Production of
Alkali Metal Sulfates from
Alkali Metal Chlorides, Sulfur, Oxygen, and Water in Molten Salt Mixtures Yu-Ren Chin’ and Lyle F. Albright School of Chemical Engineering, Purdue Uniuersity, Lafayette, Ind. 47907 A new process has been developed for the production of alkali metal sulfates, bisulfates, or pyrosulfates and hydrogen chloride from sulfur, oxygen, water, and alkali metal chlorides. Several reaction steps are included in the overall process. The first step involves the reactions between alkali metal bisulfates or sulfuric acid and alkali metal chlorides in a molten salt media. The second step occurs in an emulsion of molten alkali metal bisulfates, molten sulfur and/or sulfuric acid or alkali metal sulfates. The sulfur i s oxidized to form sulfuric acid using nitric acid and/or oxygen as the oxidant. Nitric oxide which i s produced can be reformed t o nitric acid in the third step.
T h e Mannheim and Hargreaves processes are widely used for the production of sodium sulfate and hydrochloric acid from sodium chloride and either sulfuric acid or sulfur dioxide and air (Faith et al., 1965: Shreve, 1967). The reactions between the solid salt and a fluid (liquid or gas) in these processes are relatively slow. As a result, the reactors are large and are often, if not always, operated batchwise. Attempts have been made to improve these processes. but the modified processes (Chemical Processes in Europe. 1967) apparently still have relatively slow reaction rates, high temperatures, and corrosion problems. Recent investigations indicate a possible new method for the production of alkali metal sulfates. Albright and coworkers (Albright and Haug, 1967; Haug and Albright, 1965; Pfeiffer et al., 1967) have shown that alkali metal chlorides dissolved in a salt melt react with nitric acid. I t seems quite probable that chloride salts in a melt would also react with sulfuric acid with the production of alkali metal sulfates. In addition, Schoeffel (1962) and Habashi and Bauer (1966) have indicated that sulfur, water. and oxygen react a t high pressures and temperatures to form sulfuric acid. The kinetics of the later process are too slow to be of commercial interest, but the direct reaction of sulfur may still be possible. Based on the above approach. a new process for the production of‘ alkali metal sulfates and hydrogen chloride from alkali metal chlorides, sulfur, oxygen (or air), and water was developed in the current investigation. The process includes the reaction of sulfuric acid with alkali metal chlorides dissolved in molten alkali metal bisulfates and sulfates. I n addition. a new method has been developed for the production of sulfuric acid from sulfur, water. and oxygen. The new process appears to be significantly superior to previous ones. Experimental
Techniques for Reaction of Alkali Metal Chlorides and Acid Sulfates. A standard 1-liter borosilicate glass, roundPresent address. Naugatuck Chemical, Division of Uniroyal, Inc., Naugatuck. Conn. 06770. To whom correspondence should be addressed.
bottomed flask with three top openings was used as the reactor. A thermocouple positioned in the thermocouple well of the reactor was used to measure temperatures to within +0.5” C. A Glascol heating mantle was provided for the reactor, which contained a flat-bladed glass impeller. An aqueous solution of the desired alkali metal chloride was stored in a glass reservoir provided with a reading scale on the wall of the reservoir so that the volume of the liquid inside the reservoir could be read to within 1 ml. The flow rate of the chloride solution to the reactor that contained liquid alkali metal bisulfate or sulfuric acid was measured with a rotameter. Two gas washing bottles filled with caustic solutions were connected in series to the outlet of the reactor to collect the hydrogen chloride gas from the reactor. After a weighed amount of salt melt or aqueous salt solution was added to the reactor and after the temperature was adjusted to the desired level, the chloride salt was added either in an aqueous solution or as a dry salt to the reactor. In some runs, nitrogen was used t o strip the resulting hydrogen chloride from the liquid. The amount of hydrogen chloride absorbed in the wash bottles was determined by titration. At the end of a run, the melt was sampled and analyzed for chloride ions by the Mohr method (Scott, 1962). The conversions of chloride salts were calculated by two techniques: the amount of chloride ions left in the flask (or reactor); and the amount recovered in the bottles. The calculated conversions were in all cases identical within experimental accuracy. Techniques for Oxidation of Sulfur in Molten Mixtures of Sodium Bisulfate, Sodium Sulfate, Sulfuric Acid, and Water. A 300-ml stirred autoclave, Model ABE-305 (Autoclave Engineers, Erie, Pa.), was used as the reactor for the second phase of this investigation. A large glass tube of ll./l-in. id and 6”i-in. length was used as an inner liner to protect the metal reactor from the corrosive liquids involved. A stainless steel spring positioned under the tube held it firmly against the cover of the autoclave. A solid Teflon lY,-in. od impeller was attached to the bottom of a steel stirrer shaft. Teflon tape was wrapped around the stirrer shaft, cooling coil, and thermocouple Ind. Eng. Chem. Process Des. Develop., Vol. 10, No. 1, 1971
1
wells to minimize corrosion. Stainless steel pressure gages which could be read to within i l psi were used for the pressure measurements. Thermocouples were positioned in the thermocouple wells to measure both the gas-phase and the liquid-phase temperatures. Weighed amounts of molten salts (or aqueous salt solutions) and sulfur were added to the glass tube inside the reactor before a run. Air was first removed from the system by venting as the reactor was heated to the desired temperature for the run. Yitric acid and/or oxygen were introduced as oxidants through appropriate feed lines. The feed reservoir for the nitric acid was a thick-walled glass manometer, which was pressure tested to 250 psig. Nitrogen a t about 180 psig was used to pressurize the nitric acid in the manometer. A reading scale attached on the manometer could be read to within 0.1 in., and the amount of liquid in the manometer could then be determined t o within 0.08 ml. Up t o about 6 ml of nitric acid could be added t o the reactor from the manometer within several seconds. During a run, the temperatures and pressures of the reactor system were read a t frequent intervals. The liquid mixture in the reactor was analyzed for sulfate content before and after the run to determine the conversion of sulfur to sulfuric acid or alkali metal sulfates. Results
Reaction of Alkali Metal Chlorides and Bisulfates or Sulfuric Acid. Molten NaHSOI.H20 and an aqueous solution of sodium chloride were reacted a t the following conditions: Temperature of salt melt Pressure of reactor Concentration of KaCl in aqueous solutions Moles of HaHSO,. HJO used as feed Amount of NaCl added Kitrogen purging period at end of run Agitation speed
1150 to 344' C 1 atm
175 to 262.5 gramslliter 1.8 to 5.8 moles 0.18 to 0.63 moles 30 min 850 rpm
-
The main reaction was essentially as follows:
NaCl(aq) + NaHS04(liq)
NanS04(liq)+ HCl(g) (1) At 115" to 120"C, only small amounts of hydrogen chloride were evolved and conversions of sodium chloride were obviously low. As the temperature increased from 125" to 150"C, considerably more hydrogen chloride was evolved as an aqueous solution of sodium chloride was added to the reactor andlor as the system was purged with nitrogen. The reaction was fairly fast a t temperatures higher than 160°C. About 82% or more of the chloride ions were recovered in the absorption solutions before nitrogen purgings a t 16OCC,and 98 to 100% recoveries were obtained after purging. At higher temperatures, the hydrogen chloride evolved a t almost a 100% rat,e as the sodium chloride solution was added. Gas purging obviously had little effect a t such conditions, except to strip steam from the melt or solution in the reactor. The rate of agitation or the concentration of the chloride solution had little effect on the conversion of chloride salts in the range investigated. When solid sodium chloride was added to molten hydrated sodium bisulfate, the melt tended to foam significantly unless the salt was added slowly. The molar ratio of sodium bisulfate used to sodium chloride added varied from about 3.5-to-1 to 10-to-1. After significant 2
Ind. Eng. Chem. Process Des. Develop., Vol. 10, No. 1, 1971
amounts of sodium chloride were added, a white precipitate formed in the melt. The melt and precipitate were separated using a hot Buchner funnel of medium porosity. Samples of both filtrate and residue were weighed and analyzed for chloride ion, sulfate ion, and acidity. No significant amounts of chloride ions were detected in any of the samples. The compositions of the samples were calculated using the following assumptions: The pyrosulfate, if any, in the mixture was reported as bisulfate: no hydrated water was present in the filtrate or in the residue (probably most of the water of hydration had evolved with the hydrogen chloride during the reaction). The compositions of Figure 1 are reported as if the filtrates and residues contained only sodium bisulfate and sodium sulfate. Samples of filtrates and residues a t both 142" and 320" C were analyzed using X-ray diffraction. Based on these analyses, area A in Figure 1 consisted mainly of hydrated sodium bisulfate and a complex of sodium sulfate and sodium bisulfate, and area B consisted of sodium pyrosulfate and the above complex. Line C represents the compositions of the residues recovered by filtration a t temperatures up to about 275°C. The actual precipitates formed in the reactor are thought to be mainly the complex of sodium sulfate and sodium bisulfate. The hydrated sodium bisulfate detected in the residue by X-ray analysis was caused, in part a t least, by incomplete separation of the precipitate and the filtrate. A transition area or line D existed between 250" and 300" C. Line E represents the residues obtained by filtration above 300° C. The residues here were primarily sodium sulfate 111. V. or other forms of the salt. The composition of the liquid phase in the reactor was varied over a wide range from mixtures containing mainly water and sulfuric acid, to melts of primarily sodium bisulfate, and finally to sodium bisulfate melts containing appreciable amounts of dissolved sodium sulfate. The rates of reaction of the sodium chloride with these solutions or melts were fast for all runs made a t 160°C or above. The reactions were then one of the following:
NaCl(s or aq)
+ H?SO,(liq)
-
NaHSO,(liq) + HCl(g) (2) NaCl(s or aq) + NaHSO,(liq) Na2S04(liq)+ HCl(g) (3) These reactions are probably reversible. and the rates
50 b'
400
w
a
3 I-
2
300
200
loor 0
I
10
20
I
I
30 40 50 60 MOLE YO Na2SQ
70
I
80
Figure 1. Relationship between temperatures and compositions of filtrates and residues for NaHSO4-Na601 mixture
of reaction would then depend to a considerable extent on the solubility of hydrogen chloride in the liquid. At the higher temperatures, hydrogen chloride is apparently rather insoluble in all melts or solutions. Several runs were made at atmospheric pressure in which an essentially saturated solution of potassium chloride was added to molten potassium bisulfate a t temperatures from about 220" to 250" C. Mixtures of hydrogen chloride and steam were evolved as the aqueous solutions were added. About 80% or more of the hydrogen chloride was evolved during this period. Steam stripping removed additional hydrogen chloride so that only 1 to 5% of the original chloride ions were left in the melt. Slightly higher conversions were obtained as the temperature was increased. Salt melts containing up to about 13 mole % potassium sulfate and the remainder mainly potassium bisulfate were produced. No evidence of solid precipitates was noted in these runs. Thermal Stability of Alkali Metal Bisulfates. A melt of hydrated sodium bisulfate ( N a H S 0 4 H . 2 0 ) at 100" C and atmospheric pressure slowly lost weight. I n about 20 hr, the weight loss was equivalent to that of the hydrated water. At 2OO0C, rapid evolution of steam was noted for 5 to 10 min; and a 15% loss of weight occurred in about 5 hr. The loss increased slightly in the next 20 hr. Small white crystals resulted in the melt after about 20 hr total. Sodium pyrosulfate was probably formed as follows:
2NaHS04(liq)-+ Na2S207(liq)+ HzO(g)
(4)
At 250" and 29OoC, weight losses for the NaHS04.H,0 were 18 and 2 3 5 , respectively, after 25 hr. The salt melt in these cases gradually solidified with losses of 16% or higher. Potassium bisulfate also decomposed with time to form potassium pyrosulfate and steam. Weight losses for this melt were about 2, 3, and 6% after 36 hr at temperatures of 180°, 210", and 230" C respectively. Thermal decomposition of the alkali metal bisulfates was prevented if the salts either in a solid or molten state were maintained a t a sufficiently high partial pressure of steam. The vapor pressure of NaHS04.H 2 0for example is about 13, 28, and 58 psia at 140", 170", and 270"C, respectively. With sufficient pressure, steam or water reacts with alkali metal pyrosulfates to reform the alkali metal bisulfates. Oxidation of Sulfur with Oxygen. A single-step process for the production of sulfuric acid from sulfur, oxygen, and water would be attractive for the development of the proposed new method for the manufacture of alkali metal sulfates. Schoeffel (1962), Habashi (1967), and Habashi and Bauer (1966) have investigated such a process. The latter investigators who used a titanium-lined reactor reported success in obtaining faster kinetics for the reaction when they employed certain ionic catalysts. Cupric ions obtained by adding cupric sulfate were reported to be especially good catalysts. In the present investigation employing a glass-lined reactor, no beneficial results were noted however, when cupric sulfate was used as an additive a t operating conditions similar to those employed by Habashi and Bauer. When cupric nitrate or sodium nitrate was used as an additive in runs with sulfur, water, and oxygen, the rate of sulfuric acid formation was increased significantly. Nitrogen dioxide was present a t the end of batch runs
in the oxygen-rich gas phase. Most of the sulfur oxidation occurred in the early stages of the run. Results imply that the main oxidation steps involved reactions between nitric acid and sulfur and not between oxygen and sulfur. The nitric acid was undoubtedly formed by hydrolysis. The nitrogen dioxide was produced by reaction of nitric oxide with oxygen. Oxidation of Sulfur with Nitric Acid. The following reaction occurred readily a t temperatures from 140" to 300" C.
At 100°C, which is below the melting point of sulfur, no significant reactions were noted. I n most runs, liquid sulfur was emulsified with either a salt melt or an aqueous solution containing sodium bisulfate. Liquid nitric acid was added to the emulsion. In several runs, liquid nitric acid was also contacted directly with liquid sulfur. Significant increases of temperature were noted almost immediately after the addition of the nitric acid in both the liquid and gas phases of the reactor. Such temperature increases of course depended on the overall rates of reaction. I n addition, the reactor pressure increased rapidly especially in the early stages of the run and eventually leveled off. The pressure rise in the reactor was caused almost completely if not exclusively by the moles of nitric oxide produced. The moles of nitric oxide formed were calculated based on the pressure rise of the reactor and the final gas temperature, and the calculated value was generally equal within &2% of the moles of nitric acid added. Sulfate analyses of the liquid contents of the reactor before and after a run, confirmed the production of sulfuric acid by the above reaction and the complete reaction of nitric acid in most runs. One oxidation run was made in an open, glass test tube mounted on a ring stand. The reaction mixture of hydrated sodium bisulfate and sulfur was maintained a t temperatures from about 130" to 170°C. A small glass impeller was used to mix the liquid phases maintained at atmospheric pressure. Nitric acid was added in small batches to the system. This run indicated that: The nitric acid initially dissolved primarily in the melt (or salt) phase when a melt-continuous emulsion was present; the reaction occurred primarily a t the interfaces between the salt and sulfur phases and resulted in three phases in rather intimate contact; the salt and sulfur phases were of similar densities. The relative or apparent densities of these two phases could change depending on temperature or the evolution of the gas as the reaction progressed; nitric acid reacted readily when it contacted liquid sulfur to produce nitric oxide and sulfuric acid; and sulfur did not adhere to the walls of the glass reactor at lower temperatures but did at higher temperatures. The visual run confirmed the complexity of the reaction system involving three phases. The viscosity characteristics of sulfur were also a complicating feature. Sulfur melts a t about 110" to 120°C depending on the type of sulfur used. At 130" to 150°C, sulfur has a relatively low viscosity. The viscosity increases to a very high maximum at about 210°C. The viscosity of the sulfur then decreases a t still higher temperatures. Kinetics of Sulfur Oxidation with Nitric Acid. The following first-order equation correlated most of the kinetic results obtained for the reaction between sulfur and nitric acid: Ind. Eng. Chem. Process Des. Develop., Vol. 10, No. 1, 1971
3
WEIGHT % N d S O (in N ~ H S O ~ - HM~IO XT~RE) Figure 2. Effect of salt comDosition on the reaction rate constant k , min-' at 170" C 'liquid-phase temperature and 700 rpm Ridge of maximum rate constants,
k, inin-'
The moles of nitric acid ("So?) present a t any given time during a run were calculated indirectly. The pressure rise was used to calculate the moles of nitric oxide present, and by difference the moles of nitric acid left were then determined. The value of h was essentially constant during most of a run. Somewhat higher values of h were often found during the initial stages of a run. This latter phenomenon was caused in part a t least by partial vaporization of the nitric acid while the acid was being heated and mixed with the salt (or water) phase of the system. The h values reported here are those obtained after the initial stages of the runs, and several repeat runs indicated they were reproducible within 10% or less. These h values are not true kinetic rate constants since they depend on physical factors as well as temperature. Composition of Salt Phase. Two series of runs, one a t 140°C and one at 17OoC, were made in which 5 grams of sulfur were incorporated in about 120 ml of emulsion. The composition of the salt phase (melt or aqueous solution) in the emulsion was varied significantly in these runs, and the speed of agitation was maintained a t 700 rpm. Figure 2 shows the contour map of h obtained at 170°C. This contour map and the one at 140°C (not shown here) were developed using auxiliary graphs drawn for salt phases containing constant ratios of sodium bisulfate to water but containing varying amounts of sulfuric acid or sodium sulfate. Variations of the k values with composition may have been caused in part by changes of the chemistry of reactions. The variations were probably more affected however by significant changes of the physical properties which significantly affected the character of the emulsion obtained. Since the main reactions occurred a t the interface of the liquid phases, all factors affecting the interfacial areas are important. Amount of Sulfur Used. The amount of sulfur used in 4
Ind. Eng. Chem. Process Des. Develop., Vol. 10, No. 1, 1971
an emulsion of sulfur and melt would obviously have an important effect on the interfacial area between phases. Figure 3 shows results obtained for runs at 140" and a t 700 rpm. Salt melts of two different compositions were employed with various quantities of sulfur. The results of these runs plus others not shown indicate that the k values pass through two maxima. The maxima at lower amounts of sulfur are thought to be for emulsions which have the salt melt as the continuous phase whereas the other maxima are for emulsions which are sulfur continuous. The k value at 140" C and 100% sulfur was determined by adding nitric acid directly to liquid sulfur and at the same time providing agitation. The amounts of sulfur in the salt-sulfur emulsion were also investigated in runs made a t 170" and 200°C using two levels of agitation (700 and 1200 rpm). Because of the high viscosities of sulfur at these temperatures, sulfurcontinuous emulsions such as obtained with large amounts of sulfur could not be obtained in the apparatus used. The runs had to be limited then to emulsions in which the salt phase was the continuous phase. Figures 4 and 5 show the results obtained a t 170" and 200" C, respectively. The k values for pure sulfur at these temperatures were determined in each case by adding nitric acid directly to the surface of unagitated liquid sulfur. Agitation. Increased agitation resulted in higher rates of reaction as indicated by the results of Figures 3 and 4. The higher values of h are undoubtedly caused by larger interfacial surface areas between the two liquid phases. Agitation also improved the mixing of the nitric acid feed with the melt phase and the transfer of the acid to the interface. Some agitation was also provided as the nitric oxide bubbled upward through the melt. Of interest, h values obtained near the end of a run were somewhat lower than those obtained during the main portion of the run. This difference is apparently caused primarily by decreased agitation because of less gas evolution. Temperature. The kinetics of the oxidation reaction increased with temperature in the range from 140" to 200" C as indicated by Figures 3, 4, and 5 . One run was also made a t 3OO0C in which nitric acid was contacted directly with liquid sulfur. The agitator I
I
0.ow ' 0
Figure
IO
I
I
I
I
I
t
I
40 50 60 70 80 WEIGHT % SULFUR ( BASED ON TOTAL LIQUID MELTS )
20
30
90
I
100
3. Effect of amount of sulfur on reaction rate constant at 140°C liquid-phase temperature
k, min-'
W Pure sulfur
0 58.8 wt O h NaHSOr, 5.8 wt YOHzSOr A 75.5 wt Yo NaHSOa, 1.9 wt % HzSOr V 87.0 wt Yo NoHSOI x 82.0 w t O h NoHSOI, 5.6 wt YOHzSOa
1.01
I
I
0
IO
I
I
I
I
I
I
1
40 50 60 70 BO WEIGHT % SULFUR ( BASED ON TOTAL LIQUID MELTS )
20
30
1
90
100
Figure 4. Effect of amount of sulfur and agitation on reaction rate constant k, min-' a t 170" C liquid-phase temperature X 87 w t % NaHSO, w t O h NaHSOt, 9.1 wt % H,SO,
0 78.8
V 62.5 wt
O h
NaHSO?
Pure sulfur
i
2.01
WEIGHT % ( BASED ON TOTAL
SULFUR LIQUID MELTS )
Figure 5. Effect of amount of sulfur on reaction rate constant k, min-' at 200°C liquid-phase temperature a t 700 rpm 0 65 w t Yo H i S o l A 87 wt O h NaHSO?
speed provided was 700 rpm. An essentially complete reaction occurred in about 10 sec indicating a very rapid rate of reaction. The reaction a t this temperature was too fast for accurately measuring the value of h, which is estimated to be in the range of 25 to 30 min-'. In this run, part of the sulfuric acid vaporized or decomposed to form gaseous sulfur trioxide. The h values for runs using only sulfur and nitric acid were employed to determine an apparent energy of activation of about 14 kcalig-mole of nitric acid. Reforming Nitric Acid. A series of runs were made at 140°, 170°, and 2OO0C to determine if the nitric oxide produced during the oxidation of sulfur with nitric acid could be reformed a t reaction conditions to nitric acid. Upon completion of the oxidation step with nitric acid, various amounts of oxygen were introduced into the reactor. The temperature of the gas phase increased significantly several seconds after the oxygen was introduced because of the following reaction:
(7) When less than stoichiometric amounts of oxygen were introduced, the pressure rose for a few seconds as the oxygen was introduced. The pressure then quickly decreased to less than the pressure prior to oxygen addition. These results indicate that the nitrogen dioxide was a t least partially dissolved in the liquid phase. The pressure continued to decrease for a period of time and then began
to rise as nitric oxide was once more evolved to the gas phase. During this period of time, more sulfur was oxidized and more sulfuric acid was formed. The rate of sulfur oxidation was, however, much slower, perhaps at most 20 to 30% as fast, during this period as it had been in the initial phase of the run in which nitric acid was used. Furthermore, the pressure obtained in the reactor after oxygen has been added was never as high as it had been a t the end of oxidation period using nitric acid. The results of these runs indicate that relatively little nitric acid was reformed. The nitrogen dioxide formed seems to be a much less effective oxidizing agent than nitric acid, and it may react with sulfuric acid or the sodium bisulfate to form complex groups such as nitrosylsulfuric acid or sulfonitric acid which may not be effective oxidizing agents. Although the results of the reforming runs were not particularly encouraging, further investigations are recommended to determine if an effective reforming of nitric acid in the oxidation reactor is possible. Discussion
Reaction of Alkali Metal Chlorides and Molten Alkali Metal Bisulfates or Sulfuric Acid. Reaction between alkali metal chlorides and alkali metal bisulfates or sulfuric acid is complicated by the fact that two to four phases may be present during the course of the reaction such as follows: Liquid salt phase of mainly alkali metal bisulfate (or sulfuric acid) ; gaseous hydrogen chloride plus some steam being evolved; solid precipitate of primarily alkali metal sulfates which forms at higher conversions of alkali metal chlorides; and crystals of alkali metal chloride if this method of adding the salt is used. At least two steps apparently occur when aqueous solutions of alkali metal chlorides are added to liquid phase containing alkali metal bisulfates or sulfuric acid. The first step consists of the main reaction in which the alkali metal sulfate and hydrogen chloride are produced. The hydrogen chloride may still be, however, a mixture of protons and chloride ions. The second step which is probably the rate controlling step involves desorption of hydrogen chloride from the solution. Such desorption is favored by increased temperatures. The reaction steps are, however, more complicated when the alkali metal chloride is added as solid crystals to the melt. Since no evidence was found that solid sodium sulfate deposited on the surface of solid sodium chloride crystals, a t least one of the following two series of steps was probably occurring: Sodium bisulfate (or sulfuric acid) diffused from the salt melt to the solid-liquid interface, and then reacted with solid sodium chloride. The sodium sulfate produced either was dissolved in the melt or was ejected as solid particles which were suspended in the melt; or solid sodium chloride dissolved in the salt melt and then quickly reacted with sodium bisulfate (or sulfuric acid) surrounding the particles. The hydrogen chloride then evolved from the boundary layer of the salt melt. In either case, agitation would be critical and the rates of solution of sodium chloride and/or of sodium sulfate would be important factors of the reaction. Besides, foaming or frothing may be a controlling factor when solid sodium chloride is added. Oxidation of Sulfur with Nitric Acid. Mass transfer of reactants t o the liquid-liquid interface and of nitric oxide Ind. Eng. Chern. Process Des. Develop., Vol. 10, No. 1, 1971
5
from the interface are obviously occurring in this complicated three-phase system. The effect of each operating variable on the overall reaction is highly complex. Temperature which obviously affects the true reaction rate constants also affects numerous physical properties of the phases such as viscosity, interfacial tensions, and densities. These latter factors certainly have a major effect on the character of the emulsion. Although the k values determined here can only be considered as pseudo rate constants, the values reported should be reasonable first approximations for designing larger reactors. One of the important factors to consider in any scale-up would probably be effective separation of the gas phase from the liquids. Foaming might be a problem under some conditions. A rate expression for oxidation reaction based on a more mechanistic model would obviously involve some term for the amount of sulfur present a t the interface between the sulfur and salt melt. Nitric oxide blankets portions of the sulfur droplets present in a melt-continuous emulsion. Another factor to be considered is that sulfur, a t least a t certain temperatures, may stick to the walls of the reactor. Measuring or controlling the interfacial areas would certainly be difficult and complicated. Insufficient information was obtained in this investigation to outline a detailed mechanism for the sulfur oxidation. Sulfur dioxide was, however, never detected, and it is probably not an intermediate. The nitric acid possibly attacks the sulfur chains or rings by a series of oxidation steps before the final S-S bond is broken. Proposed Processes for Production of Alkali Metal Sulfates
Results of the present investigation indicate that alkali metal bisulfates or sulfates can be produced from alkali metal chlorides, sulfur, oxygen (or air), and water. Hydrogen chloride would be a by-product. Such a process offers important advantages as compared to available commercial processes. The proposed process could be operated in a continuous-flow manner, and moderate-size equipment would be suitable for a plant with a large capacity. Details of the proposed process would depend on whether bisulfate or sulfate salts were to be produced. The following features are suggested for production of sodium sulfate: Sulfur oxidation. Sulfur is oxidized to sulfuric acid by means of nitric acid in the presence of a melt consisting primarily of sodium bisulfate and sodium sulfate. The salt melt from this reaction step is pumped to the reactor in which sodium chloride is added. Reaction o f sodium chloride with sodium bisulfate or sulfuric acid. Sodium chloride is added either as a solid or in an aqueous solution t o a melt consisting primarily of sodium bisulfate. Sodium sulfate and hydrogen chloride are formed. Hydrogen chloride and steam are evolved as gases from the reaction mixture, which is probably maintained a t elevated pressures t o minimize the evolution of water and the thermal decomposition of sodium bisulfate. Sufficient sodium chloride can be added to form a precipitate of sodium sulfate, and sodium sulfate and bisulfate complex (or sodium pyrosulfate). This precipitate can be separated by centrifuging or filtering, and the filtrate is recycled t o the reactor for the oxidation of sulfur. Production o f pure sodium sulfate. The precipitate containing sodium sulfate, bisulfate, and pyrosulfate salts possibly can be separated to obtain pure sodium sulfate. A better processing technique however may be to decom6
Ind. Eng. Chem:Process
Des. Develop., Vol. 10, No. 1, 1971
pose thermally the sodium bisulfate and pyrosulfate to produce sodium sulfate, sulfur trioxide, and steam (Boros and Lorant, 1963; Tomkova et al., 1960). The sulfur trioxide would be recovered, converted to sulfuric acid, and recycled to the reactor in which sodium chloride is reacted. If either sodium bisulfate or sodium pyrosulfate is obtained in a separation process, it too would be recycled. Nitric acid tower. Nitric oxide from the reaction vessel for oxidation of sulfur would be recovered and reformed to nitric acid using a conventional nitric acid process. The reformed acid would be used to oxidize more sulfur. With efficient recovery and use of the nitric oxide, relatively little make-up nitric acid would be required. Water which is one of the reactants can be added as part of the sodium chloride solution, with the nitric acid, and/or as a feed stream to one of the reactors. If an alkali metal bisulfate were desired as a product, it would be prepared by reacting a molten mixture of alkali metal bisulfate and sulfuric acid with the alkali metal chloride. The above mixture would be prepared by oxidizing sulfur in the presence of the molten alkali metal bisulfate. Alkali metal pyrosulfates could be prepared by the controlled thermal decomposition of molten alkali metal bisulfates. Additional details relative to the proposed process(es) are recommended by Chin (1969). He has discussed the optimum conditions for operating various sections of the process. Acknowledgment
Morton Chemical Co. provided generous financial support for this project and performed X-ray diffraction analyses of several salt mixtures. literature Cited
Albright, L. F., Haug, H. F., U.S. Patent 3,348,909 (Oct. 24, 1967). Boros, M., Lorant, M., Seifen-Oele-Fette- Wachse, 89, 531; 89, 555 (1963). Chemical Processes in Europe, British Chemical Engineering, p 111, November 1967. Chin, Y. R., Ph.D. thesis, Purdue University, Lafayette, Ind., 1969. Faith, W. L., Keyes, D. B., Clark, R. L., “Industrial Chemicals,” 3rd ed, pp 712-17, Wiley, New York, 1965. Habashi, F., Montana College of Mineral Science and Technology, Butte, Mont., personal correspondence, 1967. Habashi, F.,- Bauer, E. L., Ind. Eng. Chem. Fundum., 5 , 469 (1966). Haug, H., Albright, L. F., Ind. Eng. Chem. Process Des. Develop., 4, 241 (1965). Pfeiffer, R. J., DiFranco, V. J., Albright, L. F., J . Agr. Food Chem., 15, 949 (1967). Schoeffel, E. W., U. S. Patent 3,042,489 (July 3, 1962). Scott, W. W., “Standard Methods of Chemical Analysis,” 6th ed, Vol 1, p 332, Van Nostrand, New York, 1962. Shreve, R . N., “Chemical Process Industries,” 3rd ed, pp 274-5, McGraw-Hill, New York. 1967. Tomkova, D., Jiru, P., Rosicky, J., Collect. Czech. Chem. Commun., 25, 957 (1960). RECEIVED for review April 11, 1969 ACCEPTED October 22, 1970 Presented at the Division of Industrial and Engineering Chemistry, 158th Meeting, ACS, New York, New York, September 1969.
