Kinetics and Mechanism of Oxidation of Hydroxylamine by Peroxomonosulphate

MADHU SHARMA, D.S. N. PRASAD, and K. S. GUPTA Department of Chemistry, University of Rajasthan, Jaipur 302 004, India

Abstract

The kinetics of the reaction (1) obey the rate law (2 ) in acetate buffered solutions. A compari- son with HZOP and S Z O ~ ' shows the reactivity order to be HPOZ < HS05- > SzO8'-. 0 1992 John Wiley & Sons, Inc.

(1) 2HS05- + NH30H' - 2S04'- + N P + 3 H 2 0 + 3H'

( 2 ) -d[HSOs-] /d t = kK[PMS][NH30H'] (K + [H'])-'

Introduction

Chemical curiousity [l-51 coupled with a desire to carry out a compara- tive rate analysis [6-71 with H20z was the chief reason for investigating the oxidation reactions of peroxomonosulphate (PMS). Recently, it has been shown that peroxomonosulphate has a n important atmospheric connection [6-71; model calculations have predicted that HS05- may be present in re- mote clouds [B ] . In addition, in several heterogenous autoxidation reactions [9] of the aqueous sulphur dioxide studied in this laboratory [3] and in homo- geneous reactions [lo] studied by others, the formation of HS05- has been implicated. Because of this atmospheric connection, we decided to investi- gate the oxidation of some inorganic substrates by this oxidant in order to elicit the mechanistic features of the oxidation reaction of PMS. In the case of 2-equivalent reductants, the most prevalent peroxomonosulphate oxida- tion mechanism is of nonradical type, involving the nucleophilic attack by substrate on the peroxide bond. On the other hand, the oxidation of oxo- vanadium (1V)-al-equivalent reductant proceeds through the intermediacy of sulphate ion radicals, the evidence for which came from spin trapping experiments.

We wanted to study an oxidation reaction of PMS, which was likely to pro- ceed through a free radical mechanism. For this purpose, we chose hydroxyl- amine, the oxidation of which can proceed to different stages depending upon the [oxidant]/[NH20H] ratio. Frequently, its oxidation involves the formation of the free radicals.

Materials and Methods

Peroxomonosulphate was obtained from Aldrich in the form of a triple salt, 2KHS05 . KHSO, . K2SO4(OXONE). The iodometric assay indicated

International Journal of Chemical Kinetics, Vol. 24, 665-670 (1992) 0 1992 John Wiley & Sons, Inc. CCC 0538-80661921070665-06$04.00

666 SHARMA, PRASAD, AND GUPTA

peroxomonosulphate content to be around 85%. Hydroxylamine sulphate used was of Reidel AR make. For maintaining pH, acetate buffer was used. The reactions, which were studied in Erlenmeyer flasks, were initiated by adding temperature equilibrated PMS solution to mixture containing hy- droxylammonium sulphate and buffer in required amounts and maintained at the desired temperature (-tO.l"C).

The kinetics were followed by estimating the unreacted PMS iodometri- cally [11,12]. The reaction was quenched by adding aliquots to a titration flask containing ice cold water and cracked ice. A 10% KI solution was added and the liberated iodine was titrated against standard sodium thiosulphate solution using starch as an indicator. The kinetic results were reproducible to within 55%.

Stoichiometry

Since all kinetic runs were done with [NH30H'] in excess over [PMS], the stoichiometry was determined in this condition. The reaction mixtures were kept for more than 24h to ensure completion of the reaction. The unreacted NH30H+ was determined with the permanganate [lo]. These results showed that for each mol of hydroxylamine, 0.54? 0.05 mol PMS was required.

(1) HSOs- + BNH3OH' 4 S04*- + Nz + 3Hz0 + 3H'

The evolution of Nz was seen during the reaction. It is interesting to point out that Nz is the product in peroxytungstate [14], and Cui2 and (bipy)Cu+' catalyzed oxidation of hydroxylamine by HzO2[15], also.

Results

In all kinetics experiments, [NH30H'] was in excess over [PMS]. The pseudo-first order plots, obtained by plotting log [PMS], against time, were linear up to at least 70% reaction, showing first order in PMS. However, on increasing [PMS], the pseudo-first order rate constant, kl, were found to decrease (Table I). A similar behavior has been observed in several peroxo- disulphate oxidations [16-181, also. Incidentally, in the catalytic decomposi- tion of PMS [19] in aqueous perchloric acid solution by the dual catalyst Ag' and S202-, the pseudo-zero order rate constants defined by eq. (2) were found to increase on increasing [PMS]. This behavior was adduced as being indicative of the presence of a n unknown catalyst in PMS stock solution.

(2) -d[HSOs-]/dt = Izo

[NH,OH+] was varied in the range (0.8-3.0) x 10~-2 mol dm-3 at three dif- mol dm-3 at pH = 3.7. The pseud- ferent [PMS] of (0.75, 1.0, and 1.5) X

TABLE I. Values of pseudo-first order rate constants, k l , at different [PMSI at pH = 3.7, [NH30H+l = 4.0 X lo-' mol dm-3, and t = 33°C.

lo3 [PMSI mol dm-3 0.75 1.2 1.5 3.0 4.0 k , , 5-l 1.22 1.2 0.92 0.54 0.42

OXIDATION OF HYDROXYLAMINE 667

o-first order rate constants, k l , were found to increase in proportion to [NH30H'] at each [PMS], indicating the first order in reducing agent. The results for [NH,OH'] variation at 7.5 X mol dm-3 [PMS] are shown in Figure 1. Thus the results are in accord with the experimental rate law (3).

(3) -d[PMS]/dt = kZ[PMS] [NH,OH+]

and

(4) k , = kZ[NH30H+]

An investigation of the dependence of the rate of reaction on pH (3.7-4.5) at 1 x mol dm-3 [PMS] showed the rate to increase on increasing pH.

The values of k , were found to fit the rate law (5).

hKtNH3OH '3 K + [H']

hi =

Accordingly, the plots of k.-' and [H'] were found to be linear (Fig. 2). The values of k , k K , and K, at different temperatures, are given in Table 11.

A variation in ionic strength with sodium perchlorate did not effect the rate of reaction. The energy of activation associated with 12 was found to be 138 kJ mol-l.

When the oxidation reaction was carried out in presence of methyl methacrylate, the polymerization of latter occurred.

Discussion

The first dissociation constant of peroxomonosulphuric acid, is reported to be high [20] and the second dissociation constant [21] to be very low, being only 1.32 x lo-''. Thus, in the pH range of this study peroxomonosulphate will be almost wholly present as HS05-. The observed pH-dependence,

0 lo2 [NH,OH*], mol dni3

Figure 1. Dependence of pseudo-first order rate constant, k l , on [NH,OH'l at pH =

3.7 and I = 1.0 mol d ~ n - ~ at 33°C. [PMS] = 7.5 X mol dm-3.

668 SHARMA, PRASAD, AND GUPTA

3 16*[H+],mol dni3

Figure 2. 1 X

Dependence of pseudo-first order rate constant, kl , on [H'l a t [PMSI =

mol dm-3, [NH30H'l = 1 X lo-* mol dm-3, I = 1.0 mol dm-3, and t = 33°C.

TABLE 11. Values of k l , K, and kK at different temperatures and at [PMS] = 1 X mol dm-3.

Temp. lo5 kK k lo6 K "C S-' dm3 rnol-' s-' mol dm-3

33 35 40

0.63 0.81 2.5

0.86 1.1 6.3

7.0 7.0 4.0

therefore, cannot be linked to peroxomonosulphate and is most likely due to hydroxylammonium ion, whose pK is reported [22] to be 5.8. Consistent with kinetics results and stoichiometry applicable to the experimental con- dition of this kinetics study, i.e., [NH30H+] > [PMS], the following mecha- nism may be proposed.

K (6) NH30H' 6 NH20H + H +

(7) HS05- + NH2OH - so4-' + NH20' + HzO

(8) so,-' + NH30H' S04'- + NH20' + 2H'

k

k

(9) so4-' + NH20H SO?- + NH20. + H +

(10) 2NH20' 4 N2 + 2H20 (Fast)

Step (7) envisions the attack of NH20H on peroxide bond. This results in the formation of SO4-' and NH20' radicals, through hydrogen atom abstrac- tion as in Cu(I1) catalyzed H202-NH20H reaction [15]. The polymerization of methyl methacrylate lends support to the presence of the free radicals but not to their origin. Stoichiometric equ. (1) suggests the formation of so,-' radicals. Incidentally, the evidence for the formation of these radicals in the oxidation of oxovanadium(1V) by PMS [5] has been provided.

Subsequently, SO4-' formed in step (7) oxidizes NHBOH+/NH20H in fast steps. Fortunately, the value of k s and K, are available [23]. At ambient tem- perature, k8 has a value of 8.5 X lo8 dm3 mol-1 s-l (pH 8.2) and K, has a value of 1.5 X lo7 dm3 mol-' s-l (pH 4.1). The step (10) is a well known reac- tion [24].

OXIDATION OF HYDROXYLAMINE 669

The mechanism (6-10) leads to the experimental rate

d[HS05-1 - kK[PMSl [NHzOHl - -

d t (K + [H'I) (11)

law (11)

In the proposed mechanism, reaction (7) is written as leading to SO4-'. Since the reaction of solvated electrons with HS05- is known to lead to both possibilities (OH- + SO,-' and OH' + SO4'-) [22], it is very likely that the reaction with hydroxylamine occurs similarly. Thus, an alternative or con- current mechanism (12-14) is also considered possible.

k (12) HS05- + NHzOH - s04'- + NHz0' + OH' + H i

(13) OH' + NH30H+ --% NH,O' + H + + HzO

(14) OH' + NHzOH A NH20' + H 2 0

The values of K determined from the kinetics experiments are in good agreement with those reported by others [23]. It is of interest to compare the reactivity of HSOs- with H202 and S202'. There is no report on the un- catalyzed oxidation of hydroxylamine by H2Oz. As the rate constant for Cu(I1) catalyzed oxidation [15] by H20z is lower than the rate for HS05-- NHzOH reaction (Table 1111, it is obvious that the rate of uncatalyzed H2Oz-NH20H reaction is much less. As for S z 0 2 - , its reaction with hy- droxylamine in the absence of the catalyst is reported to be very slow [26]. Thus, the reactivity sequence is

SzOs"- < HS05- > HzOz

By an analysis of relative reactivity of HS05- and Hz02 towards several nucleophiles, Betterton and Hoffmann [7] have shown HS05- to react 3-4 orders of magnitude faster than HzOz. In the present case also, the two oxi- dants behave similarly. This is in agreement with the general finding that HS05- is more reactive kinetically than HzOz [7].

The much lower reactivity of S z 0 2 - probably arises out of steric hindrance due to the presence of two -SOa groups on both sides of peroxide bond.

Incidentally a comparison of rate data in Table I11 show Hz02, HzPzOs2- [ll] and H3P05 [27] to follow a similar reactivity trend, i.e., H2O2 <

TABLE 111. Rate constants for the oxidation of hydroxylamine with different oxidants.

React ion Rate constant Condition ref.

1. Hz02-NH20H 3 x

2. HSOs--NH20H 6.3,

3. S208"-NH30H+ very low 4. H~POS-NH~OH' 1.15 X

5. H2PzO~--NH30H' 2.65 X

(Cu"-catalyzed) dm3 mol-l S - l , a . b

dm3 mol-' s - ' , ~

dm3 mol-' s-','

dm3 mol-' s-','

t = 25" [I51 pH range 5.3-8.3 t = 40°C This work pH range 3.7-4.5

[261 t = 45" [271 [H'l = 0.1-2.0 mol dm-3 t = 60" [11l pH range 3.3-4.5

-

"Rate is defined by, -d[HzOzl/dt = k . [Cu"] [HzOzl [NH,OH'I [H'l-'. 'Rate increases on decreasing [H'I. 'Rate unaffected by [H'l.

670 SHARMA, PRASAD, AND GUPTA

H3P05 > HzPz02-. A comparison between the relative reactivities of HS05- and H3P05 is precluded as the study with H3P05 was done in strong acid medium.

Acknowledgment

The work was supported by U. G. C. and Indo-US Subcommission research projects.

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Received September 30, 1991 Accepted January 2, 1992