= chemical potential w ( r ) = spherically symmetrical Kihara potential
p
Literature Cited
(1) Azarnoosh, A., McKetta, J. J., Petrol. R@ner 37, No. 11, 275 (1958). (2) Culberson, 0. L., w[cKetta, J. J., Petrol. Trans. A T M E 1891 319 (1950). (3) Zbid., 192, 223 (1951). ( 4 ) Deaton, W. M., Frost, E. M., U. S. Dept. Interior, Bur. Mines, Monograph 8 (1946). (5) Din, F., “Thermodynamic Functions of Gases,’’ Vol. 3, Butterworth, London, 1961. (6) Hoffman, D. S., Welker, J. R., Rao, V. N. P., IVeber, J. H.,
A.2.Ch.E. J . 10, 901 (1964). (7) Hougen, 0. A., Watson, K. M., Ragatz, R. A., “Chemical Process Principles,” Part 11, 2nd ed., Wiley, New York, 1959. (8) McICoy, V., Sinanog’lu,O., J . Chem. Phys. 38,2946 (1963). (9) Nagata, I., Kobayashi, R., IND.ENG.CHEM.FUNDAMENTALS, 5, 344 (1966). (10) von Stackelberg, M., Muller, H. R., 2. Elektrochem. 5 8 , 25 (1954). (11) Waals, J. H. van der, Platteeuw, J. C., Advan. Chem. Phys. 2, l(1959).
RECEIVED for review March 21, 1966 ACCEPTEDJune 9, 1966 Work performed under the auspices of the National Aeronautics and Space Administration Research Grant NSG6-59.
A Q U E O U S OXIDATION OF E L E M E N T A L S U L F U R F A T H l H A B A S H I A N D E R W I N L. BAUER Department of Metallurgy, Montana College of Mineral Science and Technology, Butte, Mont.
+
+
The reaction SI 1l / 2 0 2 H2O + HzS04 i s greatly dependent on temperature and oxygen partial pressure. Below ,the melting point of sulfur it is extremely slow, and above this temperature the rate i s appreciable and increases rapidly with temperature. The reaction i s chemically controlled, and the activation energy is 11-75 kcal. per mole. The rate is proportional to p~:’~, thus suggesting that it takes place in two steps: S -t 1 l / 2 0 2 +- SO3 and SO3 H20 +- H2SO4, the first being the rate-determining step. Sulfur dioxide is not an intermediate reaction product. The rate i s affected by the presence of some foreign ions.
+
HE aqueous oxidation of elemental sulfur is of importance Tto the chemical industry as well as to hydrometallurgy. Thus, Schoeffel (6) and Lukas and Zizka ( 5 ) suggested that sulfuric acid can be manufactured directly and quantitatively in a n autoclave by heating an aqueous suspension of elemental sulfur to a temperature of 300’ C. under an oxygen partial pressure of 30 atm. I n hydrometallurgy, the aqueous oxidation of sulfide ores is gaining importance as a method for the recovery of nonferrous metals (7, 2). If the aqueous oxidation of sulfide ores is conducted a t a temperature below 120’ C. and in acidic medium, demental sulfur is always a reaction product. Once this temperature is exceeded, no sulfur is formed. I n spite of its importance, there is not a single published report on the kinetics a.nd mechanism of this reaction. T h e present work was undertaken to throw some light on this reaction.
Experimental
T h e reaction was conducted in a 1-gallon autoclave (Autoclave Engineers Corp.), made of stainless steel and provided with a loose liner, therniocouple well, stirrer, and cooling coil made of titanium. Temperature was automatically controlled to =t2’ C. T h e reaction mixture consisted of 200 grams of pure sulfur powder and 1900 ml. of distilled water. Speed of stirring was 600 r.p.m. and oxygen partial pressure was 30 p.s.i. unless otherwise st,ated. The mixture was heated to the desired temperature, then oxygen was admitted to the required pressure. From this moment the reaction time was measured. When the reaction time was reached, heating was stopped, the oxygen supply was shut off, and cooling water was admitted rapidly. I t tarkes about 2 minutes to cool the autoclave contents to room temperature. The apparatus was then opened, the reaction mixture filtered, and a n aliquot of the filtrate titrated with 0.1N N a O H solution using phenolphthalein indicator. 11:was then possible to calculate the number of moles of H2S04 formed under the conditions of the experiment.
Results
Rate of Reaction. T h e amount of sulfur reacting under a certain set of conditions was found to be linear with respect to time, as shown in Figure 1. The slopes of these straight lines are the reaction rates. Effect of Temperature. When the reaction was studied a t a temperature below the melting point of sulfur, the rate was extremely slow. Thus, at 60” C. and 6 0 - p s i . oxygen pressure, no acid was detected after 2 hours, and when the temperature was increased to 90’ C., only 0.002 mole of HzS04 was formed. When, however, the temperature was increased above the melting point of sulfur, the rate increased appreciably, as illustrated in Figure 1. From these data, the activation energy in the temperature range 130’ to 170’ C. was calculated (Figure 2) and found to be 11.75 kcal. per mole. Effect of Stirring. T h e effect of stirring was studied a t 130’ C. (Table I). I t is apparent that the ratk is independent of stirring in the range 400 to 900 r.p.m. Effect of Oxygen Pressure. T h e effect of oxygen pressure was studied a t 130’ C. (Table 11). These d a t a fit a straight line only when the rate is plotted against fi,,?’* as shown in Figure 3. Effect of Hydrogen Ion Concentration. This effect was studied a t 130’ C. by adding variable amounts of H2S04 to
Table I. Effect of Speed of Stirring (Temperature 130” C.,po, = 30 p.s.i.) Speed of Stirring, Rate, Mole R.P.M. H ~ S 0 4 / 2Hr. 0
0.0011
400
0.0061 0.0068 0.0061
600
900
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469
r-
0.08 I
I
I
1
I
I
-0' al
E 0
r
4
0.06
I
0
2 3 Time (Hours)
I
4
Figure 1. Effect of temperature on aqueous oxidation elemental sulfur
of
Figure 3.
Effect of oxygen partial pressure
Oxygen partial pressure = 30 p.s.1.
the aqueous phase before each run. A sample of this solution was titrated and compared with a similar sample taken after the run. The difference between these titrations gives the amount of acid generated during the reaction. T h e initial acidity has practically no effect on the rate (Table 111). Effect of Foreign Ions. When one of the experiments a t 170' C. was carried out in the autoclave equipped with allGainless steel parts, it corroded badly and a green solution was obtained which contained ferrous ion. The amount of acid formed in 4 hours in this run was twice that formed when no ferrous ion was present. When, however, FeSOr (2 grams per liter) was added initially to the reaction mixture, it was found a t the end of the run that complete oxidation had taken place and ferric hydroxide precipitated.
- 1.5 -1,6
- 1.7 -1.8 -rc
0"-
1.9
-1
-2.0 Table 111. Effect of Hydrogen Ion Concentration (Temperature 130" C., speed of stirring 600 r.p.m.,po2 = 30 psi.) Initial Acidity, Rate, Mole Mole/L. H2SO4/2 HT. 0.000 0.001
-2,l
-2.2
0.0022
0.0023
Figure 2. Table II.
0.0024 I T "K Arrhenius plot
-
Effect of Oxygen Partial Pressure
(Temperature 130" C., speed of stirring 600 r.p.m.) Rate, Mole Po,, P.S.I. H2S04/2 HT. 15 0.0046 30 0.0091 45 0.0247 60 0.0346 75 0.0460 100 0.0749
470
I&EC FUNDAMENTALS
0.010 0.100
0.0025
Table IV. Effect of Foreign Ions (Temperature 130" C., speed of stirring GOO r.p.m., pol = 30 p.s.i.)
Concn. o f Ion Litially Added, G./L. 0 1 2 3 4
cu 0.009 0.036 0.032 0.034 0.031
Rate, Mole H2SOn/2 HT. Zn +Z Ni +% 0.009 0.009 0.020 0.016
o:oia ...
o:oi7
...
co
+=
0.009
0.011 o:oi3
...
T h e effect of other ions was studied by adding variable amounts of their soluble salts to the aqueous phase and titrating samples before and after each run, using methyl orange as indicator. Frlom the results shown in Table I V it can be seen that there is a marked effect caused by C U + ~ , less effect by Z n f Z and Ni+', and practically no effect by C O + ~ . T h e catalytic effect of Cu+' is noted a t 1 gram per liter, and further increase in C U +concentration ~ has no effect.
However, because of the liberation of O H - during the reaction, a secondary reaction takes place :
M+'
-+
M(0H)y
(8)
T h e insoluble hydroxide will block the anodic area, and the reaction will stop. Therefore, an acid should be present to prevent the formation of insoluble products. The over-all reaction will then become:
+ l/zOz + 2 H +
MS
Discussion
+ 20H-
-+
M+'
+ S + Hz0
(9)
Above 120' C., the followipg reaction will be superimposed: The results of the experiments described show that the reaction is chemically controlled since the rate is independent of the speed of stirring and the activation energy is 11.75 kcal. per mole. Further, it is believed that the reaction takes place in two steps:
S
+ l 1 / z 0 z + H20
+ Hz0
+ HzS04
(2)
T h e first step is the slowest, since the rate depends on p0,3'~. If SO? were a n intermediate product-Le., if the reaction steps were
s + 02 so2 soz 4-' / 2 0 2 so3 so3 + HzO + HyS04
(3)
+
(4)
-f
(2)
one would expect the rate to depend on either PO, (if Reaction 3 were rate controlling), or p0:l2 (if Reaction 4 were rate controlling). Reaction 1 probably takes place in the following steps:
+ O2% Ss-,O + S,O Ss, + + s7-z szo+ + SZ-1
S8
(reversible)
(la)
0 2 + !so3
(Slow)
(1b)
!so3
(Slow)
(14
0 2 -+
where x may range from 1 to 4. The fact that the reaction takes place a t an appreciable rate only after the melting point of sulfur is exceeded is of importance in understantding the mechanism of the aqueous oxidation of sulfide ores. I t is believed that in the presence of acids, sulfides are solubilized by an electrochemical mechanism (3) like the corrosiom of metals (4). Thus, for a divalent metal, M, the reactions taking place are: Anodic reaction. Cathodic reaction.
+ S + 2e+ HzO + 2e- 2 0 H -
MS + Me' '/io2
+
(5) (6)
T h e final reaction products will depend on temperature. Below 120' C , elemental sulfur is very stable. Therefore, the over-all reaction will be the sum of Equations 5 and 6 :
MS
+
' / 2 0 2
+ HyO
--*
M+'
+ S + 20H-
(7)
2H+
+ Sod-'
(10 )
Thus, under these conditions, the over-all reaction is the sum of Equations 9 and 10 :
MS SO3
-+
+ 202
+
+ S04+
M+'
and metal sulfate is the exclusive reaction product. 10 is accelerated by some metal ions.
(11) Reaction
Conclusions
Although the present study was aimed at clarifying the mechanism of the aqueous oxidation of elemental sulfur and its connection to the industrial process for the hydrometallurgy of sulfide ores, yet it is felt that this reaction has a potential importance in the chemical industry. Thus, if pure sulfur is available, it would be possible to manufacture sulfuric acid directly in one step in a pressure reactor a t about 300' C. and an oxygen partial pressure of 400 p.s.i. The reaction is very fast and bypasses the conventional burning of sulfur to sulfur dioxide, the catalytic step of converting SOz to SO3, and the final absorption of the gases. Ac knowledgrnent
The authors gratefully express their thanks to V. Griffiths, Head, Department of Metallurgy, for his continuous help, and to F. L. Holdereed, Director of Metallurgy, and T. G. Fulmor, Director of Extractive Metallurgical Research, The Anaconda Co., Anaconda, Mont., for their interest in the project and in supporting it financially. literature Cited
(1) Forward, F. A., Warren, I. H., M e t . Rev. 5, 137-64 (1960). ( 2 ) Gerlach, J., Metall. 16, 1171-9 (1962). ( 3 ) Habashi, F., Econ. Geol. 61, 587-91 (1966). ( 4 ) Habashi, F., J . Chem. Educ. 42, 318-23 (1965). ( 5 ) Lukas, J., Zizka, M., Czech. Patent 101,426 (Oct. 15, 1961); C. A . 58, 7634d. (6) Schoeffel, E. W., Brit. Patent 779,506 (July 24, 1957). RECEIVED for review November 11, 1965 ACCEPTEDMay 16, 1966 Taken from a master of science thesis by E. L. Bauer, Montana College of Mineral Science and Technology.
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