Indian Journal of ChemistryVol. 23A, March 1984, pp. 192-196

Acid Catalysed Hydrolysis of Methyl & Ethyl Acetates In Presence of Urea& Substituted Ureas

P K DASGUPTA. I' K BHATTACHARYA & S P MOULIK *Department of Chemistry. Jadavpur University. Calcutta 700032

Receii cd ~7 Ju!v 1983: rerised and accepted 25 October 1983

The rates of acid catalysed hydrolysis of methyl acetate and ethyl acetate decrease.in the presence of urea. substituted ureas.formamide. acetamide, acetone and dioxane. However, thiourea does not have any effect. Analysis of rate data reveals thatthere is no regularity in the rate alteration with respect to the effective dielectric constant of the medium. In equidielectricmedium, acetone decreases the rate more effectively than dioxane. This. and the fact thai thiourea has no effect show that theadditives deactivate the catalyst, H' ion, by forming a complex. and the potential binding centre is the oxygen of the> C =0group. The amino groups in urea. substituted ureas and thiourea cannot inhibit the H + ion. Although urea and substitutedureas seriously affect the rate, they do not alter the: energetics of the process showing that the activated complex between theester and the catalyst. H + ion, is not energetically influenced by them. The kinetically evaluated activities of the H + ion areclose to those reported by Bull I!/ al. from conductance and pH measurements [Arch Biochem Biophys, 104 (1964) 297]. Theequilibrium constants of the interaction, U + H + ~UH + have been calculated and found to agree with those calculated fromBull 1'1 al.'s data and not with those reported by Schiifer [Ber dt chem Ges., 80 (1976) 529].

Urea is a very weak basel (PKb = 13.82)and accepts H +

ion with an association constant 1.5 dm ' mol -1. Ifurea is present in a fair concentration in a reactionsystem in acidic medium, such a weak interaction ofurea with H + ion is likely to seriously affect thereaction equilibrium and kinetics, In aqueous medium,urea has low viscosity B-coefficient 2

, it denaturatesprotein+", deaggregates surfactant micelles" -7,

decreases hydrogen ion conductances and increasesthe pK of a weak acid":". These properties of ureasuggest that it breaks water structure, disruptshydrophobic assembly and reduces hydrogen ionactivity. Physical evidences also show that in aqueousmedium urea is involved in self-association 10.

Although these effects can be experimentallydemonstrated, the mechanism of action is littleunderstood, The influence of urea on the pK of aceticacid was reported by Bull et al," and by Schafer? withconflicting results, Schafer tried to interpret the resultsassuming dimerisation of urea 10. It will be shown laterthat this interpretation is only apparent. Theinteraction of urea with H + ion is expected to occureither via the nitrogen of - NH2 or via the oxygen of>C =O,and a study with various substituted ureas willhelp to distinguish between these two possibilities.

Bull et a/. as well as Schafer monitored the action ofurea through the dissociation equilibrium of aceticacid. We have used a kinetic method, Acid catalysedhydrolysis of methyl and ethyl acetates has now beenstudied in the absence and presence of urea andsubstituted ureas, The reaction rate is directlyproportional to [H +] and thus could be a measure ofhydrogen ion activity in a predominantly urea

192

environment. The energetics of the process and theeffects of the polarity of the medium have also beeninvestigated.

Materials and MethodsUrea, methylurea, ethylurea, N,N'-dimethylurea,

tetramethylurea and thiourea were either AR(BDH) orGR (E. Merck) grade and were used as such. Dioxane,acetone, acetamide and formam ide (AR, BDH or GR,E. Merck) were purified by distillation followingstandard procedures. Methyl and ethyl acetates wereof AR (BDH) grade.

In a typical kinetic run, the ester (2 ml) was added tothermally equilibrated hydrochloric acid solution(20 ml) and aliquots (2 ml each) were withdrawn atregular time intervals and quickly titrated withstandard NaOH solution after 10-fold dilution tominimise hydrolysis during titration. The rateconstants were evaluated from the slopes of the linearplots of log (V2 - VoN 2 - VJ versus time by themethod of least squares. Repeat runs showed amaximum standard deviation of ± I~;~.Measurementsin the presence of urea, substituted ureas and otheradditives were carried out similarly, the initial aq.acidic: solutions contained these additives in therequired concentrations. Under the conditions of thepresent study, urea, substituted ureas, acetamide andformamide did not undergo hydrolysis.

ResultsInfluence of the additives on reaction rates

At fixed [H +] and temperature urea, substitutedureas, acetamide and formam ide strongly retarded the

~~~~~~ -....-...~-~-----.- ...- ..--

DASGUPTA et al.: HYDROLYSIS OF METHYL &: ETHYL ACETATES IN UREA

Table I-Effects of Additives on Rate Constants of Acid Hydrolysis of Methyl Acetate[HC1]=0.3 mol dm 3; temp=310K

Additive k x \03 min 1 at [additive] mol dm :"-- -- ------_ .._-------.- .---~.-

0.2 0.5 1.0 2.0 3.0 4.0

Urea 6.44 ± 0.28 5.06±0.IS 3.63 ± O.OS 2.19±0.05 0.99±0.01Methylurea 4.05±0.IS 272±0.14 1.89 ±O 17Ethylurea 4.31 ±O.O7 2.S1 ±0.13 1.80 ± 0.07N.N'-Me2urea 4.74 ± (J.()8 3S0±0.OS 2.44 ± 0.06Me4 urea 4.38±(UO 3.68±0.9 2.95±007ThioureaFormamide 2.07 ±0.36 0.86±O22 O.16±004Acetamide 4.95 ±O.22 4.26 ±O.27 3.12±O.14 2.53 ±O.29

k X 10' = 6.35 ±0.396 min 1 in the absence of additive

30,--~---~~--~~~-~--~--

... - -------~-~- ...-.~~~--------.-.~~~-

U,OM

U;O.2M

U,0,5M

20

'cE~S." 18

U,4M

5 10 15

Fig. 1=-Dependence of rate constants on simultaneous variation ofurea and acid concentrations at 310K

reaction rates. The deactivating effects were in theorder formamide > acetamide> methylurea ~ ethyl-urea ~ urea < tetramethylurea ~ N ,N' -dimethylurea,Thiourea had no effect on the rate of hydrolysis(Table I).

In order to see if urea itself had any effect on thehydrolysis of the ester in the absence of acid, the rateconstants were plotted against the combinedconcentrations of urea and acid (Fig. I). The

Limk .extrapolated rate constant values, HCI .....O at vanous

[urea] intercepted the X-axis in all the cases at pointscorresponding to the respective [urea] showing thaturea had no effect on the rate of hydrolysis in theabsence of acid.

Of all the additives, urea was the most extensivelystudied and the rate constants decreased exponentiallywith increase in [urea]. The ratio of the experimentallyobtained rate constants in the presence of urea (k) andthat in its absence (ko) was taken as a measure of thedeactivating influence of urea at the respective [acid]

Effect of dielectric constantThe effect of dielectric cons tan ts of the mixed media

on the hydrolysis of ester in the absence of the amidewas studied in acetone-water and dioxane-water media(Table 2, Fig. 2). Values of the dielectric constants ofthese mixtures and others were collected fromliterature!'!! -!3. At equidielectric constant, the effectwas more pronounced in acetone-water medium thanthat in dioxane-water. However, in both the media, therate constants increased with increase in dielectricconstant!4 as expected. A reverse trend, i.e. decrease inthe rate constants with the increase in dielectricconstant was obtained in urea-water, formamide-water and acetamide-water media (a specific effect). Tounderstand the role of specific effect, the activationparameters were evaluated at equidielectric constantfor acetone-water and dioxane-water mixtures and thevalues. are given in Table 3.

20

DiscussionUrea and substituted ureas retarded the acid

catalysed hydrolysis of simple esters by keeping thecatalyst, H + ion, away from the sphere of action via itsassociation with the amides. Urea itself could notinfluence the hydrolysis in the absence of H + ion. Theeffects of substituted ureas on the reaction rates werecomparable to those of urea (a minor difference wasobserved between disubstituted and tetrasubstitutedureas) showing that the nitrogen atoms of the - NH2groups of the amides were not the centres ofattachment with H" ion (the catalyst). The ;;C=Owasexpected to be the effective site. The fact that thioureadid not affect the rate supported the noninvolvementof -NH2 centres and the involvement of :;C=Ocentre. The small difference between urea and the di-and tetra-substituted ureas might be due toperturbation of urea molecule by the substituents.

The ion-dipole interaction [between the H + (orHtO) and ester] to form the activated complex should

19J

8

INDIAN 1. CJ-IEM., VOL 23A. MARCH 1984

ws4 :r <t

Ie ct w•.... I>:'g •... :>

u~ ct

~x:<:

2

...0ict:tg•...

16

6

26 3121.1. X 103E

36

Fig. 2-Plots of rate constants versus reciprocal 01"dielectric constant in different aquo-organic solvent mixtures at 310 K

Rate ofTable 2-EfTect of Various Environments onHydrolysis of Methyl Acetate

[HCI] =0.3 mol dm -3; temp=310K

Mol (%) 1/£x 103 k X 103 min -.

Dioxane

o3.866.80

10.7516.56

13.4816.6720.025.033.3

6.35 ±0.45.96 ±0.25.73 ±0.25.36± 1.04.69±0.3

Acetoneo 13.48 6.35 ± 0.4

6.52 15.39 5.21 ± 0.316.18 18.18 4.48±0.330.66 22.22 4.01 ±0.0742.69 25.00 3.73± 0.4

Formamide

0 \3.48 6.35 ±O.40.91 12.45 2.07±0.361.84 12.30 0.86±0.223.76 12.08 0.16±0.04

Urea

0 13.48 6.35 ±0.40.36 \3.35 6.44±0.280.91 13.21 5.06±0.181.85 12.97 3.63 ±0.083.81 12.58 2.19 ±O.O58.09 11.85 O.99±O.OI

Acetamide

0 13.48 6.35±O.40.91 12.53 4.95 ±0.221.86 12.44 4.26±O.273.86 12.27 3.12±0.146.00 12.10 2.53 ±0.29

194

be favoured in high dielectric environment. Bothacetone and dioxane lowered the dielectric constantand they decreased the rates as expected. In presence ofamides, increased dielectric constant did not increasethe rate (Fig. 2). The influence of the amides on H + ionin retarding the rate was much more powerful ascompared to the influence of the dielectric constant ofthe medium on the rate. These additives retarded therates at all compositions (mol ~J employed and atequimolar concentration magnitude of decrease of rateconstants followed the order: dioxane < acetone<acetamide <urea < formamidc. At equidielectric.dioxane was less effective than acetone and the effectsof the amides were in the order: acetamide< formamide ~ urea. This might be the consequence ofthe greater basicity of the former than that of the latter[PK~sc (acetamide) = 13.37; pK~s' (urea) = 13.82; p K;of formamide is not known to the authors and hence acomparison is not possible]. The greater effectivenessof acetone than dioxane is due to the presence of

X = 0 group in acetone.

Bu\l et al. who studied the effect of urea on thedissociation of acetic acid conductometricallysurmised that the decrease in conductance in thepresence of urea is due to the influence on the activityof H + ion. Schafer? considered binding of H + ion withboth monomeric and dimeric ureas and using theassociation constants of the monomeric and dimericureas with H + ion, calculated the apparent(concentration dependent) equilibrium constant of the

Ksystem: U + H + ¢UH + , and observed that the trendwith. increased urea concentration fairly agreed with

DASGUPTA et al.: HYDROLYSIS OF METHYL & ETHYL ACETATES IN UREA

Methyl acetateWater 70.50 66.00 99.50 101.80Water-dioxane (= 50) 68.60 66.00 100.10 109.90Water-acetone (= 50) 68.10 65.50 100.70 113.40 4.0

Urea (1 mol dm -3) 68.00 65.40 101.60 116.60Urea (2 mol dm -3) 69.40 66.80 102.90 116.30

SHAFER (CALC)Urea (4 mol dm -3) 70.80 68.20 104.60 117.40x

Urea (6 mol dm -3) 66.40 63.80 106.10 136.40 3.0

Ethyl acetateWater 68.00 65.40 100.90 114.40

BULL AND WE (08SJUrea (4 mol dm -3) 70.50 68.00 105.90 122.50

Table3--Activation Parameters+ for Hydrolysisof MethylAcetate and Ethyl Acetate

E. sn: t1G;kJ mol :' kJ mol -I kJ mol -I

-t1S:j:J mol "K--1

[Experiments were carried out at 303, 308, 313 and 318K for theacid hydrolysis of methyl acetate (0.3 mol dm -3 HCI) in water,water-dioxane and water-acetone solvent mixtures. For differentconcentrations of urea, experiments were carried out at 305, 310, 315and 323 K. For the acid hydrolysis of ethyl acetate (0.5 mol dm -3

HCI), experiments were carried out at 303,310.317 and 323K.

experimental values even after neglecting the activityeffects. Following the procedure of Schafer, wecalculated the equilibrium constant (K) from our dataand those of Bull et al. as well. These are plottedagainst [urea] in Fig. 3 along with those of Schafer.Two distinct features were observed. The theoreticallycomputed curve on Schafer's model neither fitted ourdata nor truly that of Schafer's own data. Secondly,the K values of Bull et al. differed much from those ofSchafer's and agreed closely with those of ours. Webelieve that Schafer's model overestimated thecalculation due to errors in measurements. Moreover,association of both monomeric and dimeric ureas withH + ion could not reasonably account for thisdifference in results. To fit the data, Schafer used thesecond association constant (dimer - H + complex) as18.52 dm ' mol -1 and the first association constant(monomer - H + complex) as 1.5"dm 3 mol -1. This isunacceptable since dimer-H + association is notexpected to be stronger than monomer-H + asso-ciation. A dimeric urea, if at ail formed, is expected tohave lesser tendency to form the bonding, smce in thedimer both the -.:::C = 0 centres of two urea moleculeswill be involved in intermolecular hydrogen bondingwith the hydrogens of - NH2 groups and hence thesecentres will not be free to accept H + ions. It may bementioned also that the very existence of dimers andmultimers in aqueous solutions of urea is more of aconcept than a reality!".

The K-values obtained in the present study andthose from the data of Bull et al. showed significant

6.8

SHAFER (0.'.)

5.0

1.00~----:-'2.0-:------:"4'O::-----6-!;.O::----~8-!;.O::----:1~O'O

UREA, M

Fig. 3--Equilihrium constants as a function of urea concentration[<t, present work; CD. Bull et al:s results]

concentration dependence beyond 4 mol dm -3 ofurea. Below 4 mol dm -3, the values were more or lessconstant and close to those obtained from p Ki,Considering concentration terms instead of activitiesin the equilibrium constant relation:

K =[UH+]yuw/[U]-yu[H+]yw

it is clear that for 3 or 4-fold increase in K (see Fig. 3).the activity coefficients of denominator (particularly ofH + ion) must decrease sharply in urea solution or thatof UH + ion must increase sharply. Such a sharpdecrease of Yw or increase of }'uw is veryunreasonable 1 5. It is our contention that anexplanation for the increase in K-value at [urea] > 4mol dm -3 is to be sought elsewhere and not in themodel of Schafer. It is well known that urea at andabove 4 mol dm -3 acts as a potent protein denaturantand water structure breaker 16. It is therefore possibletha t urea reacts specially with the H + ion beyond 4 moldm -3; the less ordered water favours HjO-ureacontact yielding statistically more UH + species andhigher equilibrium constant. However, more work isnecessar-y to arrive at a satisfactory explanation forthis phenomenon.

The influence of urea on the activity of H + ion wasreported by Bull et al." Since the rate constants can betaken proportional to the H + ion activity, the kineticmeasurements also yielded the activity of H + ion inurea solution relative to that in water throughkurealkwa,er (or k/ko) ratio. To do this k/ko values wereextrapolated to zero acid concentration to minimise

195

1.0...---------------------,

INDIAN J. CHEM., VOL 23A, MARCH 1984

0.4

0.2

6

UREA, M

Fig.4 Plots of k tk.; versus urea concentration. Lim H ' -.0[ •. represent Bull et at:s results]

the influence of ion-ion interaction. In Fig. 4, these

ti~ ->0 k/ko values are plotted against [ureal An

exponential trend has been noted in the H + ion activityvalues in urea solution. Further, the values closelyresembled the reported values of Bull et al. (full circles).The close similarity between our data and those of Bullet al., though following altogether differentapproaches, qualified them to be more accurate thanthose of Schafer.

196

The energetics of the process showed unperturbedactivated complex in the presence of either the solventsacetone and dioxane or urea. This meant that neitherthe association of H + ion with urea nor the waterstructure alteration by the presence of sufficient ureahad any effect on the nature of the activated complex.The rate influencing property of urea was purelythrough competition with ester for the H + ions.

References

10

\'

I Handbook 0/ chemistry and physics, 55th Ed., (C.R.C. Press,Ohio) 1974-75.

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