Jagdale, M. H.

Mane, B. S.

English

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Indian Journal of ChemistryVol. UA, December 1977, pp, 1086-1088Kinetics of Acid-catalysed Hydrolysis of Hexanohydroxamic AcidB. S. MANE & M. H. JAGDALEDepartment of Chemistry, Shivaji University, Kolhapur 416004Received 4 October 1976; revised 2 May 1977; accepted 12 August 1977Kinetics of the acld-catalysed hydrolysis of hexanohydroxamic acid has been investigated inhydrochloric, percnlorfc and sulphuric acids. The effect of catalysts, temperature, ionic strengthand solvent isotope effect have been investigated and the results show that the hydrolysis occursby an, AAC' mechanism giving hexanoic acid and hydroxylamine. The variation of reactionrate with [acid] has been satisfactorily explained by a two-parameter equatlon.·THE acid-catalysed hydrolysis of some aroma-ticl-6 and few aliphatict'? hydroxamic acidshave been studied. Though the hydrolysis ofhexanohydroxamic acid was studied-? in dil. aq.HN03, no kinetic study of the reaction has beenstudied. A detailed kinetic study of hydrolysis ofhexanohydroxamic acid in aq. mineral acids has nowbeen undertaken and a mechanism consistent withthe results has been proposed.Materials and MethodsHexanohydroxamic acid (K & K Labs, USA) wasrecrystallized from benzene, m.p. 62° (lit,!o 61·5-63·5°).HCI, HCI04, H2S04 used were of AR qua.lity. A1% solution of FeCI3.6H20 (BDH, AR) was preparedin distilled water containing 0·1M HCl.D20 (isotopicpurity 99·4%) was precured from BARC, Bombay.All other chemicals used were of AR grade.Kinetic measurement - All kinetic measurementswere carried out at 40° ± 0'05°. The hydrolysisrates were determined from the rate of disappearanceof the hydroxamic acid. Aliquots of the reactionmixture were withdrawn at regular time intervalsand the concentration of hexanohydroxamic aciddetermined colorimetrically- at 520 nm, The rateconstants were calculated using the integrated rateequation for first order reactions. Initial concen-tration of hexanohydroxamic acid in all kinetic runwas O·Ol2M.Results and DiscussionThe hydrolysis was found to be first order withrespect to both substrate and HCI. But in thepresence of excess hydronium ions, pseudo-firstorder rates vvere ensured.Effect of varying [acid] on reaciiow rate - Thehydrolysis rates were studied in aq. HCI, HCI04and H2S04 at 40° in the range 1 to 8M. The rateswere found to increase with increase in [acid] andattained optimum value at 5M HCI, 3M HCl04 and3.Ar H2S04 and then tended to decrease (Table 1).Rate maxima is also observed for acid-catalysedhydrolysis of benzohydroxamic acid and its deriva-tives2,3. It can be presumed in analogy witharnidest! that the protonation of hexanohydroxarnic1086is complete at 5M HCI, 3M HCI04 and Dr! H2S04,Activity of water goes on decreasing with increasein [acid]. Above the optimum rate, decrease inactivity of water outweighs the protonated speciesand so the rate of hydrolysis goes all decreasing athigher acid concentration.Effect of uarying temperature on reaction rate-The •• hydrolysis of hexanohydroxarnic acid wasstu.lied in 1·0.'11 HCl in the temperature range 35-55°(Table 2). The rate constants (k X 104): 0,45, 0·71,1·13 and 1·77 see'? at 35°,40°,45° and 50° respec-tively are average of duplicate runs and the averagedeviation from the mean is less than 3%. The valuesof activation parameters fall ill the range typicalof a reaction proceeding by A2-mechanism. Theentropy of activation is negative and falls in therange -15 to - 30 e.u., characteristic of bimole-cular mechanism-s. The values of ~EL ~H+ andfrequency factor A are 18·3±0·3 kcal mol-I,17·5 + 0·1 kcal mol:" and 3·1 X 108 sec-I res;Jec-tively~ .Zucker and Hamrnettw treated quantitativelyacidity dependence rates and grouped reactions intotwo categories depending on whether the logarithmof the pseudo-first order rate constants were lineareither with Ho (Ai-mechanism; rate determiningunimolecular decomposition of protonated species)or with log CH+ (Az-mechanism; rate determining at-tack on protonated substrate by a water molecule).We observed that plot of log ki against -Ho wasnot linear. However the plot of log hI against log(CH+) was linear up to 4M HCI with approximateunit slope (0·92) indicating that a bimolecularmechanism operates.Bunnett-s found empirically that for a largenumber of reactions, plots of (log ki+Ho) versusthe logarithm of the activity of water' (log aH,o) werelinear or very nearly linear and classified the reactionsinto three categories according to slopes (w values)of the above plots. The present data gave a valueof w = 7·3 in HCI. This value suggested that waterplayed the role of a proton transfer agent (w> +3,3). From w value we can say that the tran-sition state is more polar than the substrate-". Thelarge -j-ve v8.lue of w have been observed (+6·1 inMAl E & JAGDALE: HYDROLYSIS OF HEXANOHYDROXAMIC ACID:i- +TABLE 1- EFFECT OF VARYING [ACID] ON THE 0 +H30 I fast) OH/I IIRATE CONSTANT R-C-NHOH < " ?::"" + H2O[Acid] k, 10' see'? in -HpITClst I(M) OHHCl* HClO. H2SO. I [s low]R-y-NHOH1-0 0-71 (0-69) 0-68 0-77 68+2-0 1-41 (1-28) HO 1-46/ " ~ NH2OH+RCOOH3-0 1-90 (1-74) 1-30 2-17 H H4-0 2-02 (2-04) HO 1-67 : 6+ : 6+5-0 2-16 (2-12) 0-90 1-46 o· 06-0 1-97 (1-90) 0-65 1-05 /" I "-7-0 1-74 (1-81) 0-47 0-61 H H H H8-0 1-31 (1-49) 0-27 0-28 Transition complex*Values calculated with the help of Eq. 2 are given in·parentheses_TABLE 2 - VARIATION OF RATE CO:-lSTANT AT DIFFERENT[HC)] AND IONIC STRENGTHS[Ionic strength maintained by adding NaCI; temp. 40°][HCI] h, 10' sec? at fLM -----~--------------ml !-OM 2-0i11 3'OM0-10 0-07 0-090-25 0-18 0-21 0-230-50 0-36 0-41 0-45 O·sO0-75 0·51 a-58 0·62 0·691-00 0-71 081 0·901-00 1-28~1-50 HI 1·19 1·231-70 1-29 1-432-00 1-41 1-623-00 1-90HCI and +8·9 in HCI04) for acid-catalysed hydro-lysis of N-phenyl-n-butyrohydroxamic acid".Bunnett and Olsen-" suggested that the slopes(1)) of the plot of (log kl +Ho) versus (Ho+log H+)can be used as a criterion postulating for a reactionmechanism. The (1)) values obtained for hydrolysisof hexanohydroxarnic acid are 1·35 in HCI, 1·25 inHCI04 and 1·2 in H2S04 which fall clearly in thethird category implying that water acts as protontransfer agent (1) > 0·56)_ The 1> value observedfor benzohydroxamic acid and its some p-substitutedderivatives have been reported- to lie in the range0·84 to H.Bunnett and coworkers-" suggested that for A1-reactions the catalytic effect of added acids followedthe order: HCI04 > HCl > H2S04, while forA2-reactions the order was H2S04> HCl > HCI04•Thus higher rate in H2S04 c- HCI > HCl04)provides support for a bimolecular mechanism.Solvent isotope effect - The hydrolysis of hexane-hydroxamic acid in H20 and D20 using 1·0M HCl at40° showed that the hydrolysis occurs 1·8 timesfaster in D20 than in H20 (Table 2). Higher ratein D20 is suggestive of a rapid transfer of protonto the substrate in the reaction and pre-equilibrium'step therefore is not the rate determining stept?As suggested for amide hydrolysis18,19 seat ofprotonation in hydroxamic acids may be assumedto be the oxygen atom".where R=CH3(CH2)4Scheme IEffect oj ionic strength - The hydrolysis of hexa-nohydroxamic acid studied at varying [HCl] andionic strengths [1·0, 2·0, 3·0M (NaCl)], showed onlyslight increase in the rate constant (15-20%) withincreasing ionic strength (Table 2). From the plotsof rate constants against acidity at different ionicstrengths, it can be concluded that: (i) at each ionicstrength the reaction can be represented by Eq. (1)k = kH+.CU+ __.(1)where k and KHt are the observed and specific rateconstants respectively; (ii) since the slopes (kH+) ofthe curves increase with increase in ionic strength,the reaction is presumed to exhibit positive salteffect. The linearity of the plots with positiveslopes indicates an acid-ca talysed reactionw; (iii)the reaction due to neutral species is absent as theplots pass through the origin21•These results favour the reaction proceeding viathe protonated species only. The rate data maybe represented by a modified form of the Bronsted-Bjerrum equation=.k = kH~_CII+eb_lL(aJl,o)" (2)where aH,o is the activity of water23, n is the numberof water molecules and (1. the ionic strength.The values of ku+ at (1. = 1·0, 2·0 and 3·0 are 0·12,0·77 and .0·81 respectively. The values kH~ andb are obtained from the linear plot of log kH+ versus(1.. The gradient gives b while the intercept at zeroionic strength, gives kH~. The value of b is foundto be 0·01_5 and that of kHt is 0·72 X 104M-I sec-'._ lI-!echamsm oj hydrolysis - The above resultsindicate that water molecules are involved in therate determining step. The number of water mole-cules actually involved can be obtained by themethod proposed by yates and l\IcClelland24•T~e new acidity function u, (= MHo) can beobtained from protonation of some p-substitutedbenzohydroxamic acids''. The value of Jl1 was inthe range 0·58 to 0·68_ The average value of Mis taken as 0·60._The plot of (log .kltHs! versus log aH,o was linearWI th slope = 3·1, indicating that 3 water moleculesare _involved in the formation of transition complexduring th~ hydrolysis of hexanohydroxamic acid.On the baSIS of above results the probable mechanism(Scheme 1) appears to involve reversible rapid1087INDIAN J. CHEM., VOL. 1SA, DECEMBER 1977protonation of hexanohydroxamic acid followed by aslow reaction of water with the protonated form,leading to the transition complexs! and subsequenthydrolysis giving the products, hexanoic acid andhydroxylamine.AcknowledgementOne of the authors (B.S.M.) is grateful to theUGC, New Delhi, for the award of a junior researchfellowship. The authors are indebted to Dr D. C.Berndt and Dr J. G. Tillet for valuable sug-gestions.References1. BERNDT, D. C. & FULLER, R. L., J. org. Chem., 31 (1966),3312.2. BUGLASS, A. J., HUDSON, K. & TILLET, J. G., J. chem,Soc., (1971), 123.3. AHAMAD, A., SOCHA, J. & VECERA, M., csu« Czech. chem.Commun., 39 (1974), 3293.<4.BERNDT, D. C. & WARD, 1. E., J. org. Chem., 39 (1974),840, 841.5. BERNDT, D. C. & WARD, 1. E., J. org. Chem., 41 (1976),3297.6. SUBBARAO, C. & TANDON, S. G., Indian J. Chem., 14A(1976). 765.7. MUNSON, J. W. & CONNORS, K. A., J. Am. chem. Soc., 24(1973), 1979.8. BERNDT, D. C. & SHARP, J. K., J. org. Chem., 38 (1973),396.10889. MANE, B. S. & JAGDALE, M. H., J. Indian chem. Soc.,in press.10. OBUKHOV'A, ~L E., VODEN. V. G. & PUSHLENKOV, M. F.,Radiokhimiya, 12 (6) (1970), 900.11. EDWARD, J. T. & MEACOCK, S. C. R., J. chem, Soc., (1957),2000. ..12. SCHLEGER, L. L. & LONG, F. A., Adv. phys. org, Chem .•1 (1963). 24.13. ZUCKER, L. & HAMMETT, L. P., J. Am. chem, Soc., 61(1939), 2791.14. BUNNETT, J. F .• J. Am. chem. Soc., 83 (1961), 4956,4968, 4973, 4978.15. BUNNETT, J. F. & OLSEN, F. D .• Can. J. cu«; 44(1966), 1899.16. BUNTON, C. A., CRABTREE. J. H. & ROBINSON, L.. J.Am. chem. Soc., 90 (1968), 1258.17. BUNTON, C. A. & SHINNER, V. J., J. Am. chem. Soc., 83(1961), 42, 3207.18. YATES, K. & STEVENS, J. B., Can. J. Chem., 43 (1965),529.19. CHARMINA O'CONNOR, Q. Rev., 24 (1970), 553.20. BUNTON, C. A., MHALA, M. M., OLDHAM, K. G. & VERNON,C. A., J. chem. Soc., (1960), 3293.21. MHALA, M. M. & JAGDALE, M. H., Indian J. Chem., 8(1970), 147.22. LONG, F. A. & McDEVIT, W. F., Chem. Rev., 51 (1952).119.23. LEFFLER, J. E. & GRUNWALD, E., Rates and equilibriaof organic reactions (John Wiley, New York), 1963,286.2<4.YATES, K. & MCCLELLAND, R. A., J. Am. chem, Soc., 89(1967), 2686.25. KEITH YATES, Ace. chem. Res., 4 (1971), 136.

Kinetics of Acid-catalysed Hydrolysis of Hexanohydroxamic Acid

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Kinetics of Acid-catalysed Hydrolysis of Hexanohydroxamic Acid

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