624
INDUSTRIAL AND ENGINEERING CHEMISTRY
because hydrogen sulfide is the component of t.he greater oommercial interest, this is actually of little practical consequence. ACKNOWLEDGMENT
Thanks are due to the U. S. Air Force for the loan of the labor&tory space and some of the equipment used in this work.
Gilliland, E. R., and Scheeline, H. W., IND.ENQ.C H ~ M . ,32, 48 (1940).
( 5 ) "International Critical Tables," Vol. 111, p. 235, Iiew York, McGraw-Hill Book Co., 1928. (6) Kay, W. B., IND. ENG.CHEM.,30, 459 (1938). (7)Kay, W. B.,and Brice, D. B., I h i d . , 45, 615 (1953). (8) Kay, W. B., and Rambosek, G. M., I b i d . , p. 221. (9) Reamer, H. H., Sage, B. H., and Lacey, W.K.,I h i d . , 42, 140 (1950).
LITERATURE CITED
(1) Bridgeman, 0. C., J . Am. C'hem. SOC.,49, 1174 (1927). (2) Cooper, D. L., and Maass, O., Can. J . Reseaich, 4, 283 (1931). (3)
(4)
Vol. 45, No. 3
Dodge, B. F.,"Chemical Engineering Thermodynamics," 121, S e w Pork, McGraw-Hill Book Co., 1944.
( 1 0 ) I h i d . , 43, 976 (1951). (11) Steckel, F., Sz;ensk Kern. T i d s k i . , 57, 209 (1945). (12) Wright, R. H., and Maass, O., Can. J . Research, 5, 442 (1931).
p.
RECEIVED for review September 9, 1962.
ACCEPTED October 13, 1952.
Vapor-Liquid Equilibria for Ethyl Alcohol inarv Svstems ~~
d
J
LAYGLEE' Rf HELLWIG AND MATTHEW VAN WISKLE The University of Texas, Austin, Tex. APOR-liquid equilibria a t '760 mm. of mercury are reported for the systems: ethyl alcohol-n-butanol, ethyl alcohol-secbutanol, ethyl alcohol-n-pentanol, ethvl alcohol-acetone, ethyl alcohol-methyl ethyl ketone, and ethyl alcohol-methyl n-propyl ketone. The ethyl alcohol-alcohol s>stems exhibited no azeotropes, and the activity coefficients calculated from the experimental data show these systems deviate only slightly from ideal solution behavior. The ethyl alcohol-n-butanol data compare a ell v,ith those reported by Brunjes and Bogart ( I ) . The ethyl alcohol-ketone systems showed the follox~ing azcotropes: ethyl alcohol-acetone: none; ethyl alcohol-methyl ethyl ketone: 50.1 mole % ethvl alcohol a t 165.2" F. (74.0" C . ) ; ethyl alcohol-methyl n-propyl ketone: 96.2 mole % ethyl alcohol a t 172.5' F. (78.0" (2,). The ethyl alcohol-acetone data check those given by Perry ( 8 ) . ETHYL ALCOHOL-ALCOHOL SYSTEMS AT 760 O F MERCURY
D s ~ a . Of the systems inveqtigated, only the ethyl alcoholn-butanol system ( 1 ) has been previously reported in the literature ( 4 ) . The experimental data \?ere obtained a t a constant pressure of 760 mm. of mercury for the binary systerns existing in t u o-phase vapor-liquid equilibria. Figures 1 and 2 and Tables TI, 111, and I V show the results. The compositions of the liquid and vapor phases were obtained by refractive index measurements, and equilibrium temperatures were determined from the thermocouple measurements. The maximum experimental error is believed t o be 0.6 % mole fraction based on the 50y0 (composition point)-i.e., an absolute error of k0.003 mole fraction-and r t O . l 0 F. Activity coefficients from the experimental data were calculated by Equation 1 and are ehown in Tables 11, 111, and IT'.
nm1.
;~IATERIALS. The source and physical constants of the purified materials are presented in Table I. Absolute ethyl alcohol as received was used without purification. The n-butanol and sec-butanol samples were purified as outlined by Hill (Z), by taking t h e middle cut of 500 ml. from a charge of 1500 ml. The n-pentanol was the middle 300-ml. cut froin a n 800-ml. charge. All the materials employed had a boiling range of less than 0.1 C. The alcohols contacted only a dry atmosphere during storage, since suitable desiccating agents were provided for both air and t h e liquid samples. During the experimental work, frequent checks were made on the refractive indices of the compounds t o detect possible contamination. APPARATUS. The method of analysis was the same as that used by Hill ( 2 ) . The equilibrium still thermocouple vias an iron and constantan couple calibrated by a Cottrell boiling point apparatus operated a t 760 mm. of mercury. Boiling temperatures were cross-checked with a set of National Bureau of Standards thermometers, The still system was provided with connections t o a course of filtered air t o maintain the total pressure on the system at 760 mm. of mercury. The vapor-liquid equilibrium still arrangement was the same as that employed by Hill ( 2 ) with the exception that n-tetradecane was used as a manometer fluid. PROCEDURE. Samples xere prepared by the addition of ethyl alcohol t o a known volume of the second alcohol in the binary system so that a total volume of 25 ml. was obtained. The operating and analytical procedures were those reported by Hill ( 2 ) . It was necessary t o employ a special stopcock greasc composed of silicone grease and carnauba wax t o prevent leakage when the ethyl alcohol-.see-butanol system was investigated.
MOLE PERCENT ETHANOL IN LIQU1D.X
Figure 1.
Vapor-Liquid Equilibrium Diagram for Ethyl Alcohol Systems at 760 Mm.
INDUSTRIAL AND ENGINEERING CHEMISTRY
March 1953
625
TABLE I. PHYSICAL DATAFOR COMPOUNDS Refractive Index, n”no Exgtl. Lit. ( 9 ) 1.361514 1.36155 1.399050 1.39931 1.397130 1.397 1.41 1285 1.40994 1.358490 1,38880 1.378765 1.37850 1.391107 1.38946
Normal Boiling Point, O C. Exptl. Lit. (8) 78.4 78.4 117.1 117.5 99.3 99.5 137.5 137.8 56.2 56.2 79.2 79.6 102.8 108.3
Source U. S. Industry Chemical Co. J. T. Baker Chemical Co. Eastman Kodak Co. J. T. Baker Chemical Co. Commercial Commercial Mathieson Chemical Co.
Compound Ethyl alcohol n-Butanol sec-Butanol n-Pentanol Acetone Methyl ethyl ketone Methyl n-propyl ketone
VAPOR-LIQUID EQUILIBRIA AND ACTIVITY COEFFICIENT DATAFOR TABLE 11. EXPERINENTAL ETHYL ALCOHOL-?&-BUTANOL SYSTEMAT 760 MM. O F MERCURY EtOH, Mole % 96.8 88.0
EtOH, Wt. % 95.0 81.9
Y
YLOW
YHigt
Y
YLOW
YRigh
177.4
Phase Vapor Liquid
1.00
0.99
1.01
1.11
1.06
1 36
2
184.7
Vapor Liquid
91.2 70.9
86.5 60.2
0.99
0.99
1.01
1.14
1.09
i . 19
3
189.8
Vapor Liquid
87.1 61.0
80.7 49.3
0.99
0.97
1.00
1.11
1.06
1.14.
198.3
Vapor Liquid
78.6 46.8
69.5 35.3
0.97
0.96
0.99
1.09
1.06
1.11
211.3
Vapor Liquid
60.2 28.2
50.7 19.6
0.94
0.93
0.96
1.11
1.09
1.12
225.7
S’apor Liquid
38.4 12.9
27.9 8.45
0.99
0.96
1.02
1.02
1.02
1.03
199.3
Vapor Liquid
77.9 45.3
68.7 34.0
0.97
0.96
0.99
1.07
1.04
1.09
215.8
Vapor Liquid
56.3 23.3
44.5
15.9
0.98
0.96
1.00
1.03
1.01
1.04
204.4
Vapor Liquid
69.6 35.7
...
...
...
0.98
1.00
...
1.09
1.13
Run
Temp., F.
1
Ethyl Alcohol
THE
n-Butanol
EQUILIBRIA AND ACTIVITY COEFFICIENT DATAFOR THE TABLE 111. EXPERIMENTAL VAPOR-LIQUID ETHYL ALCOHOL-sec-BUTANOL SYSTEM A T 760 M M , O F MERCURY Temp., EtOH, EtOH, Ethyl Alcohol sec-Butanol
Q
Run
’ F.
1
174.0
Phase Vapor Liquid
Mole % 96.6 92.4
Wt. % 94.6 88.3
2
177.4
Vapor Liquid
91.2 82.5
3
185.8
Vapor Liquid
4
179.7
5 6
Y
YLOW
YHigh
Y
YLOW
?High
1,02~
1.02
1.03
1.57
1.39
1.79
86.5 74.5
1.00
0.99
1.01
1.57
1.48
1.66
74.3 57.0
64.2 45.2
0.986
0.97
1.00
1,38
1.35
1.41
Vapor Liquid
87.0 75.3
80.6 66.4
0.99,
0.99
1.00
1.31
1.45
1.57
177.7
Vapor Liquid
90.0 81.7
86.1 73.5
0.99
0.98
1.00
1.68
1.61
1.77
190.2
Vapor Liquid
63.6 44.2
51.7 33.0
0.98
0.97
1.00
1.30
1.28
1.33
7
195.2
Vapor Liquid
50.5 32.1
38.8 22.7
0.97
0.95
0.99
1.21
1.20
1.23
8
207.5
Vapor Liquid
12.6 5.52
1.09
1.01
1.18
1.02
1.01
1.03
8.20 3.50
The vapor phase was assumed to be ideal; such an assumption is believed to incur no errors greater than those inherent in possible experimental deviations in procedure or conditions. All vapor pressuredata were obtained from the literature [Stull (S)] and represented smoothed points which give straight lines on a semilogarithmic scale (log P us. l / T ) . Because it was necessary in some cases to extrapolate these data from 760 to approximately 4500 mm., possibly some error was introduced, and this would be reflected in the calculated values of the activity coefficients. Activity coefficients plotted versus the liquid concentration of ethyl alcohol, Figure 3, did not agree with those predicted by any of the equations, such as Margules, Duhem, or van Laar, which are usually employed for such data. A method of treating activity coefficient-concentration data has been presented by Li and Coull(6), but this method could not
be employed since i t does not allow the existence of activity coefficients less than unity-i.e., negative logarithms. Steinhauser and White ( 7 ) employed a method of calculating “limit” curves t o bound their experimentally obtained activity coefficient-composition data. I n the present work, it is estimated t h a t possible experimental errors could be represented as =l=0.003 mole fraction and f0.1’ F. On this basis limit curves were calculated from Equations 2 and 3. ?high
W O W
Y + T
= __ 5
- P-
I/-=
= -
x+P+
(3)
Figure 3 shows the experimental activity coefficient-composition curves bounded by the maxima and minima curves. Since
INDUSTRIAL A N D E N G I N E E R I N G C H E M I S T R Y
626
ure 2, are normal for the ethyl alcohol-n-butanol and ethyl alcohol-sec-butanol systems. The slight inflection in the temperature-vapor composition curve for the ethyl alcohol-n-pentanol system a t approximately 15 mole % ethyl alcohol in the liquid is based upon only one experimental point. However, the activity coefficient-composition curve for ethyl alcohol in this system shows a marked deviation from ideal solution behavior in the same region. The ethyl alcohol-n-butanol system has been previously report,ed by Brunjefi and Bogart ( I ) , and the data of this investigation closely check their results. They did not calculate activity coefficients for the ethyl alcoholn-butanol system, since they considered it to be ideal. With reference to Figure 3, the activity coefficient-composition data for the ethyl alcohol-n-butanol indicate the ethyl alcohol to
THANOL-sec-BUTANOL
ETHANOL-n-BUTANOL
tx
Vol. 45, No. 3
lAi
a 3
t0: P W
E + ETHANOL-n-PENTANOL I
'
20
30
40
50
60
70
00
90
Id0
I
I
I
I
I
I
I
I
1
I
!-
z g 2.01
I'O
W
0 MOLE PERCENT ETHANOL
Figure 2.
Equilibrium Boiling Point Diagram for Ethyl Alcohol Systems n-BUTANOL
1.3 1.1
small errors in experimental work, especially in the dilute regions, can cause large variations in calculated activity coefficients, it is
TABLE Iv. Run
Temp., F.
I
I >---'--"--
I-
' -W /
_---l---;F
l - - - __- _ - -_- -_- -___-I_ - --l - --
.
e l
-
I
EXPERIUESTSL VAPOR-LIQUID EQUILIBRIA A N D hCTIVITY COEFFICIENT ETHYL ALCOHOL-n-PENTANOL SYSTEJf AT 760 h h f . OF KIERCURY EtOH, Ethyl Aloohnl EtOH, Phase Mole % Wt. % Y 7Law Y H igh
D.LT.IFOR
THE
n-Pentanol Y
YLOW
7 8 ish
176.6
Vapor Liquid
98.8 91.7
97.8 86 3
1.00
0.98
1 .oo
1.35
1.00
1.80
180.2
Vapor Liquid
97.8 82.6
95.3 71.3
1.01
1.00
1.01
1.27
1.05
1.40
3
195.0
Vapor Liquid
91.6 56.8
85, I 40.7
1.00
0.99
1.01
1.18
1.13
1.23
4
215.2
Vapnr Liquid
79.7 32.5
67.2 20.1
1.00
0.99
1.02
1.11
1.11
1.16
5
234.2
Vapor Liquid
64.4 18.3
48.6 10.5
0.99
0.97
1.02
1.06
1.07
1.10
6
266.4
Vapor Liquid
22.4 2.83
13.1 1.50
1.26
1.11
1.42
1.03
1.02
1.04
7
247.1
Vapor Liquid
49.9 11.2
34.2 6.18
0.99
0.96
1.03
i.oe
1.06
1.08
8
270.3
Vapor Liquid
12.5 1.52
6.96 0.80
1.22
0.99
1.55
1.06
1.05
1.07
10
210.7
Vapor Liquid
79.8 32.4
67.3 20.0
0.93
0.96
0.99
1.10
1.07
1.12
11
207.7
Vapor Liquid
85.2 41.0
75,l 20.6
0.99
0.97
1.01
1.13
1.10
1.16
1 2
March 1953
627
INDUSTRIAL AND ENGINEERING CHEMISTRY.
TABLE V. EXPERIMENTAL VAPOR-LIQUIDEQUILIBRIA AND ACTIVITYCOEFFICIENT DATAFOR ETHYL ALCOHOGACETONE SYSTEMAT 760 MM. OF MERCURY Temp., RUA
O
F.
.
Phase Vapor Liquid
EtOH. Mole %
EtOH.
THE
Acebone
Ethyl Alcohol Y
YLOW
YHigh
Y
rrmw
YHth
8.10 12.5
Wt. % 6.5 10.2
1.63
1.53
1.74
1.03
1.01
1.04
3
134.9
4
137.6
Vapor Liquid
17.3 26.4
14.2 22.2
1.54
1.49
1.59
1.04
1.03
1.05
6
141.2
Vapor Liquid
27.4 42.0
23.0 36.5
1.40
1.37
1.43
1.08
1.07
1.09
7
145.2
Vapor Liquid
36.4 55.6
31.2 49.8
1.27
1.25
1.29
1.15
1.13
1.17
9
149.4
Vapor Liquid
44.5 66.1
38.9 60.7
1.18
1.16
1.19
1.21
1.19
1.23
10
152.1
Vapor Liquid
50.6 72.4
48.3 67.5
1.15
1.13
1.16
1.27
1.24
1.29
11
158.2
Vapor Liquid
62.4 82.5
56.8 78.9
1.08
1.06
1.08
1.36
1.31
1.39
12
162.0
Vapor 1.iquid
70.3 87.9
65.6 85.2
1.04
1.02
1.05
1.43
1.39
1.50
14
167.1
Vapor Liquid
82.7 94.2
79.1 92.8
1.01
1.00
1.02
1.61
1.49
1.72
TABLE VI.
EXPERIMENTAL VAPOR-LIQUID EQUILIBRIA AND ACTIVITY COEFFICIENT DATAFOR ETHYL ALCOHOPMETHYL ETHYL KETONESYSTEM AT 760 MM. OF MERCURY
Temp.,
EtOH, Mole %
Ethyl Aloohol
EtOH,
THE
Methyl Ethyl Ketone YLOW 7Hi.h
93.2 95.6
Wt. % 89.8 93.3
1.02
1.01
1.03
1.64
1.46
1.84
Vapor Liquid
74.3 80.8
64.9 72.9
1.04
1.02
1.05
1.51
1.47
1.55
166.2
Vapor Liquid
64.2 69.5
53.4 59.3
1.08
1.07
1.10
1.37
1.34
1:39
4
165.3
Vapor Liquid
51.3 52.1
40.2 41.0
1.18
1.16
1.20
1.20
1.18
1.22
5
165.5
Vapor Liquid
49.2 48.5
88.3 37.6
1.21
1.19
1.23
1.15
1.16
1.18
6
167.0
Vapor Liquid
27.9 22.3
19.8 15.5
1.44
1.40
1.48
1.06
1.05
1.07
7
168.1
Vapor Liquid
20.8 15.1
14.4 10.2
1.55
1.49
1.60
1.05
1.02
1.06
172.0
Vapor Liquid
2.21
1.99
2.49
1.00
0.99
1.01
Run
F.
1
171.2
Phase Vapor Liquid
2
167.8
3
8
8.64 4.00
5.70 2.60
7
YLO W
YHi&
Y
behave almost ideally in liquid solution, while the n-butanol is seen to deviate somewhat. The greatest deviation from ideality is that shown by the ethyl alcohol-sec-butanol system. Here again the ethyl alcohol appears to behave ideally from approximately 40 mole % to 100 mole yoethyl alcohol in the liquid. The boundary curves for ethyl alcohol have been omitted from this figure because they fall very close to the experimental curve. At concentrations lower than 40 mole %, the ethyl alcohol shows a slight deviation. The sec-butanol curve exhibits a straight-line deviation which becomes significant even a t high concentrations of sec-butanol. In the ethyl alcohol-n-pentanol binary system, ethyl alcohol behaves ideally in the range from 10 mole % to 100 mole yoethyl alcohol in the liquid. At concentrations lower than 10 mole Yo, the ethyl alcohol deviates widely. However, the boundary curves show that the ethyl alcohol could still be considered ideal in this region of concentration. Here, again, the heavier component, n-pentanol, shows approximately a straight-line deviation. CONCLUSIONS
Figure 4.
Vapor-Liquid Equilibrium Diagram for Ethyl Alcohol-Ketone Systems at 760 Mm.
Vapor-liquid equilibria data were obtained for the ethyl alcohol binary system, ethyl alcohol+-butanol, ethyl alcohol-sec-butanol, ethyl alcohollt-pentanol a t 760 mm. of mercury. No azeotropes were found in these systems. Experimental activity coefficients were calculated with the use of vapor pressure data ob-
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
628
Vol. 45, No. 3
EQUILIBRIA AND h C T I V I T Y COEFFICIENT DATA FOR TABLE VII. EXPERIMENTAL VAPOR-LIQUID ETHYL !iLCOHOL-hfETHYL 12-PROPYL KETONESYSTEM A T 760 &IM. O F MERCURY Temp., F.
Phase
1
173.0
Vapor Liquid
93.0 91.9
2
173.3
Vapor Liquid
3 4
Run
EtOH, Mole 70
EtOH. Wt. 70
THE
Methyl n-Propyl Ketone
Ethyl Alcohol Y
YLOW
YHigh
Y
87.7 85 8
1.01
1.00
1.02
1.79
89.2 86.7
81.6 77.7
1.02
1.01
1.03
173.8
Vapor Liquid
83.8 (8.9
73.4 86.6
1.05
1.03
175.5
Vapor Liquid
73.8 62.4
6'0.1 47.0
1.12
5
177.5
Vapor Liquid
68.3 53.4
53.5 38.0
6
181.4
Tapor Liquid
51.8 30.1
7
186.8
Vapor Liquid
8
172.8
9
YLOW
'7High
1.65
I .94
1.67
1.59
1.77
1.06
1.57
1.52
1.62
1.10
1.13
1.38
1.35
1.41
1.16
1.14
1.17
1.30
1.28
1.32
36.5 18.7
1.43
1.40
1.46
1.23
1.21
1.25
34.7 13 . O
22.1 7.40
1.97
1.90
2.04
1.22
1.21
1.23
Vapor Liquid
95.3 95.9
91.5 90.8
1.01
1.oo
1.02
1.92
1.69
2.17
172.7
Vapor Liquid
96.7 96.9
94.0 94.3
1.00
1 .oo
1.01
2.22
1.83
2.69
10
173.0
Vapqr Liquid
98.3 98 5
96.8 97.2
1.00
1.00
1.01
2.35
2.21
3.46
11
175.6
Vapor Liquid
73.2 61.3
59.3 45.9
1.14
1.11
1.14
1.37
1.34
1.40
12
190.2
T'apor Liquid
23.7 6 52
14.3
2.49
2.34
2.65
1.25
1.24
1.27
14
173.2
Vapor Liquid
88.5 85.4
80.4 75.8
1.04
1.03
1,08
1.62
1.55
1.71
3.60
tained from the literature. Experimental errors are believed t o be 0.6% based upon the midpoint composition in mole fraction and 1 0 . l oF.
.
Limiting activity coefficient-composition curves P ere also determined and are included in the figure. Failure of the activity coefficient-composition curve for the methyl n-propyl ketone to originate a t y = 1.00 probablv may be attributed to slight imUI t: 6.
ETHYL ALCOHOL-KETONE SYSTEMS AT 750 MM. O F MERCURY
MATERIALS. The source and physical data of the purified compounds are shown in Table I. Absolute ethyl alcohol was used without purification. Commercial grades of acetone and methyl ethyl ketone were purified by fractionation in accordance with the procedure discussed before. The sample of methyl n-propyl ketone was obtained essentially pure. All materials were stored in contact n-ith a dry atmosphere throughout the investigation.
DATA. The vapor-liquid equilibria data presented were obtained at a constant total pressure of 760 mm. of mercury. Figures 4 and 5 and Tables V, VI, and VI1 include the experimental data. Analysis of the ethyl alcohol-acetone system was obtained by means of boiling point determinations. A Cottrell boiling point apparatus was used t o obtain the boiling points a t 760 mm. of mercury of t h e liquid and vapor samples which were withdrawn from t h e equilibrium still. Percentage compositions were read from a n experimental calibration curve of the boiling points of the mixtures versus mole per cent ethyl alcohol. A maximum error of less than 1.0% is believed present in the data for this system. The experimental results agree closely with those found in the literature ( 6 ) . The ethyl alcohol-methyl ethyl ketone and ethyl alcoholmethyl n-propyl ketone systems were analyzed b y means of refractive indices. The possible experimental errors in these data are estimated t o be of the same order of magnitude as those for the ethyl alcohol-alcohol binary systems. The x - y and temperature composition data for the ethyl alcohol-acetone system show no azeotrope. The ethyl-methyl ethyl ketone system exhibits an azeotrope a t 50.1 moleyo ethyl alcohol and 166.2" F. (74.0' C.). The ethyl alcohol-methyl n-propyl ketone binary shows an azeotrope a t 96.2 mole% ethyl alcohol and 172.5" F. (78.0" C.). Activity coefficients were caIculated from the experimental data and plotted versus mole per cent ethyl alcohol in the liquid, Fig-
OL- METHYL €THY
0
MOLE PERCENT ETHANOL
Figure 5.
Equilibrium Boiling Point Diagram for Ethyl Alcohol-Ketone Systems
INDUSTRIAL A N D ENGINEERING CHEMISTRY
March 1953
629
CONCLUSIONS. Vapor-liquid equilibria data were obtained for the ethyl alcohol binary systems, ethyl alcohol-acetone, ethyl alcohol-methyl ethyl ketone, and ethyl alcohol-methyl n-propyl ketone at 760 mm. of mercury. The ethyl alcohol-acetone system evidenced no azeotrope. I n the ethyl alcohol-methyl ethyl ketone system a n azeotrope was found a t 50.1 mole% ethyl alcohol a t 165.2’ F. (74.0’ C.). The ethyl alcohol-methyl n-propyl ketone system showed a n azeotrope at 96.2 mole% ethyl alcohol at 172.5’ F. (78.0’ C.). Activity coefficients were calculated with the use of vapor pressure data obtained from the literature and were found t o indicate only slight deviations from ideal solutions. However, the data were found to be inconsistent with Gibbs-Duhem criteria for thermodynamic consistency. NOMENCLATURE
P
= vapor pressure of pure substance
yf y-
+= y + 0.003 = 21 - 0.003
at experimentally determined temperature, millimeters of mercury P f = vapor pressure of pure s u b s t a y e a t experimentally determined temperature +0.1 F., millimeters of mercury P - = vapor pressure of pure substance at experimentally determined temperature -0.1’ F., millimeters of mercury T = total pressure, millimeters of mercury 5 = experimentally determined mole fraction in liquid phaee x* = x 0.003 2- = x 0.003 y = experimentally determined mole fraction in vapor phase
R
ETHANOL * ACETONE
0
EXPERIMENTAL ACTIVITYCOEFFICIENT -,’High
y~~~
= activity coefficient at upper limit of experimental error
as defined in Equation 2 = activity coefficient a t lower limit of experimental error as defined in Equation 3 LITERATURE CITED
0
(1) Brunjea, A. S., and Bogart, hl. J. P., IND.ENG.CHEM.,35, 255
F i g u r e 6. Activity Coefficient us. C o n c e n t r a t i o n of E t h y l Alcohol purities in the methyl n-propyi ketone sample and t o the use of vapor pressure data obtained from the literature. Activity coefficient curves for the ethyl alcohol-acetone and ethyl alcoholmethyl ethyl ketone systems are normal and consistent. However, none of the three systems follows the Gibbs-Duhem criteria and, consequently, no attempt was made to smooth these data by the conventional methods.
(1943).
(2) Hill, ITT. D., Ibid., 44, 205-10 (1952). (3) Hodgman, C. D., “Handbook of Chemistry and Physics,”
Cleveland, Chemical Rubber Publishing Co., 1944. (4) Horsley, L. H., IND. ENG.CHEM.,ANAL.ED.,19, 508-85 (1947). (5) Li, Y. M., and Coull, J., J. Inst. Petroleum, 34, 692-704 (1948). (6) Perry, J. H., “Chemical Engineers’ Handbook,” p. 573, New York, MoGraw-Hill Book Co., Ino., 1950. (7) Steinhauser, H. H., and White, R. R., IND.ENG.CHEM.,41, 2912-20 (1949). (8) Stull, D. R., Ibid., 39, 617-40 (1947). RECEIVED for review October 2, 1961.
ACCEPTED November 13, 1952.
Nickel-Catalyzed Hydrogenation of Commercial Aldol C. KINNEY HANCOCKl AND DOUGLAS D. HENSON2 Bureau of Industrial Chemistry, The University of,Texas, Austin, Tex.
P
RODUCTION of 1,3-butanediol by the catalytic hydrogenation of aldol was studied in 1942 as a fourth step in the olassicfive-step synthesis of 1,3-butadienefromnatural gasthrough acetylene. This paper is concerned with the results of studies with flake Rufert nickel catalyst (1, 4 ) and with a special dry(9). reduced nickel catalyst sumended in triethvlene -glycol ” Present address, Department of Chemistry, The Agricultural and Mechanical College of Texas. College Station, Tex. Present address, Eastern States Petroleum Co., Houston, Tex. 1
In a recent article by Hancock ( 6 ) concerning the results of
a similar study with Raney nickel catalyst, the following features were discussed: properties of the commercial aldol, hydrogenation autoclaves, and procedure for analysis of hydrogenated mixtures. The commercial aldol used in these two studies ameared t o b k largely 4hydroxy-2,6-dimethyl-l,3-dioxane,the condensation product between aldol and acetaldehyde, a8 discussed by Goldstein (6). All percentage yields of 1,3-butanediol given in the
-.
