2703
Surface Tension of Saturated Anhydrous H$3 (27) C. McAuliffa, J. Phys. Chem., 70, 1267 (1966). (28) R. Archibald, J. Amer Chem. Soc., 54, 3178 (1932). (291 , , C. H. Deal. E. L. Derr, and M. N. Papadapoulos, Ind. Eng. Chem.. Fundam., 1, 17 (1962). (30) J. A. V. Butler, C. N. Ramchandani, and D. W. Thomson, J. Chem. Soc., 280 (.1935). (31) R . S. Hansen, R . A. Miller, and S. D. Christian, J. Phys. Chem., 59, 391 (19551. (32) T. Hi'guchi'and S. S. Davis, J. Pharm. Sci., 59, 1376 (1970). (33) M . S. Mamtora, MS Thesis, University of Kansas. Lawrence, Kan., 1969. (34) R . Aveyard and R . W, Mitchell, Trans. Faraday Soc,, 64, 1757 (1968). (35) L. Benjamin, J. Phys. Chem., 68, 3575 (1964). (36) J. M. Corkill. J. F . Goodman, and J. R . Tate, Trans. Faraday Soc.. 63, 773 (1 967).
(37) F. Franks, Nature (London), 210, 87 (1966). (38) J. H. Clint, J . M. Corkill, J. F. Goodman, and J. R . Tate, J. Colloid Interface Sci.. 28, 522 (1968). (39) D. J. Currie, C. E. Lough, R . F. Silver, and H. L. Holmes, Can. J. Chem., 44, 1035 (1966). (40) R. Lumry and S. Rajender, Biopoiymers, 9, 1125 (1970). (41) J. A. V. Butler, Trans. FaradaySoc., 33, 229 (1937). (42) J. C. McGowan, P. N. Atkinson, and L. H . Ruddle, J. Appi. S c i , 16, 99 (1966). (43) E. Deutsch and C. Hansch, Nature (London). 211, 75 (1966) (44) R. P. Quintana and W. R . Smithfield, J. Med. Chem., 10, 1178 (1967). (45) J. Iwasa, T. Fujita, and C. Hansch, J. Med. Chem., 8, 150 (1965). (46) C. Hansch and S. Anderson, J. Org. Chem.. 32, 2583 (1967). (47) D. J. Lamb and L. E. Harris, J. Amer. Pharm. Assoc., 49. 583 (1960),
Surface Tension of Saturated Anhydrous Hydrogen Sulfide and the Effect of Hydrogen Sulfide Pressure on the Surface Tension of Water Carlyle S. Herrick" and George L. Gaines, Jr. Genera/ Electric Corporate Research and Development, Schenecfady, New York 12307
(Received April 26, 1973)
Publication costs assisted by the General Electric Research and Development Center
The surface tension of anhydrous liquid H2S in the temperature range 25-40" varies from 11.3 to 8.7 dyn/ cm. The Guggenheim equation fits these data precisely. Measurements of the surface tension of water in contact with HzS at pressures up to 300 psi and from 25 to 40" show that H2S causes a greater reduction in surface tension with pressure than any gases previously studied. The data for H2S and for other gases suggest that the maximum lowering of surface tension corresponds to the adsorption of one close packed monolayer of gas molecules on the water surface. H2S and C02 reach monolayer adsorption at about l/2 saturation pressure. If general, this behavior may permit order-of-magnitude prediction of surface tension effects from values of the pressure required for phase change. We have also obtained an estimate for the contact angle of HzS-saturated water on the solid H2S hydrate.
Introduction In the GS process for enriching the deuterium content of natural waters,l liquid water and gaseous H2S under pressure are contacted on a large scale and surface properties probably play an important role. Slowinski, et a1.,2 have shown that pressurized gases may substantially reduce the surface tension of liquids. In a limited study (below 1 atm) Tamamushi3 found that for equal pressures hydrogen sulfide gas had a larger effect on surface tension than several other gases. Rough estimates by Heuer4 showed that additions of H2S (up to 10 mol 70in the gas) caused increased surface tension reduction in the methane-water system. The authors of ref 2 and 3 suggest that the more condensable or more soluble gases have the larger effect. Accordingly, there is both practical and scientific interest in the effect of pressurized H2S on the surface tension of water. This report presents and discusses such measurements. We also present the results of surface tension determinations on liquid H2S. Experimental Section Surface tension measurements were made using the differential capillary rise method. Two Pyrex glass capillary
tubes, of radius 0.025 and 0.0595 cm, were mounted side by side in a Teflon holder which could be inserted vertically in a two-window pressure vessel. This vessel consisted of a commercial liquid level gage with top and bottom valves and fittings suitable for sample manipulation and flow-through cleaning. Only Teflon, type 316 stainless steel, and Pyrex glass were used in the assembly. The completed vessel had a maximum working pressure of about 2000 psi. In operation it was immersed with the capillaries vertical in a 20-gal controlled-temperature water bath. A thermistor sensor working through a commercial narrow proportional band solid state controller to an electric immersion heater controlled bath temperature to k0.01". Bath temperature was measured using a mercury thermometer calibrated a t the Bureau of Standards to 0.01". Two centrifugal pumps circulated bath water around the pressure vessel at a total rate of about 10 gpm. The 22-lb mass of the pressure vessel constituted a large thermal inertia against small temperature changes so that satisfactory isothermal conditions were achieved. Working surfaces were cleaned before use by a warm (heat of mixing) 2 / 1 mixture of 98% H2S04 and 36% HN03. The capillaries were immersed in the acid for 16 The Journal of Physical Chemistry, Vol. 77, No. 22, 1973
2704
Carlyle S. Herrick and George L. Gaines, Jr.
hr and the pressure vessel was completely filled with the acid for the same time. Both were then rinsed ten times. Each rinse consisted of a complete filling with distilled water followed by complete emptying. The capillaries were emptied by wicking onto a mechanically pulped paper toweling which gave no evidence of back-contamination. In preliminary experiments, significant contamination was encountered from a presumed vacuum pump oil aerosol in the laboratory air which replaced the water during the emptying phase of each rinse. Accordingly, the replacement air was passed through an activated charcoal bed, and this eliminated the contamination. After rinsing, the component parts were either filled with HzS saturated distilled water and assembled if the succeeding measurement was to be HZO(~\-HZS,,,, or dried thoroughly with noncontaminating nitrogen and assembled if the succeeding measurement was to be H2S,,I-H2S,, . The cleaned capillaries were calibrated in six liquids, acetone, hexane, benzene, toluene, chloroform, and water, against air by determining the capillary constant C for the equation
y
= C(p,
-
p,,) Ah
(
-3-2 + 3) 3
rt
13 I'J
r
I
7t
(1)
where 7 is surface tension in dynes/cm, p1 and pv are the densities of liquid and vapor phases, respectively, in grams/cm3, Ah is the difference between heights of liquid columns in the capillaries in cm, and rl and rz are the capillary radii in cm, The average value of C for all six liquids was 22.94 with a standard deviation of' 0.17. Water-air surface tensions calculated in this wav are slightly higher than published values demonstrating that the laboratory distilled water source was free of surfacecontaminating agents. Distilled water for use in H2S measurements was degassed in a Pyrex vessel by boiling. While still boiling the liquid was transferred to a separate Pyrex storage vessel maintained under 1 atm of HZS pressure by a continuous gas bubbler system leading to a flare. After the water cooled at room temperature under the H2S atmosphere it was ready for experimental use. The H2S was Linde C P grade of 99.6% purity. This H& content was confirmed by gas chromatography on the gas actually used. COa was the major contaminant. Since H2S is stored as a liquid, a vapor takeoff from the cylinder causes COZ enrichment resulting in variable gas composition. To maintain the rated purity it was necessary to remove H2S from the storage cylinder as a liquid, then vaporize the entire sample removed. During the measurements on liquid H2S, it was noted that the liquid, initially clear and colorless. slowly developed a light brown color in the experimental chamber. The cause of this change, which did not affect the surface tension readings, is unknown. (It is possible that it resulted from the formation of some colloidal sulfur. due to traces of residual oxygen.) Pressure measurements were made using an Ashcroft 1082 test gauge with 0-300 psi range, i psi divisions, and accuracy 0.2570 of full scale. Capillary column height measurements were made to 0.0001 cm using a Gaertner traveling microscope. Liquid and vapor densities given by based on the experimental data of Janik6 Galley, ~t al and Reamer, e t al , 7 were used. Surface tension values are estimated accurate to k0.25 dyn /cm. The Journal o f Physical Chemistry, Vol. 77, No. 22, 7973
l5
7c
5
0
10
20
30
40
50
60
70
TEMPERATURE, " C
Figure 1. Surface tension of liquid HzS against its saturated vapor.
Results Surface Tension of Liquid Hydrogen Sulfide Against Its Vapor Liquid HzS was charged directly t o the experimental vessel without a phase change. Surface tension values for liquid H2S against its saturated vapor for temperatures from 25 to 40" are shown in Figure 1. Surface Tension of W a t e r Against H3'drogen Sullide Vapor. The system was equilibrated a t 1 atm H2S pressure before each set of readings. After each increase in pressure the capillary rise equilibrated within 1 min, however, the liquid columns in the capillaries required weeks to equilibrate to the new density. Consequently measurements a t pressures above 1 atm were made in rapid succession until the pressure excursion was completed. Then the 1 atm liquid density was used in eq 1 a t all pressures. It was assumed that the surface tension was fully determined by the composition of a thin layer of liquid immediately below the surface and that when this thin layer was in equilibrium with the gas phase that equilibrium values of the surface tension were obtained. After completion of a pressure excursion the initial capillary rise value was reproduced upon reequilibrating the water surface at 1atni for 60 hr. Surface tensions of water against its own vapor were estimated by correcting the accepted values against air a t 1 atm to zero pressure using the slope of the nitrogen pressure data measured by Slowinski, e t a1.,2 and these values were used for zero pressure of H2S. The experimental results are given in Figure 2. At each temperature a straight line is drawn through the zero pressure surface tension using the slope calculated by a least-squares fit of the experimental data taken in the presence of HzS. The resul-
2705
Surface Tension of Saturated Anhydrous HzS equationlo
8 o ~ ~ r - r r r m T T - T T T ; - - 7 - ~ - -
1
o 0
Y
__ EO 2
r 40,-
= io(l -
fits our data precisely, using the critical temperature Tc = 373.6"K, and yo = 80, as shown in Figure 1. Galley, et a1.,5 estimated yo = 79.02 from the critical constants. Liquid H2S surface tensions calculated from the parachor value of 80.1 recommended by Quaylell were about 20% smaller than the experimental values. Surface Tension of Water Against H2S Vapor. The effect of compressed gases on the surface tension of liquids has not received extensive study. Slowinski, et al ,2a who measured the pressure effect for six gases other than H2S interpreted their results in terms of the approximate Gibbs equation
II
HYDRATE FORMATION
"b,
E)
11 9
EXPERIMENT THEORY
35°C
5Or
-
40
40
c
0 3 0
2
4
' '
6
/
8
1
'
IO
l
l
12
l
14
,
l
16
l
l
18
l
20
~
1
1
~
22
H 2 S PRESSURE, ATM
Figure 2. Surface H2S vapor.
tension of water saturated with HzS against
tant fit is within experimental error a t all temperatures. Values for the surface tensions indicated by the lines in Figure 2 can be calculated within 0.01 dyn/cm using
+
75.828 - 0.14% - 2.6 X 10-4t2 6.38997P 0.821572Pt 3- 2.62502 X 10-'Pt2 - 2.6396 X 10-lPt'
y
=
(2) where y is in dynes/cm, P is in atmospheres of H2S, and t is in degrees centigrade. At 25" and 12.25 a t m a crystalline solid hydrate was observed. This temperature is about 0.5" above the position of the hydrate phase line a t this pressure as derived from the data of Selleck, et aL8 Hydrate formation in the capillaries was accompanied by a large decrease in column height. By suitable repeated manipulation of the degree of supersaturation solid hydrate was formed on the inner walls of both capillaries in layers so thin that no change in diameter could be detected by the microscope. Capillary rise measured against the hydrate covered surface Ahh compared with that against glass Ahg permits an estimate of the contact angle between H2S saturated water and the H2S hydrate.
r -7"l ) / ( A h , 3 3 1 The observations fell within the limits 62" < 0 < 69". Hydrate formation on the outside of the large capillary a t the immersion line limited visibility and prevented a closer estimate.
Ahh -
r2
+
$+
They observed substantially linear variations of surface tension with pressure, as we have for H2S. They pointed out that eq 5 then implies the surface excess of the gas r2t11is proportional to pressure. We have estimated in columns 1-5 of Table I the maximum values of ryl) from eq 5 evaluated a t the pressure corresponding to the first pressure induced phase change, on the premise that Tz'l' cannot continue to increase beyond such a point. By applying the monolayer assumption (i.e., A = l/rz'l!), the surface area per adsorbed molecule was estimated in column 5 . These estimated areas are close to molecular areas in close-packed monolayers arid fall in the sequence expected for the gases studied. Clearly these calculations are only semiquantitative because of the long extrapolations required to extend the hydrogen, nitrogen, and methane data to a phase change point, the ideal gas approximation implicit in eq 5, and the observable small deviations from linearity in the data for some gases. In addition Tz:" is less than the surface concentration of adsorbed gas by the bulk concentration of dissolved gas although the latter contribution is negligible in the cases treated here. Both Masterton, Bianchi, and Slowinski2b and Hough, Wood, and Rzasa12 had concluded from similar analyses of their data that maximum surface excesses approximated monolayer coverage in nitrogen-, methane-, and argon -water systems. The experimental results of these two studies (ref 2 and 12) are in qualitative agreement for the systems and pressure range covered in both; a t higher pressures, however. the experimental artifacts observed by Hough, et al., make any conclusions doubtful. Ericksson13 has analyzed the data of Slowinski, et a1.,2a in thermodynamic terms. He concluded that a monolayer model, with adsorbed gas molecule molecular areas similar to those observed in adsorption studies on solids, adequately fitted the data, but he did not consider the question of attainment of saturation adsorption. Surface tension measurements on dilute solutions are generally analyzed by another approximate Gibbs equation
Discussion
Surface Tension of Liquid H2S. Prior determinations by Devyatykh, et a1.,9 appear to be limited to temperatures below the normal boiling point (-60"). The Guggenheim
where Nz is the mole fraction of gas dissolved in the liquid. Equation 6 is equivalent to eq 5 when Henry's law is obeyed. Since the more soluble gases CUz and H2S do deThe Journal of Physical Chemistry, Vol. 77, No. 22, 1973
Carlyle S. H e r r i c k and George L. Gaines, Jr.
2708 TABLE 1: Surface Excess of Pressurized Gases at Water Surfacese
Gas
-dy/dp, erq!cm2:atm
T, '6
0.016 0.068 0.106a 0.71 0.49 I .86 1.74 1.50 1.33
Limiting pressure, atm
9200b 1 762c 432c 63.5d 41d 13,7c
22.3d 25.1d 28.1d
Max r~{'1/1O'~, Area permolec~les/cm~/10'~ moiecule, A'
35 29 1i.l 11.0 4.9 6.2 9.4 9.1 9.1
2.8 3.4 9.0 9,s 20.4 16.1 10.6 11.0 11.0
MaxIr2"' (Figure 3 ) , moieculesicmz
8.8 X 1014
4 . 9 x 1014 5.1 x 1 0 1 4
a Slope at highest experimental pressure. Formation of ice Vi, P. W.Bridgeman, J. Franklin lnst.. 177, 315 (1914). Hydrate formation; NZ and CHI, D. R. Marshall, S. Saito, and R. Kobayaskl, AIChE J., 10, 202 (1964); HzS, ref 5 and 16. Liquefaction: CO2 and C2H6; S . S. Byk and V . I . Fomina, Russ. Chem. Rev.. 37, 469 (1968). COz; 'Handbook of Chemistry and Physics," 44th ed., Chemical Rubber Publishing Co., Cleveland, Ohio. p 2428. C2HE;D. R . StLill. Ind. Eng. Chem. 39.517 (1947). H2S; ref 5. e Surface tension data for H2S, present work; ail others, ref 2a.
solubilities given by Stewart and M ~ n j a 1 . lFor ~ pressures above 45 atm, their solubility data have been extrapolated by fitting an equation of the form suggested by Zel'venskiilS
s
o\
\
60
-Pn N, Figure 3. Surface tension of water as a function of mole fraction of gas in solution, illustrating the constant limiting values of - d y / d In N2. The CO2 data are from ref 2a.
viate from Henry's law it is of interest to compare results based on eq 5 and 6. Solubility data for HzS in water have been obtained by Wright and Maass14 and Clarke and Glew15 a t low pressures, and by Selleck, et aL,8 a t higher pressures. While doubt has been expressed about some of the data,5J5 and extrapolation appears uncertain, we have extrapolated the data of Wright and M a a d * up to 5 atm, and used the equation fitted by Burgess and Germann16 to the Selleck, et ~ l . results , ~ to estimate solubilities a t higher pressures. Figure 3 shows our surface tension values at 25 and 40" plotted against the logarithm of HzS mole fraction in solution. Also shown are the data of Slowinski, et U L . , ~ for ~ the surface tension of water in the presence of CO2, using The Journai of Physical Chemistry, Vol. 77, No. 22, 1973
= 0.1416P - 1.127 X lOP3P2 for solubility in grams of CO2 per 100 grams of HzO and P in atm. As the straight lines in Figure 3 indicate, in all three cases the surface excess appears to reach a constant maximum value a t pressures above about half the phase change pressure. These constant maximum T z ' l ' values (column 7 of Table I) are consistent with essentially closepacked gas molecules in a surface monolayer. For H2S, compare the molecular area of 20 x2 (monolayer assumption) with the estimate of 17-17.5 x2 based on the 2,$ power of the molecular volume obtained from the liquid density a t 25-40" given in re€6. In summary, when the available data are analyzed by eq 5 one finds approximately a close-packed monolayer of gas adsorbed on the liquid surface a t the gas pressure corresponding to the first pressure induced phase change. By eq 6 one finds that COz and H2S reach the close-packed adsorbed monolayer state a t pressures considerably lower than the phase change pressure. Surface tension variations in other systems may be estimated by assuming a surface excess corresponding to monolayer coverage a t phase change pressure. Conversely the phase change pressure may be estimated if the surface tension has been measured. There is no adequate molecular theory for surface tensions of aqueous solutions to guide an extended analysis of the measurements. Attempts to calculate solution surface tensions from the values for H20 and H2S using the mole fraction average or several lattice-theory e q u a t i o n ~ l ~ - ~ ~ were unsuccessful.
Acknowledgment. We thank Mr. J. A. Murphy for assistance with a portion of the measurements. References and Notes (1) W. P. Bebbington and V . R. Thayer, Chem. Eng. Progr.. 55, (No. 9), 70 (1959). (2) (a) E. J. Slowinski. Jr,, E. E. Gates, and C. E. Waring, J. Phys. Chem.. 61, 808 (1957); ( b j W. L. Masterton, J. Biarlchl, and E. J. Slowinski, Jr,, J. Phys. Chem., 67, 615 (1963). (3) B. Tamamushi. Buii. Chem. SOC.Jap., 1 , 173 (1926). (4) G. J. Heuer. Ph.D. Dissertation, University of Texas, 1957: Diss. Abstr., 17, 2558 (1958). (5) M. R. Galley, A . I. Miller, J. F. Atherley, and M. Mohn. AECL-4255, Chalk River Nuclear Laboratories, 1972,
Thermochemistry of the Bromination of CC14
2709
(6) J. Janik, Acta Phys. Poion., 23, 487 (1963). (7) H . H. Reamer. B. H. Sage, and W, N. Lacey, ind. Eng. Chem , 42, 140 ( 1950) (8) F. T. Selleck, L. T Carmichael. and 8. H. Sage, Ind. Eng. Chem., 44, 2219 (1952). (9) G. G. Devyatykh, A , D. Zorin, and i. V. Runovskaya. Doki. Akad. NaukSSSR, 188,1082 (1969). 10) E. A. Guggenheim, J. Chem. Phys , 13,253 (1945). 11) 0. R. Quayle, Chem. Rev.. 53,439 (1953). 12) E. W. Haugh, B. 13,Wood, and M. J. Rzasa, J. Phys. Chem.; 56,
(13) (14) (15) (16) (17) (18) (19) (20) (21)
996 (1952). J. C.Ericksson, Acta Chem. Scand . 16,2199 (1962). R . H . Wright and 0. Maass. Can. J. Res., 6, 94 (1932). E.C.W. Clarke and D. N. Glew, Can. J. Chem.. 49, 691 (1971). M .P. Burgess and R. P. Germann.AIChEJ., 15, 272 (19691. P. B. Stewart and P. Munjal, J. Chem Eng. Data, 15, 67 (1970). Y. D. 2el'venskii.J. Chem. ind. ( U S S R ) . 14. 1250(1937). J. W. Belton and M.G. Evans, Trans. Faraday Soc.. 41. 1 (1945) I. Prigogine and J, Marechai,J. CoiioidSci. 7, 122 (1952). G. L. Gaines. Jr., Trans. Faraday Soc., 65,2320 (1969)
Thermochemistry of the Bromination of Carbon Tetrachloride and the Heat of Formation of Carbon Tetrachloride G . D. Mendenhall,' D. M. Golden, and S. W. Benson" Department of Thermochemistry and Chemicai Kinetics, Stanford Research institute, Menio Park. Caiifornia 94025 (Received May 29, 7973)
+
+
At 560°K the equilibrium constant for the reaction Brz CC14 c BrCl CBrC13 was determined as 0.0046 f 0.001 by combined optical and chromatographic techniques. From this result and literature data we calculate the heat of formation difference AHf"(CC14.g) - 1Hf"(CHC13,g)= 2.27 f 0.3 kcal/mol at 298", independent of calorimetry. Based on AHf"(CHC13.g) = -24.66 f 0.3 kcal/mol, we obtain AHf"(CC14,g) = -22.4 & 0.4 kcal/mol, a higher value than obtained from many earlier studies. The CCl3-Cl bond strength is estimated as 70.4 f: 1 kcal/mol.
Experimental Section
The Brz partial pressure was determined from the optical densities a t 550 and 570 nm, while the OD (340 nm) gave the BrCl pressure after subtracting out the contribution from Brz a t this wavelength. Mixtures of Clz and Brz were added to the cell to obtain 0: f P ( B r C 1 Torr) /OD (340 nm). In this calibration and in some of the reaction runs a correction was made for the presence of free Clz from the equilibrium 2BrC1 e Brz + Clz, which has a value of 0.375 a t 560OK.4 At the partial pressures of our experiments, pressureindependent c ( A nm) were Brz: 64.5 (550), 95.2 (570); BrC1: 48.6 (340); Clz: 49.1 (340); CBrC13: 174.4 (300). We did not relate these to published values because our instrument slit width varied over a wide range. The corrosive nature of the halogens made special handling necessary. Bromine (and probably Clz) reacted with hot silicone grease, so all stopcocks in direct contact with the hclogens were unheated and were lubricated with a silica gel-halocarbon mixture.9 The ratio CBrC13/CC14 was determined directly by condensing out the reaction mixture a t -196" followed by vpc analysis by peak height ratio (F & M Model 720, thermal conductivity detector, 2 ~1 samples on a 10 ft x 0.25 in. column of 20% di(2-ethylhexyl) sebacate on 60-80 m Chromasorb a t 140"). Mixtures of known composition of the two halocarbons were injected for calibration, and it was shown that added Br2 or Clz did not effect the ratios from these standard mixtures. The partial pressures were obtained from the ratio and the relation P'(CBrC1,) + Pe(CC14) = PO(CC1,)
The spectrophotometric procedure and apparatus ( a modified Cary Model 14) has been described previously.8
The analytical procedure was tested by adding known pressure of CBrC13 and cc14 to the cell, condensing out as
Introduction Values for the heat of formation of CC14 in the literature encompass an unfortunately wide range. Cox and Pilcher2a and Stull, Westrum. and Sinke2b have combined a number of calorimetric experiments with newer data and arrived a t Afffo298(CC14,gas) = -25.2 =k 1.5 and -24.0 kcal/mol, respectively. Many of the early values depended on 2,Ht"(CHC1s). Hu and Sinke3 recently reported results of rotating bomb calorimetry which lead to an assigned AHf"(CC14,g) = -22.9 f 0.5 kcal/moL4 Lord and Pritchard,5 however, arrived a t a value of -27.4 kcal/mol from an equilibrium study of COC12. cc14, and COz in the presence of excess Clz. Their determination required a rather elaborate analytical procedure. Sullivan and Davidson6 accurately measured the equilibrium constant for the reaction Br2
+
CHCl:,
A
e
HBr
+
CBrCl,
and the thermochemistry of the system has been evaluated by B e n ~ o n We . ~ can estimate the heat of formation difference between CHCl3 and cC14 from K A and the equilibrium constant for Br2
+
CC1,
B
-+
BrCl
+
CBrC1,
From a choice of AHf"(CHC13,g) one obtains a precisely related AHf"(CC14).
The Journai of Physical Chemistry, Voi. 77, No. 22, 7973
