k = equilibrium adsorption constant, g of pure solvent/g of inert solid K = ( c p ~ / p ~ )k, concentration relation expressed on a liquid and solid base, g of pure solvent/g of inert solid L = amount of solvent in the liquid phase, g of pure solvent P = SX/L fraction of solvent retained by the solid when passing from one extractor to another r = radial coordinate, cm ro = external radius, cm R = r/ro = dimensionless radial coordinate S = amount of inert solid in the extractor, g of inert solid t = time, s w = concentration of solute in the liquid phase, g of solute/g of pure solvent W = dimensionless concentration in the liquid phase defined by eq 7
+
Greek Letters ) = function defined by eq 13 /3( ) = function defined by eq 14 t = porosity of the solid 8 = dimensionless initial concentration of the solution X = mass of pure solvent retained per gram of inert solid, g of pure solvent/g of inert solid p~ = apparent density of the solvent, g of pure solvent/cm3 of solution p n = roots of transcendental eq 10 ps = apparent density of the solid, g of inert solid/cm3 of total solid
a(
$(
) = initial concentration profile in the solid, expressed as a function of the dimensionless coordinate R , g of solute/g inert solid
Subscripts f = refers to t equal to the time of residence in the extractor being considered i = refers to the i t h extractor I = refers to t = 0 in a given extractor L = refers to the liquid phase n = dummy index N = refers to the last active extractor of the cascade s = refers to the solid Literature Cited Kitaev, B. I.,Yaroshenko, Yu. G., Suchkov, U. D.. "Heat Transfer in Shaft Furnaces, p 31. Pergamon, London, 1967. Krasuk, J. H., Lombardi, J. L., Ostrovsky, C. D., lnd. Eng. Chem., Process Des. Dev., 6, 187 (1967). Munro, W. D.. Amundson, N. R., lnd. Eng. Chem., 42, 1481 (1950). Oplatka, G., 2.Zuckerind., 79, 471 (1954). Plachco, F. P., Krasuk, J. H., lnd. Eng. Chem., ProcessDes. Dev., 9, 419 (1970). Plachco, F. P., Lago, M.E.. Can. J. Chem. Eng., 50, 611 (1972). Rickles. R. N., Chem. Eng., 15, 157 (March 1965). Schneider. F., "Technolcgie des Zuckers", p 173, M. und H. Schaper, Hannover, 1968. Yang, H. H.. Brier, J. C., AlChEJ.. 4, 453(1958).
Received for review July 17, 1972 Resubmitted January 19,1976 Accepted March 6,1976
Ther maI Decomposition of Manganese Sulfate John P. McWilliams and A. Norman Hixson" Department of Chemical and Biochemical Engineering. University of Pennsylvania, Philadelphia, Pennsylvania 19 174
When carbon is added to MnS04 in a mole ratio between 0.5 and 1.O, the thermal decomposition of MnS04 is substantially promoted, yielding good rates (-60% conversion in 1 h) in the range 1250-1300 O F . The products of decomposition consist of almost pure manganous oxide, MnO, and a gas phase containing average values of 67 YO SO2 and 33 YOCOP.Increasing the carbon ratio above 0.5 (up to 1.O) gives an increased decomposition reaction rate but does not alter the ratio of SO2 to COP in the product gases. The overall decomposition kinetics may be described by a multiple step model which includes a sulfate pre-reduction step.
Objective The higher oxides of manganese (MnO:! and MnzO,) react readily with sulfur dioxide in aqueous solution or in a gaseous state, forming MnS04. Such diverse processes as manganese ore beneficiation (Allen, 1954; Ravitz et al., 1946; Vedensky, 1946; Wilhite and Hollis, 1968) and removing SO:! from stack gases (Bienstock et al., 1961; Ludwig, 1968) take advantage of this property. Economical recovery of the SO:! from the MnS04 would be advantageous either for recycle purposes or for further reduction to produce sulfur. MnS04 can be decomposed thermally, but on a commercial scale, the problems associated with the necessary high temperatures 2000-2300 O F are manifold. An extensive investigation of manganese ore beneficiation was carried out during World War I1 a t the Three Kids Mine in Nevada (Vedensky, 1946).Ring formation and agglomeration in the rotary kilns used for the MnS04 decomposition forced frequent shutdowns and poor process efficiency.
Pechkovsky (1955,1956,1957,1959)and others (Suchkov e t al., 1959) have reported work on an MnS04 decomposition. Apparently, manganese ore is beneficiated in Russia using an SO:! process that includes a thermal decomposition of MnS04 in horizontal kilns a t 2000-2300 O F . Pechkovsky investigated the use of carbon to reduce the temperature and found beneficial effects a t 1475 O F using up to a maximum C/Mn molal ratio of 1/1. Inasmuch as the published data were incomplete and shed no light on the reduction mechanism, this work was undertaken to study in detail the carbon-aided reduction of MnS04. I t was a further purpose to determine the maximum SO:! content obtainable in the exit gases which could be furnished to another process to produce sulfur as an integral part of an SO2 recovery system for stack gases.
Chemical Equilibrium for the System The theoretical chemical equilibria for the manganese sulfate-carbon system were developed in order to guide the Ind. Eng. Chem., Process Des. Dev.. Vol. 15, No. 3, 1976
365
Table I. MnS04-xC Decomposition Products Equilibria at 1000 K Gas phase composition” x (CIMn)
Pressure, atm
0.20
1.0
0.33
1.0
0.34
1.0
0.34
1.34
0.49
1.0
0.49
1.49
0.51
1.0
1.0
1.0
Solid products
% Mn
40.0 60.0 0.99 99.0 96.0 4.0 96.0 4.0 6.0 94.0 6.0 94.0 99.35 0.65 66.7 33.3
COS, so3, and 0 2 negligible.
6
-24
I
I \I \
i -
-
0
-32
%
-6
L OO P ‘2
’.
-10
-16
’\,
I
-8
0
I
I
liquid ~ ~ l l repion u r
4 -
n* ‘ 0
nu
@
-
I
0
d
14
-
19
-
coz
SO2
co
0.2500
0.7500
3.8 x 10-7
0.2500
0.7500
3.8 x 10-7
0.2537
0.7463
1.7 X
0.2537
0.7463
2.4 X
0.3289
0.6711
2.2 x 10-6
0.3289
0.6711
3.1 X
0.3390
0.6608
2.1 x 10-4
0.5994
0.4001
4.4 x 10-4
(1) and (2) and that the pair MnS04-MnO can coexist between (2) and (3). Alternatively stated: MnS04-Mn203 pair stable a t Pso2< 8.22 X lop3atm; MnS04-Mn304 pair stable a t 8.22 X atm < Pso2 < 2.58 atm; MnS04-MnO pair stable a t 2.58 atm < Pso2 2
which indicates that the reaction velocity is proportional to the perpendicular surface area a t the reaction front. Barret (1961) states that in the “chemical regime” each particle of the entire sample is reacting a t the same rate thus involving the total particulate surface area. If the particles are of uniformly sized isometric shapes then the 2/3 order is followed. However, he, as well as others, conducted experiments in the “diffusional regime” which gave reaction orders equal or close to the 2/3 value. This he shows is a function of the where w o = total wt of gas product theoretically removable sample bed geometry resulting from the rate of change of the and w = actual wt of undecomposed gas a t time t , this intearea of the thin reacting slice as a function of time. For exgrates to ample, a sample decomposing in a cylinder of large height t o diameter ratio open only a t one end followed a zero order, ht (uI/wl))1/3= - 3 + to while another in an inverted conical container with a large open area a t the top closely approximated the 2/3 value. Decomposition studies of ferric sulfate and aluminum Thus, a plot of ( u ; / w ~ ) ’ vs. / ~ time should yield a straight line sulfate were done by Warner and Ingraham (1962) using of slope -k/3. powders that had been compacted into regular shapes. Further Normally, only one species is decomposing to give one work using the same technique has been done by Ingraham gaseous and one solid product so that w/wois unambiguous. and Marier (1963, 1964, 1965, 1966, 1967). The method inIn the present case where the decomposition is aided by a volves the determination of the area of the changing reaction second species, one must define the unreacted fraction with interface as the decomposition proceeds. I t is based on the respect to each reactant. Hereafter, W / W O without subscript is the undecomposed weight fraction of SO2 only; ( w l w ~ ) c o , observation that with a dense pellet formed by compacting a powder, the rate of migration of the interface from the exis the undecomposed fraction of CO2. ternal surface of the pellet is uniform. The change in area can The effect of diffusion in decomposition processes has been be calculated from the geometry of the pellet and thus a rate examined by Barret (1961),who was able to show, through use per unit area can be determined. The rate measurements were of a Thiele-type modulus, t (Thiele, 1939),that under chemical made using a thermogravimetric balance. control the total reaction rate is proportional to the total particulate surface, but that under diffusion control, the rate Experimental Equipment and Sample Preparation is proportional to the surface area of a thin slice a t the reaction front. The reactor was a 3/4 in. i.d. stainless steel tube placed vertically in cylindrical heater resistance elements controlled by k7p c =La rapid-response electronic controller. The decomposition D experiments were batchwise. The sample zone was about 3 in. where L = thickness of bed, cm; D = effective diffusion coefdeep with thermocouples probing each end. T h e gases proficient, cm2/sec;h = decomposition rate constant, cm3 of gads duced by the decomposition were routed continuously through cm2 of surface; T = specific surface of solid, cm2/g; and p = bed a sample loop of a modified Beckmann GC-5 gas chromatodensity, g/cm3. graph. Separation of the polar constituents (SO2,COz, COS, Ind. Eng. Chem., Process Des. Dev., Vol. 15, No. 3, 1976
367
Table I1 Mesh
+ 100 +200
+325 -325
MnS04-Hz0, w t %
Charcoal, w t %
2.3
0.27 9.98 85.44 4.30
12.2 10.4 75.1
H2S) was done by an 8 f t 4 in. X 0.085 in. i.d. column packed with Porapak Q controlled a t 194 O F , while the noncondensable gases ( 0 2 , Ar, Nz, CO) were split a t 77 "F on a 6.5 f t X 0.085 in. i.d. column containing Grade 544, Type 13X Molecular Sieve from Davison Chemical Co. Total gas flowrate from the reactor was measured with a soap bubble flowmeter a t low flows and with a wet test meter (oil filled) a t higher flows (>150 cm3/min). Preliminary screening tests aimed a t directly decomposing MnS04 with reducing gases (CO, CH4, C3Hs) were made; however, most of the experimental work centered about decomposing manganese sulfate by heating it in admixture with charcoal powder. Attempts were made to pelletize the powder mixture in order to be able to vary the size. However, no binder could be found that would provide strength under reaction conditions unless large amounts were used. Because of the fact that the reaction involved a mixture of two powders, the introduction of a significant amount of a third component was deemed inadvisable. The reactants used were Reagent Grade MnS04mH20 and a commercial powdered wood charcoal (91.8% carbon). The particle sizes as furnished are shown in Table 11. Since the MnS04eH20 was very hygroscopic and readily formed lumps, the salt was dried a t 265 "F and maintained in storage a t this temperature. Small portions of dried charcoal were added with repeated grinding, mixing, and drying after each addition until the requisite blend had been obtained. The blend was stored a t 265 OF. Inasmuch as the MnSOpHZ0 was dehydrated in the reactor a t 525 "F a t the start of every decomposition run, it is impossible to give a definite size range for the anhydrous MnS04. However, based on the feed material and the final product it was essentially all -325 mesh (44 p). The charcoal was probably not altered greatly except 325 mesh for the larger sizes and would be in the range -200 (-74 44 p). Because of these limitations no attempt was made to vary particle size, but it was considered to be sufficiently uniform.
+
+
Experimental Procedure A 20-g sample was normally charged to the reactor tube, supported in the sample zone so that the upper and lower thermocouples partially penetrated the sample. The sample was dehydrated a t 525 O F under 50 cm3/min of N2 flow for 2 h. A decomposition run was started by raising the temperature to the desired value. In some cases N2 flow was continued throughout the run; others were done without a stripping gas. From the start of a run, until completion, upper and lower temperatures and total exit gas rate were noted and recorded every 7.5 min. Gas chromatographic samples were analyzed every 15 min. Normally a run was completed when the SO2 gas rate in the exit gases fell below 5 cm3/min. The temperature set point was lowered, and the sample was allowed to cool in 50 cm3/min nitrogen flow. Isothermal run time varied from about 1 to 2.5 h depending upon the temperature. Heating times from -550 to -1300 "F and cooling times from -1300 to -800 "F were on average about 0.5 h each. Samples were taken and gas rates measured during all periods to ensure an accurate material balance. 368
Ind. Eng. Chem., Process Des. Dev., Vol. 15, No. 3, 1976
The range of temperatures used, 1230-1340 "F, was determined by the timing of the gas analysis procedure. Fifteen minutes was required per sample analyzed. A t higher temperatures the run was too short for a sufficient number of samples to be handled-at lower temperatures the flow was either so slow or the concentration so low that the analyses were imprecise. This temperature range was about the same as that studied by Pechkovsky (1959),and the decomposition rates attained would be practical on an industrial scale. For example, reference to Figure 3 shows that 50% of the charge is decomposed in approximately 60 min.
Results Initial screening experiments attempted the direct reduction of MnS04 utilizing carbon-containing reducing gases (CO, CHI, C3H8) in the reactor. The following equations give the competing reactions with CO as the reductant:
--
+ y3co lhMn304 + SO2 + %C02 MnS04 + CO MnO + SO2 + CO2 MnS04 + 4CO MnS + 4co2
MnS04
-
(1)
(3)
The first two reactions would yield a COz/S02 ratio of 51. Reaction 3 fixes the sulfur as the sulfide making the recovery difficult. In using CO as the reducing gas, the off-gas contained high C02/S02 ratios and the solid-phase product was predominantly MnS. The equilibrium data indicate that a CO/ MnS04 ratio exceeding unity would produce the sulfide as the stable product. Apparently it is extremely difficult in a batch system to avoid local CO concentrations above this value. Methane and propane gave results similar to pure CO producing MnS as the solid phase. The introduction of hydrogen from these gases complicates the equilibria in that HzO and H2S can also be formed. In our analytical system H2 and HzO could not be determined quantitatively. The gas analyses showed high ratios of COz/SO2 with practically no H2S, indicating that all of the sulfur was converted to manganous sulfide. Because of the excessive sulfiding produced by the gaseous reductants, experiments with solid reducing agents were undertaken. For each decomposition run with solid reactants, the gas analysis and flow rate data were interpolated, normalized, and integrated for each gaseous constituent as a function of time to give a total material balance on the run. Table I11 gives the data obtained for a typical decomposition experiment, 26P. Shown a t 7.5-min intervals are the temperature, the instantaneous analysis, and flow rates of the exit gases as well as the cumulative volumes of CO2, S02, and their ratios, both cumulative and instantaneous. The temperature a t the bottom was found to be an accurate measurement of bed conditions, in this run varying from 1333 to 1345 "F. Note that during the initial stages of decomposition, COz evolution exceeds SO:! significantly, but subsequently the situation reverses. For the Mn/C ratio of 2 used in this experiment the stoichiometric ratio of S02/C02 in the exit gases should be 2 compared with the overall cumulative value of 2.33 actually found. Nitrogen was always used as a carrier gas during the heat-up and cool-down periods where product gas generation was insufficient to provide the bulk flow necessary for analysis. In some cases nitrogen flow was continued throughout the experiment, but had little influence on the results. Figures 2,3, and 4 illustrate the handling of these material balance data. Figure 2 is a plot of the instantaneous flow rates of SO? and CO2 throughout an experiment showing clearly the "head start" of CO2 evolution and the overall ratio of SO2/CO2 equals 2. (Note the different ordinate scales). Figure 3 shows
Table 111. Run 26P, 1333-1345 "F,Mn/C = 2, Nz Carrier Gas Instantaneous exit gas analysis, vol %
Temp, OF Time,min 0
15 30 45 60 75 90 105 120 135
Top
Bottom
775 1005 1175 1255 1320 1325 1330 1335 1345 1355 1360 1362 1365 1370 1240 1075 955 748 748
865 1079 1202 1286 1345 1330 1333 1336 1339 1338 1337 1337 1339 1340 1120 961 845 880 880
Total gas flow, cm3/min 68.8 70.9 77.5 112.2 160.0 136.4 120.0 115.3 102.5 91.6 85.1 81.6 77.4 74.1 64.9 64.5 64.2 63.9 63.9
con
N:!
co
cos
so:!
2.15 4.82 9.30 18.22 22.12 18.25 15.00 12.60 10.40 8.36 6.66 5.63 4.24 1.86 0.35 0.15 0.03 0.01
97.39 91.83 84.53 52.17 35.01 40.46 45.80 50.38 56.03 63.98 70.45 74.05 79.03 87.20 93.24 96.58 98.83 99.93 100.00
0.46 1.23 1.76 1.91 1.75 0.98 0.45 0.29 0.17 0.11 0.06 0.02
0.00
0.00
0.03 0.08 0.18 0.21 0.13 0.08 0.06 0.04 0.03 0.02 0.01
0.00 0.00 0.00 0.00 0.00 0.00 0.00
0.00 0.00 0.00 0.00 0.00 0.00 0.00
2.09 4.33 27.51 40.90 40.18 38.67 36.67 33.35 27.52 22.82 20.29 16.73 10.95 6.41 3.27 1.15
0.00
Flow rate of each component, cm3/min Time, min 0
15 30 45 60 75 90 105 120 135
Cumulative vol, cm3
c02
co
cos
so:!
COZ
SO:!
1.48 3.42 7.21 20.44 35.40 24.90 18.00 14.52 10.66 7.66 5.67 4.60 3.28 1.38 0.23 0.10 0.02 0.01
0.32 0.87 1.36 2.15 2.80 1.33 0.55 0.34 0.17 0.10 0.05 0.02
0.00
0.00
0
0.02 0.06 0.20 0.34 0.18 0.09 0.06 0.04 0.03 0.02 0.01
0.00 0.00 0.00 0.00 0.00 0.00 0.00
0.00 0.00 0.00 0.00 0.00 0.00 0.00
1.48 3.36 30.87 65.44 54.80 46.41 42.28 34.19 25.21 19.42 16.55 12.95 8.11 4.16 2.11 0.74 0.03
0.00
0.00
topokinetic decay curves based on SO:! for several experiments made a t conditions listed in Table IV. In Figure 4 similar plots are made for a single run (10-P) based on both SO2 and COz. In both of these decay plots the ordinate is the cube root of the undecomposed fraction (of SO2 or CO2). The points for conversion levels of 40%, 50%, and 60% are indicated and are bunched a t the top of each curve. The nonlinear portions a t the start and end of each run represent nonisothermal conditions (heat-up and cool-down). The material balance data were converted td topochemical decay form in terms of both SO2 and COS as the principal sulfur- and carbon-containing species and straight lines were fitted through the central points which represent the periods of constant temperature. Topokinetic rate constant values determined in this manner for all the runs are summarized in Table IV. Inspection of this table shows an apparent dependence of rate constant values on the Mn/C ratio used as well as on temperature. A least-squares method was utilized to determine activation energies for each of the Mn/C ratio series used.
0.05 0.00
Volume ratio SO:!/CO:! Cumulative
Instantaneous
0
0.00
0.00
56
23
0.42
353
496
1.41
665
1168
1.76
925
1840
1.99
1061
2257
2.13
1142
2540
2.23
1170
2680
2.29
1179
2735
2.32
1180
2746
2.33
0.43 0.47 1.51 1.85 2.20 2.58 2.91 3.21 3.29 3.43 3.60 3.94 5.90 18.24 21.49 43.18
Integral Data. Because of temperature drift during the course of each run, a single decomposition temperature could not be associated with the topokinetic reaction rate constant evaluated for the run as a whole. The method chosen for determining activation energy from integral slope ( h ) data was designed to minimize the effect of this temperature scatter. From each topokinetic decay plot of the type in Figure 3, the measured temperatures for each of the four or five points lying on the linear portion of the curve were associated with the rate constant for that run and treated as individual point pairs in a least-squares analysis. As indicated in Table V, there is a trend of decreasing activation energy with increasing carbon content. Discussion The calculated theoretical heat of reaction a t 1342 OF (1000 K ) for the Mn/C = 2 decomposition is: MnS04
+ l/:!C
-
MnO
+ SO:!+ l/&02 (AHlooo = 40 400 cal/mol)
Ind. Eng. Chem., Process Des. Dev., Vol. 15, No. 3, 1976
369
0
- GO2
D
- SO2 Flow
Flow
Time In Minutes
Figure 2. Instantaneous rates of gas evolution (Run 10-P).
j----L-
0
t
30
60
90
120
IS0
180
210
0.2
Time In Minuisr
Figure 4. Topokinetic decay curve Run (10-P)basis-SO2 and Con.
Table IV. Rate Constant Values Determined from Slope Data 0.6
Mn/C ratio Run no. Carrier gas k , h-' ~~
2:l
I
r
t
26-P
I
I
I
I
1
I
I
I
0
30
60
SO
120
150
180
Dl0
Tim.
In
I'
4:3
13-P 12-P 11-P 10-P 14-P
a2
Ylnulas
The true activation energy for an endothermic reaction such as this would be expected t o equal or exceed the heat of reaction. The measured E A is significantly lower and appears to be dependent on Mn/C ratio. Since MnO is the predominant product in all cases, there should be no measurable change in the experimentally determined activation energy with change in carbon ratio; Le., there was always at least the stoichiometric carbon requirement. An explanation for this discrepancy can be sought in two ways: (a) diffusion control and (b) a multiple step reaction process. Because of the limitations imposed by the on-stream analytical system it was not possible to vary the reaction rate sufficiently (i.e., use higher temperatures) t o determine whether the energy of activation would show a shift from a chemical to a diffusional regime. However, there were several results indicating that diffusion was not the controlling factor. Whether or not N p was used as a carrier gas throughout the decomposition run made little difference. In a diffusional control case a definite effect should have been noticed. In the cylindrical tube there is no variation of reaction interface area with time necessary to show a % order for the diffusional case. There is also the fact shown in Figure 4 that both reacting solids (MnS04 and C) appear t o follow the ?3 order. Most importantly, however, the diffusion control model does not explain the change in measured activation energy with increasing carbon ratio. 370
25-P 27-P 24-P 26-P
o.4
Figure 3. Topokinetic decay curves based on SO2 evolution.
Ind. Eng. Cham., Process Des. Dev., Vol. 15, No. 3, 1976
Temp range, OF
~
1:1
17-P 23-P 16-P 19-P 15-P
Nz
-
Nz Nz -
Nz NP
Nz N2
N2
-
0.176 0.310 0.327 0.464
1260-1267 1323-1344 1303-1315 1333-1345
0.134 0.210 0.235 0.299 0.322
1232-1247 1275-1295 1265-1282 1325-1360 1305-1310
0.233 0.351 0.354 0.360 0.513
1230-1240 1260-1277 1282-1290 1275-1285 1305-1315
Table V. Measured Activation Energy Mn/C ratio
E A , kcal/g-mol
211 413
36.3 ZL 3.7 28.0 f 3.1 26.7 f 1.8
1/1
Multiple Reaction Sequence In every decomposition run, i t was noted that CO2 generation always preceded SO2 until well into the run, clearly shown in Figure 2. When both SO2 and CO2 generation of a typical run are plotted on a topokinetic basis, as in Figure 4,the result suggests that pre-reduction of manganous sulfate is occurring, in a more or less shrinking interface manner, prior to SO2 evolution from a different shrinking interface. This effect, coupled with the trend of decreasing activation energy with increasing carbon ratio, suggests that the process proceeds through a series of steps: (1) partial reduction of MnS04 by CO to give an unstable intermediate, MnS03, and gaseous C o p ; (2) thermal decomposition of the MnS03 to form MnO and gaseous SO,; (3) regeneration of CO by reaction of CO2
with solid carbon; (4) diffusion of SO2 and COPout of the bed. This multiple step model is consistent with the experimental evidence in several ways. 1. The significant decrease in decomposition temperature -800 O F with carbon addition can be attributed to the lower stability of the sulfite. No thermodynamic data for MnS03 could be found (DeBerry and Sladek, 1971; Pascal, 1933); however, the average standard free energy of formation differential for a number of sulfates and their sulfites is about 60 kcal/mol. 2. The decrease in apparent activation energy with increasing carbon content reflects the dependence of the SO2 generation step on the preceding reduction step. In effect the measured value is a combined activation energy of two or more sequential processes. The overall reaction is "carbon-limited;" the CO2-carbon reaction is likely to become diffusion controlled, leading to a lower apparent activation energy. 3. Experimentally, very little CO is detected in the decomposition gas. The small amount of carbonyl sulfide detected correlated directly with this CO concentration. 4. Unlike the reduction of MnS04 with a CO gas stream, virtually no MnS is formed in the carbon system. This can only happen if the local CO concentration is controlled near the stoichiometric value a t the reaction site. 5. In the temperature range studied, CO and SO2 react rapidly to give carbonyl sulfide and sulfur vapor (Nalewak, 1974). Except as noted in (3), no COS or S2 was detected, indicating that CO and SO2 are isolated from one another in the decomposing sample.
Conclusions This study has shown the feasibility of thermally decomposing manganese sulfate in the presence of carbon. The inclusion of the appropriate fraction of carbon or carbonaceous material in the decomposition of manganese sulfate lowers the temperature necessary for useful rate of decomposition so that energy costs are reduced, the difficulty of transfer of heat to the reaction zone is less severe, and very high temperature metallurgical problems are avoided. With solid carbon in physical admixture, the decomposition proceeds without the over-reduction (to manganese sulfide) which takes place if reducing gases such as carbon monoxide are used. This results from a sequential step process within the bed where the reducing atmosphere is self-regulating. Because of the time required for gas measurement and analysis, the experimental investigation was limited to maximum reaction rates of about 60% conversion per hour. Certainly, in an application of the method, higher temperatures with resulting faster reaction rates would be used. In all like-
lihood, the primary limitation on an industrial scale would be the transfer of heat to the reaction zone, rather than the rate of reaction. T o avoid excessive dilution of the product gases some of the heat would have to be transferred indirectly, the amount depending on the SO2 concentration requirement for the succeeding step. However, in the type of equipment in which this reaction would be carried out, holding times of 60 min are not uncommon, and the rates a t temperatures used in this experimental work are practical. Finally, the off-gas of the process is sufficiently concentrated in sulfur dioxide to be fed to a modified sulfuric acid plant such as that of the Essex Chemical/Boston Edison/ Chemic0 joint venture.
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Received for revieu December 27, 1974 Accepted J a n u a r y 16, 1976
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