VAPOR-LIQUID E
RIA
Ethylene Oxide-Acetaldehyde and Ethylene Qxide-rater Systems IC. F. COLES AND FELIX POPPERL Petrocarbon Limited, Manchester 17, England T h e vapor-liquid equilibria for the systems ethylene oxide-acetaldehyde and ethylene oxide water were determined at atmospheric pressure and also at higher pressures for some compositions. The vapor pressures of ethylene oxide and acetaldehyde were determined between 0" and 35" C. The system ethylene oxide-water is discussed in terms of its activity coefficients.
I"
T H E manufacture of ethylene oxide from ethylene via ethylene chlorohydrin, it is necessary to recover the product from an aqueous solution and to separate it also from a small amount of acetaldehyde which is produced a t the same time. This separation is normally achieved by first distilling overhead a product containing about 70Yc water, the remainder being ethylene oxide and acetaldehyde. This material is then fed to a series of fractionation columns, either operating under pressure, so that water condensers can be used, or operated under atmospheric pressure with refrigeration. In making the calculations for such columns it is necewiiy to know vaporliquid equilibria for the components of the system in certain concentration ranges. Since the ethylene oxide and acetaldehyde are far more easily separated by distillation from water than from one another the information required for practical purposes is that for the three binary systems ethylene oxide-water, acetaldehyde-witer, and ethylene oxide--acetaldehyde. From the general considerations advanced by Hildebrand ( 6 ) it is to be expected that the ethylene oxide-acetaldehyde system ill not deviate widely from the ideal, whereas the ethylene oxide-water system would deviate strongly. Published data are available (16) for the system acetaldehyde-water and show clearly the deviation from ideal behavior. As there vas no reported information on the system ethylene oxide-acetaldehyde, the vapor-liquid equilibria for this system was determined. Some information (7) is available on the ethylene oxide-water system, but this is not of sufficient precision in the range, where it is chiefly required-that is, for low concentration of a ater. In the work mentioned, the authors designate as 1Wyo ethylene oxide the vapor composition in equilibrium with liquid mixtures containing more than 20.27, ethylene oxide by weight. This indicates an insufficiently precise method of analysis Vapor-liquid equilibria for acetaldehyde and water are available (16) for liquid compositions up to 38yo by weight of acetaldehyde. These data are given in Table I with the activity coefficients calculated so that they can be compared with those for the system ethylene oxide-water
Acetaldehyde was fractionated on the same column in order to separate it from polymers and acetic acid. The distillate was either used immediately after fractionation or was stored a t -10" C. in a closed bottle. Vapor pressure determinations indicated the high purity of the samples. VAPOR P R E S S U R E DETERMILTATIONS
If the vapor-liquid equilibria are to be expressed in terms of activity coefficients, the vapor pressures of ethylene oxide and acetaldehyde are required over the temperature range of the equilibrium experiments. The vapor pressure of ethylene oxide has been determined by Maass and Boomer (11) from -57" t o f12.8' C. and by Moor et al. (IS)from -5' to +40° C. The following equations can be fitted to these two sets of data. 3laass and Boomer (11): log p = Moor et az. (IS):log p =
- 1356 + 7.653 -
-1410 + 7.839
where p is the pressure in millimeters of mercury and T is the absolute temperature, Since these equations differ appreciably, experimental determinations were made with an Loteniscope and the resuks are shown in Table XI. The range over which vapor pressures were determined was from 0" to 32' C. The data can be representcd by the equation
+ 7.659
log p =
This is in close agreement with the equation fitted to the data
of Maass and Boomer (11) and as the temperature range in the present experiment differs from that of Maass, it appears probable that an equation of this type gives a good approximation over the whole range. Since the completion of the experimental part reported here,
TABLEI. VAPOR-LIQUIDEQLTLIBRIA FOR ACETALDCHYDBWATERAT 760 MM. [Data from Pascal et al. (fs)] Mole % Acetaldehyde Liquid Vapor 0.5 25 1 50 4 75 10 89 20 93
B.P., 0
c.
93.5 82 63 43 33
PURIPICATION O F COMPONENTS
Activity Coefficients CHsCHO Ha 0 5.4 0.95 7.3 1 .oo 4.6 1.15 4.0 1.43 3.0 1.75
OF ETHYLEYE OXIDE TABLE 11. VAPORPRESSURE
The ethylene oxide used contained small amounts of acetaldehyde and water. It was purified by mixing with 5% triethanolamine and fractionating on a column, 2.5 cm. in diameter and 60 cm. high, packed with glass Fenske helices. The column was operated a t a 10 t o 1 reflux. The first 5% v a s rejected and a residue corresponding to 20% of the charge was left in the reboiler. The remaining distillate vias used for the experimental work.
(-273.2'
Temp., C. 0.3 0.4 1.1 1.2 10.5 10.7 11.1 12.55 15.7
1 Present address, Coal Tar Research Association, Oxford Road, Gomersal,Leeds, England.
1434
Mm. H g 506 508 522 526 700 766 775 823 928
C.
0' K.) Temp., O C. 16.05 16.6 16.75 21.6 24.9 29.0 30.35 31.8
Mm. Hg 939 958 964 1156 1294 1518 1574 1654
INDUSTRIAL AND ENGINEERING CHEMISTRY
July 1950
Giauque and Gordon ( 4 ) have also reported data of ethylene oxide vapor pressures. Over a range from -50' to +12' C. the equation log p =
- (2045.70/T) - 0.021507 T
+
2.3328 X 10-6 T a = 14.31363
is in excellent agreement with their experimental results. Using these equations to find the boiling point at atmospheric pressure (760 mm.) the following result: Moor et al. ( I S ) : 11.2' C. Maass and Boomer (11): 10.7 * 0.1' C. Giauque et al. ( 4 ) : 10.50 * 0.05' C. Present work: 10.4 * 0.2" C. The boiling point of ethylene oxide has also been reported (19)
as 10.7"C. by Timmermans and Hennaut-Roland, but no further vapor pressure data are recorded in that paper. The freezing point of the ethylene oxide used for these experiments was -113 * 0.5' C. This may be compared with melting pointa in the literature of -111.4' C. (II), -111.7' C. (19), and -112.51 * 0.05' C. (a).
1435
The Othmer still is normally operated with liquids boiling above the ambient temperature. This results in a slight reflux taking place initially in the outer jacketing space. Under these conditions, and with liquids whose boiling points are not far apart, the temperature recorded in the vapor stream will approach closely that recorded in the boiling liquid. With the liquids under consideration here the two observed temperatures, in the vapor and liquid, respectively, differed appreciably. When ethylene oxide-acetaldehyde mixtures were being distilled at atmospheric pressure the vapor stream was always slightly superheated unless the top of the still was cooled below the dew point of the vapor. When this was done the same temperature was observed in the liquid and in the vapor. During the experiments with ethylene oxide-acetaldehyde mixtures some observations were made with and without cooling the top of the still. No significant difference in the vapor and liquid compositions was found for this change in conditions. With ethylene oxide-water mixtures of high water content in the glass equilibrium still the observed temperature in the vapor was found t o be lower than that in the boiling liquid, as a result of the very large difference between the boiling point of the liquid and the dew point of the vapor.
TABLE 111. VAPORPRESSURE OF ACETALDEHYDE (-273.2''
C.
Temp.,
-0.2 2.7 6.7 9.3 11.6
C.
Mm. Hg
0" K.)
Temp.,
332 375 443 494 531
.a
C.
13.3 17.6 20.7 30.8 34.4
Mm. H g
577 682 766 1120 1259
The vapor pressure of acetaldehyde has been determined by Gilmour (6) from its atmospheric boiling point to 28' C. and by Emeleus (3)from -97' t o 0" C. The determinations by Emeleus are in line with those of Gilmour over the temperature interval of 0" to -23' C., but below this temperature the points deviate so widely from the straight line, when plotted as log p against 1/T, that they are assumed in error. Experimental data by the present authors are recorded in Table 111. The following equations are used to express the results:
- 1412 + 7.694 T 1413 = -+ 7.694
Gilmour (6): log p = Present work: log p
From these equations the following boiling points a t 760 mm. can be calculated: Gilmour (6): 20.2' C. Present work: 20.4 * 0.2"C. The agreement in the case of acetaldehyde is excellent. Agreement on ethylene oxide is not so general. It should, however, be noted that the present results are in good agreement with the results of Giauque and Gordon ( 4 ) of recent date. In any case, it is considered that errors which might arise from the presence of impurities of the order indicated by the boiling point differences are small in relation t o those arising from the inaccuracy of the analytical methods used in the vapor-liquid equilibrium determinations, VAPOR- LI QUID E QUI LIBRIA
The vapor-liquid equilibria were determined in stills of the Othmer type. T h e experiments a t atmospheric pressure were done in a glass still of conventional design ( l a ) , and the experiments above atmospheric pressure were done in a pressure still based on a design by Othmer (16). As ethylene oxide boils a t 10.4' C., it was necessary to circulate a cooling fluid in the condenser.
a
I
0
I
I
0
MCkEoh%HzfzO Figure 1.
I
80
I
IN LIO.
Ethylene Oxide-Water System
Solid curves on activity coefficient plot represent the van Laar equations; broken lines are drawn through experimental points. triangles represent data in Table I for Acetaldehyde-water
INDUSTRIAL AND ENGINEERING CHEMISTRY
1436
The pressure still ( 1 6 ) was designed to take a charge of 500 ml. of liquid. The pressure was measured vith a Bourdon gage accurate to * 1 pound per square inch. Samples were withdrawn undrr pressure With this metal still n o difference was observed between the temperatures in the liquid and in the vapor stream. The temperatures were above atmospheric and the larger size and metal construction favored heating of the thermometer in the vapor stream to the correct temperature.
TABLE
Is'.
VAPOR-LIQUID E Q C I L I B R I A FOR
b:THYLEKE
OXlDE-
ACETALDEHYDE T
~
C. 10.4 10.4 10.6 10.6 10.6 10.6 10.6 10.7 10.7
11.6 13.5 14.6 16.0 17.3
... ...
33.5 33.5
...
34.7
~
~Mole , ,% (CHzh0 in
Liquid 99.65 99.17 99.12 98.95 98.77 98,52 97.94 96.75 96.66 88.30 64.0 52.8 41 .0 29.5 19.4 17.2 98.68 98.66 97.50 81.4
Vapor
Activity Coefficients ~ _ _ _ _ ~ CIIICHO
(CHdzO
A t 760 Mm. Pressure 99.82 1.002 = 0.005 99.56 1.005 * 0,005 99.47 0.993 0.005 99.30 0.998 * 0.005 99.12 0.993 = 0.005 99.08 0.995 * 0 . 0 0 5 98.61 0.996 = 0.005 97.78 0.995 * 0.005 97.92 1.00 * 0.01 91.80 1 . 0 0 * 0.02 72.0 1.00 * 0.04 63.6 1.02 = 0.04 51.3 1.00 * 0.04 39.0 1.03 0.04 25.4 ... 23.0 ... f
A t 34 Lb./Bq. Inch Pressure 98.99 1 . 0 3 * 0.03 99.04 1 . 0 3 * 0.03 98.20 86.2 1.06 i O . 0 5
0.76 0.12 0.78 = 0.08 0.88 * 0 . 0 8 0.98 1 0 . 0 8 1.05 * 0.08 0.92 + 0 08 0.99 * 0.08 0.99 * 0 . 0 8 0.91 t 0 . 0 8 0.99 * 0 . 0 8 1.01 1 0 . 0 4 0.96 i- 0.04 0.97 = 0 . 0 4 0.99 * 0.04
where yi P p,
Vol. 42, No. 7
activity coefficient for component 1 total vapor pressure of system vapor pressure of pure component 1 a t the temperature of the system 21 = mole fraction of component' 1 in liquid y1 = mole fraction of component 1 in vapor = = =
This calculation assumes that the vapors approximate ideal gases. I n making this calculation it is clear that the effect of the analytical accuracy can become very large, for example, where 211 is small and is found by determination of component 2. Thc estimated accuracy of the activity coefficientshas been included, therefore, in Tables I V and 17.
f
, .
1.09 1.03
* 0.1 * 0.1
1 , 0 3 ' i0 . 1
When the ethylene oxide-wat,er system was studied, the rate of hydrolysis was known to be appreciable above 50" C., if more than an hour was allowed for equilibrium to be reached. Below this temperature the glycol formation is negligible provided t>hc pH of t,he mixture is maintained in the range of 6 to 8. A s a precaution, each day's work was started with freshly prepared liquid mixtures. ANALYSES
Three methods of analysis were used in this work. For the ethylene oxide-acetaldehyde system, amounts up to 10% acetaldehyde were determined by t'hr silver oxide met,hod ( 1 7 ) . For the remainder of the range, ethylene oxide was determined by a modified Lubatti method ( 1 0 ) in which the ethylene oxide reacts with hydrochloric acid to give ethylene chlorohydrin. This method is subject to certain errors which tend to give low results. The reasons for this have been discussed hy Lichtenstein and Twigg (9) and a more detailed study has been made in t,his laboratory and is being prepared for publication. In the ethylene oxide-water system, the mixtures cont,aining less than 5% water were analyzed by the Karl Fischer reagent. The accuracy of the analyses is estimated to be 170for ethylcne oxide, 2y0for acetaldehyde, and 2% for water, the percentage in each case being on the result. The significance of the accuracy of the analyses is discussed below. RESULTS
While the curves relating liquid to vapor composition and to boiling point are sufficient for distillation calculations, the expression of the results in terms of activity coefficients is much to be preferred, as it admits an immediate appreciation of the deviation from the behavior of an ideal mixture. It also enables some evaluation of the consistency of the data and, if necessary, extrapolation from a few experimental results over the a-hole range of compositions. The significance of the activity coefficient and the application of vapor-liquid equilibrium data is discussed by Carlson and Colburn ( 2 ) . The activity coefficients have been calculated from
F7.
T'.~POR-IJIQCID EQUILIBRI.4 ATER
_X t . 70HzO in Mole_% (CHd20
11.5 11.7 11.8 11.9 12.0 18.2 13.7 14.3 15.0 15.1 16.4 31 .O 31.5 37.6 50.0
FOR E T H Y L E S E OXIDE-
IT'
Liquid
Vapor
2.05 2.85 4.2 4.95 5.85 20.1 24 5 35.5 52
0.80 0.27 0.41 0.39 0.46 0.61 0.64 0.61 0.64 0.65 0.75
57
60 79.5 82.2 85.4 90.7
1.14
1,66
2.58 6.3
Liquid
Vapor
~ .4ctivity Coefficient ~ ~(CH2)rO H20
At 760 M m . Pi' e w m 95.1 99,27 1 01 =0.005 93.3 99.34 1.02 * 0 . 0 1 91 . o 99.00 1.04 * 0.02 99.05 89.0 1.06 =t 0 . 0 2 98.88 87.5 1.06 * 0.02 98.53 61.5 1.44 * 0 . 0 5 98.45 1 . 5 6 =k 0.05 56.0 98.53 1.96 * 0.05 43.2 98.45 3 . 0 1 * 0.05 27.4 98.41 23.2 3.55 * 0 . 0 5 98.16 21 . 0 3 . 7 2 * 0.05 96. 48 4.85 * 0.08 9.5 95.95 8.2 5.54 + 0.1 6.5 93,i 5.55 * 0.1 86.0 5.5fi * 0 . 1 4.0
11.1
7.3 8.1 6.3 6.5 2.6 2.3 1.62 1.27 1.22 1.27
*
0.0
*
0.6
*o 6
*0.5 -0.5
=0.2 *lLl
* 0.1 * 0.08 0.08 * 0.08 * 0.06 * 0.06 =t
0.88 0.97 1.05 ;t 0 06 1.20 * 0.06
Pres-
~sure, b . 1 ~T~~~~. ~ ~ Wi. , , ,c/o M 2 0 Inch
C.
Liquid Vapor Above 760
34 34 34 34 65
35.4 35,4
...
37.5 56.3
1.41 1.36 2.20 4.70 7.9
0.27
0.67
(CH2)zO
Pressure 96.56 99.34 9 6 . 6 8 99.27 9 4 . 8 5 98.90 8 9 . 3 98.36 8 2 . 7 97.58
0.97 0.97 0.96 0.99 1.13
11111.
0.30 0.45 0.99
1Iole % (CHd20 Liquid Vapor
Activity CoeFi, €120
8.4 9.3 9 3 5.6 a.7
Figure 1 shows the plot of activity coefficient on a logarithmic scale against' the liquid mole composition of each component. The activity coefficients for the ethylene oxide-acetaldehyde can be seen from Table IV to be, within the limits of experimental error, unity over the whole range except for a slight decrease in the region of below 1% acetaldehyde. These data, therefore, have not been represented graphically, and it is concluded that' tjhe system obeys Raoult's lam with close approximation. This is not surprising in view of the many similarit,ies to which reference has already been made by Maass and Boomer ( 1 1 ) . The system ethylene oxide-water is shown graphically iii Figure 1. The equilibrium curve is shown only for the rango 95 to 100 mole % ethylene oxide in the vapor phase in order to show this section clearly. The Gibbs-Duhem equation may be expressed as
Assuming the vapors to behave as ideal gases this may bc rcwritten for a binary mixture as
This equation has been integrated in several ways, notably by van Laar (8), Margules (IS),and Scatchard ( I t ? ) , and it is of in-
INDUSTRIAL AND ENGINEERING CHEMISTRY
July 1950
1437
The change of activity coefficient with temperature is described by
Figure 2. Change of Activity Coefficient with Temperatule for Ethylene OxideWater System Broken lines show the experimental results corrected for temperature variation; solid lines show the van Laar equations
terest to apply the derived equations to the present data. In integrating the Gibbs-Duhem equation it is necessary to assume constant temperature over the whole range of mixture compositions and for most systems of industrial interest this does not prevent the possibility of fitting the derived equations closely t o the experimental data, since the components normally boil within 20" C. or less of one another and the temperature range for constant pressure data is relatively small. With ethylene oxide and water this is not the case. The derived equations were first fitted t o the uncorrected experimental data. The van Laar and Margules equations give essentially the same curves for these data, whereas the Scatchard equation gives a curve that is obviously not appropriate. The van Laar equation, therefore, is used here in the form
111 y2
=
-
(1
z)2
where L1 is the heat absorbed when 1 mole of component 1 is added to a n infinite amount of component 2. Bichowski and Rossini (1) give LI for ethylene oxide and water as 1500 calories per mole a t 18" C. Using this value for L1 it is possible to make an approximate calculation of the effect on In y, for ethylene oxide a t the terminal value, for a temperature drop from 100 ' to 20" C. Taking the terminal value of In y as 1.92 from the experimentally derived curves, this would become 2.40 for an 80' C. temperature fall. Similar calculations can be made for the experimental values in the range where the temperature was above 20" C. I n Figure 2 the In y curves, corrected for temperature in this manner, are shown and the "corrected" curves and those for the corresponding van Laar equations may be compared. Although the curve representing experimental results for water now approximates very closely to the van Laar curve, the deviation of the ethylene oxide curve is about the same as that found for the curve for water in the direct representation of the experimental results in Figure 1. Therefore, it appears t h a t the wide temperature range for this system is not the cause of the failure of the van Laar equations t o fit the experimental data very closely. It is of interest t o compare the equilibrium data for acetaldehyde-water with those for ethylene oxide-water. The activity coefficients have been calculated for the former system (16) and are shown in Table I, and they are compared with the ethylene oxide-water equilibria over the corresponding compositions range in Figure 1. Agreement between the two sets of results is close. This appears to be consistent with results showing t h a t the ethylene oxide-acetaldehyde system is ideal, although present theoretical treatments do not appear to permit a rigid comparison t o be made. It is proposed to discuss this point further in a later publication. Carlson and Colburn ( 8 ) quote the acetaldehyde-water system as one which cannot be fitted closely with derived equations. VARIATION O F ACTIVITY COEFFICIENT WITH P R E S S U R E
The experimental determination of the vapor-liquid equilibria a t elevated pressures was done only for compositions of immediate practical interest. Those for ethylene oxide-acetaldehyde (Table I V ) confirm the ideality of the system. Those for ethylene oxide-water (Table V) show the expected decrease of activity coefficients from the values a t atmospheric pressure. These data are not sufficiently complete for further consideration.
R
+
where A and B are the terminal values of the In y versus liquid composition curves. From the experimental curves, taking these constants to be A = 1.92 and B = 2.68, the van Laar curves have been drawn in Figure 1. The agreement is rather good and supports the view ( a ) that these equations form a useful method of smoothing or extrapolating data of this type, even for such highly irregular liquids as those under discussion. It has already been noted that the derived equations from the Gibbs-Duhem relationship assume constant temperature and pressure over the whole range of compositions. For the system under discussion, the total variation of temperature is from 10" to 100" C., although from 10 to 100 mole yoethylene oxide in the liquid the variation in the boiling point is only 20" C. It is therefore possible that the terminal value of the ethylene oxide is lower than it would be for isothermal data, since increased temperature will cause the activity coefficients to approach unity for any system.
CONCLUSIONS
Ethylene oxide and acetaldehyde behave as an "ideal" liquid mixture in regard t o their vapor-liquid equilibria. Ethylene oxide and water mixtures show high positive activity coefficients for both components. ACKNOWLEDGMENT
The authors wish t o thank H. Steiner for his assistance and encouragement in preparing this paper and also R. K. Truelove for carrying out the analytical work. LITERATURE CITED
Bichowsky, F. R., and Rossini, F. D., "Thermochemistry of Chemical Substances," p. 46, New York, Reinhold Publishing Corp., 1936. Carlson, H. C., and Colburn, A. P., IND. ENG.CHEM.,34, 581-9 (1942).
Emeleus, H. J., J . Chem. SOC.,1929, 1733-9. Giauque, W. F., and Gordon, J., J . Am. Chem. SOC.,71, 217681 (1949).
INDUSTRIAL AND ENGINEERING CHEMISTRY
1438
Vol. 42, No. 7
(6) Gilmour, R., J . SOC.Chem. Ind., 41, 293-4T (1922). (13) Moor, V. G., K a n e p , E. K., a n d D o b k i n , I. E., Trans. Ezptl. (6) H i l d e b r a n d , J. H., “Solubility of Konelectrolytes,” 2nd cd., Research Lab. K h e m g a s , Materials o n Cracking and Chemical chap. 111, 1936. Treatment of Cracking Products U.S.S.R., 3, 320-8 (1937). (14) O t h m e r , D. F., J N D .EXG.CHEM.,35, 614-20 (1943). (7) Kireev, V. A., a n d P o p o v , A. A . , J . Applied Chem. (U.S.S.E.), 7, 489-94 (1934). (15) O t h m e r , D. F., a n d Morley, F. R., Ibid.,38,751-7 (1946). (16) Pascal, P., D u p u y , Ero, a n d Garnier, Bull. soc. chim. France, (8) Laar, J. J. v a n , 2 . phusik Chem., 72, 723-51 (1910); 83, 599608 11913). 29, 9-21 (1921). . , (9) Lichtenstein, H.J., a n d Twigg, G . H., Trans. F a ~ a d a ySoc., 44, (17) P o n n d o r f , W., Ber., 64, 1913-6 (1931). 905-9 (1948). (18) Scatchard, G., a n d H a m e r , W.J., J . Am. Chem. Soc., 57, 1805(10) L u b a t t i , J., J . SOC.Chem. Ind., 51, 361-7T (1932). 9 (1935). (11) Maass, 0.. a n d Boomer, E. H.. J . Am. Chem. Soc., 44, 1709-28 (19) T i m m e r m a n s , J., a n d Hennant,-Roland, -MMme., J . chim. phys., (1922). 34, 693-739 (1937). (12) Margules, M., Sitzber. A k u d . Wiss Wien, itlath.-nutunu. K ~ S S E . 11, 104, 1243-78 (1895). RECEIVED December 28, 1949.
Solubility Diagrams for Ternary and Quaternary Li Correction and Addendum Since the article on “Solubility Diagrams for Ternary and Quaternary Liquid Systems” appeared [Smith, J. C., ISD. ENQ. CHEM.,41, 2932 (1949)], several errors and omissions have been called to the author’s attention. On page 2932 in Table I
ADDITIONALSYSTEMS Components of Aqueous Ternary Systems Acetone Chloroform Glycerol n-Heptane n-Hexane Polyvinylpyrrolidone (Kollidon) Allyl alcohol Carbon tetrachloride Trichloroethylene Ethanol Benzene Ethyl acetate Ethyl ether Toluene Trichloroethylene Trichloroethylene m-Xylene ~\‘Ha)zSOr Etkyl ether Ethanol Vinyl pyrrolidone &PO4 Glycerol Acetone tert-Amyl alcohol Aniline Benzyl alcohol %-Butylalcohol Cyclohexanol Methyl ethyl ketone Methanol iMethylmethacrylate Methyl ethyl ketone Benzene Butyl Cellosolve Chlorobenzene Glycerol n-Heptane n-Hexane 1,1,2-Triohloroethane Trichloroethylene Phenol Benzine Isopropyl alcohol Diisopropyl ether n-Propyl alcohol n-Propyl acetate Pvridine Benzoic acid Toluene Diethylamine Ethanol Trichloroethylene Allyl alcohol Ethanol Ethanol Methyl ethyl ketone Nicotine Vinylpyrrolidone Ethyl acetate Ethyl ether Methylene chloride
Temp.,
C.
25,60
the citation number for the system water-acetaldehyde-vinyl acetate should be ( 112), not (102). On page 2933 the temperature listed for the system water-ethanol-p-xylene should be 15” C., not 0’ C.; the system water-ethanol-p-nitrotoluene should be water-ethanol-0-nitrotoluene, although incorrectly listed in several places in the original reference (21). The system benzenetoluene-n-heptane-methyl sulfate should be deleted from Table IV on page 2935. Additional ternary systems for which data are available are listed, with literature citations.
25 25
LITERATURE CITED
25
25
25 and b.p.
25 and b.p. 25 70
20 25 25 20,67 25 33 20 25 0
25 7.6, 25,48.6
25.75
25,75 25 60 25 25
25, 50 25 25 25
25
25 25
25 80,90
25 20,35 50
25 25 25 and b.p. 25 20,67 25 17
25 25 2 :.
system-
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