Thermodynamic Properties of Fluorochloromethanes and -Ethanes
A. F. BENNING AND R. C. MCHARNESS Kinetic Chemicals, Inc., Wilmington, Del.
considered to be of almost equal importance. A knowledge of their complete thermal and physical properties is expected to facilitate their use, particularly with regard to the selection of refrigerant and design of equipment. Thermodynamic tables for dichlorodifluoromethane have been published, and the availability of this information has been of definite assistance in developing the refrigeration art (2-6, 7 ) . The establishment of the thermodynamic properties of these four compounds is the purpose of this investigation.
A scheme of experimental and mathematical treatment for the development of the thermodynamic properties of fluids has been applied to four fluorochloro hydrocarbons-namely, CHClF2, CHCLF, CCI,F, and CClzF-CClFz. The properties selected for experimental measurement were vapor pressure, liquid and vapor densities, liquid and vapor heat capacities, and the ratio Cp/Cvfor the vapor. The accuracy and consistency of the experimental results have been evaluated by calculation of thermodynamic networks, each involving the experimental values for different sets of properties. The information developed provides a basis for the cal: culation of the thermodynamic properties of these compounds with a degree of accuracy sufficient for normal engineering requirements.
T H E thermodynamic properties of a refrigerant which are of primary interest to the design engineer are usually presented as a table showing the heat contents, entropies, specific volumes, and vapor pressures of the fluid a t even values of temperature. The preparation of such a table of properties requires a certain minimum of information which must be developed experimentally. A variety of independent properties may be selected which will define the thermodynamic
m H E use of the fluorochloromethanes and -ethanes as refrigerants proposed by Midgley and Henne (8) in 1930 was followed by the successful commercial production of dichlorodifluoromethane. The development and more widespread use of a variety of types of compressors have created a demand for additional refrigerants having the same desirable chemical and physiological properties as dichlorodifluoromethane but with different thermal properties, particularly boiling points. The class of compounds, fluorochloromethanes and -ethanes, contains a number of individuals which have the thermal, physiological, and chemical properties desired in a refrigerant. Four compounds whose boiling points range from -40” to +47” C. were selected; this group includes one or more refrigerants in each 40’ interval and is designed to cover the entire practical refrigeration range, both as regards operating temperature and type of compressor.
1
Compound CHClFt CClzFz CHCllF
Trade Name “Freon-22’’ “Freon-12” “Freon-21”
.:
P.,
C. -40.8 -29.8
8.9
Compound CClaF CC12F-CClFn
Trade Name “Freon-11’’ “Freon-113’’
-75
-50
0 50 100 150 250 TEMPERATURE-’c.
FIGURE 1. VAPORPRESSURE OF FLUOROCHLORO HYDROCARBONS
system, and if these are determined all others can be calculated. Experimentally measured properties in excess of the minimum are advisable as checks on those properties used to define the thermodynamic system. The properties selected for experimental measurement are vapor pressure, P-V-T relations of the superheated vapor, heat capacity of the vapor a t 1 atmosphere, density of the saturated liquid, ratio of the heat capacities of the vapor (Cp/C,,), heat capacity of the liquid, and critical constants. Determination of the physical properties of these four compounds required, first, materials of high purity. Since the compounds studied were available in large quantities, purification by repeated fractionation was adopted. In each case the purity of the product was checked by methods such as freezing point, limiting vapor density, and change in vapor
5. P., C. 23.7
47.6
Although the production of these compounds was the first step necessary for their adoption as refrigerants, the measurement of their thermal and physical properties was 912
JlJLY, 1939
INDUSTRIAL AND ENGINEERING CHEMISTRY
FIGURE 2. VAPORDENSITY APPAR-~TUS
pressure during distillation. In no case was material used if the tests indicated the presence of impurities. The vapor pressure of each of the compounds was measured by a static method over the pressure range of 0.1 atmosphere to the critical point (Figure 1). The temperatures and pressures were measured on calibrated instruments which were checked against primary standards both before and after each experiment. The measurements wcre restricted to a limited number of points, since it was considered more convenient and reliable to cheek a few points a number of times rather than determine a large number of points and correlate them by means of the equations developed t,o represent the vapor pressure. Equations were formulated to represent the experimental data. It was found possible t.o do this with the conventional five-term logarithmic equation for vapor pressure: B logp = A Clog T i- DT
913
atmospheres was measured in glass equipment. The densities so determined agreed with the equations developed from the other work with a deviation of less than 0.4 per cent except in one case. The densities of the saturated vapors calculated from these equations are compared in Figure 3, where the deviation of each of them from the ideal gas equation is plotted against the reduced pressure-temperature quotient, P./T, (6). The various compounds all fall close to the same line, the maximum variation being less than 1 per cent. The heat capacity of the vapors at one atmosphere was measured in a conventionaltype flow calorimeter. Measurements were made over the range 35" to 135" C . The instrument used was checked by measurements on ammonia and sulfur dioxide and found to give resnlta in good agreement with the accepted values. The characteristics of the calorimeter as regards effect of flow rates, shield temperatures, and heat input on the values obtained for the heat capacity were carefully determined. None of these seriouslv altered the results, so that the measurements are considered to be accurate to within 5 per cent. The change of heat capacity with temperature could be represented by straight lines within the accuracy of the experimental data, as Figure 4 shows. II30 m ""
/
'LFLUOROCHLORO METHANES
I
CH CC l 3II FZFzF
, ~
CHCIzF
b20 b20 la-FLUOROCHLORO ETHANE
RT __ PV
C C I p F-CClFp
I
+T +
The pressures calculated by means of these eqiiations agreed with the exnerimenta.1 values with an average deviation of less tiian0.fpercont. The P-V-T relations of the suuerhratcd vapor were determined by direct measurement 'of the gas density. The apparatus used is shown in Figure 2. It consisted of a 4liter, polished, hollow steel ball, contained in a constanttemperature bath, to which were added the calibrated pressure gages, packless valves, and other attachments required for accurate introduction and removal of the refrigerant. After introduction of a measured quantity of the fluorochloro hydrocarbon, the pressures a t constant density were measured at various temperatures. The range covered on each isometric was from 10" to about 100" C. above saturation. At least five isometrics were measured on each gas, and generally four P-T points were determined for each isometric. Again the pressure gages and thermometers used were calibrated under the exact conditions used during the measurements. The results of these experiments were correlated by means of an equation of state of the Beattie-Bridgeman type ( 1 ) . This equation represents the experimental relations found with an average deviation of less than 0.3 per cent over the entire experimental range. In addition to the measurements made in the steel apparatus, the vapor density at 0.25 to 2.5
Pn -
TR
REDTXEDSATUKATION VOLUMESot F ~ u o n o c x ~ oHYnRocAEeom ~o
Frorrne 3
INASMCCH as the measurement of the heat capacity of a gas is difficult a t best, it was considered advisable to obtain an experimental check on this quantity by measurement of the ratio CJG by a velocity of sound method. Two resonance tubes were made up and excited at one end with frequencies of 800 to 2000 cycles by a vacuum tube oscillator. The points of resonance were determined by measurement of the sound intensity within the tubes by means of a microphone, amplifier, and output meter. One of t.he resonance tubes contained air, and the other, the gas under investigation. Both were contained within a constant-temperature bath. The air tube was used primarilyior calibration of the frequency being employed and to check the constancy of the output of the oscillator. Measurements were made at two temperatures about 50" C. apart. The heat capacity as calculated from the speed of sound in the gas and the equat,ion
INDUSTRIAL A K D ENGINEERING CHEMISTRY
914
of state was compared wit11 the direct measurements made in the flow calorimeter. Table I shows a comparison of the ratio CJC. from the two sets of experimental data. The agreement is considered good; the maximum difference between the two methods is 0.5 per cent.
TABLEI.
CoMrARlSON OF VALUES OBTAINED FOR
DIFFERENT METHOIW C d C r lrom Vclooity ui
Compound
CACIFI CliCIsF
t,
0
C.
47.3 99.7
CChF-CCIF*
%XPLI.
cp
1.178 ].ifin
0.0 0 .5
1.163 1.150 1.149
I . 166 1.155 1.148
0.8 0.4 0.1
47.7 100.1
1.12: 1.121 i .n81
1.130 1.118
0.3 0.3
1.077 1.073
0.4 0.4
77.1
mn.1
1 077
The orthobaric densities and rectilinear diameter of each of the compounds were obtained from the measurements of the vapor pressure, the P-V-T relations of the gas, and the liquid density. The graph of one of these, C C l 8 (Figure 5 ) , illustrates the accuracy with wllich the diameter is represented by a straight line. This curve is typical of that obtained 011 each of the compounds, no curvature being found on the diameter of any of them. This graph also provides the best available method of evaluating the critical density.
1iom
and Equation Deriation, of Slate %
47.7 77.0
1nn.u
CCIrF
Sound 1,178 i.iaa
c /C%
cg/cwBY
VOL. 31, NO. 7
The density of the saturated liquid completes the data required for establishing the thermodynamic system. This was measured in sealed dilatometers or by the use of standardized floats. The liquid densities can he represented in the lower temperature range by equations, but nrar tho critical points these equations arc not usable.
ALTHOUGH the thermodynamic system is identified by the properties already discussed, it was considered necessary to obtain an additional experimental measurement to check the system. The heat capacity of the liquid provides such a check, and measurements of this quantity were therefore made. The experimental procedure was not unusual, the apparatus consisting of a silver-plated steel cylinder suspended within a steel vacuum shell with mcans for heating
CI~CIFI CRCI?F
crisp
CC:I?F-CClh
-160 -136 -111 - 36
-4n 8
x
2x7 47.6
98.0 178.5 198.o
214.1
48.7 51.0 43.2 33.7
0.626 n.522
n.m 0.676
JULY, 1939 34
INDUSTRIAL AND ENGIKEERING CHEMISTRY
I
I
I
d,
10
Eo
I
1
I
=0I8 2 16 14
4 0
I
FIGURE 5 . ORTHOBARIC DENSITIES OF CC1,F Id0
I20
,!I
TEMPERATURE -"C.
F I G V R E4. H E A T CAPACITYO F FLUOROCHLORO HYDROCARBON VAPORS the former electrically and for measuring the temperature (Figure 6). The obvious corrections for, radiation loss as well as heat absorbed due to evaporation or condensation of some of the liquid within the calorimeter cylinder were made. The results are shown in Figure 7. Straight lines of minimum deviation are drawn through the experimental points and show the nearly linear relation of the heat capacity with temperature. The critical temperature was determined by heating the liquid in:. sealed glass tube until the meniscus disappeared. By successive slow heating and cooling of the bath, it was found possible to observe the critical point from
FIGURE 6. LIQUIDCALORIMETER 70
.34 ,
w
2
.
-20
-10
0
IO
20
30
40
50
60
70
80
TEMPERATURE -"C.
FIGURE 7. HEAT CAPACITIES OF LIQUIDFLUOROCHLORO
HYDROCARBONS
either direction; the temperatures for the appearance and disappearance of the meniscus agreed within less than 0.4" C. The critical pressure was estimated by slight extrapolation of the vapor pressure curve to the critical temperature. The critical density was determined by extrapolation of the rectilinear diameter. The normal boiling and freezing points of the four compounds were also determined and are shown in Table 11. Although calculation of the desired thermal properties from the fundamental experimental data can be accomplished by a number of methods, the following procedure was used: The vapor pressure, density of the liquid, and density of the vapor were obtained by direct solution of the equations representing the experimental results. The latent heats of evaporation were calculated from the vapor pressure equatjon and the saturated liquid and vapor volumes by the use of the Clausius-Clapeyron equation. The heat content and entropy of the vapor and liquid were calculated from the thermodynamic equations relating these quantities with the compressibility and thermal expansion coefficients obtained from the equation of state. A check on the over-all thermodynamic consistency of the meas-. ured properties was also carried out. Individual properties were calculated by more than one method from different sets of experimental data by taking advantage of the thermodynamic relations of the properties of fluids.
~ 1 - p I
I-
2
915
l
INDUSTRIAL AND ENGINEERING CHEMISTRY
916
For example, the latent heat of vaporization can be calculated from the vapor pressure equation and the saturated liquid and vapor densities. The same quantity can be obtained by calculation of the heat content of the saturated vapor by means of the equation of state, the heat capacity of the gas a t constant pressure of one atmosphere, and the heat capacity of the liquid. If this is done, the agreement between the two sets of heats of vaporization will be a measure of the over-all consistency of the experimental work. Figure 8 shows the change of this quantity with temperature calculated by the two largely independent methods. The solid lines represent the heat of vaporization calculated from the Clausius-Clapeyron equation; the dotted lines show the same calculation from the heat capacity measurements of both the liquid and vapor as well as the equation of state. The agreement is found to be good, considering that all the errors which may have entered into the work would tend to accumulate in this one property. The mutual agreement of the physical and thermal properties, together with consideration of the normal behavior of these compounds, as shown by the linearity of the rectilinear diameter, the Trouton constants, and the independent checks of the methods on materials whose properties are accurately known, leads to the conclusion that the information developed in this study is of a reasonably high degree of a?curacy. It appears to be sufficiently complete to more than justify use of the data for all normal engineering needs. It is expected that detaiied and complete reports of this work will appear in the journals of this society and that complete thermodynamic tables for these compounds will be
VOL. 31, NO. 7
published. The information developed in this investigation will therefore become generally available in a form readily usable for engineering work.
Aclrnowledgment The authors wish to acknowledge the assistance of W. H. Markwood, Jr., and W. J. Smith in performing some of the experimental work-the former for measurements of the heat capacity of the liquids and both for some measurements of the heat capacity of the vapors. They also wish to acknowledge the assistance of F. B. Downing, of the Jackson Laboratory, E. I. du Pont de Nemours & Company, Inc., whose advice and criticisms were available during the entire prosecution of this work.
Literature Cited (1) Beattie, J. A., and Bridgeman, 0. C.,
J. Am. Chem. SOC.,50, 3133-8 (1928). ( 2 ) Bichowsky, F. R., and Gilkey, W. K., IND. E m . CHEM..23,366-7 (1931). (3) Buffington, R. M., and Fleischer, J., Ibid., 23, 1290-2 (1931). (4) Buffington, R. M., and Gilkey, W. K., Ibid., 23, 254-6 (1931). (6) Buffington, R. M., and Gilkey, W. K., Ibid., 23, 1292-4 (1931). (6) Cope, J. Q., Lewis, W. K., and Weber, J. C., Ibid., 23, 887-92 (1931). (7) Gilkey, W.K.,Gerard, F. W., and Bixler, M. E., Ibid., 23,364-6 (1931). (8) Midgley, T.,Jr., and Henne, A. L., Ibid., 22,542-5 (1930). PRBSZNTZD before t h e Division of Industrial and Engineering Chemistry at t h e 96th Meeting of t h e American Chemical Society, Milwaukee, WIR. Contribution No. 1 from Kinetic Chemicals, Inc.
Relation between Catalvtic *Activityand Size of Particle J
E. W. THIELE Standard Oil Company (Indiana), Whiting, Ind.
A
FEW heterogeqeous catalysts (for example, the plati-
num wires used in the oxidation of ammonia) consist of dense, massive metal. In other cases the catalyst exists in the form of a sol. More commonly, however, the catalyst is in the form of more or less porous grains, ranging from powder size to good-sized pills, often artificially made. I n general, it appears to be tacitly assumed by workers in this field that the reacting fluid penetrates to the pores in the interior of the grains and maintains substantially a constant composition throughout all the pores of a single grain, which is the same as the composition of the bulk of the fluid bathing the grain at the time. It was actually demonstrated in certain cases (1, 2) that further subdivision of the grains produced no change in the catalytic activity. Qualitatively, however, it is evident that the size of the grains cannot be indefinitely increased without ultimately reaching a point a t which the reaction will produce products in the interior of the grain faster than diffusion can carry them away. The reaction will then tend to be confined to the outer layers of the grain, the interior being relatively inactive. As the grain size is further increased, the catalytic activity will
tend to become proportional to the external surface of the grains (or lumps). There appears to be little or no published information on this point. Rideal and Taylor (4) assume that a reduction in grain size will regularly be accompanied by an increase in catalyst activity per unit weight of catalyst. Since the size of the catalyst grains is a practical matter of some importance, it seemed worth while to treat the matter mathematically, with a view to determining the factors that will be of importance and to developing a means of predicting the effect of varying grain size on activity. Although a number of simplifying assumptions were necessary, the results obtained seem to give a correct idea of the influence of various factors, although it was not found possible to determine the behavior of catalysts in this respect independently of experiment. The treatment is also applicable to cases like the water-gas reaction where a gas reacts with a porous solid. In this case, however, the porosity changes with time, so that some additional factors are introduced, and the results apply only during a short period or where there is a countercurrent flow of fluid and solid. The modifications introduced in this case
