Mechanism of the oxidation of hydrazoic acidby tetrachloroaurate(III) ion

Vimal Soni Æ Raj N. Mehrotra

Received: 19 October 2007 / Accepted: 11 December 2007 / Published online: 26 January 2008

� Springer Science+Business Media B.V. 2008

Abstract The oxidation of hydrazoic acid in perchloric

acid in the absence of added chloride under pseudo first-order

conditions ([HN3] » [AuCl4-]) is first order in [Au(III)].

Michaelis–Menten type of dependence (linear plots of kobs-1 vs

[HN3]-1) is observed with respect to [HN3]. The kobs is

independent of ionic strength and the plot between kobs-1 and

[H+] is linear. The inner-sphere mechanism is consistent with

the formation of an axial complex (K = 25 dm3 mol-1)

between AuCl3(HO)- ion and HN3 prior to its rate deter-

mining decomposition (k = 0.0182 s-1). It is inferred that

the free radicals N3• do not oxidise Au(II). The reaction

becomes outer-sphere in the presence of added Cl- ions

which are inferred to form a cage around the hydronium ion

surrounding the AuCl4- ions. The penetration of N3

- through

the cage is rate controlling and within the cage, the electron

transfer from N3- ion to AuCl4

- is fast. The value of the rate

determining constant k2 is 0.547 dm3 mol-1 s-1 and the

equilibrium constant KCl for the cage formation is 5 dm3

mol-1 at 25 �C. It is calculated that the minimum HN3

concentration required before the reaction exhibits zero-

order dependence in HN3 is 0.31 mol dm-3 when

[H+] = 0.18 mol dm-3 at 25 �C.

Introduction

The chemistry of azide ion has been reviewed more than once

[1–4]. Depending on the nature of the oxidant, both inner and

outer-sphere mechanisms have been observed. The

oxidations by one-electron oxidants such as Co3+(aq) [5],

Mn3+(aq) [6], [Mn(bipy)2OH]2+ [7], [Mn(edta)(H2O]- [8],

Ce4+(aqV) [9], [Ni(bipy)3]3+ [10] and [Co(III)O4W12O36]5-

(abbreviated to [Co(III)W]5-) [11] in acidic medium, and

by two-electron oxidants viz. [Ag(OH)4]- [12], [CuIII

(H2TeO6)2]5- and [AgIII(H2TeO6)2]5- [13] in alkaline

medium and [(CH2)2(C2N5H6)2Ag]3+, ethylenebisbigua-

nidesilver(III) in acidic medium [14] are inner-sphere.

A reinvestigation of an earlier study [15] with IrCl62-, IrBr6

2-

and [Fe(bpy)3]3+ ions made in the presence of N-tert-butyl-a-

phenylnitrone (PBN) and 5,5-dimethyl-1-pyrroline N-oxide

(DMPO), traps for the free radical N3•, suggested that the

reactions are outer-sphere [16]. The outer-sphere nature of

the reactions is supported by the application of the Marcus

cross relation [17].

Concerted efforts have been made to understand the

oxidations of organic and inorganic substrates, such as

oxalic acid [18], hydroxylamine [19], histidine [20], gly-

cine [21], Pt(CN)42- [22], I- [23], 4-thio-20-deoxyuridine

and 4-thiouridine 50-monophosphate [24], dimethyl sul-

phide [25], thiocyanate [26], HNO2 [27], HSO3- [28] and

As(III) [29] by various gold(III) complexes. It is suggested

that Au(III) complexes are reduced to the corresponding

Au(I) complexes by soft nucleophiles [30] either by a

direct two-electron transfer without observable intermedi-

ates [23] or by rapid and/or consecutive substitutions

followed by reductive elimination through attack on the

complex by the reducing agent [31].

In the light of the above observations, we have studied

the oxidation of azide ion, a relatively hard nucleophile, by

acidic AuCl4- to know whether the mechanism of this

reaction has any resemblance to the oxidation by

Ag(OH)4-, since AuCl4

- is isoelectronic with Ag(OH)4-. It

was also of interest to know whether the oxidation involves

two-electron transfer in a single step, or one-electron

V. Soni � R. N. Mehrotra (&)

Department of Chemistry, JNV University,

Jodhpur 342 005, India

e-mail: mehrotra_rn@yahoo.co.in

123

Transition Met Chem (2008) 33:367–376

DOI 10.1007/s11243-007-9052-9

transfer in two steps as in the case of [CuIII(H2TeO6)2]5-

oxidation [13]. The results indicate an inner-sphere

mechanism, which is consistent with the behaviour of azide

ion as a ligand for the transition metals [32].

Experimental

Sodium azide (Fluka, purum) was crystallised once and its

solution was standardised with sulphatocerate(IV) solution

using ferroin as indicator. Though the azide solution is

stable over several days [16], fresh solutions were prepared

and used. The solution of NH4AuCl4 (Johnson Matthey)

was prepared in perchloric acid (e = 4.82 9 103 against

literature value = 4.86 9 103 dm3 mol-1 cm-1 [27(b),

33]), or in distilled water (e287 = 3,133 dm3 mol-1 cm-1).

LiClO4, prepared as described elsewhere [34] was used to

adjust the ionic strength. Mixtures of LiCl (Sigma) and

LiClO4 solutions were used to study the effect of [Cl-] on

the observed rate at constant ionic strength and [Li+].

[Ru(NH3)5(H2O)]2+ was prepared as described [35]. Twice

distilled water was used for the preparation of the solutions.

Rate measurements

The rates were measured, under pseudo first-order condi-

tions ([HN3] » [AuCl4-]) at constant ionic strength

(l, LiClO4), in terms of disappearance of Au(III) at

360 nm, at which Beers law is obeyed at different hydro-

gen ion concentrations and HN3 is transparent, using a

Spectrochem Mk (II) colorimeter. The pseudo first-order

rate constant kobs was calculated as 2.303 9 the slope of

the linear plots of log (At - A?) against time, which were

linear for more than two half-lives of the reaction. A? was

taken as zero because the spent reaction mixture was col-

ourless. Some representative plots are shown in Fig. 1. The

EXCEL program was used to draw the plots and the least-

square values of the slopes and intercepts. The reproduc-

ibility of the pseudo first-order rate constant, kobs, was

within ±5%. The program Act-e by H. Strahlow was used

to calculate the activation parameters.

Stoichiometry

A number of reaction mixtures (0.5 C [H+] C 0.1 mol

dm-3) with different initial concentrations of AuCl4- and

HN3 were prepared such that [AuCl4-] was always in

excess. The optical density of the reaction mixture was

measured at 360 nm, with the same colorimeter that was

used for the rate measurements, until it became constant.

The concentration of unreacted AuCl4- in all the reaction

mixtures indicated that D[AuCl4-]/D[HN3] = 0.51 ± 0.02,

which suggested that the stoichiometry of the reaction is

given by Eq. 1.

AuCl�4 þ 2HN3 �! 3N2 þ AuCl�2 þ 2Hþ þ 2Cl� ð1Þ

Characterisation of the reaction products

No quantitative measurements were carried out for the N2

evolved during the reaction though its formation was

qualitatively established. The reaction was carried out at

room temperature. Pure argon gas was passed through the

reaction mixture for sufficient time and the outgoing gas

was passed through a solution of [Ru(NH3)5(H2O)]2+ from

which a red complex, [(NH3)5Ru(N2)Ru(NH3)5]4+, was

isolated [13]. The Raman spectra of this complex shows a

strong m(N:N) band around 3,100 cm-1 [13] which is due

to the coordinated N2.

The spent reaction mixture was evaporated and the dry

residue was dissolved in nitric acid and evaporated again.

The process was repeated several times to ensure complete

decomposition of hydrazoic acid. The residue was finally

dissolved in the minimum volume of hot aqua regia. The

cooled solution on treatment with benzidine produced a

blue colour confirming the presence of Au(III) in the solu-

tion [36] and therefore AuCl2- as the product of the reaction.

Test for free radicals and spectroscopic studies

The solutions of AuCl4- and N3

- were purged with N2 and

to each solution was added 1–2 ml of acrylonitrile. There

appeared no cloudiness or precipitate over several minutes

0

0.15

0.3

0.45

0.6

0.75

0 100 200 300 400 500 600Time (s)

1 +

log

(At -

A∞)

Fig. 1 The plots of 1 + log (At - A?) against time under

experimental conditions: (d) 102[HN3] = 0.1, [H+] = 0.03,

l = 0.087 mol dm-3 at 29 �C; (�) 102[HN3] = 0.5, [H+]

= 0.1, l = 0.7 mol dm-3 at 15 �C; (h) 102[HN3] = 0.1, [H+]

= 0.015, l = 0.087 mol dm-3 at 29 �C

368 Transition Met Chem (2008) 33:367–376

123

before the solutions were mixed to initiate the reaction.

There resulted no cloudiness or thick precipitate in the

reaction mixture even after a prolonged time.

Since the reaction was fast, the reactant solutions of

desired concentrations were cooled in an ice–salt mixture for

sufficient time and so was the quartz cell of 1 cm path-length.

The reactant solutions were mixed and the reaction mixture

was transferred to the quartz cell in quick succession. The

cell was immediately returned to the cell compartment of a

Hewlett-Packard 8452A UV-vis diode array spectropho-

tometer1 having a band width of 2 nm over 350–500 nm.

The entire process was competed within 20–25 s and the

spectral measurement was triggered immediately. A red shift

in the kmax of AuCl4- solution from 310 to 327 nm indicated

the formation of a complex by HN3 with AuCl4- ions.

Fractional distribution of Au(III) and HN3 species

in acidic solution

The fraction of each species (=concentration of the species/

total analytical concentration of the reagent), calculated over

0.005 B [H+] B 0.3 mol dm-3 using KH = 4.17 9 10-5

[37], Khy = 1.6 9 10-5 [38] and Ka = 0.23 mol dm-3 [39]

for the equilibria (2)–(4) respectively, are shown in Fig. 2.

It is seen that the N3- fraction is almost negligible

compared to that of HN3 even at the lowest H+ concen-

tration. Hence, HN3 is considered as the reactive entity of

azide and consequently this equilibrium has nothing to do

with the rate retardation by H+ ion.

The fractional distribution of Au(III) species indicates

that all the three species are present in concentrations hard

to ignore. The concentration2 of AuCl4- increases and that

of AuCl3(OH)- decreases approaching that of AuCl3(OH2)

which increases continuously with [H+] and finally exceeds

that of AuCl3(OH- around [H+] % 0.25 mol dm-3. The

consequences of considering one or a combination of these

species are discussed below.

Results

The plots of kobs against [HN3], Table 1, Fig. 3, are curves

and do not show any sign of saturation at high [HN3] within

the concentration range used. The plots of kobs-1 against

[HN3]-1, shown in Fig. 4, are linear with intercepts on the

rate ordinate within the range of concentrations used. These

Michaelis–Menten plots indicate that Au(III) and hydrazoic

acid complex(es) intervene in the mechanism. The forma-

tion of the complex(es) is supported by the red shift in the

spectrum of Au(III) solution on adding hydrazoic acid as

described above.

To summarise our experimental observations, the kobs

decreased with increasing [H+], Table 2, and the plots of

kobs-1 against [H+], Fig. 5, are linear with positive intercepts on

the rate ordinate. The dependence of kobs on [Cl-], Table 3,

indicates that kobs decreases with increasing [Cl-] and the

plot of kobs-1 against [Cl-], Fig. 6, is linear with positive

intercept on the rate ordinate. However, the kobs is inde-

pendent of the ionic strength, Table 4, indicating that the

reaction partners are either molecular species or an ion and a

molecular species. The present reaction in this respect is

different from the similar oxidation with 12-tungstocobalt

(III) ion [11].

0

0.2

0.4

0.6

0.8

1

0 0.1 0.2 0.3

[H+] (mol dm-3)

frac

tion

of th

e sp

ecie

s

Fig. 2 The plot of the fraction of AuCl4-(D), AuCl3(OH2) (�) and

AuCl3(OH)- ion (d), (m) N3- and (9) HN3 against [H+] at 25 �C for

the initial concentration of gold(III)-complex = 1 9 10-4 mol dm-3

and that of azide = 0.035 mol dm-3

0

5

10

15

20

25

30

35

0 0.01 0.02 0.03 0.04

[HN3] (mol dm-3)

103 k ob

s (s-1

)

Fig. 3 The non-linear plots of kobs against [HN3] at temperatures 25

(d), 30 (h), 35 (D) and 40 �C (�) 104 [Au(III)] = 1.0, [HClO4] =

1.0 and I = 1.5 M

1 The instrument having no facility for maintaining the constant

temperature was not used for the kinetic study.

2 In making the calculations [Cl-] equal to the initial gold(III)

complex is assumed to be present.

Transition Met Chem (2008) 33:367–376 369

123

Mechanism and discussion

Chloroauric acid, HAuCl4, like manganic acid, HMnO4, is

a strong acid and is, therefore, completely ionised even in

dilute solutions of strong mineral acids. AuCl4- ion,

however, exists in equilibrium with other Au(III) species

as shown in equilibria (3) and (4). The consequence of

considering AuCl4- ion as the sole reactive species would

require the reaction to be independent of H+ and Cl- ions,

which is contrary to our observations. The equilibrium (3)

is helpful in explaining the rate retardation by Cl- ions on

the assumption that AuCl4- ion is less reactive than

AuCl3(H2O). Similarly, the equilibrium (4) can be instru-

mental to explain the retardation of the rate by H+ ions by

considering AuCl3(OH-) to be more reactive than

AuCl3(H2O). In conclusion, the reactivity of Au(III) species

is in the order: AuCl3(OH)- [ AuCl3(H2O) [ AuCl4-. The

0

200

400

600

800

1000

1200

1400

1600

1800

0 200 400 600 800 1000 1200

[HN3]-1 (dm3 mol-1)

k obs-1

(s)

Fig. 4 The linear plots of kobs-1 against [HN3]-1 at temperatures at

temperatures 25 (e), 30 (h), 35 (D) and 40 �C (�). 104 [Au(III)] =

1.0; [HClO4] = 1.0 and l = 1.5 mol dm-3

Table 1 Dependence of kobs on the initial [HN3] at different tem-

peratures 104 [AuCl4-] = 1.0, [H+] = 0.18 and l = 0.215 mol dm-3

[HN3] (mol dm-3) 0.001 0.005 0.010 0.015 0.020 0.030 0.035

103kobs

(s-1, 25 �C)

0.566 2.52 4.42 5.92 7.14 8.93 9.62

103kobs

(s-1, 30 �C)

0.951 4.15 7.19 9.43 11.2 13.9 14.9

103kobs

(s-1, 35 �C)

1.54 6.54 11.0 14.3 16.7 20.2 21.5

103kobs

(s-1, 40 �C)

2.48 10.3 17.0 21.6 25.0 29.7 31.1

Table 2 Dependence of kobs on [H+] at different temperatures 104

[AuCl4-] = 1.0, [HN3] = 0.001 and l = 0.037 mol dm-3

[H+] (mol dm-3) 0.001 0.005 0.01 0.015 0.020 0.030 0.035

103kobs

(s-1, 17 �C)

12.1 5.68 3.68 2.91 2.30 1.59 1.37

103kobs

(s-1, 21 �C)

14.8 7.39 4.96 3.83 3.05 2.14 1.92

103kobs

(s-1, 25 �C)

19.2 9.59 6.92 4.63 4.09 2.99 2.59

103kobs

(s-1, 29 �C)

23.5 12.1 8.10 6.17 4.93 3.71 3.14

103kobs

(s-1, 33 �C)

30.2 14.9 10.0 7.84 6.44 4.54 3.85

0

100

200

300

400

500

600

700

800

k obs-1

(s)

0 0.01 0.02 0.03 0.04

[H+] (mol dm-3)

Fig. 5 The plots of kobs-1 against [H+] at temperatures 17 (�), 21 (h),

25 (d), 29 (D) and 33 �C (j) 104 [AuCl4-] = 1.0; 103 [HN3] 1.0 and

l = 0.1 mol dm-3

2

2.5

3

3.5

4

4.5

5

5.5

0 0.05 0.1 0.15 0.2 0.25 0.3 0.35 0.4

[Cl-] (mol dm-3)

10-3 k

obs -1

(s)

Fig. 6 The linear plot of kobs-1 against [Cl-] at 25 �C. 104 [AuCl4

-] = 1.0;

103 [HN3] 1.0 and l = 0.1 mol dm-3

370 Transition Met Chem (2008) 33:367–376

123

rational consequence of AuCl4- being the least reactive is to

assume it to be inactive. The AuCl3(H2O) is also unlikely to

be the reactive species of choice since H2O is a stronger

nucleophile than OH- ion. Therefore, the substitution of

H2O in AuCl3(H2O) by HN3 is unlikely. Thus, AuCl3(OH)-

is the most likely reactive species.

The reaction is first-order in Au(III)-complex, fractional

order in HN3, and each molecule of AuCl4- consumes two

molecules of hydrazine. The formation of an intermediate

complex between Au(III)-complex and HN3 has spectral

evidence, the rate is independent of the ionic strength and

the free radical test is negative. This leads us to the sequence

of reactions in (2)–(8) (Scheme 1) as part of the most likely

mechanism in the absence of added Cl-. The mechanism in

the presence of Cl- is different as discussed later.

KH

HN3 H+ + N3– (2)

Khy

AuCl4– + H2O AuCl3(H2O) + Cl– (3)

Ka

AuCl3(H2O) AuCl3(OH)– + H+ (4)

K

AuCl3(OH)– + HN3 AuCl3(H2O)N3– (5)

K1

AuCl3(H2O) + N3– AuCl3(H2O)N3

– (6)

kAuCl3(H2O)N3

– AuCl3(H2O)2– + N3+ (7)

fastN3

+ + N3– 3N2 (8)

Scheme 1

Consistent with the known substitution of heterocyclic

amines L in AuLCl3 by N3- and the nucleophilic behaviour

of N3- in kinetics with Pt(II)-complexes [40], the formation

of an intermediate through the reaction (5) or its alternate

(6) with the rate determining step (7) involving two-elec-

tron transfer within the intermediate in a single step

resulting in the formation of N3+ ion is suggested.

The alternate reactions (5) and (6) are not kinetically

distinguishable because of the identity3 KKa = K1KH.

However, reaction (5) is likely to have precedence over

reaction (6) because it involves the fast process of

neutralisation of the OH- ion by the proton of hydrazoic

acid itself. The formation of such an intermediate has been

reported in the oxidation by [Ag(OH)4]- [12], [CuIII

(H2TeO6)2]5- and [AgIII(H2TeO6)2]5- [13] and [(CH2)2

(C2N5 H6)2Ag]3+ ions [14].

The rapid formation of a cyclic hexazine, which has

been supported by the spectral evidence in the decompo-

sition of cis-Pt(N3)2(PPh3)2 [41] through the reaction of N3+

with N3- is assumed in the oxidations by [Ag(OH)4]- [12],

[CuIII(H2TeO6)2]5- and [AgIII(H2TeO6)2]5- ions [13]. The

intermediate hexazine decomposes to N2 in the fast step.

Reaction (8) is thus consistent with the stoichiometry of the

reaction.

The proposed mechanism is similar to those proposed

in the similar oxidations by [CuIII(H2TeO6)2]5- and

[AgIII(H2TeO6)2]5- ions [13] and ethylenebisbiguanide-

Ag(III) [14] but is noticeably different from that by

[Ag(OH)4]- ion [12]. It might be mentioned that the plots in

Fig. 3 are curved towards the concentration axis whereas

similar plots in the oxidation by [Ag(OH)4]-, which has a

second-order dependence in N3-, were curved towards the

rate ordinate [12]. Two kinetically indistinguishable paths

were suggested to account for the second-order dependence

in N3-. In one path, the redox occurs within a five-coordinate

intermediate formed by azide attack on Ag(OH)3N3- and its

conjugate acid. The formation of both cis- and trans-diazido

complexes, Ag(OH)2(N3)2-, constituted the second path and

there was OH--catalysed transformation of the trans-

complex to cis-complex before redox occurred. Since the

ligands, tellurate in the Ag(III) and Cu(III) complexes [13],

ethylenebisbiguanide in the other Ag(III) complex [14] and

Cl- in AuCl4-, are not labile like OH- in Ag(OH)4

- and

there is no second-order dependence in N3-, therefore cis-

and trans-complexes have no existence in these oxidations.

The rate of the reaction deduced from the reactions

(2)–(8) is expressed by Eq. 9

Table 3 Dependence of kobs on [Cl-] at 25 �C 104 [AuCl4-] = 1.0,

[N3-] = 0.001, [H+] = 0.1 and l = 0.687 mol dm-3

LiClO4 (mol dm-3) 0.586 0.500 0.400 0.300 0.200

LiCl (mol dm-3) 0.0 0.086 0.186 0.286 0.386

103kobs (s-1) 2.11 0.381 0.284 0.219 0.186

Table 4 The effect of the ionic strength on the rate of the reaction at

20 �C 104 [AuCl4-] = 1.0, [HN3] = 0.01, [H+] = 0.1 mol dm-3

[LiClO4] (mol dm-3) 0.014 0.064 0.164 0.264 0.364 0.464

l (mol dm-3) 0.114 0.164 0.264 0.364 0.464 0.564

103 kobs (s-1) 8.33 8.20 8.20 8.24 8.43 8.39

3

½AuCl3ðH2OÞN�3 � ¼KKa½AuCl3ðH2OÞ�½HN3�

½Hþ�¼ K1KH½AuCl3ðH2OÞ�½HN3�

½Hþ�

Transition Met Chem (2008) 33:367–376 371

123

�d½AuIII�odt

¼ kKKa½AuCl3ðH2OÞ�½HN3�½Hþ� ð9Þ

The mass balance for the species N3- and [AuCl3(H2O)]

in terms of the distribution of the respective species given

in Eqs. 2 and 3–5 is expressed by the Eqs. 10 and 11,

respectively.

½N�3 � ¼KH½HN3�oKH þ ½Hþ�

ð10Þ

½AuCl3ðH2OÞ�

¼ Khy½Hþ�½AuIII�oKhy½Hþ� þ KaKhy þ ½Hþ�½Cl�� þ K1Khy½Hþ�½N�3 �

� �

ð11Þ

Substituting the values of [N3-] and [AuCl3(H2O)] from

Eqs. 11 and 12 into Eq. 9 gives the Eq. 12.

Since [H+] » KH, Eq. 12 is reduced to Eq. 13.

�d½AuIII�odt

¼ kKKaKhy½AuIII�o½HN3�oKhy½Hþ�þKaKhyþ½Hþ�½Cl��þKKaKhy½HN3�o

ð13Þ

Reaction in the absence of Cl- ions

In the absence of Cl- ions, [H+][Cl-] is neglected and so is

the product KaKhy because of the low value of Khy. The

Eq. 13 is therefore modified to Eq. 14, the inverted form of

which is given in Eq. 15, which is consistent with the linear

plots of kobs-1 against [HN3]-1, Fig. 4, and that of kobs

-1

against [H+] (Fig. 5).

kobs ¼�d½AuIII�odt½AuIII�o

¼ kKKa½HN3�o½Hþ� þ KKa½HN3�o

ð14Þ

1

kobs

¼ ½Hþ�kKKa½HN3�

þ 1

kð15Þ

The close agreement between the k values at 25 �C,

0.0182 s-1 from the plot in Fig. 4 and 0.0190 s-1 from the

plot in Fig. 5, is a strong support for the proposed

mechanism. The so-calculated k values at different

temperatures are reported in Table 5 and the linear plot of

log k against 1/T is shown in Fig. 7. The activation

parameters DH� and DS� are also reported there. The

relatively small value of DH� could be attributed to the

formation of an intermediate in the equilibrium (5).

The formation of an axial bond between the positive AuIII

and N3- helps to stabilise the intermediate.

It is apparent from Eq. 15 that I/S = KKa/[H+] where I

and S are the intercept and slope of the plots in Fig. 4.

Therefore K is given by Eq. 16

K ¼ ½Hþ� � I

Kað¼ 0:23Þ � Sð16Þ

Using Eq. 16 and substituting the values of I = 54.8 s,

S = 1.7 mol dm-3 s, [H+] = 0.18 mol dm-3 and Ka =

0.23 mol dm-3, a value of K = 25 dm3 mol-1 at 25 �C is

obtained. Thus the value of K is within the range (0.1–100)

of formation constants of azido complexes [6(a)–8, 12, 13].

The values at other temperatures are not obtainable because

the values of Ka at other temperatures are not known.

The intercept on the rate ordinate of Fig. 3 is an indi-

cation that the hypothetical rate at infinite concentration of

HN3 has a finite value. This demands that the plot falls off

asymptotically to a constant kobs at high [HN3]. This situ-

ation in terms of Eq. 14 is achieved when

KKa[HN3] » [H+] so that the Eq. 14 changes to Eq. 17.

kobs ¼ k ð17Þ

The least value of [HN3] required under which the

inequality KKa[HN3] » [H+] holds good is calculated from

the simple relation KKa[HN3] = 10[H+]. For [H+] =

0.18 mol dm-3 and the known value of Ka = 0.23 mol

dm-3 [39] since K has a calculated value of 25 dm3 mol-1

(see above) at 25 �C, the minimum HN3 concentration is

calculated to be 0.31 mol dm-3 which is about nine times the

maximum concentration of HN3 used in Table 1.

The N3- ion is considered to coordinate axially because a

negative ion is generally preferably coordinated to the axial

position of a metal. Such coordination has been proposed in

the oxidations of hypophosphate [42], iodide [43] and

mandelate [44] by the square planar Ag(OH)4- complex.

The formation of a similar axial intermediate complex has

been proposed in the oxidation of histidine [20], SO32- ion

[45] and oxalic acid [18] by AuCl4- ion. This possibility is

most likely considering that AuCl3(OH)- ion is also

square-planar like Ag(OH)4- ion and that AuIII and AgIII are

�d½AuIII�odt

¼ kK1Khy½Hþ�½AuIII�oKhy½Hþ� þ KaKhy þ ½Hþ�½Cl�� þ K1Khy½Hþ�KH½HN3�o

KHþ½Hþ�

� �� KH½HN3�oKH þ ½Hþ�

ð12Þ

372 Transition Met Chem (2008) 33:367–376

123

isoelectronic. The graphical representation of the formation

of the axial complex and the electron transfer within this

complex are illustrated in Eqs. 18 and 19, respectively.

Equations (18) and (19) correspond to Eqs. 5 and 7,

respectively. The product side of Eq. 19 could be suitably

modified to correspond to Eq. 20.

It has been pointed out that the activated pathway for

ligand substitution in four-coordinate planar d8 metal

complexes requires a five-coordinate transition state [46].

The six-coordinate species are unknown in solution though

these are known in the solid state [46]. The processes are

likely to require a high degree of orientation in the tran-

sition state as indicated by the large negative DS� value

[46].

The negative free radical test does not conclusively

eliminate the possibility of the formation of such species

which may decay rapidly before reaching the bulk sol-

vent for their detection. Thus the possibility of

successive one-electron transfer cannot be effectively

ruled out in view of the fact that Au(II) has been

reported as an intermediate in several redox reactions

[47–52], and its dithiolene and dicarbolyl compounds are

known [53].

Accordingly, Scheme 1 can be suitably modified by

replacing (7)–(8) by (20)–(22) (Scheme 2), the rate law

will have the same form except that k is replaced by

k1

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

3 3.1 3.2 3.3 3.4 3.5 3.6

103/T

2 +

log

k

Fig. 7 The linear plot of log k against 1/T where k is the rate

determining rate constant

Table 5 The k (s-1) values obtained from the plots in Figs. 4 (a) and 5 (b) at different temperatures with associated activation parameters

Temp (�C) 17 21 25 25 29 30 33 35 40

102k 1.37(b) 1.57(b) 1.82(a) 1.90(b) 2.42(b) 2.62(a) 3.23(b) 3.47(a) 4.78(a)

DH� = 40 ± 2 kJ mol-1

DS� = -126 ± 7 J K-1 mol-1

Au

ClCl

Cl

Au

ClCl

Cl−

+ HN N NK

OHOH 2

N

N

N −

(18)

Au

ClCl

Cl OH2

N

N

N −

kAuCl3(H 2 O)2− + N3

+(19)

Transition Met Chem (2008) 33:367–376 373

123

k1

AuCl3(H2O)N3– AuCl3(H2O) – + N3

• (20)

fast( k )′

AuCl3(H2O)– + N3– AuCl3(H2O)2– + N3

• (21)

fast( k )2N3

• 3N2 (22)

Scheme 2

The reaction (21) is proposed because the oxidation of

N3• radical either by Au(III) or Au(II) in the fast step

would violate the observed stoichiometry of the reaction.

However, on the other hand N3• happens to be a strong

oxidant, because the potential of N3-/N3

• couple is 1.33 V

[16], therefore it is likely that Au(II) may be oxidised back

to Au(III). Ideally, the measurement of the rate constant

for the reaction AuCl3- + N3

• would have provided the

right answer. Since this is not easily feasible4 the fol-

lowing considerations were helpful to conclude against the

possible feasibility of the reversibility of reaction (21).

The relatively high value of rate constant for the decay of

N3• radical to N2, 4.5 9 109 dm3 mol-1 s-1 [16] over that for

the disproportionation of Au(II), 1.4 9 109 dm3 mol-1 s-1,

generated from AuCl4- in the absence of Cl- ion in aque-

ous solution at pH 2, suggests that the build up of the

necessary concentration of N3• radical required to make the

reaction (21) reversible is most unlikely. Further, if the reac-

tion (23) happens to be reversible, then the usual first-order

plots for the disappearance of Au(III) would show a definite

deviation from linearity. Since this is not the case as shown by

the plots in Fig. 1, therefore the reversibility of the reaction

(21) is highly improbable.

Consideration of the potentials of Au+/Au++,\1.29 V vs

NHE. [54] and N3-/N3

• (1.33 V [16]) which are similar,

suggests that N3• can in no way oxidise Au++ ion to any

appreciable extent. For the reversibility of the reaction

(21), it is important that there is a necessary build up of N3•

concentration, which is highly unlikely in view of the high

rate of decay of N3• radical to N2 [16].

N�3 þ N�3 �fast

N��6 ð23Þ

The reaction (23) is another path for the decay of the N3•

radical into N2. The consideration of the equilibrium

constant, 0.33 dm3 mol-1 [55] suggests that it will be

a minor species, and since the forward rate constant,

1 9 106 dm3 mol-1 s-1 [16] is about 103 times smaller

than the rate constant k00 therefore it would be unable to

compete with the reaction (22).

Mechanism in the presence of Cl- ions

For the reaction in the presence of Cl-, the term [H+][Cl-]

is no more negligible, KaKhy is neglected as before and

Khy[H+] « [H+][Cl-], the Eq. 13 would change to Eq. 24

the inverted form of which is given in Eq. 25.

kobs ¼�d½AuIII�dt½AuIII�o

¼ kK1KHKhy½HN3�o½Hþ�½Cl�� þ K1KHKhy½HN3�o

ð24Þ

1

kobs

¼ ½Hþ�½Cl��kKKaKhy½HN3�

þ 1

kð25Þ

Equation (25) predicts that the intercept values of the plots

of kobs-1 against [X] ([X = H+ or Cl-]) should be the same.

However, the intercept derived from Fig. 4 (X = H+,

intercept = 52.7 s; k = 0.019 s-1) differed widely from

that obtained from Fig. 5 (X = Cl-, intercept = 1.82 s,

k = 0.55 s-1). The k value derived from Fig. 5 is about 30

times the value obtained from Fig. 4. This clearly indicates

that either the reactive species or the mechanism or both

are different in the presence and absence of chloride ions.

Although no systematic study on the effect of [HN3] on

the observed rate constant in the presence of chloride ions

was carried out, an indication of an outer-sphere mecha-

nism was provided by the rate constant which increased

almost by a factor of two on increasing the concentration of

hydrazoic acid by the same factor. The change from the

inner-sphere to outer-sphere mechanism, in the presence of

excess Cl- ions, is not caused by any change in the Au(III)

species considered in equilibria (3) and (4) because the

coordination shells in these species are full. The excess

Cl- ions may push the equilibrium (3) to the far left

affecting the equilibrium (4) to the extent that AuCl3(H2O)

and AuCl3(OH)- become negligible minor species leaving

AuCl4- as the only bulk species. The equilibrium (2),

however, remains unaffected.

The Cl- ions, in addition to affecting the equilibria (3)

and (4) as described above have another role in the system.

The excess chloride ions form a cage around hydronium ion

surrounded AuCl4-. The caged AuCl4

- species, ½AuCl�4 �Cl� ,

and AuCl4- are in equilibrium as shown in reaction (26). The

N3- and ½AuCl�4 �Cl� are the redox partners, and the reaction

is initiated by the penetration of the chloride cage by N3-,

which faces repulsive forces. Thus, the penetration is a slow

process or the rate determining step. The attack of AuCl4- by

N3- inside the cage is fast and may involve either a net two-

electron change with the formation of N3+ followed by the

reaction (8) as considered in the Scheme 1 or single-electron

transfer in two successive steps resulting in the formation of

N3• radical as considered in Scheme 2. The details of the

reactions in the latter possibility are considered below

where ½AuCl�4 ;N�3 �Cl� represents the presence of N3

- inside

the cage.

4 The rates reported for the corresponding reaction with IrCl63- in

[11] were measured at the Centre for Fast Kinetics Research at the

University of Texas at Austin.

374 Transition Met Chem (2008) 33:367–376

123

½AuCl�4 � þ Cl� �KCl ½AuCl�4 �Cl� ð26Þ

½AuCl�4 �Cl� þ N�3 �!k2 ½AuCl�4 ;N�3 �Cl�

chloride cage chloride cageð27Þ

½AuCl�4 ;N�3 �Cl� �!

fastAu(II)þ N�3 ð28Þ

Au(II)þ N�3 �!fast

Au(I)þ N�3 ð29Þ

The nature of Au(II) and Au(I) species in the present

context is not properly understood. The reaction (27) is rate

limiting, and the reactions (28) and (29) are followed by

the reaction (22) for the decay of N3•. The rate of the

reaction is expressed in Eq. 30.

�d½AuIII�odt

¼ k2½AuCl�4 �Cl� ½N�3 � ð30Þ

The proper substitution of the values of ½AuCl�4 �Cl� from

the consideration of equilibrium (26) and that of [HN3]o

given in Eq. 10 into Eq. 30 would change it to Eq. 31, the

inverted form of which is given in Eq. 32.

�d½AuIII�odt½AuIII�o

¼ kobs ¼k2KH½HN3�

½Hþ� þ KCl½Hþ�½Cl�� ð31Þ

1

kobs

¼ ½Hþ�k2KH½HN3�

þ KCl½Hþ�½Cl��k2KH½HN3�

ð32Þ

Equation (32) is consistent with Fig. 6 from which k2 =

0.546 dm3 mol-1 s-1 and KCl = 5 dm3 mol-1 are obtained

at 25 �C.

Conclusion

In conclusion, the oxidation of hydrazoic acid, in the

absence of added chloride ions, by tetrachloroaurate(IIII)

ion is inner-sphere in which hydrazoic acid is axially

coordinated to Au(III) ion, the free radicals N3• do not

oxidise Au(II) and the H+ ions retard the kobs. In presence

of added Cl- ions, the hydronium ion surrounded AuCl4- is

caged by chloride ions and the reaction is outer-sphere. The

penetration of the chloride cage by N3- is rate controlling.

The electron transfer within the cage is fast.

Acknowledgements Thanks are due to the UGC, F.12-59/1997 and

F.12-147/2001, for the financial support of the work. Thanks are also

due to Kriti Mehrotra, my grand daughter, at Cornell University, in

helping with certain references.

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