Electrical Conductivity of Aqueous Ethanol Solutions Containing

Article pubs.acs.org/jced

Electrical Conductivity of Aqueous Ethanol Solutions Containing Ammonium Salts under High Pressure at 298 K Kouji Maeda,*,† Katsutoshi Maeno,† Keisuke Fukui,† Masato Moritoki,‡,∥ and Hidetoshi Kuramochi§ †

Department of Mechanical System Engineering, University of Hyogo, 2167 Shosha, Himeji, Hyogo 671-2201, Japan Research Center of High Pressure Technology, 3-11-6 Modorigaokahigashi, Miki, Hyogo 673-0533, Japan § Research Center for Material Cycles and Waste Management, National Institute for Environmental Studies, 16-2 Onogawa, Tsukuba, Ibaraki 305-8506, Japan ‡

ABSTRACT: The molar equivalent electrical conductivity of aqueous ethanol solutions containing ammonium chloride or ammonium nitrate was measured under high pressure, up to 400 MPa, at 298 K. The concentration range of the ammonium salts in solution was broad because their solubility is relatively high compared with that of other salts. A common result was that the molar equivalent electrical conductivity decreased as a function of solute concentration. Additionally, the molar equivalent electrical conductivity decreased considerably when the ethanol content was increased and decreased slightly as the pressure was increased. The molar equivalent electrical conductivity was treated as a function of the square root of electrolyte molality, in accordance with the Debye−Hückel− Onsager equation, but could not be represented by one function because of dependence on ethanol concentration and on pressure. Therefore, use of the electrolyte−nonrandom two-liquid (NRTL) model is proposed for extending the Debye−Hückel−Onsager equation to these two electrolyte solution systems.

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Cruz and Renon9 first considered the nonrandom two-liquid (NRTL) model combined with the Debye−Hückel model for industrial applications. Since the 1980s, Chen and Evans10 have also applied the electrolyte−NRTL model to electrolyte solutions. Anderko and Lencka11 have proposed the unrestricted primitive model of the mean spherical approximation (MSA) to represent electrical conductivity in a relatively wide range of concentrations. Vila et al.12 have also reported another ionic friction model to express electrical conductivity across a large concentration range. Usobiaga et al.13 have measured the electrical conductivity of mixed electrolytes at high concentrations and provided an empirical model to fit the experimental data. However, the electrical conductivity of an electrolyte in mixed solvents and of multiple electrolytes is still not well understood. The electrolyte−NRTL model is useful in industrial applications because of the limited ionic species used in electrolyte solutions. We have correlated the activity and solubility of electrolytes in aqueous solution by using ion-specific NRTL parameters14,15 and reported solubility data.6,16 To develop new electrolyte materials and find good operative conditions, the electrical conductivity of electrolytes under extreme, nonideal conditions, such as polymer solutions and solutions under high pressure or high temperature, should be studied.

INTRODUCTION Industrial processes, such as precipitation, crystallization, and electrophoresis for inorganic substances, use nondilute electrolyte solutions. For example, electrolyte solutions used in the battery industry must have high electrical conductivity to improve charge− discharge properties as well as to prevent dissolution of the electrode. However, most studies on the electrical conductivity of electrolyte solutions have been carried out under dilute conditions. Ho et al.1,2 have studied the limiting molar equivalent electrical conductivity of dilute aqueous solutes under extreme conditions and provided a function of the association constant with a good fit to electrical conductivity data for dilute solutes. Ueno et al.3 have also investigated ion−solvent interaction at high pressure through electrical conductivity measurements of dilute electrolytes. Recently, they have examined the Hubbard−Onsager theory in regard to molecular transport processes by performing the measurements using highly condensed dilute solutions under high pressure.4 For dilute solutions of ideally dissociated electrolytes, in which the ionic strength and concentration are equivalent, the Debye−Hückel theory can represent the activity coefficient or electrical conductivity as a linear function of the square root of molarity.5 Robinson and Stokes6 explained the theory of electrolyte solutions and provided the activity data for many electrolytes. However, the electrical conductivity of concentrated or partially dissociated electrolytes cannot be represented as a linear function. Pitzer7,8 developed a thermodynamic equation to determine activity coefficients of electrolyte solutions under conditions where the Debye−Hückel model is inadequate. In 1978, © 2013 American Chemical Society

Received: July 24, 2012 Accepted: December 6, 2012 Published: January 17, 2013 264

dx.doi.org/10.1021/je3008233 | J. Chem. Eng. Data 2013, 58, 264−270

Journal of Chemical & Engineering Data

Article

Figure 1. High pressure apparatus used to measure the electrical conductivities of the solutions. 1, high-pressure cell; 2, pressure promoter; 3, 4-way valves; 4, plunger pump; 5, water; 6, solution cell for electrical conductivity; 7, electrical conductivity meter; 8, computer; PL, pressure gauge (low); PH, pressure gauge (high); VH1,2,3, high-pressure valves; VL1,2, low-pressure valves; MT, thermometer.

High pressure is the operative variable that achieves the fastest and most uniform change, as compared with concentration and temperature. However, the properties of condensed materials are not greatly altered through the application of pressure because the liquid molar volume does not respond to the increased pressured. Extremely high pressures, greater than 1000 bar, can alter some properties of electrolyte solutions. For example, the solubility of electrolytes may increase or decrease as pressure is increased.17 The objectives of this study are to measure the electrical conductivity of two ammonium salts in aqueous ethanol solutions using a wide range of concentrations and to understand the effects of ethanol content and high pressure on the electrical conductivity of the solution. There are few salts which have a relatively large solubility in alcoholic solvent, but ammonium salts have a larger solubility than the other salts in alcohol. The electrolyte solution such as aqueous alcohol solution dissolving a high amount of ammonium salts could be useful for developing new electrolyte solutions for battery industries using high pressure. We consider the relation between the molar equivalent electrical conductivity and the activity of electrolyte solutions by using a combination of the electrolyte−NRTL model and the Debye−Hü ckel−Onsager equation for the aqueous ethanol solutions of a relatively large concentration of ammonium salts.

The aqueous solution, with a specific concentration of an electrolyte (NH4Cl or NH4NO3), was packed into the electrical conductivity cell (2−3 mL) by using a free piston cap at atmospheric pressure. The aqueous solution was completely separated from the pressurizing water. Two electrodes, made of titanium rods with porous platinum plating, extended from the pressurized cell into atmospheric pressure through a special seal. The aqueous solution in the electrical cell can be isolated at 400 MPa in the absence of the two electrodes. The electrical cell was then placed inside the high-pressure cell. The electrical conductivity of the solution in the cell was measured with a conductivity meter (DS-52, Horiba Co.). For each sample, the cell constant was calibrated against 0.1 mol·kg−1 KCl aqueous solution. The electrical conductivity of each solution was statically measured twice at (0.1, 100, 200, 300, and 400) MPa. The measurements were taken with both increasing and decreasing pressure to account for any hysteresis. The electrical conductivity of the solutions was found to have a constant value at a given pressure. The solutions used to dissolve the ammonium salts were aqueous ethanol, in which ammonium salts are highly soluble. However, the concentration of the ammonium salts [m (mol·kg−1)] should be prepared at less than the saturated concentration at atmospheric pressure. The mole fractions of ethanol used in the aqueous solutions (xe) were 0 to 0.78 in salt-free solution.

EXPERIMENTAL METHODS Materials. Ammonium chloride (NH4Cl) and ammonium nitrate (NH4NO3), used as electrolytes in aqueous ethanol solutions, were purchased from Kanto Chemicals Co., and their purities certified as being greater than 99 mass %. Ethanol was used as an antisolvent and was also purchased from Kanto Chemicals Co. Its purity was certified as more than 99 mol %. The aqueous ethanol solutions were prepared by mixing ethanol with distilled water. High-Pressure Apparatus. Figure 1 shows the experimental setup used to measure the electrical conductivity of aqueous solutions at high pressure. The high-pressure cell has four sapphire windows at right angles. The high-pressure cell is rectangular and has sapphire windows on four of its sides, excluding the top and bottom. Water is used as the pressurizing fluid. The plunger pump supplies water pressurized at 20 MPa directly to the high-pressure cell. Adjusting a six-way valve allows the plunger pump to supply the pressurized water to the pressure promoter, which brings the high-pressure cell to 400 MPa. The temperature in the high-pressure cell was monitored during the experiment.

RESULTS AND DISCUSSION The electrical conductivity of aqueous ethanol solutions was obtained as the molar equivalent electrical conductivity Λ (S·cm2·mol−1) in this study. The molar equivalent electrical conductivities of aqueous ethanol solution containing NH4Cl are listed in Table 1, and those for solutions containing NH4NO3 are listed in Table 2. The values are listed as a function of electrolyte molality at different ethanol mole fractions and at different pressures. Here molality was used for the electrolyte’s concentration because the volume of solution at high pressure may change from that at atmospheric pressure. The error in the conductivity and in the pressure was less than 5 %. It was determined that the molality of the electrolytes had an error of less than 0.1 %. Figure 2 shows the molar equivalent electrical conductivity of electrolyte solutions under a range of high pressures as a function of the square root of the molality of NH4Cl (see Table 1 for values). Figure 3 shows the results when NH4NO3 was used as the electrolyte (see Table 2). Generally, the electrolyte solubility decreased as the ethanol content was increased. Therefore, the lower electrolyte molality was used at higher ethanol content (Figures 2 and 3). In addition,

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Table 1. Molar Equivalent Electrical Conductivity of the Aqueous Ethanol Solutions Containing NH4Cl under High Pressure at 298 Ka mole fraction of ethanol in solution, xe 0

0.11

0.28

a

temperature, T/K

pressure P/MPa

molality of electrolyte, m/mol·kg−1

298

0.1 0.1 0.1 100 100 100 200 200 200 300 300 300 400 400 400 0.1 0.1 0.1 100 100 100 200 200 200 300 300 300 400 400 400 0.1 0.1 0.1 100 100 100 200 200 200 300 300 300 400

0.500 1.000 1.500 0.500 1.000 1.500 0.500 1.000 1.500 0.500 1.000 1.500 0.500 1.000 1.500 0.701 1.402 2.104 0.701 1.402 2.104 0.701 1.402 2.104 0.701 1.402 2.104 0.701 1.402 2.104 0.460 0.920 1.380 0.460 0.920 1.380 0.460 0.920 1.380 0.460 0.920 1.380 0.460

298

298

molar equivalent electrical conductivity, Λ/S·cm2·mol−1 114.7 109.1 106.6 117.2 111.6 106.2 118.1 111.8 105.6 117.1 107.1 105.1 116.8 109.8 103.1 96.3 89.3 84.9 95.6 89.1 84.0 95.6 88.6 83.6 95.0 88.0 83.4 93.5 87.5 82.7 86.5 80.2 77.2 85.9 78.6 76.0 84.7 77.3 73.3 83.2 76.4 71.9 81.9

mole fraction of ethanol in solution, xe

temperature, T/K

0.48

298

0.78

298

pressure P/MPa


molar equivalent electrical conductivity, Λ/S·cm2·mol−1

400 400 0.1 0.1 0.1 0.1 100 100 100 100 200 200 200 200 300 300 300 300 400 400 400 400 0.1 0.1 0.1 0.1 100 100 100 100 200 200 200 200 300 300 300 300 400 400 400 400

0.920 1.380 0.230 0.459 0.689 0.918 0.230 0.459 0.689 0.918 0.230 0.459 0.689 0.918 0.230 0.459 0.689 0.918 0.230 0.459 0.689 0.918 0.049 0.097 0.146 0.194 0.049 0.097 0.146 0.194 0.049 0.097 0.146 0.194 0.049 0.097 0.146 0.194 0.049 0.097 0.146 0.194

75.6 71.1 82.3 76.5 72.5 65.5 79.9 74.7 70.6 64.4 77.6 72.5 68.9 63.1 75.1 70.3 67.0 61.8 72.2 68.1 65.0 60.9 79.0 74.2 73.2 70.8 75.5 71.7 70.9 68.7 72.2 67.5 67.5 66.4 69.9 66.6 66.1 64.2 67.4 64.2 63.9 62.2

Standard uncertainties u are u(T) = 0.1 K, u(P) = 0.01 MPa for P < 1 MPa, u(P) = 1 MPa for P > 100 MPa, u(m) = 0.001 mol·kg−1, u(xe) = 0.01, and u(Λ) = 1 S·cm2·mol−1 (0.95 level of confidence).

electrolytes in the aqueous ethanol solutions was found. When the pressure increased, the molar equivalent electrical conductivity essentially decreased. When the ethanol mole fraction is small in the solution, the decrease of the molar equivalent electrical conductivity by increasing of pressure was not considerable. When the concentration of electrolyte becomes large in the solutions especially containing much ethanol, the decrease of the molar equivalent electrical conductivity by increasing of pressure was significant even though the molality of electrolytes could not be increased more than the molality in less ethanol solutions.

the molality of NH4Cl was set to be about half the molality of NH4NO3 because its solubility is less than that of NH4NO3. As for the molar equivalent electrical conductivities for the two electrolytes in pure water, our data were similar to the literature data,18 and an effect of high pressure was not clearly observed. The molar equivalent electrical conductivity of electrolytes decreased considerably, and the effect of high pressure became stronger, as the ethanol mole fraction was decreased. As the pressure was increased, the electrical conductivity of the electrolyte solutions decreased. A function to describe the linearity of molar equivalent electrical conductivity of 266



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Table 2. Molar Equivalent Electrical Conductivity of Aqueous Ethanol Solutions Containing Ammonium Nitrate under High Pressure at 298 Ka mole fraction of ethanol in solution, xe 0

0.11

0.28

a

temperature, T/K

pressure P/MPa


298

0.1 0.1 0.1 100 100 100 200 200 200 300 300 300 400 400 400 0.1 0.1 0.1 100 100 100 200 200 200 300 300 300 400 400 400 0.1 0.1 0.1 100 100 100 200 200 200 300

0.500 1.000 1.500 0.500 1.000 1.500 0.500 1.000 1.500 0.500 1.000 1.500 0.500 1.000 1.500 1.429 2.858 4.287 1.429 2.858 4.287 1.429 2.858 4.287 1.429 2.858 4.287 1.429 2.858 4.287 1.137 2.275 3.412 1.137 2.275 3.412 1.137 2.275 3.412 1.137

298

298

molar equivalent electrical conductivity, Λ/S·cm2·mol−1 113.5 106.5 102.5 114.4 106.7 105.1 114.1 105.5 106.0 113.1 104.5 105.3 111.6 103.3 104.3 92.3 87.4 83.7 91.4 86.5 82.6 90.1 85.4 81.4 88.9 84.2 80.2 87.0 82.8 78.9 82.4 77.5 72.6 81.3 76.6 71.0 79.0 74.8 69.1 77.5

mole fraction of ethanol in solution, xe

temperature, T/K

0.48

298

0.78

298

pressure P/MPa


molar equivalent electrical conductivity, Λ/S·cm2·mol−1

300 300 400 400 400 0.1 0.1 0.1 100 100 100 200 200 200 300 300 300 400 400 400 0.1 0.1 0.1 0.1 100 100 100 100 200 200 200 200 300 300 300 300 400 400 400 400

2.275 3.412 1.137 2.275 3.412 0.837 1.675 2.512 0.837 1.675 2.512 0.837 1.675 2.512 0.837 1.675 2.512 0.837 1.675 2.512 0.235 0.470 0.705 0.939 0.235 0.470 0.705 0.939 0.235 0.470 0.705 0.939 0.235 0.470 0.705 0.939 0.235 0.470 0.705 0.939

72.7 67.4 75.4 71.2 66.0 77.4 66.7 60.2 75.1 65.0 57.3 72.6 62.9 56.5 70.2 61.5 55.3 68.7 60.6 54.4 70.1 69.0 66.0 63.8 67.8 66.1 64.2 62.1 65.1 63.8 62.2 59.9 62.6 61.7 60.1 58.2 60.3 59.7 58.4 56.5

Standard uncertainties u are u(T) = 0.1 K, u(P) = 0.01 MPa for P < 1 MPa, u(P) = 1 MPa for P > 100 MPa, u(m) = 0.001 mol·kg−1, u(xe) = 0.01, and u(Λ) = 1 S·cm2·mol−1 (0.95 level of confidence).

Ion mobility influences electrical conductivity.4,12 The viscosity of the solution can affect ion mobility. The kinetic model such as Hubbard−Onsager theory can well represent the electrical conductivity of extremely dilute solutions.4 Theoretical studies on the effects of molecular behavior on electrical conductivity can be expected to improve the usefulness of the Debye− Hückel−Onsager theory when applied to many different dilute electrolyte solutions. The Debye−Hückel−Onsager theory posits a linear relationship between molar equivalent electrical conductivity and the square root of electrolyte molality in a dilute solution. In other words, an ideal ionic species has a molar equivalent electrical conductivity that is a linear function of the square root of its molality. If the activity of the electrolyte is known, the molar equivalent electrical conductivity is a linear

function of the square root of the activity of any salts, even if it does not behave ideally like highly concentrated electrolytes in the solution. Under ideal conditions, the molar equivalent electrical conductivity of a solution of strong electrolyte is relatively large and is a linear function of the square root of electrolyte molality, that of a solution of moderately strong electrolyte is relatively small and is a nonlinear function of the square root of its molality, and that of a solution of weak electrolyte shows no linear correlation to the square root of its molality. However, if the activity is known, most molar equivalent electrical conductivities should be expressible as a linear function of the square root of its activity. Activity represents actual ionic behavior, which is different from the molality of the ions. Therefore, the activity should account for any nonideal behavior of ions in 267



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represented by the combination of the Pitzer−Debye−Hückel equation and the NRTL equation used for industrial applications. ln γi* = ln γi*,pdh + ln γi*,NRTL

(4)

We have extended the description of the activity coefficient in the electrolyte−NRTL model by using ion-specific parameters. The solubility of several complicated electrolyte systems in aqueous solution was represented by using these parameters. We have determined the ion-specific NRTL parameters for ammonium salts in aqueous ethanol solutions (Table 3). The molar equivalent Table 3. Electrolyte−NRTL Parameters (Δgij) for Aqueous Ethanol Solutions Containing Ammonium Salts

Figure 2. Molar equivalent electrical conductivity of different aqueous ethanol electrolyte solutions at high pressure as a function of the square root of NH4Cl molality: ● 0.1 MPa; ■, 100 MPa; ▲, 200 MPa; ○, 300 MPa; □, 400 MPa; , calculated values as same order as pressure; − −, reference data18 (ICT, 1949).

water ethanol NH4+ Cl− NO3−

water

ethanol

NH4+

Cl−

NO3−

0 −527 2820 −9280 −13300

5427 0 5290 −8220 −6700

−22500 −20000 0 −23300 −41500

−13200 −3500 −27800 0 0

−1540 −10900 −19800 0 0

electrical conductivity, Λ, can be expressed by a modified version of the Debye−Hückel−Onsager equation as a function of the activity, mγ±,ac * . Λ = Λ 0 − S mγ±*,ac

(5)

Λ0 and S are the limiting electrical conductivity and the Onsager limiting slope, respectively. γ*±,ac is calculated by using the electrolyte−NRTL model with parameters14 from Table 3. The parameters of the electrolyte−NRTL model were newly obtained by fitting the solubility of ammonium salts in aqueous ethanol solution as shown in Figure 4. Λ0 and S could be a function

Figure 3. Molar equivalent electrical conductivity of different aqueous ethanol electrolyte solutions at high pressure as a function of the square root of NH4NO3 molality: ●, 0.1 MPa; ■, 100 MPa; ▲, 200 MPa; ○, 300 MPa; □, 400 MPa; , calculated values as same order as pressure; − −, reference data18 (ICT, 1949).

solution even if the electrolytes are highly concentrated or dissolved in a mixed solution at high pressure. To calculate the mean activity of an electrolyte, the mole fraction of the ion species should be used and can be derived from the molality of the salt (m) as shown by the following equations, assuming complete dissociation. υ+m x+ = 1000 + m (υ+ + υ−) M s

x− =

υ−m 1000 Ms

+ m(υ+ + υ−)

Figure 4. Solubility of NH4NO3 and NH4Cl molality in aqueous ethanol solution at 0.1 MPa, 298 K. ●, NH4NO3; , electrolyte NRTL model; ○, NH4Cl; − −, electrolyte NRTL model.

(1)

Since the activity coefficient of an individual ion i (γi*) cannot be measured, the activity coefficients of aqueous ions are usually expressed as the mean ionic activity coefficient of a neutral electrolyte. The mean activity coefficient (γ±,ac * ) is given by ln γ±*,ac =

of ethanol mole fraction at different pressures, as shown in Figures 5 and 6. In this study, the values of these parameters Λ0 and S were empirically determined using the mole fraction of ethanol in solution and the pressure for two different electrolytes. The average deviation of molar equivalent electrical conductivity for each aqueous ethanol solution for two ammonium salts was listed in Table 4. Λ0 and S might depend on the viscosity and the density of aqueous ethanol solutions because the viscosity of aqueous alcohol solutions reported by Tanaka et al.19 show the similar shape as Figures 5 and 6. In future work, we will measure the density and viscosity of aqueous ethanol solutions at high pressure to account for any effects on Λ0 and S. However, these empirical correlations by Debye−Hückel−Onsager equation

νa ln γa* + νc ln γc* νac

(2)

where νa is the stoichiometric coefficient of the anion, that is, the number of moles of the anion produced during dissociation of one mole of the salt (ac); νc denotes that of the cation, and νac is νac = νa + νc (3) In the electrolyte−NRTL model proposed by Chen et al.,9,10 the activity coefficient of an ion, i, in an aqueous solution is 268



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agreed with the experimentally determined molar equivalent electrical conductivities of NH4Cl and NH4NO3 in aqueous ethanol solutions at high pressure.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes ∥

Retired from Kobe Steel Co. The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We are grateful to Professor Masakatsu Ueno from Doshisha University for advice on electrical conductivity measurement in a high-pressure system, as well as for his comments regarding the experimental and theoretical aspects of the research.

Figure 5. Limiting electrical conductivity in the Debye−Hückel− Onsager equation for aqueous ethanol electrolyte solutions as a function of ethanol mole fraction at three different pressures. , 0.1 MPa; − −, 100 MPa; ···, 200 MPa.

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(1) Ho, P. C.; Palmer, D. A.; Wood, R. H. Conductivity Measurements of Dilute Aqueous LiOH, NaOH, and KOH Solutions to High Temperatures and Pressures Using a Flow-Through Cell. J. Phys. Chem. B 2000, 104, 12084−12089. (2) Ho, P. C.; Palmer, D. A.; Gruszkiewicz, M. S. Conductivity Measurements of Dilute Aqueous HCl Solutions to High Temperatures and Pressures Using a Flow-Through Cell. J. Phys. Chem. B 2001, 105, 1260−1266. (3) Ueno, M.; Shimizu, K.; Osugi, J. The Electrical Conductivity of [Co(NH3)5NO2]SO4 in Aqueous Solution under High Pressure. Rev. Phys. Chem. Jpn. 1973, 43, 33−43. (4) Hoshina, T.; Tsuchihashi, N.; Ibuki, K.; Ueno, M. Electric Conductivities of 1:1 Electrolytes in Liquid Methanol along the Liquid-Vapor coexistence curve up to the Critical Temperature. I NaCl, and CsCl solutions. J. Chem. Phys. 2004, 120, 4355−4365. (5) Pitzer, K. S.; Brewer, L. Thermodynamics, 2nd revised ed.; Lewis, G. N., Randall, M., Eds.; McGraw-Hill: New York, 1961. (6) Robinson, R. A.; Stokes, R. H. Electrolyte Solutions; Butterworth: London, 1965. (7) Pitzer, K. S. Thermodynamics of Electrolytes. I. Theoretical Basis and General Equations. J. Phys. Chem. 1973, 77, 268−277. (8) Pitzer, K. S.; Mayorga, G. Thermodynamics of Electrolytes. II. Activity Coefficients for Strong Electrolytes with One and Both Ions Univalent. J. Phys. Chem. 1973, 77, 2300−2308. (9) Cruz, J.-L.; Renon, H. A New Thermodynamic Representation of Binary Electrolyte Solutions Nonideality in the Whole Range of Concentration. AIChE J. 1978, 24, 817−830. (10) Chen, C.-C.; Evans, L. B. A Local Composition Model for the Excess Gibbs Energy of Aqueous Electrolyte Systems. AIChE J. 1986, 32, 444−454. (11) Anderko, A.; Lencka, M. M. Computation of Electrical Conductivity of Multicomponent Aqueous Systems in Wide Concentration and Temperature Ranges. Ind. Eng. Chem. Res. 1997, 36, 1932−1943. (12) Vila, J.; Rilo, E.; Segade, L.; Cabeza, O.; Varela, L. M. Electrical Conductivity of Aqueous Solutions of Aluminum Salts. Phys. Rev. E 2005, 71, 031201. (13) Usobiaga, A.; Diego, A.; Madariaga, J. M. Electrical Conductivity of Concentrated Aqueous Mixtures of HCl and KCl in a Wide Range of Compositions and Temperatures. J. Chem. Eng. Data 2000, 45, 23−28. (14) Kuramochi, H.; Osako, M.; Kida, A.; Nishimura, K.; Kawamoto, K.; Asakuma, Y.; Fukui, K.; Maeda, K. Determination of Ion-Specific NRTL Parameters for Predicting Phase Equilibria in Aqueous Multielectrolyte Solutions. Ind. Eng. Chem. Res. 2005, 44, 3289−3297. (15) Maeda, K.; Safaeefar, P.; Ang, H.-M.; Kuramochi, H.; Asakuma, Y.; Tade, M. O.; Fukui, K. Prediction of Solid-Liquid Phase Equilibrium in the System of Water (1) + Alcohols (2) + MgSO4·7H2O (3) + MnSO4·H2O (4) by the Ion-Specific Electrolyte NRTL Model. J. Chem. Eng. Data 2009, 54, 423−427.

Figure 6. Onsager limiting slope for the Debye−Hückel−Onsager equation for aqueous ethanol electrolyte solutions as a function of ethanol mole fraction at three different pressures. , 0.1 MPa; − −, 100 MPa; ···, 200 MP.

Table 4. Average R2 of Molar Equivalent Electrical Conductivity of Aqueous Ethanol Solutions Containing Two Ammonium Saltsa mole fraction of ethanol in solution, xe

R2 for NH4Cl salt

R2 for NH4NO3 salt

0 0.11 0.28 0.48 0.78

11 0.6 1.8 1.9 3.2

3.2 1.2 1.8 6.4 3.7

REFERENCES

a Λ0 = 129.5 − 241.2xe + 534.5xe2 − 372.3xe3 − 0.0174Pxe; S = 31.7 − 164.8xe + 535.0xe2 − 355xe3 + 0.0439Pxe.

with electrolyte−NRTL model could consistently represent the molar equivalent electrical conductivity of mixed solvent solutions with high concentrations of NH4Cl or NH4NO3 at high pressure, as shown in Figures 2 and 3.

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CONCLUSION The molar equivalent electrical conductivity of ethanol solutions containing a concentrated ammonium salt (NH4Cl or NH4NO3) was measured under extremely high pressures at 298 K. As the electrolyte molality was decreased, the molar electrical conductivity also decreased. The presence of ethanol greatly affected the electrical conductivity of the solutions: the conductivity decreased as the ethanol content was increased. However, exposing the electrolyte solutions to high pressure had a minimal effect on electrical conductivity. The Debye−Hückel− Onsager equation, modified with the electrolyte−NRTL model, 269



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(16) Mullin, J. M. Crystallization, 3rd ed.; Butterworth-Heinemann: Oxford, 1993. (17) Moritoki, M.; Kitagawa, K.; Onoe, K.; Kaneko, K. Industrial Crystallization; Elsevier: Amsterdam, 1984; p 377. (18) International Critical Tables, VI; McGraw-Hill: New York, 1949; pp 231−255. (19) Tanaka, Y.; Matsuda, Y.; Fujiwara, H.; Kubota, H.; Makita, T. Viscosity of (Water + Alcohol) Mixtures under High Pressure. Int. J. Thermophys. 1987, 8, 147−163.

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Electrical Conductivity of Aqueous Ethanol Solutions Containing

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