VOL. 29, NO. 7
INDUSTRIAL AND ENGINEERING CHEMISTRY
846
in fair agreement with the results of the equilibrium studies. This fact is also brought out by the comparison of the equilibrium constants derived from the three free-energy equations, as tabulated in columns 5, 6, and 7. -4direct comparison of the results obtainable from Equations 3 and 5 with the experimentally measured equilibrium constants of Bliss and Dodge provides still another way of testing the exactness of the thermodynamic calculation of the equilibrium. At 623 O K. Bliss and Dodge found K p = 0.00149, whereas Equation 3 gives 0,00190 and Equation 5, 0.00167. At 651 K. their experimental value was 0.00126 as against 0.00129 by Equation 3 and 0.00114 by Equation 5 . I n the writer’s opinion this agreement, especially with Equation 5, is very satisfactory and serves to indicate quite fairly the potentialities of the thermodynamic calculation of equilibria. Literature Cited Beeck, O., S. Chem. Phys., 4, 680 (1936). Bliss, R. H., and Dodge, B. F., IND.ENQ.CHEM.,29, 19 (1937). Eucken, A., and Partz, A., 2. physik. Chem., ZOB, 184 (1933). Fiock, E. F., Ginnings, D. C., and Holton, W. B., Bur. Standards S.Research, 6, 881 (1931). Giauque, W. F., and Stout, J. W., S.Am. Chem. SOC., 58, 1144 (1936).
Gordon, A. R., J . Chem. Phys., 2, 65 (1934). Kelley, K. K., S.Am. Chem. Soc., 51, 779 (1929). Parks, G. S., Chem. Rev., 18, 325 (1936). Parks, G. S., and Huffman, H. M. “Free Energies of Some Organic Compounds,” A. C. S. Monograph No. 60, pp. 23-5, New York, Chemical Catalog Go., 1932. lbid., p. 125. Rossini, F. D., Bur. Standards J. Research, 13, 189 (1934). Tbid., 17, 636 (1936). Sanders, F. J., and Dodge, B. F., IND.ENG.CHEM.,26, 208 (19x41. \ - - - - I .
(14) Stanley, H. M., Youell, J. E., and Dymock, J. B., J . SOC.Chem. Znd., 53, 205T (1934).
I would be one of the last t o belittle the value of thermodynamics in the prediction of chemical equilibrium, and I am glad t o have Parks correct any false impression that our paper may have created. We may conclude that the equilibrium conditions for the ethylene hydration reaction, a t least a t low pressures, are well established as a result both of the recent work on heats of reaction and absolute entropy and of the direct experimental measurements of the equilibrium constants. It would appear desirable t o continue the attack on organic reaction equilibria along both of these lines. I hope that Parks and his co-workers and Rossini, to mention only a few of the investigators on the thermodynamic side, will continue to supply accurate thermal data. I should also like to point out that, even if we can calculate from thermal data the equilibrium constant for a given reaction, i t will still be desirable to carry out direct experimental studies of the same reaction. This is not only because of the desirability of having check results by independent methods, but also because we learn a great deal about the reaction from the direct experimental study t h a t cannot be obtained from the thermodynamic calculations. For example, the important practical questions of pressure, temperature, and catalyst to realize a rate of reaction t h a t permits an approach t o equilibrium, and of the extent of side reactions, are answered only by a n experimental investigation. Literature Cited (1) Kistiakowsky, G. B., Romeyn, H., Jr., Ruhoff, J. R., Smith, H. A., and Vaughan, W. E., S.Am. Chem. SOC.,57, 65-75 (1936). (2) Parka, G. S., Chem. Rev., 18, 325-34 (1936). (3) Rossini, F. D., Bur. Standards S. Research, 13, 21 (1934).
BARNETTF. DODGE Y A L UNIVERBITY, ~ NEWHAVEN,CONN. May 17,1937
GEORGES. PARKS STANFORD UNIVERBITY, CALIF. -4pril 2, 1937
Correction
,....
Attention is called t o a n article by C. D. West [ J . Am. Chem. SIR: Parks points out that the free-energy change for the ethylSoc., 59,742 (1937) ] on “Optical Properties and Polymorphism of ene hydration reaction can now be calculated from thermal data Paraffins” where the author points out an error in the article by withconsiderable confidence. This has been made possible largely ENG.CHEM.,28, 856 (1936)j on “Commercial ParafPage [IND. by the new value of the heat of combustion of ethylene obtained fin Waxes.” by Rossini a t the Bureau of Standards. Our work on the experiI n Table I1 of Page’s article the ordinary and extraordinary mental study of this reaction was completed more than a year berefractive indices are reversed, so that the column no - ne is fore this value was published, and our paper was submitted nearly actually ns - no,and the mean refractive index n is incorrect. 6 months prior t o its publication. Consequently a t the time that Table VI is therefore incorrect also. Table VII, corrected, is our calculations for the paper were made, our pessimistic stategiven here. ment about the value of the thermodynamic calculation may have been justified. It is true, however, t h a t prior to the submission TABLE VII. TRUEDENSITYOF, AND VOLUME PERCENTAIR IN, SOLID COMMERCIAL PARAFFIN WAXES of our paper the data on heat of combusion of
ethane Obtained by Rossini )(’ and On heat Of hy41.P. of Wax drogenation of ethylene by KistiakoTTsky ( 1 ) perF. o e. mitted a more accurate calculation of the free 121 49.4 energy of hydration than had been possible a t the tirne our work was started. These data were used 126 52.2 by Parks (2) to compute a value for the free energy of ethylene formation-namely, 15,820 in the usual 131 55.0 units’ If this is combined with the best 136 57.8 data on ethyl alcohol and water (the same data used by Parks in his communication), the value of 141 60.6 -1790 is obtained for the standard free energy of the hydration reaction in the vapor phase. This is to be compared with Parks’s value of -2030 a s given above, The equilibrium constant, K p , a t 700” K., based on the former A F ” value, is 1.62 x 10-3 compared to 2.36 X 10-8 given by Parks. Considering the circumstances, this is good agreement and on this basis the statement of Bliss and myself was certainly too pessimistic.
r D at
200 C.
Temp. OF.
0
di
d2 (Calod.)
(Obsnvd.) (Obsvd.)
dz
- dl
Vol. % Air
O C .
0.3343
50 10.0 60 16.6
1.5153 1.5136
0.900 0.897
0.902 0.899
0.002 0.002
$0.23 $0.23
0.3342
50 10.0 60 15.6
1.5179 1.5155
0.911 0.909
0,906 0.902
-0.005 -0.007
-0.55 -0.77
0.3339
50 16.6 60 50 10.0 15.6 60 15.6 70 21.1
0.912 0.909 0,917 0.914 0.914 0.911
0.907 0.904 0:911 0.909 0.913 0.909
-0.006 -0.005
0.3341
1.5182 1.6160 1,5213 1.5197 1.6214 1.5192
-0,005 -0.001
-0.56 -0.65 -0.55 -0.41 -0.10
-0.002
-0.22
10 .o
60
0.3337
-0,006
The difference between the calculated and observed densities now becomes so small as to be attributable to experimental error.
oILcoMPANY ( I ~ CASPER,WYO.,May 13, 1937
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