Excess molar enthalpies for binary mixtures of cyclopentanone, cyclohexanone, or cycloheptanone with n-nonane at T = 298.15 K and atmospheric pressure

J. Chem. Thermodynamics xxx (2014) xxx–xxx

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Excess molar enthalpies for binary mixtures of cyclopentanone, cyclohexanone, or cycloheptanone with n-nonane at T = 298.15 K and atmospheric pressure Dongwei Wei ⇑, Sainan Han, Baohe Wang Key Laboratory for Green Chemical Technology of Ministry of Education, Research and Development Center for Petrochemical Technology, Tianjin University, Tianjin 300072, PR China

a r t i c l e

i n f o

Article history: Received 14 January 2014 Received in revised form 28 February 2014 Accepted 7 March 2014 Available online xxxx Keywords: Molar excess enthalpy Cyclopentanone Cyclohexanone Cycloheptanone n-Nonane Redlich–Kister equation

a b s t r a c t Excess molar enthalpies, HE, for the binary mixtures of cyclopentanone, cyclohexanone, or cycloheptanone with n-nonane were measured at T = 298.15 K and 0.1 MPa, by means of a Setaram Tian-Calvet MS80 microcalorimeter. All the binary systems investigated show endothermic behaviour (positive values) over the whole mole fraction range. The molar excess enthalpies decrease as the size of the cycloalkanone increases. The experimental results of HE are fitted to the Redlich–Kister equation to correlate the composition dependence. The experimental HE data are also used to test the suitability of the Wilson, NRTL, and UNIQUAC models. The correlation of excess enthalpy data in these binary systems using the UNIQUAC model provides relatively the most appropriate results. Ó 2014 Elsevier Ltd. All rights reserved.

1. Introduction From a theoretical point of view, the excess molar enthalpies can be used to study the energetic interactions between the molecules present in a mixture, such as the three van der Waals forces (orientation, induction, dispersion), and hydrogen bonding interactions, etc. Moreover, in modern separation design, an important part of many phase-equilibrium calculations is the mathematical representation of pure-component and mixture enthalpies. Mixture enthalpy data are important not only for determination of heat loads, but also for the design of distillation units. Furthermore, mixture enthalpy data, when available, are useful for extending (vapour + liquid) equilibria to higher (or lower) temperatures, through the use of the Gibbs–Helmholtz equation [1,2]. This work reports the experimental excess molar enthalpies of cyclopentanone, cyclohexanone, or cycloheptanone with n-nonane, respectively, at T = 298.15 K and 0.1 MPa. As far as we know, no previous HE measurements for the binary mixtures investigated were found in the literature despite their great theoretical and technical interest. The measured values were fitted by the Redlich–Kister equation. Thermodynamic models (the Wilson, NRTL, and UNIQUAC) based on the local composition theory were also

examined for the suitability by correlating experimental HE data with compositions. 2. Experimental 2.1. Materials All chemicals of cyclopentanone, cyclohexanone, cycloheptanone, and n-nonane were supplied by J&K Scientific Company (Beijing, China). The mass fraction purity of substances, checked by gas chromatography, was not less than 0.996. Evidence of chemical purity was also provided by comparison of measured refractive indices, n298:15K and densities, q298.15K with the literature D values, in table 1. Densities were measured using a vibrating-tube densimeter and a sound analyser, Mettler-Toledo model DM40. Refractive indices were measured using Abbe refractometer, Sanghai Shenguang model WYA-2S. The liquids were dried and stored over 4A molecular sieves and filtered through Millipore filter (0.45 lm). Before being used, all chemicals were degassed by evacuation. 2.2. Apparatus and procedure

⇑ Corresponding author. Tel.: +86 22 27406959; fax: +86 22 27406581. E-mail address: [email protected] (D. Wei).

The excess enthalpies were measured directly by calorimetry at T = 298.15 K and atmospheric pressure, using a standard system of

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D. Wei et al. / J. Chem. Thermodynamics xxx (2014) xxx–xxx

TABLE 1 Refractive indices (nD), and densities (q) of the pure components at T = 298.15 K and p = 0.1 MPa.a

a b c d

nDc

Components

Formula

Mass fraction purity

qb/(g  cm3) Exp.

Lit.

Exp.

Lit.

Cyclopentanone

C5H8O

0.9999

0.9443

0.944 [3] 0.944107 [4] 0.94435 [5] 0.9453 [6,7]

1.4347

1.4347 [6] 1.43476 [4] 1.43494 [8] 1.4352 [9]

Cyclohexanone

C6H10O

0.9999

0.9422

0.9410 [10] 0.9416 [11] 0.942039 [4] 0.9424 [12] 0.94276 [13]

1.4481

1.4480 [12] 1.4482 [6] 1.44837 [4]

Cycloheptanone

C7H12O

0.9997

0.9502

0.951 [14]

1.4605

n-Nonane

C9H20

0.9960

0.7140

0.713834 [15] 0.71402 [16] 0.71405 [17] 0.7148 [18]

1.4022

d

1.4608 [20]d 1.4021 [18] 1.4035 [19]

Temperature uncertainty, u(T) = 0.05 K, and pressure uncertainty, u(p) = 10 kPa. Density uncertainty, u(q) = 4  104 g  cm3. Refractive index uncertainty, u(nD) = 4  104. Measured and compared at T = 293.15 K.

the Tian-Calvet MS80 model by Setaram. A Sartorius A200S analytical balance (precision ±0.1 mg) was used to weigh the pure component masses. The uncertainties in the temperature and mole fractions were less than 0.05 K and 0.0002, respectively. The detection of MS80 is based on a three-dimensional fluxmeter sensor. The thermal effect, Q (J), in the calorimetric process can be determined by:

Q ¼ KA;

TABLE 2 Experimental excess molar enthalpies, HE, for the binary mixtures of cycloalkanones with n-nonane at T = 298.15 K and p = 0.1 MPa.a

ð1Þ

where A is the area of the corresponding measuring curve, m2; and K is the calorimeter constant, J  m2. The excess molar enthalpy for the binary solution, HE (J  mol1), at the temperature T of the experiment, is calculated from: E

H ¼ Q =n;

ð2Þ

where n is the total moles of mixtures. The apparatus and the procedure were tested by determining excess molar enthalpies, for the standard system (hexane + cyclohexane) at T = 298.15 K and atmospheric pressure. The results were found to differ by less than 1% from those of the literature [21] near x1 = 0.5. The uncertainty of HE value was estimated to be (0.01  HE) J  mol1. The experiment procedure and the reliability of the apparatus have been described in detail elsewhere [22–24]. 3. Results and discussion

x1

HE (J  mol1)

0.0455 0.098 0.1522 0.1989 0.2932 0.3864 0.4335 0.4912

Cyclopentanone (2) + n-nonane (1) 293 0.5382 529 0.5854 750 0.6878 920 0.7946 1142 0.8452 1308 0.8888 1360 0.9223 1414 0.9652

0.0508 0.1022 0.1495 0.2001 0.3005 0.3599 0.4152 0.4494

Cyclohexanone (2) + n-nonane (1) 260 0.5399 489 0.6204 670 0.7021 845 0.7681 1102 0.8559 1198 0.8825 1278 0.9435 1310

0.0453 0.1042 0.1544 0.2077 0.3021 0.4045 0.4571

Cycloheptanon (2) + n-nonane (1) 202 0.5010 443 0.5518 595 0.6065 793 0.7012 1036 0.8053 1195 0.9005 1241 0.9428

x1

HE (J  mol1) 1428 1430 1333 1098 902 718 520 276 1332 1296 1188 1043 770 645 356

1249 1241 1213 1110 870 519 302

a

Experimental values of molar excess enthalpies at T = 298.15 K and atmospheric pressure are given in table 2 and shown graphically in figure 1. In all the binaries, x1 is the mole fraction of the n-nonane. The excess enthalpies for cyclopentanone, cyclohexanone, or cycloheptanone with n-nonane are large and positive over the entire range of mole fractions. The curves of excess molar enthalpies vs. composition vary almost symmetrically and maximum positive values are (1430 J  mol1 at x1 = 0.5556), (1334 J  mol1 at x1 = 0.5253) and (1255 J  mol1 at x1 = 0.5252) respectively for mixtures: {cyclopentanone (2) + n-nonane (1)}, {cyclohexanone (2) + n-nonane (1)}, {cycloheptanone (2) + n-nonane (1)}. The excess enthalpy decreases as the size of the cycloalkanone increases. The same trend is found in the results of Kiyohara et al. [25] for mixtures of aliphatic ketones with n-alkanes. It was suggested by Kiyohara et al., that the large positive excess enthalpy for aliphatic

Standard uncertainties u are u(x1) = 0.0002, u(T) = 0.05 K, u(p) = 10 kPa and the combined expanded uncertainty Uc is Uc(HE) = (0.01  HE) J  mol1 (0.95 level of confidence).

ketones/n-alkanes arise mainly from the breaking of strong (dipole + dipole) interactions between ketone molecules. The same explanation can be extended to mixtures of cycloalkanone/nalkanes. The decreasing trend in HE of a cycloalkanone with n-nonane, as the size of the cycloalkanone increases, is apparently due to a weakening of the interactions between carbonyl groups (C = O) of higher cyclic ketones. 3.1. Fitting to the Redlich–Kister equation The composition dependence of experimental molar excess properties of the binary systems HE is described by the following Redlich–Kister equation [26]:

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D. Wei et al. / J. Chem. Thermodynamics xxx (2014) xxx–xxx 1600

TABLE 4 Constants of pure component: v, the molar volume; and qi, molecular area parameter; in the models of Wilson and UNIQUAC.

1400

1200

1000

a

-1

H /(J mol )

b

Component

v/(cm3  mol1)

Cyclopentanone Cyclohexanone Cycloheptanone n-Nonane

89.08 104.16 118.05 179.63

a

qb 2.800 3.340 3.880 5.476

Calculated from table 1: v = molar mass/q. Calculated from UNIFAC tables [30].

E

800

TABLE 5 Parameters of the Wilson equation with max relative deviation (RD) and average relative deviation (ARD).

600

400

Mixtures

200

0 0.0

C5H8O (2) + C9H20 (1) C6H10O (2) + C9H20 (1) C7H12O (2) + C9H20 (1) 0.2

0.4

0.6

0.8

FIGURE 1. Plot of excess molar enthalpy HE as a function of mole fraction x1 for the binary mixtures {C5H8O (2) + C9H20 (1)}, {C6H10O (2) + C9H20 (1)}, and {C7H12O (2) + C9H20 (1)} at T = 298.15 K: (j), C5H8O; (d), C6H10O; (N), C7H12O. The curves were calculated using the Redlich–Kister equation (parameters taken from table 3).

ð3Þ

i¼1

where x1 is the mole fraction of n-nonane; Ai is the adjustable parameter and m is the number of fitted parameters. The quality of the fit is assessed by means of the standard deviation, given as:

( SD ¼

n  2 X 1 HEcal;j  HEexp;j n  m  1 j¼1

)1=2 ð4Þ

;

where n is the number of experimental points. Values of the coefficients Ai are given in table 3, along with the standard deviations for the representation. Selection of the appropriate value of m in equation (1) was based on application of the F-test to the variations of SD. Figure 1 shows the experimental values of HE plotted against x1 together with fitted curves by using the Redlich–Kister polynomial. 3.2. Correlation by thermodynamic models The experimental HE data were also used to test the suitability of thermodynamic models (the Wilson, NRTL, and UNIQUAC equations) for representing experimental HE data over the entire range of compositions. The excess enthalpy, which indicates the temperature dependence of the excess Gibbs free energy, can be got through the Gibbs–Helmholtz equation:

" # E E 2 @ðG =TÞ H ¼ T : @T

Deviations/%

Dk12

Dk21

Max RD

ARD

3495.65 3506.09 3476.23

4269.25 3998.42 3873.10

47.0 45.5 43.5

44.4 40.2 35.5

1.0

x1

m   X 1 HE J  mol ¼ x1 ð1  x1 Þ Ai ð1  2x1 Þi1 ;

Parameters/(J  mol1)

ð5Þ

P;x

From equation (3), the resulting expressions for the excess enthalpy can be derived by substituting GE for the corresponding excess Gibbs free energy equations. The Wilson, NRTL and UNIQUAC expressions of HE are given by the following equations, respectively. Wilson [27]

    K12 x2 1 ðk12  k11 Þ HE J  mol ¼ x1 x1 þ K12 x2   K21 x1 ðk21  k22 Þ; þ x2 x2 þ K21 x1

ð6Þ

with

K12 ¼

v2 exp v1

  k12  k11 ;  RT

K21 ¼

v1 exp v2

  k21  k22 ;  RT

ð7Þ

where Dk12 ð¼ k12  k11 Þ and Dk21 ð¼ k21  k22 Þ are proportional to the interaction energy between molecules 1 and 2; v1 and v2 are the molar volume of pure component 1 and 2, calculated by their densities from table 1, were given in table 4. NRTL [28]

"   G21 ðg 21  g 11 Þðx1 þ x2 G21  x1 s21 a12 Þ 1 H J  mol ¼ x1 x2 ðx1 þ x2 G21 Þ2 # G12 ðg 12  g 22 Þðx2 þ x1 G12  x2 s12 a12 Þ ; þ ðx2 þ x1 G12 Þ2 E

ð8Þ

with

G12 ¼ expða12 s12 Þ;

s12 ¼

g 12  g 22 ; RT

G21 ¼ expða12 s21 Þ;

s21 ¼

ð9Þ

g 21  g 11 ; RT

ð10Þ

where Dg 12 ð¼ g 12  g 22 Þ and Dg 21 ð¼ g 21  g 11 Þ are adjustable binary parameters; and a12 is the non-randomness parameter.

TABLE 3 Coefficients Ai and standard deviations (SD) for representations of molar excess enthalpies at T = 298.15 K by equation (1). Mixtures

C5H8O (2) + C9H20 (1) C6H10O (2) + C9H20 (1) C7H12O (2) + C9H20 (1)

Ai/(J  mol1)

SD/(J  mol1)

A1

A2

A3

A4

A5

5704.5 5313.1 4995.2

953.7 495.6 149.3

616.2 869.1 727.0

326.9 270.9 2328.0

1573.8 673.6

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A5

2330.5

15.30 5.73 9.71

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D. Wei et al. / J. Chem. Thermodynamics xxx (2014) xxx–xxx

TABLE 6 Parameters of the NRTL equation with max relative deviation (RD) and average relative deviation (ARD). Mixtures

Parameters

Deviation/% 1

C5H8O (2) + C9H20 (1) C6H10O (2) + C9H20 (1) C7H12O (2) + C9H20 (1)

1

Dg12/(J  mol )

Dg21/(J  mol )

a12

Max RD

ARD

15218.70 14025.86 13817.08

7232.51 2737.51 1923.53

0.12 0.11 0.12

9.57 8.62 2.69

7.51 3.70 1.78

1600

TABLE 7 Parameters of the UNIQUAC equation with max relative deviation (RD) and average relative deviation (ARD). Mixtures

Deviations/%

Du12

Du21

Max RD

ARD

2481.94 1833.90 1457.04

89.00 43.99 22.48

8.79 7.49 3.15

6.38 3.34 1.59

1200

1000 -1

H /(J mol )

C5H8O (2) + C9H20 (1) C6H10O (2) + C9H20 (1) C7H12O (2) + C9H20 (1)

1400

Parameters/(J  mol1)

800

E

1600

600

1400 400 1200 200

0 0.0

-1

H /(J mol )

1000

800

0.2

0.4

0.6

0.8

1.0

E

x1

600

FIGURE 3. Plot of the excess molar enthalpy as a function of mole fraction x1 for the system {C6H10O (2) + C9H20 (1)} at T = 298.15 K: (d), experimental data; (s), Wilson; (4), NRTL; (h), UNIQUAC.

400

Three expressions of HE thus have two adjustable parameters that are optimally fitted by minimising the objective function (obj) between the calculated and the experimentally determined HE:

200

0 0.0

0.2

0.4

0.6

0.8

1.0

F obj ¼

x1

j¼1

FIGURE 2. Plot of excess molar enthalpy as a function of mole fraction x1 for the system {C5H8O (2) + C9H20 (1)} at T = 298.15 K: (j), experimental data; (s), Wilson; (4) NRTL; (h), UNIQUAC.

UNIQUAC [29]

    h2 1 HE J  mol s21 ðu21  u11 Þ ¼ q1 x1 h1 þ h2 s21   h1 s12 ðu12  u22 Þ; þ q2 x2 h2 þ h1 s12

ð11Þ

with

s12

h1 ¼

x1 q 1 ; x1 q1 þ x2 q2

h2 ¼

s21

 u u  21 11 ¼ exp  ; RT

x2 q 2 ; x1 q1 þ x2 q2

ð12Þ

ð13Þ

where Du12 ð¼ u12  u22 Þ and Du21 ð¼ u21  u11 Þ are the interaction energy parameters; q1 and q2 are the structural parameter of pure component 1 and 2, calculated from van der Waals area data, were given in table 4.

HEexp;j

!2 ;

ð14Þ

where the summation is over all jth data point. The adjustable parameters of each model, Dk12 ð¼ k12  k11 Þ and Dk21 ð¼ k21  k22 Þ in the Wilson equation, Dg 12 ð¼ g 12  g 22 Þ and Dg 21 ð¼ g 21  g 11 Þ, and a12 in the NRTL equation, and Du12 ð¼ u12  u22 Þ and Du21 ð¼ u21  u11 Þ in the UNIQUAC equation are obtained by nonlinear least-square fit, and summarised in tables 5–7 for each systems together with max relative deviation (RD) and average relative deviation (ARD) based on following equation:

RD ¼

 u u  12 22 ¼ exp  ; RT

n X HEcal;j  HEexp;j

   E  E Hcal;j  Hexp;j  HEexp;j

;

   E  E n H cal;j  H exp;j  1X ARD ¼ : n j HEexp;j

ð15Þ

ð16Þ

The correlation curves for molar excess enthalpies of {cyclopentanone (2) + n-nonane (1)}, {cyclohexanone (2) + n-nonane (1)}, and {cycloheptanon (2) + n-nonane (1)} are plotted in figures 2–4. The scatter map of residual deviations of the molar excess enthalpies for the models was plotted in figures 5–7.

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D. Wei et al. / J. Chem. Thermodynamics xxx (2014) xxx–xxx 1400

200

1200 0

/ J mol

800

-H

E

exp

-200

E

cal

600

H

E

-1

H /(J mol )

-1

1000

-400

400

200 -600

0 0.0

0.2

0.4

0.6

0.8

0.0

1.0

0.2

0.4

FIGURE 4. Plot of the excess molar excess enthalpy as a function of mole fraction x1 for the system {C7H12O (2) + C9H20 (1)} at T = 298.15 K: (N), experimental data; (s), Wilson; (4), NRTL; (h), UNIQUAC.

1.0

200

0

/ J mol

-1

-1

0

-200

-200

H

E

E

cal

cal

-H

-H

E

E

exp

exp

/ J mol

0.8

FIGURE 6. Plot of residual deviation DHE as function of mole fraction x1 for the system {C6H10O (2) + C9H20 (1)} using the thermodynamic models: (s), Wilson; (4), NRTL; (h), UNIQUAC.

200

H

0.6

x1

x1

-400

-400

-600

0.0

0.2

0.4

0.6

0.8

1.0

-600 0.0

FIGURE 5. Plot of the residual deviation DHE as a function of mole fraction x1 for the system {C5H8O (2) + C9H20 (1)} using the thermodynamic models: (s), Wilson; (4), NRTL; (h), UNIQUAC.

As can been seen from tables 5–7 and figures 1–4, Redlich– Kister equation can give a good correlation of enthalpies for the systems in this work. NRTL, and the UNIQUAC model can give a relatively good correlation of enthalpies, while the Wilson model has an obvious deviation. It may be due to the fact that the Wilson model is simpler, compared to NRTL and UNIQUAC models. Their GE models could have similar accuracy when they are used to correlate the VLE data. But, after the partial derivative is taken with respect to temperature (T) using the Gibbs–Helmholtz equation, the Wilson model (HE) may not be suitable for correlating the HE data.

0.2

0.4

0.6

0.8

1.0

x1

x1

FIGURE 7. Plot of residual deviation DHE as function of mole fraction x1 for the system {C7H12O (2) + C9H20 (1)} using the thermodynamic models: (s), Wilson; (4), NRTL; (h), UNIQUAC.

4. Conclusions Excess molar enthalpies for several binary mixtures involving cycloalkanones with n-nonane have been measured at the temperature 298.15 K and atmospheric pressure, over the whole mole fraction range. It was found that mixing process is endothermic for all binary systems. The values of HE for the mixtures containing n-nonane increase in the sequence:

cyclopentanone > cyclohexanone > cycloheptanon:

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D. Wei et al. / J. Chem. Thermodynamics xxx (2014) xxx–xxx

The Redlich–Kister equation successfully correlated the experimental excess molar enthalpy values of binary mixtures. The experimental excess enthalpy data were also correlated and predicted by means of solution theories, viz. the Wilson, NRTL and UNIQUAC models for binary mixtures, respectively. For all mixtures, relatively good agreement was achieved between experimental HE and calculated HE values as a function of mole fraction derived from NRTL and UNIQUAC models with the exception of the Wilson model. Of the two models, the UNIQUAC equation was found to be more appropriate for correlating the enthalpy of mixing results compared to the NRTL model for all systems with only two parameters. References [1] H. Dong, R. Zhang, W. Yan, S. Li, J. Chem. Thermodyn. 38 (2006) 113–118. [2] S.K. Jangra, S.K. Neeti, J.S. Yadav, J.S. Dimple, V.K. Sharma, Thermochim. Acta 530 (2012) 25–31. [3] J.-L.M. Abboud, R. Notario, Pure Appl. Chem. 71 (4) (1999) 645–718. [4] C. Bermúdez-Salguero, J. Gracia-Fadrique, J. Chem. Eng. Data 56 (2011) 3823– 3829. [5] O. Ciocirlan, M. Teodorescu, D. Dragoescu, O. Iulian, A. Barhala, J. Chem. Eng. Data 55 (2010) 3891–3895. [6] D. Dragoescu, A. Barhala, M. Teodorescu, Fluid Phase Equilib. 267 (2008) 70– 78. [7] A. Barhala, D. Dragoescu, M. Teodorescu, I. Wichterle, J. Chem. Thermodyn. 38 (2006) 617–623. [8] P. Brocos, A. Piñeiro, R. Bravo, A. Amigo, J. Chem. Eng. Data 47 (2002) 351–358. [9] M.V. Rathnam, S. Mohite, M.S. Kumar, J. Chem. Eng. Data 55 (2010) 5946–5952. [10] S. Singh, V.K. Rattan, S. Kapoor, R. Kumar, A. Rampal, J. Chem. Eng. Data 50 (2005) 288–292.

[11] M.N. Roy, B.K. Sarkar, B. Sinha, J. Chem. Eng. Data 54 (2009) 1076–1083. [12] J.N. Nayak, M.I. Aralaguppi, T.M. Aminabhavi, J. Chem. Eng. Data 48 (2003) 628–631. [13] O. Ciocirlan, M. Teodorescu, D. Dragoescu, O. Iulian, A. Barhala, J. Chem. Eng. Data 55 (2010) 968–973. [14] Cycloheptanon. Sigma-Aldrich [Online], , (accessed 2013). [15] D.C. Landaverde-Cortes, G.A. Iglesias-Silva, J. Chem. Eng. Data 53 (2008) 288– 292. [16] M.M. Piñeiro, J. García, B.E. de Cominges, J. Vijande, J.L. Valencia, J.L. Legido, Fluid Phase Equilib. 245 (2006) 32–36. [17] E. Cisneros-Pérez, C.M.R.-S. Germán, M.E. Manríquez-Ramírez, A. ZúñigaMoreno, J. Solution Chem. 41 (2012) 1054–1066. [18] J.G. Baragi, M.I. Aralaguppi, M.Y. Kariduraganavar, S.S. Kulkarni, A.S. Kittur, T.M. Aminabhavi, J. Chem. Thermodyn. 38 (2006) 75–83. [19] M. Huang, Z. Chen, T. Yin, X. An, W. Shen, J. Chem. Eng. Data 56 (2011) 2349– 2355. [20] W.M. Haynes, D.R. Lide, CRC Handbook of Chemistry and Physics, 92nd ed., Taylor & Francis, London, 2011. [21] K.N. Marsh, R.H. Stokes, J. Chem. Thermodyn. 1 (1969) 223–225. [22] J. Ortega, F. Espiau, R. Dieppa, Fluid Phase Equilib. 215 (2004) 175–186. [23] A. Amigo, J.L. Legido, R. Bravo, M.I. Paz-Andrade, J. Chem. Thermodyn. 22 (1990) 1059–1065. [24] A. Sanahuja, Thermochim. Acta 90 (1985). 9–4. [25] O. Kiyohara, Y.P. Handa, G.C. Benson, J. Chem. Thermodyn. 11 (1979) 453–460. [26] O. Redlich, A.T. Kister, Ind. Eng. Chem. 40 (1948) 345–348. [27] G.M. Wilson, J. Am. Chem. Soc. 86 (1964) 127–130. [28] H. Renon, J.M. Prausnitz, AIChE J. 14 (1968) 135–144. [29] D.S. Abrams, J.M. Prausnitz, AIChE J. 21 (1975) 116–128. [30] Aa. Fredenslund, G. Gmehling, P. Rasmussen, Vapour–Liquid Equilibria Using UNIFAC, Elsevier, Amsterdam, 1977.

JCT 14-147

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