Manjrekar, S. K.

Soman, S. V.

English

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Indian Journal of ChemistryVol. 15A. September 1977. pp. 781-784Kinetics of Aldehyde Oxidation on Platinum Anode In AqueousPerchloric Acid & Sulphuric AcidS. K. :\lA;'\JREKAR & S. V. SOllIANDepartment of Physical Chemistry, Institute of Science. Bombay 400032Reu!ived 13 October 1976; accepted 10 Alarch 1977Kinetics of aldehyde oxidation on platfnurn anode surface has been studied in aqueous HeIO.and sulphuric acid. The correlation between first order rate constants and acidities of thesolutions has been worked out. The values of the parameters Wo and IjI (wo and ljIo in this case)have been evaluated on the lines prcvious ly suggested. A reaction mechanism involving proto-nation of the aldehyde, reaction of the protonated form with oxygen and finally reaction withwater to yield the products has been suggested,INthe presence of mineral acids like HCl, H2S04or HCl04 many chemical reactions exhibit achange in the reaction rate or in the equilibriumas the [acid] is changed. Hammett ct ap·2 suggest-ed that these effects can be explained on the basisof the concept of acidity functions. Accordingto these authors, if the rate constant Ie varied linearlywith the acidity fnn ct ion , Hu' t hr- cnllv('rsinll of t hr-protonated substrate SH+ to the transition statedoes not involve interaction with water. If lug kvaried linearly with log [HX] the same transforma-tion required several molecules of water. Thepossiblity that the transition state in case of someacid-catalyscd reactions may involve participationof several molecules of water was also suggested byListen3.In a different approach Bunnct.t? studied a largenumber of hvdrolvt ic reactions (over 150) andfound that th(~ plots of log !?'¥+ H 0 versus log aH,Owere linear. Slopes of these plots provided a newparameter TV and the different numerical w valueswere suggestive of different reaction mechanisms.Bunnett and Olsen5,6 further observed that theeffects of changing the acidity of the medium on theequilibria or the rate constants can also be viewedfrom the free energy relationship. They suggestedanother approach where log k +Ho was plottedagainst Ho+ log [H+] instead of log aH,O' Theslope in this case was called rp, a parameter whosevalues appear to acquire more important signi-ficance than the I'v value".The equilibrium protonation of a substra~e. [~] isrepresented as S+H+~SH+ and the equIlIbrIumquotient Q=[SH+]/[S][H+]. Q is the reciprocal of theequilibrium constant for acid dissociation of SH+,log Q thus becomes the thermodynamic pKSH+ forthat substrate in that medium.It was shown by Bunnett and Olsen5,6 that ~inearfree energy relationship exists botweeen logarithmsof equilibrium quotients and rate .constants. Theslope rp of the linear plots charactenzes the .respon:,eof equilibrium quotient Q to the changing acidconcentrations whereas its intercept representsthe thermodynamic PK o 0 f the ba-e. It thus pro-vides a general method for estimating PKa of anyba-.« which undergoes protonation in moderatelyconcentrated mineral acids.This work has been planned to st udy the effectof the extent of prot onation on the rate of electro-lvt ic oxidation or reduction of aldehydes on polish-1"{1 platimll11 anod,> or ""thode in solut iort s of differ-ent acidities. Some important findings of thisin vestigation arc reported here.Materials and MethodsThe aldehydes studied were: (i) formaldehyde,(ii) acetaldehyde and (iii) propionaldehyde (all ofBDH, AR grade) and were purified if necessary byfractionation. Investigation was carried out ina cell where the anode and the cathode compart-ments were separate but were joined to one centralcompartment through sintered glass partition.Polished platinum electrodes (area 1 cm2) wereplaced in the compartments and once fixed wouldmaintain a constant position. Pure nitrogen bub-bling very slowly in each compartment maintainedan atmosphere of nitrogen and also stirred the solu-tions. The platinum foil electrodes were carefullypolished with a suspension of alumina in ethanoland then washed thoroughly with distilled waterhefore each experimental run. The reactions werestudied at various acid concentrations of (i) HCl04and (ii) H2S04, A direct current in a range of50 to 250 ma/cm? was passed through the electrolyticsolut ion ,Reaction mixture (,-...,100 ml] was taken in a cellfor every experiment. The system was thermostat-ed throughout the run. The rate of disappearanceof aldehyde in each electrode compartment wasmeasured separately. However, the resoponse ofthe reaction to the changing acid concentration and~he temperature, only in the anodic compartmentIS reported here. At known intervals, an aliquotsolution (1 ml) was removed from the electrodecompartment and analysed for its aldehyde con-tents. These were estimated colorimetric ally with781INDIAN ]. CHE;'\1., VOL. 15A, SEPTEMBER 1977MBTH (3-methyl-2-benzothiozolinone hydr.a~onehydrochloride monohydrate) which forms .a bnlh~ntblue cationic dye in the presence of ferric chl~r~deand sulphamic acid. This method is very se!lsltl\:efor determination of water-soluble aliphaticaldehvdes", The reaction was carried out at threediffer~nt temperatures, viz. 30°, 40° and 50° respec-t ivclv. The rate constants 'k' were calculatedconsidering the reaction to he of first order withrespect to [aldehyde].Results and DiscussionAn analv .-is of the results brings out the followingimportant" features. .(i] On passing current through the solutionaldehvde disappeared from both. the electrodecompartments and the rates of dlsappea~ance ofaldehyde (in both the compartments) was found tofollo\~' first order rate law.(ii) The rate constants were .found. to dependupon the [acid], Wit~ incrc<l:s~ ~n [aClCI]the ratewas found to increase m the initial stages onl~ .uPto a certain specific concentration of the acid whichdepended on the nature of the acid (Fig. 1). .The rate constant, thereafter, decreased Withfurther increase in [acid]. . .(iii] The rate constants ir:creased \'YIth increasein temperature but the [acid] at which the ratemaximum "vas observed was found to be the same.(iv) The energy of activation and the entropychanges of the reaction were also calculated and arelisted in Table 1.ACETALDEHYDE IN HCI04- ACETALDEHYDE IN H2SOl,." 1G~ ! I~I - -?-RO.LP-IO-N...•.AL-D-E-H~Y-DE~I."-· -H-2S~O:-4---t!=PR=O~P:;::IO-::NA;-;-LDEHYOEIN HClO,..~ i8 i'.lJ sOr~1/'\30 0! 01/t5612MOLES OF ACIO/UTI'IEFig. 1 - Dependence of rate constant on acid concentrationso782TABLE 1 - ACTIVATION PARAMETERS FOR THE ELECTRODEREACTIONS OF ALDEHYDESAcid t1Ht kcal/rnol at t1st e. u. at40° 50°FORMALDEHYDE13-16811·62510·37611·265-38,00-41·85-47·30-43·177ACETALDEHY DE10·5309·5849·8258·713-42·579-45·99+-44'917-48·795PROPION ,\LDEHYDE9·7639·7689·1649·292-37·758-39·193-39·778-40'154For explaining the reaction mechanism Bunnet t?:and othcrs-? considered t he change in the hydrationas a very important factor. The support for thischange in the hydration postulate was derived fromthe interrelationship between the entropy valuesand the w values. However, factors not relatedto solvation change are known to affect the entropyof activation. The rate maxima for specific concen-tration of the acid is also observed in several othercases like the carboxylic amide hydrolysis--. Insuch reactions it is recognized that the substrategets progressively protonated and the protonationis almost complete where the acid concentrationshows highest k value. Similar results have alsobeen reported by Rosenthal and Taylor-!A good deal about the mechanism of the reactioncan be learnt from its kinetics under various experi-mental conditions. Besides, a lot of additionalinformation can be obtained by determining therate of reaction for a series of reaction compounds.The results obtained during the anodic oxidationof three aldehydes (i.e. reaction taking place only.in the anode cornpartment) are presented in Table 2.In the analysis that follows the first order rateconstant is represented by symbol I.'. Its dependenceon the [H+], the acidity and the amount of water,is brought out in two independent \\'ays by plotting(i) log k + Ho against aH,O (Fig. 2) and (ii) log k+ n, against Ho+ log [W'J (Fig. 3). The slopesof these plots in the present case afe designated asWo and CPo (signifying the oxidative process) respec-tively which in the case of hydrolytic reactions weredesignated as It' and cp by Bunnett and 0Isen5•6.With increasing [acid] the k values rise smoothly,reach a miximum and then gradually fall off. Theabove-mentioned plots can thus be shown to consistof two distinct straight lines with a point of inflexionmarking the maximum k value (Fig. 1). Each ofthese plots will, therefore, have two slope values,One pertaining to the relation where h increasedwith [acid] and the other where k decreased withincreasing [acid]. Following earlier convention;these values of slopes are also designated as WOA,WOB and CPOA, CPOB, values for the oxidative reactions.The results are included in Table 3.MAKJREKAR & SOMAN: ALDEHYDE OXIDATION ON Pt A~ODE[Acid]M0·00·;;1·01·52·02·53·04·05·06·07·00·01·02·03·04·05·06·0TABLE 2 - DEPENDEKCE OF THE REACTIOK RATES OK ACID CONCENTRATIOKS(Current density, 50 mAjcrn-]Formaldehyde in H2S04(k X to-6) atAcetaldehyde in H2SO,(11 ,{ to-5) at30° 40° 500 30° 40° 50C22·5 45·1 88·1 10·8 21-16 32·524·8 53·3 89·3 13'7 24·0 35·231·1 60·0 93·5 16·1 26·3 47·722·0 46·6 75·0 18·5 32·6 61·718·6 36'6 65·1 15·0 26·3 45·017-6 35·0 55·6 13-6 22·7 38·515·0 32·6 52·3 12·1 19·7 32·59·0 26·0 41·0 6·5 14·0 25·04·0 18·0 32·0 1·0 10'5 19·012-0 22·0 4·0 14·55·0 12·0Formaldehyde in HClO. Acetaldehyde in HCIO.22·5 45·1 88·1 10·8 21·2 32·520·5 57·1 92·8 14·2 24·2 18·534·8 60·0 113-1 15·6 29·8 51·237-1 70·0 120·5 13·8 22·7 35·240·1 75·0 138'3 12·2 19·5 26·336·1 61·8 109'3 10·1 15·6 20·1330·8 60·0 101·1 g·7 13-8 17·0"801Propionaldehvdc in H.SO,(11 y. 10-5) at30°10·614·216·219515·8nil12·37·02·040°19·224·326·331·526·223·519·814'510·56·050°32·235·239·247·742·038·528·3no17·012·0Propionaldehvd e in HClO.10·612·515·013-711·610·08·,19·024·029·821·317·515-0no32-036·350·134·726·3no21·2I8'0r-!!I65~[I1- FORMALDEHYDE IN H2 salt2-ACET~.LDEHYDE IN H 2SCl~3-FORMALHYDE IN HelOt,.4--ACETALDEHYOE IN HCI04-3·S'------_-'·':-1 -------.-=-2-------'.73---'log QH20Fig. 2 - Variation of log k + Ho with aH,OBunnett and his coworkers+" on the basis ofgenerally accepted mechanisms for certain calibra-tion reactions considered certain values of wand4> as suggestive of the role of water in the reactionsequence. These values are given in Table 4.Reaction mechanism -- Based on our observationsthe mechanism of the reaction appears to involvethe steps shown in Scheme 1.(i) The rate of the reaction increases initially with[acid]. This suggests that the protonated form of26'S 30 J,J:+3-~01.ss-ol-FORMALDEHYDEIN H:1S04-2-FORMALDEHYDE~ HCI043-ACETALDEHYDE IN H2S04!4-ACE TA LDEHYDE IN HCI04-1I3 5~----n';:-------:-''::--------::'-::---'O'S 1·0 1·5Ho+109 [H+]Fig. 3 - Variation of log k + Ho with Ho + log [H+]the aldehyde whose concentration increases withincreasing concentrations of the acids, is one of thereaction species. Two tautomeric structures (A) or(B) of the protonated aldehyde are possible.(ii) The protonat od form reacts with oxygen,which could be either at the electrode surface orin the bulk of the solution if a suitable carrier foroxygen is present-".(iii) Since there is a smooth fall in the rateconstant after the specific acid concentration isreached the reaction appears to be influenced bythe decreasing concentration of water. Hence itis suggested that for the conversion of the product783INDIAN J. CHEM., VOL. 15A, SEPTEMBER 1977TABLE 3 - VALUES OF Wo AND .pOAT 300Acid .pOA .pOB WOA WOBFORMALDEHYDEHCIO. 1·725 1·25 17·2 6·6H2SO. 3·075 1·41 33·0 12·3ACETALDEHYDEHClO. 1·85 1·26 15·0 6·3H2SO4 2·40 1·56 25·0 13·8PROPIONALDEHYDEH2SO4 2·55 1-48 22·7 11·2TABLE 4 - WAND .p VALUES FOR SUBSTRATESPROTONATED ON OXYGEN OR XITROGENW values .p values Role of water in the ratedetg step-2·5 to 0·0 Water is not involved<0'0+ 1·2 to 3'3 +0'22-0·56 Water acts as a nu cleophile>+3·3 >0·58 Water acts as a proton transferagentin stage (ii) to the final product, water must be anessential constituent.It will be observed that the Wo and <Po valuesobtained by us in the anodic oxidation of alde-hvdes are numerically very different from values ofw' and <p reported by Eunnett for the hydrolyticreaction. It is likely that for the oxidative processwhich we have studied these norms of wand <Pvalues suggestive of the role of water may be higher.Further experimentation with new systems andsimilar analysis of the cathodic disappearance ofaldehydes are necessary to throw more light on theutility of this approach.784+ O~-H~R-C-:/""H[AJ+ ..~R-c~R-HHtB] .... (i I(i-HR-C"'" +.......HO-H~~:~ R-C~-·':-·--- ...o"'-w -···(ii)+ 6.....-0- H -6 ----"R-C -----_0 +H 0<,H-- 2"'-+ ...,-o-HH30 + R--C~OSchemel.., (iii blJl.eferences1. ZUCKER, L. & HAMMETT, L. P., J. Am. chem, Soc., 61(1939), 2791.2. HAMMETT, L. P., Pbysicat organic chemistry (McGraw-Hill, New York), 1940, 273-77.3. LISTEN, J. B., Chemy Ind., (1959), 397.4. BUNNETT, J. F.,J. Am. chem. Soc., 83 (1961). 4956.5. BUNNETT, J. F. & OLSEN, F. P., Can. J. Chem., 44(1966). 1899.6. BUNNETT, J. F. & OLSEN, F. D., Can. J. cie«, 44 (1966).1917.7. jVIACAGNO,V. A. & GIORDANO, M. C., Electrochem. Acta.14 (1969). 335.8. SAWICKI, E., HAUSER, T. R., STANLEY, T. \V. & ELBERT.W., Analyt. cu-«, 33(1) (1961). 93.9. BUNNETT, J. F., J. Am. chem, Soc., 83 (1961). 4973.10. MANJREKAR, S. K. & SOMAN, S. V., Indian J. Chem .•12 (1974). 307.11. EDWARD, J. T. & YIEACOCK, S. C. R., J. chem, Soc., (1957).2000.12. ROSENTHAL, D. & TAYLOR, 1. T., J. Am. chem, Soc., 79(1957), 2684.

Kinetics of Aldehyde Oxidation on Platinum Anode In Aqueous Perchloric Acid & Sulphuric Acid

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Kinetics of Aldehyde Oxidation on Platinum Anode In Aqueous Perchloric Acid & Sulphuric Acid

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