*
171
INDUSTRIAL AND ENGINEERING CHEMISTRY
January 1950
per square foot. At low flow rates, the 0.25-inch saddles resulted in about 10% higher nicotine extraction than O.$inch saddles, but a t higher flow rates, extraction was practically the same with either size. Extraction was essentially constant over the range of 0.1 to 0.5% nicotine in water as the feed stock. Extraction was increased about twofold by tripling the ratio of kerosene t o water. When high flow rates were used, extraction was increased about 16% by dispersing the kerosene rather than the water; a t low flow rates, however, extraction was the same with either phase dispersed. The extraction capacity coefficient a t 63' C., based either on the water or kerosene phase, increased with increasing flow rates, but the coefficient based on the water phase was 3.5 times that based on the kerosene side-that is, kerosene offered the major resistance in this system. The coefficient based on the water phase increased uniformly with temperature, increasing sevenfold over the temperature range 21 to 84" C. , whereas the coefficient based on the kerosene side remained constant over this temperature range. The height of the transfer unit varied with temperature, flow rate, and other operating conditions but covered the range of 1.1 to 6.9 feet when based on the water side, and from 2.8 t o 11.4 feet when based on the kerosene side. O
ACKNOWLEDGMENT
c. o. Willits and associates Of this laborator ~ ~ r i ~ ~ $ $ , k ~ . ~ ~ y ~ i d~ $ & ~~ ~ ~~ ~$ ' c w~ ~ ~~ ~ Salem, N. c., for their generous cooperation in supplying inThanks are due
formation relative to the industrial extraction of nicotine from tobacco. LITERATURE CITED
(1) Aitken, R.G., Can. Patent 401,143(Dec. 2,1941). (2) Chem. & Met. Eng., 42,373(1935). (3) Elgin, J. C.,and Browning, F. M., Trans. Am. Inst. Chern. Engrs., 31,639 (1935). (4) Fritzsche, R.,Syracuse Chemist, 35,3 (1942). (5) Johnson, H. F., and Bliss, H., Trans. Am. Inst. Chem. Eng., 42, 331 (1946). (6) Kingsbury, A. W., Mindler, A. B., and Gilwood, M . E., Chem. Eng. Progress, 44,497 (1948). (7) McConnell, H . K.,U.8. Patent 1,927,180(June 28,1924). (8) Norton, L.B.,IND. ENG.CHEM.,32,241 (1940). (9) Sata, Naoyasu, BUZZ.Chem. SOC.Japan, 2,139 (1927). (10) Sherwood, T.K.,Evans, J. E., Longcor, J. V. A., IND.ENG. CHEM., 31, 1144 (1939). (11) Skalweit, J. J., Rer., 14,1809 (1881). ENG.CHEM.,34,251 (1942). (12) Smith, C.R.,IND. (13) Willits, C. O.,et al., E a s t e r n Regional Research L a b o r a t o r y ,
private communication.
RECEIVED June 7, 1949. Report of a study made, in part, under the Research and Marketing Act of 1946.
Solubility and Density Isotherms SODIUM SULFATE-ETHY L ALCOHOL-WATER RAYMOND E. VENERl AND A. RALPH THOMPSON University of Pennsylvania, Philadelphia, Pa. Solubilities and saturated solution densities are reported at intervals of 5" from 25" to 75" C. for sodium SUIfate in aqueous ethyl alcohol solutions. Only one liquid phase was found to occur with the stable solid phases studied. The lowering by the ethyl alcohol of the sodium sulfate anhydrous-decahydrate transition temperature in saturated aqueous solutions was determined from 32.4' to just below 20' C. With solvents containing more than % ' the so'ubility Of anhydrous 'Odium sulfate increases with temperature.
T H E solubility of sodium sulfate in aqueous ethyl alcohol solu-
li
tions was required to permit a thorough small-scale study of the crystallization of the salt from aqueous solutions by saltingout with alcohol using a general technique suggested by Thompson and Molstad (7). Very few data are reported for this ternary system for any appreciable temperature or solvent range. The data of de Bruyn ( I ) , Flatt and Jordon (d), Schiff (4), and Schreinemakers and de Baat ( 5 ) are cited or reproduced by Seidell (6). Some values for the solubilities of the anhydrous, heptahydrate, and decahydrate salt were reported by de Bruyn for 15", 25", 36O, and 45' C. Practically all of the work by the other authors was restricted t o 25' C. The materials used in these studies were Baker's C.P. anhydrous sodium sulfate, commercial grade (95%) ethyl alcohol from Publicker Industries, Inc., and distilled water from a Barnstead still. Tests to show the adequate purity of these grades of salt and water were described in an earlier paper (8). Commercial 1
Present address, Drexel Institute of Technology, Philadelphia, Pa.
grade ethyl alcohol was deemed suitable for this work when it was found that salt solubilities in test solutions were the same as those obtained with solutions of identical composition prepared from highly purified ethyl alcohol. EXPERIMENTAL
The equipment and techniques employed for this system were essentially those used for the sodium sulfate-ethylene glycol-water system (9). Sodium sulfate was equilibrated with a series of aqueous ethyl alcohol solutions over a temperature range of 25' to 75" c. Densities of the saturated solutions were determined with a Precision of *0.0003. gram per ml. The salt solubility values were ascertained with a precision of *0.01 weight ?6 salt for solubilities of over 10% sodium sulfate and considerably better for
OF AQUEOUSETHYLALCOHOL TABLEI. ANALYSES SOLVENTS
(In vacuo) Solvent No. E-3 E-6 E-10 E-15 E-20 E-25 E-30 E-40 E-50 E-60
E-80
Density a t 35.3
c
G./Mi: 0.98847 0.98234 0.97029 0.96951 0.96154 0.95315 0 * 94443 0.92297 0.90150 0.87940 0.83111
Ethyl Alcohol in Solvent Wt. % Mole % 3.01 1.20 6.62 10.35 14.74 19.88 24.95 29.79 40.41 50.08 59.62 79.07
2.70 4.32 6.33 8.84 11.50 14.23 20.95 28.17 36.59 60.50
Vol. 42, No. I
INDUSTRIAL AND ENGINEERING CHEMISTRY
I72
11. AQUEOUSETHYLALCOHOLSOLUTIOSS IK EQI-II.TBRIUM WITH SOLIDPHASE OF BOTH SODIUM SULFATE AND SODIUM SULFATE DECAHYDRATE
TABLE
(At transition points) Transition Density of Saturated Solution, Tempenature. G./hll. = c.
Wt. % Ethyl Alcohol in Solvent
Wt. ? ' & NarSOr in Saturated Solution
EXPERIMENTAL POINTS
32.38 31.60 31.01 30.65 30.00 27.44 25.17 19.97
0.00 3.09 11.33 25.34 33.17 45.85 52.25 56.56
33.27 29.38 20.14 7.247 3.366 0.943 0.409 0.194
1.3324 1.2808 1.1626 1.0139 0.9671 0.9257 0.9071 0.9009
utilizing the density-composition data of Osborne, blcKelvy, and Bearce (5)as described previously (7, 8). Table I contains the densities of these solvents at 35" C. and their corresponding compositions. TRANSITION POINTS.The transition line between regions of anhydrous and decahydrate salt was determined from 32.4' (water as solvent) to 20" C. Experimental points are plotted in Figure 1 with transition temperature plotted against solvent composition and salt content, respectively. The values in Table I1 were interpolated from the experimental points at solvent compositions corresponding to those in Table I. SOLUBILITIES AND DESSITIES. Smoothed values for solubility and density isotherms are presented in Table I11 at 5" intervalfi from 25' to 75 C. They were obtained from experimental valO
XNTERPOLATBD VALUES
lower salt content values. The evaporations for salt content determinations were carried out in an oven at atmospheric pressure. These studies were restricted to stable solid phases-namely, the anhydrous and the decahydrated salt. Although the heptahydrate was detected occasionally below 25" C., it was unstable in the presence of small seed crystals of the decahydrate. The analyses of the aqueous ethyl alcohol solvents were run by Wl
wt.
X
% Ethonol in Solvent
No2S04
Figure 3
i n S o f u r o l e d Solution
Figure 1. Variation of Sodium Sulfate-Sodium Sulfate Decahydrate Transition Temperature with Aqueous Ethyl Alcohol Solvent Composition and Salt Content
'0
4
8
12 16 20 24 28 W t % N o 2 5 0 4 in S o l u r o t e d Solution
Figure 4
80 75 70 65 60 i)
*
55
2 50
p 45
;4 0 ly
2
35
30 25
0
4 Wt
8 9.
12 16 20 N O ~ S O In ~ Solurolsd
Figure 2
24 28 Solution
32
Density of
Saturated
Figure 5
Solution
-
g./ml,
32
36
INDUSTRIAL AND ENGINEERING CHEMISTRY
January 1950
TABLE 111.
3.01WT. % '
ETHYL ALCOHOL I N SOLVENT
18.50 26.74 29.10 28.69 28.29 27.91 27.60 27.34 27.12 26.90 26.70
40.00
45.00 50.00 55.00 GO.00 65.00 70.00 76.00
6.62WT. Yu ETHYL ALCOHOL 15.38 22.87 24.83 24.44 24.10 23.80 23.55 60.00 23.33 65.00 23.15 70.00 22.98 22.81 75.00
1.1684 1.2535 1.2752 1.2674 1.2601 1,2530 1,2465 1.2406 1.2349 1.2291 1.2235
60.00
65.00 70.00 75.00 25.OOa
30.00"
35.00 40.00 45.00 50.00 55.00 60.00
65.00 70 no 75.00
10.35WT. To E T H Y L ALCOHOL 11.97 19.05 20.65 20.31 20.03 19.79 19.59 19.43 19.29 19.17 19.06
19.88W t . 25.ODa 30.00" 35.00 40.00 45.00 50.00 55.00 60.00 65.00 70.00 75.00 25.OOa 30,OOa 35.00 40.00 45.00 60.00
55.00
60.00
65.00 70.00 75.00
GO.00
65.00 70.00 75.00
1.1258 1.2006 1,2177 1.2125 1.2060 1,1998 1,1941 1.1885 1.1834 1.1783 1.1733
0.9710 0,9834 0,9820 0.9803 0.9784 0,9760 0,9733 0.9704 0.9677 0.9635 0.9595
7 0 ETHYL ALCOHOL I N SOLVENT
50.08WT.
7 0 ETHYL ALCOHOL I N SOLVENT
* 1.0867
1.1632 1,1669 1.1606 1,1548 1.1495 1.1447 1.1400 1.1353 1.1307 1.1261
25.005 30.00 35.00 40.00 45.00 50.00 55.00 60.00 65.00 70.00 75.00
1.0490 1.1059 1,1115 1.1067 1.1020 1.0974 1.0930 1.0888 1.0848 1,0810 1.0773
25.00 30.00 35.00 40.00 45.00 50.00 55.00
1,0168 1.0553 1.0678 1.0541 1.0507 1.0473 1.0439 1,0405 1.0371 1.0337 1.0303
25.00 30.00 35.00 40.00 45.00 50.00 55.00
0.9412 0,9374 0.9355 0.9330 0.9299 0.9268 0.9235 0.9200 0.9160 0.9117 0,9072
1.155 1.615 1.741 1.854 1,969 2.070 2.160 2,237 2,303 2,360 2.413
0.9143 0.9098 0.9061 0.9024 0.8986 0.8947 0.8905 0.8862 0.8818 0.8772 0.8727
0.510 0.593 0,637 0.680 0.722 0,760 0.798 0.832 0.862 0.886 0.907
SOLVENT
60.00
65.00 70.00 75.00
7 0 ETHYL ALCOHOL I N SOLVENT
8.00 10.52 11.36 11.34 11.33 11.32 11.32 11.33 11.34 11.34 11.36
24.95 WT. Yo ETHYL ALCOHOL 3.902 7.130 7.521 7.678 7.818 7.930 8.019 8,086 8.138 8 182 8.227
I S SOLVENT
40.41WT. 25.00" 30.00 35.00 40,OO 45.00 50.00 55.00 GO.00 65.00 70.00 75.00
I N SOLVENT
14.74WT. YUETHYL ALCOHOL I N 8.56 14.79 16.04 15.85 15.69 15.54 15.42 15.31 15.25 15.19 15.15
29.79WT. 7 0 ETHYL ALCOHOL 2.672 4.513 4.822 5.032 5 229 5.387 5.519 5 626 5.704 5.761 5.817
Density of Saturated Solution, G./MI.
25.00" 30.OOa 35.00 40.00 45.00 50.00 55.00
I N SOLVENT
25.OOa 30.005 35.00 40.00 45.00 50.00 55.00
25.o o a 30.OOa 35.00 40.00 45.00 50.00 55.00
VALUESFOR SOLUBILITY AND DENSITY ISOTHERMS
(Solid phase is anhydrous sodium sulfate unless otherwise indicated) Wt. % NazS04 Wt. Yo NrtzsO4 Density of in Saturated Temperature, in Saturated Saturated Solution, So1uti on c. Solution G./M1.
Temperature, "C.
25.OOa 30.OOa 35.00
SMOOTHED
173
ues by an interpolative method ( 7 ) . Anhydrous sodium sulfate was the stable solid phase for all values listed with the exception of those marked by the superscript, a, for which cases the decahydrate was the solid phase. Experimental points for sodium sulfate solubility, with solvent composition as parameter, are plotted against temperature in Figure 2, and against saturated solution density in Figure 3 . The solvent compositions corresponding to the various curves may be identified by referring t o Table I. The curves for water as solvent were obtained from data reported in a previous paper (9). The sharp break points in most of the curves represent transition points with both anhydrous and decahydrate salt as the solid
I N SOLVEXT
79.67WT. T ' o ETHYL ALCOHOL 0.002 0.004
I N SOLVENT
60.00
65.00 70.00 75.00
I N SOLVENT
0.9907 1.0149 1.0144 1.0121 1.0097 1.0069 1.0038 1.0007 0.9977 0.9945 0.9916
59.62WT. yo E T H Y L ALCOHOL 0.170 0.187 0.201 0.216 0,230 0.241 0.252 0.262 0.269 0.277 0.286
a
0.006
0.008 0.0011 0,0013 0.0015 0.0017 0,0018 0.0019 0.0021
0.8878 0.8842 0.8806 0.8768 0.8726 0.8682 0.8638 0.8592 0.8546 0.8499 0.8453 0.8399 0.8356 0.8312 0.8266 0.8220 0.8175 0.8129 0.8083 0.8036 0.7990 0.7942
Solid phase NazSO~.lOHzO.
phase. The open circle points above the transition points indicate anhydrous sodium sulfate as the solid phase; at temperatures below the transition points the solid phase is the decahydrate. Several isotherms are presented on rectangular coordinates with solvent composition plotted against salt solubility in Figure 4 and against saturated solution density in Figure 5 . The 35" isotherm on the former and the 25" isotherm on the latter are indicated by broken lines merely for greater clarity. Solubility isotherms between 35" and 75" with the anhydrous salt as the stable solid phase all cross at the same point as shown in Figure 4. From an examination of Figure 2 it is apparent that with over 20 weight ethyl alcohol as solvent the salt solubility increases with temperature. Exploratory runs on the system, sodium sulfate-methanol-water, indicated very similar results with regard t o both the depression of the transition temperature and the reversed slope of the solubility curves. Additional tests using either acetone or isopropyl alcohol as the organic constituent re-
I N D U S T R I A L A N D E N G I N E E R I N G CHEMISTRY
174
sulted in two liquid phases over most of the solvent composition range. LITERATURE CITED
(1) B r u y n , B. R. de, 2. physik. Chem., 32,63-115 (1900). (2) F l a t t , R., and Jordon, A., Helv. Chim.Acta, 16, 37-53 (1933). (3) Osborne, N. S.,McKelvy, E. C . , and Bearce, H. W., Bull. Bur. Standards, 9, 327-474 (1913). (4) Schiff, H., Ann. Chem., 106, 108-18 (1858); 118,362-72 (1861). ( 5 ) Schreinemakers, F. A . H., and B a a t , W. C. de, 2. physik. Chem., 67, 551-60 (1909).
Vol. 42, No. 1
(6) Seidell, il., "Solubilities of Inorganic and M e t a l Organic Compounds," 3rd ed., Vol. I, p p . 1300-02, 1312-14, New Y o r k , D. Van Nostrand Co., 1940. (7) Thompson, A. R., a n d Molstad, M. C., IND.ENG.CHEM?., 37, 1244-8 (1945).
(8) Thompson, A. R., a n d Vener, R. E . ,I b i d . , 40, 478-81 (1948). (9) Vener, R. E., and Thompson, A . R., I b i d . , 41, 2242-7 (1949). RECEIVED June 28, 1949. Based on a dissertation presented by Raymond E. Vener to the Graduate School, University of Pennsylvania, in partial fulfillment of the requirements for the degree of doctor of philosophy.
erature on active In S. S. KURTZ, JR., SENTA A330N, AND ALBERT SANIIIN, S u n Oil
In order to emphasize the fact that the Eyliman equation has a wide range of usefulness, the authors have summarized available data, including the data of Eyliman, showing that the Eykman equation represents experimental data for the effect of temperature on the density and refractive index of organic compounds quite accurately over a wide range of temperature. PAPER published recently by Dreisbach ( 4 )shows that over the limited range, 20' to 25" C., the Eykman equation (7, 8)
more accurately represents experimental data for the effect of temperature on the density and refractive index of organic compounds than does the Lorentz-Lorenz equation (3, I S , 14,15)
The conclusions reached by Dreisbach are entirely in accord with the previous conclusions of Kurtz and Lipkin ( l l ) ,Ward and Kurtz ( l 7 ) ,Bauer and Fajans ( d ) , and Gibson and Kincaid (9). The purpose of this paper is to summarize available data, including the original data of Eykman, showing that the Eykman equation represents experimental data quite accurately over a wide range of temperature. I n Table I Eykman's experimental data for 42 liquid hydrocarbons are compared with values calculated from the Eykman equation, the Lorentz-Lorenz equation, the Gladstone and Dale n-1 equation (IO), - CB, and a simple empirical approximation, ~
An = 0.6Ad, recommended by Ward and Kurtz (27, 18) for correcting the refractive index of hydrocarbons for small changes in temperature. The minimum temperature interval in this table is 31.0" and the maximum, 122.3" C. Table I shows that the observed change in refractive index agrees with the calculated change in refractive index with an average deviation of *3 x per degree if the Eykman equation is used. The simple equation, An = 0.6 Ad, agrees with the data with an average deviation of 1 8 X lo+, and the Gladstone and Dale equation with an average deviation of -26 X while the Lorenta-Lorenz equation shows an average deviation of +38 X 10-8 per degree change in temperature. In the tables of Dreisbach's paper the most reliable data are quoted from the tables of A.P.I. Project 44 ( I ) for seven
Company, Norwood, Pa.
hydrocarbonsTnamely, hexane, heptane, benzene, toluene, ethylbenzene, isopropylbenzene, and styrene. If one calculates the refractive index a t 25" C. from the refractive index at 20" C. using the constant C1, the average deviation for these 7 compounds between the experimental and calculated refractive index is * 5 X 10-5 for the 5 ' intervals, or =t1 X for a 1O C. temperature change. The average deviation per degree in Table I is, therefore, only one third of the deviations indicated by the limited temperature interval considered by Dreisbach. Table I1 presents hydrocarbon data from the literature other than Eykman data for the effect of temperature on density and refractive index, but considers only data which cover a temperature range of 15" C. or more. (The A.P.I. Project 44 tabulations a t present cover only 5" C. trmperature intervals and are omitted for that reason.) These data do not agree with the Eykman equation as well as do Evkman's original data. The deviations per degree are no worse, however, than are those shown by Dreisbach. The deviations from A n = 0.6Ad are still smaller than those from the Gladstone and Dale and the Lorentz-Lorenz equations; although the diff rrence between equations is much less than it is for Eykman's hydrocarbon data. Table I11 presents corresponding data obtained by Eykman for nonhydrocarbons, This table also shows excellent agreement between experimental data and values calculated from the Ekyman equation. These data are of interest in connection with the theory of refraction, because the regularity of associated liquids in regard to the Eykman function must mean that association between molecules has little effect on refraction. This is not surprising if one recognizes that refraction in the visible part of , the spectrum for hydrocarbons and other simple organic compounds depends upon the valence electrons (6, 11, l a ) . The frequency of vibration of these electrons is probably not appreciably modified by association effects between molecules. The failure of An = O.6Ad to represent these data is not surprising as this simple equation was derived for hydrocarbons. Consideration of the currently available data for the change of refractive index and density with temperature leads one to conclude that changes in refractive index calculated by the Eykman equation from experimental data for density a t two or more temperatures are likely to be more accurate than experimentally determined changes in refractive index unless the experimental work is very carefully done. The Eykman equation is entirely empirical; however, it has been shown by Kurtz and Ward ( l b , l 7 ) that the Sellmeier-Drude dispersion equation (6) can be modified by adding one constant SO that i t will give data in quantitative agrecment with the Eykman equation.
