Indian Journal of ChemistryVol. 20A, April 1981, pp. 347-352

Ion-Solvent Interactions II} Aqueous 2-MethoxyethanolABHDIT BHATTACHARYA, ASIM K. DAS & KIRON K. KU1'.'DU·

Department of Chemistry, Jadavpur University, Calcutta 700 032

Received 10 March 1980; revised and accepted 8 September 1980

Standard free energies of transfer (6.G~ ) of MCI [M = Li, Na, K, Rb and Cs] and KX [X=Br and I] bavebeen determined in aqueous mixtures of 10,30,50 and 70 wt % of 2-methoxyethanol (ME) at 25°C. These valuesbave also been apportioned into individual ionic contributions using the reference electrolyte (tetraphenylarsoniumtetraphenylboride) method. It is observed that while the halide ions are increasingly destabilized due to decreasedacidity of the mixed solvents, alkali metal and H+ ions are increasingly stabilized due to combined effects of increasedbasicity, the possible formation of electrostatically bound bidentate complex of the type M+(ME)2' soft-soft interactionsand/or structural changes of the solvents. The observed pronounced stabilization oftetraphenylarsonium tetraphenyl-boride ions is attributable to the combined effects of increased dispersion and cavity-forming interactions. But thedistinctly less destabilization of picrate ion appears to result from the opposing effects of decreased acid-base inter-action of hydrophilic NO; and 0- groups and tbeincreased dispersion interaction on the bydrophobic benzenoid ring.

As a part of our systematic studies for gaininginsight into ion-solvent interactions on theone hand and for establishing a universal

scale of activity, acidity and basicity as well aselectrode potentials on the other, We have recentlyreported the free energies of transfer, /'::,G,o , ofH+X- (refs 1-4a), H+OH- (ref. 4b, Sa, b), M+X-(ref. 4b, 5c, d, 6,7) and of their individual ions indifferent aquo-organic solvents, using the widelyuseds-loa reference electrolyte (RE) method andemploying tetraphenylarsoniumtetraphenylboride(Ph4 AsPh4B) as the RE, which has been recentlyprovedll,12 to be most sound, self-consistent andreasonable among the different extra-thermodynamicmethods lOb •

The results in aqueous ethylene glycol (EG)1,5eappear to suggest that despite Bom-rype-" destabiliz-ing effect due to decreased dielectric constant, theobserved stabilization of alkali metal cations inEG-rich region is the result of possible formation=of electrostatically bound bidentate complex of thetype M+ (EGh. Such a possibility, however, willbe predominant if the cosolvent EG is replaced by2-methoxyethanol (ME) or 1,2-dimethoxyethane(DME) where the hydroxylic H-atoms of EG aresuccinctly replaced by electron donating methylgroups imparting increased electro negativity on thecomplexing Ovcentres of the cosolvents. In thepresent paper we report our studies on the freeenergies of transfer, /'::,G~ , ofMCl [M =Li, Na, K,Rb and Cs] and KX [X = Br and I] in the aqueousmixtures of 2-methoxyethanol. The results of thepresent study, it is hoped, will help elucidate the truenature of the solvation of these cations in aqueousmixtures of glycols and their mono-alkyl ethers.Since ME is quasi-aprotic in nature due to the possibleintramolecular hydrogen-bonding, the results arelikely to reflect solvating propensities which will be

I

characteristic partly of protic EG and partly ofdipolar aprotic DME.

In the present studies, double cells of the type (A)have been employed, as before4,6,7,14

Ag-AgX/MX(m), S/M(Hg)/MX(m), W/AgX-Ag(A)

Here S, Wand m denote aqueous mixtures of 10, 30,50 or 70 wt % ME, pure water and the molalitiesof the electrolyte MX respectively. /'::,G~ (i) valueswere estimated by the reference electrolyte (RE)assumptions-loa,

/'::,G~ (Ph4As+) = /'::,G~ (Ph4B-)= 1/2 /'::,Gf (Ph4AsBPh,,) ... (1)

/'::,G~ values of the RE (Ph4AsBPh4) were obtainedusing Eq. (2)

/'::,G~ (Ph4AsBPh4) = /'::,G~ (Ph4AsPi) +/'::,Gf (KBPh4) - /'::, G~ (KPi) ... (2)

where Pi = picrate and the corresponding /'::,G~values of Ph4 AsPi, KBPh4 and KPi being obtainedfrom solubility studies.

Materials and Methods2-Methoxyethanol (E. Merck) was distilled twice

before uSle15• Solvents of different compositions(wt %) were prepared using triply distilled water.The origin and the pretreatment of alkali halides(MX) used, the preparation and treatment of themetal amalgams including precautions involvedtherein, the construction of amalgam electrode,Ag-AgX electrodes and the cell used, and the pre-paration and estimation of the concentrations of thecell solutions used were essentially similar to thosedescribed in the previous works4c,5a,d,16.

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INDIAN J. CHEM.. VOL. 20A. APRIL 1981

Emf's were measured with the help of a Leeds solvents were taken from literature>. The plots ofand Northrup K-4 potentiometer in conjunction 6£::' 1 versus m were linear in all the cases and whenwith a moving coil galvanometer (Cambridge Instru- extrapolated to m = 0 yielded the values of 6£::'ment Co.). The general experimental procedure has values as intercepts. The 6£':,,' values in eachbeen described earlier4c,5a,d,14. case when fitted to linear plots by the method of

Methods of preparation and purification of KPi,"_ least squares al~o. yielded the 0same 6£~ va~ue.s.Ph4AsPi and KBPh" as well as the procedure for the The stadard deviations III 6£m values he withindetermination of saturated solubilities in these mixed ±0.5 mY.solvents were essentially similar to those described 6G~ (MX) values from water to the mixed solventearlier4c,5e,r,17b. The dilution of the aliquots of the have been computed on the mole fraction scale bysaturated KBPh4 solution was made with acetonitrile the relation (7)+ water (1.: .1) mixture, whereas that for KPi 6m (MX) = F6E~, -2RT In M./M ••and Ph.Asf'i With acetone + water (1 : 1) mixture. .The appropriately diluted solutions were analyzedspectrophotornetrically using a Perkin-Elmer (model200) spectrophotometer. As in earlier studies4b,5e,f,6,7triplicate sets were used and the estimated averageerrors are of the order 1% for Pi- and 2 % forPh4B- ion.

ResultsThee.m.f. values (6£)t of the double cell (A) for

different molalities (m) of MX in different aquo-MEmixtures were obtained and utilized 4c,5e,r,14 toconstruct the following function6£ = 6£°' - 2 k [Is (m) - fw (m)]

= 6E::' - 2 k 6 b m ... (3)

In Eq. (3) f (m) = AoBo m1/2 (1 + ~Bo mI12)-1 ~In (1 + 0.002 M, m) ... (4)

The subscripts m, sand w denote the molal scale,mixed solvent and water respectivly, k = RT/F.M, is the appropriate molecular weight of the mixedsolvents or water, ao is the ion-size parameter, AoBoand Bo are the Debye-Huckel constants in S1 unitsas given by Eqs (5) and (6) respectively

AoBo = 4.2006 X 106 d.lI'!. (E.T)-3f2 mol-l l!lkg1f2... (5)

s, = 50.29 X 1010 d.112(E.n-1Il m> mol-1f2 kg112... (6)

where E. is the dielectric constant and d. is the densityof the solvent. As in earlier studies4c,5e,rsince altera-tion .of a« values caused no significant change in6£::' values, the value of ao was taken to be zeroin all the cases. The required E. and d, values for the

. .. (7)

where M•• is the molecular weight of water. The6G~ (MX) values are given in Table 1 and arecorrect within ± 0.05 kJ mor=.

The observed saturated solubilities (S) in moldm-a of KBPh4, KPj and Ph4AsPi are presented inTable 2.

Assuming complete ionization, as observed inethanol-water mixturesl8a,b and used in other aquo-organic mixtures of comparable dielectric cons-tants4c,5e,f pKc values of the salts on molar scalespwere obtained by the relation (8)

pKc = -2 log S - 2 log y±sp

The required mean activity coefficient (y±) of thesalts were obtained as before4c,5e,rusing the extendedform of Debye-Huckel equation-log y± =A B d 1I2S112_0_0 8 __ [(1+ao B d 1/2S1/2)-1I (aOB d 1/2S1/2)-I]2 In 10 + o· T _ 0 s

- (log (d ~ 0.001 SM + 0.002 SM2)/d.

... (8)

... (9)

TABLE 1 - STA,NDARDFREE ENERGIESOF TRANSFER/'>,. G~ OFMX FROMWATERTO 2-METHOXYETHANOL+WATERMIXTURES

AT 25°C

[Values in kJ mol=']

Wt% LiCl NaCl KCl RbCl CsCl KBr KI nci-ME10 0.88 1.14 1.00 0.86 0.89 0.59 0.34 0.2630 2.09 2.85 2.73 2.29 2.37 1.65 0.26 0.8150 4.03 5.03 4.85 4.70 4.42 3.16 -0.04 1.6670 6.36 7.10 6.85 6.66 6.55 4.41 -0.22 4.29

(a) Ref. 14

TABLE2 - SOLUBILITIES(S), MEANACTIVITYCOEFFICIENTS(Y±) ANDTHESOLUBILITYPRODUCTS(K~p) OFKBPh., Ph.AsPi ANDKPi IN AQUEOUS2-METHOXYETHANOLAT 25°C

S(mol dm-3) Y± pKs~ (molar scale)Wt%

Ph.AsPl KPi Ph.AsPi KPiME KPh.,B Ph.AsPi KPi KPh.B KPh.B(x 103) (x 104) (x 102)

lOS 7.53 8.92 3.4110 034 0.82 210 0977 0.994 0.843 6.96 818 3.5030 098 2.27 2.53 0.953 0.976 0.813 6.06 7.31 3.3750 0.90 11.04 435 0.877 0.929 0.697 4.93 5.98 3.0470 34.97 68.49 9.41 0.580 0.756 0.472 3.39 4.39 2.71

(a) Ref. 5e

+Derailed data can be had from the authors on request.

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BHATIACHARYA et al, : ION-SOLVENT INTERACTION IN 2-METHOXYETHANOL

where at and a~ are the ion-size parameters for thecation and anion of the salt respectively, M is themolar mass of the salt, and d is the density of thesolutions (assumed to be approximately equal to thatof the pure solvent, d.). The required Debye-Huckelparameters were obtained from Eqs. (5) and (6) andthe values for at and a~were taken as before4c,5e,f,loa.

a~+ = a;1-= 0.3nm; a;h As+ = a;h B- = 0.5nm4 4

The y± and pK~pvalues are presented in Table 2.Free energies of transfer, 6G? of the salts on the

mole fraction scale were evaluated using the relation(l0) and the values are given in Table 3

6G~ = 2.303RT rp(K.,,)~- p(Ksp)~] +4.606 RT log MlDd~

M.dw.. (10)

The values are accurate within ± 0.1 kJ mol-l andwere utilized to compute free energies of transferof the RE (Ph4AsBPhJ with the help of Eq. (2).These were then apportioned into individual ioniccontributions by the RE assumption expressed inEq. (1). 6G~ (i) values of other ions were thenestimated by the additivity rule from the 6G~values of the appropriate electrolytes and are listedin Tables 3 and 4. The estimated uncertainties ofthese values lie within ± 0.3 kJ mol ".

DiscussionThe variations of 6G~ (MX) with wt % cosolvent

for LiCI, NaCI, KCI, KBr, KI, KPi, KBPh4 and

TABLE 3 - STANDARDFREE ENERGIESOF TRANSFEROFREFER-ENCE ELECTROLYTESFROM WATERTO 2-METHOXYETHANOL+

WATERMIXTURESAT 25°C.

[Ll G; values in kJ rnol=']

Wt % Ph,AsPi KPh,B KPi Pi- Ph,B-ME

10 --4.22 -3.25 0.51 -0.2 ---4.030 -9.19 -8.39 -0.23 -0.5 -8.7

50(-S.O)

-16.78 -14.84 -2.11 -2.0 -14.8

70(-9.1)

-2S.86 -23.65 --4.00 -3.6 -22.7(13.3)

Values in parentheses refer to those in EG-water mixtures[vide ref. Se]

Ph4AsPi are shown in Fig 1. The behaviour ofRbCI and CsCI being nearly similar to that of KCI,6G~ (KCI) - composition profile represents thatfor those two salts as well.

Significantly enough, the observed behaviour ofthese salts and their relative order are essentiallysimilar to those observed in aqueous mixtures ofethanolv», glycol= as well as glycerol=. Thus, whilethe observed relative order, - 6 G~ (KPi) ~ - 6 G~(KBPh4) < - 6 G~ (Ph4 AsPi), is indicative of theresult of combined effects of dispersion as well ascavaity forming interac~ions6,7,10b,19of the large sizedtetraphenyl ions in particular, the relative order6G~ ~KCl» 6G~ (KBr) > 6G~ (KI) is probablydue to increased destabilization of the simple anionspartly as a result of decreased acidity of solventsresulting from quasi-protic nature of the cosolventand partly to the Born-type'> destabilization causedby decreased dielectric constant of the solvents. But

10

NoCIKCI_-::::::::::::!~~::::LiCI_ KBr

KI

'- KPi0

E -10..,.x

'::::-x~.-C>4

-20

KPh.a

-301

L---,-,':- -L. -L... ..!l- Ph.AaPio 10 .30 50 70

wt % C'osolvent

Fig. 1 - Variation of free energies of transfer (LlG~ ) of someelectrolytes with wt % 2-methoxyethanol (ME) in aquo-ME

solvent system at 2SoC

TABLE4 - STANDARDFREEENERGIESOFTRANSFEROFINDIVIDUALIONSFROMWATERTO 2-METHOXYETHANOL-WATERMIXTURESAT2SoC

[LlG; values in kJ rnol=']

Wt% Li+ Na+ K+ Rb+ Cs+ H+ Cl- Br 1-ME10 0.4 0.7 0.6 0.4 0.4 -0.2 O.S 0.0 -0.230 -1.0 -0.2 -0.3 -0.8 -0.7 -2.3 3.1 2.0 0.6SO -2.1 -1.0 -1.3 -1.4 -1.7 --4.4 6.1 4.4 1.270 -3.3 -2.S -2.8 -3.0 -3.1 -S.3 9.6 7.2 2.5

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INDIAN J. CHEM., VOL. 20A, APRIL 1981

the orderofl',Gt's ofMCI, namely LiCI< <NaCI>KCl > RbCI z CsCI being the result of variouscompetitive effects's, is apparently intriguing as insome other aquo-organic solvents4c,5C,G,7. Conse-quently, the analysis of the individual ion contribu-tions should be more rewarding.

l', GOt (i) versus composition profiles for differentions are depicted in Fig. 2. The profile for K+ isrepresentative of all the alkali metal cations ascorresponding l', G~ (i) values for other cations areonly slightly off form each other. l',GOt(i) valuesfor H+ were obtained from Sadek et al.'s data foruci».

Interestingly enough, as in aqueous mixtures ofethanolv=, l', GOt{H+) -composition profile exhibits aslight hump at water-rich region followed by a distinctminimum at higher cosolvent composition. Thus, asin the cases of cosolvent ethanol, the slight hump isindicative of the effect of the cosolvent-inducedpromotion of water structurew, whereas the distinctminimum is the result of opposing effects of increasedbasicity on the one hand and of the decreased diel-ectric constant as well as the possible replacementofHaO+ by alkoxonium ion on the other. The increasedbasicity of the mixed solvent: arises partly from themore basic monomeric water molecules as releasedby packing imbalance of waterl3>li and partly fromthe more basic methoxy oxygen of the cosolventcompared to that of water. In fact, IR studies indi-cate that as in EG21'22the predominant configurationof ME22'23 molecules in the liquid state is in thegauche-form to facilitate intramolecular hydrogen-bonding as shown in structure (I).

Simple consideration of molecular structure indicatesthat due to the presence of electron repelling inductiveeffect of CH3 group in -OCH3 ME is more basicthan EG and water. Also, the restricted availabilityof the H-bonded proton and the absence of a fullyavailable acidic H-atom imparting a quasi-aproticcharacter, make ME molecules less acidic than EGand water. Moreover, since this basicity and proti-city are likely to be relayed through co-operativestructure of Hvbondingw between the cosolvent andH20 molecules in the aqueous solutions of ME,similar' to that proposed in other aquo-organicsolvents-":" the possible H-bonded cosolvent-watercomplex of the type (II) is also more basic and lessacidic than water and EG-water mixtures. And thisexplains the larger basicity and lower acidity of thesesolvents compared to water and EG-water mixtures.

Similarly, the observed hump in l', G~-compositionprofiles for the alkali metal cations in water-richregion being similar to that for H + as well as to thatobserved in some other aquo-organic solvents5d,?c

it is attributable chiefly to the structural effects".But in ME-rich regions despite Born-type destabiliza-tion effect, the observed increasing stabilization ofM+ is due partly to the increased acid-base type15'24ion-solvent interactions guided by cation-o-centreelectrostatic interactions'< and partly to the possibleformation of electrostatically bound bidentate comp-

"lex of the type (III) which is similar to that proposedearlier in the cases of EG5e as well as DME27.Expectedly, the loss of stability of ME moleculesaccompanying the rupture of the H-bonds duringMf--ion solvation should be compensated by thepossible formation of four essentially electrostaticbonds in solvents of the type (III).

The observed increased destabilization of the halideions reflects the combined effects of Born-type-sdestabilization and the decreased anion-H-centretype acid-base interaction arising from the decreasedproticity of the mixed solvent. And their relativeorder of destabilization CI- > Be >1- also conformsto that expected from the decreasing acidity effect ofthe mixed solvent and the increased soft-soft interac-tion28 down the group of soft anions.

10 Cl

.<5'<l

K+

li+-_-=::'~r-H·

Pi-

'0 -10E...,.x.",.

-20

-3oQL----1LO----~~~--------~5~O------~70wt % Co solvent

Fig. 2 - Variation of free energies of transfer, t:" 0; (i), ofsome individual ions (obtained by Ph.AsBPh4 referenceelectrolyte assumption) with wt % 2-methoxyethanol in aquo-

ME solvent system at 25°C.

::o/H\ *H ~H3.YH > H .--H-O

H 0 H-O H

H R=CHl for ME H(iJ • H for EG [ID

350

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BHATIACHARYA et al. : ION-SOLVENT INTERACTION IN 2-METHOXYETHANOL

Unlike the halide ions, 6G~ (i) values of Ph4B-/Ph4As+ and Pi- are increasingly negative and themagnitude for Ph4B- exceeds that of Pi-. SinceBorn-type energies of transfer are likely to be smallfor the large-sized Ph4B-jPh4As+ at least, the obser-ved pronounced stabilization of tetraphenyl ions canbe attributed as in other solvents= 7,11>12, to thecombined effects of dispersion interactions, 6G~ -(int) and the cavity-forming interactions 6 G~ (cav),Now, as has been indicated earlierv? the less associa-ted is the cosolvent the less energy is required tocreate a cavity in it. ME being less prone to H-bond-ing than the structurally similar EG, aqueous mixturesof ME should be less associated than that of EG.Thus, one would expect the contribution of 6 G~(cav) to the total 6G~ in the former solvent systemto be more than that in the latter. Also, the interactionterm for ME-water mixture should also exceedthat of EG-water mixture due to the larger polariza-bility of the quasi-protic cosolvent ME comparedto EG. The combined effects should, therefore,make the 6 G~ of tetraphenyl ions in ME-watersystem fairly larger than that in EG-water systemas observed (vide Table 2).

The observed less stabilization of Pi- compared toPh4B- in ME-water system is evidently the result of theopposing effects, as in EG-water system=. Withincreasing proportions of the cosolvent whilst thedispersion interactions of the rr-electron cloud of thephenyl nucleus with the quasi-aprotic ME molecules(approaching in a perpendicular direction) increasesthe H-bonding interactions through all or some ofits Seven O-atoms each of which bears the share ofnegative charge delocalized in the aromatic nucleusdecrease.

The 6G~ (M+) versus (rJvl+)-l profiles in ME-water system are shown in Fig. 3. The observedsequence of 6 G~ (i) values of alkali metal cationswith a maximum at Na+ is typically similar to thatobserved in the aqueous mixtures of some othercation-solvating solvents, namely methanolw,DMS05d and DMP. Evidently, this should have acommon genesis. Thus, as has been stated earlier,6,7with increasing sizes of the cations the intensity ofshort-range interactions of the cations, be of acid-base type and/or of bidentate complexation, shoulddecrease. But according to HSAB principle due toPearson 28 the intensity of soft-soft interactions shouldincrease with the size of the cations. Evidently,these two opposing effects result in a maximum in6G~ (M+) versus (rJvl+)-l profiles as observed.Notably, as the softness of these mixed solvents islikely to increase with the more polarizable cosolventME compared to water, the superimposed soft-softinteractions result in larger deviations from the datapoints of Li+ and Na+ ions. This is partially trueas shown in Fig. 3. In view of the fact that the nega-tive slopes ofthe guiding lines dictated by the cationo-philicity of the mixed solvents as opposed by theBorn-type electrostatic effect6 G~ t,.L(i),the observeddeviations may not be in consonance with the truemeasure of soft-soft interactions. This of coursenecessitates normalization of 6 G~ (i) values by

(

1-0No+,

10",,'10 ME

o

--30 ",'"10 ME-1·0

7_ - 2·0oE...,'";:: -30

~C£<l

50",,% ME

70 wt % ME

-4.0

L---~0~6--~07~--~'~0---+'.2~--~'4~--1~6~--71~.(rM' i/A+-'

Fig. 3 - Plots of l::.G~ (M+) versus (rM+)-l for some alkalimetal cations (M+) in aqueous mixtures of 2-methoxyethanol

(ME).

subtracting 6G~ ," (i) of the ions. Unfortunately,this is difficult to assess as simple Born equation ishardly adequate for these small cations. However,the general trend of the observed results are fairlyindicative of the plausability of the above contention.

Acknowledgement

The authors thank the CSIR, New Delhi forfinancial assistance.

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