Indian Jouma! of ChemistryVol. 28A, June 1989, pp. 501-504 -

Kinetics and mechanism of uncatalysed & Ir(III)-catalysed oxidation ofoxalate ion by N-bromosuccinirriide in basic medium

P Saroja. B Kishore Kumar & Sushama Kandlikar*

Department of Chemistry, Nizam College (Osmania University), Hyderabad 500 001

Received 22 February 1988; revised 16 May 1988; accepted 8 September 1988

The title reaction in the presence of sodium hydroxide is first order in [NBS] both in the presenceand absence of Ir(lII) catalyst. However, the order in [C204J2- is unity in the absence and fractional inthe presence of Ir(IlI) catalyst. Increase in [OH-l accelerates the reaction while the added succinimideretards the reaction rate. Increase in the ionic strength of the reaction medium increases the reactionrate in the case of uncatalysed reaction. The rate increases with decrease in dielectric constant of thereaction medium. All these observations show that the rate-determining step involves two negative ions.A mechanism involving the hypobromite ion as the reactive species of the oxidant has been proposed.Individual rate constants (k), the formation constant (Kf) of the complex between oxalate ion and thecatalyst at different temperatures and the corresponding thermodynamic parameters have been comput-ed.

N-Bromosuccinimide (NBS) has been extensivelyused as an oxidant in the study of kinetics andmechanism of oxidation of both organic and inor-ganic substrates in acid medium'. However, thereis no report on the oxidation of any substrate byNBS in basic medium. Herein we present the re-sults of title investigation. An attempt has beenmade to identify the most probable reactive spe-cies of the oxidant in basic medium. The values offormation constant (Kf) of the complex formed be-tween the oxalate ion and Ir(III) have been deter-mined at different temperatures. In order to pro-vide support for the formation of such a complex,oxidation of oxalate ion by the hypobromite ionproduced from Br2 in NaOH has also been carri-ed out in the presence of the above catalyst andthe formation constants (Kf) obtained by both themethods have been compared.

Materials and MethodsAll the chemicals used were of AR grade. The

solution of IrCl3 (Johnson-Matthey) was standar-dised by the method given by Singh et al.' z-Butylalcohol (J T Baker, NJ) was distilled before use. A0.1 mol drn? aqueous solution of NBS (E Merck)was prepared and standardised '. A 0.1 mol dm-3

solution of sodium oxalate (E Merck) was alwaysprepared afresh in doubly distilled water. A solu-tion of sodium hypobromite was prepared by dis-solving bromine (E. Merck) in sodium hydroxidesolution at O°C and its strength was verified io-dometrically. This solution was kept in a dark

,(

container at O'C. All the reactions were carriedout in reaction flasks blackened from outside.

Kinetic runsRequisite amounts of sodium oxalate, sodium

hydroxide, sodium perchlorate and doubly dis-tilled water (to keep the total volume constant forall the runs) were mixed and the reaction mixturewas equilibrated at 25° ± 0.1"C. Allowance wasmade for the amount of sulphuric acid used in thepreparation of the catalyst, while preparing solu-tions for the kinetic runs. A measured amount ofthe oxidant, also equilibrated at the same tempera-ture was rapidly added to the above reaction mix-ture. The progress 'of the reaction was followed byiodometric determination of the unreacted oxidantin aliquots (5 ml each) of the reaction mixturewithdrawn at regular time intervals. The rate con-stants calculated were reproducible within ± 3-5%.

Results and DiscussionThe results can be summarised as follows:(i) A mixture containing 5.0 X 10-2 mol dm-3 of

NaOH, 5.0 x 10-2 mol dm ? of Na2C204 and alarge excess of NBS was kept for 36 hr at25° ± 2°C. The results of stoichiometric runs, afterapplying a small blank correction for the self-de-composition of NBS, indicated that 1 mole of so-dium oxalate reacted with 1 mole of NBS to af-ford 1 mole of sodium carbonate and 1 mole ofcarbon dioxide gas. CO2 gas was identified as

501

INDIAN J. CHEM., SEe. A, JUNE 1989

usual while estimation of Na2C03 was done" asfollows:

In one portion of the stoichiometric mixture thetotal alkali (carbonate +unutilised NaOH) was de-termined by titration with standard sulphuric acidusing methyl orange indicator. This concentrationwas found to be 0.048 ± 0.002 mol dm":', In an-other portion of the above mixture, the carbonatewas precipitated with a slight excess of BaCl2 solu-tion and without filtering, the solution was titratedwith the same standard acid using phenophthaleinas indicator. This titration gave the hydroxide con-tent left behind at the end of the reaction. Theconcentration of Na2C03, calculated from the dif-ference of the two titre values. was found to beO.044±0.002 mol dm-3. The end points were notsharp when the same estimation was carried outunder experimental conditions where the alkaliused was 0.001 mol dm-3•

(ii) Urider the experimental conditions[NBS] ~ [S] (where S denotes the substrate) and inthe presence and absence of Irflll) the plots oflog(rate) versus 10g[NBS] were linear with unitslopes indicating the order in [NBS] to be unity forboth catalysed and uncatalysed systems. As one ofthe products, succinimide had a retarding effect onboth catalysed and uncatalysed reactions; the rateswere obtained by the differential rate methodthroughout this work.

(iii) The rate increased with increase in [S]. Theplot of log(rate) versus 10g[S] in the absence ofIr(IlI) was linear with unit slope indicating first or-der dependence in [S].

(iv) At fixed [Ir(Ill)], the rate increased consider-ably but the order in [S] changed from unity tofractional, viz. 0.6 (Table 1).

(v) The reaction rate increased with increase in[Ir(Ill)] and order in [Ir(IlI)] was found to be frac-tional, when the gross rate Vg was considered.However, when the uncatalysed rate Vu was de-ducted from Vg the order in [Ir(Ill)] was found tobe unity from a linear plot of log Vc (= Vg - VJversus log [Ir(Ill)] (Table 2).

(vi) Keeping the concentration of all otherreactants fixed, the ionic strength was varied in therange of 0.002 to 0.02 mol dm" ' using NaCI04and Na2S04. The pseudo-first rate constants in-creased with increase in ionic strength (Table 3).When log k values were plotted against 1.\.1/2 forvarious concentrations of the salt in each case theslopes were found to be 2 ± 0.2.

However, addition of sodiumperchlorate in therange of 0.002 to 0.02 mol dm-3 had a negligibleeffect on the rate of catalysed reaction.

(vii) Addition of ten-fold excess of the product

S02

(

Table l=-Effect of Variation of [Substrate], on rate constant103 [NBS]= 5.00 mol dm -3; 103 [NaOH] = 1.00moldm-3;

105 [lr(III)] = 1.2 mol dm - 3; T= 298 K

102 [Substrate] Uncatalysed Catalysedmol dm-3

104 k' 10-4 X 11k' 104 k' 10-4 X 11k'(s-~ ) (s) (s -I) (s)

2.00 0.16 6.3 1.4 0.713.00 0.24 4.2 1.8 0.554.00 0.34 2.9 2.2 0.465.00 0.40 2.5 2.50 0.406.00 0.50 2.0 2.67 0.37g.OO 0.66 1.5 3.2 0.31

Table 2 - Effect ofYariation of[lr( III)]on rate constant103 [NBS] = 5.00 mol dm - 3; 103 [NaOH] = 1.00 mol dm - 3;

102 [Oxalate] = 5.00 mol dm-3; T= 298 K; 106yu = 0.20 moldm-3s-1

1Q5[Ir(I1I)] 5 +Iog 106Yg 106 Yc= 6+logYcmol dm " ' Ir(I1I) moldm-3

S-I (Yg-Yu)

1.2 0.0792 1.25 1.05 0.02122.4 0.3802 2.33 2.13 0.32843.6 0.5563 3.16 2.97 0.47274.8 0.6812 4.16 3.97 0.59886.0 0.7781 5.08 4.88 0.68849.6 0.9822 8.00 7.80 0.8921

Table 3-Effect of Variation of [NaCl04] on rate constant103[NBS] = 5.00 mol dm - 3; 103 [NaOH] = 1.00 mol dm - 3; 102[Ox-

alate] = 5.00moldm-3;T= 298K

11112 106y 104 k' 5 + log k'(rnol dm=') (mol dm Ys") (S-I)

OD447 0.25 0.500.0632 0.29 0.580.0894 0.33 0.660.1095 0.37 0.750.1265 0.42 0.83•

103 [NaCI04]

(mol dm r t)

2.04.08.0

12.016.0

0.69900.76340.81950.87510.9191

Table 4-Effect of variation of [Succinimide] on rate constant103[NBS]=5.00 mol dm >'; 102[Oxalate]=5.00 mol dm>';

103[NaOH]= 1.00moldm-3;T=298K

103[Succinimide] 106y 104k' 104 x 11k'(moldm-3) (mol dm r+s" ') (S-I) (s)

1.0 0.15 0.33 3.3

2.0 0.12 0.23 4.34.0 0.08 0.16 6.06.0 0.06 0.13 8.08.0 0.05 0.10 10.0

10.0 0.04 0.08 12.0

(snccinimide) at constant ionic strength decreasedthe rate of both catalysed and uncatalysed reac-tions (Table 4).

(viii) The rate of reaction increased with increasein [OH-] at constant ionic strength (NaCI04) in thepresence as well as in the absence of Ir(Ill) catalyst.

SAROJA et af.: KINETICS OF OXIDATION OF OXALATE ION BY N-BROMOSUCCINIMIDE

(ix)Considerable acceleration in rate was observedwith decrease in dielectric constant of the medium, i.e.by increasing the percentage of r-butyl alcohol in thesolvent medium in both catalysed and uncatalysedreactions.

(x)There was no induced polymerization when ac-rylamide was added to the reaction mixture indicatingthe absence of free-radicals.

(xi) The reaction was carried out at five differenttemperatures in therangeof298-318 Kand the activa-tion parameters were computed using Arrheniusequation.

Uncatalysed reaction~crease in rate with increase in [OH - J may be ex-

plained by assuming an equilibrium

fastNBS + NaOH r!- NaOBr + succinimide

K

The observed first order dependence each in [NBS]and [S],and the positive salt effect, indicating the reac-tion between two negative ions, are suggestive of a me-chanism involving hypobromite ion as the reactivespecies of the oxidant in basic medium

_ k"Na2C204 + OBr---- NaZC03 + COz + Br ".

slow

This leads to rate law (1)

_ d[NBS]= k"[S][OBr-]dt

where [OBr-] is given by relation (2)

[OBr-] = K[NBSlr[OH, - NHJ[NH]

... (1)

... (2)

In Eq. (2) NH denotes succinimide, f and t stand forfree and total. respectively. [NBS]r is given by the rela-tion (3)

[NBS] = lNBSMNH]r [NH]+K[OH,-NH]

... \3)

Considering Egs (2) and (3), the rate law (1) assumesthe form (4)

_ dlNBS] = k" KlS][NBSllOH, - NH]dt [NH]+K[OH,-NH]

or

dlNBS] x __ 1 _ = k' k" KlSllOH, - NHJdt [NBSj, [NH]+ K[OH,-NH]

or

1 [NH] 1-= +--k' k"K[OH,-NH][S] k"[S]

... (4)

The plot of 11k' against [succinimideJ was linear, theintercept of which gave the value of k" as to be8.0 x 10-4. When this value was substituted in the va-lue of slope, K[OH, - NH] was obtained. When 1/ k'was plotted against 1/[S] a linear plot passing throughthe origin was obtained. The value of [succinirnide]was obtained by substituting k" and K{ OH, - NH] va-lues in the slope of the latter plot. The value of K wasdetermined as 4.5 by substituting this value of[succini-mide] in K[OH, - NH].

The increase in reaction rate with decrease in die-lectric constant of the medium indicate that the acti-vated complex was less polar than the reactants andcan be accounted for by Laidler and Eyring':" equa-tion (Eg. 5).

e2 [1 1]lnk=lnk +-- ---

o 2DkT r r'... (5)

where k'« are the rate constants and r's the radii ofthereactant species and the activated complex. It is evi-dent from Eq. (5) that the radius ofthe activated com-plex is higher than the radii of the reactants as the rate-determining step involves the interaction between twonegative ions.

Catalysed reactionFirst order dependence in [NBS], fractional order

each in [S]and [Ir(III)] and the ability of the platinumgroup metals to form complexes with organic sub-strates and highly enhanced rates substantiate the for-mation ofIr(III) - S complex (Scheme 1)which slowlyreacts with NBS in the rate-determining step to giveproducts.

fastNBS+ NaOH"K NaOBr+ NH(CH2CO)z

KrIr(lIl) + S r!- [Ir(Ill) - S] complex

k .[Ir(IlI)-S]+ -OBr----+[Ir(V)-S]

slow

fast[Ir(V)-S]- Na2C03 +C02 + Ir(IlI) + Br :

Scheme 1

Scheme 1 leads to rate law (6 )

d[NBS]- dt = klOBr-][complex]

503

INDIAN J. CHEM., SEe. A, JUNE 1989

or

~ d[NBS] = kKKJS][Ir(III)][NBSlt[OH, - NH]dt {l + KJS]}{[NH] + K[OH, - NH]}

ord[NBS] 1

dt [NBSlk' = kKKJS][Ir(III)][OH, - NH]

{l + KJSJl{[NH]+ K[OH, - NH]}

or

1 1 [[NH] l)k' = [S] kKKJIr(III)][OH, - NH] + kKr[Ir(III)]

+ [NH] + K[OHI - NH]kK[Ir(III)][Og-NH] ... (6)

where k' is the observed pseudo-first order rate con-stant, kthe bimolecular rate constant for the slow stepand K c the formation constant of the complex betweenthe catalyst and the substrate. Equation (6) explains allthe experimental results. The plot of 1/ k' against 1/[S]at constant [Ir(III)] was linear with an intercept on the1/ k' axis confirming the derived rate law. Dividing theintercept value by toe slope value the values of forma-tion constants (K c) were evaluated at different temper-atures. The values of k; the bimolecular rate constantswere calculated from the values of intercept after sub-stituting for K[Og - NH] and [succinimide] takenfrom theuncatalysedreaction. It was observed that theformation constants (Kr) decreased with increase intemperature indicating that the reaction betweenIr(III) and oxalate ion is exothermic in nature. This isalso indicated by the /)"H value corresponding to theformation constant (Table 5).

In order to verify the rate law derived, a direct de-termination of the formation constant was also carriedout by the oxidation of oxalate ion by the hypobromiteion produced from bromine in sodium hydroxide inthe temperature range of 268-278 K. As this studycould not be conducted at room temperature, the va-lues of Kcin the temperature range of298-3l3 K were

504

Table 5=-Kinetic and thermodynamic parameters for catalysedreaction

(Kr values determined by Br/NaOH are given in parentheses)

Temperature K, -.l'1H -l'1G -l'1S(K) (mol dm r t] (kJmol-l) (kJmol-l) (JK-I mol-I)

298 29.0 (25.1) 16.1 8.3 26.2303 26.7 (24.0)308 23.4 (23.2)313 20.8 (22.2)

obtained by extrapolating the plot oflog K r versus lffdrawn for tlie temperature range of 268- 278 K. Thesevalues agreed with those determined by NBS as an oxi-dant within ± 5-10% error (Table 5).

As regards the effect of dielectric constant of thesolvent on the rate of catalysed reaction, the ratio ofcatalysed/uncatalysed reaction rates was found to in-crease with increase in I-BuOH in the reaction mix-ture. This may be due to r~a,> r~ncal'in the Laidler-Eyring (Eq. 5) as the charge gets more dispersed overthe activated complex formed between the Ir(III) - Scomplex and -OBr ion in the catalysed reaction with% increase in I-BuOH in the reaction medium.

AcknowledgementThe authors are thankful to Prof T Navaneeth Rao

andProf B Sethuram for constant encouragement.One of the authors (SP) is thankful to the Principal, SNV M V College for laboratory facilities.

ReferencesI De Mattos Regina T, Lachter, Elizabeth R, Neto Ade\ina Costa

Quirn Nova 1<)85, [(I), \7-26 (Port). A review with JX refs.Chern Abstracts 107, 197105v.Sharrna J P,Singh R N P,Singh A K & Singh Bharat, Tetrahedron,(1986) 42 p. 2739, and the references therein.

2 Singh A K, Mohan Katyal & Singh R P, } Indian chem Soc, 53(1976)671.

3 Mathur N K & Narang C K, The determination of organic com-pounds with Nrbromosuccinimide and ullicd rl'l/gmtl (!\C'lti-emic Press. NY) 1975.

4 Vogel T B, Quantitative inorganic analysis (ELBS) 197H.5 Laidler K J, Chemical kinetics(McGraw Hill, New York) 1965.6 Entelis SG& Tiger RP, Reaction kinetics in the Iiquid phasetvvv-

ley, New York) 1976.