Nadar, P. Ananthakrishna

Shanmugasundaram, A.

English

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NOTESThe R"NI salts (Mis Distillation Products Indust~ies,USA) were purified in the usual manner!': Solutionswere prepared on molal basis and tile density wasdetermined with a dilatometer of about 40 ml capa-city, the stem of which was graduated in 0·01 ml sothat volume could be read with ea.se up to 0·003 mlwithin an error not exceeding 0'002 ml with amagnifying glass. Rest of the experimental l;ro-cedure was the same as given in earlier communica-tions from this laboratory. From the density dataat 40°, apparent molal volume was obtained in theusual manner. Only Me4NI, Er4NI, Pr4 NI andBu"NI could be studied since the higher h~mologueshave little solubility in this solvent medium.The linear plots of cpo versus y'C (Fig. 1) indicatethe applicability of the Mason's equation, cpo = CPo+Svv'C , to these solutions within the concentrationrange studied. The slope of the curves for. Me4NIand Et NI is positi ve in all the five mixtures.Howeve~, for Pr4NI, it is negative in mixtures con-taining 9·47 and 19'84% methanol and beco?;espositive at higher concentrati.ons (~fmethanol. . ~heslope is negative for Bu4NI in mixtures containing9·47 .19·84 and 34'48% methanol and then becomespositive. Comp.aring th~ results ?btained. in th2present study With those in water-dioxane rr~lxtures ,it appears that change over of the negative sloJ~eto a positive one, occurs at E c::: 60 for Bu~NI mboth water-dioxane and water-methanol mixturesand for Pr4NI, it appears at E c::: 63-64 in water-methanol and at E ~ 69 in water-dioxane mixture.In view of the approximate nature of these estimates,it is justified to assume that chan~e over from anegative to a positive slope for a given sal~ occu~saround the same approximate value of the dielectricconstant. The change in the nature of the mole-cular species (dioxane is non-pola~- and with?ut hydro-!Yenbondinz whereas metbanol is polar WIth hydro-gen bond interaction). does not se:em to. have ,anyappreciable effect. It IS only the (lIe~ectnc c0!ls .antof the medium which matters. 'WIth the limiteddata in a few solvent systems, it is not possible tomake any generalization.The authors thank Head, Chemistry Department,for laboratory facilities. They are also th~nk~ulto the CSIR, New Delhi, for the award of a Jumorreasearch fellowship to one of them (R.P.) and .tothe Society of the Sigma XI, USA, for financialassistance.References1. R!\JENDRA. K., Ph.D. thesis, Lucknow Un ivor sj tv , 1973.1. REDLICH, O. & MEYER, D. M .. Chern, Rev:. 64 (1964), 221.3:/ HARNED, H. S. & OWE~. B. 13., Phys}cal chemistry ofelectrolytic solution (Reinhold. New york). 1958. 362.4. DUNN. L. A., Trans. Faraday sc«, 64 (1968). 1898.5. DUNN, L. A .. Trans. Faraday ss«, 67 (197}l. 2525.6. GOPAL, R. & SINGH, K, Z. phvs. Chem. .• ,5 (1971). 217.7. PADOVA, J. & ABRAHAMEE .. [, phy«. Chem .. 71 (1967),2112. _8. GOPAL, R .. AGARWAL, D. K. & KlDI.\R. R .• Z. phys. Chc.n.,84 (1973). 141 ,9. GOPAL, R., AGARWAL. D.1\:. & KUMAR. R, Bull. cuem,Soc. (japan). 46 (1973). 1973.. .10. ANDERSON, K. P .. BUTLEH, E. A. & "·OOLLEY. E. 1\1..J. phys. Chon .. 77 (1973), 2652'11. CONWAY, B. E.. VERRALL. R. }<,. & DEI'WYERS, J. E.,Trans. Faraday Soc .• 62 (1966), 2738.Kinetics of Reaction Betweeno-Substituted Benzoate Ions &Phenacyl BromideP. ANANTHAKRISHNA NADAR & A. SHANMl:GASUNDARAMDepartment of P.G. Chemistry & Research CentreV.H.N.S.N. College. Virudhunagar 626002Received 31 October 1975; accepted 4 March 1976The rate of reaction of phenacyl bromide with variouso-substituted benzoate ions has been measured in 90%(v/v) acetone-water mixture at 30°, 35° and 40°. Arr-henius parameters as well as the enthalpies and theentropies of activation have been evaluated. The re-action does not obey the Bronsted relationship. Theisokinetic temperature unexpectedly is found to be wellbelow the ambient temperature. Though the behaviourof o-substituents is correlated with their polar effectsby the Taft equation, it appears that some substituentsbehave quite unexpectedly. This may be due to thefact that the values of k are all obtained above the iso-kinetic temperature. The polar substituent constants,(J~, for SCH3 and SO,CH3 groups are evaluated by inter-polation in the Taft plot. Variation in ASt has beenexplained in terms of the bulk effect of the groups.KINETICS of reaction of m-, p- and a fewa-substituted benzoate ions with phenacylbromide have been recently reported by Rout andcoworkers", But a systematic study on theinfluence of o-subst ituents has not so far been madein this reaction and thus a study to understand theeffect of o-snbstituents in this reaction appearednecessary.The kinetics of the reaction between a number ofa-substituted benzoate ions and phenacyl bromideare presented in this paper.The kinetic results are summarized in Table 1.Based on the results obtained by AnanthakrishnaNadar and Gnanasekaran- in the reaction of6-substituted w-bromo-2-acetonaphthones withbenzoate ion, we have proposed transition state (I)in this reaction.H H8- """/ 3-RCOO C BrICOIPhIThis shows that the attack is at a point remotefrom o-substituents in the benzene ring. Thismeans that the primary steric effect" is absent inthis reaction and its contribution to the total effectof the substituents is also absent. But the secon-dary steric effect+ of the o-substituents may beexpected to make contribution. Presumably conju-gation between ring and carboxy side-chain willstabilize the ester product and some sign of thismight appear in the transition state and besubject to steric inhibition which will lead toretardation. Therefore, the major effects of805INDIAN ]. CREM., VOL. 14A, OCTOBER 1976-0·8o 0:*oFig. 1 - Taft plot at 4000'8o-substituents that arc expected to influence the ratemay be their polar and secondary steric effects.A plot of log klkMe at 40° against (j~ values islinear (Fig. 1) giving a value of -0'63 for the polarreaction constant (P*) with correlation coefficientr = 0·970. This indicates that the reaction con-forms to the Taft linear polar energy relationship".The SCH3 and S02CH3 groups are assigned O"ri valuesof 0·18 and 0·89 respectively by interpolation in theabove plot.It is difficult to say with certainty that the abovementioned two effects alone operate in t.his series.A pattern in which the o-chloro compound reactingalmost close to the parent. compound and theo-methyl compound reacting much faster than theparent compound cannot possibly fit with this.The Taft plot is effectively a three point correlationand its validity can also mean that all the substi-tuents in question show approximately the samesteric effect as methyl, and this is a steric accelera-tion for the replacement of H by any other group.Indeed, it is difficult to imagine that this reactionis subject to steric acceleration. That the substi-tuents do not exhibit their normal steric effect isknown from the very poor correlation in the plot oflog klkMe against Es, the steric parameters ofo-substituents. Therefore, it is a situation in whichthe substituents (at least some) behave quite un-expectedly and this prevents us from making amore quantitative approach to the problem.The cause for this unusual behaviour may pos-sibly lie in the isokinetic t emperaturcv B. As ex-plained by Exners-", the isokinetic relationship inthis reaction has been examined through a plot oflog kl versus log k2 where k) refers to the rate con-stant at temperature T1 and k2 to the constant attemperature Tz, T', being greater than T2• If wetake T', = 313°K and T2 =303°K in this series, aplot of log kl versus log k2 is fairly rectilinear with a .slope b 0= 1'18~ (by least squares). This gives~= 258 K, WhIC~ IS below the lowest experimentaltemperature. ThIS may be due to the negative p*value and the negative relative enthalpies of activa-'tion (AAHt) observed for many compoundss. Whenthe rate constants are obtained above the isokinetictemperature, their pattern for the different substi-t,:ents is difficult to interpret and it would be quitedifferent at temperatures well below the isokinetictemperatures.The plot between rates and PK values of theconjugate acids is not linear indicating the failureof Bronsted relationship'? in this reaction. Thisfailure may be due to different effects of o-substi-tuents separately influencing rate and acidity.Substituent effects on ASt - Though the patternof 109 k values poses problem to attribute it to anydefinite effect, the variation in ASt can be bestexplained by the bulk effect of the substituents.A bulky substituent near the reaction centre willreduce the number of energy levels available tothe transition state relative to those in the initialstate and this renders L\st more negative-t. Thedata in Table 1 reveal that the AS: values for allthe o-substituted compounds are lower than thatfor the unsubstituted compound. As the size ofthe group increases AS: becomes more negative.The methyl and the bromo substituents haveapproximately the same van der Waals radius-"and have nearly the same bulk. The AS: . valuefor o-~romo compound is found to be slightly morenegative as. compared to that for o-methyl com-pound. ThIS may be due to the increase in effec-tive radius of the reaction site as a result of in-creased solvation caused by the negative end of thecarbon-bromin~ dipole thereby increasing the bulkeffect of bromine. Even though chlorine has sameven del' Waals radius'" as methyl, it exhibits un-expectedly lesser bulk effect than Me.B.enzoic acid and o-toluic acid were commerciallyavailable and were recrystallized from gla. aceticacid. o-Iodobenzoic acid was prepared from anthra-nilic acid and purified by repeated sublimationunder reduced pressure and crystallization fromwater. o-Chloro, o-brorno and o-nitrobenzoic acids:-vere prepared by the oxidation of the correspond-mg toluenes. o-Methylsulphonylbenzoic acid wasprepared by hydrogen peroxide oxidation ofo-methylmercaptobenzoic acid. All the acids hadTABLE 1 - RATES OF REACTION BETWEEN O-SUBSTITUTED BENZOATE IONS AND PHENACYL BROMIDEo-Substituent loa X k (litre mole-I see-I) Ea 6.Ht 6.St log PZkeal mole-' keal mole"! e.u.300 350 40°R 8·5 12·6 23·7 20·0 19·4 3·8Cl 8·0 11·5 22·3 20·0 19·3 4·4Br 8·4 12-4 22·6 18·7 18·1 8'4I 10·8 18·6 27·5 17·3 16·7 12·3CRa 16·5 30·2 45·0 18·7 18·1 6·9SCRa 12·1 22·3 33·1 19·2 18·6 5·8SO,CR, 5·2 10·7 11·0 14·6 14·0 22·7NO. 4-6 7·1 10·4 15·8 15·2 19·280612·412·311-410·511·712-08'39·0m.ps. identical with, or very close to those recordedin the literature.Analar acetone was purified by the method ofSachs-! and diluted with triply distilled water.Phenacyl bromide and sodium salts of the acidswere prepared by well-known procedures.Rate measurement - The solvent used was ace-tone-water 90% (v/v). The temperature was con-trolled to within ± 0'05°. The kinetic procedurewas as reported in Iiteratnrc'j'. It was convenientto use low concentrations (0·02M) of both thereactants for following the kinetics satisfactorily.The equat ion used to calculate the rate constants k(litre mole"! sec-1) was the second order rate expi essionwith the reactants present in equal concentration.The rate constants were reproducible within ± 2%.The authors thank the managing board of theV.H.N.S.N. College, Virudhunagar, for financialsupport and Dr J. Shorter, Department of Chem-istry, The University, Hull, England, for usefulsuggestions.References1. MISHRA, P., NAYAK, P. L. & ROUT, M. K., Indian J.Chem., 11 (1973), 452.2. ANANTHAKRISHNA NADAR, P. & GNANASEKARAN, C.,Indian J. Chem., (communicated).3. CHAPMAN, N. B., SHORT~R, J. & UTLEY, J. H. P., J.chem, Soc., (1962), 18~.4. BOWDEN, K., BUCKLEY, A., CH.AP~IAN, N. B. & SHORTER,J., J. chem, Soc., (1964), 3380.5. TAFT, R. W. (Jr), Steric effects in organic chemistry, editedby M. S. Newman (Wiley, New Yqrk) , 1956, chap. }3.6. EXNER, 0., Nature, 201 (1964), 488"; 227 (1970), 360":_/7. EXNER, 0., Colin Czech. chem. Conimun., 29 (1964), 1094.8. RITCHIE, C. D. & SAGER, 'V. F., Progress in physicalorganic chemistry, Vol. 2, edited by S. G. Cohen, A.Streitwieser (Jr.) & R. W. Taft (John Wiley & SonsInc., New York), 1964, 354.9. SHORTER, J., The University, Hull, England (personalcommunication).10. BENSEN, S. W., The foundations of chemical kinetics(McGraw-Hili, New York), 1960, 564. /11. IVANOFF, N. & MAGAT, M., J. chim. Phys., 47 (1950), 9rt.12. MOSER, c. L. & KOHLENBERG, A. 1., J. chem, Soc., (1951),W~ •13. BONDI, A., J. phys. Chem., 68 (1964), 441(14. SACHS, F., Ber. dt. chem, Ges., 34 (1901), 494.15. ANANTHAKRISHNA NADAR, P. & GNANASEKARAN, C.,Indian J. Chem., 13 (1975), 1237.Kinetics & Mechanism ofAlkaline Hydrolysis of SalicylamideRAIS AHMAD, M. NIYAZ KHAN & A. AZIZ KHANDepartment of Chemistry, Aligarh Muslim UniversityAligarhReceived 15 July 1975; accepted 24 February 1976The kinetics of the hydrolysis of salicylamide hasbeen studied under strongly alkaline medium. Theobserved pseudo first-order rate constants vary accord-. 1 B.Ing to the empirical equation kobs = B1+[OH-j atlower [OH-1 (O·7-1·9M) and an independent rate isobserved at hlgher [OH-] (1·9-3·0M). On the basis ofthese results, a plausible mechanism has been suggest-ed. The existence of dianionic reaction intermediateIs ruled out,NOTESKINETICS ?f hydrolysis of sali~ylami~e in acidicand alkaline media has been mveshgated by anumber of workers>e. In view of the postulateregarding the evidence of additional dianionic inter-rnediatcv" in acyl transfer reactions, it was consi-dered worth while to investigate the kinetics. ofhydrolysis of salicylamide in strongly alkalinemedium to ascertain the existence of additionaldianionic intermediate.Salicylamide (E. Merck) was used as such. Itsfreshly prepared solution in 0'5M NaOH was usedfor each set as no appreciable reaction occurredunder these conditions at room temperature. Otherchemicals used were of reagent grade. The reactionwas followed by estimating ammonia evolved spectre-photometrically by Ncsselerization mcthodlo-15.The effect of [salicylamide] on reaction rate wasstudied at 95° in 3'OM NaOH and the values ofrate constant kobs X 103 (min ") were found to be4'49, 3·70, 3'98, 3·35 and 3·62 at [salicylamide]= 0·008, 0'007, 0·006, 0·005 and 0·004111 respec-tively. Studies were also carried out at differenttemperatures and activation parameters were evalu-ated using Eyring and Arrhenius equations. Thevalues are: E; = 14·9 ± 0·6 kcal mole'", ASt= - 39·7 ± 1·8 e.u.; f).Ht = 14·2 ± 0·6 kcal mole:".The rate increased slowly with the increase in ionicstrength.The rate dependence on [OH-] was studied at95° in NaOH solutions of concentration rangingfrom 0·7 to 3·0M keeping ionic strength constantat 3 by adding NaN03 solution. It was found thatthe rate increased at [OH-] = 0·7 to 1·9M and wasindependent of [OH-] at higher concentrations(-1·9 to 3·0M). In the lower range the rate con-stants were found to follow the empirical Eqn (1)(Fig. 1).k~bS = n,+ [g~_] ...(1)The mechanism consistent with the observed resultsmay be represented by scheme l.As pKa value of salicylamide is of the order......,7-8,it will remain in ionized form in highly alka-0·60·5U<IIIII 0 t."'521} 0·30·20·'o 0·5 '·0 '·5[' -OH"jPlot ot _L vs J:.Kobs [OHjThe line is drawn to'om thecalculated values of rate constantsfrom equation (3). The points ar",the observed vatues of_rateconstants at vanoue [OHj.Fig. ,807INDIAN J. CHEM., VOL. 14A, OCTOBER 1976Ka1 Ka,OC.II.CONH.+OH- .= -OC8H,CONH2 .=-OC.H.COXH-+H20 .)/).,.,)1.Scheme 1line medium. The observed pH-rate profile indi-cates that the contribution to the rate through thedianionic tetrahedral addition intermediate is al-most negligible probably due to the presence ofionized phenolic group. The steady state treatmentto the reactive tetrahedral addition intermediate(D) leads to the kinetic Egn (2).k k 1(a1 [OH-121 J Kw .IKobs= (k_1+k?)(1+ ~: [OH-] + ~:1.~~[OH-]2)... (2)The Eqn (2) is further reduced to the Eqn (3)by applying the condition that1«~:[OH-] + ~:. ~: [on-J2)which is satisfactorily acceptable as Ka1is of theorder ,.....,10-8.k1kz(OH-]Kobs = (k_1+kJ( 1+~:-=-'[-O-H--]-"-)Eqn (3) is similar to the best fitted empiricalEqn (1) withB - (k1+k2) tc.; d1- kk 'K-an1 2 wAt relatively higher [OH-]1< Kat [OH-] is found toKwpronounced and henceto Eqn (4),K _ klk2obs - (k-l +k2)Ka.IKwwhich shows that the rate is independent of [OH)The authors thank Prof. "V. Rahman for provid-ing facilities, and the CSIR, New Delhi, for financialassistance to tW0 of them (KA. and M.N.K.).... (3)B _ (k_l+k~2 - klk2the conditionbe kinetically moreEqn (3) is further reduced... (4)References1/REID, E. E., Am. chem. j., 21 (1899), 284; 24 (1900),397.2. KOSHY, K. T., j. pharm. Sci., 58 (1959), 560.3. BRuI~E, T. C. & TAN:<ER, D. W., J. org . Chem., 30 (1965),166'8.4. DARKE, D., SCHOWEN, R. L. & JAYARAMAN, H., j. Am.chem. Soc., 95 (1973), 454.5. BIECHLER, S. S. & TAFT (Jr), R. "V., j. Am. chem, Soc.,79 (1957), 492ii.6. POLLACK, R. M. B: BENDER, ~1. L., j. Am. chem . Soc.,92 (1970), 71 CWo ,7. STAUFFER, C. E., J. Am. chem, Soc., 94 (1972), 1791.8. ERIKSSON, S. 0., Acta pharm. Suecica, 6 (1969), 1J9.808J,4.15.KHAN(_1. N. & KHAN, A. A., j. org. Chem., 40 (1975),179J.BUNTON, C. A., N,WAK, B. & O'CONl'!OR, C., J. org,Chem., 33 (1 '.168;, 51·2. .BENDER" lVI, L., GINGER, ~. D. & KElVIP, K c.. j. Am.cheni. Soc., 76 (1954). 3~.)0; 80 (19S8), ID44. /YAKOYAMA, F.,.f. phar ni. Soc. Japan, 63 (1943),,5.HALE, C. H., HALE, M. x. & JONES, "V. II., Anolyt,ci-.« .. 21 (1949), 15,49.LY:-1E, x. R., j. phys. cu;«, 69 (1965), 6137.BU:<To:-1, C. A., FARRER, S. J., M[LBANK, A. J. G.,O'CO:-1XOR, C. j. & 'feRNEY, J. A., J. chem, Sac, PerkinII, (1972), 1869,Kinetics of Oxidation of Cinnamyl Alcohol withChloramine- T in Hydrochloric Acid MediumD. S. l\IAHADEVAPPA & H. M. K. NAIDUDepartment of Postgraduate Studies & Research inChemistry, University of Mysore, Mysore 570006Received 15 July 1975; accepted 6 Ma'Ych 1976The kinetics of oxidation of cinnamyl alcohol bychloramine-T in the presence of hydrochloric acid(0·1-0·3M) has been studied at 0°. The reaction showsfirst order dependence on [chlorarntne-T], second orderin [H'r] and is independent of [substrate). The energyand entropy of activation of the reaction have beenfound to be 57·2 KJ mole= and -75·01 .TK-I respec-tively. The Arrhenius factor is 1·203 x 10· sec=. Ionicstrength and presence of p-toluenesulphonamide havevery little influence on the rate of reaction. The pro-posed mechanism assumes a stepwise addition of twoprotons to release the precursor chloriniwn Ion, whichon hydrolysis yields H.OCI+. The latter attacks thesubstrate in a fast step to form the reaction products,cinnamic aldehyde and p-toluenesulphonamide.thatIT was earlier! noticed that cinnamvl alcohol in50% aq, ethanol (v/v) underwent ready oxidationby chloraminc-T (CAT) in the presence of HCI(0·1 to O'3M'). The results of kinetic studies of thisreaction are reported in this note. Since the re-action is fairly rapid at room temperature (25°),the kinetic runs were made at 0°.Chloramine-T (E. Merck) was purified by themethod of Morris at al.2 and its aqueous solutionstandardized by the iodometric method. Stocksolutions of cinnamyl alcohol (Naarden, b.p. 256·6°)were prepared in 50% aq. ethanol (vJv). A concen-tratcd solation of sodium perchlorate was used tokeep the ionic strength constant. All other chemi-cals used were of analytical reagent grade. Theexperimental procedure has been described earlier'.Stoichiometry- Reaction mixtures containing ex-cess CAT over cinnamyl alcohol in the presenceof HCI (overall concentration, 0·5 to 1·0M) were keptat room temperature (25°) for 24 hr. Estimationof the unreacted CAT showed that one mole ofcinnamyl alcohol consumed one mole of CATcorresponding to a two electron change. Theproducts were identified as cinnamic aldehydeand p-toluenesulphonamide3.Kinetics of the reaction was studied at severalinitial [cinnamyl alcohol] and [CAT]. When the

Kinetics of Reaction Between o-Substituted Benzoate Ions & Phenacyl Bromide

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Kinetics of Reaction Between o-Substituted Benzoate Ions & Phenacyl Bromide

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