Article pubs.acs.org/jced
Solubility, Density and Solution Thermodynamics of NaI in Different Pure Solvents and Binary Mixtures Ramesh R. Pawar,* Chandrakant S. Aher, Jitendra D. Pagar, Sonali L. Nikam, and Mehdi Hasan P. G. Department of Physical Chemistry, M.S.G. College Malegaon Camp, Pin 423105, India
ABSTRACT: Solubility of NaI in water, methanol, ethanol, propan-2-ol, and also in water + methanol, water + ethanol, and water + propan-2-ol binary mixtures have been experimentally measured using a gravimetric method at temperatures (298.15.303.15, 308.15, and 313.15) K. The combined NIBS (nearly ideal binary solvent)−Redlic−Kister equation is used to fit experimental solubility data at constant temperature. The densities of the saturated solutions are also reported in pure and binary solvent mixtures at temperatures mentioned above. Thermodynamic functions including ΔH0soln, ΔG0soln, and ΔS0soln of solution of NaI are obtained from the modified van’t Hoff equation. A comparison of the relative contributions by enthalpy (ζH) and entropy (ζTS) is made which indicated that the main contributor to the positive standard molar Gibbs energy of solution of NaI is the entropy for solution of NaI in water + methanol having x0C < 0.5 and for all other solutions it is enthalpy.
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INTRODUCTION NaI is used for the treatment of iodine deficiency1,2 and as a reactant in the Finkelstein reaction.3 Solubility studies of electrolytes have applications in diverse fields such as the pharmaceutical industry, agriculture, biology, medicine, etc. Important information can be obtained about interactions in the solution. Inorganic salts can be recovered from aqueous solutions by addition of alcohols as co-solvents, which has several advantages over the standard crystallization technique.4−7 Solubility data for many inorganic salts in aqueous systems are available.8−10 Solubility data of sodium iodide in water−ethanol and water−propan-2-ol mixed solvents for a few compositions and at 303.15 K is available in ref 11. However there is no data available for solubility of sodium iodide in water−methanol, water−ethanol, and water−propan-2-ol for the complete binary composition range at temperatures mentioned above. Therefore in continuation of our work12,13 we have undertaken systematic measurements of solubility and densities of sodium iodide in water + methanol/ethanol/propanol-2-ol binary solvents over the entire composition range from 0 to 1 mass fraction of methanol, ethanol, and propan-2-ol at (298.15, 303.15, 308.15, and 313.15) K. The experimental data of solubility at constant temperature in pure and mixtures of solvents were fitted in a combined nearly ideal binary solvent (NIBS)−Redlic−Kister equation. Solubility of NaI depends on not only the temperature but also the composition of the solvent mixture. The thermodynamic functions of solution of NaI were calculated by using van’t Hoff equation. © 2012 American Chemical Society
EXPERIMENTAL SECTION Material. In all experiments, triple distilled water was used. Sodium iodide was supplied by MERCK with purity (NaI) 99.5 %. Methanol 99.9 %, ethanol 99.9 %, and propan-2-ol 99.5 % were supplied by Jiangyin Huaxi International Trade Co. (China). Apparatus and Procedure. The apparatus and procedures used for solubility and density measurement have been described earlier.12−15 An excess amount of NaI was added to the binary solvents mixtures prepared by weight (Shimadzu, Auxzzo) with an uncertainty of ± 0.1 mg, in a specially designed 100 mL double jacketed flask. Water was circulated at constant temperature between the outer and inner walls of the flask. The temperature of the circulating water was controlled by thermostat to within (± 0.1) K. The solution was continuously stirred using a magnetic stirrer for long time (about 1 h) so that equilibrium is assured, no further solute dissolved, and the temperature of solution is same as that of circulating water; the stirrer was switched off; and the solution was allowed to stand for 1 h. Then a fixed quantity of the supernatant liquid was withdrawn from the flask in a weighing bottle with the help of pipet which is hotter than the solution. The weight of this sample was taken and the sample was kept in an oven at 343 K until the whole solvent was evaporated and the residue was completely dry. This was Received: July 6, 2012 Accepted: October 29, 2012 Published: November 7, 2012 3563
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Table 1. Experimental Values of Mole Fraction Solubility, xB, and Density, ρ, for Various Initial Mole Fractions, x0C, of Methanol at Temperatures (298.15, 303.15, 308.15, and 313.15) K and Pressure 101.32 kPaa T/K
x0C
xB
ρ·10−3/kg·m−3
T/K
xB
ρ·10−3/kg·m−3
298.15
0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0
0.1813 0.1773 0.1727 0.1718 0.1726 0.1754 0.1914 0.1819 0.1734 0.1635 0.1533 0.1918 0.1873 0.1877 0.1857 0.1889 0.1956 0.1922 0.1820 0.1728 0.1634 0.1513
1.9169 1.8258 1.7368 1.6566 1.5927 1.5450 1.5563 1.4814 1.4109 1.345 1.2807 1.9520 1.8369 1.7773 1.6944 1.6468 1.6127 1.5502 1.4709 1.3985 1.3332 1.2624
303.15
0.1865 0.1828 0.1808 0.1794 0.1784 0.1846 0.1925 0.1821 0.1732 0.1636 0.1532 0.1991 0.1938 0.1962 0.1953 0.1997 0.2012 0.1923 0.1823 0.1731 0.1635 0.1503
1.9312 1.8366 1.7535 1.6786 1.6047 1.5747 1.5550 1.4760 1.4050 1.3392 1.2770 1.9732 1.8479 1.7836 1.7251 1.6745 1.6334 1.5452 1.4658 1.3934 1.3279 1.2544
308.15
a
313.15
Standard uncertainties u are u(T) = 0.1 K, u(x0C) = 0.0002, u(xB) = 0.003, and u(ρ) = 10 kg·m−3. The relative uncertainty in pressure ur(p) = 0.05.
Table 2. Solubility, xB, and Density, ρ, of Sodium Iodide for Various Initial Mole Fractions, x0C, of Ethanol at (298.15, 303.15, 308.15, and 313.15) K and Pressure 101.32 kPaa T/K
x0C
xB
ρ·10−3/kg·m−3
T/K
xB
ρ·10−3/kg·m−3
298.15
0.0000 0.0416 0.0891 0.1435 0.2068 0.2811 0.3697 0.4771 0.6100 0.7787 0.8953 1.0000 0.0000 0.0416 0.0891 0.1435 0.2068 0.2811 0.3697 0.4771 0.6100 0.7787 0.8953 1.0000
0.1806 0.1744 0.1693 0.1687 0.1643 0.1610 0.1572 0.1470 0.1375 0.1206 0.1131 0.1094 0.1894 0.1873 0.1861 0.1835 0.1821 0.1805 0.1750 0.1680 0.1633 0.1541 0.1491 0.1480
1.8996 1.8493 1.7275 1.6379 1.4941 1.3700 1.2384 1.1140 1.0040 0.9544 0.8965 0.8744 1.9469 1.8783 1.7484 1.66611 1.5263 1.3959 1.2610 1.1386 1.0314 0.9769 0.9412 0.8953
303.15
0.1858 0.1840 0.1818 0.1786 0.1744 0.1710 0.1640 0.1566 0.1499 0.1418 0.1351 0.1334 0.1972 0.1936 0.1981 0.1914 0.1942 0.1902 0.1850 0.1793 0.1740 0.1685 0.1611 0.1545
1.9223 1.8660 1.7356 1.6488 1.5095 1.3785 1.2443 1.1227 1.0124 0.9647 0.9385 0.8818 1.9489 1.8960 1.7697 1.6825 1.5551 1.4215 1.2846 1.1570 1.0589 1.0020 0.9851 0.9228
308.15
a
313.15
Standard uncertainties u are u(T) = 0.1 K, u(x0C) = 0.0002, u(xB) = 0.003, and u(ρ) = 10 kg·m−3. The relative uncertainty in pressure ur(p) = 0.05.
confirmed by weighing two or three times until a constant weight was obtained after keeping the sample in an oven for another 30 min every time. The solubility has been calculated using weight of solute and weight of solution. Each experimental value of solubility is an average of at least three different measurements
and the standard uncertainty of the experimental mole fraction solubility (xB), value is ± 0.003. The saturated mole fraction solubility (xB), initial the mole fraction of methanol/ethanol/ propan-2-ol (x0C), and initial the mole fraction of water (x0C) were calculated using usual equations. The standard uncertainty for x0C 3564
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Table 3. Solubility, xB, and Density, ρ, of Sodium Iodide for Various Mole Fractions, x0C, of Propan-2-ol at (298.15, 303.15, 308.15, and 313.15) K and Pressure 101.32 kPaa T/K
x0C
xB
ρ·10−3/kg·m−3
T/K
xB
ρ·10−3/kg·m−3
298.15
0.0000 0.0254 0.0554 0.0914 0.1353 0.1901 0.2603 0.3538 0.4842 0.6786 0.8233 0.9000 1.0000 0.0000 0.0254 0.0554 0.0914 0.1353 0.1901 0.2603 0.3538 0.4842 0.6786 0.8233 0.9000 1.0000
0.1806 0.1627 0.1459 0.1280 0.1144 0.1056 0.0982 0.0843 0.0685 0.0440 0.0324 0.0156 0.0067 0.1918 0.1786 0.1651 0.1493 0.1362 0.1282 0.1252 0.1193 0.1133 0.0925 0.0883 0.0722 0.0596
1.8996 1.7509 1.6291 1.5395 1.3957 1.2716 1.1400 1.0155 0.9054 0.8560 0.8125 0.7701 0.7760 1.9467 1.7889 1.6590 1.5717 1.4369 1.3065 1.1716 1.0492 0.9420 0.8875 0.8341 0.8198 0.8059
303.15
0.1858 0.1704 0.1549 0.1376 0.1257 0.1161 0.1108 0.1024 0.0922 0.0679 0.0591 0.0411 0.0289 0.1972 0.1853 0.1728 0.1579 0.1476 0.1374 0.1355 0.1315 0.1284 0.1105 0.0978 0.0883 0.0785
1.9223 1.7766 1.6462 1.5594 1.4201 1.2891 1.1549 1.0333 0.9231 0.8753 0.8436 0.7900 0.7925 1.9489 1.7976 1.6713 1.5841 1.4567 1.3231 1.1862 1.0585 0.9605 0.9036 0.8654 0.8200 0.8244
308.15
a
313.15
Standard uncertainties u are u(T) = 0.1 K, u(x0C) = 0.0002, u(xB) = 0.003, and u(ρ) = 10 kg·m−3. The relative uncertainty in pressure ur(p) = 0.05.
Figure 1. Mole fraction solubility (xB) variation with initial mole fraction (x0C) of methanol at temperatures (298.15, 303.15, 308.15, and 313.15) K. 3565
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Figure 2. Mole fraction solubility (xB) variation with initial mole fraction (x0C) of propan-2-ol at temperatures [298.15 (■), 303.15 (▲), 308.15 (×), and 313.15 (⧫)] K and ethanol at temperatures [298.15 (•), 303.15 (+), 308.15 (-), and 313.15 (*)] K.
is 0.0002. Densities were determined using a 15 cm3 bicapillary pycnometer as described earlier14,15For calibration of pycnometer triply distilled and degassed water with a density16 of 0.99705 g·cm−3 at 298.15 K was used. The pycnometer filled with air bubble free experimental liquids was kept in a transparent walled thermostat (maintained at constant temperature ± 0.1 K) for (10 to 15) min to attain thermal equilibrium. The heights of the liquid levels in the two arms were measured with the help of a traveling microscope, which could read to 0.01 mm. The estimated standard uncertainty of the density measurements of the solvent and binary mixtures was 10 kg·m−3.
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RESULTS AND DISCUSSION The experimental values of solubility (xB) and density (ρ) of the saturated solutions at (298.15, 303.15, 308.15, and 313.15) K for sodium iodide + water + methanol/ethanol/propan-2-ol are given in Tables 1, 2, and 3, respectively. Variation of solubility with x0C is visually shown in Figures 1and 2. Figure 3 shows a comparison of our experimental values of solubility of sodium iodide in pure water with literature values.8−10 It can be seen from the figure that our experimental values agree well with literature values. It can be seen that the solubility of NaI in both pure solvents and binary solvent mixtures with given initial compositions increases with temperature (except in water + methanol with x0C >0.8). At the same temperature, the solubility trend in pure solvent is water > methanol > ethanol > propan-2-ol and in binary mixtures water + methanol > water + ethanol > water + propan-2-ol.This trend implies that solubility of NaI increases with increasing polarity of solvent. Ion-dipole interactions play an important
Figure 3. Solubility of NaI in water as a function of temperature, ●, this study; ○, ref 5; ■, ref 7; ▲, ref 8. 3566
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Table 4. Values of Coefficients of NIBS−Redlich−Kister Equations T/K 298.15 303.15 308.15 313.15
298.15 303.15 308.15 313.15
298.15 303.15 308.15 313.15
range of x0C
M0
0.1−0.5 0.6−0.9 0.1−0.5 0.6−0.9 0.1−0.5 0.6−0.9 0.1−0.5 0.6−0.9
0.165 0.843 0.350 0.815 0.694 0.56 0.611 0.629
0.042−0.207 0.28−0.90 0.042−0.207 0.28−0.9 0.042−0.207 0.28−0.9 0.042−0.207 0.28−0.9 0.02−0.90 0.02−0.90 0.02−0.90 0.02−0.90
−21.57 0.166 1.054 −0.049 8.826 0.005 0.673 0.085 2.858 0.888 0.139 0.020
M1 Water + Methanol + NaI −0.768 −2.438 1.82 −1.036 −7.830 3.512 1.552 0.486 Water + Ethanol + NaI −51.15 0.154 3.358 −0.733 22.14 −0.490 1.724 −0.427 Water + Propan-2-ol + NaI 0.254 −0.256 −0.267 −0.531
M3
R2
−0.336 3.436 2.736 2.003 −4.678 −5.702 1.635 0.096
0.553 −1.805 1.470 −1.320 −0.87 2.479 1.054 −0.397
0.996 1.000 0.995 1.000 0.994 1.000 0.983 1.000
−118.4 −0.502 7.182 0.269 49.26 0.149 3.825 0.397
−88.77 −0.176 4.800 0.016 36.20 −0.071 2.793 0.071
1.000 0.970 1.000 0.988 1.000 0.985 1.000 0.925
0.312 −0.706 −0.632 −0.965
3.988 2.311 1.723 1.397
0.965 0.961 0.960 0.978
M2
Figure 4. Mole fraction solubility (xB) of NaI in water + methanol solvent mixtures with temperature.
maximum at x0C = 0.6 for temperatures (298.15 and 303.15) K whereas for temperatures (308.15 and 313.15) K maximum is observed at x0C = 0.5. Temperature effect on solubility of NaI in water−methanol mixture is more pronounced up x0C = 0.6 afterward it becomes negligible (Figure 1). This variation of
role in determining solubility of NaI. As NaI is ionic compound and like dissolves like, its solubility increases with increasing polarity of solvent. The solubility of NaI in water−methanol mixture shows an interesting variation with x0C. It decreases first then reaches to 3567
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Figure 5. Variation of mole fraction solubility (xB) of NaI water + ethanol with temperature.
Figure 6. Variation of mole fraction solubility (xB) of NaI with temperature for water + propan-2-ol. 3568
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Table 5. Slope and Intercept of Plot of xB versus T water + methanol
water + ethanol
water + propan-2-ol
x0C
m
c
R2
x0C
m
c
R2
x0C
m
c
R2
0.000 0.100 0.200 0.300 0.400 0.500 0.600 0.700 0.800 0.900 1.000
0.001 0.0006 0.001 0.001 0.001 0.001 5·10−5 2·10−5 −3·10−5 −2·10−5 −2·10−5
−0.168 0.0016 −0.289 −0.287 −0.375 −0.351 0.177 0.175 0.180 0.169 0.218
0.992 0.990 0.998 0.993 0.982 0.985 0.950 0.992 0.961 0.770 0.909
0.0000 0.0416 0.0891 0.1435 0.2068 0.2811 0.3697 0.4771 0.6100 0.7787 0.8953 1.0000
0.001 0.001 0.001 0.001 0.001 0.001 0.002 0.002 0.002 0.003 0.003 0.003
−0.138 −1.94 −0.371 −0.267 −0.416 −0.418 −0.449 −0.495 −0.573 −0.825 −0.827 −0.779
0.976 0.983 0.969 0.984 0.993 0.999 0.989 0.995 0.995 0.987 0.979 0.936
0.0000 0.0254 0.0554 0.0914 0.1353 0.1901 0.3538 0.4842 0.6786 0.8233 0.9000 1.0000
0.001 0.001 0.001 0.002 0.002 0.002 0.002 0.003 0.004 0.004 0.004 0.005
−0.152 −0.29 −0.396 −0.476 −0.535 −0.542 −0.654 −0.859 −1.126 −1.308 −1.461 −1.469
0.999 0.998 0.997 0.996 0.997 0.999 0.996 0.992 0.990 0.962 0.992 0.986
where x0A is initial mole fraction of water and x1 and x2 are solubilities of NaI in pure methanol/ethanol/propan-2-ol and water, respectively. Mi is curve fit parameters (four parameter equation). It is found that for water + methanol + NaI and water + ethanol + NaI systems there are two rigions of fitness whereas for water + propan-2-ol system single set of Mi values obtained throught the region studied. For water + methanol + NaI system two regions are x0C = 0.1to 0.5 and x0C = 0.6 to 0.9, for water + ethanol the first region is from x0C = 0.04 to 0.20 and the second region is from x0C = 0.28 to 0.89. All values of Mi along with R2 values are listed in Table 4. The solubility data varies linearly with temperature for a fixed composition of solvent mixture (Figures 4, 5, and 6) obeying a straight line equation
x B = mT + c
The values of slope and intercept are listed in table 5 along with correlation coefficient. Thermodynamic functions of solution are important to study the dissolution behavior of the solute in different solvents. The temperature dependence of the solubility allows a thermodynamic analysis that permits insight into the molecular mechanisms involved in the solution processes.21 In this work the thermodynamic functions in the process of solution of sodium iodide are calculated on the basis of the solubility of sodium iodide in different solvents. According to the van’t Hoff equation, the standard molar enthalpy change of solution ΔH0soln is generally obtained from the slope of the ln xB vs 1/T plot. Average temperature Tmean is introduced to obtain a single value of ΔG0soln and ΔS0soln in the temperature range studied
Figure 7. Plot of ln xB versus 10000(1/T − 1/Tmean) for NaI + water + methanol system.
solubility implies that ion-dipole interaction is maximum, at x0C = 0.6 for temperatures (298.15 and 303.15) K and at x0C = 0.5 for (308.15 and 313.15) K.The little effect of temperature on solubility of NaI beyond x0C = 0.6 shows thermal energy in this particular temperature range studied is insufficient to break the association of solvent molecules. The solubility of the sodium iodide in water ethanol mixture decreases appreciably with increasing mole fraction (x0C) of ethanol. Whereas the solubility of the sodium iodide in water propan-2-ol decreases sharply with increasing mole fraction (x0C) of propan-2-ol up to 0.19 and after x0C = 0.19, the decrease in xB is slow for all four temperatures. This implies that solvent−solute interaction is more and solvent−solvent interaction is less in NaI + ethanol + water system than in NaI + propanol-2-ol+ water system. The density of the saturated solution of NaI is found to decrease with increase of mole fraction of methanol/ethanol/ propan-2-ol. However there is a slight increase in density with increase of temperature. The solubility data at constant temperature is fitted in to combined NIBS−Redlic−Kister model17−20
Tmean =
ln x B =
ln x1 +
xA0
ln x 2 +
xC0xA0
∑ i=0
Mi(xC0
−
xA0 )i
n n ∑i = 1
( T1 )
(3)
where n is the number of experimental points. In the present work, Tmean = 305.65 K and the temperature range is (298.15 to 313.15) K in both pure solvents and binary solvent mixtures. Heat capacity of the solution can be assumed as constant. Hence values of ΔH0soln are derived using eq 4. 0 ΔHsol
3
xC0
(2)
⎡ ⎛ ∂ ln x B ⎞ ⎢ ∂ ln x B = −R ⎜ ⎟ − R⎢ ⎝ ∂(1/T ) ⎠ ⎢∂ 1 − 1 Tmean ⎣ T
(
)
⎤ ⎥ ⎥ ⎥ ⎦
(4)
The ln xB vs 10000 (1/T − 1/Tmean) plot of different solutions including pure solvents and binary solvent mixtures are displayed
(1) 3569
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Figure 8. Plot of ln xB versus 10000(1/T − 1/Tmean)for NaI + water + ethanol.
Figure 9. Plot of ln xB versus 10000(1/T − 1/Tmean) for NaI + water + propan-2-ol.
in Figures 7, 8, and 9. From these figures, it can be seen that a trend of increasing solubility with temperature is observed. The slope and the intercept for each solvent are listed in Table 6. Thus the modified van’t Hoff equation can be thought to be fit to calculate the enthalpy change of solution. The standard molar Gibbs energy change for the solution process ΔG0soln, can be calculated in the way similar to Krug et al22 as
0 ΔGsoln = −RT × intercept
(5)
In which the intercept used is that obtained in plots of ln xB as a function of (1/T − 1/Tmean). The standard molar entropy change ΔS0soln is obtained from 0 ΔSsoln =
3570
0 0 ΔHsoln − ΔGsoln Tmean
(6)
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Table 6. Slope and Intercept of the ln xB vs 10000(1/T − 1/Tmean) Plot water + methanol
water + ethanol
water + propan-2-ol
x0C
m
c
R2
x0C
m
c
R2
x0C
m
c
R2
0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0
−0.057 −0.029 −0.078 −0.078 −0.092 −0.087 0.002 −0.0009 0.001 0.0004 0.013
−1.662 −1.696 −1.691 −1.698 −1.688 −1.665 −1.649 −1.703 −1.753 −1.811 −1.884
0.992 0.946 0.998 0.994 0.983 0.985 0.482 0.552 0.468 0.262 0.904
0 0.0416 0.0891 0.1435 0.2068 0.2811 0.3697 0.4771 0.61 0.7787 0.8953 1
0.052 0.063 0.092 0.076 0.101 0.103 0.11 0.123 0.144 0.206 0.217 0.218
−1.669 −1.689 −1.694 −1.712 −1.722 −1.739 −1.762 −1.818 −1.86 −1.926 −1.975 −1.998
0.976 0.982 0.969 0.983 0.994 0.999 0.991 0.998 0.993 0.979 0.967 0.992
0 0.0254 0.0554 0.0914 0.1353 0.1901 0.2603 0.3538 0.4842 0.6786 0.8233 0.9 1
0.055 0.081 0.106 0.132 0.157 0.166 0.203 0.278 0.391 0.574 0.696 1.079 1.524
−1.667 −1.748 −1.836 −1.946 −2.037 −2.11 −2.149 −2.226 −2.323 −2.598 −2.752 −3.101 −3.48
0.999 0.998 0.996 0.996 0.999 0.996 0.993 0.983 0.972 0.971 0.924 0.926 0.91
Table 7. Thermodynamic Functions Relative to Solution Process of NaI at Tmean = 305.65 K x0C
ΔH0sol/kJ·K−1·mol−1
ΔG0soln/kJ·K−1·mol−1
0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1
4.7390 4.4896 6.4849 64.8492 7.6489 72.3318 0.1663 0.0715 −0.0831 −0.0355 −1.0808
4.2234 4.2819 4.2971 4.3149 4.2895 4.2311 4.1904 4.3287 4.4547 4.6021 4.7876
0.0416 0.0891 0.1435 0.2068 0.2811 0.3697 0.4771 0.6100 0.7787 0.8953
5.2378 7.6489 6.3186 8.3971 8.5634 9.1454 10.2262 11.9722 17.1268 18.0414
0 4.3028 4.3485 4.3739 4.4171 4.4755 4.6178 4.7244 4.8921 5.0166
0 0.0254 0.0554 0.0914 0.1353 0.1901 0.2603 0.3538 0.4842 0.6786 0.8233 0.9000 1
4.5727 6.7343 8.8128 10.9745 13.0530 13.8012 16.8774 23.1129 32.5077 47.7224 57.8654 89.7081 126.7054
4.2342 4.4400 4.6635 4.9429 5.1740 5.3595 5.4585 5.6541 5.9005 6.5990 6.9902 7.8766 8.8393
ΔS0soln/KJ·K−1·mol−1 Water + Methanol 0.0017 0.0007 0.0072 0.1981 0.0110 0.2228 −0.0132 −0.0139 −0.0148 −0.0152 −0.0192 Water + Ethanol 4.3233 0.0109 0.0064 0.0132 0.0136 0.0153 0.0183 0.0237 0.0400 0.0426 Water + Propan-2-ol 0.0011 0.0075 0.0136 0.0197 0.0258 0.0276 0.0374 0.0571 0.0871 0.1345 0.1664 0.2677 0.3856
Both ΔG0soln and ΔS0soln pertain to the mean temperature Tmean =
TΔS0soln/KJ·K−1·mol−1
ζH
ζTS
0.5155 0.2077 2.1878 60.5343 3.3594 68.1007 −4.0241 −4.2572 −4.5378 −4.6376 −5.8684
0.9019 0.9558 0.7477 0.5172 0.6948 0.5151 −0.0431 −0.0171 0.0180 0.0076 0.1555
0.0981 0.0442 0.2523 0.4828 0.3052 0.4849 1.0431 1.0171 0.9820 0.9924 0.8445
4.2393 3.3461 1.9701 4.0232 4.1463 4.6699 5.6085 7.2477 12.2347 13.0248
0.9160 0.6957 0.7623 0.6761 0.6738 0.6620 0.6458 0.6229 0.5833 0.5807
0.0840 0.3043 0.2377 0.3239 0.3262 0.3380 0.3542 0.3771 0.4167 0.4193
0.3385 2.2944 4.1494 6.0316 7.8789 8.4418 11.4189 17.4588 26.6073 41.1234 50.8753 81.8314 117.8661
0.9311 0.7459 0.6799 0.6453 0.6236 0.6205 0.5965 0.5697 0.5499 0.5371 0.5321 0.5230 0.5181
0.0689 0.2541 0.3201 0.3547 0.3764 0.3795 0.4035 0.4303 0.4501 0.4629 0.4679 0.4770 0.4819
by enthalpy and entropy respectively, which are calculated by eqs 7 and 823,24 toward the solution process
305.6 5 K. The results are shown in Table 7, together with ζH and ζTS.
ζH =
ζH and ζTS represent the comparison of the relative contributions 3571
0 ΔHsoln 0 0 |ΔHsoln | + |T ΔGsoln |
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dx.doi.org/10.1021/je300754n | J. Chem. Eng. Data 2012, 57, 3563−3572
Journal of Chemical & Engineering Data ζTS =
Article
0 |T ΔGsoln | 0 |ΔHsoln |
+
0 |TΔGsoln |
(6) Fleischmann, W.; Mersman, A. In Industrial Crystallization 84; Jancic, S., de Jong, J., Eds.; Elsevier Science Publisher: Amsterdam, 1984; pp 165−168. (7) Carton, A.; Sobron, F.; Bolado, S.; Tabares, J. Composition and Density of Saturated Solution of Lithium Sulphate + Water + Propan-1ol. J. Chem. Eng. Data 1994, 39, 61−62. (8) Solubilities of Inorganic and Organic Compounds, vol-1, Binary Systems; Part-1; Stephan, H., Stephen, T., Eds.; Pergamon Press: Oxford, England, 1963; p 110. (9) Dean, J. A. Langes Handbook of Chemistry, 13th ed.; McGraw-Hill: New York, 1987; p 10. (10) Lide, D. R. CRC Handbook of Chemistry and Physics, 79th ed.; CRC Press: Boca Raton, FL, pp 8−108. (11) Stephan, H.; Stephen, T. Solubilities of Inorganic and Organic Compounds; Pergamon Press: Oxford, England, 1964; Vol. 2, Ternary Systems, Part-1, p 133. (12) Pawar, R. R.; Nahire, S. B.; Hasan, M. Solubility and density of potsium iodide in binary water-ethanol solvent mixtures at(298.15, 303.15, 308.15, and 313.15) K. J. Chem. Eng. Data 2009, 54, 1935−37. (13) Pawar, R. R.; Golait, S. M.; Hasan, M; Sawant, A. B. Solubility and density of potsium iodide in binary water-propan-1-ol solvent mixtures at(298.15,303.15,308.15 and313.15)K. J. Chem. Eng. Data 2010, 55, 1314−16. (14) Kadam, U. B.; Hiray, A. P.; Sawant, A. B.; Hasan, M. Density,Viscosities, and Ultrasonic Velocity Studies of Binary Mixtures of Chloroform with Propan-1-ol and Butan-1-ol at (303.15 and 313.150) K. J. Chem. Eng. Data 2006, 51, 60−63. (15) Hasan, M.; Shirude, D. F.; Hiray, A. P.; Kadam, U. B.; Sawant, A. B. Densities, Viscosities and Speed of Sound Studies of Binary Mixtures of Methylbenzene with Heptan-1-ol, Octan-1-ol and Decan-1-ol at (303.15 and 313.15) K. J. Chem. Eng. Data 2006, 51, 1922−26. (16) Marsh, K. N. Recommended Reference Materials for the Realisation of Physicochemical Properties; Blackwell Scientific Publications: Oxford, U.K., 1987. (17) Jouyban, A.; Khoubnasabjafari, M.; Chan, H.; Clark, B. J.; Acree, W.; E., Jr. Solubility prediction of anthracene in mixed solvents using a minimum number of experimental data. Chem. Pharm. Bull. 2002, 50 (1), 25−25. (18) Joyce, R.; Powell, B.; Miller, J.; Acree, W. E., Jr. Solubility of Anthracene in binary alcohol + 1,4-dioxane solvent mixtures. J. Chem. Eng. Data 1995, 40, 1124−1126. (19) Acree, W. E., Jr.; Zvaigzne, A. I. Thermodynamic properties of nonelectrolyte solutions: Part 4. Estimation and mathematical representation of solute activity coefficients and solubilities in binary solvents using the NIBS and modified Wilson equations. Thermochim. Acta 1991, 178, 151−167. (20) Acree, W.e., Jr. Mathematical represesentation of thermodynamic properties: Part 2. Derivation of combined nearly ideal binary solvents NIBS/Redlich-Kister mathematical representation from a two body and three body interactional mixing model. Thermochim. Acta 1992, 198, 71−79. (21) Pacheco, D. P.; Martínez, F. Thermodynamic analysis of the solubility of naproxen in ethanol + water cosolvent mixtures. Phys. Chem. Liquids 2007, 45, 581−595. (22) Krug, R. R.; Hunter, W. G.; Grieger, R. A. Enthalpy-Entropy Compensation. 2. Separation of the Chemical from the Statistical Effect. J. Phys. Chem. 1976, 80, 2341−2351. (23) Perlovich, G. L.; Kurkov, S. V.; Bauer-Brandl, A. Thermodynamic of solutions II Flurbiprofen and diflunisal as models for studying salvation of drug substances. Eur. J. Pharm. Sci. 2003, 19, 423−432. (24) Perlovich, G. L.; Kurkov, S. V.; Kinchin, A. N.; et al. Thermodynamic of solutions III: comparison of the salvation of (+)-naproxen with other NSAIDs. Eur. J. Pharm. Biopharm. 2004, 57, 411−420.
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From Table 7 it can be concluded that the standard Gibbs free energy of sodium iodide is positive in all the studied pure and binary solvents as is the enthalpy of solution(except for x0C = 0.8,0.9,1.0 for water + methanol + NaI system), Therefore the process is always endothermic, and the entropy of solution is also positive (except >0.5 in the water + methanol + NaI system) in all of the solvents, indicating the entropy as driving the overall solution process for both pure solvents and mixtures. Moreover, the main contributor to the positive standard molar Gibbs energy of solution of sodium iodide is the enthalpy up to x0C = 0.5, and beyond x0C = 0.5 is entropy in water + methanol + NaI system. For other systems, the main contributor to the positive standard molar Gibbs energy of solution of sodium iodide is the enthalpy.
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CONCLUSION The solubility of sodium iodide was found to increase with increasing temperature in all (pure and mixed) solvents except in water + methanol with x0C > 0.8. In the pure solvents used in this work, the solubility of sodium iodide depends on the polarity of the solvents to some degree; as polarity decreases, the solubility also decreases. In the binary mixtures methanol + water, ethanol + water, and propan-2-ol + water, increasing the initial content of water results in an increase of its solubility. The experimental data are very well represented by the equations used. The solution process is always endothermic, and the entropy is driving the overall solution process. The main contributor to the positive standard molar Gibbs energy of solution of sodium iodide is the enthalpy up to x0C = 0.5 and beyond x0C = 0.5 is entropy in the water + methanol + NaI system. For other systems the main contributor to the positive standard molar Gibbs energy of solution of sodium iodide is the enthalpy.
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AUTHOR INFORMATION
Corresponding Author
*E-mail: rameshpawar_09@rediffmail.com. Funding
The authors are thankful to U.G.C. (New Delhi) for providing financial assistance in the form of minor research project for this work. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS The authors express their sincere thanks to the general secretary M. G. Vidyamandir and Principal M.S.G. college, Malegaon Camp, and for providing laboratory facility and encouragement.
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REFERENCES
(1) The Lancet, 2008, 372 (9633), 88. (2) Lyday, P. A. Iodine and Iodine Compounds. In Ullmanns Encyclopedia of Industrial Chemistry; Wiley-VCH: Winheim, Germany, 2005; Vol.15, pp 382−390. (3) Finkelstein. Ber. Dtsch. Ges. 1910, 43, 1528. (4) Lozano, J. A. F. Recovery of Potassium Magnesium Sulfate Double Salt from Seawater Bittern. Ind. Eng. Chem. Process Des. Dev. 1976, 15, 445−447. (5) Mydlarz, J.; Jones, A. G. Solubility and Density Isotherms for Magnesium Sulfate Heptahydrate-Water-Ethanol. J. Chem. Eng. Data 1991, 36, 119−121. 3572
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