Thermodynamic Study of (Perﬂuoroalkane Alkane) Mixtures

7/24/2019 Thermodynamic Study of (Perfluoroalkane Alkane) Mixtures

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Thermodynamic study of (perfluoroalkane + alkane) mixtures:Excess and solvation enthalpies

Celia Duce a, Maria Rosaria Tine a, Luciano Lepori b, Enrico Matteoli b,*

a Dipartimento di Chimica e Chimica Industriale, Universitadi Pisa, Via Risorgimento 35, 56126, Italyb Istituto per i Processi Chimico-Fisici del CNR, Area della Ricerca di Pisa, Via Moruzzi 1, 56124 Pisa, Italy

Received 2 March 2007; received in revised form 4 April 2007; accepted 4 April 2007Available online 14 April 2007

Dedicated to Professor J. Simoes Redinha on the occasion of his 80th birthday

Abstract

A newly designed calorimetric technique and calculation procedure have been used to obtain partial molar enthalpies, Hi, and excessenthalpies,HE, for binary mixture of hexane + perfluoro-n-alkanes (C5C8) and perfluorohexane +n-alkanes (C5C8), + cyclohexane,and + 2-methylheptane. All mixtures are endothermic, and the heat effects increase with the size of the second component. The HE andHi values found are the largest ever observed for mixtures of non-polar compounds. An estimate of the excess heat capacity for (perflu-orohexane + 2-methylheptane) has been obtained from Hi at two different temperatures. From Hi at infinite dilution and from theknown enthalpies of vaporization, the enthalpies of solvation, DsolvH

, have been evaluated either for alkanes and perfluoroalkanesin both hexane and perfluorohexane solvent. Solutesolvent interactions have been examined by describing the DsolvH

with an additivescheme of surface interactions and by applying the Scaled Particle Theory. The effects of chain lengthening, branching, and cyclizationhave been discussed. Perfluoroalkanes proved to be inert molecules that interact weakly with themselves as well as with alkanes. 2007 Elsevier Ltd. All rights reserved.

Keywords: Excess enthalpy; Partial molar enthalpy; Solvation; Alkanes; Perfluoroalkanes; Molecular interactions

1. Introduction

Fluorinated compounds have been widely investigatedduring the fifties and the anomalous behaviour of theirsolutions (vapour pressures, solubilities, presence of morethan one azeotrope, heat of mixing, and volume changesof mixing) was reviewed by Scott [1]. A renewed interest

in these substances has been arisen in recent years due tothe importance they have in both industrial (e.g.their useas refrigerants, propellants, and fire extinguishers) and bio-medical applications (e.g.their use as blood substitutes andin eye retina surgery). To make the state of the art and topromote research on thermochemical, thermodynamic andtransport properties of halogenated hydrocarbons and

their mixtures, three workshops were held under theIUPAC auspices in Pisa (1999), Paris (2001) and Rostock(2002).

In previous works [2,3], (liquid + liquid) equilibria,(vapour + liquid) equilibria and excess volumes, wereexamined for mixtures of perfluoro-n-alkanes CnF2n+2(Fn,n= 5 to 8) with hexane (H6) and ofn-alkanes CnH2n+2

(Hn, n= 5 to 8) with perfluorohexane (F6). In the presentwork we report excess enthalpies, HE, for the same systemsas well as for F6 + cyclohexane (cy-C6H12) and 2-methyl-heptane (MeC7H15). The systems (F5 + H6), (F6 + H5),and (F6 + H6) have been investigated in the whole compo-sition range, while the other mixtures only in the miscibilityregion at 298.15 K. Among the systems here studied, HE

data are reported in the literature only for (F6 + H6) [4].To obtain an estimate of the excess heat capacity, CEp,

the heat of solution for one system, (F6 + MeC7H15), hasbeen determined at two different temperatures.

0021-9614/$ - see front matter 2007 Elsevier Ltd. All rights reserved.

doi:10.1016/j.jct.2007.04.002

* Corresponding author. Tel.: +39 050 315 2515; fax: +39 050 315 2442.E-mail address:[email protected] (E. Matteoli).

www.elsevier.com/locate/jct

J. Chem. Thermodynamics 39 (2007) 13461353


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The analysis of the Gibbs free energies of solvationrevealed that perfluoroalkanes interact very weakly withalkanes and with themselves [2]. The aim of the presentwork is to investigate the enthalpy contribution to Gibbsfree energy in order to give a more precise answer on therelative interaction strength of CH2 and CF2 with the sol-

vent (F6 or H6). Moreover, the effect of the shape of thesolute alkane molecule on solutesolvent interactions wasinvestigated. To this purpose, the enthalpy of solvation,DsolvH

, in F6 and H6 of linear alkanes was compared withthat of a cyclic molecule such as cy-C6H12 and of abranched structure such as MeC7H15.

2. Experimental

2.1. Materials

All chemicals were commercial products of the bestgrade quality, and were used without further purification.

Their purity, in mass fraction, was checked by gas chro-matographic analysis. No significant difference was foundwith respect to the impurity content claimed by the factory:perfluoropentane, Fluorochem, 0.97; perfluorohexane, Flu-orochem, 0.99; perfluoroheptane, Fluorochem, mixed iso-mers, 0.98; perfluorooctane, Fluorochem, 0.98; pentane,Fluka, P0.995, water 60.005%; hexane, Fluka, P0.995,water 60.01%; heptane, Fluka, P0.995 GC, water60.05%; octane, Fluka, P0.995, water60.02%; 2-methyl-heptane, Aldrich, P0.99; cyclohexane, Fluka, P0.995.

2.2. Instrumentation

The heat effects were determined by using a heat-flowcalorimeter, the Thermal Activity Monitor (TAM) Mod.2277 by Thermometric (Sweden). The apparatus wasequipped with a stainless steel ampoule of 20 cm3 and witha dispensing unit composed of a Lund Syringe Pump Mod.6120 and a number of Hamilton gas-tight syringes ofcapacity ranging from 100 to 1000 ll. We made accurateinjections starting from a minimum of 1 ll and we mea-sured accurate heat effects as low as 0.1 J. To convert thevolume injected into mass, density values from our labora-tory were used for all substances. These are reported asmolar volumes intable 1.

2.3. Procedure and data treatment

The ampoule was filled with a weighed amount of one ofthe components in such a quantity as to leave no more than1 cm3 of vapour space. After the thermostatting of the dis-pensing unit and the thermal and (vapour + liquid) equili-bration of the ampoule, we carried out a series ofprogrammed injections of the components and we mea-sured the heat effect involved in each addition.

The procedure followed to obtain excess enthalpies fromheats of solutions is described and validated in [5]. In short,

the heat effect Q involved in each addition of a neat com-

ponent can be related to the partial molar enthalpy andthe number of moles of the components initially presentin the calorimetric ampoule, in the titration unit, and inthe final mixture. The presence of a vapour space and theheat involved in the (vapour + liquid) equilibration Qvlefollowing each injection is also taken into account(Q= Qexp+ Qvle). The parameters of the model equationfor the description of the excess enthalpy are finally deter-mined by minimizing the objective function OF by meansof a least-squares routine:

OF X

Qexp Qcalc2; 1

where Qexp and Qcalc are the experimental and calculatedheat effect, respectively.

As model equation, the RedlichKister (RK) expression,with three or four parameters, was found adequate:

HE x 1 x Xn

k1

ck2x 1k1; 2

wherex is the perfluoroalkane mole fraction. The parame-ters, ck, provide values of the molar excess enthalpy, H

E,and of the partial molar enthalpy of each component, Hi.Since in this technique the direct experimental measure-ments and the fitted quantities are heats of solution, theHi (and the H1i ) values obtained are by far more precisethan those obtained from the direct HE measurements bymeans of mixing-flow calorimetry.

3. Results

Due to the large amount of the experimental data (about500), each point (injection) being characterized by four

numbers (mass of 1 and 2 in the cell, mass of 1 or 2

TABLE 1Standard enthalpies of vaporization, DvapH

, vapour pressures, p*, molarvolumes, V*, of neat liquid compounds

Compound T/K DvapH/(kJmol1) p*d/kPa V*d/(cm3 mol1)

H5 293.15 57.65 115.2298.15 26.75a 68.34 116.1

H6 293.15 16.7 130.7

298.15 31.73a 20.17 131.5H7 298.15 36.66a 6.09 147.5H8 298.15 41.53a 1.86 163.5MeC7H15 288.05 40.6

b 1.15 162.64e

298.15 39.72a 2.76 163.65e

cy-C6H12 298.15 33.12a 13.01 108.65e

F5 293.15 71.03 177.8298.15 27.3b 86.21 178.9

F6 288.05 18.45 198.16e

293.15 23.98 200.5298.15 31.73c 29.32 201.6

F7 298.15 36.36c 10.18 224.3F8 298.15 40.96c 3.55 249.6

a Majer and Svoboda[6].b

Estimated by us from vapour pressures in reference[7].c Marchionni et al.[8].d Reference[2].e Data from our laboratory.

C. Duce et al. / J. Chem. Thermodynamics 39 (2007) 13461353 1347


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injected, andQ), we do not report here this extensive tab-ulation. These supplementary data may be obtained fromthe authors upon request. The procedure mentioned aboveto obtain the parameters of the fitting equation (equations(1) and (2)) requires some physico-chemical properties foreach component, such as molar volume, V*, heat of vapor-

ization, DvapH

, vapour pressure,p*

, as well as the activitycoefficients, ci, of the mixture components. This informa-tion, by us used, is collected intable 1, with the exceptionofc i, available in reference[2].

The parameters of the RK equation for HE (equation(2)) are reported intable 2together with the standard devi-ation,r(Q), of experimental points.Figures 1 to 3show theexperimental partial molar enthalpies Hi and the best fit-ting curves for all investigated mixtures. The extrapolatedH1i values of both components can be found in table 3together with the smoothed or extrapolated values ofHE

at x= 0.5. The uncertainties on H1i and HE, calculated

from those of single parameters, do not exceed

0.3 kJ

mol1 for H1i , and 0.02 kJ mol1 for HE.

Heat of mixing data are reported in the literature onlyfor (F6 + H6). Our HE, at x= 0.5, 2.10 kJmol1 com-pares well with the value 2.14 kJmol1 by Williamsonand Scott[4].

The standard molar enthalpies of solvation, DsolvH,

associated to the process: ideal gas ! infinitely dilute solu-tion, are also reported in table 3. DsolvH

were obtainedthrough the equation:

DsolvH H1 DvapH

3

where DvapH

is the standard enthalpy of vaporization(table 1). The magnitude ofDsolvH, and its variation with

molecular structure, mostly depend on DvapH since H1

contribute less than 40% to DsolvH. Consequently the

accuracy of this quantity is practically the same as thatofDvapH

which is usually better than 0.5% and does notexceed 0.3 kJmol1 for alkanes, and twice as much forperfluoroalkanes.

4. Discussion

4.1. Excess enthalpies

A glance atfigures 1 to 3and table 3shows that excessenthalpies HE as well as partial molar enthalpies H1i havevery large positive values for all examined mixtures. Theseenthalpy effects are the largest ever observed for mixturesof non-polar compounds. It can be noticed from table 3

TABLE 2RK coefficients ckfor H

E of equation(2), and standard deviation for Q, r(Q)

Mixture T/K c1/(kJmol1) c2/(kJmol

1) c3/(kJmol1) c4/(kJmol

1) r(Q)/mJ

(F6 + H5) 293.15 7.965 0.411 2.247 1.401 8.2(F6 + H6) 298.15 8.417 2.824 1.348 9.8

(F6 + H7) 298.15 8.640 3.378 0.668 10.9(F6 + H8) 298.15 7.665 5.257 0.264 7.8(F6 + MeC7H15) 288.05 8.213 3.091 0.312 8.0

298.15 8.695 3.014 0.261 4.3(F6 +cy-C6H12) 298.15 8.873 5.189 3.499 13.5(F5 + H6) 293.15 7.620 0.598 2.710 1.170 22.8(F7 + H6) 298.15 8.325 3.661 2.578 2.0(F8 + H6) 298.15 7.371 1.484 4.996 4.562 2.0

FIGURE 1. Plot of the molar excess enthalpy,HE, and of partial molar enthalpies, Hi, of the components vs. x of F6 for (F6 + H5) (a), (F6 + H6) (b),

and (F6 + H7) (c). (s) Experimental from TAM calorimeter; () calculated from parameters oftable 2.

1348 C. Duce et al. / J. Chem. Thermodynamics 39 (2007) 13461353


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that for both F6 + Hn and Fn+ H6 (n= 5 to 8) systemsHEs have similar values ranging from 1.9 to 2.2 and gener-ally increasing with increasingn. For alkanes Hnin F6,H12increases sensibily (8.6 to 12.7 kJ mol1) with increasingsize of dilute component. Similarly for perfluoroalkanesFnin H6,H1

1

increases withnfrom 10.9 to 15.4 kJ mol1.On the contrary, H11 for F6 in different alkanes slightlyincreases from 12.0 to 13.2 kJmol1, whereas H12 forH6 in different perfluoroalkanes slightly decreases from10.0 to 9.2 kJmol1. Table 3 also shows the effect ofbranching and cyclization on the excess enthalpies. Themixture (F6 + MeC7H15) displays lower H

1i than

(F6 + H8). For (F6 + cy-C6H12),HE andH1i are markedly

more positive than for (F6 + H6).It is seen fromfigures 1 to 3 that HE curves are nearly

symmetric. This pattern is exhibited either by mixturescompletely miscible in the entire composition range(F6 + H5, F6 + H6, and F5 + H6) and mixtures not com-

pletely miscible. When HE is plotted against the volume

fraction of perfluoroalkane, /, the maximum of the curveshifts to the range 0.55


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(F6 + MeC7H15) were used to estimate HE and H1i at

298.15 K from HE at 293.15 K for (F6 + H5) and(F5 + H6) (table 3).

According to the simple interpretative scheme proposedby Patterson [9], the positive HE and H1i values for theinvestigated mixtures can be qualitatively explained witha net destruction, during the mixing process, of the struc-tural order present in the neat liquids, perfluoroalkanesand especially alkanes. The positive GE [2]and VE [3] for

the same systems can be explained likewise. However, posi-tive CEp and CE;1p values indicate a more ordered structure

in solution. These apparently conflicting results are hard toexplain. The basic hypothesis that order destruction/for-mation brings about positive/negative YE (Y= H, S, G,V, and Cp) values is evidently an oversimplification, ordernot being a well defined property. Therefore, the Pattersonscheme is inadequate in some cases, though being an usefultool for a qualitative interpretation of the sign of the excessquantities for most mixtures.

Insight into the structure of the investigated solutionscame from calculation of the fluctuation or KirkwoodBuff integrals, Gij [2]. In summary, like-like interactions

(HH and FF) and the tendency to self-associationare very strong also in the limit of infinite dilution.The systems are near a critical solution temperature.The self-association and the (alkane + perfluoroalkane)antipathy, which is a direct consequence of the like-likeaffinity, increases with increasing the size of eithercomponent.

4.2. Enthalpies of solvationDsolvH

The transfer properties of the solute molecules from thegaseous state to infinitely dilute solution, besides an

amount related to the formation of the cavity to accommo-

date the solute, are determined by solutesolvent interac-tions only. Since we are interested in getting informationabout solutesolvent interactions, mostly related to thecontact among molecular surfaces, we choose to apply tothe enthalpy of solvation DsolvH

in hexane or perfluoro-hexane a simple scheme of surface based group contribu-tions. A similar approach gave satisfactory results whenapplied to DsolvG

and DsolvH of different organic solutes

in heptane[1012]perfluorohexane[2], and chlorinated sol-

vents [10,11,13]. In this model, the solute molecules aresubdivided into surface groups,j, each of which is assumedto contribute a constant amount Bj to DsolvH

, so thatDsolvH

can be expressed as:

DsolvH A

XnjBj 4

whereA is a constant term that depends on the solvent andnjis the frequency of the jth group in the solute molecule.TheA term is suggested by the non-zero intercept of a plotofDsolvH

of the homologous series ofn-alkanes or n-per-fluoroalkanes against any molecular descriptor related tosize (seefigure 4). This feature has been observed for differ-ent thermodynamic properties and the advantage due to itsuse have been described in detail elsewhere [14,15].

The CH2and CF2 contributions were first calculated asa mean value of the increments ofDsolvH

in the homolo-gous series ofn-alkanes and n-perfluoroalkanes. The CH3group has been assumed to contribute 1.57 times theBCH2 value on the basis of Bondi surface increments Sj[16]. Similarly, the CF3 was assumed to contribute 1.5times theBCF2 value. TheAvalue was evaluated from equa-tion(4)using DsolvH

and the known Bjcontributions. TheDsolvH

of alkanes in H6 were obtained from the knownH1i [17], which are usually negligible. TheDsolvH

of perflu-oroalkanes in F6 were assumed identical to DvapH

, sup-

posing H1i 0 in any case.

TABLE 3Excess enthalpy of equimolar mixtures, HE(0.5), partial molar enthalpies at infinite dilution, H1j , with standard deviations, r, for {perfluoroalkane(1) + alkane (2)} mixtures, and enthalpies of solvation of perfluoroalkane in alkane,DsolvH

1, and of alkanes in perfluoroalkane, DsolvH

2

Mixture T/K HE(0.5)/(kJ mol1)

H11 =kJ mol1 H12 =kJ mol

1 rH1i =kJ mol1 DsolvH

1=kJ mol

1 DsolvH2=kJ mol

1

(F6 + H5) 293.15 1.99 0.02 12.02 8.40 0.18298.15 2.05a 12.18a 8.63a 19.55 18.12

(F6 + H6) 298.15 2.10 0.02 12.59 9.89 0.18 19.14 21.842.14b

(F6 + H7) 298.15 2.16c 12.69 11.35 0.34 19.04 25.31(F6 + H8) 298.15 1.9c 13.19 12.66 0.33 19.54 28.87(F6 + MeC7H15) 288.05 2.05

c 11.62 10.99 0.08298.15 2.17c 11.97 11.45 0.10 19.76 28.27

(F6 +cy-C6H12) 298.15 2.23c 17.56 10.56 0.18 14.17 22.56

(F5 + H6) 293.15 1.91 0.01 10.90 9.76 0.15298.15 1.97a 11.06a 9.99a 16.24 21.74

(F7 + H6) 298.15 2.08c 14.56 9.41 0.12 21.80 22.32(F8 + H6) 298.15 1.9c 15.44 9.29 0.28 25.52 22.44(F7 +iso-C8H18) 303.15 2.11

b 9.2b

a Estimated from the data at 293.15 K using CEp or CE;1p;i of (F6 + MeC7H15).

b Williamson and Scott[4].c

Estrapolated using parameters in table 2.



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The results of the calculations of the group contribu-tions to DsolvH

are collected in table 4together with vander Waals surface increments S taken from Bondi [16].The standard deviation for BCH2 and BCF2 contributionsand for the constant term A is less than 0.5 kJmol1

and consequently the predicted DsolvH values of alkanes

and perfluoroalkanes have the same uncertainty. The Avalue for alkanes,AH, was found larger than that of perflu-oroalkanes AF in both H6 and F6. The value of BCF2

(4.41 kJ

mol

1) is more negative than BCH2

(3.58 kJmol1) in perfluorohexane, while it is less nega-tive in hexane BCF2 3:09 kJ mol

1; BCH2 4:96

kJ mol1. This indicates that CF2 interacts with hexanemolecules more weakly that CH2, though having a largersurface area.

A graphical representation of the different solvationbehaviour of alkanes and perfluoroalkanes both in hexaneand perfluorohexane is given infigure 4. The experimentalDsolvH

, together with calculated straight lines according to

equation(4), are plotted against the van der Waals surface

FIGURE 4. Enthalpies of solvation,DsolvH, cavity formation,Hcav, and interaction, Hint, ofn-alkanes (h), perfluoroalkanes (s), cyclohexane (n) and

branched alkanes () in perfluorohexane (a) and hexane (b) as solvents vs. the van der Waals surface area S. () calculated from equation(4); ()calculated according to[18].

TABLE 4Group contributions to DsolvH

and Hint in hexane and perfluorohexane solvents

Group S/(109 cm2 mol1)

DsolvH/

(kJ mol1)Hint/(kJmol1)

Hint(j)/Sj/(kJ109 cm2)

DsolvH/

(kJmol1)Hint/(kJmol1)

Hint(j)/Sj/(kJ 109 cm2)

Solvent: hexane Solvent: perfluorohexane

AH 3.8 3.8CH2 1.35 4.96 8.0 5.9 3.58 5.7 4.2CH3 2.12 7.79 12.6 5.9 5.62 8.9 4.2AF 2.5 1.1CF2 2.30 3.09 7.3 3.2 4.41 7.2 3.1

CF3 3.45 4.64 11.0 3.2 6.62 10.8 3.1



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area S[16]of the solutes in order to allow comparison ofthe interaction strengths at the same surface area. It is seenfromfigure 4thatDsolvH

of perfluoroalkanes are less neg-ative than those of alkanes in both solvents, the differencein DsolvH

being small (2 kJmol1 to 4 kJmol1) in per-fluorohexane and very large (>30 kJmol1) in hexane.

These differences in DsolvH

reveal that FF interactionsare slightly weaker than HF ones in perfluorohexane,and that FH interactions are much weaker than HH onesin hexane.

Figure 4 shows that the slope of the line representingperfluoroalkanes is smaller than that of alkanes in both sol-vents, the slopes and their differences being higher in hex-ane than in perfluorohexane. The values of these slopesare a measure of CH2 or CF2 interaction with solvent perunitS. They suggest, as already said, that perfluoroalkanesinteract with solvent molecules more weakly than alkanesin both solvents. The difference in slope and, consequentlyin interaction, is very large in hexane and small in perflu-

orohexane, as also argued from DsolvH values in the two

solvents.The branched 2-methylheptane shows a higher less

negative DsolvH than its linear isomer octane in both

solvents (see figure 4), the difference being larger in hex-ane (1.8 kJmol1) than in perfluorohexane(0.6 kJmol1). A similar and more pronounced behav-iour is exhibited by three-branched 2,2,4-trimethylpen-tane (the DsolvH

in F6 has been assumed to be thesame as in F7). Branching is likely to bring about a lar-ger cavity as well as a looser packing and consequently aweaker interaction with solvent molecules. This view is

supported by a similar difference ofDsolvH

between lin-ear and branched alkanes in heptane and CCl4 [11]. Theeffect of branching is smaller in perfluorohexane sincethis solvent interacts with any solute molecule more fee-bly than hexane.

It is evident fromfigure 4that the DsolvH of cyclohex-

ane in hexane is about 5.7 kJmol1 lower than thestraight line of n-alkanes. The difference reduces to4.4 kJmol1 in perfluorohexane. Similar differences wereobserved for different size cycloalkanes in CCl4 and hep-tane [11]. It can be argued that cyclic molecules, whichare more compact and have less degree of freedom than lin-ear molecules, need a smaller space to accommodate in thesolvent structure and consequently interact attractively in astrong manner with the solvent. The effect of cyclization,like that of branching is smaller in perfluorohexane thanin hexane. The reason is the same: the solutesolvent inter-actions are weaker in perfluorohexane, and consequentlythe difference in DsolvH

between cyclic and linear mole-cules is smaller in perfluorohexane.

4.3. Application of the Scaled Particle Theory

In order to obtain an estimate of the comparativeimportance of the effects connected to the formation of

the cavity and to the solutesolvent interactions in deter-

mining the enthalpy of solvation, we applied the ScaledParticle Theory (SPT)[18]to our systems. The SPT expres-sion ofDsolvH

is

DsolvH Hcav Hint RT aRT

2; 5

where Hcav and Hint being the cavitational and interac-

tional enthalpy terms, respectively, and a the coefficientof thermal expansion of the solvent. By using the equationof Pierotti [18], we have calculated Hcav for all solutes inboth H6 and F6. The hard sphere diameters, d, for eachsolute and solvent were taken from the literature [18] orevaluated from the heats of vaporisation according toSPT. For the solvents we employed the values: d(H6) =6.02108 cm and d(F6) = 6.59108 cm. The contribu-tion to DsolvH

due to solutesolvent interactions, Hint,has been obtained from equation(5).

The results of the above calculations are illustrated infigure 4, where Hcav and Hint are plotted, together withDsolvH

, against the solute surface area S. It is seen that

Hcav values are systematically positive and increase aboutlinearly with S in the size range here considered indepen-dently of the chemical nature of the solutes. This suggeststhat the negative values of experimental DsolvH

are exclu-sively due to the negative interaction terms Hint whosemagnitudes are nearly twice as much.

The cavitation enthalpyHcavis higher in hexane than inperfluorohexane. As suggested by SPT expression for Hcav,this is due to the loose packing (low number density) ofperfluorohexane solvent molecules. As a consequence, theHint values for both kinds of solutes are more negative inhexane than in perfluorohexane. In this latter solvent, Hint,

likeDsolvH

, is about the same for both Hnand Fnwith thesame surface area S, whereas in hexane the Hint values ofHnare sensibly more negative than those of Fn.

The above-described behaviour ofHint, as well as that ofDsolvH

(figure 4), indicate that perfluoroalkanes, Fn, arevery inert molecules, that interact weakly both with them-selves and with alkanes, Hn, the molecular interactionforces being in the order FF < FH < HH.

The almost strict linearity of both Hcav and Hint withrespect to the number of solute carbon atoms allows toobtain the contributions to Hcav and Hint of CH2 andCF2groups as mean increment ofHcavand Hint in homol-ogous series of alkane and perfluoroalkane solutes, respec-tively. The values ofHint(CH2) and Hint(CF2) in both H6and F6 are collected in table 4. They indicate a slightlyweaker interaction of CF2in hexane, and a stronger inter-action in perfluorohexane, Hint(CF2) being more negativethan Hint(CH2). In the assumption that the interactionsare due to contact between the solute and solvent surfacesand therefore that are proportional to the surface area, amore suitable comparison can be made by normalizingthe Hint contributions with respect to the group surfacearea S. These relative enthalpies, Hint(j)/Sj, reported intable 4, indicate that CF2 group has a lower interactioncapability than CH2 in both solvents, as discussed in the

previous section.



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JCT 07-85


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