Indian Journal of ChemistryVol. 21A. November 1982. pp. 1035-1039

Ion Association & Specific Solvent Effects on Rates ofAlkaline Hydrolysis of Monomethyl Succinate

BIKAS CHANDRA BAG & MIHIR NATH DAS·

Department of Chemistry. Jadavpur University. Calcutta 700032

Receired 22 January 1%2; reriscd 19 May 1982; accepted 19 August 1982

Salt effects on rates of alkaline hydrolysis of monomethyl succinate ion ha ve been studied in water, and the relative rates inthe presence of different cations follow the order Ba 2 + > Li + > Na + > K ". In the presence of Ba 2 + and Li + ions, theseparate rate constants for the reactions between the two free ions and between the respective cation-OH - pair and the freeester anion ha ve been evaluated using the corresponding values of the association constants. The catalytic action of the cationsarises from the fact that the activated chelate complex becomes more stabilised when a cation is more strongly bound to thenegatively charged ester and hydroxide ions. The effect of changing dielectric constant (D) on the rate has also been studied atdifferent temperatures in ethanol-water and ethylene glycol-water mixtures. Fairly good linear plots are obtained by plottinglog z , (ko being the rate constant at zero ionic strength) against liD, upio about a dielectric constant of - 55. The radius of theactivated complex as obtained from the negative slope of the plot for each solvent system is found to increase with increasingtemperature. Specific solvent effect has been studied at 30 at a constant dielectric constant (D = 50.38) using different aquo-organic mixtures. and the rates lie in the order acetone > dioxane > ethanol > ethylene glycol.

The solvent is known to exert pronounced effects onthe rates of reactions, particularly of the ion-ion type 1 .

The rates of such reactions are also considerablyaffected by ionic strength. In some cases, the rates arefound to follow Bronsteds theory of primary salteffect. but in very many cases, the rates, which havebeen explained on the basis of Bjerrum ion-pairformation. are found to depend neither on thestoichiometric ionic strength nor on the concentrationof ions of the charge opposite to the reactants. To getfurther information on the effects of solvent and ionicstrength on ion-ion reactions. the alkaline hydrolysisof monomethyl succinate has been chosen for thepresent investigations.

Materials and MethodsSuccinic acid half-ester (methyl), m.p. 58' was

prepared and purified according to the standardmethod. The purity of the ester. as checked by titrationwith alkali. was found to be 98,S'j;,. It was alwaysstored in a vacuum desiccator. Ethylene glycol (BDH).dioxane (May and Baker). acetone (E. Merck) andethanol were purified by standard methods. NaOH.LiOH. KOH. Ba(OHh. HC!. and all the salts usedwere of AR grade (BDH). Doubly distilled CO2-freewater was always used for preparing the solutions.Solutions of the ester and the alkali were alwaysprepared afresh.

The rates were measured at 20-, 30- and 40"±0.05~C. Initial [alkali] and [ester] used were0.01 moldm -3 each in the reaction mixture. For thekinetic studies in different solvent mixtures, the ionicstrengths (I) were varied from 0.02 to 0.11 mol dm -3

by adding calculated amounts of standard NaCisolutions. For the study of specific salt effects inaqueous medium. salt concentrations were varied from0.02 to 0.1 N using the corresponding bases. Thecourse of the reaction was followed titrimetrically".The rate constants (k) obtained graphically from thesecond order plots, were reproducible within ± 2%.The rate constants (ko) at / = 0 were obtained byextrapolating the linear plots of logk against /1/2 tozero ionic strength. The rate constants on the mol-fraction scale (k~) were calculated from the molar rateconstants (ko) from the relation k~ = kon where n is thetotal number of mol of the solvents per litre of thesolution. as calculated from the densities of the solventmedia.

To study the effect of the dielectric constant of themedium. the rates were measured at 20", 30e and 40°Cusing water-ethanol and water-ethylene glycolmixtures containing upto 70'X, and 60% by weight ofethanol and ethylene glycol respectively. The values ofdielectric constants were obtained from theliterature+" and interpolated values were usedwhenever necessary.

Rates were also measured at 30'C in isodielectricmedia (0 = 50.38) comprising acetone-water. dioxane-water, ethanol-water and ethylene glycol-watermixtures to study the specific solvent effects.

Results and DiscussionEffect of ionic strength-The rate constants on

molar and mol fraction scales measured at 30°C inwater are presented in Table I. It is seen from Table Ithat the rate constants are almost independent of the

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INDIAN J. CHEM., VOL. 21A, NOVEMBER 1982

Table I-Effect of Added Salts on the Reaction Rate at 300{[Ester ion] =0.01 M; [OH-]=O.OIM}

Added Salt Stoichiometric True ionic k logkCone. (N) ionic strength strength (I) (litre mol -lmin -1)

With NaOH

None 0.02 0.0188 5.77 0.7612NaCI 0.02 0.04 0.0388 6.00 0.7782

0.06 0.08 0.0788 6.30 0.79930.10 0.12 0.1188 6.58 0.81820.14 0.16 0.1588 6.70 0.8261

Na2S04 0.02 0.05 0.0420 6.03 0.78030.06 0.11 0.0774 6.32 • 0.80070.10 0.17 0.1050 6.60 0.8195

Na-Citrate 0.02 0.06 6.10 0.78530.06 0.14 6.34 0.80210.10 0.22 6.65 0.8228

With KOH

None 0.02 0.02 5.75 0.'1597KCl 0.02 0.04 0.04 5.95 0.7745

0.06 0.08 0.08 6.20 0.79590.10 0.12 0.12 6.40 0.8062

K2SO. 0.02 0.05 0.0415 5.99 0.77740.06 0.11 0.0743 6.18 0.79100.10 0.17 0.1025 6.37 0.8041

K-citrate 0.02 0.06 5.97 0.77600.06 0.14 6.22 0.79380.10 0.22 6.35 0.8028

With LiOH

None 0.02 0.0186 5.85 0.7627LiCI 0.02 0.04 0.0384 6.10 0.7852

0.06 0.08 0.0780 6.52 0.82090.10 0.12 0.1167 6.89 0.8451

Li2S04 0.02 0.05 0.0430 6.24 0.79520.06 0.11 0.0835 6.55 0.81620.10 0.17 0.1133 7.05 0.8482

With Ba(OH),

None 0.03 0.0278 6.15 0.7889Bact2 0.02 0.06 0.055 6.82 0.8338

0.06 0.12 0.110 7.64 0.8831- ----- --_._- - ----

nature of the anions, when a particular cation at agiven concentration is considered. But at a particularionic strength, when anions are the same, the rateconstants are distinctly influenced by the nature of thecation, in accordance with the general theory of thespecific ionic interaction 5. The relative rates in thepresence of the different.cations follow the order Ba Z +

> Li + > Na + > K ".Davies's equation (I) was used to correlate the race

constants with F(l)

-logfi=AZ;F(l) ... (I)

where F(l)=/1iZ/(l +/lIZ)_0.2/. Due to ion-association, the stoichiometric ionic strength does notreflect the true ionic strength of the solution. The trueionic strengths, therefore, were calculated from theconcentrations of the ion-pair computed from the

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F(l)

0.1.1680.15680.20340.23250.25310.16160.20210.2237

0.120o 1580.2040.2330.1610.1990.222

0.11620.1560.20270.2310.1630.20750.229

0.13730.17890.2270

respective association constants" at zero ionic strengthby the method of successive approximations 7, usingDavies's equation (I). When the values of theassociation constants for a particular ion-pair at 30were not available, vant Hoff's equation was usedwhere data for atleast two other temperatures wereknown. Ion-association has been considered forNaOH, Na2S04, K1S04• LiOH, LizS04 andBa(OH)l' For other electrolytes, either there was noappreciable formation of the ion-pair or the values ofthe association constants were not available so that asa first approximation the electrolytes were consideredto be fully dissociated. The plot of log k against F(l),where 1is the true ionic strength, shows good linearity.The slopes of the linear plots obtained with Ba 2 +- , Li + ,

Na + and K + are 1.05,0.80,0.50 and 0.42 respectively,the intercepts. being 0.65, 0.68, 0.70 and 0.71

BAG & DAS ALKALINE HYDROLYSIS OF MONOMETHYL SUCCINATE

respectively. These observations are however hardlyamenable to a simple and straightforwardinterpretation. Two alternative approaches have beenmade to correlate the results as discussed below.

Besides the reaction between the negatively chargedester ion and OH -, the reaction may also occursimultaneously between the ester-cation ion-pair andthe free OH -, as well as between the free ester ion andthe cation-OH - pair. The overall reaction rate (RT)

should therefore be given by Eq. (2).

RT=k[OH-1L[E Js=kl[OH-J[E J

+k2[MOH(n I)+J[E -J+ k3LOH· J[ME(n I)+J

... (2)

where [OH -Js and [E -Js are the stoichiometricconcentrations of the reactants, and [MOH(n -I)+Jand [ME(n -I) +J the concentration of the ion-pairsformed between the cation Mn+ and OH - and esterion (E) respectively. Due to the non-availability of thevalue of the association constant for the cation-esterion-pair, no exact treatment of the Eq. (2) could bemade. Under the condition, the observed rate constant(k), as obtained from the Eq. (2), is given by the Eq. (3). -

k={kl[OH -J+k2[MOH(n I)+J}/[OH -1 ... (3)

The value of k I has been taken to be the interpolatedvalue of k obtained with KCI at the correspondingionic strength, assuming that no ion-pairing occursbetween the OH and K +. In the presence of lithiumand barium chlorides, the concentrations of the freeOH - and cation-OH pairs have been calculated fromthe association constants of the correspondinghydroxides, using the method of successiveapproximation 7. The results in the presence of lithiumand barium chlorides are presented in Table 2. It isobserved that in the presence of LiCI, k 2 remainsremarkably constant (8.0 1-8.06) inspite of a substantialchange in the ionic strength, as expected for a reactioninvolving a neutral species, the ion-pair Li +OH . The

k2 values in the presence of BaCI}., however, show adistinct, though small, trend of decreasing withincreasing ionic strength, as expected for a reactionbetween two ions of opposite charges, the anion E andthe cation BaOH ".Similar results were obtained for thealkaline hydrolysis of the acetyl mandelate ion 2.

An alternative approach may be made in respect ofthe catalytic action of the cations for the presentreaction, by postulating the participation of a cation inthe transition state (A) through chelation 7. On the

basis of this postulate":", the following Eq. (4) can beobtained

k = k~l[Mn +Jfl - fl - f, + /f(2 -n) - ... (4)

where k~ is the catalytic rate constant for the cationM" + and k is the observed second order rate constant.Eq. (4) can be written in the following forms for Na+and Ba2 ". using Davies's equation (I) for activity co-efficients.

(1112 \

logk +6A --;-- - 0.2/ )= logk~A2+ + log [Ha2+J1+112 ... (5)

(/1/2 \

logk+2A ---. -0.21 )=logk~a++log[Na+J.1+/12 (6)

Equations (5) and (6) predict that if the activatedcomplex entirely remains in association with thecation, the plot of the quantities on the left hand sideagainst cation concentration should be linear with aunit slope. In fact, for the cation Na +, the slope isabout 0.18 which is far from unity, but for Ba2 +, theslope has been found to be 0.56.

Table 2 Rate Constant for Reactions of Monomethyl Succinate with Free Hydroxyl Ion (k I) and Cation-OH - Pair (k 2) atDifferent Ionic Strengths at 30

[LiCI] k(obs) [Li +OH ] F(/) k, k2N (litre mol'min ') (litre mol 'rnin ')

0.00 5.~5 0.0004992 0.1162 5.73 8.010.02 6.10 0.0007580 0.1560 5.943 8.020.06 6.52 0.001670 0.2027 6.209 8.050.10 6.89 0.002542 0.2310 6.500 8.06

[BaCI2] k(obs.) [BaOH +] F(/) k, k2N

0.00 6.15 0.000829 0.1373 5.82 9.80.02 6.82 0.001736 0.1789 6.15 9.430.06 7.64 0.003443 0.2270 6.47 9.17

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

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INDIAN J. CHEM .. VOL. 21A. NOVEMBER 1982

From this. it can be argued that while in the presenceof sodium salt, most of the activated complex remainsfree from direct cation participation, in the presence ofbarium salt, the activated complex seems to besubstantially in association with the cation, Ba2 ". Themore strongly a cation is bound with the ester ion andOH -, the more stabilised will be the chelate complex inthe transition state and hence faster the reaction rate inthe presence of that cation.

Both the concepts based on ion association andchelation respectively are, however, basically similar,the catalytic action of the cations essentially arisingfrom a reduction in the electrostatic repulsion betweenthe two reacting ions of like charge, and hencefacilitating their close approach, and increasing theprobability of the formation of the transitionstate 7.10.11.

Effect of dielectric constant of the medium-According to Bronsted-Christiansen-Scatchardequation, the dependence of the rate constant (ko) atinfinite dilution for a second order reaction on thedielectric constant of the medium is given by Eq. (7)

ZAZBe2Ink =lnkL----

o 0 DKTr ... (7)

where ZA and ZB are the charges of the two ions A andB, e the electronic charge, K the Boltzmann constant, Tthe temperature C'K) and r the radius of the activatedcomplex, kf; being the value of ko in a hypotheticalmedium of infinite dielectric constant.

The rate constants (ko) for ethanol-water andethylene glycol-water at 30' are given in Table 3 onmolar and mol-fraction scales. The plots of log k 0

versus I/O at 20",30',40° are linear for both solventmixtures (Figs I and 2), as expected from Eq. (7), up todielectric constant values of 55, beyond which thelinear plots either band upwards or remain almost flat.The plots obtained for ethanol-water mixtures on themol-fraction scale are less curved. From the slopes of

the linear portions of the plots the values of the radius(r) of the activated complex in ethanol-water andethylene glycol-water mixtures have been calculated at20", 30' and 40°. In ethanol-water mixture the valuesof r(A) obtained from ko values were 3.99,4.6 and 7.0and those obtained from k~ values were 2.34, 2.99 and3.76 respectively. Similarly in ethylene glycol-watermixture, the r (A) values obtained from ko values were1.17,1.51 and 1.85 and those obtained from k~ were1.03, 1.19 and 1.70 respectively. It is striking that the rvalues for ethanol-water and ethylene glycol-watermedia increase with increase of temperature, the rateof increase of r values with temperature in the mol-fraction scale being rather small. Such variations arerarely observed for the size of the activated complexwith temperature and the causes seem to be ratherobscure.

'·00·9

0·8

0·7

0·6

o- 5" '="2-c,"".3-;',•.• :---:':'.5'-----:-'.6;:--7'1-7:;-7'.8:;--;"'. 9

t)t10~

'·2,.,"0"'-:;.2-';':;" 3---;'-;1-4--+'.5c--f'.6....:'.7 --:1-6'0-.,'0.9-;2.0

ill 102

Fig. Iia) -Rate constant as a function of dielectric constant onmolar scale at (I) 20', (2) 30 and (3) 40 in ethylene glycol-water

Fig. I(b) -Rate constant as a function of dielectric constant on molefraction scale at (I) 20, (2) 30' and (3) 40 in ethylene glycol-water

D

310.84252.50196.01152.17106.2968.9049.1136.22

76.7573.8471.0467.9864.8961.5257.8153.22

Table 3--Rate Constants (ko and ki/) for Alkaline Hydrolysis of Monomethyl Succinate in Ethanol-Water and Glycol-WaterMixtures at 3~'

Ethanolethylene

glycol(,/", w!w)

Ethanol-water

D ko(lit mol -, min -1)

k~(mol fr -, min -')

o10203040506070

76.7571.0465.3159.5053.6847.7742.1936.98

5.625.014.313.752.982.241.901.73

----_. __ ._-------------------

\038

AD(lit mol' min -')

5.623.603.002.401.901.661.54

Ethylene glycol-water

k~(mol fr -, min -')

310.84186.91145.59108.3379.3463.6953.87

BAG & DAS: ALKALINE HYDROLYSIS OF MONOMETHYL SUCCINATE

2·'

2·06

+0 t-s.>t;

'"~

'·2,.,'·0e-s0·8L---L..--1.-.L..--L-L.....L-1-:-:l-:--,L--::'-:--:l-:c----:-'L--:'7--:-'-::--c:'::-~:_::'::__='

'·2 '·3 '·4 '·5 '·6 '·7 1·8 1·9 2·0 2·1 2·2 2-3 2·4.2·5 2·6 2·7 2·8 2·9 3·0

.J....lO'o

Fig. 2··-Rate constant as a function or dielectric constant at 20 .30 and 40 on molar scale: [curves: (I). (2), (3)]. and mole fraction scale:[curves: (4), (5). (6)], in ethanol-water.

Table 4-Rate Constants (ko and k~) for Alkalyne Hydrolysis of Monomethyl Succinate in Isodielectric vledia (D = 50.38) ofDifferent Aquo-organic Mixtures at 30

Solvent Solvent Mol fro ko k~",,(WIW) of water (litre mol' 'rnin ') (mol fr-'min -I)

Ethyleneglycol 75 0.534 0.95 28.41Ethanol 45.6 0.758 2.66 87.58Dioxane 30.25 0.918 3.78 160.15Acetone 44.25 0.802 4.26 163.86

Specific solvent eflects--The values of the rateconstants (ko and ki{) for four isodielectric media onmolar and mol-fraction scales are given in Table 4. Thevalues lie in the following order: acetone> dioxane>ethanol> ethylene glycol. The relative order of ratesin the isodielectric aqua-organic mixtures may beinterpreted as largely arising from specific solvation ofthe reacting ions. as discussed in details in an earlierpaper+'.

AcknowledgementWe thank the UGC. New Delhi. for the award of a

Teacher-fellowship to one of the authors (BCB) and tohis employer. B.N. Mahavidyalaya , ltachuna, WestBengal. for granting him leave.

References

I Arnis E S. Solient effects on reaction rates and mechanisms(Academic Press, New York) 1966.

2 Roy A K & Das M N. 1 chem Soc (A), (1970) 464.3 Akerlof G, 1 Alii chem ss« 54 ( 1932) 4125.4 Akerlof G & Short A O. 1 Am chem Soc. 58 (1936) 1241.5 Guggenheim E A, Thermodynamics (North-Holland,

Amsterdam) 1950.316.6 Stabilit v constants of metal-ion complexes. Special Publication

No. 17 (The Chemical Society. London) 1964.7 Hoppe J I & Prue J E. 1 chem Soc. (1957) 1775.8 Kershaw M R & Prue J E, Trans Faraday SOl', 63 (1967) 1198.9 Basu M K & Das M N. 1 chem SodA), (1968) 2182.

10 Indelli A. Molan G &Amis ES,1 Am chem Sac, 82(1960) 3237.

II Smith Lennart & Olin Bror, 1 phys Chern (A), (1936) 131. 177.12 Roy A K & Das M N, J chem Soc (A). (1970) 1576.

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