur for S-nitrosocysteine. However at much lower concentrations, in the micromolar range, and below, (which is more akin to the in vivo levels) then the statement is not true and both these RSNOs decompose by the Cu^"^-catalysed route to give nitric oxide at about the same rate.

204


Whilst the question of nitric oxide release from RSNOs in vitro is now reasonably well understood, this is not the case for reactions in vivo and of possible enzymatic reactions. It is known that some copper-containing enzymes will also catalyse the decomposition of RSNOs. An example is copper-zinc superoxide dismutase which will decompose GSNO in the presence of GSH [582]. The reactive species is again believed to be Cu"^, generated by reduction of the enzyme-bound Cu^"^ by GSH. Superoxide (generated from xanthinexanthine oxidase) will also catalyse the decomposition of GSNO and Snitrosocysteine. Some of the biological properties of RSNOs were known before the recognition of nitric oxide as the EDRF. S-Nitrosothiols are particularly potent anti-platelet aggregation agents and also will induce vasodilation. The former effect is achieved at levels well below those required to generate the latter. It has been generally assumed that both of these effects arise from the prior release of nitric oxide, although this is not known with certainty. The circumstantial evidence currently available probably points to the conclusion that the biological properties of RSNOs do stem from their ability to generate nitric oxide. There is an increasing number of papers in the recent literature which supports the view that the biological properties of RSNOs derive from nitric oxide itself and that Cu(I) is involved in the nitric oxide release. Here are some examples :(a) The anti platelet aggregation effect of administered RSNOs is dramatically reduced when a specific Cu"^ inhibitor such as neocuproine is present [583]. (b) Similarly the vasodilation effect produced by RSNOs is reduced in a concentration-dependent way when applied to isolated rat muscle, whereas the Cu(II) chelator cuprizone has little or no effect [584]. (c) Interestingly, neocuproine also inhibits the relaxation of mouse gastric fundus for both the UV-promoted and RSNO-promoted relaxations [585], Use is now being made of the copper-catalysed decomposition of RSNOs, in potentially important clinical areas:(a) Catalytic generation of nitric oxide has been achieved from RSNOs at the surface of polymer films doped with a lipophilic Cu(II) complex. It is hoped that these materials have a more blood-compatible property [586]. (b) Copper bound to the protein PrP23-98 catalyses the release of nitric oxide from RSNOs, which again is inhibited by copper chelators [587].


205

In common with nitric oxide, RSNOs also have a cytotoxic action, again it is beUeved by prior nitric oxide formation. This effect is enhanced in oxygen-depleted situations. This in interpreted [588] in terms of competition between the RSNO and oxygen for reaction with Cu"^ (leading to Cu^"^ in the latter case). There are a number of pieces of evidence that suggest that RSNOs in biological situations are more involved than just as nitric oxide donors. These are discussed in more detail elsewhere [589]. The most promising aspect of the clinical use of RSNOs is undoubtedly derived from their very powerful anti-platelet aggregation property, which occurs without a significant lowering of blood pressure. Here, there is considerable potential for use during coronary angioplasty operations, in which GSNO is undergoing clinical trials. There is also one reference [590] to its application for the treatment of one of the rare forms of pre-eclampsia, suffered by some pregnant women. There has been much emphasis on the synthesis of sugar-derivatives of RSNOs. Some of these structures are given in Chapter 8. The reasoning behind this approach is to make use of the better transport across membranes which occurs for monosacharride derivatives, which in part derives from a higher solubility in aqueous media. It transpires that many of these so-called sugar-SNAPs are more stable in solution than is SNAP itself. In this context a number of S-nitroso peptides have been synthesised [591], which again have increased stability relative to SNAP and also greater biological activity. A relative novel approach has been adopted elsewhere, in which existing drugs have been linked to SNO groups to make the drugs bifunctional. A number of anti-inflammatory drugs such as ibuprofen have been treated in this way with the objective that the SNO function might reduce the gastric problems suffered by some patients, from the use of anti-inflammatories, by virtue of the vasodilatory effect of RSNOs, or of nitric oxide derived from them. Another example of the synthesis of a bifunctional drug is the Snitrosated derivative of yohimbine, 56, known as NMI-187, which has major

XXl^-^

206


potential in the treatment of male impotence. Another class of recently synthesised RSNOs are the fluorophore-labelled compounds, e.g. structure 57, N-dansyl-S-nitrosohomocysteine. The NHAc SNO N

O 57

fluorescence emission spectra of these derivatives are enhanced when the -NO group is removed, either by photolysis or by transnitrosation to a free thiol group. This property can be used to probe thiols/RSNOs within cells [592]. In an attempt to stabilise RSNOs in their pure state for biomedical applications, S-nitroso-N-acetylcysteine has been synthesised directly in a polyethylene glycol matrix, using the NO/O2 route. This increases the stability in solution and allows storage for many weeks without significant decomposition at -20°C in the dark [593]. There are indications both from in vivo and in vitro experiments that RSNOs can act in some physiological processes directly, rather than through prior release of nitric oxide. For example it is believed that during platelet activation, NO is transferred from administered RSNOs to thiol groups associated with platelets by a transnitrosation reaction rather than by reactions involving free NO [594] (and references therein). Further, RSNOs appear to be able to bring about bronchodilator effects in a mechanism which does not involve NO [595]. These findings could have important clinical consequences for the treatment of some lung diseases and pulmonary hypertension. Experiments with two iron-thiol complexes 2,3-dimercapto-l-propanesulfonic

"O.S 58


207

"3C\

/ \ y\ /CH3 ^N—C. Fe >—N HOCH2(CHOH)4CH2^ ^^Z \ ^ \ CH2(CHOH)4CH20H 59 acid, 58, and N-methyl-D-glucamine dithiocarbamate, 59, show that reaction occurs rapidly with a number of RSNOs, in which the NO group is transferred to the iron atom. The final UV-visible spectra are identical with those generated from a reaction of 58 and 59 with nitric oxide gas. Reaction is much faster than that leading to nitric oxide production via the copper-catalysed reaction, and is also first-order in both 58 or 59 and RSNO [596]. The results very much point to a mechanism whereby the NO group is directly transferred to the iron atom without any prior formation of nitric oxide. One oddity is that the reaction of 58 with RSNO gives the thiol RSH, whereas the reaction of 59 yields the disulfide RSSR, implying that the NO group is transferred in the NO sense for the reaction of 58 and in the NO sense for the same reaction of 59. In the chemistry of metal nitrosyls it is not unknown for the formal oxidation state of the NO ligand to be different in the product from that in the reactant. Alternatively, in the case of the reaction of 59 it is possible that RSH is first formed, which undergoes a rapid oxidation to give RSSR. Whatever the nature of that aspect of the reaction it is clear that RSNOs will rapidly and directly transfer the NO group to the iron atom of the thiol complex. Even though this is by no means a reasonable model for the haem part of the enzyme guanylate cyclase, it does raise the possibility that RSNOs might in certain circumstances be carriers of the NO group which can be directly transferred to the enzyme thus bringing about the various biological properties, which thus far have been attributed to the reaction of nitric oxide itself. This is an important area and further work is clearly needed to establish whether this is indeed so. There have been a number of reports particularly from the work of Vanin and co-workers, that Fe^"^ will also decompose RSNOs leading to nitric oxide production [597-8]. In contrast to the proposals for the copper-catalysed reactions the Fe^'^-induced reactions are not simple one-electron transfer processes but involve a more complex series of reactions involving the

208


c 2+

r^

NO^RS^ 60

intermediate 60, where the iron binds two molecules of RSNO, which then break up giving thiolate anions (RS~) and thiyl radicals (RS*), leaving the NO groups bound to the iron. 12.4. Inorganic nitrosyl complexes 12.4.1. Sodium nitroprusside Na2[(CN)^FeNO](SNP) This is by far the most widely studied and most used in biological systems of metal nitrosyl complexes in the context of nitric oxide donors. It has a powerful vasodilation action and as such has had a significant clinical use particularly during hypertensive emergencies. Inhibition of platelet aggregation can also occur in the presence of vascular tissue and there are reports that penile response can be effected by transurethral application. The in vitro chemistry of SNP has been much studied, particularly with regard to possible nitric oxide release. Solutions of SNP in water are quite stable over long periods of time. Photochemical decomposition occurs readily releasing cyanide ion, and reactions with nucleophiles particularly thiolate anions have been examined in detail. Some of these reactions are covered in Chapter 9. There has been no significant reaction in these in vitro experiments which leads to quantitative nitric oxide formation, although some quite low yields have sometimes been reported. Nevertheless it has been shown that SNP does generate nitric oxide in vivo and this is believed to account for its pronounced vasodilatory action. The mechanism is not fully understood although both enzyme-catalysed and non enzyme reactions have been reported. Both these require the presence of vascular tissue or a one-electron reducing agent. The photolytic reaction is not a possibility in biological systems, though it could play a part under operating theatre conditions, if precautions are not taken to exclude light prior to administration. In all of these situations cyanide release accompanies nitric oxide formation and this is a serious drawback when working with living cells. However, since the vasodilation effect is so powerful, the required dose to effect substantial effects is quite low, and under these conditions biological systems are able to deal with the low level of cyanide ion released.


209

12.4.2. Ruthenium nitrosyls Since ruthenium complexes have a powerful affinity for nitric oxide, there has been much activity concerning the possibility of the use of ruthenium nitrosyls as nitric oxide donors. Nitric oxide can be released by a photochemical process and also by a reduction process, from, for example the octahedral complex shown in Eq. (389), [599]. The rate of NO release is /raw5-[Ru(L)(NO)(NH3)4] 2+

trans- [Ru(L)(NO)(NH3)4] 3+ + e"

H2O 2+ + NO /raf7^-[Ru(L)(H20)(NH3)4f""

(389) affected by the nature of the ligand L, and there is an expectation that a slowrelease long lasting nitric oxide donor might be developed using this finding. A ruthenium nitrosyl has been shown to be part of a remarkably simple model for the biosynthetic pathway for nitric oxide synthesis from L-arginine [600]. The ruthenium(III) complex [Ru(Hedta)Cl]~ reacts with L-arginine and hydrogen peroxide to generate a ruthenium(II) nitrosyl complex together with Ru(III) + arginine + H2O2 — ^

[Ru(II)NO]'^ + citruUine + H2O (390)

L-citrulline, Eq. (390). The evidence points conclusively to a reaction pathway where nitric oxide is generated from L-arginine and hydrogen peroxide, which is then complexed by Ru(II). Possible mechanisms are discussed [600]. 12.4.3. Dinitrosyl-iron-thiol complexes (DNICs) Since the discovery that DNICs exist in biological systems, their chemistry has been widely examined, particularly with regard to their ability to act as NO donors. Structure 61 shows their general formula. The most widely RS.

NO

RS^

^NO 61

210


studied examples studied in the laboratory are for the derivatives of cysteine and of glutathione. In biological systems the complexes are formed with high molecular weight proteins containing free thiol groups. They can be prepared readily from aqueous solutions of ferrous sulfate, the thiol and nitric oxide under anaerobic conditions, and also from S-nitrosothiols and ferrous salts. The low molecular weight examples can exist in both paramagnetic and diamagnetic forms and the former has been examined by EPR. At room temperature they decompose quite rapidly generating nitric oxide. The proteinbound examples are more stable. DNICs in general have been shown to have many of the biological properties of nitric oxide itself, such as inhibition of platelet aggregation, vasodilation and the lowering of blood pressure, and so the general assumption is that they accomplish these by prior decomposition to nitric oxide. This is supported by the observation that the biological properties are inhibited by haemoglobin, presumably by the scavaging of nitric oxide by oxyhaemoglobin. It has been demonstrated [601] that DNICs and also RSNOs activate the enzyme soluble guanylate cyclase, in the same way as does nitric oxide itself It has been suggested that DNICs (protein-bound) might act as storage and transport vehicles of nitric oxide in vivo, given the known very short half life of nitric oxide itself This is a role that has also been ascribed to RSNOs (also protein-bound). It is not surprising given the complexity of the chemistry in vivo, that it has not yet proved possible to distinguish between these (and other) possibilities. The ready NO group transfer between S (in RSNOs) and Fe (in DNICs) may indicate that both have a role in this regard [602]. 12.4.4. Iron-sulfur cluster nitrosyls The best known examples here are Roussin's black and red salts described in Chapter 9 as structures 32 and 33. Apart from being something of a chemical oddity there has been much interest generated in their chemistry since the discovery that iron-sulfur clusters form an important part of a number of enzymes and co-enzymes. Their synthesis is described in Chapter 9. When thiols are used instead of inorganic sulfides then the so-called neutral esters are formed e.g. [Fe2(SR)2(NO)4]. The esters can also be generated by alkylation of the salts with alkyl halides. The structures of these compounds together with aspects of their chemistry, particularly with regard to NO release and their biological properties, have been well established. Photolysis of both the salts and their so-called esters in aqueous solution leads to quantitative nitric oxide formation. This is shown in Eq. (391) for the reaction of the red salt, when the other product is the black salt [603].

Nitric Oxide Releasing Compounds (NO donors) [Fe2S2(NO)4]^-

^^ >

[Fe2S2(NO)3]^-

+ NO

211 (^91)

[Fe2S2(NO)4]^-

[Fe4S3(NO)7]^- ^

[Fe4S4(NO)7]^~

Similarly a number of red ester compounds yield four mol of nitric oxide [604]. The application of these and earlier results to the delivery of nitric oxide to specific biological targets is currently being developed. Experiments with cell cultures have shown that Roussin's red salt is a good candidate. Roussin's salts, particularly the black salt 32, are rather surprisingly soluble in organic solvents - a property which could assist nitric oxide liberation in biological systems by its lipid solubility. Nitric oxide can readily be obtained in vitro from 32 by oxidation e.g. using Cu(II) salts. These clusters may also act as NO"^ donors and thus bring about nitrosation of e.g. amines. Evidence for this comes from a detailed analysis of NMR studies [518]. There are a number of examples where the methyl ester of Roussin's red salt will nitrosate secondary amines, such as diethylamine, morpholine and pyrrolidine, Eq. (392), particularly in the presence of air [605], which suggests [Fe2(SMe)2(NO)4]

+

^NH

^

^NNO

(392)

that nitric oxide is first released which in air is a potent nitrosating system. These iron-sulfur cluster nitrosyls show many of the biological properties of nitric oxide itself, i.e. vasodilation, anti platelet aggregation, antimicrobial and antitumor activity etc. The inhibition of platelet aggregation is itself inhibited by haemoglobin, which reinforces the view that in biological systems the physiological properties derive from the action of nitric oxide itself produced from these nitrosyls [606]. There is a recent comprehensive account of the chemistry and biology of these non-haem iron nitrosyls [607]. 12.5. NONOates (Diazeniumdiolates) In 1960 Drago and Paulik [608] obtained a complex by passing nitric oxide into a solution of an amine in an organic solvent, Eq. (393). The structure was established as an adduct in which nitric oxide dimer is formally bound to the nitrogen atom of the amine as the anionic part of a salt with the

212


2R2NH + 2NO

•

R2N^ ^N-0~R2NH2

(393)

dialkylammonium cation. These are known as Drago complexes and other salts can be synthesised readily. The original ammonium salts are hygroscopic and the sodium salts are more stable and easier to handle in the laboratory. Reaction is quite general for a large range of R groups. There has been some confusion in the literature regarding the nomenclature of these salts, the trivial name NONOate is now being replaced by the more descriptive diazeniumdiolate. It turns out that these salts are part of a wider family of compounds in which bonding occurs in the anion from one of the nitrogen atoms of the N2O2 grouping with a variety of other elements. The general formula is given in the resonance forms 62 and 63. Examples include bonding to aromatic systems

i 62

63

such as in 64 (the cupferron reagent), via aliphatic carbon, 65, oxygen 66 (Angeli's salt), and sulfiir, 67. An excellent comprehensive review of the

^N-O-NH/

R2

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