Measurement and Prediction of Excess Enthalpies for Ternary

Jan 15, 2014 - The excess enthalpies of binary and ternary solutions composed of ionic liquid, 1-ethyl-3-methylimidazolium dimethylphosphate [Emim][DM...
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Measurement and Prediction of Excess Enthalpies for Ternary Solutions 1‑Ethyl-3-methylimidazolium Dimethylphosphate + Methanol or Ethanol + Water at 298.15 K and at Normal Atmospheric Pressure Xiaodong Zhang,* Dapeng Hu, and Zongchang Zhao Faculty of Chemical Engineering, Environmental and Biological Science and Technology, Dalian University of Technology, 2 Linggong Road, Dalian 116024, China S Supporting Information *

ABSTRACT: The excess enthalpies of binary and ternary solutions composed of ionic liquid, 1-ethyl-3-methylimidazolium dimethylphosphate [Emim][DMP], water, and methanol (or ethanol) were measured with an adiabatic calorimeter at 298.15 K and under normal atmospheric pressure. The experimental data for three sets of binary solution [Emim] [DMP] (1) + H2O(2)/CH3OH(2)/C2H5OH(2) were correlated with the Gibbs−Helmholtz, the nonrandom two liquid (NRTL), and the universal quasi-chemical (UNIQUAC) equations. The binary interaction parameters in the NRTL and UNIQUAC models were obtained and used to predict excess enthalpies of ternary solutions [Emim] [DMP](1) + CH3OH(2) or C2H5OH(2) + H2O(3). The results show that the excess enthalpies of ternary solutions are negative across the whole range of concentrations, which indicates that the [Emim] [DMP] mixed with H2O, CH3OH, or C2H5OH is exothermic. Meanwhile, the average relative deviations (ARD) between the experimental and correlated excess enthalpies for the three sets of binary solutions are 2.49 %, 1.64 %, 2.44 % (NRTL) or 2.98 %, 3.00 %, and 3.43 % (UNIQUAC). The ARD between the experimental and predicted excess enthalpies for two set of ternary solution are 7.89 %, 5.17 % (NRTL) and 1.86 %, and 1.86 % (UNIQUAC), respectively.

1. INTRODUCTION With the rapid growth of the world economy, the shortage of energy resources has grown into an increasingly serious problem in recent years. Developing new energy-saving technologies has become an urgent task for many industrial sectors. Compared to the mechanical vapor compressive refrigerator, the absorption refrigerator consumes much less electric power as it is driven by low-grade heat, such as solar energy, industrial waste heat, and so forth. The performance of the absorption refrigerator depends on the thermodynamic properties of the working fluids employed. Commonly used working fluids are the aqueous solution of lithium bromide (H2O + LiBr) and the aqueous solution of ammonia (H2O + NH3). However, these two solutions suffer from shortcomings. For example, the working fluid H2O + LiBr corrodes iron-steel equipment and crystallizes easily; furthermore, refrigeration temperature is limited above the freezing point of water. The other working fluid, H2O + NH3, can work at lower refrigeration temperature. However, its problems include its toxicity and explosive volatility. Furthermore, when H2O + NH3 is used, both a higher operation pressure and complex distillation equipment are also required. Ionic liquids (ILs) have emerged as new varieties of solvents. They have excellent physical and chemical properties, such as negligible vapor pressure, nonflammability, thermal stability, © XXXX American Chemical Society

and low melting points, they also have a wide liquid-state temperature range, from room temperature to 200 °C or 300 °C and good solubility with many organic or inorganic solvents. Therefore, searching for new alternatives for absorbents of common refrigerants by examining various kinds of ILs has been of considerable interest within recent years. Keeping those goals in mind, some researchers already began to look for suitable ILs and examine their thermodynamic properties, such as the vapor pressure, specific heat capacity, and excess enthalpy. Kim et al.1 measured the vapor pressure of the aqueous solutions of some ILs. Shiflett and Yokozeki2−4 examined the solubility of CO2 and NH3 in some ionic liquids and also predicted the water solubility and excess enthalpy in some ionic liquids by using the Redlich−Kwong cubic equation of state. Wang et al.5 measured and correlated the vapor pressures of some binary solutions composed of water and some ionic liquids. The authors6−8 measured and correlated the vapor pressure, excess enthalpy, and specific heat capacity of binary solutions composed of ionic liquid [Emim][DMP] and water/ethanol/ methanol. It is found that these binary solutions exhibit a great negative deviation from Raoult’s Law and show obvious Received: February 4, 2013 Accepted: January 6, 2014

A

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the resultant mixture with ether. Then the raffinate containing [Emim] [DMP] is evaporated under vacuum conditions and at a temperature 403.15 K using a rotary evaporator for 24 h in order to remove all volatile components such as water and residual ether. The purity of the prepared IL is above 98.0 % (wt) as determined by H1NMR (400 MHz, D2O, ppm). The water content in the IL is 0.024 % (wt) measured by a 756 Karl Fisher coulometer. 2.2. Apparatus and Procedure. The apparatus for measuring excess enthalpy is schematically detailed in Figure 2. It consists of two water baths with a temperature uncertainty of 0.05 K, an adiabatic calorimeter with a stirrer, a heater, and a thermometer with an uncertainty of 0.05 K. To prepare the sample solutions, a certain amount of IL and water/ethanol/methanol are weighed using a counter balance with an uncertainty of ± 0.01 g. A certain amount of [Emim] [DMP], deionized water and ethanol or methanol, (the total mass of three samples is about 80 g) are poured into the three graduated containers, and the mouths of graduated containers are sealed with rubber plugs. These graduated containers are placed in the water bath (part 3a in Figure 2) with temperature set at 298.15 K and kept in the water bath for at least 30 min to ensure that the temperature of the samples reached 298.15 K. Next, the IL, water, and ethanol (or methanol) in the three containers are quickly poured into an adiabatic calorimeter in the water bath (part 3b in Figure 2) whose temperature is set approximately at the determined final mixing temperature of the samples in order to avoid heat loss during the mixing process. The final temperature can be estimated by performing an experiment in advance by mixing the sample in the same adiabatic calorimeter and measuring the temperature of resultant solution. Mixed samples are stirred in the adiabatic calorimeter until their temperature peaks, and then the temperature of the resultant solution is recorded. The adiabatic calorimeter is placed in a slurry of ice and water for a long time until the temperature of the solution falls to 298.15 K. Then the adiabatic calorimeter is placed in the water bath (part 3b) again and the electric heater is turned on to heat the solution in the adiabatic calorimeter until the solution temperature reaches the final mixing temperature. The

negative excess enthalpies and good thermodynamic cycle performance. Enthalpy of the working fluids is one of the most important thermodynamic properties, which is used in the performance analysis and the design of refrigerators. Common working fluids used are nonideal solutions. Note that an excess enthalpy is defined as the enthalpy change when two or three different components are mixed to form a solution, a property which cannot be ignored in the performance analysis and the design work of refrigerators. It should be also noted that data for the excess enthalpy, as one of the basic thermodynamic properties of solutions, of some binary solutions which contain ILs have also been published in the past decade.9−14 In this paper the ternary solutions [Emim][DMP](1) + CH3OH(2) or C2H5OH(2) + H2O(3) are proposed as new working fluids for absorption refrigerators. Here the [Emim][DMP] is used as an absorbent for the binary mixed refrigerants CH3OH(2) or C2H5OH(2) + H2O(3). The binary mixed refrigerants combine the advantages of water and alcohol, which permits a refrigeration temperature below 0 °C and higher vaporization latent heat than pure methanol or ethanol. Therefore, measuring and predicting the excess enthalpies of ternary working fluids are of great importance to analyze the cycle performance of refrigerators which employ these ternary working fluids.

2. EXPERIMENTAL SECTION 2.1. Chemicals. N-Ethylimidazole (≥ 99 %) was purchased from Tianjin Fuchen Reagents Company. The ethanol, methanol, and trimethylphosphate (all of which were 99.8 % pure) were purchased from Sinopharm Chemical Reagent Company and used without further purification. The ionic liquid, [Emim] [DMP], was prepared according to the method given by Zhou et al.15 The reactive formula for [Emim] [DMP] is shown in Figure 1. The method is briefly introduced as follows.

Table 1. The Excess Enthalpies for Diethyleneglycol(1) + Water(2) at 298.15 K and at Normal Atmospheric Pressure Figure 1. Reactive formula for [Emim] [DMP].

HElit x1

First, an appropriate amount of N-ethylimidazole is poured into a flask with a reflux condenser and then mixed with equimolar trimetylphosphate. After reacting for 10 h at a temperature 423.15 K, the resulting mixture is cooled to room temperature. Unreacted agents are extracted many times from

HEexp −1

0.200 0.401 0.501 0.804 ARD =

RD −1

kJ·kmol

kJ·kmol

−1200 −1245 −1157 −459

−1219 −1217 −1149 −463

% 1.59 2.23 0.67 0.82 1.33 %

HElit values from ref 16. Standard uncertainties of temperature T and mole fractions x are as follows: u(T) = 0.05 K, u(x) = 0.0005. The expanded uncertainty with 0.95 level of confidence for the excess enthalpies HE is 10 kJ/kmol.

Table 2. Surface Parameters in the UNIQUAC Model

Figure 2. Schematic diagram of experimental apparatus for excess enthalpy: 1, thermometer; 2, stirrer; 3, thermostatic water bath; 4, adiabatic calorimeter; 5, heater; 6, stirrer; 7, constant voltage DC power supply.

a

B

species i

qi′

group k

Qk

[Emim][DMP] H2O CH3OH C2H5OH

6.384 1a 0.96a 0.92a

[Mmim][DMP] CH2

5.844b 0.54b

Data coming from ref 17. bData coming from ref 18. dx.doi.org/10.1021/je400158p | J. Chem. Eng. Data XXXX, XXX, XXX−XXX

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Table 3. Experimental and Correlated Excess Enthalpies HE for the Binary Solutions [Emim][DMP](1) + H2O(2)/ CH3OH(2)/C2H5OH(2) at 298.15 K and at Normal Atmospheric Pressurea HEexp x1

kJ·kmol

HENRTL −1

0.10 0.18b 0.20 0.31b 0.40 0.45b 0.50 0.64b 0.70 0.82b ARD =

−3686 −5291b −5439 −6304b −6120 −6021b −5777 −4807b −4216 −2529b

0.10b 0.20 0.30b 0.40 0.50b 0.65b 0.70 0.79b ARD =

−1572b −2815 −3360b −3572 −3351b −2594b −2329 −1820b

0.10 0.15b 0.20 0.31b 0.40 0.59b 0.70 0.80b ARD =

−1301 −1792b −2169 −2743b −2814 −2184b −1813 −1195b

−1

kJ·kmol

HEUNIQUAC −1

kJ·kmol

RDNRTL

RDUNIQUAC

%

%

[Emim][DMP](1) + H2O(2) −3360 −3913 8.84 −5216 −5115 1.42 −5499 −5323 1.11 −6275 −6008 0.47 −6263 −6105 2.33 −6087 −6011 1.09 −5818 −5822 0.72 −4669 −4850 2.87 −4036 −4256 4.26 −2573 −2792 1.74 2.49 % [Emim][DMP](1) + CH3OH(2) −1585 −1658 0.83 −2781 −2736 1.21 −3394 −3330 1.01 −3543 −3523 0.81 −3358 −3386 0.21 −2659 −2700 2.51 −2348 −2374 0.82 −1716 −1703 5.71 1.64 % [Emim][DMP](1) + C2H5OH(2) −1270 −1382 2.38 −1817 −1859 1.40 −2225 −2216 2.58 −2695 −2653 1.75 −2750 −2724 2.27 −2269 −2280 3.89 −1767 −1776 2.54 −1227 −1223 2.68 2.44 %

Table 5. Binary Interaction Parameters in UNIQUAC Model

6.16 3.33 2.13 4.70 0.25 0.17 0.78 0.89 0.95 10.40 2.98 %

i

j

αij

αji

ARD/%

[Emim][DMP] [Emim][DMP] [Emim][DMP] CH3OH C2H5OH

H2O CH3OH C2H5OH H2O H2O

−377.7 830.8 666.3 −356.9 −298.0

−418.2 −425.5 −393.5 252.7 201.2

2.98 3.00 3.43 4.87 33.69

is the excess enthalpy (kJ·kmol−1). The uncertainty of U, I, and t are 0.01 V, 0.001 A, and 0.01 s, respectively, in this paper. On the basis of eq 2, the absolute error and relative error of excess enthalpy can be given as follows: ΔHE = ε=

∂HE ∂HE ∂HE ∂HE ΔU + ΔI + Δt + Δn ∂U ∂I ∂t ∂n

(3)

ΔHE ΔU ΔI Δt Δn = + + + = 0.002 = 0.2 % U I t n HE (4)

5.47 2.81 0.89 1.37 1.04 4.09 1.93 6.43 3.00 %

The concentrations of the ternary sample solutions are designed as follows. First, the molar fractions of [Emim] [DMP], x1, is 0.2, 0.4, 0.5, and 0.7, respectively. Then the molar fraction ratio of methanol (or ethanol) to water, x2:x3, is taken as 1:4, 2:3, 3:2, and 4:1, respectively. Therefore, there are 16 solutions with different concentrations to be measured. The validity of the experimental apparatus was first checked by measuring mixing heat of diethylene glycol (> 99.8 wt %) with deionized water at 298.15 K and at normal atmospheric pressure, the molar concentration of diethylene glycol in the mixing solution ranges from 0.2 to 0.8. The experimental data are listed in Table 1, which show good agreement with those in literature16 and with an average relative deviation of 1.33 %.

6.23 3.74 2.17 3.28 3.20 4.40 2.04 2.34 3.43 %

3. THERMODYNAMIC MODELS The excess enthalpy of the solution can be obtained from the following Gibbs−Helmholtz equation.17 N ⎛ ∂ ln γi ⎞ HE = −RT 2 ∑ xi⎜ ⎟ ⎝ ∂T ⎠ i=1

a

Standard uncertainties of temperature T and mole fractions x are as follows: u(T) = 0.05 K, u(x) = 0.0005. The expanded uncertainty with 0.95 level of confidence for the excess enthalpies HE is 10 kJ/kmol. b Data obtained from our previous works in ref 6.

The activity coefficient of the component, γi, in eq 5 can be obtained from some thermodynamic models such as WILSON, NRTL, and UNIQUAC. NRTL and UNIQUAC are employed in this paper. NRTL Model. In the NRTL model,17 the activity coefficient of the component i is given as follows.

heat added to the adiabatic calorimeter, which is equal to the excess enthalpy of the solution, is given as follows Q = UIt

(1)

N

Q UIt H = = n n

(5)

p,x

E

ln γi =

(2)

where Q is the heat added into the adiabatic calorimeter. U, I, and t are electric voltage (V), electric current (A), and time (s), respectively. n is the mole of the sample solution (mol), and HE

∑ j = 1 τjiGjixj N

∑l = 1 Glixl

N

+

∑ j=1

N ⎛ ∑ xτ G ⎞ ⎜τ − r = 1 r rj rj ⎟ ij N N ∑l = 1 Gljxl ⎜⎝ ∑l = 1 Gljxl ⎟⎠

xjGij

(6)

Gij = exp( −αijτij), Gji = exp(−αjiτji)

(7)

Table 4. Binary Interaction Parameters in NRTL Model i

j

αij

gij − gjj

gji − gii

ARD/%

[Emim][DMP] [Emim][DMP] [Emim][DMP] CH3OH C2H5OH

H2O CH3OH C2H5OH H2O H2O

0.0978 0.2691 0.3971 −0.3201 −0.2711

−7617 −7161 −4993 −2570 −2237

72933 13954 11159 −7027 −11671

2.49 1.64 2.44 2.22 8.81

C

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Table 6. The Experimental and Predicted Excess Enthalpies for Ternary Solutions [Emim][DMP](1) + CH3OH(2)/ C2H5OH(2) + H2O(3) at 298.15 K and at Normal Atmospheric Pressurea HEexp x1

x2

HENRTL −1

kJ·kmol

0.200 0.397 0.501 0.701

0.160 0.121 0.099 0.062

−4625 −5556 −5140 −3860

0.201 0.404 0.500 0.695

0.318 0.240 0.200 0.125

−4173 −4932 −4515 −3410

0.200 0.401 0.502 0.695

0.480 0.359 0.298 0.184

−3684 −4406 −4118 −3098

0.201 0.401 0.500 0.699 ARD =

0.638 0.481 0.398 0.240

−3143 −4033 −3644 −2779

0.200 0.399 0.500 0.702

0.159 0.121 0.100 0.065

−4407 −5292 −4936 −3695

0.201 0.400 0.499 0.701

0.319 0.240 0.200 0.116

−3852 −4683 −4398 −3240

0.200 0.401 0.500 0.700

0.477 0.359 0.300 0.180

−3380 −4119 −3862 −2620

0.201 0.401 0.500 0.700 ARD =

0.638 0.479 0.400 0.241

−2798 −3383 −3135 −2230

−1

kJ·kmol

HEUNIQUAC

RDNRTL

RDUNIQUAC

kJ·kmol−1

%

%

12.02 5.60 5.56 3.78

0.06 2.66 0.78 2.38

14.07 9.10 10.61 1.17

2.18 2.09 2.75 0.03

14.33 10.37 9.32 0.74

2.69 1.27 1.29 1.48

14.00 5.08 9.11 1.44 7.89 %

0.48 3.12 3.38 3.17 1.86 %

11.75 5.97 5.33 4.19

3.56 0.27 2.82 1.22

11.86 5.59 4.46 2.22

1.32 2.03 0.00 0.59

8.20 3.13 1.92 3.21

2.28 4.66 2.95 3.44

6.61 4.02 4.27 0.00 5.17 %

1.87 2.25 0.35 0.13 1.86 %

[Emim[DMP](1) + CH3OH(2) + H2O(3) x2: x3 = 1:4 −5181 −4628 −5867 −5408 −5426 −5180 −3714 −3768 x2: x3 = 2:3 −4760 −4082 −5381 −4829 −4994 −4639 −3450 −3411 x2: x3 = 3:2 −4212 −3585 −4863 −4350 −4502 −4171 −3121 −3052 x2: x3 = 4:1 −3583 −3158 −4238 −3907 −3976 −3767 −2739 −2691 [Emim][DMP](1) + C2H5OH(2) + H2O(3) x2: x3 = 1:4 −4925 −4564 −5608 −5304 −5199 −5075 −3540 −3650 x2: x3 = 2:3 −4309 −3903 −4945 −4588 −4594 −4398 −3168 −3221 x2: x3 = 3:2 −3657 −3303 −4248 −3927 −3936 −3748 −2704 −2710 x2: x3 = 4:1 −2983 −2748 −3519 −3307 −3269 −3146 −2230 −2227

a

Standard uncertainties of temperature T and mole fractions x are as follows: u(T) = 0.05 K, u(x) = 0.0005. The expanded uncertainty with 0.95 level of confidence for the excess enthalpies HE is 10 kJ/kmol.

τij = (gij − gjj)/RT , τji = (gji − gii)/RT

⎛ ϕ ϕ θ z ln γiC = ⎜⎜ln i + qi ln i + li − i xi 2 ϕi ⎝ xi

(8)

where gij, gjj, gji, and gii are interaction parameters between the species i and j and αij, αji (αij = αji) are nonrandom parameters; these parameters can be obtained by fitting experimental excess enthalpy data of binary solutions. UNIQUAC Model. In the UNIQUAC model the activity coefficient of the component i is given as follows.17 C

ln γi = ln γi + ln γi

R

⎛ n ln γi R = qi′⎜⎜1 − ln(∑ θj′τji) − ⎝ j=1

ϕi = (9) D

n

∑ j=1

n

⎞

j=1

⎠

∑ xjlj⎟⎟

⎞ ⎟ n ∑k = 1 θkτkj ⎟⎠

(10)

θj′τij

qixi qi′xi rx i i , θi = , θi′ = ∑j rjxj ∑j qjxj ∑j qj′xj

(11)

(12)

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⎛ Δuij ⎞ ⎛ αij ⎞ τij = exp⎜ − ⎟ ≡ exp⎜ − ⎟ ⎝ T⎠ ⎝ RT ⎠ ⎛ Δuji ⎞ ⎛ αji ⎞ τji = exp⎜ − ⎟ ≡ exp⎜ − ⎟ ⎝ T⎠ ⎝ RT ⎠

(13)

lj =

z (rj − qj) − (rj − 1) 2

(14)

qi′ =

∑ νkiQ k

m k=1

γiC

(15)

γiR

where and are the combinatorial term and residual term of γi, respectively, ri, qi, or qi′(in most cases qi′ = qi, ) are the molar volume and surface parameters of component i, respectively, vki and Qk are the number of the group k in component i and the corresponding surface parameter, respectively. Δuij and Δuji are the energy parameters and are independent of composition and temperature. z is taken as 10. Only qi′ appears in eq 5, as γiC is independent of the temperature. For ionic liquid [Emim][DMP] the qi′can be obtained from the eq 15.18 The Qk of each group in [Emim][DMP] and the qi′ of each component are listed in Table 2.

Figure 3. The experimental and predicted excess enthalpies of [Emim][DMP](1) + CH3OH(2) + H2O(3) at 298.15 K and at normal atmospheric pressure: right pointing triangle, x2 = 0; left pointing triangle, x2:x3 = 1:4; ▽, x2:x3 = 2:3; △, x2:x3 = 3:2; ○, x2:x3 = 4:1; □, x3 = 0; , predicted by NRTL.

4. RESULTS AND DISCUSSION The Binary Systems. The experimental excess enthalpies of the three set of binary solutions [Emim][DMP](1) + H2O(2)/ CH 3 OH(2)/C 2 H 5 OH(2) at 298.15 K and at normal atmospheric pressure are listed in Table 3. Some of the experimental data in Table 3 are taken from our previous article6 while the others are new supplementary data in order to get more precise binary interaction parameters. The excess enthalpy data of binary solutions CH3OH(2)/C2H5OH(2) + H2O(3) can be found in ref 19. The binary interaction parameters in NRTL and UNIQUAC can be obtained by fitting the excess enthalpy data of the five set of binary solutions [Emim] [DMP] (1) + CH3OH(2)/C2H5OH(2)/H2O(3) and CH3OH(2)/C2H5OH(2) + H2O(3), which are listed in Table 4 and Table 5, respectively. The correlated excess enthalpies of these binary solutions are also listed in Table 3. It should be emphasized that in our previous article6 the experimental excess enthalpies of binary solutions [Emim][DMP](1)+CH3OH(2)/C2H5OH(2)/H2O(2) were fitted by Redlich−Kister equations rather than NRTL and UNIQUAC models. Redlich−Kister equations are merely the fitting equations and do not have the ability to predict excess enthalpies of ternary solutions, while the NRTL and UNIQUAC models together with Gibbs−Helmholtz equation have, that is why these two activity coefficient models are employed in this paper. From Table 3 it can be seen that the average relative deviations between the experimental and the correlated excess enthalpies obtained from the NRTL model are 2.49 %, 1.64 %, and 2.44 %, respectively. Meanwhile those obtained from the UNIQUAC model are 2.98 %, 3.00 %, and 3.43 %, respectively. This indicates that the excess enthalpies can be well correlated by NRTL and UNIQUAC models. The excess enthalpies of the binary solutions are all negative and their magnitudes range in the order of [Emim][DMP] + H2O > [Emim][DMP] + CH3OH > [Emim][DMP] + C2H5OH.

Figure 4. The experimental and predicted excess enthalpies of [Emim][DMP](1) + C2H5OH(2) + H2O(3) at 298.15 K and at normal atmospheric pressure: right pointing triangle, x2 = 0; left pointing triangle, x2:x3 = 1:4; ▽, x2:x3 = 2:3; △, x2:x3 = 3:2; ○, x2:x3 = 4:1; □, x3 = 0; , predicted by NRTL.

The Ternary Systems. The excess enthalpies of ternary solutions were calculated or predicted by the NRTL and UNIQUAC models in which the binary interaction parameters were obtained by fitting the excess enthalpy data of the binary solutions. The predicted excess enthalpy together with experimental excess enthalpy data of the two set of ternary solutions [Emim][DMP](1) + CH3OH(2)/C2H5OH(2) + H2O(3) at 298.15 K and at normal atmospheric pressure are listed in Table 6. As shown in Table 6, the average relative deviation between the experimental and predicted excess enthalpies are 7.89 % and 5.17 % (NRTL) or 1.86 % and 1.86 % (UNIQUAC), respectively. The variations of the excess enthalpies of the ternary solutions which have different IL molar fractions are shown in Figures 3 to 6. From these figures it can be seen that the curves HE−x for ternary solutions are within those of the corresponding binary solutions containing [Emim][DMP]. For example, the curves HE−x for the ternary solution [Emim][DMP](1) + CH3OH(2) + H2O(3) lie between those of [Emim][DMP](1) + CH3OH(2) and [Emim][DMP](1) + H2O(3). Meanwhile the E

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experimental excess enthalpies and the predicted ones are 7.89 %, 5.17 % (NRTL) or 1.86 %, and 1.86 % (UNIQUAC), respectively. This indicates that NRTL and UNIQUAC models can well predict excess enthalpies of the ternary solutions which contain IL, while UNIQUAC model exhibits better accuracy than NRTL model in this research.

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ASSOCIATED CONTENT

* Supporting Information S 1

H NMR spectra for [Emim][DMP]. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Tel.: +86 411 84986171. Fax: +86 411 84986171.

Figure 5. The experimental and predicted excess enthalpies for [Emim][DMP](1) + CH3OH(2) + H2O(3) at 298.15 K and at normal atmospheric pressure: right pointing triangle, x2 = 0; left pointing triangle, x2:x3 = 1:4; ▽, x2:x3 = 2:3; △, x2:x3 = 3:2; ○, x2:x3 = 4:1; □, x3 = 0; , predicted by UNIQUAC.

Funding

This work was supported financially by the National Natural Science Foundation of China (51076021, 50876014). Notes

The authors declare no competing financial interest.

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REFERENCES

(1) Kim, K. S.; Park, S. Y.; Choi, S.; Lee, H. Vapor pressure of the 1butyl-3-methylimidazolium bromide + water, 1-butyl-3-methylimidazolium tetrafluoroborate + water, and 1-(2-hydroxyethyl)-3- methylimidazolium tetrafluoroborate + water system. J. Chem. Eng. Data 2004, 49, 1550−1553. (2) Shiflett, M. B.; Yokozeki, A. Solubility and diffusivity of hydrofluorocarbons in room-temperature ionic liquids: [Bmim][PF6] and [Bmim][BF4]. AIChE J. 2006, 52, 1205−1219. (3) Yokozeki, A.; Shiflett, M. B. Vapor−liquid equilibria of ammonia + ionic liquid mixtures. Appl. Energy. 2007, 84, 1258−1273. (4) Yokozeki, A.; Shiflett, M. B. Water solubility in ionic liquids and application to absorption cycles. Ind. Eng. Chem. Res. 2010, 49, 9496− 9503. (5) Wang, J. Z.; Zheng, D. X.; Fan, L. H.; Dong, L. Vapor pressure measurement for water +1,3-dmethylimidazolium chloride system and 2,2,2-trifluoroethanol + 1-ethyl-3-methylimidazolium tetrafluoroborate system. J. Chem. Eng. Data 2010, 55, 2128−2132. (6) Ren, J.; Zhao, Z. C.; Zhang, X. D. Vapor pressures, excess enthalpies, and specific heat capacities of the binary working pairs containing the ionic liquid 1-ethyl-3-methylimidazolium dimethylphosphate. J. Chem. Thermodyn. 2011, 43, 576−583. (7) Zhang, X. D.; Hu, D. Performance simulation of the absorption chiller using water and ionic liquid 1-ethyl-3-methylimidazolium dimethylphosphate as the working pair. Appl. Therm. Eng. 2011, 31, 3316−3321. (8) Zhang, X. D.; Hu, D. Performance analysis of the single-stage absorption heat transformer using a new working pair composed of ionic liquid and water. Appl. Therm. Eng. 2012, 37, 129−135. (9) Kato, R.; Krummen, M.; Gmehling, J. Measurement and correlation of vapor−liquid equilibria and excess enthalpies of binary systems containing ionic liquids and hydrocarbons. Fluid Phase Equilib. 2004, 224, 47−54. (10) Liebert, V.; Nebig, S.; Gmehling, J. Experimental and predicted phase equilibria and excess properties for systems with ionic liquids. Fluid Phase Equilib. 2008, 268, 14−20. (11) Miaja, G. G.; Troncoso, J.; Romaní, L. Excess enthalpy, density, and heat capacity for binary systems of alkylimidazolium-based ionic liquids + water. J. Chem. Thermodyn. 2009, 41, 161−166. (12) Miaja, G. G.; Troncoso, J.; Romaní, L. Excess properties for binary systems ionic liquid + ethanol: Experimental results and theoretical description using the ERAS model. Fluid Phase Equilib. 2008, 274, 59−67.

Figure 6. The experimental and predicted excess enthalpies for [Emim][DMP](1) + C2H5OH(2) + H2O(3) at 298.15 K and at normal atmospheric pressure: right pointing triangle, x2 = 0; left pointing triangle, x2:x3 = 1:4; ▽, x2:x3 = 2:3; △, x2:x3 = 3:2; ○, x2:x3 = 4:1; □, x3 = 0; , predicted by UNIQUAC.

ternary excess enthalpies are all negative and their magnitudes range in the order of [Emim][DMP](1) + CH3OH(2) + H2O(3) > [Emim][DMP](1) + C2H5OH(2) + H2O(3).

5. CONCLUSIONS The excess enthalpies of the three set of binary solutions [Emim][DMP](1) + CH3OH(2)/C2H5OH(2) + H2O(3) are negative, which indicates that the mixing processes of [Emim][DMP] with H2O, CH3OH, or C2H5OH are all exothermic, and the heats of mixing are ranged in the order of [Emim][DMP]+H2O > [Emim][DMP] + CH3OH > [Emim][DMP] + C2H5OH. The average relative deviations between the experimental excess enthalpies and the correlated ones are 2.49 %, 1.64 %, 2.44 % (NRTL) or 2.98 %, 3.00 %, and 3.43 % (UNIQUAC), respectively. The excess enthalpies of the two set of ternary solutions [Emim][DMP](1) + CH3OH(2)/C2H5OH(2) + H2O(3) are negative, the heats of mixing for ternary solutions are ranged in the order of [Emim][DMP] + CH3OH + H2O > [Emim][DMP] + C2H5OH + H2O. The average relative deviation between the F

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