Solubility and Density Isotherms J
J
POTASSIUM AND AMMONIUM NITRATES IN ETHANOL SOLUTIONS A. RALPH THOMPSON AXD R. E. VENER University of Pennsylvania, Philadelphia, Pa.
A study
of crystallization of inorganic salts by means of organic precipitants has been continued. Experimental solubility and density data are necessary to permit technical and economic consideration of such a process. Solubilities and saturated solution densities are reported at intervals of S o , from 25' to 75' C., for potassium and ammonium nitrates in aqueous ethanol solutions. Only one
liquid phase w-as found for potassium nitrate over the entire range, but for the ammonium salt, separation into two liquid phases occurred at temperatures above 65.7' C. For this region, compositions and densities of the phases in equilibrium are given. A discussion of earlier solubility data for these salts in water and in aqueous ethanol solutions is included.
T
weight mere determined, using solutions prepared from commercial alcohol and from purified alcohol. The purification was carried out by the method recommended by Lund and Bjerrum (13). To remove any possible volatile impurities, the procedure suggested by Coplep, Murray-Rust, and Hartley (5)was followed. Values obtained by using the commercial and purified alcohol solution for two different temperatures showed excellent agreement. At both temperatures the differences were not greater than 0.002% salt by weight. Since these checks were within the limits of the experimental measurements, it was decided that data obtained using aqueous solutions prepared from commercial alcohol would be satisfactory. Initially, ethanol solutions of approximately 10, 20, 30, 50, and 70% by weight were used in addition to the commercial (95y0by volume) product. In the case of the system involving the ammonium salt, it was found necessary to employ, in addition, solutions of about 82 and 88% by weight for the purpose of determining the approximate solvent compositions in the two-phase regions. hnalyses of these ethanol solutions were made by density measurements a t 35" * 0.02" C., using the density-composition data of Osborne, PvIcKelvy, and Bearce (16). For all ethanol solutions used, the average values of the density in vacuo at 35" C. are given in Table I, with the corresponding compositions in mole and weight per cent alcohol.
HE program to investigate the commercial possibilities of crystallization by means of organic precipitants, which was announced in a previous paper (19), has been continued. As stated in that article, complete solubility data for a number of systems of the type salt-water-organic liquid are needed in order to evaluate the economic feasibility of the method. In the present work the solubilities, phase equilibria, and densities of saturated solutions were determined for potassium and ammonium nitrates in aqueous solutions of ethanol over the temperature range 2 5 " to 75" C. Similar data have been presented for the same salts in isopropanol solutions (19). Most previous work involving similar ternary systems (inorganic salt-water-aliphatic alcohol) has been concerned chiefly with the effect of the presence of small amounts of salt on the mutual solubility of the alcohol and water. Considerable solubility data have been obtained for various inorganic salts in absolute ethanol or aqueous solutions of ethanol, but most of this work was done a t one temperature only (frequently at 25" (3.). Comparatively little information is reported for salts dissolved in aqueous ethanol over the full range of compositions and for more than one temperature. Mention might be made of those systems for which the ranges of temperature and composition have been well covered. Gerardin (9) studied the solubility in aqueous ethanol of the nitrates of potassium and lead as well as potmsium chloride. Data for ammonium chloride a t several temperatures were reported by Armstrong and Eyre (1). Fleckenstein ( 7 ) , Schreinemakers ( 1 7 ) , and de Wahl (BO) have reported some values for ammonium nitrate. Ammonium sulfate data of several investigators are reproduced by Seidell (18). Ketner (12) gave data for sodium carbonate; those for sodium chloride, by numerous investigators, are reported by Seidell (18). Sodium sulfate was investigated by de Bruyn ( 3 ) . Sodium thioantimonate was studied by Donk (6). The solubilities of potassium halides in alcohol-water mixtures at 10" and 20" C. were studied by Zeitlin ( 2 1 ) . Recently Gee (8) reported data for the system aluminum sulfate-ethanol-water for 30' and 80" C. As indicated earlier, data for other systems are fragmentary in nature.
TEMPERATURE, DENSITY, AND SALT CONTENTS
The constant temperature bath and auxiliary equipment employed in the work were described in detail in the previous paper (19). The temperature was ascertained by means of a calibrated thermometer, and it is considered that the temperature was known to 0.02" C. The density vias determined by weighing a definite volume of the solution in weighing pipets designed especially for handling saturated salt solutions a t elevated tem-
SOLUTIOXS TABLEI. ANALYSESOF AQUEOUSETHANOL (IN VACUO)
Solvent
PREPARATION OF ALCOHOL SOLUTIONS
KO.
E-10 E-20 E-30 E-50 E-70 E-82 E-88 E-92
As in the case of isopropanol, it was decided that commercial (95%) ethanol would be used for making up the aqueous solutions, provided the solubility in a test solution of given composition was the same as that in a solution of the same composition prepared from purified alcohol. The solubilities of the salts in approximately 50% alcohol by
478
Density a t 35' C . , G./Ml. 0.97714 0.96150 0,94424 0,90191 0.86510 0.82511 0.80996
Weight % Aloohol 9.82 19.90 29.89 49.90 69.83
0.79769
88.03 92.63
82,08
lllole % Alcohol 4.09 8.86 14.29 28.03 47.51 64.17 74.20 83.09
INDUSTRIAL AND ENGINEERING CHEMISTRY
March 1948
70
60
ci
-,u I
50
E
40
F
30
20 0
10
20
30
40
50
60
70
Wt 9- KN03 In Saturated Solution
Figure 1
479
compositions) is better than "0.05% for salt content and *0.0005 gram per milliliter for densities. SOLUBILITY OF SALTS IN WATER, It might be of some value t o compare the experimental data for the solubility of these salts in water obtained in this work with those of previous investigators. The data accepted in most solubility tables show some variance with those recently reported, including those found in the current work. For the most part, the data for the solubility of potassium nitrate in water show fair agreement (Table 11). I n the case of ammonium nitrate, a comparison with older values shows the present data, and also those of two other contemporary investigators, to be slightly lower throughout the temperature range than the values generally accepted in solubility tables. hlillican, Joseph, and Lowry (14) reported that errors inherent in their freezing point method tended to give slightly higher values. The data of Cohen and Bredee (4) check well with those found in this work, and those obtained by Hoeg (10) agree over the entire range (Table 111).
peratures. These pipets and the technique for their use have been described fully (19). The salt content was found by evaporating the solution to dryness in an oven a t 100-105' C. The composition of the solvent in each layer, when two liquid phases were present, was determined by two indirect analytical methods. The procedures, which involved extrapolations of large scale plots of density against solvent composition, and solubility against solvent composition, were described in the aforementioned article (19). EXPERIMENTAL RESULTS
'
In all figures, open circles represent analyses of samples taken when only one liquid phase was present; solid black points are used to give the same information for regions where there were two liquid phases. All density and solubility data were interpolated to give values at intervals of 5' C. by a frequently used graphical method, which involved such magnification that the values could be read easily to the same number of decimal places as the original data. The equation of a straight line which passed close to the experimental points was determined. For all experimental temperatures the deviation from the straight line, A (data value minus equation value), was calculated, and a plot was made of A against temperature. For temperatures a t intervals of 5" C., A was found on the plot and the interpolated value obtained by adding A to the equation value for the temperature in question. The over-all precision of these interpolated values (as determined from experimental measurements a t several temperatures and
TABLE11. SOLUBILITY OF POTASSIUM NITRATE IN WATER Temp., C. 25 30 40 50 60
70
(Weight % KNOs in saturated solution) Seidell Seidell Present (28), Av. (28), Av. I.C.T. Work No. 1 No. 2 (22) 27.24 27.5 28.00 27.2 31.49 31.4 31.2 31.70 39.02 . 39.0 38.0 39.,11 46.11 44.0 45.5 46.39 52.33 52.0 51.5 52.62 57.87 58.0 57.5 58.07
-
TABLE111. SOLUBILITY
Temp., C. 30 40 50 60 70
OF
Mulder (16)
27.2 30.8 39.0 46.2 52.6 58.1
AMMONIUM NITRATEI N WATER
(Weight % ' NHaNOs in saturated solution) Millican. Joseph, Cohen, Present Seidell Lowry Bredee Work (18) (14) (4) 69.86 70.73 70.1 69.9 73.58 74.82 74.2 73.7 76.99 77.49 77.8 77.0 80.14 80.81 80.7 83.20 83 32 83.5
.. ..
Hoeg (10) 89.9 73.3 76.8 80.2 83.2
POI
0.7
I 0.8
1
I I I 0.9 1.0 1.i 1.2 Density d Soturoted Solution
J
I
1.3 gjml.
1.4
1.5
Figure 2
POTASSIUM NITRATEIN ETHANOL SOLUTIONS.Data for this same system, from 30" to 60" C., obtained by Gerardin (9) in 1865, show fair agreement with the values presented here, but the older data show some scattering of points at the lower alcohol compositions. The authors' data agree excellently with those of Schreinemakers (17) for 30" C., the only temperature he employed. Some data reported by Bathrick (2) for 30" and 40" C. are higher than others throughout the range of alcohol composition. Smoothed values for solubility and density isotherms are presented in Table IV for the potassium nitrate-ethanol-water system. Pjots of the experimental points for salt content against temperature, with alcohol composition as parameter, are shown in Figure 1. Similar plots for density of saturated solutions are presented in Figure 2. The composition isotherms, from 30" to 70" C., for thi$ system are shown on the triangular diagram in Figure 3. AMMONIUM NITRATE IN ETHANOLSOLCTIONS.The only previous data, covering a wide temperature range, reported for this system were those of Fleckenstein ( 7 ) . His experimental points were sew, and he admitted large deviations a t high alcohol concentrations in the solvent. He noted the presence of two liquid phases under some conditions but gave no specific solubility data in this region. De Wahl (BO) reported solubilities in several aqueous ethanol solutions a t O", 30°,and 70' C., and observed the formation of two liquid phases but omitted solubility measurements in that section. Schreihemakers (17) gave solubilities for several alcohol concentrations, all a t 30" C. In this system, separation into two liquid phases occurs over part of the concentration and temperature range. A plait point (determined from large scale plots similar to Figure 4 and
INDUSTRIAL AND ENGINEERING CHEMISTRY
480
T-1BLE
---KKOa-Ethanol-Tater-Density Wt. 70 KKOs in of satd. Temp., satd. soha, O C. so1n.a g./ml. Water as Solvent 25.00 27.24 1.1864 30.00 31.49 1.2167 35.00 35.40 1.2458 40.00 39.02 1 ,2740 45.00 42.64 1,3059 50.00 46.11 ' 1.3358 55.00 49.28 1,3640 60.00 52.33 1,3905 65.00 55.20 1.4137 70.00 57.87 1,4390 75.00 60.50 1.4629
S M O O T H E D VALUES FOR SOLUBILITY AND
K t . 70 K N 0 8 in Temp., satd. c. soln. E-50as Solvent Q
Density of satd. soin. a , g./ml.
E-70as Solvent 1.1049 1,1285 1,1524 1,1771 1 ,2026 1.23'02 1.2596 1.2892 1.3180 1.3461 1 3738
E-20as Soll-ent
25.00 30.00 35.00 40.00 46.00 50.00 55.00 60.00 65.00 70.00 75.00
E-10as Sol vent 65.45 1,2859 67.46 1.2968 69.40 1.3063 71.34 1.3160 73.24 1.3248 75.10 1.3331 1.3410 76.90 78.68 1,3489 80.41 1,3666 82.12 1.3641 83.81 1.3716
25 00 30 00 35.00 40.00 45.00 50.00 55.00 60.00 65.00 70.00 75.00
E-20as Solvent 60 90 1.2532 63.73 1.2669 66.22 1,2790 68.59 1.2904 70.89 1,3014 73.11 1.3121 75.27 1.3222 77.34 1.3322 1.3418 79.18 1.3513 80.71 1.3608 81.75
E-92as Soivent
26.00 30 00 35.00 40.00 45.00 50.00 55.00 60 00 65.00 io.00 75,oo
1 0469 1 0607 1 0763 1 0938 1 1136 1.136'7 1,1628 1 1910 1 2190 1 2496 1 2735
25.00 30.00 35.00 40.00 45.00 50.00 55.00 60.00 65.00 70.00 75.00
E-30as Solvent 9.02 1.001 10.85 1.009! 12.86 1.0193 15.18 1,0324 17.87 1,0473 21.00 1,0636 24.09 1,0802 27.39 1,0987 30.77 1.1199 34.45 1.1433 38.38 1.1685
DENSITY IbOTHER\iS
--~-HaSOa-Ethanol-~-aterWt. % Density NHaSOa of satd. Temp In satd. soln.Q, 0 c." so1n.a g./ml. Water as Solvent
4.20 5.11 6.07 7.11 8.34 9.80 11.58 13.50 15.60 17.94 20.60
E-10a s Solvent 25.00 30.00 35,OO 40.00 45.00 50.00 55.00 60.00 65.00 io.00 75.00
ITr.
_-_ IiNOa-Ethanol-Water-?
a
I n vacuo.
Vol. 40, No. 3
'
E-30a s Solvent 56.38 1.2110 25.00 59.37 1.2265 30.00 62.25 1.2413 35.00 64.89 1.2546 40.00 67.27 1.2667 45.00 69.57 1.2784 50.00 71.79 1.2908 55.00 73.97 1.3032 60.00 76.11 1.3159 65.00 78.22 1.3286 70.00 80.30 1.3413 75.00
SO.-Ethanol-KatcrDensity Wt. % SHaSOa of satd. Tenip., in satd. solma, g./ml. soh.a E-50as Solvent 25.00 43.50 1.1052 30.00 46.90 1.1222 50.26 1.1380 35.00 53.51 1.1535 40.00 56.55 1.1688 45.00 50.00 59.49 1.1842 62.39 1.1996 55.00 65.20 1.2143 60.00 65.00 67.94 1.2276 70.00 70.66 1.2402 75.00 73.33 1.2524
--"a
c.
25.00 30.00 35.00 40.00 45.00 50.00 55.00 60.00 65.00
E-70as Solvent 26.05 29.16 32.27 35.36 38.48 41.63 44.92 48.39 51.93
60.00 65.00 70.00
E-82as Solvent 29.73 0.9434 32.98 0,9576 36.24 0.9718
60.00 65.00 io.00 75.00
E-88as Solvent 19.98 22.36 24.74 27.12
0.8795 0.8848 0,8902 0.8954
25.00 30,OO 35.00 40.00 46.00 50.00 55.00 60.00 65.00 70.00 75.00
E-92a s Solvent 7.29 8.07 8.88 9.72 10.63 11.65 12.75 13.91 15.24 16.69 18.19
0.8382 0.8377 0.8374 0.8372 0.8372 0.8375 0.8379 0.8385 0.8400 0.8435 0.8407
Figure 5 ) was found to occur at 65.7 ' C. At the plait point the salt content is 54% (determined from the equal distribution point on a plot of per cent ammonium nitrate in the lighter liquid phase against per cent ammonium nitrate in the heavier
0.9758 0.9882 1.0007 1.0134 1.0267 1.0406 1.0561 1.0724 1.0888
liquid phase) and the alcohol content is 33% of the total solution or 73% of the solvent (determined from the equal alcohol point on a plot of the alcohol content of one phase against that of the other). With water as the solvent, it is known that a transition point for the crystalline form of t'he solid phase occurs at. about 32.3 C., with a change in the slope of t'he solubility and density curves. KNO3 30att,empt was made t o determine t,his transition point for any of the solutions.' Experimental points for density and solubility, including those for both 30.00' and 35.OO0C., when plotted on a large scale against temperature were found to lie on a smooth curve. Plots of the experimental points for salt cont,ent against temperature, with alcohol composi.tion as parameter, are shown in Figure 4. Similar plots for density of saturated solutions are presented in Figure 5 . These figures show the region in which two liquid phases are possible. The triangular composition diagram, Figure 6, also indicates the presence of two liquid phases at the higher temperatures. Smoot,hed values for solubilit,y and density isot,hcrms for t,his systcm are given in Table IT'. 20 40 60 00 E T-0 ti Table V preeenk similar data obtained in the W t . Ye E T - O H region where two liquid phases were present, as Figure 3. Composition Isotherms for System of Potassium Nitrate w~~~as comPositions determined by with Ethanol and Water the methods indicated earlier.
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March 1948
INDUSTRIAL AND ENGINEERING CHEMISTRY
I
S O L U B I L I T Y OF N H ~ N OIN~ AQUEOUS ETOH SOLUTIONS
,
80
481
70
6 60
-e
g 50
040 ID
30
I
20 0
IO
20
30 wt.
e&
50
40 NH.,NO~
60
70
80
90,
in Soturoted Solution
REQIONFOR SYSTEX TABLE V. DATAFROM TWO-PHASE NH4NO3-E~~~~~~--H20 Temp.,
C. 66.91 69.01 70.41 71.84 73.98
% 2 satd. soln.
Phase-Density Wt. % of satd. alqohol soln., in g./ml. solvent
----Lower Wt. % salt in satd. soh.
Phase---Density Wt. % of satd. alophol soln., in g./ml. solvent
46.60 42.89 40.53 39.21 37.67
1.0479 1.0180 0,9995 0.9875 0.9700
60.75 64.95 67.29 69.39 72.41
1.1594 1.1972 1.2129 1.2229 1.2341
76.3 78.9 80.4 81.4 82.7
I
I
I
I
1.2 Density of Soturoted Solution 10
1.3
I1
1 1.4
o./mi.
Figure 5
Figure 4
-----Upper
0.9
65.9 60.3 56.8 54.7 51.9
An attempt to explain the points of inflection, occurring at a solvent concentration of approximately 20y0 alcohol, shown in the isotherms, has been made (Figure 6). It was thought that perhaps the inflection was sharp enough to indicate the presence of compound formation or a change in crystal form. The absence of solvation or hydration of the solid phase was proved by gravimetric determinations. Also, x-ray diffraction measurements
NH~NOJ
made by the powder method failed to substantiate the possibility of compound formation or change in crystal form at the points of inflection. X-ray diffraction patterns were obtained for the ammonium nitrate salt in equilibrium with solutions in four different conditions: distilled water at 29.5' and 65.0' C., and 50% aqueous ethanol a t 29.5" and 85.0" C. Patterns from both 29.5" C. samples were identical and indicate the orthorhombic IV form which is that normally stable in the pure dry state at temperatures below 32.3" C., the transition temperature. The 65.0' C. samples also g w e identical patterns indicating the same crystalline formin this case, orthorhombic 111. It was concluded from this evidence that no solvation or transition in the solid phase occurs, at a given temperature, as the ratio of alcohol to water is changed. ACKNOWLEDGMENT.
The authors wish t o thank K. A. Krieger for making the x-ray diffraction measurements on ammonium nitrate. LITERATURE CITED
(1) Armstrong, H. E.,and Eyre, J. V., Proc. Roy. SOC.(London), A, 84,123-36 (1910-11). (2) Bethrick, H. A.,J.Phys. Chem., 1, 15749 (1896). (3) Bruyn, B. R. de, Z.physik. Chem., 32,63-115(1900). (4) Cohen, E.,andBredee, H. L., Ibid., 117,143-55(1925). (5) Copley, E. D., Murray-Rust, D. M., and Hartley, Harold, J. Chem. SOC.,1930,2492-8. (6) Donk, A.D., Chem. Weekblad., 5,529(1908). 6,419-22(1905). (7) Fleckenstein, A.,Physik. Z., 67,179-82(1945). (8) Gee, E.A., J . Am. Chem. SOC., (9) Gerardin, M. A.,Ann. chim. et phys., [4]5, 129-60 (1865). (10) Hoeg, F. M.A., 2. anal. Chem., 81,114-16(1930). (11) International Critical Tables, Vol. 4, p. 240, New York, MCGraw-Hill Book Co., 1928. (12) Ketner, C. H.,2.physik. Chem., 39,641-90 (1901). and Bjerrum, J., Ber., 64,210(1931). (13) Lund, H., (14) Millioan, I. L., Joseph, A. F., and Lowry, T. M., J . C h m . SOC., 121,959-63 (1922). (15) Mulder, G. J., "Scheikundige Verhandlingen en Onderzoekingen," Vol. 3, Part 2, p. 89, Rotterdam. 1864.
