Determination of Solubility of Sodium Salts in Aqueous Surfactant

Article pubs.acs.org/jced

Determination of Solubility of Sodium Salts in Aqueous Surfactant and STPP Solutions Using an Ion Selective Electrode Xin Wang,† Ya-Na Zhang,†,‡ and Yang-Xin Yu*,†,§ †

Laboratory of Chemical Engineering Thermodynamics, Department of Chemical Engineering, Tsinghua University, Beijing 100084, P. R. China ‡ College of Chemical Engineering, Beijing University of Chemical Technology, Beijing 100029, P. R. China § State Key Laboratory of Chemical Engineering, Tsinghua University, Beijing 100084, P. R. China S Supporting Information *

ABSTRACT: Solubility of inorganic salts in the mixed systems containing surfactants and sodium tripolyphosphate (STPP) is very important for the production of laundry detergents. In this work we explored the solubility of NaCl, Na2SO4, and Na2CO3 in aqueous surfactant and STPP solutions using an ion selective electrode at 298.15 K. In the experiments, electromotive forces (EMF) for the ternary systems were measured using an ion-meter with a sodium ionselective composite electrode and the relationship between the solubility and EMF was analyzed by a potentiometric analysis. The proposed method has advantages in avoiding the solid− liquid separation operation and solution supersaturation. The experimental results indicated that the salting-out effect caused by surfactant and STPP can decrease the solubility of sodium salts in the aqueous systems.

1. INTRODUCTION Laundry detergent, or washing powder, is one of the most common detergents added for cleaning laundry. It generally comprises surfactant, builder, bleaching agent, filler, and so forth. The main decontamination role in detergent is surfactants. Surfactants are a class of compounds that lower the surface tension between two liquids or between a liquid and a solid.1 To increase the efficiency of the surfactant, builders are added to laundry detergent. Builders can remove calcium ions through chelation or precipitation. Typical builders are sodium tripolyphosphate (STPP), sodium carbonate, sequenstering agents, soap, and zeolites. The laundry detergent manufacturing process in industry involves a complicated liquid−solid transformation during the drying process. The solid−liquid phase ratio directly affects the size of detergent particle in the product. Thus, research on the solid−liquid phase equilibria of surfactant-sodium salt-water and STPP−sodium salt−water ternary systems is of great significance for production of detergent, and can be applied to control the index of solid−liquid phase ratio. The phase diagrams of solid−liquid ternary systems have been extensively constructed by many methods, such as equilibrium method,2,3 kinetic method,4,5 synthetic method,6 wet residue method,7,8 and potentiometric method.9,10 Although the traditional methods fully describe the phase diagrams, they still have some disadvantages. For example, they take hours to achieve the solid−liquid equilibrium (SLE), need to analyze the solid phase by X-ray diffraction, and can hardly avoid the influences © XXXX American Chemical Society

due to the solid−liquid separation process and temperature fluctuations. Ion selective electrodes (ISEs) are known as a sensor to covert activities of ions into electrical potentials.11 ISEs come in a variety of shapes and sizes. Nevertheless, ISEs have certain common features. All the ISEs contain a cylindrical plastic tube. One end is fixed by an ion-selective membrane making the exterior solution only contact with the outer surface, and the other end is equipped with a gold plated pin to connect to the millivolt measuring device. Four main types of ion-selective membrane used in ISEs are solid state, liquid-based, glass, and compound electrode. ISEs are used in a wide variety of applications for the measurement of the ionic concentration in the tested solution. The ionic concentration is associated with the electrode potential by the Nernst equation. Some of the main areas that ISEs have been used are the pollution monitoring, agriculture, food processing, detergent manufacture, paper manufacture, explosives, electroplating, and so on.12−15 The most widely used and simplest method of using ISEs is direct potentiometry. The electrode response in solution is simply measured and the concentration is read directly from the meter display on a self-calibrating ion meter. This method has advantages that it is relatively inexpensive and simple to measure different ranges of concentrations without changing range, recalibrating or making any complicated calculations. Received: November 1, 2015 Accepted: May 25, 2016

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DOI: 10.1021/acs.jced.5b00926 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

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Table 1. Chemical Substances Employed in This Work chemical name

molecular formula

CAS number

purity (% mass)

7647-14-5

99.5

sodium chloride

NaCl

sodium carbonate sodium sulfate sodium dodecyl benzenesulfonatea sodium dodecyl sulfate sodium tripolyphosphate

Na2CO3 Na2SO4 C18H29SO3Na

497-19-8 7757-82-6 25155-30-0

99.8 99.0 95.0

C12H25SO4Na Na5P3O10

151-21-3 13573-18-7

99.0 95.66

a

purification method

source Beijing Modern Oriental Technology Development Company Beijing Chemical Reagents Company Beijing Chemical Reagents Company Beijing Modern Oriental Technology Development Company Tianjin Fuchen Chenical Reagents Factory Procter & Gamble

none none none none none none

A mixture of various isomers.

In this work, we used the ISE method to measure the solubility of sodium salts in aqueous surfactant or STPP systems. In addition, the relationship between the solubility and electromotive force (EMF) was analyzed by the potentiometric analysis. First, the proposed method was used to measure the solubility data of single salts in aqueous systems at 298.15 K and 303.15 K and compared them with the literature reported data in order to verify the accuracy of the method. Then the method was used to determine solubility data of sodium salts in the surfactant + sodium salt + H2O and STPP + sodium salt + H2O ternary systems. The surfactants used in the experiments are sodium dodecyl benzenesulfonate (SDBS) and sodium dodecyl sulfate (SDS). The sodium salts are sodium sulfate, sodium carbonate, and sodium chloride. This paper is divided into three parts. In the first part the procedures of the experiments were described. In the second part, the solubility data of the sodium salt in water were measured to verify the reliability of the method. Finally, the method was used to determine the solubility data of sodium salts in the surfactant− sodium salt−water and STPP + sodium salt + water ternary systems. In the experiments, we found that the salting-out effect made by the surfactant and STPP decreases the solubility of sodium salts and the influencing strength varies with both types of surfactants and sodium salts.

Figure 1. Molecular structures of STPP and surfactants used in our experiments.

2. EXPERIMENTAL SECTION 2.1. Materials. Table 1 lists the materials used in this work along with the molecular formula, CAS number, purity, and source. All reagents were purchased from commercial sources and used without any further purification. In addition, SDBS is a technical product being a mixture of various isomers with different degrees of branching. STPP is also a technical product, including 95.66% STPP, 2.99% pyrophosphate, 0.46% orthophosphate, and 0.4% metaphosphate. Molecular structures of STPP and surfactant used in experiments are shown in Figure 1. 2.2. Measurements. The solubility of sodium salts in the aqueous surfactant and STPP solutions were measured using a sodium ion-selective composite electrode at 298.15 K. The EMFs of the ternary system were measured directly using an Orion 720A ion-meter with a 701 sodium ion-selective composite electrode at atmospheric pressure. The pure component masses were weighed using an Ohaus digital balance. The uncertainties in the measurements of the EMFs and masses were 0.1 mV and 0.1 mg, respectively. 2.3. Solubility of the Single Salt System. The solubility of each sodium salt was measured as follows, and the equipment for solid−liquid equilibration and EMF measurements was shown in Figure 2. The procedure consists of five steps: (1) 200 g deionized water was injected into a constant

Figure 2. Experimental apparatus: (1) thermostatic bath, (2) magnetic stirrer, (3) constant temperature water jacket, (4) sodium ion-selective composite electrode, and (5) Orion 720A ion-meter.

temperature water jacket in a thermostatic bath, in which the temperature was controlled; (2) 5 g salt was added into the water jacket at every turn and stirred. The result was recorded until the EMF value became constant, which indicated that dissolution equilibrium was achieved. It took 10 to 20 min for reaching steady state condition before recording the EMF value; (3) step 2 was continuously repeated until the EMF reached a certain fixed value; (4) the EMF−logarithmic molality (ln msalt) diagram was plotted according to the measured data; (5) the solubility can be obtained from the EMF−ln msalt diagram (the detail was shown in results and discussion). Because the complicated ion-selective electrode was used in this work and the EMF is only an indicator for the saturated point of the solution, no effect of “liquid junction” B



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potential between reference electrode and solution was considered. And for the same reason, there is a one-to-one correspondence between a value of EMF and the concentration of the added salt. The activity coefficient has no effect on the measurements, though it is a function of ionic strength. 2.4. Solubility of Sodium Salts in Aqueous Surfactant or STPP Systems. The measurement procedures are similar to those described in 2.3 and the only difference is in step 1. In the measurement of solubility data in aqueous surfactant or STPP systems, 200 g aqueous surfactant or STPP solution was injected into the constant temperature water jacket.

3. RESULTS AND DISCUSSION The ISE method was used to measure the solubility of sodium salt in aqueous ternary systems in our research. The EMF− ln msalt diagram of the single salt system and sodium salt + surfactant + H2O ternary system can be obtained. In the EMF− logarithmic molality diagram, the horizontal axis was used in the logarithmic molality. The solubility was obtained and discussed in detail as below. The relationship between the electrode potential and the ionic concentration (activity) is presented by the Nernst equation E = E0 +

⎛ RT ⎞ ⎜ ⎟ln a ⎝ nF ⎠

Figure 3. EMF of solution as a function of logarithmic molality of Na2SO4 in water at 303.15 K. Filled squares refer to the experimental measured data and the lines represent linearly fitted values.

(1)

where E is the measured EMF at a certain temperature, E0 is the standard EMF, R is the idea gas constant, T is the absolute temperature, n is the electron transfer number in the electrode reaction, F is the Faraday constant, and a is the activity. From eq 1, one can figure out that there is a linear relationship between the measured EMF (E) and logarithmic activity (ln a). The concentration is in proportion to the activity. Before the solution reached saturation, the added salt weight can be considered approximately in proportion to the concentration. Therefore, there is a linear relationship between the measured EMF (E) and ln msalt. EMF−ln msalt curves are plotted to analyze experimental results. The experiment data show that when the solution is close to saturation, there appears a significant inflection point in the diagram. This phenomenon is because the concentrations of ions in the solution increase with the addition of salts into the solution and when the solution is close to the saturation, the concentrations of ions are almost remain stable, which leads to the unchanged value of EMF as shown in Figures 3−7 and S1−S11. The data of the EMFs in Figures 3 to 7 are also tabulated in Tables S1−S5. Then, we fitted the points on the left side of the inflection point and determined that the x-coordinate value of the intersection point is the solubility. In order to verify the accuracy of the ISE method, we first measure the solubility of single salts in aqueous system at temperatures T = 298.15 K and 303.15 K. We have determined the composition of the solid phases by measuring the weight loss on heating and found that sodium sulfate and carbonate form Na2SO4·10H2O and Na2CO3·10H2O crystal hydrates in equilibrium with aqueous solutions at T = 298.15 K. Figure 3 shows the EMF−ln msalt diagram for Na2SO4 + H2O system at 303.15 K. For abscissa axis, logarithmic coordinate is used. From the figure, we can directly find an intersection point of the two fitting lines. The corresponding value of x coordinate is the solubility of Na2SO4 in water. The measured solubility for the Na2SO4 + H2O system is listed in Table 2, which is in

Figure 4. EMF of solution as a function of logarithmic molality of Na2CO3 in water at 303.15 K. The lines and symbols have the same meaning as in Figure 3.

Figure 5. EMF of solution as a function of logarithmic molality of NaCl in water at 303.15 K. The lines and symbols have the same meaning as in Figure 3. C



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excellent agreement with the previously literature data. To further proof the reliability of the method, we also measured two other systems, that is, Na2CO3 + H2O and NaCl + H2O systems at 303.15 K. The results are plotted in Figures 4 and 5, respectively. The corresponding EMF−ln msalt diagrams at 298.15 K are plotted in Figures S1−S3. As we can see from these figures, the EMFs increase as the logarithmic molality of salts is increased before the solution reaches its saturation concentration. That is due to that the ion concentration in the solution increases as the salts are added into the solution. After reaching the saturation concentration, the EMF almost keeps a constant because the concentration in the solution remains unchanged. In order to avoid the effect of a phenomenon of supersaturation, a horizontal line after saturation was adopted to draw in the figures to eliminate the influence caused by the supersaturation. The mass fractions obtained from the two figures are also given in Table 2. As we can see from Table 2, the solubility data obtained in this work is extremely close to the data of literature at the two temperatures. The solubility data of Na2SO4 received in this work are both 0.04% lower than those from literature at two temperatures. Those of Na2CO3 received in this work are 0.5% lower at 298.15 K and 0.03% lower at 303.15 K than those from literature. For NaCl, the experimental data are 0.11% higher at 298.15 K but 0.01% lower at 303.15 K than literature data. In conclusion, the comparisons between the experimental data obtained in this work and those from literature16−18 indicate that the largest error is 0.5%. Thus, it can be concluded that the method proposed to measure solubility data of sodium salts is simple and credible. However, the ISE method has limitations that it is unable to measure the solubility of system with inadequate ionization such as surfactant or STPP. This method is restricted to the strong electrolyte system such as NaCl, Na2CO3, and Na2SO4. Therefore, we were only able to measure half of the ternary phase diagram, which is the solubility of NaCl, Na2CO3, and Na2SO4 in aqueous surfactant or STPP solution. In addition, we need to measure many EMF data (as shown in Figures S1−S11, Supporting Information) for one solubility point. Therefore, as an introduction of a new method this work provides only limited experiment data points of ternary system under specified temperature. Using the validated ISE method, we obtain the experimental solubility data for the ternary systems from the EMF−ln msalt plots (Figures S3−S11, Supporting Information). Two EMF− ln msalt diagrams of the ternary systems were selected as representatives and are shown in Figures 6 and 7. Figure 6 depicts the EMF increases up to saturation. Furthermore, we can see a phenomenon of supersaturation in this figure. To exclude the effect resulting from the supersaturation, we draw a horizontal line. The intersection point is the solubility of the sodium salt. As the salt is added to aqueous SDS solution, the solution becomes turbid, which indicates that the solution separates into a surfactant-rich and a surfactant-poor phase at this salt composition. This is clouding behavior of aqueous surfactant solutions. At clouding point, there is no solid salt phase and precipitate of surfactant formed. Even when the turbid solution is saturated with the salt, surfactant only exists in the turbid liquid mixture, and not in the solid phase, as verified by determining the surfactants with the two-phase titration with two dyes (thymol blue and methylene blue). Tables 3 to 5 present the experimental solubility data in weight fraction (w1) for each ternary system. The uncertainties

Figure 6. EMF of solution as a function of logarithmic molality of Na2SO4 in aqueous SDBS solution with mSDBS = 0.1435 mol/kg at 298.15 K. The lines and symbols have the same meaning as in Figure 3.

Figure 7. EMF of solution versus logarithmic molality of NaCl in aqueous STPP solution with mSTPP = 0.0544 mol/kg at 298.15 K. The lines and symbols have the same meaning as in Figure 3.

Table 2. Experimental Solubility of Sodium Salts in Water at Temperatures T = 298.15 and 303.15 K and Pressure p = 101.3 kPaa solubility (wt %) salt

this work

Na2SO4 Na2CO3 NaCl

21.90 23.00 26.56

Na2SO4 Na2CO3 NaCl

29.18 28.67 26.51

literature

equilibrium solid phase

T = 298.15 K 21.94b Na2SO4·10H2O 23.50c Na2CO3·10H2O 26.45d NaCl T = 303.15 K 29.22b Na2SO4·10H2O 28.70c Na2CO3·10H2O 26.52d NaCl

error (%) 0.04 0.50 0.11 0.04 0.03 0.01

a

Literature data are included for comparison. Standard uncertainties u are u(T) = 0.1 K, u(p) = 0.3 kPa, and u(w) = 0.0007. bFrom Potter et al.16 cFrom Sohnel and Novotny.17 dFrom Young.18

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Table 3. Experimental Solid−Liquid Equilibrium Mass Fractions w for the Systems of Na2SO4 (1) + SDBS (2) + H2O (3), Na2CO3 (1) + SDBS (2) + H2O (3), and NaCl (1) + SDBS (2) + H2O (3) at Temperature T = 298.15 K and Pressure p = 101.3 kPaa w1 0.2190 0.2142 0.2071 0.2021 0.1953 0.1901 0.1784 0.1670 0.2300 0.2249 0.2207 0.2165 0.2123 0.2075 0.2016 0.1973 0.2656 0.2583 0.2529 0.2466 0.2397 0.2331 0.2273 0.2170

w2

w3

Table 4. Experimental Solid−Liquid Equilibrium Mass Fractions w for Systems of Na2SO4 (1) + SDS (2) + H2O (3), Na2CO3 (1) + SDS (2) + H2O (3), and NaCl (1) + SDS (2) + H2O (3) at Temperature T = 298.15 K and Pressure p = 101.3 kPaa w1


Na2SO4 (1) + SDBS (2) + H2O (3) 0.0000 0.7810 Na2SO4·10H2O 0.0078 0.7780 Na2SO4·10H2O 0.0156 0.7773 Na2SO4·10H2O 0.0233 0.7746 Na2SO4·10H2O 0.0310 0.7737 Na2SO4·10H2O 0.0386 0.7713 Na2SO4·10H2O 0.0502 0.7714 Na2SO4·10H2O 0.0615 0.7715 Na2SO4·10H2O Na2CO3 (1) + SDBS (2) + H2O (3) 0.0000 0.7700 Na2CO3·10H2O 0.0077 0.7674 Na2CO3·10H2O 0.0154 0.7639 Na2CO3·10H2O 0.0228 0.7607 Na2CO3·10H2O 0.0304 0.7573 Na2CO3·10H2O 0.0377 0.7548 Na2CO3·10H2O 0.0488 0.7496 Na2CO3·10H2O 0.0595 0.7432 Na2CO3·10H2O NaCl (1) + SDBS (2) + H2O (3) 0.0000 0.7344 NaCl 0.0075 0.7342 NaCl 0.0147 0.7324 NaCl 0.0221 0.7313 NaCl 0.0292 0.7311 NaCl 0.0366 0.7303 NaCl 0.0437 0.7290 NaCl 0.0581 0.7249 NaCl

0.2190 0.2032 0.1892 0.1752 0.1629 0.1536 0.1440 0.1315 0.1131 0.2300 0.2314 0.2343 0.2336 0.2304 0.2248 0.2194 0.2138 0.2082 0.2656 0.2631 0.2611 0.2594 0.2565 0.2542 0.2509 0.2465 0.2435

a

Standard uncertainties u are u(T) = 0.1 K, u(p) = 0.3 kPa, and u(w) = 0.0007.

w3


a

Standard uncertainties u are u(T) = 0.1 K, u(p) = 0.3 kPa, and u(w) = 0.0007.

in the measurements of the temperatures and compositions were found to be 0.1 K and 0.0007, respectively. To observe regularities directly, the corresponding data are depicted in Figures 8 to 10. Figure 8 displays the solubility data for the sodium salt + SDBS + water ternary systems. From this figure, we can discover that the slopes of the tie lines are negative, which means the mass fraction of sodium salts decreasing as the mass fraction of surfactants increasing. Thus, we deduce that a salting-out effect exists in the solution. Although the solubility of salt in mixed electrolyte solutions can be well described by the modified mean spherical approximation19,20 and semiempirical activity coefficient of Pitzer type,21 we determined the influence of added SDBS on the three salts using a modified Setschenow equation22 for simplicity. The well-known Setschenow equation that describes the effect of a salt on the solubility of a nonelectrolyte in water is written on the mole fraction scale as follows22

⎛x ⎞ lg⎜ 0 ⎟ = kxxs ⎝x⎠

w2

Na2SO4 (1) + SDS (2) + H2O (3) 0.0000 0.7810 Na2SO4·10H2O 0.0081 0.7887 Na2SO4·10H2O 0.0161 0.7947 Na2SO4·10H2O 0.0241 0.8007 Na2SO4·10H2O 0.0322 0.8049 Na2SO4·10H2O 0.0405 0.8059 Na2SO4·10H2O 0.0485 0.8075 Na2SO4·10H2O 0.0611 0.8074 Na2SO4·10H2O 0.0806 0.8063 Na2SO4·10H2O Na2CO3 (1) + SDS (2) + H2O (3) 0.0000 0.7700 Na2CO3·10H2O 0.0077 0.7609 Na2CO3·10H2O 0.0151 0.7506 Na2CO3·10H2O 0.0224 0.7440 Na2CO3·10H2O 0.0296 0.7400 Na2CO3·10H2O 0.0372 0.7380 Na2CO3·10H2O 0.0479 0.7327 Na2CO3·10H2O 0.0586 0.7276 Na2CO3·10H2O 0.0722 0.7196 Na2CO3·10H2O NaCl (1) + SDS (2) + H2O (3) 0.0000 0.7344 NaCl 0.0080 0.7289 NaCl 0.0144 0.7245 NaCl 0.0219 0.7187 NaCl 0.0287 0.7148 NaCl 0.0368 0.7090 NaCl 0.0450 0.7041 NaCl 0.0570 0.6965 NaCl 0.0682 0.6883 NaCl

Table 5. Experimental Solid−Liquid Equilibrium Mass Fractions w for the Systems of Na2SO4 (1) + STPP (2) + H2O (3) and NaCl (1) + STPP (2) + H2O (3) at Temperature T = 298.15 K and Pressure p = 101.3 kPaa w1 0.2190 0.2141 0.2127 0.2102 0.2017 0.1966 0.2656 0.2629 0.2603 0.2573 0.2550 0.2518

(2)

where x is the mole fraction of the sodium salt in aqueous phase, the subscripts 0 and s denote water and STPP or surfactant, respectively. kx is the salting effect constant. The nature of the salting effect depends on the sign of kx: positive for salting-out and negative for salting-in. Salting effect constants kx in ternary systems are shown in Table 6. In the

w2

w3


Na2SO4 (1) + STPP (2) + H2O (3) 0.0000 0.7810 Na2SO4·10H2O 0.0081 0.7778 Na2SO4·10H2O 0.0155 0.7718 Na2SO4·10H2O 0.0232 0.7666 Na2SO4·10H2O 0.0305 0.7678 Na2SO4·10H2O 0.0379 0.7655 Na2SO4·10H2O NaCl(1) + STPP (2) + H2O (3) 0.0000 0.7344 NaCl 0.0078 0.7293 NaCl 0.0145 0.7252 NaCl 0.0217 0.7210 NaCl 0.0287 0.7163 NaCl 0.0356 0.7126 NaCl

a Standard uncertainties u are u(T) = 0.1 K, u(p) = 0.3 kPa, and u(w) = 0.0007.

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Figure 10. Solubility of sodium salts in the systems of Na2SO4 (1) + STPP (2) + H2O (3) and NaCl (1) + STPP (2) + H2O (3) at 298.15 K. Data are presented in mass fraction and dashed lines are regressed using eq 2.

Figure 8. Solubility of three sodium salts in the systems of Na2SO4 (1) + SDBS (2) + H2O (3), Na2CO3 (1) + SDBS (2) + H2O (3), and NaCl (1) + SDBS (2) + H2O (3) at 298.15 K. Data are presented in mass fraction and dashed lines are regressed using eq 2.

Table 6. Salting Effect Constants kx in Ternary Systems ternary system

kx

Na2SO4 (1) + SDBS (2) + H2O (3) Na2CO3 (1) + SDBS (2) + H2O (3) NaCl (1) + SDBS (2) + H2O (3) Na2SO4 (1) + SDS (2) + H2O (3) Na2CO3 (1) + SDS (2) + H2O (3) NaCl (1) + SDS (2) + H2O (3) Na2SO4 (1) + STPP (2) + H2O (3) NaCl(1) + STPP (2) + H2O (3)

21.9461 14.2729 19.6780 56.6337 −14.7714 1.4419 8.2454 4.1344

three sodium salts, the kx of Na2SO4 + SDBS + H2O system is the positively most remarkable indicating that the solubility of Na2SO4 suffers the most significant effect of the added SDBS. The kx of NaCl + SDBS + H2O system is slightly smaller than that of Na2SO4 + SDBS + H2O and larger than that of Na2CO3 + SDBS + H2O. Through comparing the values of salting-out constants, we can determine the order of the influence of added SDBS on the three salts is Na2CO3 < NaCl < Na2SO4. Three regression lines calculated by the salting effect constants are also plotted in Figure 8. Comparing the solid lines with the dash lines, we find that there are tiny differences between them, which further confirm the reliability of the experimental results. Figure 9a shows the solid−liquid equilibrium data for the sodium salt + SDS + H2O ternary systems. The slopes of NaCl and Na2SO4 are both negative; however, the concentration change of Na2SO4 is larger than that of NaCl. Thus, it can be deduced that in the ternary system, the solubility of Na2SO4 is more sensitive to the mass of SDS. The slope of Na2CO3 changes from slightly positive to negative, as shown in Figure 9a, which could be explained by the character of SDS. When the solution is in a state of dilution, SDS in the solution may promote the dissolution of salt. As the concentration of solution gradually increases, SDS may lead to the competition with the adding CO32− in the solution to reduce the dissolution capacity of Na2CO3. By use the values of kx listed in Table 6, the competitions in the three systems can be evaluated. It can be easily concluded that the most obvious salting-out system is Na2SO4 + SDS + H2O. The value of kx for Na2CO3 + SDS + H2O system is negative representing salting-in effect in dilute solution, which is consistent with the deduction we put

Figure 9. (a) Salt solubility and (b) mass fraction at clouding points for the systems of Na2SO4 (1) + SDS (2) + H2O (3), Na2CO3 (1) + SDS (2) + H2O (3), and NaCl (1) + SDS (2) + H2O (3) at 298.15 K. Data are presented in total mass fraction in one or two liquid phases. The dashed lines are regressed using eq 2 and the stars represent the data from Zhou and Hao.23

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forward. Using these kx values for the Setschenow equation, we derived the dashed solubility lines in Figure 9a in comparison with the experimental data. When we compared our solubility data with those from Zhou and Hao23 for the system of NaCl + SDS + H2O, wide discrepancies were found. Thus, we checked their method for the measurements and confirmed that what they reported are the composition of NaCl at the clouding point (where turbidity takes place). Figure 9b displays our measured results for the mass fractions of NaCl, Na2SO4, and Na2CO3 at the clouding points. For comparison, we also included the data of Zhou and Hao23 in Figure 9b. As can be seen, our results are in good agreement with the limited data of Zhou and Hao.23 It should be pointed out that the solubility measured by the ISE represents that when the crystal salts are in equilibrium with the single liquid phase or several liquid phases. Therefore, although the two groups of data are all solubility, they represent different meanings of thermodynamics. Figure 10 depicts solubility curves of NaCl + STPP + H2O and Na2SO4 + STPP + H2O systems. The fitted line of NaCl + STPP + H2O system is almost a straight line with a negative slope, but the line of Na2SO4 + STPP + H2O system is a curve taking a “saddle pattern”. This warped phenomenon for curve of the STPP + Na2SO4 + H2O may be caused by the competition between the ions of STPP and Na2SO4. The values of kx in Table 6 demonstrate the effect of STPP on the solubility of NaCl and Na2SO4 in water. They reveal that the effect of STPP on Na2SO4 is much more than on NaCl. In general, Figures 8−10 show that as the mass fraction of surfactants increases, the mass fractions of sodium salts decrease. From the figures, we deduce that the solubility of sodium salts in the aqueous system decreases with increasing mass of surfactants and STPP. This deduction is demonstrated by the value of the salting effect constants kx. The added surfactants give rise to the salting-out effect and suppress the dissolving capacity of sodium salts in water, that is, decreasing the solubility of sodium salts in aqueous system. Moreover, through analyzing the kx values in Table 6, we conclude that in the three sodium salts, the dissolving capacity of Na2SO4 is most suppressed by the salting-out effect, and in the surfactant and STPP, SDS produces the greatest salting-out influence on the Na2SO4.

the solubility of Na2SO4, in which the effect caused by SDS is the most remarkable.

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ASSOCIATED CONTENT

* Supporting Information S

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jced.5b00926. Data tables of Figures 3−7, and EMF−ln msalt diagrams used to determine the solubility values. (PDF)

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Funding

This work was supported by the National Natural Science Foundation of China (No. 21376131) and the China’s Postdoctoral Science Foundation (No. 2015M581109). Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We would like to acknowledge Yong-Jun Du, Wen-Dong Wu, and Kun Tian for their great help during the experimental measurements.

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REFERENCES

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4. CONCLUSIONS We have proposed an ISE method to measure solubility of sodium salts in aqueous surfactant and STPP solutions. The proposed method has been validated by comparing the measured solubility data of sodium salts in pure water with those reported in literature. We found that the ISE method has advantages in convenience, sensitivity and high accuracy. And the sensitivity is the advantage and also the drawbacks of this method. Because of the sensitivity, it can be easily estimated that the dissolution equilibrium is arrived when the EMF value becomes constant. However, also because of that, the EMF value is easily affected by the surrounding environment, which leads to measured data unstable. Therefore, multiple measurements need to be provided to ensure the accuracy of the data. In addition, the ISE method is unable to measure the solubility of incomplete ionization system and therefore is restricted to the strong electrolyte systems. The experimental solubility data demonstrate that the salting-out effect caused by the surfactant can suppress the dissolving capacity of sodium salts. In addition, the salting-out effect has the greatest influence on G



Article

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Determination of Solubility of Sodium Salts in Aqueous Surfactant

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