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Polyhedron 24 (2005) 1167–1174

Kinetics of oxidation of L-histidine by tetrachloroaurate(III) ionin perchloric acid solution

Vimal Soni, R.S. Sindal, Raj N. Mehrotra *

Department of Chemistry, JNV University, Pali Road, Jodhpur 342 005, India

Received 14 January 2005; accepted 9 March 2005

Available online 25 May 2005

Abstract

The oxidation of L-histidine by tetrachloroaurate(III) ion in HClO4 is first-order in both AuCl4� and histidine. The stoichiometry

ratio, D½AuCl4��=D½Histidine�, is 1.07 ± 0.10 and the oxidation product is b-imidazolylpuruvic acid. The kobs decreased with increase

in [H+] and [Cl�] at the constant concentration of the other. The rate dependence on [Cl�] is due to the equilibrium between the

reactive AuCl4� and AuCl3OH� species of Au(III); the reactivity being in the order: AuCl3OH� � AuCl4

� ion. The rate depen-

dence on [H+] is attributed to the reactive histidine species, RCH2-CHðNH3þÞCOO� which is in equilibrium with the non-reactive

R-CHðNH3þÞCOOH where R is the imidazole ring. An inner-sphere mechanism based on the NMR study on the complex formed

between [AuCl4]� with glycine is suggested. The outer-sphere mechanism is ruled out because the DS� values are very different for

the similarly charged species AuCl4� and AuCl3OH�.

� 2005 Elsevier Ltd. All rights reserved.

Keywords: Mechanism; Kinetics

1. Introduction

The reduction of Au(III)-complexes has been a sub-

ject of interest [1,2] in recent years. Substrates such as

dimethylsulfide [1c], thiocyanate [1d,2a], thiosulfate

[1e] and iodide [2d] (Y) ions attack the coordinated li-

gand X of the Au(III)-complex in the rate determining

step [3,4] which is followed by a fast bridged-two-elec-

tron transfer to the metal centre with the elimination

of XY. The [Pt(CN)4]2� and [Pt(NH3)4]

2+ ions are sim-ilarly oxidised [5]. There is a direct interaction between

the reducing nucleophile and the metal centre in the oxi-

dations of the sulfite [4] and thiosulfate ions [6]. In the

oxidation of thiocyanate [3b], dimethyl sulfide [7] and

thiosulfate [6] the substitution is faster than the reduc-

tion of [AuCl4]�, whereas for the iodide ion the redox

reaction is faster than the ligand substitution [8]. How-

0277-5387/$ - see front matter � 2005 Elsevier Ltd. All rights reserved.

doi:10.1016/j.poly.2005.03.057

* Corresponding author. Tel.: +912912721478.

E-mail address: raj_n_mehrotra@yahoo.co.uk (R.N. Mehrotra).

ever, in the oxidation of thiocyanate by AuðNH3Þ43þthe rate controlling step is the substitution of an amineligand by thiocyanate, whereas for the oxidation by

AuðNH3Þ2X2þ, thiocyanate replaces halide (X) in two

rapid consecutive and reversible substitution steps prior

to the slower reduction [9].

The disulfide bond in the oxidation of cystine cleaved

to sulfonic acid [10–12], and the sulfur in methionine

was oxidised stereospecifically to the sulfoxide [13–15].

The study of the crystal structure of the complexes ob-tained by the slow reaction of AuCl4

� with dipeptide

glycyl-L-histidine [16], glycylglycyl-L-histidine [17] and

1,2-diaminoethane-N,N,N 0,N 0-tetra-(N-methylacetamide)

[18] suggested that these complexes are square planar.

There seems to be no kinetic study available on the

oxidation of amino acids, though NMR investigation

with [15N]glycine [1a], suggested the formation of the

[AuCl3N15H2CH2COOH] complex via the amino group

and subsequently a chelate in which N and O are

bonded with Au(III) is formed. The chelate, probably,

1168 V. Soni et al. / Polyhedron 24 (2005) 1167–1174

is inert to the oxidation of glycine. Two electrons within

the complex are transferred to Au(III) from glycine to

form a Au(I)-imine intermediate, which being unstable

undergoes hydrolysis to NH4þ and glyoxylic acid. An

excess of [AuCl4]� oxidizes glyoxylic acid to CO2 and

Au(o) through the intermediate formation of formicacid. Since the rate law was not given (dependence of

the rate on [glycine], [H+] and [Cl�]), it was of interest

to study the kinetics of the oxidation of histidine, an-

other amino acid.

2. Experimental

2.1. Chemicals and solutions

The solutions of AuCl4� (Johnson and Matthey) were

freshly prepared in perchloric acid (E. Merck, GR),

though such solutions are stable over 24 h. The concen-

trations were determined from the optical density

measured at 310 nm on a HP 8452A diode array spec-

trophotometer using carefully determined e = 4.82 ·103 dm3 mol�1 cm�1, which compared well with the lit.

value [19–21] 4.86 · 103 dm3 mol�1 cm�1. The solutions

of L-histidine (BioChemika Micro Select, Fluka) and

D-histidine (Puriss, Fluka) were freshly prepared by

weighing the samples. LiClO4 solution for adjusting

the ionic strength (l) was prepared and standardised

as described earlier [22]. All other chemicals were used

as received. The solutions were prepared in water, dis-tilled twice, and purged by nitrogen.

2.2. Reaction products

At the end of the reaction, the metal ion from the

reaction mixture was removed using an ion exchange

method. A portion of the filtrate was treated with a sat-

urated solution of 2,4-dinitrophenylhydrazine and thesame was left overnight in a refrigerator. The precipi-

tated 2,4-dinitrophenylhydrazone was separated by fil-

tration and recrystallised from a mixture of ethyl

acetate and light petroleum. The melting point of the

precipitated hydrazone, 190 ± 0.5 �C, was characteristicof b-imidazolyl pyruvic acid, indicating that the seat of

the reaction is not the imidazole ring since it remained

intact. Another portion of the eluent on treatment withNessler�s reagent confirmed the presence of the NH4

þ

ion.

2.3. Stoichiometry

Several reaction mixtures having excess of AuCl4�

over known histidine concentrations were prepared.

The optical density of the reaction mixtures was mea-sured at 310 nm at room temperature until a constant

value was recorded. The mean of several such estimates

yielded D½AuCl4��=D½Histidine� ¼ 1.07� 0.10. Hence, in

view of the characterised products and the stoichiometry

ratio, the equation of the reaction is expressible by

Eq. (1).

H2Oþ C3H3N2CH2CHðNH2ÞCOOH þAuCl4�

! C3H3N2CH2COCOOHþAuCl43� þNH4

þ þHþ

ð1Þ

2.4. Kinetics

The reaction mixtures were prepared so as to con-

form to the pseudo-first-order conditions (excess of

histidine over Au(III) ion) for studying the rate ofthe reaction that was followed at 360 nm (Beer�s law

was obeyed) with a Spectrochem MK II colorimeter

fitted with a thermostated reaction cell. The pseudo

first order rate constants kobs were calculated from

the slopes of the plots between log (At�A1) against

time ‘‘t’’, which were linear for more than two half-

lives (80–90%) where At and A1 are the optical

densities of the reaction mixture at any time ‘‘t’’and infinite time, respectively. These linear plots indi-

cated the monophasic character of the reaction and

that the products did not interfere with the rate pro-

file of the reaction. There was no trace of metallic

gold in the reaction mixture at the time of measuring

A1 after several half-lives. The EXCEL program was

used for plotting the log (At � A1)-time data and

evaluating the least square values of the slopes andintercepts of the linear plots. The pseudo first order

rate constants, kobs were generally reproducible

within 5%.

The repetitive spectra of the reaction mixtures for

several concentrations of histidine, recorded on a

HP8451A spectrophotometer, had not shown any

change from the spectral characteristics of the control

solution of AuCl4�. The absorbance of the spectra de-

creased with time and no isosbestic points appeared,

indicating that the spectra of the products of the reac-

tion do not overlap with that of the AuCl4 solution.

The absence of any transient absorbance peaks in the

spectra over a period of time, therefore, suggests that

the intermediates, if formed, are weak or only weak

bridged complexes are formed.

2.5. Test for free radical

The addition of acrylonitrile (6% v/v) to the reaction

mixture, degassed with nitrogen, produced cloudiness

due to the formation of polyacrylonitrile, indicating

the formation of free radicals. A similar solution of

acrylonitrile when added to similarly degassed blank

solutions of AuCl4 and histidine did not produce anycloudiness.

V. Soni et al. / Polyhedron 24 (2005) 1167–1174 1169

3. Results

kobs increased proportionately with [Histidine] at con-

stant [H+], [Cl�] and ionic strength (Table 1). The plot

of kobs against [Histidine]0 was linear and passed

through the origin, suggesting a first order dependencein histidine. The decreasing kobs with increasing [Cl�]

and [H+], at constant concentrations of the other, is re-

ported in Tables 1 and 2, respectively. The correlations

between kobs and [H+] or [Cl�] at constant concentration

of the other, are shown in Figs. 1 and 2, respectively.

These plots indicate the absence of a linear correlation

between kobs and the respective concentrations of H+

ð4Þ

ð5Þ

ð6Þ

and Cl� ions. A later analysis of the respective data indi-

cated that a linear correlation between kobs and [H+] or

[Cl�] is consistent with the empirical rate law, Eq. (2),

where k2 = kobs/[Histidine] and X is a constant. These re-

sults suggested the presence of an equilibria between

species that differ by a single proton and a [Cl�] ion

k2 � X ¼ aþ b½Hþ��1½Cl���1. ð2Þ

kobs decreased with an increase in the ionic strength,

Table 3, and the plot, in accordance with Eq. (3)

[23,24], is linear with a negative slope, Fig. 5, suggesting

the reactive species have opposite charges

log kobs ¼ log k0 þ 1.02ZaZb

ffiffiffil

p

1þ ffiffiffil

p � 0.05l

� �. ð3Þ

4. Mechanism and discussion

The following equilibria between various histidinespecies, A2+, B+, C and D� (the charge shown is the

net total charge on that species), and the respective

equilibrium constants [25] are known. The retarding ef-

fect of the H+ ion could be due to any one or all the

equilibria shown in Eqs. (4)–(6). However, these equi-

libria have no bearing on the retarding effect of the

Cl� ion. For an initial [Histidine] = 0.002 and[H+] = 0.02 mol dm�3, the calculated [B+] is

8.5 · 10�4 mol dm�3 using the given Ka1 value The con-

centrations of C and D�, using the respective Ka2 and

Ka3 values, can be shown to be negligibly small com-

pared to [B+]. The equilibria (7)–(9), involving Au(III)

species [19,26,27], can also explain the retarding effect

of both H+ and Cl� ions.

AuCl4� þH2O �

Khy

AuCl3H2Oþ Cl� ð7Þ

AuCl3H2O �Ka

AuCl3OH� þHþ ð8Þ

AuCl3ðOHÞ� þH2O �

Kaq

AuCl2ðH2OÞðOHÞ þ Cl� ð9Þ

The equilibrium (9) is excluded from consideration be-

cause the calculated [AuCl2(H2O)(OH)] is negligibly

small. The equilibria (7) and (8) can be combined to a

single equilibrium (10). A consequence of equilibrium

(10) is that the likely reactive Au(III) species are

AuCl4� and AuCl3OH� ions. The equilibrium constant

K (=KhyKa) can be calculated from the known Khy andKa values. The rate constants tend to limiting values at

the higher end of the [H+] (Fig. 1) and [Cl�] (Fig. 2) be-

cause the equilibrium constant K has a very small value.

The reactivity of AuCl4� and AuCl3OH� ions is consis-

tent with the analysis of the rate data that indicated two

parallel reactions. The reactivity of AuCl3OH� is �103

times the reactivity of AuCl4�ðaqÞ (Table 4). The

12

14

1170 V. Soni et al. / Polyhedron 24 (2005) 1167–1174

participation of a hydroxo species, AuCl3OH�, is con-

sistent with the inverse dependence of the rate on [H+],

that is suggestive of an inner-sphere mechanism [28].

Scheme 1.

Table 1

The dependence of kobs (s�1) on [L-Histidine] and [Cl�]at 31 �C104½AuCl4

�� ¼ 2.0, [H+] = 0.026 and l = 0.4 mol dm�3

S. No. [L-Histidine] [Cl�] 103kobs

1 0.002 0.013 0.234

2 0.005 0.013 0.569

3 0.011 0.013 1.23

4 0.020 0.013 2.25

5 0.050 0.013 5.68

6 0.100 0.013 10.4

7 0.150 0.013 16.1

8 0.200 0.013 23.9

9 0.250 0.013 28.3

10 0.300 0.013 33.1

11 0.400 0.013 46.4

12 0.020 0.013 2.34

13 0.020 0.020 1.83

14 0.020 0.025 1.66

15 0.020 0.030 1.49

16 0.020 0.050 1.31

17 0.020 0.062 1.21

18 0.020 0.074 1.18

19 0.020 0.084 1.13

20 0.020 0.094 1.10

21 0.020 0.150 1.05

22 0.020 0.200 1.02

0.8

1

1.2

1.4

1.6

1.8

2

2.2

2.4

2.6

2.8

3

0 0.03 0.06 0.09 0.12 0.15 0.18 0.21 0.24

[Cl-] (m ol dm-3)

103 k

obs

(s-1

)

Fig. 2. The plot of kobs against [Cl�] at constant [H+] at 31 �C.

0

2

4

6

8

10

0 0.02 0.04 0.06 0.08 0.1

[H+] (mol dm-3)

kob

s (s

-1)

Fig. 1. The plots of kobs against [H+], at constant [Cl�], at 27 �C (d),

31 �C (s), 35 �C (¤), 40 �C (n) and 45 �C (h).

Table 2

The dependence of kobs (s�1) on [H+] at different temperatures

[H+] (mol dm3): 0.022 0.024 0.026 0.028 0.030 0.040 0.050 0.060 0.070

103kobs, 27 �C 2.02 1.81 1.67 1.47 1.35 0.943 0.731 0.582 0.494

103kobs, 31 �C 2.79 2.53 2.25 2.07 1.92 1.31 1.02 0.821 0.696

103kobs, 35 �C 4.50 3.98 3.62 3.26 2.97 2.01 1.50 1.15 0.946

103kobs, 40 �C 8.93 7.86 7.04 6.23 5.67 3.76 2.72 2.06 1.66

103kobs, 45 �C 12.2 10.4 9.44 8.49 7.85 5.03 3.66 2.71 2.21

104½AuCl4�� ¼ 2.0, [L-Histidine] = 0.02, [Cl�] = 0.013 and l = 0.4 mol dm�3.

0

0.2

0.4

0.6

0.8

1

1.2

1.4

1.6

1.8

0 10 20 30 40 50 60

[H+]

-1 (dm

3 mol

-1)

k2(

Ka1

+ [

H+ ])

/Ka1

Fig. 3. The linear plot of k2ðKa þ ½Hþ�ÞK�1a against [H+]�1 at constant

[Cl�] = 0.13 mol dm�3, consistent with the Eq. (21), at temperatures

27 �C (h), 31 �C (M), 40 �C (s) and 45 �C (d).

Table 4

The values of k and k1 at different temperatures and the related

thermodynamic parameters

Temperature (�C): 27 31 35 40 45

kc (dm3 mol�1 s�1) 0.0813 0.114 0.125 0.151 0.189

10�2kc1 (mol dm�3 s�1) 1.10 1.38 2.33 4.59 5.62

107 K (mol2 dm�6) 4.12 4.51 4.91 5.46 6.04

DH zk ¼ 32� 2 kJ mol�1; DH z

k1¼ 76� 5 kJ mol�1; DSzk ¼ �159�

8 J K�1 mol�1; DSzk1 ¼ 48� 17 J K�1 mol�1.

Table 3

The effect of the ionic strength (LiClO4) on the rate of the reaction at

35 �C

l (mol dm�3) 0.1 0.20 0.40 0.60 1.00 1.20

103kobs (s�1) 3.49 3.29 3.20 3.10 3.01 2.96

104½AuCl4�� ¼ 1.0, [Histidine] = 0.005, [H+] = 0.01 mol dm�3.

V. Soni et al. / Polyhedron 24 (2005) 1167–1174 1171

Though the formation of intermediates between

Au(III) and histidine is not indicated by the spectral

study of the reaction mixture in the UV region, their for-

mation in the proposed mechanism would be reasonable

in view of the NMR studies on glycine [1a], glycylglycyl-

L-histidine [17]. The observation that the imidazole ring

is not the seat of reaction because it is present in the final

oxidation product of the reaction is consistent with theview that the binding of a proton or a metal ion is very

unfavourable at the pyrrole nitrogen of histidine [17].

The results of the spectral studies would not be in con-

tradiction to the assumption of the formation of tran-

sient intermediates in the rate-determining step. The

mechanism in Scheme 1 is based on this assumption.

The probable free radical is R-CH2C�(NH2)COOH

where R is the imidazole ring. The presence of a freeradical is indicative of the one electron reduction of

Au(III) in the rate determining step. Transient Au(II)

is reported in several redox reactions [29–35]. The unsta-

ble Au(II) can oxidise the free radical instantaneously to

the imine (15). Imines with a proton on N are seldom

stable and undergo fast hydrolysis to the NH4þ ion

and the final product (17).

Alternately, Au(II) may disproportionate to Au(III)andAu(I) instantaneously, reaction (17), because the rate

of disproportionation [36], 1.67 · 108 dm3 mol�1 s�1, is

very near to the rate of diffusion. Au(II) is, therefore, very

likely removed through disproportionation that makes

the reactions (15) most unlikely. The rate law, deduced

from reactions (10)–(17), is expressible by Eq. (18) which

changes to Eq. (21) on substituting the proper values of

½AuCl4�� and [B+] (R–CH2CH(N+H3)COOH), given by

Eqs. (19) and (20), respectively, and applying the inequal-

ity K � [H+][Cl�] in view of [H+]and [Cl�] used where

k2 = kobs/[Histidine]0

�d½AuðIIIÞ�dt

¼ k þ k1K½Hþ�½Cl��

� �AuCl4

�� �½Bþ�; ð18Þ

AuCl4�� �

¼ ½AuðIIIÞ�½Hþ�½Cl��K þ ½Hþ�½Cl�� ; ð19Þ

½Bþ� ¼ Ka1½Histidine�0Ka1 þ ½Hþ� ; ð20Þ

k2ðKa1 þ ½Hþ�ÞKa1

¼ k þ k1K½Hþ�½Cl��

� �. ð21Þ

The reactivity of the species B+ in outer-sphere

oxidation with [Co(III)W]5� (E0 = 1.01 V [37],) and

[Mncdta]� (E0 = 0.81 V [38],) ions [39] is of the same

order as are the respective E0 values of these ions, and

so is the reactivity of the species D� with [Mncdta]�

and FeðCNÞ63� (E0 = 0.36 V) ions [40]. Though E0

(=0.95 V [41],) for the AuCl4�=AuCl2 couple is between

the E0 values of [Co(III)W]5� and [Mncdta(OH2)]� ions,

the reactivity of the species B+ with the Au(III) ion is

greater than that with the [Co(III)W]5� ion. This is

probably because the present reaction is inner-sphere.

The slope value, �0.3, of the plot in Fig. 5 is less than

the expected slope of �1 considering that the reactive

histidine species B+ has a net single positive charge

and the reactive Au(III) species have a single negative

charge. The deviation from the expected value couldbe due to the spreading of the net single positive charge

over the entire histidine molecule.

The Eq. (21) is consistent with the linear plots of the

left hand side against [H+]�1 at constant [Cl�] (Fig. 3),

and [Cl�]�1 at constant [H+] (Fig. 4) with intercepts on

the rate axis. The values of k (0.120 dm3 mol�1 s�1) and

k1 (139 dm3 mol�1 s�1) obtained from the plot in Fig. 4

are in close agreement with those (k =0.114 dm3 mol�1 s�1; k1 = 15.6 dm3 mol�1 s�1) obtained

from the plot in Fig. 3 at 31 �C. This close agreement

0.12

0.14

0.16

0.18

0.2

0.22

0.24

0.26

0.28

0.3

0.32

0 10 20 30 40 50 60 70 80[Cl-]-1 (dm3 mol-1)

k2 (

Ka1

+ [H

+ ])/K

a1 (

dm3 m

ol-1

s-1)

Fig. 4. The linear plot of k2ðKa þ ½Hþ�ÞK�1a against [Cl�]�1 at constant

[H+], consistent with the Eq. (21), at 31 �C.

0.4

0.42

0.44

0.46

0.48

0.5

0.52

0.54

0.56

0.58

0.6

0 0.1 0.

(√µ(1+√µ) − 0.05µ)2 0.3 0.4 0.5

3+lo

gk

obs

Fig. 5. The linear plot of log kobs against fl=ð1þplÞ � 0.05lg has an

intercept of �0.3. The ionic strength is varied using LiClO4.

KAuCl4

−(aq) + H2O AuCl3OH− + Cl− + H+ (10)

KcAuC l4

− + R-CH2CH(N+H3)COOH [AuCl4− R-CH2CH(NH2)COOH] + H+ (22)

C

Kc1AuCl3OH− + R-CH2CH(N+H3)COOH [AuCl3R-CH2CH(NH2)COOH] + H2O (23)

C1

kcC

AuCl42- +

NH2

COOH + H+CH2

.

C

N

NH

.

NH2CH(COOH)CH2-RCl

Cl ClAu

Cl

(24)

CH2 COOH

HN H

C

H

AuClCl

Cl

AuCl3- +

NH2

COOH + H+CH2

.

C

N

NH

N

NH

.

C1

kc1

(25)

fastR-CH2C

−(NH2)COOH + Au(II) R-CH2CH(N+H)COOH + Au(I) (15)

fast

R-CH2CH(N+H)COOH + H2O R-CH2COCOOH + NH4+ (16)

fast 2Au(II) Au(III) + Au(I) (17)

Scheme 2.

1172 V. Soni et al. / Polyhedron 24 (2005) 1167–1174

between the values of k and k1 obtained from two differ-

ent plots, not only support the validity of the rate law butthat of the proposed mechanism too. The values of k and

k1 at different temperatures are in Table 4. The values of

Ka1 [25], and K [20] (Table 4) at different temperatures,

were estimated from the published data. The large differ-

ence in the values of k and k1 (k1 � 103k) reflect the ease

of formation of the respective transient species.

The two intermediates are formed by equatorial coor-

dination of histidine with the respective Au3+ species.The rate of formation of the intermediate formed by

elimination of a water molecule is obviously faster than

that formed through a chloride bridge. This mechanism

is unlikely to show any difference in the rates of oxida-

tion of H1-histidine and H2-histidine (deuteriated mole-

cule). Since an authentic sample of H2-histidine was

unprocurable from a reliable source, the rates of H2-his-

tidine were not measurable. However, the most disturb-

ing aspect of the mechanism is the value of k1(110 dm3 mol�1 s�1 at 27 �C) which is inadmissible fora process in which a water molecule is eliminated be-

cause of the reaction between H+ and OH� ions. The

rate of such a process is diffusion controlled. This indi-

cated the presence of some other factor with k1 and thus

a modification in the mechanism is required.

The alternate mechanism presumes the formation of

the intermediates in equilibrium steps rather than in

the rate-limiting steps. The intermediates so formedcould be of a transitory nature defying spectrometric

detection if the rates of the formation and decomposi-

tion of the intermediates to the products approach the

rate of the backward reaction of the equilibrium, so that

there is no accumulation of the intermediates over a per-

iod of time. Yet another reason for the intermediates

defying spectrometric detection could be that the molar

absorptivity coefficient of the transition complexesmight not be very different from that of tetrachloroau-

rate(III) solution. The mechanism so modified is consid-

ered in Scheme 2 in which the reaction (15A) is

eliminated for the reasons given earlier.

The rate law, based on Scheme 2, is in Eq. (26). The

transformed form of (26) is Eq. (27) which is identical

with Eq. (21) where k = kcKc and k1 = kc1Kc1

�d½AuðIIIÞ�dt

¼ kcKc þkc1Kc1K½Hþ�½Cl��

� �AuCl4

��� �

½Bþ�;

ð26Þ

k2ðKa1þ½Hþ�ÞKa1

¼ kcKcþkc1Kc1K½Hþ�½Cl��

� �¼ kþ k1K

½Hþ�½Cl��

� �.

ð27Þ

V. Soni et al. / Polyhedron 24 (2005) 1167–1174 1173

The square planar AuCl4� is isoelectronic with the

square planar [Ag(OH)4]� ion. The latter is reported

to form intermediates with the H2PO2� ion [42] and

several others [43–50] in which the ligand is axially

coordinated. The high rates of S(IV) oxidation with

the AuCl4� ion is attributed to the axially coordinated

SO22� and HSO4

� ions in the intermediates [1e]. There-

fore histidine can coordinate with Au(III) either axially

(C1) or equatorially (C). In the equatorial coordina-

tion, the amide nitrogen can coordinate with AuCl4�

through a chloride bridge [1e] though glycine

substitutes for one of the chlorines. It may be men-

tioned that it is not clear if glycine had two simulta-

neous paths of oxidation as observed presently. Sincethe nature of coordination in the two intermediates is

different, their rate of decomposition are obviously ex-

pected to be different and there is no direct removal of

OH� by a proton as in Scheme 1.

In conclusion, histidine forms transitory intermedi-

ates with AuCl4� and AuCl3OH�, preferring an equato-

rial coordination with the AuCl4� ion, and the

substitution is likely to be through a chloride bridgewhereas an axial coordination is preferred with the

AuCl3OH� ion despite the sterically hindered large his-

tidine molecule. Since this coordinated species may be

under strain it obviously has an overall high rate of

decomposition compared to strain free equatorially

coordinated species. The intermediates defy spectromet-

ric detection either because the total rates of formation

and decomposition to the products have approachedthe rate of the backwards reaction of the equilibrium

step so that there is no accumulation of the intermedi-

ates over a period of time or the molar absorptivity coef-

ficient of the transition complexes is not very different

from that of the tetrachloroaurate(III) solution.

Acknowledgement

Sponsorship of this work by the University Grants

Commission (F.12-59/97 and F.12-147/01) is gratefully

acknowledged.

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