Indian Journal of Chemistry Vol. 40A. July 2001 , pp. 7 14-7 19

Kinetics and mechanism of silver (I) catalyzed oxidation of 1,3-propanediol by peroxodiphosphate in acetate buffers

G Singh, Savita Bansal, Divya Gupta, Indu Sharma, C L Khandelwal & PO Sharma*

Department of Chemistry, University of Raj as than , Jaipur 302 004. India

Received 13 Novell/ber 2000; revised 9 April 200 I

The kinetics of si lver (I) catalysed oxidati on of 1,3-propanediol by peroxodiphosphate (pdp) has been studied in acetate buffers. 3-Hydroxypropanal has been identified spectrally to be the ox idation product of 1,3-propanediol. However. the rate is independen t of di ol concen tration but first order each in the cata lyst and ox idant. Rate is significantly rewrded by acetate ions. The ki netics of the reaction account for the rate law.

_ dlpdp] = (k", + k",.KIMeCO; ])[pdp][Ag(I)] dl 1+ KIMeCO; ]

Rate parameters have been evaluated under varied ex perimental conditions. The role of the catalyst has been suggested II ,

th rough Ag / Ag redox cycle.

Peroxodiphosphate and peroxodisulphate despite being isoelectronic and isos tructural differs distinctly in their mode of ox idation of the substrates in aqueous ac id medium'. These ox idants have almost simil ar oxidation potenti als , but do not follow the same pattern of reacti vi ty and the ir reactions are signifi-cant ly slow despite their high values of oxidation potentials2.3. [n the case of peroxodiphosphate, hydrolysis step was limiting but in the case of peroxodisulphate' homolysis step was the rate limiting and these are considered to be the important reasons for their slow reactivity .

Silver (I) is considered to be a useful catalyst in reactions of peroxocompounds in aqueous acid soluti ons5.6 and the role of silver (I ) in reactions of peroxodisulphate is more or less established6. Peroxodiphosphate undergoes facile hydrolysis in aq ueous acidic solutions but the latter is almost completely checked in acetate buffers. Therefore, ex tensive analys is of the kinetic studies o f peroxodiphosphate with vari ous substrates in acetate buffers in the presence of Ag(1) as catalyst is required before arriving at any conclusion regarding reaction mechanism. Thi s was the reason that prompted us to undertake the detailed kinetic study of the title reaction.

The role of silver (I) as a catalyst in oxidation of non-vicinal diols has not yet been ex plored ex tensively. Since dial in the title reaction does not have hydroxy groups on the vici nal carbons, the probe

of the reactivity pattern will not o nly be interesting but useful too in gaining better insight of silver (1) chemistry as a cata lyst.

If the reaction takes place th rough ~he interaction of the diol and one-equivalent ox idant, the substrate free radical is expected to be formed. Whether the free radical has any role in the reaction mechani sm is a subject for furth er investi gation .

Materials and Methods Potass ium peroxodiphosphate was llsed as received

(a g ift sample from FMC Corporation, New York, USA). Although the salt contained some impurity of phosphate, the latter did not interfere in the kinetics of the reaction. The preparation and standardization of other reagents and solutions are given elsewhere4 .

Doubly di stilled water was e mployed throughout the study, the second distillation was from alkaline potass ium permanganate solution in an all glass apparatus.

The reactions were conducted in g lass stoppered Erlenmeyer fl asks suspended i a water-bath thermostated at 50 ± 0.1 DC unless stated otherwise. All other ingredients of the reaction mixture except peroxodiphosphate (pdp) were taken in these fl asks. Required volume of pdp was withd rawn from the thermally pre-equilibrated solu tion and then di scharged into the reaction mixture. The time of initiation of the reaction was recorded when the pipette was half empty of the contents . However,

SINGH el al.: OXIDATION OF 1,3-PROPANEDIOL BY PEROXODIPHOSPH ATE 7 15

reac tions were not in iti ated by di ol due to ensuing sil ver(l ) catalysed decompos ition of peroxodiphos-phate5.

An aliquot sample (5 cmJ ) of the reaction mi xture was withdrawn peri odi call y and then di scharged into a freshl y prepared potassi um iodide solution ( 10%) which contained 0.5 mol dm-J HClO~. A solution of the mi xed catalys t (I cmJ ) (Fe(ll ) + Cu(lI )) was added in the end .

The liberated iodine was titrated against sodium thiosulphate solution using starch as an indicator7 . Diol did not interfe re in iodometri c analysis of peroxodiphosphate. However, iodine liberated by the mi xed catalyst was accounted fo r in the subsequent calcu lations of the peroxod iphosphate concentrati on.

Initi al rates were computed by employ ing plane mirror method. Pseudo- first order plots were also made wherever reacti on conditions permitted. Rate measurements in tripli cate were reproducible to within ± 5%.

Sroich iOlllel ry The stoichiometry was not determined in those

reacti ons in which excess of peroxodi phosphate over diol concentrati on was taken as the results were not reproducible due to ensuing decomposition of the ox idant in the presence of sil ver (I ). Stoichiometry was determined in an excess of [diol] over [pdp] . The reactions were allowed to occur in a thermostated water-bath at 50°C and the ox idati on product of the diol was identified spectrall y.

Product idelllijicalioll Reaction mi xture containing [diol] (0.1 mol dm-3)

in excess over [pdp] (0.05 mol dm-3) in the presence of sil ver (I ) (1.0 x IO-J mol dm-3) catalyst was thermally equilibrated at 50°C fo r 6 h. When reaction was over, sil ver(l ) was precipitated by adding NaCI solution. A solution of 2,4-dinitrophenylhydrazine was added in to the filtrate in the presence of 2.0 mol dm-J HCI and then react ion mixture was left overni ght at refrigerated temperature (- SOC). The solution was centri fuged and the brown-orange residue was washed th rice with ice-cold water and then air-dried before carrying out MR spectral analysis.

NMR results alongwith the m.p. of hydrazone deri vati ve ( 152°C) bei ng in agreement wi th the reportedS value (l50°C) adequ ately confirmed the

ox idation product of the diol to be 3-hydroxypropanal.

It is, therefore, clear that only one hydroxy group of the di ol was attacked and other remained intact aivin a stoichiometry of the reacti on as represented by b b Eq. ( I )

Ag(l) CH20H-CH2- CH20H+H4P20 S ~

CHOCH2CH20H+2H 3PO~ . . ... (1) It is, however, worth mentioning that 4-hydroxy-

butanal9 was found to be the ox idation product of the butanedi ol by peroxodiphosphate in the presence of sil ver (I ) as catalyst and also in an excess of diol over peroxodi phosphate.

Results and Discussion The concentrati on of peroxodiphosphate was

varied in the range (1.0 x 10-] to lOx 10-3) mol dm-J

keeping concentrati ons of other reacti on ingred ients fixed, viz. [diol] = 2.0 x 10-2 mol dm-3, [Ag(l ) = I x IO-J

mol dm-J , pH= 4.0 and 1=0.2 mol dm-J (adjusted by sodium nitrate). In itial rates (k j ) were calcul ated and a plot of initial rate (kj ) versus [diol] was linear passi ng through the onglI1 confirming first order in peroxodiphosphate. Pseudo-first order plots were also constructed and the pseudo first order rate constants (k') were found to be independent of the ini tial [oxidant] (Table I).

The [diol] was varied from 5.0 x IO-J to 1.0 x 10- 1

mol dm-3 at fixed concentration of other reaction -3 I d -3 ingredients viz. [pdp] = 2.0 x 10 mo m ,

[Ag(l)] = 1.0 x IO-J mol dm-3 and pH = 4.0 at 1 = 0.2 mol dm-J . The rate is independent of [diol] ex hi biting zero order in the substrate (Table I).

[Silver(l )] was varied from 3.0 x I O~ to 2.0 x 10-] mol dm-3 at fixed [pdp]= 2.0x 10-3, [diol] = 5.0 x 10-2

mol dm-J and 1=0.2 mol dm-J at p H = 4.0. Plot of calculated pseudo first order rate constants (k ') against [silver(I)] was linear passing through the ori gin indicating first order in the catal yst (Table! ).

Hydrogen ion concentration was varied by chanaina the pH of the reaction mi xture from 3.63 to

b b ?

5.87 at [pdp]= 2.0 x 10-3 dm-3 ,[diol] =5.0x 10-- mol dm-3, [Ag(l)]= 1.0 x 10-3 mol dm-J and 1= 0.2 mol dm-J . However, p H before and after the reacti on exhibited a maximum change of 0.01 unit. Nevertheless, pH range was limited because of two main experimental di ffi culti es. Firstly, the pH if lowered below 2.5, slow hydrolysis of pdp started. Secondly if Ph is raised above 5.6, sil ver phosphate

716 INDIAN J. CHEM., SEC A, JULY 2001

Table 1-Pseudo-first order. second order rate constants and initial rates for the reaction of peroxod iphosphate (pdp) and propane- I ,3-diol (dial) in acetate buffers

pH = 4.0. temp = 45°C

103[pdp] 102[diol] 10J[Ag(l )] 107ki 10~k' 102k mol dm-J mol dm-J mol dm-3 mol dm-3 S- I S- I drn3 mol- I S- I

1.0 2.0 1.0 1.11 11.1 2.0 2.0 1.0 2.20 11.0 3.0 2.0 1.0 3.40 11.3 4.0 2.0 1.0 4.48 11.2 5.0 2.0 1.0 5.57 11.2 6.0 2.0 1.0 6.68 11.2 7.0 2.0 1.0 7.78 11.1 8.0 2.0 1.0 8.88 11.1 9.0 2.0 1.0 10.20 11.3 10.0 2.0 1.0 11.20 11.2 1.0 5.0 1.0 11.2 11.2 2.0 5.0 1.0 11.0 11.0 3.0 5.0 1.0 11.2 11.2 4.0 5.0 1.0 11.2 11.2 5.0 5.0 1.0 11.1 11.1 6.0 0. 1 1.0 11.4 11.4 7.0 0. 1 1.0 11.4 11 .4 8.0 0. 1 1.0 11.5 11.5 9.0 0. 1 1.0 11.3 11.3 2.0 5.0 1.0 11.2 11.2 2.0 5.0 2.0 22 .8 11 .4 2.0 5.0 3.0 34.2 11.4 2.0 5.0 4.0 45.8 11.5 2.0 5.0 5.0 56.7 11.3

Table 2 - Effect of sall s on the rate of the reaction

Ipdp] = 2.0 x IO-J mol dm-J. [dia l] = 5.0 x 10-2 mol dm-3, [Ag(l )] = 1.0 x 10-3 11101 dm-3, pH = 4.0, temp=45°C

INaNO~] 10~k ' [LiCI04 ) 10~k' [NaCI041 lO~k' 11101 dm-3 5- 1 mol dl11-'~ S- I mol dm-3 S- I

0.2 8.7 0.2 0.4 6.8 0.4 0.6 5.5 0.6 0.8 5.1 0.8 1.0 5.1 1.0

started precipitating. The rate, however, did not

depend upon pH of the reaction. Sodium nitrat e was e mployed to vary the ionic

strength of the reaction mixture from 0.2 to 1.0 mol dm-·I at fixed [pdp] = 2.0 x 10-3 mol dm-3,

[d iol]=5 .0x 10-2 mol dm-J , [Ag(l)]= 1.0 x 10-3 mol

dm-3 and pH =4.0. The rate decreased with increase in

[sodium nitrate].

The effec t of ionic strength on the rate of the

reacti on was also studied employ ing both sod ium perchlorate and lithium perchlorate respectively under

the same conditions. The rate decreased with increase in [sodium perchlorate] and [lithium perchlorate]

8.8 0.2 9. 1 7.68 0.4 7. 1 6. 1 0.6 5.8 4.9 0.8 4.8 4.6 1.0 4.4

respectively (Table 2). This decrease in the rate can

be assigned to the interaction of oppositely charged

ions in the rate controlling step. Sodium acetate concentration was a lso varied from

1.0 x 10-2 to 8.0 X 10-2 mol dm-3 under si milar

conditions and at 1 = 0.2 mol dm-3 ancl al so at three te mperatures viz. 45 , 50 and 55°C respect ive ly . The

rate increased with increase in concenLration of

sodium acetate and tended to attain limiting value at

higher concentrations of this salt.

It has earlier been establi shed that peroxodi-phosphate species in acidic solution are governed by the equilibria 'o (2) to (5)

SINGH el al.: OXIDATION OF 1,3-PROPANEDIOL BY PEROXODIPHOSPHATE 717

K,

H 4 P 20 8 ::::;;ooc;=::== ... ~

H 2 P2 0 ~ - ::;;OOC;=::="'~ HP20 ~- + H+

K4

... (2)

... (3)

.... (4)

HP20 ~- OOC ... P20 :- +H + .. . (5)

K, and K2 to be 2.0 and 0.3 mol dm-3 respectively estimated earlier were less by a factor of 20 as reported by Venturini el a/." and also by Gupta and co-workers5. Rao el a /. 12 reported these values to be ionic strength-dependent. Edwards el al." estimated the values of KJ and K4 spectrophotometrically to be (6.6 ± 0.3) x 10-6 and (2. 1 ± 0.1) x IO-

g mol dm-J

respectively . If the effect of pH on the rate of the reaction is taken into account, all the species viz. H3P20 g- , H2P20 g

2- and HP20 g

3-are reactive with varying degrees of reactivity towards the substrate. However, H2P2ol- appears to be the highly reactive form of peroxodiphosphate and thi s species is predominantly present in the reaction system in the pH range employed.

Since the rate is independent of [diol] , the rate controlling step must be an interaction of the catalyst with the oxidant 13 . Also, the complexation of silver(l) by peroxodiphosphate on the pattern of known complexation of Li+, Na+, K+ and Mg2+ by peroxo-diphosphate in solution is now well es tablished 'o. Further, sil ver(l) is also known to catalyze peroxodi-sulphate decomposition through the rate-limiting decomposition of si lver(l)-peroxodisulphate complex formed in the preceding fast equilibrium step. Since peroxodiphosphate and peroxodisulphate are isoelectronic and isostructural, it is expected that the decomposition of the fo rmer should occur in a similar manner. Further. the rate increases with increase in concentration of sodium acetate. It is, therefore, apparent that the acetato-si lver(I) species catalyzed decomposition of peroxodiphosphate is more facil e than the aquo-silver(l ) species catalyzed decomposition.

Considering these observations along with other experimental results, the fo ll owing mechanism in Scheme I can be envisaged to account for all the reaction events.

H2P20 S2-

Ag' ooc ... [Ag. H2P20 gr K' Ag

K MeCO;-

+ other products

k' Ac

Fast Agil + CH20H. CH2• CH20H )

Ag' + CHOH.CH2.CH20H + H+ ... (6)

Fast Agil + CHOH.CH2.CH20H )

Ag'+CHO. CH2. CH20 H + H+ .. . (7)

Scheme 1

If this mechanism is coupled with the equilibrium step (3) , the loss of pdp leads to the rate law (8)

_d.::..:..[ p--.:d p~] = dl

(k:g K:g + k:c K:c K[MeCO ; ])[Agi ]T [pdp] T [H +] (I + K[MeCO ; ]) (K J + [H +]

.. . (8) Since K3 is reported 'O to be significantly small , it is

much less as compared to [H+] in the denominator of Eq. (8). Thus the inequality [H+]» K3 reduces the rate law (8) to (9) or (10).

d[pdp] (k:g K:g + k:cK:cK[MeCO; ]) [Agi ]T[pdp ]T ---=

dt ( 1 + K[MeC0 2 ]) ... (9)

or

k:g K: g + k :cK :c K [MeCO; ] k = --=----=----- - --- ... (10)

1 + K [MeCO; ]

where k is an observed second order rate constant. Also, the equilibrium constants K 'Ag and K' Ac, are small. Hence rate law ( 10) can be written as in Eq . (1\ )

kAg +k Ac K [MeCO; ] k = .. . (I I )

1+ K [MeCO ; ]

where kAg = k'Ag K'Ag and kAc = k'Ac K'Ac Since the formati on constant K of the acetato-

silver (I ) species is reported'S to be 5.4, a plot of

718 INDIAN 1. CHEM., SEC A, lUL Y 2001

k( I+K[MeC02- ] versus [MeC02- ] was made. It was linear with a non-zero intercept (Fig. I). The values of kAg from the intercepts were found to be (2.S±0.2)x 10-2, (4.0 ±0.2)x 10-2 and (6.4± 0.4)x 10-2, and kAc K from the gradien ts were to be ( 1.1 ± 0.1), (2.2 ± 0.2) and (3.2 ± 0.2) dm 6 mol-2 s- ' at 4S, SO and SsoC respecti vely at 1= 2.0 mol dm-J . Such plots justify our earli er proposition that the contribution of HP20 g

J- species as compared to that of H2P20 S

2-

species is negligible. Had it not been the situation, the fu nctional plot as in Fig. 1 would have certainly shown deviations from linearity.

The catalytic role of sil ver(l ) in redox systems has been reported either via Ag'/AgO (ref. IS ), AgII/Ag' (ref. 16) or Ag" '/Ag' (ref. 17). In case of Ag'/Ag", the reac ti on system exhibits transition from homogeneous to heterogeneous state which has rarely been observed. Since the system remains homogeneous throughout the reaction , thi s can be considered as an evidence to negate the operation of Ag'l AgO redox cycle of the catalyst. Further, Ag'1I1 Ag' redox cycle can be rul ed out on the premise that Ag'lI in aquo-state does not ex ist in acidic solutions unless stabilized by a complex ing agent. As no such complex ing agent is

. h . . A 'lilA' d I present In t e reaction mi xture, g g re ox cyc e is also rul ed ou t.

H ' f A III· '11 'd d ' 9 b owever, I g IS Stl consl ere to e an intermediate based on Eq. ( 12).

Ag' + P20 g ~- ) AglIl + 2PO/- .... (12)

it can immediately change to AglI via fast interaction with Agi in solutions of low [H+] or disproportionate

25.0

20.0

Itb ~10.0

5.0

00 2.0 4.0 6.0 8.0 10.0

1 cf[CHj:OONo). mol dm3

Fig. 1- Va ri:lIi on of sod ium acct

SINGH el 01.: OXIDATION OF I,3-PROPANEDIOL BY PEROXODIPHOSPHATE 719

5 Rao R J N & Rao P V K. Illdiall J Chem. 2 1 ( 1982) 1066. Gupta S, Gupta K S & Gupta Y K. J chelll Soc Dalloll Trails, ( 1984) 1873 and the references cited therein .

6 Anderson J M & Kochi J K, J Alii chelll Soc, 92 (1970) 165 1 and references cited therein .

7 Kapoor S, Sharma P D & Gupta Y K. Tolallfa 22 ( 1975) 765. 8 Hei1born I, Cook A H, Bunbury H K M & Hey D H.

Dicliollary of orgallic cOlllpollllds (Ey re and Spotlis Woode. London) 1865. p. 1775.

9 Gupta C. Mishra S K & Sharma P D. Tn llls lIIel Chelll. 19, ( 1994) 6569 and the references cited therein .

10 Crutchfied M M & Edwards J 0 , J Alii chelll Soc, 82 ( 1966) 3533.

11 Venturini M. Indelli A & Raspi G, J e1eclro- lIIol Chelll . IlIle/.facial £ Iecrrochelll. 33 ( 197 1) 99.

12 Rao I, Ph. D. Dissertation, University of Rajas than. Jaipur. 1992.

13 Huyser E S & Rose L G , J org Chelll, 37 (1972) 85 1. 14 McDougall F H & Topel L E. J phys Chelll , 56 ( 1952) 169.

McDougall F & Peterson S. J phys Col/aid Chelll, 5 1 (1947) 1346.

15 Agrawal A, Mishra S K & Sharma P D, o.ridll COI/II/IlIII . 5. (1992) 3.

16 Walling C & Camaion D N, J org Chelll . 43 (1978) 3226. 17 McMillan G A, Chelll Rev, 62 ( 1962) 65. 18 Gordon B N & Wahl A C, J Alii chelll . Soc. 80, ( 1958) 273.

Jirsa F Z, Allorg AI/gem Chelll. 148 (1925) 130; 188 (1926) 61.

19 Morgan G T & Burstale F H , J chelll Soc. (1960) 2594. 20 Gangwal C M, Ph .D. Dissertation, University of Rajasthan.

Jaipur, 1995.