364
INDUSTRIAL AND ENGINEERING CHEMISTRY
T,may be predicted if the molecular weight, M , normal boiling point, T., and the density at the normal boiling point, d,, are known. This method is applicable only to non-polar liquids. The procedure is as follows:
Vol. 23, No. 4
in conjunction with the curve of Figure 2, be used for the estimation of latent heats of vaporization of non-polar liquids a t any desired temperatures. Acknowledgment
(a) Estimate the critical temperature of the material by
means of Figure 1 and Equation 7 as previously described. (b) Calculate the reduced temperature, T,,corresponding to temperature T. (c) On Figure 3, by interpolation between the curves, locate the point corresponding to T , and the value of M/d, of the material. ( d ) Read the corresponding value of M L / T from the scale of ordinates.
It is desired to acknowledge the important encouragement and assistance given by the Universal Oil Products Company in the initiation and development of the work herein reported, and particularly the criticisms and helpful suggestions of W. F. Faragher, 0. L. Kowalke, and W. H. McAdams.
I n Table I1 are values of ( M L / T ) ’ which were calculated on the basis of the Hildebrand principle from Figure 3. The agreement of these calculated values with those experimentally observed is in general good, but not so good as that of the values calculated from the equation of Kistyakowsky. The difference between the results of the two methods becomes marked when they are applied to liquids having high molal volumes such as the higher hydrocarbons. For these substances the latent heats predicted by this application of the Hildebrand principle are considerably higher than those based on the Kistyakowsky equation. Unfortunately reliable experimental data in this range are not available. However, in view of its scientific basis and the fact that it is in good agreement with practically all observed data, it would seem that the equation of Kistyakowsky should offer the more reliable basis. It is therefore suggested that this equation,
(1) Audubert, Rene, Compt. rend., 172, 375-8 (1921). (2) Friend, N.,Chem. News, 123, 219-20 (1921). (3) Herz, W., Z.anorg. allgem. Chem., 116, 202-6 (1922). (4) Hildebrand, J . A m . Chem. Soc., 37, 970-8 (1915). (5) Kistyakowsky, Z . physik. Chem., 107, 65-73 (1923). (6) Laar, van, J. J., “Zustandsgleichung,” 1924. (7) Laar, van, J. J . , Z . anorg. allgem. Chem., 146, 263-80 (1925) (8) Lewis and Weber, J. IND. ENG. CHEM.,14, 485 (1922). (9) Lewis and Weber, I b i d . , 14, 486 (1922). (10) Longinescu, I. N . , J . chim. phys., 26, 312-16 (1929). (11) McAdams and Morrell, IND. END. CHEM.,16, 375 (1924). (12) Moles, E . , J. chim. phys., 17, 415-24 (1919). (13) Mortirner, J . A m . Chem. Soc., 44, 1429-35 (1922). (14) Nutting, IND.ENG. CHEM.,22, 771 (1930). (15) Prud’bornme, M.,J . chim. phys., 21, 243-6 (1924). (16) Prud’hornme, I b i d . , 21, 461-65 (1924). (17) Schultz, IND. END. CHEM.,21, 557-9 (1929). (18) Schultz, I b i d . , 22, 785-8 (1930). (19) White, I b i d . , 22, 230-2 (1930). (20) Winter, J . Phys. Chem., 32, 576-82 (1928).
Literature Cited
Thermodynamic Properties of Dichlorodifluoromethane, a New Refrigerant 11-Vapor Pressure’ W. K. Gilkey, Frank W. Gerard, and Milo E. Bixler FRIGIDAIRB CORPORATION, DAYTON, OHIO
The vapor pressures of dichlorodifluoromethane apparatus by distillation inHE vapor pressure of a sured the absence of oil, and from -70” C. to the critical temperature (111.5’ C.) refrigerant is of more have been measured by a static method. An equation, the removal of successive obvious practical value small fractions by evaporaeuitable for thermodynamic calculations, is given: than any of the other thermotion a t c o n s t a n t pressure, dynamic properties, and its Iog,,p = 31.6315 - T - 10.859 log,, T 0.007175 T both before and d u r i n g a determination has repeatedly series of measurements, made where p is the pressure in absolute atmospheres and occupied the attention of this certain the a b s e n c e of air. Tis the temperature in degrees Kelvin. The calculated l a b o r a t o r y since an early normal boiling point is -29.8” C.; the Trouton and The remote possibility of an stage in the d e v e l o p m e n t u n k n o w n impurity which Hildebrand constants are, respectively, 20.0 and 27.1. of dichlorodifluoromethane. was not readily removable A table of calculated vapor pressures at 10’ C. intervals Three series of m e a s u r e by fractional distillation was menta, hereafter designated is given. checked in the case of 5 as A, B, and C, have been made by the static method, using six samples of the re- and 6 by fractional freezing. This was found to have no observable effect on the vapor pressure. As a final test for frigerant. purity, samples 5 and 6 were slowly evaporated almost to Purity of Samples dryness a t constant temperature without perceptible change The crude material as manufactured at the semi-works in vapor pressure. plant was known to contain only impurities removable by Measurement of Temperatures and Pressures fractional distillation and readily detectable if present. All laboratory samples were obtained from the middle cut of the Temperatures were measured with copper-constantan final fractional distillation as practiced in the regular plant thermocouples using an ice junction. The wire from which operation. Dryness was established in samples 1, 2, 3, and these couples were made was checked for uniformity and 4 by bubbling the vapor through sulfuric acid. Samples was calibrated a t the melting point of mercury, the transition 5 and 6 were dried over phosphorus pentoxide. Filling the point of sodium sulfate, and the temperatures of condensing steam and condensing naphthalene. 1 Received January 26, 1931.
T
x5
+
I S D USTRIAL A X D ENGINEERING C H E X I S T R Y
April, 1931
The obsen7ed microvolts a t these fixed points determined an almost linear deviation plot from the microvolts calculated from the copper-constantan tables in the International Critical Tables. These tables and the deviation plot, were used. to convert potentiometer readings into temperature. For series A a Leeds and Northrup type 8657C potentiometer with self-contained galvanometer was used; for series B the same potentiometer was used with a Leeds and Northrup type 2420b external galvanometer; a Leeds and Northrup
0
e o
-----
'
-8
I
I
I
I
I
28
30
32
34
36
I
Derivation of Equation
equation l o g 9 = 4.321
-
1047/T
as quoted by lllidgley and Henne (3) was derived from the data of series A only. Although this equation satisfactorily represents the data from which it was derived, it does not extrapolate correctly and is obviously inadequate for thermodynamic calculations involving derivatives. Soon after it was published the data of series B were obtained and the combined data represented by an equation in which the logarithm of the pressure was expressed as a function of a series in 1/T. The Jackson Laboratory of E. I. du Pont de Nemours and Company pointed out that this equation gave too much curvature in regions above and below those covered by the data from which it was derived, so a third series, C, covering the range from - 70' C. to the critical temperature, was made. For final representation of the combined data, anfequation of the form
-----.-
.
-T
365
1
I
I
I
I
I
40
42
44
46
48
'
Figure 1-Deviations of Observed P o i n t s from T h o s e C a l c u l a t e d b y E q u a t i o n 3 AA--Deviations corresponding to 1.0 per cent error in pressure BB-Deviations corresponding to 0.1 C. error in temperature
type K potentiometer and the type 242013 galvanometer were used in series C. The instrumental error in temperature measurement is estimated to be of the order OF -0.3" C. for series A, -0.2' C. for series B, and 1 0 . 1 " C. for series C. Nearly all pressures below 5 absolute atmospheres were measured on a mercury manometer, the usual barometric and temperature corrections being observed. Measurements of higher pressures were made with Bourdon tube gages calibrated before and after use against a dead-weight gage, whose constant was determined by direct comparison with the mercury manometer. The average error in the pressure measurements is estimated to be *0.5 per cent, although in a few instances it may he as great as 1.5 per cent. Apparatus The apparatus was essentially the same for all measurements, though certain changes were made to suit the varying requirements imposed by the wide range of pressures. A container, partially filled with the liquid, was connected to the pressure gage by small-bore tubing. For most of the lowtemperature measurements the container was a Pyrex bulb connected to the manometer by a short section of rubber tubing; a Dewar flask containing gasoline cooled by solid carbon dioxide furnished a satisfactory bath. At the higher pressures small welded steel cylinders, fitted with valves, held the liquid, and connection to the pressure gage was made with '/sinch copper tubing; an oil bath of large heat capacity with a manually controlled electric heater was used. The baths were thoroughly stirred and thermal equilibrium was established by violently shaking the containers. Temperatures and pressures could be held constant within the sensitivity of the instruments for 3 or 4 minutes. Data Space does not permit reproduction here of the ninety-eight individual measurements made. The essential information is contained in Table I, in connection with the final equation 3 and deviation plot (Figure 1).
log 9 = A
+ TB + Clog T + D T
(1:)
was chosen because it was known to represent other vapor pressure data satisfactorily and also because its derivative
is a parabolic function of the temperature, a matter of considerable practical convenience. OBSERVATION
T a b l e I-Experimental Data SAMPLE OBSERVER
TEMP.RANGE
c.
SERIES I
F. W. G. F. W. G. F. W. G. W.K. G.
1
1-19 20 21-24 25-30
3 4
31-49 50-63
4 5
64-98
6
7
SERIES B
W. K. G. W. K . 0.
0 t o 70 0
0 to 36 0 to 49
-
-35 to s -57 to -29
SERIES C
M. E . B.
-71 to 100
T a b l e 11-Vapor Pressures C a l c u l a t e d from E q u a t i o n 3 TEMP. PRESSUREf ABSOLUTE) TEMP. PRESSURE (ABSOLUTE) c. Alm. Lbs./sa. in. .4fm. Lbs./sa. i n. c. 30 7.337 107.9 0 1217 1.79 - 70 9.456 139.0 40 3.29 - 60 0 2240 176.3 11.99 5 69 - 50 50 0.3870 220.2 14.98 60 9.32 - 40 0 6340 271.7 18.48 70 14.6 0,9915 30 2 2 . 5 4 3 31.3 2 1 . 9 80 - 20 1 490 27.20 90 399,s 31.8 2.162 10 4 7 8.9 3 2 . 5 8 4 4 . 8 100 3.045 0 38.58 567.1 61.4 110 4.174 10 5 8 1.5 3 9 , 5 6 82 2 1 1 1 . 5 20 5.592
-
The data were plotted as loglo p against 1/T on a large scale and a smooth curve was drawn through the points. of this graph were plotted against The slopes, d log p / d absolute temperatures and approximately fitted by an equation having the form of (2). This equation was then integrated, the boiling point supplying the constant of integration. As was to be expected, this first approximation to the integral did not fit the vapor-pressure data accurately. The final equation was obtained by successive approximations using the Trouton and Hildebrand (2) relations to indicate the general trend of the slope curve in the regions to which they apply. The final equation is
(k)
loglo p = 31.6315 -
1816.5 7 - 10.859 loglo T + 0.007175T
(3)
1-VDUSTRIAL AND ENGINEERING CHEMISTRY
366
where p is the vapor pressure in absolute atmospheres; T is the temperature in degrees Kelvin (" C. 273.1).
+
Vol. 23, No. 4
and the equation of state (1) is 20.0. The Hildebrand constant is 27.1.
Discussion
Acknowledgment
Figure 1 shows the deviations of the observed data from the equation. The data of each observer are well represented by the equation. These deviations are random with respect to observer, sample, and order of observations on each sample. The dotted lines A A show the deviations from the equation assuming a 1 per cent error in pressure measurement. Lines BB show those due to an error in 0.1' C. in temDerature. The boiling point as calculated from the equation is f29.8" C. The 'Onstant as from the Of log p ,/ d a t the boiling point, the density of the liquid,
(+)
The writers gratefully acknowledge the assistance and under whose direction this suggestions of R. work was done, and of F. Bichowsky. They also wish to thank F. B. D o w i n g and A. F, Benning, of the Laboratory of E. 1,du pant de Nemours and Company, for criticisms offered during the course of this investigation. Literature Cited (1) Buffington and Gilkey, IND.ENO. CHEM.,23, 254 (1931) (2) Hildebrand, J . Am. Chem. Sot., 37, 970 (1915). (3) hlidgley and Henne, IND.END. CHEM.,22, 542 (19501
I I I-Cri tical Constants and Orthobaric Densities' F. R. Bichowsky and W. K. Gilkey FRIGIDAIRE CORPORATIOX,DAYTON, OHIU
I
N ORDER to calculate the heats of v a p o r i z a tion and other thermodynamic properties of dichlorodifluoromethane from the vapor p r e s s u r e s , the o r t h o b a r i c d e n s i t i e s are needed. Density of Saturated Vapor
The orthobaric densities of dichlorodifluoromethane have been measured up t o the critical temperature, and the critical constants determined. The saturated vapor densities below 50" C. were calculated from the vapor-pressure equation and the equation of state. Above 50" C. they were obtained by determination of dew points. Liquid densities from -40' to 50' C. were determined by a dilatometric method, and above 50" C. by Faraday's method, using glass floats. The critical temperature, pressure, and density are, respectively, 111.5' C., 39.56 atmospheres, 0.555 gram per cubic centimeter.
The density of the saturated vapor may be calculated from the equation of state ( I ) ,
p where A = 23.7 (1
=
RT
(V+ B)
A - v*
-
Table 11-Densities
p
= pressure in atmospheres V = volume in liters per gram mol T = absolute Centigrade temperature
by introducing the corresponding pressures and temperatures from the vapor-pressure equation (2): 1816.5
-T - 10.859 loglo T + 0.007175 T
and solving for the volume. The computed values are given in Table I. Table I-Volume and Density of Saturated Vapor, as Calculated from Equation of State and Vapor-Pressure Equation TEMP. DENSITY VOLUME TEMP. DENSITY VOLUME Cc./gram Cc./gram C. Grams/cc. Grams/cc. C. 10 0.02379 42.04 244.1 0.004097 -40 20 0,03149 31.75 161.3 0,006199 -30 30 0.04111 24.33 110.7 0.009037 -20 40 0.05313 18 82 78.13 0.01280 -10 0 0.01765 56.67 50 0.06856 14 59
These values have a probable error of k l . 0 per cent. At temperatures above 50' C. this equation of state fails and we depend on two determinations of the dew point. The values are d = 0.202 a t 94.4" C. and 0.168 a t 86.7" C. These dew-point determinations were made in small glass tubes and are probably accurate to * l per cent. Direct observations of the critical point showed that the IReceived
January 31, 1931.
Density of Liquid
Preliminary determination of the density of the liquid gave the density a t 0" C. as 1.40 and the coefficient of expansion, 0.00233. These values have been superseded by very careful determinations, by a dilatometric method, of the density by F. B. Downing and A. F. Benning, of the Jackson Laboratory of the E. I. Du Pont de Nemours and Company. Their values are given in Table 11. TEMP. O c .
loglo p = 31.6315
critical phenomena were unusually sharp and that the critical t e m p e r a t u r e mas 111.5 * 0.5' C.
-37.8 -28.4 -23.9 -17.8 - 9.60
of Liquid
DENSITY Grams/cc.
TEMP. c.
1.5095 1.4822 1.4689 1.4511 1.4254 1.3946
13.8 25.6 35.1 35.2 46.9 56.5
DENSITY Grams/cc.
1.3521 1,3081 1.2725 1.2722 1.2260 1.1834
The densities a t higher temperatures were determined by Faraday's method-that is, by the use of small floats in a sealed tube. The values obtained were 0.9785 at 91.1" C., 0.910 a t 98.9" C., 0.814 at 106.7" C. The estimated error was +0.5 per cent. Computed Densities-Rectilinear Constants
Diameter-Critical
These various determinations, together with the points determining the rectilinear diameter, are shown graphically in Figure 1. Table I11 gives interpolated densities at even temperature intervals. Table 111-Density of Liquid DENSITY Liquid Vapor Gram/cc. Grams/cr. 0.004097 1.517 0.006199 1.486 0.009037 1.456 0.01280 1.425 0.01765 1.393 0.02379 1.362 0.03149 1.329 0.04111 1.293 0.05313 1.155
and Saturated Vapor
TEMP.
TEMP.
* c.
* c.
- 40 - 30 -20 - 10
0 10 20 30 40
50 60 70 80 90 100 110 111.5
DENSITY Liquid Vapor Grams/cc. Gram/cc.
1.213 1.165 1.115 1.056 0.988 0.898 0.725 0.555
0.06856 0.0875 0.111 0.142 0.182 0.242 0.380 0.555
