September 1959
INDUSTRIAL AND ENGINEERING CHEMISTRY
Goring, G. E., Curran, G. P., Zielke, C. W., and Gorin, Everett, I b i d . , 45, 2586 (1953). Johnstone, H. P., Chen, C. Y., and Scott, D. S.,Ibid., 44, 1564 (1952). Key, A., Gas Research Board, London, Pub. 40 (July 1948). Laidler, Keith J., “Catalysis,” p. 75, Reinhold, New York, 1954. Long, F. J., and Sykes, K. W., J. chim. phys., 47, 361 (1950). Long, F. J., a n d Sykes, K. W., Proc. Roy. SOC.(London), 193A, 377 (1948). Lowry,‘ H. H., “Chemistry of Coal Utilization,” p. 904, Wiley, New York, 1945. Mayers, M. A., J. Am. Chem. Soc., 56, 79 (1934). Petersen. E. E., Walker, P. L., Jr., and Wright, C. C., IXD.ENG. CHEM..47. 1629 (19551. Schaeffer, W’. D., Srnith; W. R., and Polley, R.1. H., I b i d . , 45, 1721 (1953). Wagman, D. D., Kilpatrick, J. E., Taylor, W. J., Pitzer, K . S.,
1749
and Rossini, F. D., Natl. Bur. Standards, Research Paper RP 1634 (1945). (19) Walker, P. L.,Ji., Foresti, R. J., Jr., and Wright, C. C., IND. ENQ.CHEM.,45, 1703 (1953). (20) Walker, P. L., Jr., McKinstry, H. d.,and Pustinger, J., Ibid., 46, 1651 (1954). (21) (22) (23) (24)
Warner, B. R., J. Am. Chem. SOC.,65, 1447 (1943).
Ibid., 66, 1306 (1944). Wheeler, A,, Advances in Catalysis, 3 , 249-327 (1951). Zeldowitch, J., Acta physicochim. (U.R.S.S.),1 , 961 (1935).
RECEIVED for review July 2 6 , 1954. ACCEPTED March 21, 1955. Presented beforr the Division of Gas and Fuel Chemistry a t the 114th Meet8t. Louis, Mo. From a thesis by ing of the AXERICAK CHEMICAL SOCIETY, J. M. P. in partial fulfillment of the requirements for the Ph.D. degree in fuel technology. Contribution No. 54-48 from the College of Mineral Industries. The Pennsylvania State University.
Vapor-Liquid Equilibria by Radioactive Tracer Techniques SYSTEM CARBON DIOXIDE-HYDROGEN SULFIDE-SODIUM CARBONATESODIUM BICARBONATE-SODIUM SULFIDE-WATER I 0.50N). perimental accuracy, K3 is pressure-independent over the range of pressure applied. For C = 0.01 to 0.501y Figure 6 indicates that the temperature dependence of K Smay (24) be given by K1 = 2.984 X lop6 C-0.128exp (2586/T) Table IV.
Values Obtained for Equilibrium Constant, K1 Dev. 9% Dev.
O
. . I
K3(C)=
For C = 0.50 to 2.00N
K1 = 1.591 X 10-5
exp (2729/T)
(25)
100
It was found that all the experimental equilibrium constants given in Table I V could be predicted by Equations 24 and 25 within 3.4y0.These correlating equations apply well over tem-
eo
C-o.362
perature ranges of 20" to 65" C. and concentrations of 0.01 to 2.00N and i t is recommended that they be used in estimating the vapor-liquid equilibrium behavior of the systeni sodium carbonate-bicarbonate-water. For the purpose of setting up equilibrium diagrams that may be used in design computations, the general correlating equations given in Equations 24 and 25 were rearranged in the following manner to yield practical expressions. If the quantity f is defined as the fraction of sodium ions present as sodium bicarbonate in a sodium bicarbonate-carbonate solution, Equation 12 defining K,may be rewritten as
%'
exp ( m " / T )
60
40
I5
I f 20
I
dj 0 3
10
$ 8 W
96 A
where C again represents the total sodium ion concentration. Combining Equations 24 and 26 and solving for Pcol yields
9 5
4
For C = 0.01 t o 0.50N 2
Similarly, a combination of Equations 25 and 26 gives
For C
=
0.50 to 2.OON I
These equations give Pco2,in millimeters of mercury, as a function of the sodium bicarbonate-carbonate fraction in tlhe solution for the system sodium carbonate-bicarbonate-water a t any tem-
FRACTION NoHC03
Figure 8. Partial pressure of carbon dioxide over 1.00N sodium carbonatebicarbonate solutions
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INDUSTRIAL AND ENGINEERING CHEMISTRY
Vol. 47, No. 9
chemical absorption equipment in which aqueous solutions containing sodium sulfide, sodium bicarbonate, and sodium carbonate are to be carbonated with carbon dioxide the equilibrium partial pressures of carbon dioxide and hydrogen sulfide may be calculated from Equations 27, 28, and 33. For a complete design, however, a knowledge of mass transfer rates is necessary, The kinetics of the gas-liquid reactions defined by Equations 1 and 2 as well as the mass transfer rates are now being studied in these laboratories. SUMMARY
02
0.4
0.6
1.0
20
C, SODIUM ION NORMALITY Figure 9. Dependence of Ka on sodium ion concentration
where m" is a constant having a value equal to the slope, (dlogK3/dT-'), or - ( A H 3 / R ) . The dependence of K a on the sodium ion concentration, C, is illustrated in Figure 9. The points utilized in establishing the isotherms on this plot represent the arithmetic averages of the experimental values for K 3 for each fixed temperature and for sodium ion concentrations up t o and including l.OOLVv.The 2.00X points were obtained by an extrapolation of the data presented on Table I1 coupled with related Ka values obtained by Berl and Rittener ( 1 ) for 2.00M solutions. The solubility limit line, above which Ka cannot be defined in practice, is also shown. This graph indicates a relationship of the form
K q T ) = a'' ( C b " )
To facilitate vapor-liquid equilibrium measurements of the system carbon dioxide-hydrogen sulfide-sodium carbonate-sodium bicarbonate-sodium sulfide-water over a temperature range of 20' to 65' C., a pressure range of 50 to 3600 mm. of mercury, and a sodium ion concentration range of 0.10 to 1.00N, analytical techniques were developed for the continuous analysis of the vapor and liquid phases. These techniques involved a scintillation counting method whereby the partial pressure of hydrogen sulfide tagged with sulfur-36 could be determined continuously in the vapor phase, and p H and conductance measurements, which, coupled with the use of calibration diagrams, gave a continuous indication of the ratio of sodium carbonate, sodium bicarbonate, and sodium hydrosulfide in solution in the liquid phase. These techniques were found to be accurate to within 3%.
(30)
where b" is a constant representing the slope (dlog Ka/dlog C). h combination of Equations 29 and 30 results in the general relationship for the dependence of K S on temperature and sodium ion Concentration
Ka = n" Cb" exp ( m " / T ) where
1z"
(31)
is an arbitrary constant.
By utilizing the method of least squares this equation was fitted to the arithmetic mean values of K 3 for any concentration and temperature listed in Table 11, and resulted in the following correlating equation:
K3 = 9.746 X
lo-* Co.125exp (2275/7')
(32)
T h e values of KBcalculated from this equation agreed with the experimental data within 2 % on the average, with the maximum deviation being 5.5%. The use of this equation for the determination of equilibrium data for the system carbon dioxide-hydrogen sulfide-sodium carbonate-sodium bicarbonate-sodium sulfidewater over temperature ranges of 20" t o 6 5 " C. and 0.10 to 1.00 sodium ion normalities may therefore be recommended. A combination of Equations 15 and 32 yields the following expression for the ratio ( P H 2 s / P c o 2 )
where C again represents the total sodium ion concentration, normality, and T the absolute temperature, O K. A typical plot of this expression is shown in Figure 10 for sodium ion concentrations of 0.10, 0.50, 1.OO, and 2.00Ar. Application to Carbonation Processes. For the design of
0
2
4
6
e
IO
12
( NaHSV (NaHC03)
Figure 10. Partial pressure ratio of hydrogen sulfide and carbon dioxide over sodium hydrosulfidebicarbonate solutions
The equilibrium constants which were experimentally obtained by an application of these analytical methods 'were found to be reproducible to within 3%. Enthalpy changes accompanying the reactions described by these equilibrium constants were calcylated from the experimental data and agreed with the heats of reaction calculated from published heats of formation within 6%. By treating the vapor-liquid equilibrium data statistically, enipirical correlations were obtained for the equilibrium constants
September 1955
INDUSTRIAL AND ENGINEERING CHEMISTRY
and KS as functions of temperature and sodium ion concentration. These correlating equations are presented in Equations 24, 25, and 32 and were found to reproduce all evperimental data available t o within 4%. The correlating equation developed by Harte and others ( 1 0 ) for the equilibrium constant defined by Equation 1 was found to yield acceptable results only over sodium ion concentrations of 0.50 to 2.00N and temperatures of 30” to 50” C. *Itlower concentration ranges and higher temperatures equilibrium constants predicted by this equation disagreed b y as much as 73.5’35 with published experimental constants. The results of this work indicated that two correlating equations were necessary for this equilibrium constant, one for a sodium ion concentration range of 0.01 to 0.50N and the other for a range from 0.50 to 2.00N. In both these equations the temperature was expressed as an exponential function of the equilibrium constant rather than as the linear function presented by Harte. Practical expressions derived irom these correlating equations are presented in Equations 27, 28, and 33 and will be of use in the design of equipment for industrial carbonation and the recovery of hydrogen sulfide. However, additional data on mass transfer rates are necessary to make final design calculations possible. K1
ACKNOWLEDGMENT
The authors wish to thank Albert Wakefield for his help with the electrical circuitry and the University of Washington Engineering Experiment Station for financial support during the course of this work. NOMENCLATURE
sodium ion concentration, normality counts per minute heat of reaction, calories per mole eauilibrium constant partial pressure of COz,mm. of mercury partial pressure of H& mm. of mercury partial pressure of HzO, mm. of mercury gas constant, calories per O C. per mole absolute temperature; O K . Bunsen solubilitv of COZ in water, liter S T P Der liter total molal Concentration of CO, a t 1 atm. total molal concentration of €IzSa t 1atm. fraction sodium ion as NaHCOJ in NaHCOs-NazCOa solutions temperature, O C. ionic activity coefficient of ?r’aHCOt ionic activity coefficient of ?JazC03 activity coefficient of H&03 ~
activity coefficient of H,S
y’
=
6
= ionic activity coefficient of SaHS = total pressure, mm. of mercury
T
1757
LITERATURE CITED (1) Berl, E.. and Rittener, A., Z. angew. Ghem., 20, 1637 (1907).
(2) Bichowsky, F. R., “International Critical Tables,’’ 1st ed., vol. V, p. 169, McGraw-Hill, New York, 1929. (3) Birks, J. B., “Scintillation Counters,” McGraw-Hill, New York, 1953. (4) Chem. Eag., 59, 165 (January 1952). (5) Comstock, C. S., and Dodge, B. F., IND.ENG.CHEM.,29, 520 (1937). (6) Datar, D. S., Indian Patent 45,217 (March 1, 1952). (7) Dryden, I. G. C., J . SOC. Chwm. Ind., 66, 59 (1947). (8) Gamson, B. W., and Elkins, 1%.H., Chem. Engr. PTogr., 49, 203 (1 953). (9) Guggenheim, E. A , , “Modern Thermodynamics by the Method of Willard Gibbs,” p. 27, hlethuen and Co., London, 1933. (10) Harte, C. R., Baker, E. M., and Purcell, H. H., IND.ENG. CHEM., 25,528 (1933). (11) Jones, G., and Bradahaw, B. C., J . Am. Chem. Soc., 55, 1780 (1933). (12) Koeller, R. C., and Drickamer, H. G., J . Chem. Phys., 21, 267 (1953). (13) McCoy, H. N., Am. Chenz. J., 26,437 (1903). (14) Mai, K. L., Ph.D. thesis, University of Washingtoil, Seattle, Wash., 1954. (15) 3Iai, K. L., and Babb, A. L., ‘Vucleonics, 13, No. 2, 52 (1955). (16) . , Nakarov, S.2.. and Krasnikov. S.N., Bull. acad. sci. U.R.S.S.. Classe sci. math. nat., Ser. chim., 1937, 363-81. (17) Repa, A . G., and Danilchenco, E. E’., SfekEo i Keram, 6, KO.9, 10 (1949). (18) Robb, W. L., and Drickamer, H. G., J . Chem. Phys., 19, 1504 (1952). (19) Sawyer, F. G., and Hader, R. N., IND.END. CHEM.,42, 1938 (1950). (20) Scott, W. W.,“Standard lLIet.hods of Chemical Analysis,” 5th ed.. vol. 11. I). 218: Van Nostrand, New York. 1946. (21) Timmerhaus, K..D., and Drickamer, H. G ., 6.Chem. Phys., 19, 1242 (1951). (22) I b i d . , 20,981 (1952). (23) Treadwell, F. P., and Hall, W. T., “Analytical Chemistry,” 9th ed., vol. 11, p. 615, Wiley, New York, 1948. (24) Tung, L. H., and Drickamer, H. G., J . Chem. Phys., 20, 6, 10 (1952). (25) Walker, -4.C., Bray, U. B., and Johnston, J., J. Am. Chem. Soc., 49,1235 (1927). (26) Weissberger, A., “Physical IMethods of Organic Chemistry,” 2nd ed., vol. I, Part 11, p. 1654, Interscience. New York, 1949. RECEIVED for review December 6 , 195C. ACCEPTED April 4, 1955. Presented before the Division of Industrial and Engineering Chemistry at the 127th Meeting of the .!!hrERIC.4U CHEMICAL SOCIETY, Cincinnati, Ohio, 1955.
Vapor-Liquid Equilibria for the System Carbon Tetrachloride-Furfural ROBERT E. WINGARD, WILLIAM S. DURANT, HAROLD E. TUBBS’, AND WILLIAM 0. BROWN2 Chemical Engineering D e p a r t m e n t , Alabama Polytechnic I n s t i t u t e , A u b u r n , Ala.
T
HE widespread use of furfural in the chemical industry led to a n investigation b y Long and Kuerner (3)into the feasibility of extracting furfural from aqueous solutions using carbon tetrachloride as the treating agent, and recovering the furfural from the extract b y distilling off the carbon tetrachloride. The deaign of distillation equipment to recover the solvent requires a 1 1
Present address, Hercules Powder C o . , Bessemer, Ala. Present address, Humble Oil Co., Houston, Tex.
knowledge of the vapor-liquid equilibrium compositions. T h e purpose of this investigation was to obtain the necessary data for the construction of the vapor-liquid equilibrium diagram. MATERIALS
Technical grade carbon tetrachloride was dried and purified by distillation. The purified material was found to have a refractive index of 1.4603 =k 0.0002, in good agreement with published values ( 1 ) . Technical grade furfural mas purchased from t h e
