Kinetics and mechanism of oxidation of

Kinetics and mechanism of oxidation of hydroxylamine by hexachloroiridate(1V) ion in buffer solutions M s . REKHA,ADITYAPRAKASH,AND RAJ N . MEHROTRA' Department of C h e ~ n i s t r Universit)~ ~~, ofJodhpl(r, Jodhp~lr342 005, India

Can. J. Chem. Downloaded from www.nrcresearchpress.com by 14.139.230.12 on 03/20/14 For personal use only.

Received June 23, 1993 Ms. REKHA, ADITYA PRAKASH, and RAJ N. MEHROTRA. Can. J. Chem. 71, 2164 (1993). The oxidation of hydroxylammonium ion by [IrCl,]' ion in acetic acid - acetate buffer solutions, studied by stopped ~ ~I H . The ] oxidation involves the species [IrCI,]'- and NH,OH+ flow, has the stoichiometric ratio A [ I ~ C ~ ~ ] ' - / A [ N H = although spectral analysis of the spent reaction mixture indicates [IrC15(OHz)]' to be the main product (almost to the . rate is retarded both extent of 80%). This anomaly arises because of the aquation of the reduced product [ I r ~ l ~ ] ' -The by H t and C 1 ions and the plots of k;' against respective concentrations are linear. The proposed mechanism is given by reactions [i]-[v]. Kipd

[i]

NH30H+C1-

NH,OH+

+ CI

. k1

+ N H ~ O H 'z[1rCI6(NH20H)]'- + H+

[ii]

1r~1,'-

[iii]

[IrCI6. NH,OH]?-

[iv]

[ ~ r ~ l i. -. NH?OH'+] .

k- I kc,

[ I ~ c I , ' . . .NH,OH"]

k-CI

-

kd z [1rC16]' + NH,OH" k-d

fast ~NH?OH'+ N2 2 H 2 0 + 2H' [v] The values of the rate constants at 25OC are as follows: kl = 147 dm3 mol-' s-' and 10'k,,~,,= 2.8 s-I. The related activation parameters are AH:, = 29 2 kJ mol-I, and AS;, = - 115 2 6 J K - I mol-' and, AH:=,,,>, = 15 k 2 kJ mol-' and = -233 i 3 J K-' inol-', respectively. The value of Kip, is 2.91 k 0.03 mol dm-' and that of Kip;, (= l/Kipd)is 0.344 0.004 dm3 mol-'; both values are almost independent of temperature.

+

*

*

ADITYA PRAKASH et RAJN. MEHROTRA. Can. J. Chem. 71, 2164 ( 1993). Ms. REKHA, L'oxydation de I'ion hydroxyammonium par l'ion [ I C I ~ ] 'dans des solutions tamponnCes d7acCtateet d'acide acCtique a CtC CtudiCc par la technique du flux stoppC a un rapport stoechiomCtrique A [I~CI,]'-/A[NH~OH]= I. L'oxydation ilnplique les espkces [1rCI6]' et NH'OH' m&me si I'analyse spectrale du melange rkactionnel residue1 indique que le produit principal de la reaction (pratiquement 80%) est I'ion [I~CI,(OH,)]'-. Cette anomalie est due h l'aquation du produit rCduit, le [ I ~ c I ~ ] ' -La . vitesse de la rkaction est retardee par les ions H ' ainsi que CI- et les courbe de k2-l en fonction des concentrations sont lineaires. Le mecanisme proposCs est represent6 par les Cquations (i)-(v) : r , .

Kipr~

[i]

NH,OH+CI-

Z N H 3 0 ~ ++ C1

kI 1rc16?- + NH,OH+ S[ I ~ c I ~ ( N H ~ O H+ ) ]H ' ' k- I kc, [iii] [IrC16. NH?OH]'[1rCl6'-. . .NH20H't] k-c, kd [iv] [1r~16)-.. . NH~OH"] [1rCln13-+ NH30He+ k-d rapide ~NH?oH'+ Nz + 2 H 2 0 + 2H+ [v] Les valeurs des constantes de vitesse a 25"CAsont: kl = 147 dm3 mol-' s ' et 10'k,,~~,, = 2.8 s-'. Les parametres = 15 d'activation apparent& sont respectiveinent AH;, = 29 k 2 kJ moll et AS:, = - 115 i 6 J K ' mol-' et 2 kJ moll et AS^,,,,, = -233 ? 3 J K ' mol-'. La valeur de Kip,,est Cgale i 2,91 k 0,03 mol dm-%t celle de K,,;, (= l/Kipd)est tgale h 0,344 2 0,004 mol dm-3; les deux valeurs sont pratiquement independantes de la temperature. [Traduit par la redaction] [ii]

-

*

Introduction Hydroxylamine is a n intermediate in the enzymatic oxidation of ammonia to nitrite and (or) nitrate o r in the reverse reduction of these ions to ammonia (1). This and the fact that it acts both as a n oxidizing and reducing agent with ' ~ u t h o to r whom correspondence may be addressed. Reprints are not available.

transition metals a n d their complexes has stimulated a keen interest in investigating its redox and ligand properties (2). Amongst several oxidations, the oxidation by hexachloroiridate(IV) ion in the presence of ~ e r c h l o r i cacid has been studied (3). The rate, studied in low [HCIO,] (0.01-0.15 mol dm-3), showed an inverse dependence on [H'], and k,,, [H'l was constant. T h e simple two-step mechanism considered the formation of a n [IrC16(NH,0H)12 complex in an equi-


REKHA ET A L .

librium step prior to its disproportionation to [ I r c l , ] ' and NHOH' free radical in the rate-determining step. The rate was found to become complex at [HCIO,] > 0 . 1 5 rnol dm-' and the reason w a s attributed to the probable formation of [HIrC16]-. However, [I~cI,]'- is considered to have a negligible base strength and is therefore unlikely to b e protonated (4). T h e interestin?& feature w a s the detection and estimation of [IrClj(OH2)]- ( < l o % ) in the spent reaction mixtures but no explanation for its probable formation w a s given. T h e redox chemistry o f NH,OH with transition metals in solution is complicated. A variety of oxidation products, viz., N2, N,O, N O , NO,, NO,-, o r NO,-, depending o n the fate of the H z N O . radical that undergoes reactions [1]-[3], are formed under different conditions (2). Reaction [3] is the dimerization of H N O . to hyponitrous acid ( 3 , which exists in cis and trans forms. T h e trans form, m o r e stable than the cis form, decomposes into N 2 0 (6). I n the case of strong q ) ~ ~ " ( a q( 7) ) , the free radioxidants such as ~ n ~ + ( a and cal is oxidized to NO,- by excess oxidant. T h e formation of the H,NO. radical in oxidations by one-electron oxidants has been demonstrated b y e s r spectroscopy (8). [I]

N H 2 0 . + 0.5N2 + HZ0

[2]

H'NO

+ HNO

+ H+ + e-

This study was, therefore, undertaken in the p H range 3.42-5.23 (acetic acid - sodium acetate buffer) t o check if NH,OH, which could b e easily produced a s a result of increased dissociation of NH,OH+ at such low [H'], happened to b e reactive and whether [IrCIs(OHZ)]- is the reactive Ir(1V) entity in view of the fact that [I~c~,(oH,)]~- w a s reported to b e o n e of the oxidation products (3). W e were further inclined to investigate whether the reaction proceeded through a n Ir(1V)-hydroxylamine complex, and the s a m e could b e substantiated by a rapid-scan technique that might help to characterize its character as an outer o r inner sphere. Preliminary studies indicated that the rates are retarded b y H+ and C1- ions; retardation by the latter ion w a s observed for the first time in an Ir(IV) oxidation although an acceleration of the rate in hydrazine oxidations b y Ir(1V) (9) and [F~(cN),]'- (10) was reported without any explanation being given.

Fresh solutions of Ir(IV), using either H,IrC16 (Soekawa Chemicals) or Na,IrC16. 6 H 2 0 (JM), were standardlzed spectrophotometrically (Shimadzu 240 Graphicord) using €488 = 4050 dm3 mol-I cm-I (1 1) and were used within a few hours. Solutions of hexachloroiridate(II1) (JM) were similarly prepared and standardlzed using ejS8 = 74 dm3 mol-I cm-I) (12). However, we could more closely reproduce the E,,, literature value of 2080 dm7 mol-' cm-' (12), more often, in a given sample in which eAsxshowed considerable variation from the literature value of 4075 2 25 dm3 mol-' (see Spectrophotometry section). Solutions of hydroxylamine hydrochloride (E. Merck, puriss), prepared in distilled water purged with nitrogen, were standardized bromometrically (13). Solutions of lithium and sodium perby neutralizing lithium hydroxide (G.F. chlorate were Smith) and sodium carbonate (E. Merck, GR) by perchloric acid. The solutions were cooled, the pH was adjusted to 7 , and the concentration was determined by evaporating the water from weighed aliquots and drying to a constant weight. The buffer solutions of

2165

desired pH values were prepared by mixing solutions of sodium acetate (E. Merck, GR) and E. Merck's GR acetic acid (14) ([CH3COO-] = 0.06 rnol dm-') and the pH was checked with a pH meter. While adjusting the ionic strength (I), CH3COONa was assumed to be completely ionized. One of the reviewers has drawn our attention to the catalysis by trace levels of Cu2+ion (15); moreover, addition of Fez+ at the level of lo-, rnol dm-3 is stated to accelerate the rates by more than a factor of 1000. Although the samples were not purified and were used as received, we feel in the light of observed linear correlations between the respective rate and concentration that these samples did not contain any Cu2+ or ~ e ions ~ as + impurity, at least at levels that could affect the rates. Twice-distilled water, once from alkaline permanganate from an all-glass still, was purged with nitrogen before use. Spectrophotometry The uv-visible spectra of the various solutions, described below, were recorded using an HP8452 diode array spectrophotometer having a band width of 2 nm and using an integration time of I s. The reference in each case was a water solution having the other constituents of the same concentration. Although the spectra were recorded between 190 and 820 nm, the spectra in each of the various figures are shown over the relevant range of wavelength. The [IrC161Z-sample was about 3-4 months old. A 25 mL stock solution of Ir(IV) was prepared by weighing 0.044 g (0.003 14 rnol dm-3) of a Na,IrCl,. 6 H 2 0 (JM) sample in [HCIO,] = I rnol dm-3. The spectrum of the diluted solution (1.57 x lo-, rnol dm-') in 3 rnol dm-3 perchloric acid showed an absorbance of 0.497 and 0.263 at 488 (A,;,,) and 460 (A,,,) nm, respectively, from which E488 = 3163 and E460 = 1678 dm3 mol-' cm-' were calculated. These values, smaller than those obtained for a solution similarly diluted with chlorine-saturated perchloric acid of the same concentration (see below), indicate that the sample had some [IrCl6I3- as the impurity. The fact that the characteristic peak of [IrC16j3- at 358 nm was not observed could well be due to its relatively small concentration present as impurity and to the sinall value of E',~,74 dm' mol-I cm-' (12b). The absorbance of the solution deteriorated slightly after 24 h. The visible spectrum of a similarly diluted stock solution in chlorine-saturated perchloric acid (3 rnol dm-3) is shown in Fig. I . The reference solution was 3 rnol dm-3 perchloric acid saturated with chlorine. The absorbance at 488 and 460 nm was 0.554 and 0.328, respectively, iving E,,, = 3525 against 4050 and E460 = 2086 against 2080 dmg nlol-I c m ' (120). The increase in the absorbance at both wave lengths could be ascribed to the oxidation of [Irc1613-, present as an impurity, by chlorine. The extent of [IrC1613- present as impurity, calculated on the basis of the amount of the sample weighed and calculated on the basis of the €488 value, is about 13%. The E, = 2086 dm3 mol-' cm-', closer to the literature value, was obtained more often, whereas a significant difference in the €488 value existed in the same sample, fo;which we have no explanation at the moment. One of the reviewers felt that hydrolysis of Ir(IV) chloride could be the probable reason for the discrepancy. Although it was not investigated by us, the fact that the acidic [IrCl,]'- solutions were spectrophotometrically shown to be stable for at least 7 days (4) negates such a possibility. Next, the reaction mixture ( l O ~ I r ~ l , " ]= 3.14, IO"NH~OH+] = 4, and [HCIO,] = 1 rnol dm-') was left overnight for completion as the reaction was slow at this [H+] value. The spectrum of a twice-diluted reaction mixture in [HClO,] = 3 rnol dm-3 had a peak at 342 nm and an absorbance = 0.0227. The peak probably corresponds to [ I ~ C I ~ ( O H ~ ) ]which '-, has a A,,, at 347 nm (120). The peak at 358 nm corresponding to [IrC1613- was not obtained. This indicates that the expected [1rc1,13- was aquated and converted to [IrC1,(OHZ)]'-, since the specific rate of aquation of [I~cI(,]~is ca. 400 times that of [I~c~,I'-at 50°C (1 I ). The spectral characters of [IrC1613- and [I~c~,(oH,)]'- are very similar, making it difficult to distinguish one from the other. The spectra of the corresponding Ir(IV) species are, however, differ-

CAN. J . CHEM. VOL. 71. 1993


400

450

500 550 WAVELENGTH (nm)

600

FIG.1 . The visible spectrum of 104[1r(1V)]= 1.57 mol dm-' treated with a calculated volume of chlorine-saturated perchloric acid to give a solution in 3 mol dm-' perchloric acid. The reference solution was 3 mol dm-' chlorine-saturated perchloric acid.

WAVELENGTH (nm) FIG.2. The spectrum of the spent reaction mixture diluted with an equal volume of chlorine-saturated perchloric acid. The initial concentrations present in the diluted reaction mixture were 1O4[1rCI6'-] = 1.57 and [HCI04] = 3 mol dm-', along with unreacted hydroxylamine. ent. Since Ir(II1) complexes are rapidly and quantitatively oxidized to the corresponding Ir(1V) complexes by chlorine (1 l ) , the spent reaction mixture was oxidized with perchloric acid saturated with chlorine. The details are given below. The spent reaction mixture was twice diluted with chlorine-saturated perchloric acid (= 3 mol dm-'). The reference was a solution of the same perchloric acid of the same strength. The spectrum of this solution, Fig. 2, showed a peak absorbance of 0.477 at 448 nm, which is same as the known A,,,,, at 450 nm for [IrCIS(OH2)]- (12). Using eli0 = 3320 dm3 mol-' cm-', the [IrC1,(OH2)-] was calculated to be 1.44 X 10-"01 dm-3, indicating that it is present to the extent of almost 92% of the initial [IrCI:-]. And this includes the conversion of 13% of [I~cI,]'- present as impurity in the sample. Rapid scan of the reaction mi.rtzrr The successive rapid-scan spectra of the reaction mixture, recorded using 482 nm as the central wavelength on an Union-Giken RA-415 stopped-flow spectrophotometer with an interval (the time lapsing between the two consecutive scans) of zero time and a gate time (time taken to complete the scan) of 2 ms, are shown in Fig. 3. The reference was a spent reaction mixture allowed to remain in the mixing cell of the stopped-flow instrument for a period much greater than 9 half-lives of the reaction. The scan (-0-0-0-) corresponds to pure [IrCl,]". It is to be noted that the absorbance of the scan recorded after 4 ms of mixing was greater than that of the pure solution. The absorbance of the successive scans, recorded at time intervals of 8, 12, 16, 20, 26, and 32 ms, decreased and tended to reach a saturation value. That the gradual

decrease in the absorbance, indicative of the formation of new species believed to be an outer-sphere complex, is not due to the decrease of [1rC12-] used in oxidizing of hydroxylamine during the period of measurement is supported by the rate constant at much higher [NH,OH+]. The equilibrium for the complex formation is probably attained in about 32 ms when the absorbance assumes a constant value. Kinetic mectsurements The kinetics were studied in acetate buffer under pseudo-firstorder conditions ([NH20H]2 10[IrC16'-1) at constant ionic strength (I) maintained with NaCIO,. The solutions of NH'OH+ in the desired buffer and of Ir(IV), taken separately in the stock cylinders of the Union-Giken RA-401 stopped-flow spectrophotometer, were brought to thermal equilibrium at the desired temperature (+O. 1°C). The temperature was maintained by circulating water from a Haake D8G refrigerated circulatory water bath. The instrument has a dead time of about 0.5 ms for a 2 mm flow cell. The rate was followed in terms of disappearance of [I~CI,J'- at 488 nm, at which wavelength other components of the reaction mixture are transparent. The averaged volt-time curves, at least from 10-12 runs, were analysed for the values of the pseudo-first-order rate constant, k,~,,, by an in-built program in the computer used for collecting the data. The reproducibility was within ? 5 % . Test,for ,pee radical The reactant solutions were purged with nitrogen and 1 mL of acrylonitrile was added to each solution. No polymerization of the monomer was noted over ca. 5-10 min. However, a curdy white

REKHA ET AL.

W

0

Z

a

g!

0.10

S: m Can. J. Chem. Downloaded from www.nrcresearchpress.com by 14.139.230.12 on 03/20/14 For personal use only.

a

0.05

0 4 34

482

530

WAVELENGTH (nm) FIG.3. The rapid-scan spectrum of the reaction mixture having 10'[1r~l,'-] = 2.5, IO~[NH,OH+]= 5.0 mol dm-', pH 4.27, and 20°C. corresponds to pure [ I ~ c I , ] 'solution of the same concentration (absorbance = 0.0925, hence calculated E = The spectrum ( - 0 - 0 - 0 - ) 3700 dm3 mol-' cnl-I). The solid line spectra, top to bottom, were recorded at the beg~nningof 4, 8, 12, 16, 20, 26, and 32 ms after mixing.

0 1 0.00

I

I

I

I

0.05

0.10

0.15

0.20

C CI-1 (mol Frc. 4. Plots of k ? ' vs. [CI-] at temperatures 25°C ( O ) , 30°C (0).35°C (0)and 40°C (A). ( ~ o ~ [ I ~ ( I=v )1 .O, ] ~ O ' [ N H ~ O H= + ]1.0, [Nail = 0.4, and I = 0.4 mol dm-'). precipitate appeared on mixing the two solutions, indicating the formation of a free radical in the reaction mixture.

Results Stoichiometry The stoichiometry is dependent upon the concentratior~s of the reactants. The A[IIc~,'-]/A[NH,OH'] increased from 1.03 t 0.04 to 2.13 + 0.07 with decreasing [NH,OH']

(Table SI'). The results could be expressed by the stoichiometric equations [4] and [5] under the two conditions. [4]

2[1rCl,12

+ ~NH?OH+ -,2[1rc1,13-

[5]

4[IrCI(,]'-

+ Nl + 2H10 + 4H

+

+ ~NH'OH' -,4[1rc1613-

+ N 2 0 + H1O + 6 H '

CAN. J. CHEM. VOL. 71. 1993

TABLE1. The effect of variation of pH on the rate constant kz (= k,,,,,/[NH,OH]) at different temperatures' k,(dm3 mol- ' s - ' )


pH

Obsd.

Calcd.

"10"[1rCl,'-] = 1.0, 10'[NHIOH]

Obsd.

=

1.0, I

=

Calcd.

Obsd.

Calcd.

Obsd.

Calcd.

0.3 niol dm-'

A variable stoichiometry as a function of excess [oxidant] is known for one-electron oxidants (8, 9, and 16). Since the primary interest in the kinetics is the characterization of the initial product, the rates were measured in excess of [NH,OH+], limiting the study represented by eq. [4]. Kinetic measurements were made over a broad range of [Ir(IV)] to confirm first-order dependence in [Ir(IV)]. The rate increased proportionately with [NH~OH'] at different pH, and the plots of kobsVS.[NH~OH']were linear, passing through the origin. The slope (~,,,/[NH~oH'] of the plot is taken as the measure of k', the second-order rate constant. The k,,, was unaffected by the outside addition of [1rc1,l3(Table ~ 2 ' ) . It was noted that changing the ratio of CH,COOH and CH,COO- in the buffer did not affect the rate. The k,,, was independent of [LiCIO,] but increased by 23% when NaClO, was increased by almost fivefold (Table ~ 3 ' ) .The increase .in the rate is, perhaps, due to medium effects of the sodium ion. The medium effect is an important factor in redox and substitution reactions (17), and is usually not so marked for intramolecular redox reactions (18). The effect of [Cl-] on k,,, (Table ~4'1, was studied at constant [Na'] = 0.4 mol dm--?, being the sum total concentrations of NaCl, NaClO,, and CH3COONa present in the reaction mixture, to avoid variations in the rate due to the medium effects of the ion by using mixtures of NaCl and NaClO,. Since the ionic strength and [Na+] are constant and [Clod-] has no effect on the rate, the decrease in the rate is considered to be due to chloride ions. At the suggestion of one of the referees, the effect of [Cl-] on the rate was briefly checked using mixtures of LiCl and LiClO,. These results confirmed that the inhibition of the rate was due to the C1ions. The linear plots of kz-' against [Cl-] have intercepts on the rate ordinate at each temperatures (Fig. 4). The data in Tables S1-S4 are presented as supplementary material.' Dependence on p H The effect of pH variation on the rate at different temperatures is given in Table 1 . The values of k, increase with the pH. The plots of k2-' vs. [H'] are linear with positive intercepts on the rate ordinate at different temperatures (Fig. 5 ) . ' ~ a b l e s of supplementary data (Tables S1-S4) can be purchased from: The Depository of Unpublished Data, Document Delivery, CISTI, National Research Council Canada, Ottawa, ON K l A OS2, Canada.

Thus the empirical rate law for the dependence of both [H'I and [Cl-1 is expressed by eq. [6].

where X represents H' and C1- ions.

Discussion The evidence, provided by the inertness of [I~cI,]'- to substitution,-?the stoichiometry, the induced polymerization of acrylonitrile, and the empirical rate law, can be simply explained by an outer-sphere mechanism in which the intermediate NH'0H.I is not oxidized further by the deficient Ir(IV) present in the reaction mixture. The lack of rate dependence on [Ir(III)] indicates that it does not appear on the product side of any equilibrium involved, either at the ratedetermining step or at any step preceding it. The rate is retarded by both C1- and H + ions in an identical manner. It suggests the existence of equilibria in which these ions appear on the product side. For chloride ion, such an equilibrium could be as in eq. [7]:

However, the formation of an itztzer-sphere complex, [Kls(NH30H)], is considered improbable because [K1612is substitution inert (19). The retarding effect of C1- ion could therefore be attributed to either of the equilibria shown in eqs. [8] and [9] where NH,OH+Cl- is an ion-pair. These two equilibria offer scope for proposing two different mechanisms, each leading to a rate law that is consistent with the empirical rate law in eq. [6]. The mechanism based on equilibrium [8] will be discussed in detail whereas the mechanism based on equilibrium [9], in which [Kl,(OH,)Iwould react with NH30H' as in reaction [13], is less likely, in view of the fact that the spectra of freshly prepared s o h tions and those aged for 24 h are the same, indicating that 3 ~means t that substitution is not complete in less than 1 min. See ref. 10. Hence the formation of such an inner-sphere complex for a reaction taking place on the stopped-flow time scale is simply inadmissible.


REKHA E T A L .

103 FIG.5. Plots of kZ-' 0.3 mol dm-').

VS.

C H + I (mol

dm'3

)

[H'] at 25°C (0130°C , (U), 35°C (a),and 40°C (A). (104[1r(1~)] = 1.0, IO'[NH,OH+]

=

1.0, and I

=

no species other than [I~cI,]'- was initially present in the solution.

constant for the dissociation of the product after the electron transfer. Reaction [16], with an excess of Ir(IV), is replaced by the reactions [16a] and [16b], leading to the formation of N 2 0 , and thus explaining the high stoichiometric ratio with excess of Ir(1V).

With support from the rapid-scan spectrum for the formation of an outer-sphere complex, the following mechanism is postulated. It might be mentioned that a chloride-bridged analogous to a cyanide-bridged intermediate [1rC1,~1~0,]~-, intermediate [F~(CN),CNSO,]'- (19), was suggested in the Lr(1V) oxidation of sulphite ion (4). The formation of the intermediate was confmed in a recent reinvestigation of the reaction (20). The formation of an outer-sphere complex is further probable on the basis of the analogy that the sulfur atom in SO,'- and the nitrogen atom in NH,OH each have an 's pair of unshared electrons, and [1rc1612- and [F~(cN),~]'-each have one unpaired d electron.

[16n]

NH~OH" + Ir(IV) + Ir(II1) + 2H'

+ HNO

From the application of the steady-state treatment, one obtains:

[ 181

[I~cI,NH,oH]'-

=

k, [ I C ~ ~ ] '[NH,OH'] k-,[H+]

+ k,,

With the assumption that k, 9 k-,,, the k,,, is expressed by eq. [20].

Equation [20] can be rearranged to eq. [21] where [NH'OH . HCl]o/k,bs = kz-l.

[16]

-

fast 2NH20He+ N2 + 2HZ0+ 2H'

where K,,, is the equilibrium constant for the dissociation of the NH,OH+Cl- ion pair, KO,(= kI/k-,) is the equilibrium constant for the formation of the outer-sphere complex [I~CI,NH,OH]'-, k,, is the rate constant for the electron transfer within the outer-sphere complex, and k, is the rate

Equation [2 11 is consistent with the linear plots between k2-I and [Cl-1, Fig. 4, at the different temperatures investigated. The estimated values of K,,,, at different temperatures, from the intercepts and slopes of these plots are given in Table 2. It is interesting to note that the Kip, values, 2.91 t 0.03 mol dmp3, are almost independent of temper-

C A N . J . CHEM. VOL. 71, 1093

TABLE2. The values of the rate constants k, and k,,K,, and the equilibrium constant K,,, at different temperatures, and the respective activation parameters. Temperature ("C)


k, (dm' mol-I s - ' ) 1O ' ~ ~ , K(s< ,'1~ K,,,, (mol dm -') K,,,,,(dm' mol-I)

147 2.80 2.94 0.340

ature. This could be rationalized if the AH,,, has a small value. The value of K,,! (= l/Ki,,) was computed from the Fuoss equation (2 1) uslng 1.40 and 1.8 1 A as the radii of NH,OH+ and C1- ions, respectively. Again, in the absence of [CIP] eq. [20] can be written as eq. [22], which explains the dependence of the rate on [H']. kc,k , [NH20H . HClIo [221

kc,,, =

k-,[H+]

[NHIOH. HCl],

1

- - -

k2

kubs

[H'] k,lK,s

3. 4. 5. 6. 7.

+ k,,

Equation [22] can be rearranged to eq. [23] k , / k - , = K,,, which is consistent with the plots of kz-' vs. [H'] (Fig. 5). [231

168 3.18 2.92 0.342

+

8.

1 -

kl

The computed values of k , and k,,K,,, are given in Table 2 together with the associated activation and thermodynamic parameters. In conclusion, N H 3 0 H T rather than NH,OH is the reactive hydroxylamine species and the skeleton mechanism proposed by Sen Gupta et al. (3) is confirmed. However, it must be noted that the retardation in the rate by H' is due to reaction [13] and is not due to the disNHzOH + H + . Thus the mechanism sociation NH,OH+ of the oxidation of hydroxylamine by [Ir~l,]'- is different ' in the from its oxidation by isostructural [ ~ t ~ l , ]studied same pH range (22).

=

Acknowledgments The authors thank the Department of Science and Technology for the grant to purchase the stopped-flow instrument and the Council of Scientific and Industrial Research, New Delhi, for financial support of the research work. Thanks are also due to the referees and Professor D.M. Stanbury for the constructive suggestions. 1. M.N. Hughes. The inorganic chemistry of biological processes. 2nd. ed. Wiley, New York. 198 1. pp. 204-2 1 1. 2. K. Wieghardt. bl Advances in inorganic and bioinorganic

9. 10. 1 1. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22.

209 3.46 2.91 0.344

249 3.77 2.87 0.348

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