'Sulfur Dioxide as a Raw Material CHLORINATION OF METAL OXIDES WITH SODIUM CHLORIDE AND SULFUR DIOXIDE H. F. Johnstone and R. W.
Darbyshire'
University of Illinois, Urbana,
Th.
e reactions of metal oxides w i t h sodium chloride, sulfur dioxide, and air were studied to
determine if they might serve as possible outlets for the large quantities OF sulfur dioxide that may be recovered from waste gases. A t temperatures between 300' and 700" C. the products are gaseous chlorine and a residue containing metal chlorides and sulfates. In the case of ferric oxide the initial reaction i s apparently the formation of a compound between the oxide and sulfur trioxide, of either the valence type or an adsorption complex, which reacts w i t h the sodium chloride at the point of contact. Among the products of the secondary reaction are chlorine and sulfur dioxide which readily attack iron oxide. Within the range of conditions studied, the reaction i s negligible when the t w o solids are not i n close contact. The vapor pressure o f ferric chloride over molten solutions containing this salt and sodium chloride i s very l o w so that volatilization from such mixtures i s prevented. The industrial applications of the reactions are discussed, and a n e w technique for the production of anhydrous sulfates i s described.
A
NOTHER series of papers presented the experimental results and estimates on several methods of recovering sulfur dioxide from dilute waste gases (IO,11, 12). According to the best available estimates of the production costs, by-product pure sulfur dioxide can be produced for less than that made from crude sulfur a t the present price of brimstone. On the other hand, while the pure gas has somewhat greater value than the sulfur dioxide in burner gases for the manufacture of sulfuric acid or wood pulp, in view of the present small consumption of liquid sulfur dioxide, the production of large tonnages will inevitably reduce the selling price to near the value of the sulfur equivalent. At this point transportation and storage costs become serious, and reduction to elemental sulfur must be resorted to for a t least a part of the product. These are the economic conditions that require smelters, which are now recovering sulfur dioxide from dilute fumes primarily to eliminate an atmospheric nuisance, t o convert the product to sulfur by means of coke 1 Present address, Pennsylvania Salt Manufaoturing Company, Wyandotte, Mich.
111.
(1, 1 7 ) . The additional cost of reduction would be unnecessary if new uses of this cheap product could be found in the heavy chemical industry. I n this paper the reactions of sulfur dioxide on metal oxides in the presence of sodium chloride are presented. Later papers will contain a review of other possible uses in both inorganic and organic chemistry. It will be shown that there is an excellent opportunity for the gas to become an important raw material. HISTORICAL ( 5 )
The reaction of sulfur dioxide with sodium chloride in the presence of air and water vapor was patented as early as 1850 (18). The products are salt cake and hydrogen chloride. After several modifications the method was developed successfully by Hargreaves (7) and was used extensively in Europe during the latter part of the nineteenth century (16). Although numerous improvements mere introduced in the equipment and in the chemistry, such as the use of metal oxides as catalysts, the process proved expensive to operate and ceased t o be of importance. I n this country one company, located on the Gulf Coast where the raw materials originate, continues t o find use for it. The mechanism of the reaction is unfavorable, as shown by Xeumann and Kuna (19), since it requires a slow preliminary conversion of the chloride to sulfite by the action of sulfur dioxide and steam. I n the presence of air the solid sulfite is rapidly oxidized. Apparently, however, the intermediate reaction is responsible for the poor yields, large kilns, and expensive labor which have caused the process to become outmoded as a source of salt cake and hydrogen chloride. I n the absence of steam the reaction of sulfur dioxide and air on sodium chloride produces chlorine (2, 20), a reaction which is also catalyzed by iron oxide. The sulfur dioxide itself may act as the source of oxygen in the absence of air, yielding chlorine, sulfur, and sodium sulfate (24). The Hargreaves reaction employs an excess of salt and a small quantity of metal oxide. Under the reverse conditions the same reactants are used in the chloridization of ores for the purpose of producing soluble sulfates or of volatilizing small quantities of chlorides. Chloridization is an important process in metallurgy and has been studied extensively (4, 8, 16,23). Both sulfides and oxides are converted to sulfates in the presence of sulfur dioxide and sodium chloride. The gas, of course, is not required with the former ores since it is produced during the roasting process. Formation of the sulfate takes place through the intermediate metallic chloride, which is easily converted to the sulfate by means of sulfur dioxide (3). Chlorine gas is known to be present during the reaction, and apparently the combination sf the two gases produces the result. The use of an oxygen acceptor to pro280
INDUSTRIAL AND ENGINEERING CHEMISTRY
March, 1942
mote chlorinations, such as sulfur dioxide, carbon monoxide, as well as sulfur and carbon, is well known. The direct chemical production of chlorine from sodium chloride appears t o be a possibility either through the reactions mentioned above or by the formation first of the anhydrous sulfates which liberate chlorine from salt when the two are heated together in air (IS,16,9626).
281
dioxide added at any time was known accurately from the weight of oil displaced. The air flow, which was, in general, five times that of the sulfur dioxide, was measured by the calibrated orifice meter, N . When water vapor was used in the reaction, the y e s were passed through humidi ers M , maintained at constant temperature. When dry gases were used, the phosphorus entoxide tube was inserted in the ne directly ahead of the reaction tube in place of the humidifier. The exit gases passed by wa of the three-way cock into absorbers F which contained stanJard alkali solution. The residual nitrogen and oxygen were collected in rubber balloons from which the volume was measured and the composition determined. The volume of nitrogen found was used to check the flowmeter readings for total volume of air passed. The alkaline absorbing solution was analyzed for sulfite, sulfate, chloride, and excess hydroxide, From them the amounts of chlorine evolved, of unreacted sulfur dioxide, and of sulfur trioxide formed by oxidation with either oxygen or chlorine were calculated. The sulfate formed by direct oxidation consumed oxy en from the reacting gases, and a correction for this was mate in the volume of ox gen consumed in the reaction tube. When water vapor was a&ed to the reaction gases hydrogen chloride was formed with a resulting consumption of chlorine and liberation of oxygen. The reaction residue was analyzed for the constituents of the ores in the soluble and insoluble portion, and for chloride, sulfate, and sodium. I n most cases the sample was first extracted with absolute alcohol to remove all of the iron chlorides and part of the sodium chloride. The material insoluble in alcohol was extracted with 5 per cent acetic acid, leaving the unreacted oxide. In es~
EXPERIMENTAL
In the present work the reaction of sulfur dioxide and air on mixtures of various oxide ores and sodium chloride was studied. While the products obtained were of definite interest, the principal object was t o study the mechanism by which the reactions proceed. With this information available, a better selection of the type of equipment t o be used on a large scale and of the optimum conditions required for a given product in order to avoid side reactions can be made. The work will be followed by space-time-yield studies of several of the reactions having commercial possibilities. They will be reported later. The ap aratus used is shown in Figure 1. It was constructed entirely of Pyrex glass. Reaction tube A was heated in electric furnace B which had automatic temperature control. The mixture of sulfur dioxide and air was brought into the bottom of the tube at a rate.of 200 to 600 cc. per hour. The flow of sulfur dioxide wm maintained constant by displacement from the reservoir by mercury, which was in turn dis laced by oil through the precision gear pump, C. The rate of 8ow could be maintained within 2 per cent of constancy, and the total amount of sulfur
1
Figure 1. Apparatus for Studying the Reaction of Sulfur Dioxide and Air on M i x tures of Sodium Chloride and Metal Oxides
A. B. C. D.
E. F.
Reaction tube Electric Furnace Oil pump Gas traps A i r- p r e s s u r e regulator Gas absorbers
G. Rubber balloon
H. 1. J. K.
L.
Thermocouple C o l d junction Potentiometer Galvanometer Temperature controller
M. N. 0. P.
Humidifiers A i r orifices Mercury seals Liquid SO2 tank Q. CaCh tube R. Constant-temperature baths
INDUSTRIAL AND ENGINEERING CHEMISTRY
282 TABLE I.
RE.4CTION O F SULFUR
(Charge:
Vol. Length % Reacted Run Temp., Ratio. of Run, Hr. FezOa NaCl No. O C. Air/SOz 1 1
DIOXIDE AND AIR
ON MIXTURES OB PURE
4.10 grams FepOa, 9.00 grams NaCI; gas rate:
Vol. 34, No. 3
FERRIC OXIDE AND
SODIUM CHLORIDBI
100 cc. SO? per hour, air as indicated)
$%&
Ratio of Products t o Sodium ( X 100) in Solid Moles SO1 ( X 100) R e a c t e d ~ ~ s $ ~ d , 0-12 12-24 24-36 36 hr.Fez08 % of Evolution, F e C h FeChZFez(S0a)a FeSO4 2NazSO4 NaCl hr. hr. end Na Na Na Na Na Na hr. Total Hr. Na
...
... ...
...
32 33 28 31
350 350 350 350
1
1
12 24 46 83
41 79 87 67
11.4 23.2 37.0 49.3
2.8 3.1 3.5 3.3
2.4 2.3 2.3
35 26 34
350 350 350
5 5 6
12 26 53
56 75 60
24.2 42.6 40.1
4.4 4.7 4.1
1.7 2.0
36 30 29
400 400
1
1
12 132 48
37 12.3 88 77.6 400 5 64 50 4' 500 5 35 66 27b 360 6 138 39 93 0 Charge WEB 1.5 times t h a t of other runs. b Water vapor added t o entering gases. 0 HC1 was formed instead of Clz.
2.8 4.5 4.6
1.9 2.6
0.9 1.1
5.0 0.6
4.8 1.7
4.6 1.6
3.7 1.4
6.3
...
...
...
0.8 1.3
...
...
0.9
...
...
0.3 1.7
... ...
0.9
...
timating the probable compounds present, the sodium in the alcohol extract was calculated as sodium chloride, and the remainder of the chloride was divided between the ferrous and ferric iron (for the reaction with iron oxide). In the water extract, the chloride was calculated as sodium chloride, and the remainder of the sodium was assumed to be sodium sulfate. The remaining sulfate was divided between the ferrous and ferric sulfate. This distribution is based on the assumptions that there is no exchange of ions during the alcohol extraction, that no ferric oxide is dissolved, and that all iron chlorides are soluble in absolute alcohol. The accuracy of the ratio of ferrous to ferric iron was not great when only small amounts of one or the other were present, so that very large or very small ratios should be disregarded.
1.0 2.3 4.3 4.2
0.7 1.5 1.4 0.7
0.4 4.2 4.0 8.3
4.3 5.4 1.0 1.9
5.7 11.6 18.5 24.7
44.3 38.4 31.5 25.4
29.5 10.5 6.5 6.5
6.2 5.2 1.9
0.6
2.1
0.0
4.5 7.7 9.0
1.1 0.6 3.4
12.2 21.3 22.1
37.9 28.7 27.9
22.0 12.5 20.0
24 5
1.1 0.8 0.8
1.2 0.2 0.2
2.2 13.5 10.2
0.0 0.0
1.7
6.2 38.8 25.0
43.9 11.2 26.0
31.5 6.0 18.0
0 0
0.0
0.0
7.7
0.0
46.6
3.5
30.5
28 36
1.2 6.8 23.7
12 10
0 65 41.5 34 80.5C
...
.~ ,.
0 0 2 16.6
..
..
...
. ..
...
...
...
. ..
.. .
In general, the charge in the reaction tube was 9 grams (0.154 mole) of 10-20 mesh sodium chloride and the equivalent amount of oxide or ore, based on the rincipal constituent. The ore was ground to pass SO-mesh stancfard Tyler screen. The charge was dried in the reaction tube at 600" C. in a stream of nitrogen. After drying, tlhe temperature waq lowered to the reaction temperature and the gas started, The runs were continued either until absorption of the sulfur dioxide had nearly ceased or for a definite time interval. The course of the reaction W R S followed by comparison of the quantity of the gaseous products and reactants and by analyses of the solid residue from different heights in the reaction tube. RESULTS WITH VARIOUS
0.14
I
3
q MO ES FE IN
sbz
TIME
Figure 2.
P
OLIISITE
- HOURS
Reactions with Various Ores
ORES The results of the reaction with mixtures of sodium chloride and nine ores at various temperatures are shown in Figure 2. The ratio of air to sulfur dioxide in every case was 5 to 1. I n the graphs the difference between the moles of sulfur dioxide retained a n d m o l e s of c h l o r i n e evolved indicates directly the amount of the oxide reacting to form chloride or sulfate. This is true since acid salts, such as pyrosulfates or chlorosulfonates, were never found in the residue. On this basis the following conclusions may be drawn from the curves and from the analyses of the products : Hematite ( B r a z i l i a n ) , Fe203, reacted readily at 350" C., with absorption of sulfur dioxide and liberation of small amounts of chlorine. In 22.5 hours 7 per cent of the total chlorine a v a i l a b l e was liberated. At higher temperatures chlorine evolution was increased and reached 61 per cent in 27 hours a t 500" C. The reaction products were
March, 1942
INDUSTRIAL AND ENGINEERING CHEMISTRY
liquids a t 400' and above. A considerable quantity of the oxide was rendered soluble. Bauxite (AlzOa Fez03) reacted slowly below 400' C., but a t higher temperatures chlorine evolution was rapid. About 50 per cent of the ore was rendered soluble. Fusion did not take place up to 600'. Chromite (FeO.CrZOs) reacted slowly below 450' C. Above this temperature sulfur dioxide was readily absorbed and chlorine was liberated. Liquid formed a t 550°, and the residue showed a considerable amount of soluble chromium
+
salts.
283
as in the case of rutile. In the former case the yield of chlorine is only incidental to other reactions which, if allowed to go to completion by longer contact, would consume all of this gaseous product. Thus, a t low temperatures chlorine evolution follows the absorption of sulfur dioxide only after a considerable lapse of time.
REACTION W I T H PURE FERRIC O X I D E I n order to study the chlorination reaction further, pure ferric oxide was substituted for the ore and runs were made under several conditions. A summary of the results is shown in Table I. I n two series of runs (32, 33, 28, 31, and 35, 26, 34) all experimental conditions were kept constant and only the length of the run was varied. The residues, which showed some sintering but no fusion, were separated into four sections according to height. They were completely analyzed, and the course of the reaction was determined by the distribution of the products. One series was made with dilute gases and the other with concentrated. The results of the over-all analyses of gases and solids for these runs are shown in Figures 3 and 4. The curves showing the mole ratios of the gases represent the actual ratios at any given time and not ratios of cumulative quantities. The reproducibility of the work is best shown by the over-all analyses where the data for the several runs in each series agree closely. The sectional analyses of two of the residues are shown in Figures 5 and 6. The results show that ferrous sulfate is formed in the first part of the reaction and is the predominant iron compound a t this point when concentrated gases are used. This is changed to ferric sulfate as the reaction proceeds, owing to the increased oxygen concentration present in the lower portion after the reaction zone has passed up the tube. The ratio of chlorine to sodium in the different zones of the residue shows that chlorine migrates up the reaction tube, either as free chlorine- or as -a chlorine-sulfur compound. The ratio of sodium sulfate to sodium a t the top of the charge indicates that, a t least in the low oxygen *OLE RATIO - Cl&02 concentrations, the chlorine and sulfur dioxide react ! ? WQLE RATIO - OJS02 -b"Q% with ferric oxide to give ferrous chloride and 0.g 0.1e
