Das, A. R.

Moulik, S. P.

Ray, S.

English

NOPR

Indian Journal of ChemistryVol. 14A, December 1976, pp. 921-924Interaction of p-Nitrophenol & Ethylenediamine inAcetone-Water Medium: Effect of Solvent Composition onProton Donor-Acceptor EquilibriumS. P. MOULIK & S. RAYDepartment of Chemistry, ]adavpur University, Calcutta 700032andA. R. DASDepartment of Physical Chemistry, Indian Association for the Cultivation of Science, Calcutta 700032Received 15 October 1975; accepted 16 January 1976The proton transfer complex formation involving p-nitrophenol and ethylenediamine inacetone-water mixtures has been investigated spectrophotometrically. From the evaluatedthermodynamic parameters of the process, it is observed that while AHo and ASo vary, AGoremains fairly constant in all the solvent compositions. Specific solvation of the complex aswell as solvent structural effects are considered to be the vital factors as compared with theminimal contribution of dielectric constant of the medium.C0:'fPLEX formation between a phenol andan amine in aqua-organic solvents involvingproton-transfer may serve well as a modelfor the hydrogen-bonded interactions commonlyoccurring in biological systems having aqueous-nonaqueous snrroundings-. Though the investiga-tions in this area are growing in number, a few veryrecent works have become critical in approach-:".In these, the general features of mixed solvents incontrolling the complex formation for systems in-volving p-nitrophenol and ethylenediamine as wellas p-nitrosalicylic acid and ethylenediamine havebeen presented in detail. Specific solvation andsolvent structural effects have been considered toplaya dominant role in the hydrogen-bonded ionpairs or proton transfer complexes+", The presentpaper mainly deals with this particular aspect ofp-nitrophenol (PNP) and ethylenediamine (en) in-teraction in acetone-water mixtures of varyingcomposition. The choice of PN'P was favouredby its appreciable interaction with the diamine inwide proportions of the mixed solvent. The dipolaraprotic nature of acetone is expected to help inunderstanding the specific role of the hydro xyliccosolvent, water. The thermodynamic parametersof the PNP-en interaction have been evaluated.Materials and MethodsThe details of the materials used and the experi-mental methods followed were similar to those re-ported in our earlier papcrs+".Results and DiscussionSolvent effects - The solvent dependent complexformation is shown in Fig. 1, where a comparisonof the effects of acetone, l,4-dioxane and n-propanolreveals aJ1.initial increase in absorbance for a smalladdition of each of the solvents. The observedabsorbance decrease past the maximum is rathersharp for acetone and l,4-dioxane. However, theeffect of the hydro xylic solvent, n-propanol is muchless pronounced. If these solvents are comparedat iso-rnole fractions in the light of dielectrics of themedia, the log A against D-l plots are observed tobe excellently linear giving three separate linescorresponding to the three nonaqueous solvents(Fig. 2). This is expected in view of our recentfindings? that the hydroxylic solvents when comparedin a similar manner fall on one line from which thenonhydroxylic solvents sharply deviate. The specificsolvent effects, similar to those observed in the caseof p-nitrosalicylic acid-en system, are thus envisaged.It is quite likely that a significant part of thisspecific effect depends on how the nonaqueous mole-cules interact with aqueous ones. In the presentstudy acetone has been chosen as a probe solvent,which, besides being nonhydroxylic, can impartappreciable dielectrics to the mixed media.o 0.2 0.4 0.5c..XsFig. 1 - Plots of absorbance of the complex at 405 nmagainst the mole fraction (Xs) of added cosolvents(1, n-Propanol, 2, acetone, 3, dioxane). For curves 1 and2: [PNP] = 4'92 X lO-sM and [en] = 3·60xl0"'M; for curve3: [PNP] = 4·92 x 10-sM and [en] = 7·7 X 1O-3M921INDIAN J. CHEM., VOL. 14A, DECEMBER 1976E,<]uiLibrium constant of the complex - The forma-tion constant (J{) of the complex in various acetone-water mixtures has been evaluated following theprocedure described earlier+t, The absobance dataat 405 nm obtained with increasing addition of theamine to PNP have been processed to give the J(values. The absorbance of PNP alone in themedium has been subtracted from the values in thepresence of amine to obviate the effect of solvent-induced proton transfers-s on the equilibrium betweenthe donor phenol and the acceptor amine. Thesignificant increase in absorbance by the additionof amine indicates only minimal proton donation tothe solven t in presence of the base2,6.0·90.150·7..C7'.9I 0.1 0·50.050.1ono1Fig. 2 - Plots of log A versus D:» (I, n-Propanol; 2, acetone;3, dioxane). For curves 1 and 2: [PNP] = 4·92 X 10-5Mand [en] = 3·60 X 10-'M; for curve 3: [PNP] = 4'92 X 10-5Mand [en] = 7·7 x10-3MThe complex formation process may be representedby Eq. (1)KAH+B ~ A- H+B (1)(a-x) (b-x)where a, b and x are the concentrations of thedonor, acceptor and the complex, respectively (thedonor being non-absorbing at Amax of the complexhaving extinction coefficient, €c)'When b,?>a. it can be shown! in a manner ana-logous to the treatment of the modified Benesi-Hildebrand equations thata/A = 1/€c+1/€cKb ... (2)where A is the absorbance of the complex.The plots of a/A versus lIb, from which the valuesof K and €c can be evaluated, are linear at differenttemperatures as demanded by Eq. (2). The valuesof K and €c thus evaluated are given along withtheir appropriate standard deviations in Table 1.Effect of temperature - The absorbances have beenmeasured at four different temperatures and the J(values have been computed for va'fious compositionsof the solvent. The AHo values have been evaluatedfrom th : least square plots of log K versus T>, andthe AGo and ASo calculated using the well knownthermodynamic relationship of Gibbs. The standardstates of the reactants have been considered to bethe hypothetical states of ideal Solution of unit mola-rity. The thermodynamic parameters at 25° forvarious acetone-water mixtures are given in Table 2.The variation of J{ with temperature may as well beconsidered to be partly due to the electrostatic effectinherent in the medium, keeping in view thetemperature dependence of dielectric constant ofa medium. The dielectric constants of varioussolvent mixtures at different temperatures are givenin Table 3. The marked dependence of log Kversus D-1 relationship on the solvent composition(Fig. 3) is thus corroborated by our earlier findings".Comparison of thermodynamic parameters - Thenature of the plots of !1Ho and (1So versus the molefraction of acetone (Fig. 4) are similar to theabsorbance versus mole fraction plots in Fig. 1.TABLE 1 - MOLAR EXTINCTION AND EQUILIBRIUM CONSTANT OF THE COMPLEX lN ACETOr(E-VVATER MIXTURESAcetone <c X 10-4 litre mole+ ern'? (K X 10-3 litre mole=) atwt %(mole fraction) 25° 30° 35° 45°0·0 2·95±0·24 2·67±0·15 2-22±0-07 1-80±0·09(0-0) (7·70±0·03) (S'57±0'02) (11'30:l: 0-01) (19'50±0'03)16'37 3'46±0'14 3'04±0-lS 2'5S±0'09 2·11 ±0·06(0'05S) (5·03±0·01) (6·17±0·02) (S'531: 0'01) (15'00±0'01)34·32 3'53±0-62 3'15±0'2S 3·01±0'26 2'53±0-14(0,140) (4'S5±0'37) (5·97±0·19) (6'72 ±0-17) (10'45± 0'09)43·92 3·1O±0·09 2·69±0·05 2'4S±0-05 2·21±0·03(0,195) (H7±0'03) (5-60±0'02) (6·30..l0·03) (9'20±0'07)4S'92 2·56±0·05 2·34±0·05 2·0S1:0-06 2·15±0·07(0'229) (5'63 ±0'04) (6·35±0·10) (7·07:1::0'17) (7'65 ±0·2S)54·04 2·31 +0·05 2'24±0'02 2'16±0-04 2'lS±0'05(0'26S) (S'21±0'25) (7·41±0'10) (7'11 ±0·16) (S·Sl±0·17)59·23 2'32±0·04 2'23±0'09 2-16±0-06 2·04±0·07(0'310) (4-65±0·04) (4'23 ±O'OS) (4'11±0·03) (3'99±0-03)64'62 2·36±0·03 2·39...L0·05 2-49±O·10 2·51 ±0·1O(0'362) (3'37 ±0·06) (2'72±0'02) (2·14±0·01) (1'99±0'01)922MOULlK et al.: INTERACTION OF p-NITROPHENOL & ETHYLENEDIAMINEIt is observed that the plots of energy parametersagainst either the mole fraction (Xs) or inverse ofdielectric constant (D-l) hardly show any differencein the trend. The !:iCo values, however. remainfairly constant up to acetone composition as highas 65 wt per cent. The above observations clearlyindicate the marked effect of acetone on !:iHo andll.So values for the overall process in which thespontaneity of the complex formation remains practi-cally unaffected. The mild increase in !:is° andTABLE 2 - THERMODYNAMIC PARAMETERS OF THECOMPLEX FORMATION AT 25°±0·1 ° IN DIFFERENTACETONE- \VATER MIXTURESAcetonewt %(molefraction)tJ.GO(kcal mole=)tJ.HO ss:(kcal mole'"] (cal deg" mole<)0·0 (0,0)16·37 (0'058)34·32 (0'140)43·92 (0,195)48·92 (0,229)54-04 (0'268)59·23 (0'310){i4-62 (0,362)+9'27±0'45+10'19±0'21+7-44±0'21+6·41 ±0'25+3·44±0·46-1·60±0·30-1·72±0·40-9'53±0'15+48'87±1-61+51-11 ±0'71+41·81 ±0'84+38'20±0'84+28·71±1·56+12·54±1·08+11-O0±1'36-15'83±0'52-5'30±0'03-5·05±0·01-5·03±0·05-4·98±0·01-5·12±0·01-5'34±0'02-5'00±0'01-4-81 ±0'01TABLE 3 - TEMPERATURE INDUCEDDIELECTRIC CONSTANTS OF THE SOLVENT MIXTURESAcetone Dielectric constant at{mole fraction)25° 30° 35° 45°0·000 78·5 76·7 74·9 71'50·058 69·0 67·5 66·0 63·00·140 58·0 56·5 55·0 52·50·195 51-5 50·5 49·5 47·00·229 48·5 47·5 46·5 44·00·268 45·5 44·5 43·5 41·00·310 42·0 41·0 40·0 38·00·362 39·0 38·0 37·0 35·04.51IIII.x: "I(J>"j9!:iHo for an initial small addition of acetone seemsto be related to the breaking of water structure.An initial small addition of acetone breaks downthe normal structure of water which is manifestedby the increase of !:iSo and !:iHo for the overallprocess. Further addition of the nonaqueousI60r---------------------------------~II1\9407~fs20\"., \3... \.~% !I!\i. -eCD 10CI \\0\0.05 0.10 0.15 0.20I0.'5-IXs __0.2 0.6 1.0 1.4 1.8 n-10 YOXlot_-3-20-7- - •.- '- --"--"2 .--.-~ ~ __ ._-c-.:J="----- 1-9Fig. 4 - Acetone induced thermodynamic parameters forthe complex formation process [1 and 2: tJ.Go versus Xs;3: sn- versus X s; 4: tJ.so versus X s; 5: tJ.Ho versus D-l]36B3.01~.0::-------:,-;5-------!2.~O-------'2.-5------,J".o1'0',0'Fig. 3 - Temperature induced electrostatic effects on the equilibrium constant pl'?tted as log K versus D-l [Curves1-8 represent 0, 0·058, 0'140, 0·195, 0'229, 0·268, 0·310 and 0·362 mole f~actlOn of acetone, respectively]923INDIAN J. CHEM., VOL. 14A, DECE:\1BER 1976through the negative oxygen centre. Such animpaired solvation is absent when the contributingmolecules are all water. The crowding of acetonenear the complex as well as PNP and the aminemay, however, contribute a share to the observed-entropy decrease with its increasing addition into-the system.It has been shown in Fig. 3 that for a constantsolvent composition, log K is a linear function ofthe reciprocal of temperature dependent dielectricconstant of the medium and is displayed by thereversal of slopes. Taking into account that variousphenomena may appear in the proton transferprocess+", only a qualitative explanation can beoffered. At a constant dielectric constant the forma-tion of complex depends on pK of both the dcnorand acceptor-'.". Increasirg the temperature affectsthe PK value, normally the dissociation constantsare increased. This will favour the transfer of protonwhich at the same time will be disfavoured by thetemperature effected decomposition of the complex.The decreasing dielectric constant of the mediumwith increasing temperature should also disfavourcomplex formation. The observed magnitude is,therefore, the resultant of all these effects. So longthe mole fraction of acetone is >0·23, the com-plex formation decreases with decreasing dielectricconstant which is a physical manifestation of thechanging temperature. Below this mole fraction,the complex formation is favoured by the increasein temperature. Therefore, in this range the pK-controlled complex formation overrides the decom-position as well as dielectric induced restriction ofthe transfer process, which, in the other range, isjust the reverse. Both solvation and solvent struc-tural effects are considered to have little scope toobliterate this explanation since in Fig. 3 comparisonof the results has been made at constant solventcomposi tions,1612lOx I<I e ~I450 -20 -10 0 30 40 ~o/os·Fig. 5 - Linear enthalpy-entropy relationship at 25° forthe complex formation process. Signs 0, • and 0represent the results on p-nitrosalicylic acid-en complexin water, 12·15 wt% ethanol, and 8·2 wt% n-propanol,respectively?solvent subsequently leads to structure formationcausing decrease in the overall entropy change.The systematic variation of I).Ho and 1).5° at nearlyconstant I).Go manifests energy utilization throughsolvent structure modification and consequentlya case of enthalpy-entropy adjustment. Such aprocess must then reflect a linear relationshipbetween I).Ho and 1).5° in an extra thermodynamicway9.10. This is strongly supported by the resultsof Fig. 5 where the linear enthalpy-entropy variationis almost exact. In this figure three pairs of datacollected on p-nitrosalicylic acid-ethylenediaminesystem in water, 12 wt% methanol in water, and8·2 wt% n-propanol in water, respectively havebeen shown to fit in I).Ho versus 1).5° plot. Theextra thermodynamic relationship has thus thepossibility of being a feature of the proton transferprocesses reported herein. Such a relation is knownto hold good for a given reaction involving sub-stituted derivativest-, and we have seen that thisrelation also holds good under the changed environ-mental conditions. Fig. 5 shows that at 0·273 molefraction of acetone the enthalpy change for thereaction is zero. It has been observed that thehydrogen-bonded dimerization of acetic acid in aq.medium ends up with a zero enthalpy change12.Since the dielectric constants of acetone-watermixtures used are not low, stabilization of thecomplex by solvation--e" and the structure of theoverall environments are expected to playa vitalrole in the complex formation reaction. The twoprocesses compete with each other, and destabiliza-tion of the complex by de solvation may eventuallybe offset by the stabilizing influence of the neigh-bouring surroundings. The two adjacent methylgroups in acetone may sterically hinder the approachof the whole molecule to the negative centre of thecomplext to solvate it through the positive centreof acetone on its electron deficient carbon atom,[(CH3)2c+ z: 0-]. But it can e~sily solvate at thepositive centre of the complex Via hydrogen bonding924AcknowledgementThe authors thank Prof. M. N. Das, Hpad of theDepartment of Chemistry, Jadavpur University,for providing facilities.References1. Vn.;QGRADOV,S. K., HUDSON, R A. & SCOTT, R M.,.Biochim, biophys, A cia, 214 (1970), 6.2. SCOTT, R M. & VINOGRADOV,S. N., J. pbys, Chem,Ithaca, 73 (1969), 1890.3. SCOTT, R 1\1., J. phys, Chem, Ithaca, 76 (1972).587. •4. MOULIK, S. P., SAHU, C. R & MITRA, A. K, Carbobyd.Res., 25 (1972), 197.5. MOULIK,S. P., CHATTERJEE,A. K. & SEN GUPTA, K. K.,.Spectrochim. Acta, 29 (1973), 365.6. MOULIK, S. P., CHATTERJEE, A. K. & SEN GUPTAK K, Indian]. Chem., 12 (1974), 92. ~7. MOULIK, S. P., RAY, S. & DAS, A. R, J. pbys, Chem.(in press, 1976).8. BENES!, H. A. & HILDEBRAND, J. M., J. Am. chem, Soc.71 (1949), 2703. '9. PIMENTEL, G. C. & MCCLELLAN,Amlu. Rev. Phys. Chem.,22 (1971), 347.10. LUMRY. R, Biopolymer. 9 (1970), 1125.11. EON, c., POMMIER,C. & GUIOCHON,G., J. phys. cu«.•.75 (1971), 2632.12. DAVIS, M., Hydrogen bonding, edited by D. Hadzi &H. "V. Thompson (Pergamon Press, London), 1959,.393.

Interaction of p-Nitrophenol & Ethylenediamine in Acetone-Water Medium: Effect of Solvent Composition on Proton Donor-Acceptor Equilibrium

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Interaction of p-Nitrophenol & Ethylenediamine in Acetone-Water Medium: Effect of Solvent Composition on Proton Donor-Acceptor Equilibrium

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