I N D U S T R I A L A N D E N G I N E E R I N G CHEMISTRY
1968
change with an increase of molecular weights whereas the density definitely approaches an asymptotic value. Therefore, substituting log y z s o = 1.00 0.0123 in the two Smith equations, thus eliminating t,he viscosity terms, and keeping in mind that, the viscosit,y relation holds reasonably well only for liquids of molecular. weights greater than 2500, it was possible t o mathematically differentiate the two equations (keeping the specific heat and density constant) and find a relation between the thermal conductivity and molecular weight,. Proceeding on this basis, Smith's 1930 equation indicated a maximum thermal conduct,ivityat a molecular weight of approximately 12,900 or a viscosit,y around 220 cs. a t 25" C. This region for a maximum checks reasonably well with experimental results, although thermal conductivities calculated by this equation for dimethylsiloxane polymers would be of little value. Smith's 1936 equat'ion, when tested mathematically as outlined above, indicated a minimum thermal conductivity a t a molecular weight of approximately 6800 or a viscosity around 100 cs. at 25" C. This does not in any way conform to the experimental results but was calculated simply for general interest.
+
den
ACKNOWLEDGMEKT
Acknowledgment is made t o the Dow Corning Corporation for their general interest and assistance; the research was performed under their sponsorship.
Vol. 41, No. 9
LITERATURE CITED
Barry, A4.J., J . Applied Phys., 17, KO.12, 1020-4 (1946). (2) Barry, A. J., Dow Corning Corp., Midland, hfich., private conimunication (August 1948). (3) Bass, S. L . , I b i d . (April 1945). (4) Bass, S. L., Hyde, J. F., Britton, E. C., and McGregor, R. R . , M o d e r n Plastics, 22, 124 (November 1944), (5) Bates, 0. K., IND.ESG.CHEX.,25, 431 (1933). (6) Ibid., 28, 494 (1936). (7) Bates, 0. K., Hazzard, G . , and Palmer, G., IND. ENG. CHEM., ANAL.ED.,10, 314 (1938). (8) Chem. Eng. X e w s , 22, 1134 (1944). (9) Chem. C Met. Eng., 51, 66. 138 (1944). (10) D o w Corning Corp., Midland, Mich., tech. pamphlet, "Dow Corning Fluids." (11) Ibid., "Dow Corning Silicones." (12) Hunter, hl. J., Dow Corning Gorp., Midland, Mich., private communication (November 1945). (13) Hunter, M . J., Warrick, E. L., Hyde, J. F., and Currie, C. C., J . Am. Chem. SOC.,68, 2284 (1946). (14) Lipka, "Graphical and Mechanical Computations," pp. 120-70, New York, John Wiley & Sons, 1915. (15) Rochow, E. G., "Chemistry of the Silicones," pp. 64-70, NewYork, John Wiley &- Sons, 1946. (16) Rochow, E. G . , Sci. American, 179, No. 4, 50 (1948). (17) Smith, J. F. D . , IND.ENG.CHEM.,22, 1246 (1930). (18) Smith, 6. F. D., Trans. Am. SOC.Mech. E n g r s . , 58, 719-25 (November 1936). (19) Weber, H. F., Wied.A?~al.,10, 103,304,472 (1880). (1)
RECLIVEU Deceiiiber 1, 1048.
e
Thermodynamics oz Carbon Disulfide Production I,
J
D. R. STULL The Dour Chemical Company, Midland, Mich. Heat capacity, entropy, and free energy data have been calculated for the equilibrium sulfur vapor S * (g) containing 1 gram atomic weight of sulfur at 1 atmosphere total pressure distributed among the forms SSi -Sa e Sz. This data has been coupled with the appropriate thermodynamic data on graphite, carbon disulfide g a s , methane gas, and hydrogen sulfide gas to calculate the equilibria for the reactions:
C (graphite) CH, (g)
+ 2S* (g)
+ 4s' (g)
7 - L 2HzS
C S Z(g) (9)
+ CSz (g)
The calculated data is in fair agreement with the measurements of Koref (19), but predicts a higher yield of carbon disulfide than the measurements of Holtz (12) indicate.
I
K 1796, Lampadius heated pyrites with charcoal and acci-
dentally produced a new compound which is known today as carbon disulfide. This material has many unusual properties which have caused it to develop into an important article of present day commerce. Recent production figures by Hibben (11) show- the quantity of carbon disulfide produced in the United States has increased 45% within the last 5 years and now stands a t ioughly 195,000 tons per year. The rapid expansion in this item shows it t o be the subject of much technological interest. Gaps in the fundamental data of a material of this importance indicate one of the first places the research worker should spend his time. I n spite of the voluminous literature on carbon disulfide ( 1 4 ) there is yet much t o learn. This discussion
will confine itself to the thermodynamic aspects of the equilibria encountered in the commercial production of carbon disulfide from carbon or methane with elemental sulfur. Owing to the experimental difficulties involved, the corrosive action of sulfur, and the general complexity of the problem, there has been no direct measurement of the heat capacity of sulfur above its boiling point. The only available information is t h e spectroscopic calculations of the heat capacity of diatomic sulfur gas SZby Godnev and Sverdlin (8). These calculations have been checked by Kelley (16) who has done a splendid piece of work iiL collecting and correlating the thermodynamic data on sulfur a n d related compounds. Calculations based on this diatomic sulfur gas information permit a n approach to the thermodynamic problems involving sulfur, b u t give a very unrealistic view of these problems a t temperatures below about 650 O C. because of the considerable quantity of heat involved in the dissociation of the equilibriuni gas above the boiling point. At the boiling point, the equilibrium gas consists mostly of the species SSand Sa,which subsequently dissociate as the temperature is raised into S2. This dissociation is substantially complete to the diatomic form a t 850" C. Further dissociation t o monatomic bulfur requires temperatures above 1500" C. Consequently an attempt has been made to, calculate the heat capacity of equilibrium sulfur vapor in order to come closer to the facts in the temperature range from t h e normal boiling point (445" C.) to 650" or 700" C. Throughout this discussion, the symbol S* (g) will be used to represent t h e equilibrium sulfur vapor containing 1 gram atomic weight of sulfur a t 1 atmosphere total pressure distributed among t h e forms SS Sg F= SO.
==
INDUSTRIAL AND ENGINEERING CHEMISTRY
September 1949
1969
DISSOCIATION OF THE POLYATOMIC FORMS OF SULFUR
The vapor density of sulfur has been t h e subject of a number ( I ) , Riecke (39), Bleier and Kohn ( S ) , of investigations-Biltz
c
Biltz and Preuner (,?.?)-which were climaxed by the extensive measurements of Preuner and Schupp (35). The data of Preuner and Schupp from 300" to 850 O C. has more recently been called into question by Klemm and Kilian (18) who have repeated the work from 450" to 850" C. with more modern facilities, and have obtained average molecular weights somewhat higher (as much as 10% in some places) than Preuner and Schupp, although Klemm and Kilian regard their work as unfinished and give only a preliminary report. (Klemm and Kilian worked at the Institute for Inorganic Chemistry a t the Technische Hochschule in Daneig-Langfuhr, and were subsidized by the Leverkusen Works of the I. G. Farbenindustrie to repeat this work on the vapor density of sulfur vapor. They were forced t o give up their pioject early in World War 11.) Preuner and Schupp found t h a t the average molecular weight of sulfur vapor in the region of the normal boiling point is between S6 and Sa,and that the observed vapor densities can be exSS SP. plained by a stepwise dissociation of the type Sa Their work seems t o rule out definitely the presence of S*, since t h e resulting equilibria were not consistent with their findings. This is in agreement with the data of Stafford and von Wartenberg (34) who studied the thermal conductivity of sulfur vapor and failed t o find any evidence for the existence of a n S,form. Presumably a four-membered ring of sulfur atoms is not sufficiently stable a t these temperaturefi.
* =
By assuming t h a t a t 300 and 350 O C. the concentration of SZis
.so low that i t can be neglected (which the later calculations bear
out) the only equilibrium to be considered is: 3ss 4Sa a n d the equilibrium constant K l for this is:
K~ =
(s6)4i(s8)3
(1) (2)
This system is conditioned by the two equations: PS f PS = pt
ptu
=
mRT
-
M.32
(3)
(4)
where pa is the partial pressure of octatomic sulfur as, p~ is the partial pressure of hexatomic sulfur gas, pt Is the t o t 8 pressure on t h e system of volume 0 and of mass m having on the average M atoms per molecule, R is the gas constant and T is the absolute temperature. Insertion of the experimental data into Equations 3 and 4 and solving simultaneously, one obtains p a and 1)6 and can then evaluate K1 for the 300" and 350' C. isotherms. Table I gives the experimental data listed by Preuner and Schupp. The variation in K1leaves more to be desired, but i t is the only data in this region. Substitution of these values of K I into the van't Hoff reaction isochor leads to the value 29,300 calories for the heat of dissociation for the reaction in Equation 1. By assuming that the heat of
TABLE I. EQUILIBRIUM CONSTANT FOR EQUILIBRIUM 3Sa = 4Sa [KI = ( p e ) 4 / ( p s ) S , calculated from data of Preuner and Schupp (86)] Pressure, Molecular pt (mm.) Wt., M Pfi PS K1 300' C. Isotherm 7 5 7.03 3.64 3.86 3.05 10.5 6.98 5.35 5.15 6.00 13.5 7.09 6.15 7.35 3.60 28.5 7.12 12.54 15.96 6.08 7.20 49.0 15.60 32.40 1.74 4.09 av. 350' C. Isotherm 32.0 6.89 16.32 15.68 18.40 55.0 6.97 28.33 26.67 33.95 81.0 7.03 39.69 41.31 35.20 121.0 7.03 58.68 62.32 48.57 34.03 av.
TABLE 11. CALCULATED EQUILIBRIUM CONSTANT K1 FOR REACTION 3Ss = 4Sa T , C.
Ki = ( p e ) ' / ( p d a 182 769 2,680 8,020 21,000 49,600 107,000 213 000 397:OOO 698,000
400 450 500 550 6 00 650 700 750 800 850
this dissociation fis constant, the ratio of Sa to SSis fixed for all temperatures to be considered, and has been used to calculate the Kl values found in Table 11. At tem eratures from 350 O to 850 O C., two other equilibria accompany Equation 1-namely,
ss -_L
452
(5)
S6 e382 with their accompanying equilibrium constants :
(6)
KZ = (Sd4/(Sd
(7)
KB
(8)
(S2)'/(Se.) The system:is now:conditioned by three equations:
Pa
+ Ps + Pz = Pl + 2pz = Mpt
(9)
8 p s +"bps
(10)
K~ = (s6)4/(s8)3 (2) where p2 represents the artial pressure of diatomic sulfur. A simultaneous solution of Equations 2, 9, and 10 gives the partial pressures of all the constituents over the temperature range of interest, which are listed in Table 111. Figure 1showing the partial pressures of ss, SO,and sBpresent in the equilibrium sulfur vapor at a total pressure of 1 atmosphere from 400' to 850' C. has been obtained by the methods outlined above from the experimental data of Preuner and Schupp (86)and of Klemm and Kilian (18) (Table IV). The calculations presented here are in substantial agreement with similar calculations made by Preuner and Schupp ( 8 5 ) and more recently by Kelley (16). Klemm and Kilian (18) presented only basic data with no further calculations. In this present
TABLE 111. PARTIAL PRESSURES OF VARIOUS CONSTITUENTS OF THE EQUILIBRIUM SS e SS e 52 (Partial pressure in mm. of Hg; one atmosphere total pressure) c-Preuner and SchuppY-Klemrn and KilianT , C. PI Pa PP P8 P6 PP 400 366 37 1 23 327 450 i ii 28 28 420 405 500 237 61 87 435 327 37 1 550 149 189 402 207 209 364 600 367 292 398 288 70 104 650 589 22 535 149 183 41 700 6a 667 686 10 68 83 750 4 la 709 729 30 47 800 748 2 12 30 Oa 728 850 0 752 734 25 8 1 0 Interpolated, paralleling Klemm and Kilian's data in this region.
.*.
~~
~~~
-
~~
~
~
TABLE Iv. WEIGHT FRACTIONS I N GRAMS OF THE LrAR1OUS CONSTITUENTS O F THE EQUILIBRIUM SS S6 S S? T, K. 720 750 800 850 900 950 1000 1050 1100 1150 1200
(One atmosphere pressure) --Preuner and Schupp--Klemm W8
16.62 14.71 11.11 7.56 3.94 1.42 0.44 0.14 0.00 0.00 0.00
and Kilian-
WE
WP
W8
We
WZ
15.03 16.66 18.83 19.06 15.57 9.54 4.85 2.15 1.22 0.64 0.30
0.34 0.68 2.11 5.57 12.53 21.09 26.76 29.76 30.84 31.41 31.76
20.56 18.74 14.76 10.22 5.96 2.64 0.85 0.32 0.14 0.00 0.00
11.15 12.83 15.65 17.03 15.60 10.83 6.52 4.13 3.17 2.74 2.51
0.33 0.49 1.63 4.81 10.49 18.58 24.69 27.61 28.74 29.32 29.54
I N D U S T R I A L A N D E N G I N E E R I N G CHEMISTRY
1970
Vol. 41, No. 9
700
2 600 3
e v
:
500
U
400
f
I u .
s
300
#e 900 n
IO0 0
400
450
Figure 1.
500 550 600 TEMPERATURE,
650
700
750
800
850
900
' CENTIGRADE
TEMPEQATURE,
Calculated Partial Pressures for the System
s8 e s6
' KELVIN
Figure 2
SZ
pressure, sulfur is more dissociated than Klemm and Kilian find it to be. HEAT CONTENT O F SULFUR VAPOR A T ONE ATMOSPHERE TOTAL PRESSURE
I
I
iI
if II/
From the consideration of a number of reactions, Kelley ( 1 6 ) has arrived a t a value of 15,510 calories for the conversion of 1 gram atomic weight of rhombic sulfur a t 298.16" K. to 1 gram atomic weight of diatomic sulfur gas also a t 298.16" I(. Kow Godnev and Sverdlin (8) and Kelley ( 1 6 ) have calculated the heat capacit,y of Sz (9) from spectroscopic data. The linear equation given by Kelley is used here:
i
I
I
S2 (g): C,
300
500
700 900 1100 TEMPERATURE, K E L V I N
1300
1500
Figure 3
paper, however, data related to a constant inass of sulfur are desired, and so the partial pressure of each constituent was multiplied by its molecular weight giving a set of weight fractions t o g , m,, and wg,which were then converted t o fractions of the grain atomic weight of sulfur (32.06 grams). These are shown in Figure 2 where the weight fractions of 1 gram atomic weight of sulfur are plot,ted from 400' to 850 "C., and have been read baclrat 10" K. intervals. The principal difference between the results of Preuner and Schupp and of Kleinm and Kilian is easily observed in Figure 2. The data of Iilemm and Kilian indicate t h a t a given degree of dissociation is reached a t a higher temperature than shown by the data of Preuner and Schupp. P u t another way Preuner and Schupp find that a t a given temperature and
TABLEV. HEATCONTENT AND EXTROPY CALCCLATIONS FOR SOLIDAND LIQCIDSULFUR
JG Reference temperature 298.16' K. Integrate rhombic f o r 4 298 16O to 368 6O K. Transition rhombic to Aonoclinic, 368.6O K. Integrate monoclinic form, 368.6' t o 392' K. Transition monoclinic t o liquid, 3Q2O K. Integrate equilibrium liquid form, 392O to 717.8" K. Transition t o equilibrium gas form, 717.8O K. Equilibrium gas form a t 717.S0 K.
- H&.M, Cal. cr So, Gram f t o r n E.G. 0.0 7.624 395.2 1.218 86.0 0.233 145.1 0.374 295.0 0.752 2790,O 6.178 2228.0 5939.3 18.482
3.103
=
7.76
+ 0.000888 T (calories/gram formula u-cight)
(11) Taking rhombic sulfur a t 298.16" K. as the reference aero, the heat content, of diatomic sulfur gas can be computed with considerable confidence (Figure 3). Eastman and hlcGavock ( 7 ) have measured the heat capacity of rhombic and monoclinic sulfur from liquid hydrogen temperatures to about 375' K., and Mondain-l\ionval ( 2 4 )has given the heat of transition of rhombic to monoclinic sulfur a t 368.6 ' K. as 68 calories per gram atom and also 295 calories per gram atom for the heat of melting monoclinic sulfur t,o the liquid form (S+) a t 392" K. Kelley (15)has kindly furnished the heat content for the equilibrium liquid ( S i p ) from the melting point t o the boiling point. West and Menzies (39) give 2228 ca,lories per gram atom for the heat, of vaporization at 717.8" K., t,he normal boiling point. Integration of this heat, shown in Table V, brings the author's figure for heat content of liquid sulfur, a t the normal boiling point, t o 3711 calories per gram atom plus 2228 calories for vaporization, or 5939 calories per gram atom for the heat content in the gas phase a t the normal boiling point. This gives t,he starting point and the upper boundary into which the calculated heat content of equilibrium sulfur gas S"( g ) must fit. Kelley (16) has estimated the difference between thc heat capacity of Ss (9) and 481 (g) to be AC, = 6; and for the difference in the heat capacity of S g (g) and 3S9 (9) as AC, = 4. This leads to the equations:
+ 0.003552 T (calories/gram formula n-eight) (12) S6 (g) : C , = 19.25 + 0.002664 T (calories/gram formula weight)
Ss (g) : C, = 25.00
(13)
Multiplication of W E , Wg, and wz by the appropriate specific heat from Equations 11, 12, and 13 and totaling leads to the gaseous heat capacity of the equilibriuin sulfur gas S E shown in Table VI and Figure 4. From a careful analysis of Preuner and Schupp's data, Kelley (16) has given the heats of the following reactions:
10
1971
INDUSTRIAL AND ENGINEERING CHEMISTRY
September 1949
,
TABLE VI. HEAT CAPACITIES IN EQUILIBRIUM SULFURVAPOR (One atmosphere pressure) Calories per Gram Atom Calories per Gram Atom T , "!K. Sia S E ~ S*(g)c T,' K. 62 SB s*(g) 4.38 4.35 6.50 720 7.60 4.19 1150 4.41 4.38 5.17 750 11.05 4.21 1200 4.43 4.41 4.85 22.10 4.23 800 1250 4.45 4.44 4.57 4.25 1300 850 40.15 4.47 4.47 4.47 900 61.65 4.27 1350 4.49 4.49 4.49 950 61.20 4.29 1400 4.52 4.52 4.52 34.25 4.32 1450 1000 4.54 4.54 4.34 4.54 1.500 17.50 1050 4.36 1100 9.85 a S z = Diatomio sulfur gas caloulated from spectroscopic data. b SE e Equilibrium gas (Sa SB Sz) qnly. C S* ( g ) = Equilibrium gas plus heat of dissociation.
e
phere total pressure, instead of the hypothetical gas S,,which is only obtained completely a t temperatures above say 1200" K. 700
I
I
eo0
BOO
I
I 1100
1000
TEMPERATURE
I
I
I
1*00
1000
1.00
I FORMATION OF CARBON DISULFIDE FROM ITS ELEMENTS
,100
*KELVIN
Brown and Manov ( 4 )have measured the low temperature thermal data for carbon disulfide from liquid hydrogen temperatures to the boiling point. Their measurements lead to a value of 57.1 * 0.5 e.u. for the entropy of carbon disulfide gas a t 298.16" K. Cross (6) has calculated the entropy, heat capacity, and free energy for carbon disulfide from spectroscopic information. His value of 56.84 e.u. at 298.16' E(. for the ideal gas state is well within the experimental errors of the calorimetric values of Brown and Manov (4). Coupling this data on carbon disulfide by Cross (6) and the thermodynamic data on graphite published by Rossini and co-workers (50) with the data on equilibrium sulfur gas S* (g) presented in this discussion, equilibrium data for the reaction
Figure 4
, 500
600
700
800 900 1000 1100 1200 1300 1400 1500 TEMPERATURE, a K E L V I N
Figure 5
Sg (g) S6(9)
= 452 (g) A€€ = 95,200
= 3Sz (9) AH = 64,090
-b 67' (calories/Ss mole) + 477 (calories/~amole)
(14)
(15) Average values of these heats of dissociation from 700 €0 1000' K. are : 452 (9) 391.0 calories per gram of s8 dissociated (16) Sg (9) = 352 (g) 350.8 calories per gram of s6 dissociated (17) Over each 10" K. interval, the gram atom fractions WE,W6,and zu~change slowly. The increment in weight for a 10" interval multiplied by the proper heat of dissociation (15, 16) gives the contribution to the heat content. Addition of the contribution of the gaseous heat capacity S E to the heat content resulting from the dissociation gives the apparent heat content of equilibrium sulfur gas S* (g). Figure 3 shows the heat content on the basis of the data of Preuner and Schupp, and of Klemm and Kilian. The average of these two sources is shown in Table VII, and has been calculated to the apparent heat capacity of equilibrium sulfur gas S* (g) and is shown in Table VI and in Figure 4.
SS (g)
f
*
C (graphite)
+ 2S* (g)
CS2
(g)
(18)
has been calculated and is given in Table VI11 and Figure 5. Figure 6 gives the heat content picture for this system. Cross (6) has also calculated this same equilibrium using S, (g) as the basis for his calculations. On the basis of his data the free energy is negative a t all temperatures, which fact is contrary t o experience. As the equilibrium SS 86 SZ
=
THERMODYNAMIC DATA FOR EQUILIBRIUM SULFUR GAS
By using the heat data given in the last section, and in Tables V and VII, one has thermodynamic data a t hand t o make equilibrium calculations with the actual form of sulfur a t 1 atmos-
AND ENTROPYTALCULATIONS FOR GASEOUS TABLE VII. HEATCONTENT EQUILIBRIUM SULFUR,S E ~PER GRAMATOM(32.06 GRAMS)
-
-
Hi88, Av. C i a , H g H:a 8, 50 SO, T, E.U. Cal. Cal. E.U. K. Cal. 16,375 5,939 41.00 18.48 990 717.8 36.50 16,740 18.50 1000 720 7.50 5,955 17,060 6,035 18.61 1010 32 00 730 7.95 740 8.50 6,120 18.73 1020 28.00 17,340 6,223 24.70 17,587 18.87 1030 750 10.30 760 6,341 19.02 io40 17,803 11.80 21.60 770 13.60 6,477 19.20 1050 18.60 17 999 18:163 19.40 1060 780 15.60 6,633 16.40 18,306 790 18.00 6,813 19.63 1070 14.30 800 7,019 12.50 18,341 20.60 19.89 1080 810 23.60 7,255 20.18 1090 11.20 18,543 820 18,646 26.90 7,524 20.51 1100 10.30 830 7,829 30.50 20.88 1110 9.40 18,740 840 34.00 8,169 21.29 1120 8.50 18,825 850 38.10 8,550 21.74 1130 7.70 18,902 860 42.20 8,972 22.24 1140 7.20 18,974 870 9,436 19,041 46 40 22.77 1150 6.70 880 23.35 1160 50.80 9,944 6.30 19,104 890 6.00 19,164 55.10 10,495 23.98 1170 900 5.70 19,221 59.30 11,088 24.64 1180 910 64.00 11,728 25.34 1190 19,276 5.50 920 67.00 12,398 26.08 1200 5.20 19,328 930 67.80 13 076 26.81 1250 5.00 19,578 940 19,813 67.60 13:752 27.53 1300 4.70 950 4.45 20,036 64.40 14,396 28.21 1350 960 20,259 58.00 14,976 28.82 1400 4.46 970 20,482 52.40 15,500 29 37 1450 4.46 980 46.50 15,965 29.84 1500 4.47 20,705 Average heat capacity from inidicated temperatui' e to the temperature immediately below.
T,
O
K.
Av. C i a ,
Cal.
I
H;
INDUSTRIAL AND ENGINEERING CHEMISTRY
1972
fide with its elements derived from these sources is in poor agreement wit,h the free energy changes computed here, and are not in agreement wit,h each other. C,ross (6) and t'he writer hold the opinion that the computed values are closer t o the truth. The equilibrium constant data for thc equilibrium ( 1 7 ) have been solved for concentration, and results in the mole fractions given in Table 1-111 and in the diagram given in Figure 7. The yield of carbon disulfide from this reaction reaches a maximum a t about 750" C. and decreases slowly as the temperature is raised above this
+
TABLE VIII.
EQUILIBRIUM DATAFOR C (GRAPHITE) 2S* (g) = [Equilibrium constant K = (CS%)/(S)*]
T,
K.
Cal./Gfw.
Cal./Gfw.
+25
298 300 400 500 600 700 800
+27,580
+15,594 15.522 +11,661 +7,929 +4,441 +1,09i -2,079
L/
127 227 327 427 527
627 727
900
1000 1100 1200 1300 1400 1500 88 SS
A27
927 1027 1127 1227 a * =
S*,
4F;,
T, C.
AH:,
+
+27,576
+26,517 $25,679 +24,920 +24,085 +17,569
-3.708
+10,791 1.378 -2,548 -3,025 -3,118 -3,130 -3,149
-4;674 -5,012 - 6,197 - 5,403 -5,594 -5,774
&/T
LogloK
62,301 51.740
28.902
15.858 7.403 1.567
-2.600
-11.4312 - 6,3171 - 3.4660 - 1.6181 - 0,3425
- 11 ,3087
0.5680 0,0005
-4,120 -4.674 -4.556 -4.330 -4.156 -3.995 -3.840
1.0216 0.9959 0,9465 0,9084 0.8733 0.8413
Atm.
:
1 000 1,000 0.999 0.976 0,688 0.213
0.112 0.087 0.091 0.102 0,110 0.118 0,126
Vol. 41, No. 9
:
0 000 0.000 0.001 0 024 0,312
0.787 0,888 0.913 0,909
0.898 0,890 0 882 0,874 I
pornt
SZ.
I
FORRZ4TION OF CARBON DISULFIDE FROM M E T H h l E .4hD SULFUR 50
D e Sinio ( S j , Rakovskil and Kanineva (as), and Thacker and Miller (38)have proposed the production of carbon disulfide from methane and suliui To test the feasibility of this ieaction its thermodynamic equilibrium has been computed. Free energy data on methane by Rossini and eo-workers (30) has been used with the data on hLdrogen sulfide b? Cross (6)together with the sulfur and carbon disulfide already mentioned to calculate the equilibrium for the reaction:
40
i 4
v 3 30 c '
CHd (g)
5
c
z
+ 48* (9) =--'2HzS (g) + CSz (g)
The iesults are given in Table S and Figuie 8.
8 PO c
4 r 10
0 300
5 00
700 900 1100 TEMPERATURE, K E L V I N
1300
I500
Figure 6
approaches the Sa state, the free energy values approach each other and finally become the same. Koref ( 1 9 ) passed sulfur vapor over Acheson graphite in a flowing system and arrived a t the equilibrium constants shown in Table IX and Figure 5 . They are in qualitative agreement with the predictions. Holtz (13) passed sulfur over charcoal obtained by charring sucrose. His equilibrium data is also given in Table IX and in Figure 5 His data is farther from the prediction than the data of Koref, and agrees better a t the lower temperatures. Other workers have measured equilibria of other reactions involving carbon disulfide [Kleinenc ( l 7 ) , Lewis and Lacey ( $ I ) , RassoLv and Hoffnianri ( 2 8 ) , Stock and Seelig ( 3 6 ) , Stock, Siecke, and Pohland (56)) and Terres and Wesemann ( S T ) ] . Randall (2'7) has discussed these results but no record of this discussion is to be found in the chemical literature. Data on the equilibrium of caiboii disul-
0
PO0
400
TEMPERATURE,
600
800
1000
19.00
CENTIGRADE
Figure 7 . Equilibrium Data for Reaction C (Graphite) 2s" ( g ) CSz ( g )
+
+
TABLE IX. EXPERIMENTAL VALCESOF EQL-ILIBRIWN c 2s*(g) = CS2
+
T, K.
R In K
1096 1179 1282 1383 1089
-4.05
1089 1089 1144 1200
-4.79 - 5 76 -6.40 -3.44
-3.36
-3.20 -1.03 -0.89
Reference Koref Koref Koref Koref Holtz Holta Holtz IIoltz Holtz
(19) (19) (19) (29) (13)
(I,$?)
(Id) (1%) (Id)
0
PO0 400 600 800 1000 TEMPERATURE, O C E N T I G R A D E
1200
Figure 8. Equilibrium Data for Reaction CHd ( g ) + 4S* ( g ) e 2HzS (g) CS2 (8)
+
(19)
September 1949
T,
T,
C.
OK. 298 3 00 400
+
2
127 227 327 427 527 627 727 827 927 1027 1127 1227 Q
(7) Eastman, E. D., and Mc-
+
EQUILIBRIUM DATA FOR CH4 (g) 45* (9) = 2H2S (g) -I[Equilibrium constant K = (HzS)z(CSz)/(CH3(S)41 AH; AF; HzS CSt, Cal./Gkw. Cnl./Gkw. AFl/T Loglo K Atm: Atrn. t40.223 - 8.7915 0.000 0.000 +11,993 +35,869 - 8,6342 0.000 0.000 +39.503 +11,851 +35,863 -2,1578 0.000 0.000 +9,872 33,845 +3,949 1.4548 0,646 0,323 -3,328 -6.656 +32,267 3.7510 0.333 - 10,297 - 17.161 0.667 +30,863 5.3215 0,333 - 24,347 - 17,043 0.667 +29,290 - 29,315 6,4073 0.333 -23,452 0.667 16,348 6,5014 0.333 -26,771 - 29.745 0.667 +2,884 0.333 6.2945 - 28.799 -28,799 0.667 - 17,869 0.333 -29,561 0.667 - 26.864 5,8717 - 23,644 0.333 - 25.006 5,4654 -30,007 -24,511 0.667 0.333 5.1484 - 23,555 0.667 -30,622 -24,733 - 22.194 0,333 -31,072 -24,702 4.8609 0.667 0,333 4.5918 0.667 -21.008 - 31,673 -24,673
TABLE X. O
+
500
600
700 800
+
900
1000 1100 1200 1300 1400 1500
1913
INDUSTRIAL AND ENGINEERING CHEMISTRY
* = S8 e SB
Gavock, W. C., J . Am. Chem. SOC.,5 9 , 145-51
CS2 k ) "
(1937).
S*
CHI Atm: 0.333 0.333 0.333 0.011
0.667
0.000
0.000
0.000 0.000 0.000 0.000 0,000 0.000 0.000 0.000 0.000
At&. 0.667
0.667
0.020 0.000 0.000
0.000
0.000 0.000
0.000 0.000 0.000
0.000
(8) Godnev, I. N., and Sverdlin, A., Z . Physik., 97, 124-30 (1935).
(9) Griawold, T., Jr., U. S.
Patent 1,918,033 (July 11, 1933). (10) Harkness, A. M., Ibid., 2,046,818 (July 7 , 1936).
(11) Hibben, J. H.. IND.ENG. CHEM., 40, 979-91 (1948). (12) Holtz, J. C., dissertation,
Johns Hopkins Universitv, Baltimore, Md.
Sl.
(1930). (13) Iddings, C . , U. S. Patent 2,079,017 (May 4 , 1937).
__---.-.-
100
Kausch, O., "Der Schwefelkohlenstoff," Berlin, Julius Springer, 1929.
HEAT CONTEN7 RELdTlONSHlPS FOR THE REACTION
1
Kelley, K. K., private communication (Dec. 23, 1946). Kelley, K. K., U.8. BUT.Mines, Bull. 406 (1937). Klemenc, A . , Z . anorg. u. allgem. Chem., 191,246-82 (1930). Klemm, W., and Kilian, H., 2. physik. Chem., 49B, 279-83
I
(1941).
Koref, F., Z . anorg. Chem., 66, 73-92 (1910). Legler, E., U. S.Patent 1,793,181 (Feb. 17, 1931); Ibid., 2,200,4 7 5 (May 14, 1940). Lewis, G . N., and Lacey, W. N., J . A m . Chem. SOC.,37, 1976-83 ( 19 15).
McElroy, K. P., U. S. Patent 1,369,825 (March 1, 1921). Merriam, H. F., Ibid., 2,141,758 (Dec. 27, 1938). Mondain-Monvsl, P., BUZZ.soc. chim., 39, 1349-68 (1926). Preuner, G., and Schupp, W., 2. physik. Chem., 68, 129-56 (1909).
Rakovskii, E. V., and Kamneva, A. I., J . Applied Chem.
(U.S.S.R.), 13, 1436-41
(1940).
Randall, M., presented before the Division of Phykcal and Inorganic Chemistry a t the 90th Meeting of the AN. CBEM.SOC., San Francisco, Calif., August 1935. Rassow, B., and Hoffmann, K., J . prakt. Chem., 104, 207-40 300
500
100
900
TEMPERATURE,
O
1100
1300
Figure 9
b
-
The data indicate that theoretical yields are t o be expected at temperatures as low as 300' or 400' C., but Thacker and Miller (38)report their best conyersions in the vicinity of 700"C, Two factors are possibly the reason for this. First, this higher temperature is required because of the low reaction rate obtained a t the lower temperatures indicated by the thermodynamic data. Secondly, the true reactant in both of these reactions (17, 18) may really be SZ (g). Thus at the lower temperatures its concentration will be dependent upon its rate of formation from the equilibria 1, 5, and 6. Industrial experience with these s y s t e m shows t h a t these equilibria which form SZ (g) are comparatively slow a t temperatures below 660' C. Figure 9 shows the heat content for the foregoing reaction. The information in both Figures 6 and 9 indicate graphically that one of the cardinal problems of carbon disulfide production is t h a t of supplying the requisite heat at the proper temperature level to enable the reaction t o proceed. T h a t this fact is true is amply demonstrated by the patent literature (9, 10,13, 90,22, 23, Si, 3 2 ) which includes methods of superheating sulfur vapor and schemes for introducing heat into the reacting masses. LITERATURE CITED (1) (2) (3) (4)
(1922).
1500
Riecke, E., 2.physik. Chem., 6, 430-6 (1890). Rossini, F. D., et al., Natl. Bur. Standards (U.S.), Circ. C461
KELVIN
Biltz, H., 2. physik. Chem., 2, 920-47 (1888). Biltz, H., and Preunei, G., Ibid., 39, 323-41 (1901). Bleier, O., and Kohn, L., Monatsh., 21, 575-620 (1900). Brown, 0 . L. I., and Manov, G. G., J. Am. Chem. Soe., 59, 500-2
(1937). (5) Crdss, P: C., J . Chem. Phys., 3, 168-9 (1935). (6) Ibid., 3, 825-7 (1935).
(1947). (31) Saladin, O., U. 9. Patent 2,258,367 (Oct. 7, 1941). (32) Silsby, C. S., Ibid., 2,141,766 (Dec. 27, 1938); 2,141,768 (Dec. 27, 1938). (33) Simo, M. de, U. S. Patent 2,187,393 (Jan. 16, 1940). (34) Stafford, 0.J., and Wartenberg, H. von, Z . physik. Chem., 77, 66-72 (1911). (35) Stock, A., Siecke, W., and Pohland, E., Ber., 57B, 719-35 (1924). (36) Stock, A., and Seelig, P., Ibid., 52, 681-94 (1919). (37) Terres, E., and Wesemann, H., Z. angew. Chem., 45, 795-802 (1932). (38) Thacker, C. M., and Miller, E., IND. ENG.CHEM.,36, 182-4 (1944). (39) West, W. A., and Menzies, A. W. C., J . Phys. Chem., 33, 1880-92 (1929). RECEIVED July 31, 1948. Presented before the meet& of the Division of Physical and Inorganic Chemistry of the AMERICAN CHEMICAL SOCIETY at Syracuse, N. Y . , June 28 to 30, 1948.
* * * i
Last year [IND. ENG.CHEW,40, 87A (November 1948) and 69A (December 1948)] Fontana devoted two of his monthly columns t o a description of the particular corrosion problems encountered in contact sulfuric acid plants. Corrosion causes and corrosion-resistant materials of construction are discussed, starting with the piping used in the pits where the sulfur is melted and continuing on through the process covering blowers, heat exchangers, coolers, absorbers, valves and pumps, and storage and shipping facilities.
