ds 133 (2007) 111–115www.elsevier.com/locate/molliq

Journal of Molecular Liqui

The dilution enthalpies of formamide in mixtures of waterand propanol at 298.15 K

Xu Wang, Rui Zhang, Li Xu, Ruisen Lin ⁎

Department of Chemistry, Zhejiang University, Hangzhou 310027, China

Received 18 December 2005; accepted 15 July 2006Available online 30 November 2006

Abstract

The dilution enthalpies of formamide in aqueous n-propanol and i-propanol solutions have been determined using a CSC-4400 isothermalcalorimeter at 298.15 K. The homogeneous enthalpic interaction coefficients in the range of propanol concentration (0%–30%) have beencalculated according to the excess enthalpy concept. The results are interpreted in terms of solute–solute and solute–solvent interactions.Hydroxyl groups present different efficiency in promoting hydrophobic interactions in dependence of their position in the interacting molecules.© 2006 Elsevier B.V. All rights reserved.

Keywords: Formamide; Propanol; Dilution enthalpy; Enthalpic pairwise interaction coefficients

1. Introduction

The study of the thermodynamic stability of the nativestructure of proteins has proved quite challenging and stillremains a subject of extensive investigation. But becauseproteins have complicated structure and some intricate effectson its structure, it is very difficult to study the interactionsbetween proteins directly. Therefore, a useful approach is tostudy their model compounds [1–3] (amino acids, small peptidesand their derivatives) and by using additivity schemes, it may bepossible to estimate thermodynamic properties of the completelyunfolded polypeptide chain of proteins. As model compounds ofpolypeptides [4,5], the hydrogen bond link with amide is themost important secondary structure in proteins. It plays animportant part in understanding the conformational stability ofproteins and providing insights into physiochemical phenomenain life [6].

In biological environment the majority of proteins exist inaqueous mixed solvents containing many organic substances.Mono- and bi-functional alkanols affect the conformationalstability of proteins [7]. Alcohol–water mixtures as solvents areimportant because of their mixed hydrophobic–hydrophiliccharacter [8]. The study of hydrophobic hydration and hydro-

⁎ Corresponding author. Fax: +86 571 87951895.E-mail address: wangxu.linda@163.com (R. Lin).

0167-7322/$ - see front matter © 2006 Elsevier B.V. All rights reserved.doi:10.1016/j.molliq.2006.07.008

phobic ‘non-bonding’ solute–solute interactions is an importantstep for understanding several phenomena in biological systems.The occurrence of non-bonding interactions in aqueous solutionswas largely emphasized, but the influence of the mutual positionof the –OH groups was not completely elucidated. So the presentstudy is aimed at examining the homogenous enthalpic interactioncoefficients of formamide in aqueous n-propanol and i-propanolsolutions of different compositions.

2. Experimental

2.1. Reagents

Formamide (AR) was kept over CaCO3 overnight and thendistilled twice at reduced pressure. n-propanol (AR) and i-propanol (AR) were distilled, respectively, under atmosphericpressure with middle fraction collected. All of the purifiedproducts were stored over P2O5 in a desiccator before use. Thewater used for the preparation of solutions was deionized anddistilled using a quartz sub-boiling purifier.

Both the aqueous solutions, which were used as mixedsolvents (water+propanol), and the formamide solutions(formamide+propanol+water ) were prepared by mass usinga Mettler AE 200 balance precise to ±0.1 mg. The mass percentof alcohols are from 0% to 30%. All the solutions weredegassed and used within 12 h after preparation.

112 X. Wang et al. / Journal of Molecular Liquids 133 (2007) 111–115

2.2. Calorimetric procedure

The enthalpies of dilution for formamide in aqueouspropanol solution were measured with an isothermal calorim-eter (model 4400 IMC, Calorimeter Science Corporation, USA)at 298.15 K. The flow-mixing system is comprised of twoCSC4442 flow mixing cells [9,10] and two syringe pumps(model: 260D, ISCO Inc., USA). The variation in flow rates ofthe syringe pumps is less than 0.2%. The relative meandeviation of thermal powers determined was 0.3%. The flowrates were determined by weighing the mass of liquids through

Table 1Enthalpies of dilution of formamide in aqueous n-propanol solution at 298.15 K

w (n-propanol) mi/(mol·kg−1) mf /(mol·kg−1) Δdil Hm/(J·m

0.0000 2.9581 2.3134 52.622.9581 2.2427 58.742.9581 2.1568 66.272.9581 2.0481 75.402.9581 1.9020 88.342.9581 1.6970 106.792.9581 1.3989 132.93

0.0500 3.0118 2.3547 90.033.0118 2.2826 99.513.0118 2.1950 114.043.0118 2.0843 130.323.0118 1.9355 154.433.0118 1.7267 186.813.0118 1.4231 238.13

0.1000 3.0037 2.3468 136.443.0037 2.2748 153.353.0037 2.1873 170.963.0037 2.0768 196.523.0037 1.9282 230.373.0037 1.7199 280.663.0037 1.4171 357.74

0.1500 2.9948 2.3444 161.502.9948 2.2730 179.682.9948 2.1861 205.872.9948 2.0764 235.792.9948 1.9287 270.892.9948 1.7213 331.892.9948 1.4195 415.95

0.2000 3.0010 2.3446 139.843.0010 2.2727 160.713.0010 2.1853 175.603.0010 2.0749 205.753.0010 1.9264 240.683.0010 1.7183 296.163.0010 1.4157 377.83

0.2500 3.0044 2.3492 131.633.0044 2.2773 148.013.0044 2.1900 170.063.0044 2.0796 194.843.0044 1.9312 227.383.0044 1.7230 277.543.0044 1.4201 350.96

0.3000 2.9893 2.3394 127.192.9893 2.2680 143.032.9893 2.1813 159.692.9893 2.0716 182.402.9893 1.9242 214.892.9893 1.7171 260.472.9893 1.4159 330.52

the pump during 5 min. The calorimeter was tested on severaloccasions by performing dilution experiments of urea solutionsand the results obtained from these agreed very well with bestvalues available in the literature [11].

The enthalpies of dilution ΔdilHm were calculated from theequation

DdilHm ¼ Pð1þ miMÞ=mi f2 ð1Þ

in which P is the dilution thermal power (μW), mi is the initialmolality of the formamide solution (mol·kg−1), M is the molar

ol−1) mi/(mol·kg−1) mf /(mol·kg−1) Δdil Hm/(J·mol−1)

2.9581 1.1071 159.622.9581 0.9137 176.992.9581 0.7792 187.362.9581 0.6810 197.462.9581 0.6043 206.542.9581 0.5419 211.60

3.0118 1.1261 286.133.0118 0.9293 323.233.0118 0.7924 341.563.0118 0.6925 361.243.0118 0.6145 374.283.0118 0.5511 384.31

3.0037 1.1210 432.093.0037 0.9249 483.013.0037 0.7886 519.783.0037 0.6891 547.473.0037 0.6114 567.053.0037 0.5483 589.59

2.9948 1.1239 508.602.9948 0.9278 570.062.9948 0.7914 613.692.9948 0.6918 641.962.9948 0.6140 681.052.9948 0.5506 692.49

3.0010 1.1199 461.243.0010 0.9239 521.653.0010 0.7878 568.813.0010 0.6884 600.303.0010 0.6108 624.843.0010 0.5477 658.79

3.0044 1.1238 424.653.0044 0.9274 479.353.0044 0.7909 519.193.0044 0.6912 552.783.0044 0.6133 564.543.0044 0.5500 597.91

2.9893 1.1209 404.082.9893 0.9252 454.642.9893 0.7892 490.562.9893 0.6898 517.132.9893 0.6122 536.682.9893 0.5490 557.21

113X. Wang et al. / Journal of Molecular Liquids 133 (2007) 111–115

mass of formamide (kg·mol−1) and f2 is the flow rate offormamide solution (mg·s−1). The uncertainties of all ΔdilHm

values owing to duplicate runs at each initial molality and theslight variations of flow rates are within 1%. The final molalitymf was calculated from the equation

mf ¼ mi f2=½ f1ðmiM2 þ 1Þ þ f2� ð2Þ

in which f1 is the flow rate of diluent (aqueous propanolsolution).

Table 2Enthalpies of dilution of formamide in aqueous i-propanol solution at 298.15 K

w (i-propanol) mi/(mol·kg−1) mf /(mol·kg−1) Δdil Hm/(J·m

0.0000 2.9581 2.3134 52.622.9581 2.2427 58.742.9581 2.1568 66.272.9581 2.0481 75.402.9581 1.9020 88.342.9581 1.6970 106.792.9581 1.3989 132.93

0.0500 2.9807 2.3323 96.332.9807 2.2611 106.032.9807 2.1746 121.892.9807 2.0652 138.412.9807 1.9181 163.772.9807 1.7117 196.762.9807 1.4113 249.84

0.1000 2.9974 2.3440 143.882.9974 2.2724 158.662.9974 2.1852 181.672.9974 2.0752 206.122.9974 1.9271 246.042.9974 1.7194 296.742.9974 1.4172 375.35

0.1500 2.9839 2.3338 182.182.9839 2.2625 202.482.9839 2.1758 230.512.9839 2.0663 263.012.9839 1.9189 312.722.9839 1.7122 377.052.9839 1.4114 483.03

0.2000 2.9966 2.3431 195.132.9966 2.2714 215.172.9966 2.1843 246.292.9966 2.0742 279.622.9966 1.9262 333.302.9966 1.7185 402.072.9966 1.4164 508.24

0.2500 2.9975 2.3438 169.742.9975 2.2721 188.812.9975 2.1849 215.232.9975 2.0748 245.322.9975 1.9267 290.972.9975 1.7190 352.352.9975 1.4168 449.26

0.3000 2.9971 2.3424 149.972.9971 2.2706 165.662.9971 2.1834 188.352.9971 2.0732 216.592.9971 1.9250 254.962.9971 1.7172 309.482.9971 1.4150 394.15

3. Results and discussion

The information about the interaction mechanism can beinferred from the pair-wise interaction coefficients of thevirial expansion of an excess thermodynamic property.According to the McMillan–Mayer theory [12,13], all thethermodynamic properties of multi-components solutions canbe expressed by using a virial expansion in m which relatesthe non-ideal contributions of any total thermodynamicfunction to a series of interaction parameters. If aqueouspropanol solution is regarded as solvent, the excess enthalpy

ol−1) mi/(mol·kg−1) mf /(mol·kg−1) Δdil Hm/(J·mol−1)

2.9581 1.1071 159.622.9581 0.9137 176.992.9581 0.7792 187.362.9581 0.6810 197.462.9581 0.6043 206.542.9581 0.5419 211.60

2.9807 1.1171 304.442.9807 0.9221 340.622.9807 0.7865 366.702.9807 0.6874 385.332.9807 0.6100 397.762.9807 0.5471 415.06

2.9974 1.1216 456.142.9974 0.9256 516.162.9974 0.7894 553.522.9974 0.6899 592.262.9974 0.6122 605.952.9974 0.5490 632.55

2.9839 1.1170 583.302.9839 0.9219 665.892.9839 0.7862 707.342.9839 0.6871 748.952.9839 0.6098 785.892.9839 0.5468 805.15

2.9966 1.1209 620.152.9966 0.9250 704.152.9966 0.7888 754.252.9966 0.6894 807.632.9966 0.6117 829.412.9966 0.5486 858.91

2.9975 1.1212 540.032.9975 0.9252 613.562.9975 0.7890 651.442.9975 0.6895 694.172.9975 0.6119 722.632.9975 0.5487 744.21

2.9971 1.1195 475.552.9971 0.9237 539.452.9971 0.7877 572.242.9971 0.6883 607.982.9971 0.6108 631.672.9971 0.5477 649.07

Table 4Enthalpic interaction coefficients of formamide in aqueous i-propanol solutionat 298.15 K

w(i-propanol)

h2/(J·kg·mol−2)

h3/(J·kg2·mol−3)

h4/(J·kg3·mol−4)

mi/(mol·kg−1)

r

0.0000 −104.68 9.30 −1.69 2.9581 0.99980.0500 −209.90 10.61 0.25 2.9807 0.99990.1000 −371.35 47.80 −5.51 2.9974 0.99980.1500 −444.46 37.95 −2.09 2.9839 0.99980.2000 −494.43 52.60 −4.07 2.9966 0.99980.2500 −423.90 50.82 −6.19 2.9975 0.99980.3000 −324.66 12.46 1.36 2.9971 0.9999

114 X. Wang et al. / Journal of Molecular Liquids 133 (2007) 111–115

per kg of solvent (HE) of a solution containing formamide atmolality m is given by

HE ¼ h2m2 þ h3m

3 þ h4m4 þ L ð3Þ

in which h2, h3, h4, etc. are enthalpic pair-wise, triplet,quadruplet interaction coefficients, respectively. The molarenthalpy of dilution (ΔdilHm) of the solution from an initialmolality (mi) to a final molality (mf), is therefore given by

DdilHm ¼ HEmðmf Þ−HE

mðmiÞ ¼ h2ðmf−miÞ þ h3ðm2f−m

2i Þ

þ h4ðm3f−m

3i Þ þ N ð4Þ

The experimental values of ΔdilHm of formamide in aqueousn-propanol and i-propanol solutions together with the initial andfinal molalities are given in Tables 1 and 2. This data was fittedto Eq. (4) using a least-squares procedure to obtain the hcoefficients (Tables 3 and 4). As there are difficulties associatedwith the interpretation of higher coefficients [14], we havechosen to deal only with the coefficient representing pair-wisesolute interactions. The h2 value of formamide in pure water is−104.68 J·kg·mol−2, which is in good agreement with thatobtained by other workers [15]. From Fig. 1, it can be seen thatthe enthalpic pair interaction coefficients of formamide are allnegative in aqueous propanol solutions and pass through aminimum at about 0.20 mass fraction of propanol, while the h2coefficients of formamide in aqueous i-propanol solutions aremore negative than those in aqueous n-propanol solutions.

The physical meaning of the pair interaction coefficient inaqueous solutions is related to the changes in the thermody-namic property when two hydrated molecules are brought froman infinite distance, where only solute–solvent interactions areoperating, to a finite distance where hydrated solute–hydratedsolute interactions occur [16]. Formamide is able to formhydrogen bonds with water through both the carbonyl andamide moieties [17], which can be classified as a hydrophilic‘structure-breaking’ solute and gives negative contribution toh2. Moreover, formamide has no alkyl groups, which caninteract through the hydrogen bonds in pure water. It gives alsoexothermic contribution and is expected to result in a negativecontribution to h2 [18]. Hence, h2 of formamide in pure water isnegative.

Since the h2 coefficient relates closely to the solvent-mediated interactions between the two solvated molecules [19],

Table 3Enthalpic interaction coefficients of formamide in aqueous n-propanol solutionat 298.15 K

w (n-propanol)

h2/(J·kg·mol−2)

h3/(J·kg2·mol−3)

h4/(J·kg3·mol−4)

mi/(mol·kg−1)

r

0.0000 −104.68 9.30 −1.69 2.9581 0.99980.0500 −158.85 −10.84 3.76 3.0118 0.99990.1000 −297.48 15.41 0.33 3.0037 0.99990.1500 −396.30 47.30 −5.37 2.9948 0.99980.2000 −471.06 100.01 −14.68 3.0010 0.99980.2500 −408.97 88.80 −14.78 3.0044 0.99970.3000 −303.04 22.19 −0.03 2.9893 0.9999

the energetic effect arising from changes in the solvent structurein the vicinity of the dissolved particles leads to the variations inh2 coefficients in the mixed solvents. In the ternary solutionsunder investigation (formamide+propanol+water), the overalleffect on h2 reflects the equilibrium among the followingsuperimposed processes: (1) Hydrogen bond interactions be-tween two formamides (an exothermic process), which lead to anegative contribution to h2. (2) Interaction between formamideand alcohol (an exothermic process), which lead to a negativecontribution to h2. An endothermic process between formamideand the alkyl groups of alcohol is counteracted by an exo-thermic process between formamide and the hydroxyl groups ofalcohol. (3) A partial dehydration of the hydration shell offormamide and alcohol molecule (an endothermic process),which leads to a positive contribution to h2. The positive valuesof the h2 coefficients were interpreted as due to the prevailingrelease of structured water from the hydration co-spheres to thebulk.

In propanol solutions, the exothermic process is predominantover the endothermic processes, resulting in negative values ofh2 for formamide. The predominance increases gradually withthe increasing concentration of propanol in the mixtures. Theeffect is the most prevailing one at about 0.20 mass fraction ofpropanol, leading to a minimum of h2 coefficients. It isnoteworthy that the position of the observed minima of the h2

Fig. 1. The variations in h2 of formamide with the mass fractions of n-propanoland i-propanol in mixed solvents at 298.15 K: (▪) n-propanol; (•) i-propanol.

115X. Wang et al. / Journal of Molecular Liquids 133 (2007) 111–115

coefficients correspond very well with the maxima of thedissolution enthalpies of simple inorganic electrolytes both inwater–propanol and water–isopropanol mixtures [20,21]. Thephenomenon is related to the fact that alcohols may stabilize, orrigidify, the structure of aqueous solutions [22] up to somecritical alcohol concentration. Furthermore, the alcohol mole-cule including the hydroxylic group can be easily built into thethree-dimensional network of hydrogen bonds in water and theinfluence of alcohols on water depends mainly on the size andshape of the alkyl group in an alcohol molecule [23].

A general chemical principle is postulated namely that thestronger solutes are solvated less will be their propensity tointeract with other solutes [24]. In i-propanol solutions twomethyl groups and methane groups (CH) seem to screen thehydroxyl group in i-propanol, which lead to a stronger effect ofthe hydrophobic hydration than that of the propyl-group in n-propanol solutions. Consequently, this causes the phenomenonthat interactions between i-propanol and formamide are weakerthan that between n-propanol and formamide, which give a lessnegative contribution to h2. Similar differences are noticeable inthe case of the interactions of electrolytes such as NaCl and NaIwith the alcohols under consideration [25].

The preferential configuration model proposed by G.Castronuovo tries to rationalize the difference among positionalisomers of alkanols through the hypothesis that preferredinteractions exist in solution between groups having the sameaction on water structure (hydrophobic–hydrophobic, hydro-philic–hydrophilic), while mixed interactions (hydrophobic–hydrophilic) are unfavoured [26]. According to preferentialconfiguration theory, when formamide interacts with alcohol,they associate in the side-by-side manner [27]. For alkanols,assuming –OH as the functional group, nα is the number ofequivalent CH2 in the α-position to CH2 bearing a functional –OH groups (an alkyl CH2 group which corresponds to CH3=1.5CH2 and CH=0.5 CH2 [28]). In the side-on configuration α-CH2 are remote, meaning that their contributions to thehydrophobic interactions are small with respect to the otherCH2 groups in the molecule [29]. Nα of i-propanol is more thanthat of n-propanol, which lead to less effective towards theoverlap of the hydration co-spheres and give a less positivecontribution to h2 in i-propanol solutions.

Desonyers et al. [30] gave a general discussion forstructural interactions, and thought that structural interactionsmake a large contribution to the enthalpic function, andsometimes surpass the effect of other interaction [31,32]. Inour experiments the h2 coefficients of formamide in aqueous i-propanol solutions are more negative than those in aqueous n-propanol solutions, which testify that structural interactions

make a larger contribution to h2 than interactions betweenformamide and i-propanol.

Acknowledgement

The authors are grateful to the National Natural ScienceFoundation of China and to the National Education Committeeof China for financial support (No. 20273061).

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