Chlorine and Salt Cake from Salt and Sulfur ARTHUR W. MIXSON -4ND ALVAN H. TENNEY Columbia University, New York, N. Y.
Increased demand for chlorine and salt cake emphasizes the need for new sources of these materials. Anhydrous reaction of sulfur trioxide on sodium chloride offers a possible means of producing sodium sulfate and chlorine simultaneously from cheap chemicals. Previous work on the several reactions involved is inadequate. A laboratory method for their study has been developed. The chemistry of a portion of the process may be elucidated as follows: (1) Sodium chloride and gaseous sulfur trioxide react to form addition compounds of 1,2, and perhaps 3 moles of sulfur trioxide per mole of sodium chloride. This reaction is impeded by crust formation. (2) The higher ad-
T
HE demand for chlorine today exceeds existing capacity for production. Even under normal conditions, present electrolytic chlorine processes are not economically suited for expansion t o meet requirements because of the necessity for disposing of by-product caustic soda. Consequently search for feasible methods of chlorine production, without parallel formation of caustic, has assumed unusual importance. This article presents the basic chemistry and some of the kinetics of a process which not only produces chlorine without caustic but also yields anhydrous sodium sulfate as its only other final product. The over-all process from salt and sulfur to chlorine and salt cake is exothermic, as are most of the individual steps, Neither electric power, except for handling, nor appreciable amounts of process heat are required. As in the case of chlorine, there is an economic shortage of salt cake. Prior to the present conflict, more than a third of the United States' consumption of salt cake was supplied by import, principally from Germany (19). Like chlorine, transportation of sodium sulfate is uneconomical. With sodium chloride available cheaply at widely scattered points and ample supplies of easily transportable sulfur in Texas and Louisiana, the process may be utilized to produce chlorine and salt cake economically a t or near the point of use. Notable in this respect is the coincidence of use areas for both products to supply the kraft pulp industry, the largest single industry using salt cake. The following equations indicate the over-all chemical changes required by the process and the stoichiometric relations between raw materials and products:
+ 202 2Soz(g.) + '/zOz 2S(rh.)
2SO&.)
dition compounds are unstable a t room temperature. Above 115' C., NaC1.2SOs decomposes to SO, and NaSOaCl. (3) Between 225' and 300" C., 3 moles of sodium chlorosulfonate liberate 1 mole each of Clz and SOz. (4) The solid residue of this reaction is empirically an equimolar mixture of sodium pyrosulfate and chloride which reacts to form sodium sulfate, chlorine, and sulfur dioxide above 330' C. The reaction kinetics of (3) and (4) have been investigated, and reaction orders and rate constants determined at several temperatures. Comparisons of a possible process using these reactions with existing and abandoned processes for chlorine and salt cake have been made.
The process as embodied in Equations 1 and 2 parallels the familiar contact sulfuric acid process. It will be shown that a more complete statement of the chemical changes required to bring about the over-all reaction indicated in Equation 3 may be represented as: 2
Sodg.)
+ NaCl
-
NaC1(S03), [z 2 2 at room temp.) (4)
NaCl(SO& + iSaCl(SO&
+
(2
- 2)SOa (below 110'
C.) (5)
+ SO8 (above 110" C.) + NaCl f SO2 + Cle (above 225'
NaCl(S0S)z -+ NaSO&l
3NaSOsC1+ NasSz07
may be compound
+ 2 N a C l d 3NatS04-ISO2 +
2NazS.201
/
(6)
C.)
(7) Cle (above 330" C.) (8)
Indications of the presence of a pair of side reactions a t lower temperatures will be discussed. They may be tentatively represented as follows: NaC1(S03). -+ Na(S0&-1 SO&l 2Na(SO3),-1 SO&l-+ Na&07
(9)
+ (22 - 3)SOs + SOL+ Clz
(10)
-+
2SOn
(1)
Apparently the rearrangement shown in Equation 9 is slow and takes place in the presence of excess sulfur trioxide a t or near room temperature with the subsequent oxidationreduction, as shown in Equation 10, liberating sulfur dioxide and chlorine. It is to be expected on the basis of the experiments reported that these complicating reactions (9 and 10) will be absent if sulfur trioxide and sodium chloride are reacted above 120' C.
+
2so3
(2)
SUMMARY O F WORK BY OTHER AUTHORS
+ 2 N a C l d Na2SOl +
SO8
SOn recyoled to Equation 2
+ C4
Rose ( l a ) mentions that sulfur trioxide and sodium chloride form addition com ounds. Schultz-Sellack (26) states t h a t the (3) addition compouncf formed (presumably at room temperature) is NaCl(SO8)4. He also states that the compound AgCl(S0a)r 1472
. .
1473
INDUSTRIAL AND ENGINEERING CHEMISTRY
December, 1941
formed from silver chloride and sulfur trioxide partially decomposes, giving the sulfate, chlorine, and sulfur dioxide. Traube (18)claims to have re ared a compound, to which he assigns the formula Na-OS&, %y the interaction for 3 to 4
/
0
ObCl days a t room temperature of the vapors from 70 per cent oleum and powdered salt ( reci itated from alcohol). He resents anal ses (which are s?ightry hi her than the calculatef values for E$04-- and lower for Naf a n t C1-) to prove the composition. He cites the reaction of this product with diamylamine to give diamylamine sulfamic acid, in accord with his roposed formula. Salley (14) investigated the reaction of s&ur trioxide with sodium chloride at temperatures between 295’ and 350’ C. and states that “1.5 moles of sulfur trioxide react with 1.0 mole of sodium chloride to produce 0.5 mole each of chlorine, sulfur dioxide, and sodium pyrosulfate”, and that “in flow runs the decomposition of a low temperature sulfur trioxide-sodium chloride addition com lex has been shown to yield sulfur dioxide and chlorine in exactyy equivalent proportions; this holds at all temperatures between 220’ and 440 in a nitrogen or an oxygen atmosphere.” These conclusions, while in part true, in the light of the present investigation are inadequate to describe the process. Patosz and Rabek (11) erroneously mention that the sodium salt of chlorosulfonic acid melts at about 150’ C. The patent literature discloses the following information: Deacon (4) states that sulfur trioxide, mixed with air or oxygen and passed over or through sodium or potassium chloride with all reactants “heated”, yields chlorine with sulfur dioxide always present. He mentions that the formation of the sulfate is a kind of cementation from the exterior to the interior of the articles or masses. Schmidt (16) claims that only sodium s d a t e and chlorine are formed by the reaction of mlfur trioxide and sodium chloride, while Clemm (3)suggests that the presence of ox gen is a factor. Iler (7) claims a method for the formation ofalkali metal chlorosulfonates by contacting sulfur trioxide and MC1 a t temperatures between 30’ and 100’ C. when the salt is subjected to attrition. A patent ( 6 ) claims the formation of “metal chlorosulfonates” var ing in composition between 2MCl.SOa and MC1.2SOs, when is a univalent metal, by treating powdered metal chlorides without the application of heat in a mixing a paratus with liquid sulfur trioxide or 96 per cent oleum in the agsence of water. Laury (10) suggests that the decom osition of sodium chlorosulfonate between 150’ and 350’ C. is gest expressed as:
&
+ NsSlOi + SOCln + Cla
4NaC1.508 + Na2S04
He claims the invention of the cyclic process of treating sodium chloride with gases containing sulfur trioxide to form sodium chlorosulfonate; heating the sodium chlorosulfonate, passing the resulting gases through a contact sulfuric acid converter together with oxygen to obtain a mixture of chlorine, oxygen, and sulfur trioxide, and removing the sulfur trioxide by passing these &ases, separated from sodium sulfate, over more sodium chloride. he se aration of the chlorine from the nitrogen andpxygen is then egected by other means such as liquefaction. EXPERIMENTAL PROCEDURE
Preliminary Experiments.
Attempts were made t o react sulfur trioxide (liquid), distilled from oleum in a Soxhlet extractor, with sodium chloride contained in a crucible in the bowl of the extractor’. The product was hygroscopic and so reactive with water that no practical quantitative data on its composition were obtained. Passing sulfur trioxide through salt distributed on glass wool in the body of a Turner absorption bottle at 176” C. resulted in the formation of a brown, extremely viscous liquid, which was not quantitatively investigated because of contamination by the joint lubricant and the difficulty of transfer. Powdered salt placed in a n evacuated glass bomb with liquid sulfur trioxide in excess showed such a tendency t o “gum up” that it was thought reaction would still be incomplete after any reasonable time. Powdered sodium chloride mixed with eight or ten times its volume of Celite analytical filter aid was placed in a loosely packed layer 5 om. deep in the bottom of a 1-cm.
i. d. glass tube. The tube containing this mixture and a sealed ampoule of sulfur trioxide was evacuated and sealed. The liquid sulfur trioxide had not wet the mass completely after a period of 6 months from the time the ampoule of sulfur trioxide was broken, although considerable excess liquid sulfur trioxide remained above the solid material. The sulfur trioxide and sodium chloride had, upon reaction, apparently effectively cemented the particles together so that the permeability of the mass t o sulfur trioxide became practically zero. These results indicated the absolute necessity for exclusion of moisture and stopcock lubricant, and the desirability of avoiding masses of sodium chloride which could be “cemented” by the reaction with sulfur trioxide. The use of stopcocks lubricated with modified pyrophosphoric acid ( 1 , 1 7 ) in an apparatus similar t o the one finally adopted was found unsatisfactory since some of the reagents used combined with the lubricant and made a material balance impractical.
Apparatus. The equipment consisted essentially of an apparatus for dry distillation at constant volume. Details are shown in Figure 1. All parts of the apparatus in the working system were made of Pyrex glass, sealed t o one another with blown joints, with the exception that the spiral tube of the differential manometer Q was fused silica. The iron weights in the magnetic break-seals were encased in Pyrex. During the actual tests, no stopcocks remained in contact with the system. Pyrolysis tube A was a 22-mm. i. d. Pyrex tube sealed at the lower end; its length was 84 cm. to the flare a t bulb B. A s ecial 33-cm., immersion, -in-glass, nitrogen-filled tfermometer TI(over-all length??cxmarked in 2“ C. divisions C. was laced in this tube on a bed of glass from 50’ to 500’ chips. At the bottom of tute A was a cylindrical reel of glass yarn; it consisted of 90 turns of double-twist, I-mm., fiber-glass yarn, wound basket fashion on a reel constructed by sealing seven 1-mm. rods onto and perpendicular to the plane of a 2-mm. small diameter, 15-mm. large diameter glass torus. The reel had an axial wound length of 7 cm. Sodium chloride was distributed on this reel prior to each run. Pyrolysis tube A was placed axially in a tubular furnace F. This furnace was mounted in vertical guides and counterpoised to facilitate raising and lowering for visual observation of the solid in the bottom of tube A . Exploration of the furnace with a Chromel-Alumel thermocouple showed a variation of not more than 2” C. in the central 30-cm. section, at about 500’ C. Control of the furnace, operated on 110-120 volts, direct current, was obtained by a tapered series rheostat. Bulb B was made from the body of a 500-cc. P rex flask to which were sealed side tubes C and C‘. Side tube 8 w a s used to hold am oules of sulfur trioxide inserted through C’. When chlorosulfonic acid was used, hooked ampoules were inserted through tube D (shown dashed) and hung on a Pyrex rod, lying in tube D’ and extending across tube A. This rod was mani ulated by an attached, Pyrex-incased, iron wei ht which sliain with a magnet tube D‘. Withdrawal of this rod back into allowed the chlorosulfonic acid ampoule to fall to the bottom of the pyrolysis tube. !Pa was an ordinary 150” C. laboratory thermometer graduated in 1 O C. divisions, suspended in the air with its bulb touching the surface of bulb B. A thick paperboard shield, Y , protected bulb B from the heat of the furnace. Tube E 4-mm. i. d., 63-cm. length connected tube A from the flare o/ bulb B to the magnetic kreak-seal manifold, G. A 2-mm. capillary tube K, 70 cm. long, connected tube E to coil W. W consisted of a single turn, 60 cm. in circumference, of thin-walled 0.6-1 mm. capillary, and was used to protect the differential manometer from sudden large pressure changes and from vibration. Coil W was connected to the inside member of the differential manometer with 30 cm. of 2-mm. capillary. Manifold G consisted of six break seals built in 10-mm. i. d. tubing sealed to a header of the same size. Figure 1 shows details o! the break seal. Side tube ER, 20 cm. of 4-mm.-bore Pyrex, connected the system temporarily to I-mm. capillary R which, in turn, communicated with calibrated bulb V through stopcock &. At J and successive similar oints on the other break-seal tubes, trap L, containing a seakd ampoule of water M in a
8’
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I N D U S T R I A L A N D E N G I N E E R I N G C H E MISTRY
separate bulb, was connected and opened, as required, t o the system by operating the corresponding magnetic break seal. Differentialmanometer Q was a quartz-spiral Bodenstein gage of familiar design. The position of the pointer of this gage with respect to an opposed stationary needle was determined with a low-power microscope having a 100-division eyepiece scale. Movement of the tip of the pointer in a horizontal plane, as measured on the scale, showed a straight-line calibration against pressure difference (1 division = 0.870 mm. mercury). The jacket of the gage was connected by glass tubing through a stopcock to a rotary vacuum pump and to a carefully constructed, closed Utube, mercury manometer (not shown) of 900-mm. range. The menisci in this manometer were capable of being read precisely to 0.04 mm. of mercury. Reference volume V was calibrated by weighing its content of boiled distilled water at 25" C. Volume of V between XI and St was found to be 541.8 cc.; bore of cock 81, 0.1 cc.; and tube from SI to 2,0.4 CC. The system was evacuated through V and two-way cock Sa or through trap L, without exception. Whenever the system was opened to the atmosphere, except when checking the differential manometer zero, air entering the system was caused to pass through the two U-tubes, UI and, UZ. UI was filled with a 1:5 mixture by volume of PzOb and diatomaceous earth, and Ue with 8-mesh Drierite. These U-tubes each had a packed length of 14 cm. and an inside diameter of 12.5 mm. Air would pass through them at a rate not exceeding 3 cc. per second at a pressure difference of one atmosphere, assuring dryness of the gases in the apparatus. Chemicals. SODIUMCHLORIDE.Eimer & Amend c. P. crystal. SULFURTRIOXIDE. Ampoules of SO8 were prepared by distilling the gas from c. P. fuming sulfuric acid (30 per cent oleum), by the method suggested by Brown (2). The ampoules contained from 0.5 to 3 grams of SOs weighed t o the nearest 0.1 mg. Analysis: 99.9 per cent SO8 by acidity; 99.83 per cent by Bad04 determination. CHLOROSULFONIC ACID. Commercial acid was distilled in an all-glass sealed apparatus through two packed fractionating columns and collected in ampoules with hooked capillary tips. Analysis: acid content, calculated, 25.75 milliequivalents per gram; found, 25.92 milliequivalents. Chlorine, calculated, 30.43 per cent; found, 29.61 per cent. Sulfur, calculated, 27.51 per cent. found, 27.87 per cent. STANDARD hLK.mI. Approximately 0.5 M sodium hydroxide (carbonate-free) was standardized against hydrochloric acid prepared from the constant-boiling mixture. This acid was in turn checked against the mercuric nitrate solution used for chloride determination: STANDARD MERCURIC NITRATE. Approximately 0.01 M prepared by the method of Roberts (12). SODIUM PYROSULFATE. Preparation described in detail under run 5.
Manipulations. The apparatus was oleaned with chromic acid and distilled water when almost completely assembled. Then the necessary final joints were made, and the whole was dried by evacuation. The cleanliness of the operations was attested by absence of discoloration, after exposure t o sulfur trioxide, in the apparatus o r products. I n two instances three or four almost microscopic specks of black were found in the tube E , but no other evidences of extraneous material were noted. The dryness of the apparatus during the runs was amply assured by the technique detailed below. Manipulations forming part of the typical procedure follow: Glass-yarn reel X was washed in boiling chromic acid, boiled five times in fresh distilled water, dried in a covered beaker in an oven at 110" C. for 12 hours, placed in a desiccator, and weighed. It was then immersed in sodium chloride
Vol. 33, No. 12
solution of a strength calculated to give approximately the desired weight of salt clinging to the glass when dried by evaporation. (The reel absorbed about 10 cc. of water.) It wa8 then dried in the oven and weighed. The increase in weight was taken as the sodium chloride charged to the system. The glass reel with its attached salt was carefully lowered on the end of thermometer TI into dry tube A . Tube A , containing salt, glass reel, and thermometer TI, was then sealed t o B a little below the junction with D'. The balance of the system as shown in the diagram with break seals intact had previously been assembled. Through tube C' into C was inserted a n ampoule containing a weighed amount of sulfur trioxide, and the end of C' was closed. I n the run using chlorosulfonic acid, this procedure was altered slightly by lowering a n ampoule containing the acid in weighed quantity through tube D and hanging it on a rod crossing from D' t o D" and then sealing D. The furnace was raised to surround tube A in such a manner that glass reel X carrying the sodium chloride was centered radially and vertically in the furnace, and thus brought the immersion mark on thermometer TI on a level with the top of the furnace. While evacuating through V , the sodium chloride was heated t o about 200" C. for several hours and the apparatus tested for leaks with a spark coil. After cooling t o room temperature, the volume of the system was measured by means of calibrated bulb V and pressure measurements, Checks of this volume were obtained t o 1 part in 1000. The system was dried by twice evacuating and admitting PzOs-dried air, evacuated to less than 1mm. pressure, and sealed at R. If not more than 0.2 mm. pressure increase
FIGURE 1. APPARATUS FOR REACTION OF SULFURTRIOXIDE WITH SODIUM CHLORIDE AND SUBSEQCENT DRYDISTILLATION OF ADDITIONPRODUCTS A . Pyrolysis t u b e
R. Capillary connection t o measured
E . Gas reservoir C, C', D, D'. Tubes for inserting am-
S. Connection for vacuum pump
E.
F. G. H.
I. J. K.
L. M.
N. 0. P.
Q.
poules Tubing, connection Furnace Glass break-seal manifold Break seal Pyrex incased iron weight T r a p oonnection Capillary conneotion ( 2 nim., i. d.) Gas trap, cooled portion Gas trap, sealed-in water ampoule Constriction for sealing off t r a p Capillary t i p Constriotion for sealing off p u m p Quartz spiral (Bodenstein) differential manometer
volume flask
Si, Sz, Sa.
Stopcocks Ti. Thermometer f o r measuring N G l temperature Tz. Thermometer for measuring apparatus (gas) temperature U I . Phosphorus pentoxide drying tube U2. Drierite drying tube V . Calibrated volume bulb for measuring volume of apparatus T V . Capillary (thin wall) conneotion X. Glass reel carrier f o r solid material Y. Paperboaid convected heat deflector 2 . Capillary-tipped side tube for opening system
I N D U S T R I A L A N D ENGIN-ERRING CHEMISTRY
December, 1941
was observed after 12 hours, the system was assumed tight and the run proper started. The furnace was brought to the desired temperature as measured by thermometer TI,pressure measurements were made, and the ampoule of sulfur trioxide in tube C was broken by light flaming of the outside of tube C. I n the chlorosulfonic acid runs the rod across tube A at D' was withdrawn by using a magnet to pull its attached iron weight out into D' and allowing the ampoule of chlorosulfonic acid to fall into the heated portion of tube A , where it broke spontaneously a t the higher temperature prevailing. Timing, heating, and pressure measurements for each run will be discussed later. The sampling technique was not varied. Trap L containing water ampoule M was sealed on at J or at corresponding points on the other break-seal tubes as each was used. Then the trap system, as far as the break seal, was evacuated, dried with air passed over P205,and sealed off under vacuum a t the constriction in S. Portion L of the trap was immersed in a Dewar flask of liquid nitrogen. Opening the trap to the system was effected by dropping the Pyrex-incased iron weight inside tube J onto the hook forming the seal. In several cases where the pressure of the whole system then fell to less than 1mm. absolute, it was necessary only to seal off trap L a t constriction N and a t P. If this low pressure did not occur by condensation in a reasonable time, a vacuum pump was connected to the trap at the capillary tipped tube 0. After evacuation of the rubber connection, the rubber tube was bent so that the glass capillary tip a t 0 broke off inside the rubber tube. Evacuation through the liquid nitrogen trap was continued until an absolute pressure of less than 0.5 mm. was reached. Tube 0 was then sealed off at P and the system a t N . The trap was removed while still in the Dewar flask. To return the system to its original volume, tube H was collapsed by heat at the break seal and sealed off, so that only the tiny volume contained in the tip of the hook of the break seal was lost from the original volume. The trap, now removed to a convenient place, was withdrawn from the liquid nitrogen for visual examination and quickly returned to the Dewar flask. Heating the tube protruding from the Dewar flask and containing the ampoule of water, M ,caused the ampoule to break. The trap could then safely be allowed to come to room temperature for weighing and washing out of the sample. Each run was invariably finished by removing a sample, and therefore with an evacuated system. The apparatus was opened by covering the tip of side tube Z with a rubber tube, evacuating the rubber tube and refilling it with dry air through U,breaking the tip of tube Z inside the rubber tube, and allowing dry air to re-enter the system through Ul and Ut. Tube A was then cut off about 8 inches from its lower end, and the thermometer was drawn out. (In none of the runs did any visible material cling to the thermometer.) This end of tube A containing the glass reel and pyrolysis residue was washed with water, and the glass reel was boiled with distilled water twice after the first washings had been removed. These washings, together with those from the washing of the thermometer, bulb B, and the balance of tube A were combined and immediately neutralized with standard alkali.
Analytical Procedure. The trapped gaseous products of the reaction or the pyrolysis residues were dissolved in water, neutralized to phenolphthalein with standard sodium hydroxide, filtered through a Gooch crucible, and evaporated on a steam bath to small volume. These filtrates were then made up to 100 cc. in volumetric flasks, and duplicate determinations of chloride were made on 20-cc. aliquots, and of sulfate on 25-cc. aliquots. I n cases where a 25-cc. aliquot
\
1475
would have given less than 100 mg. of barium sulfate on precipitation with barium chloride, a single sulfate determination was made on a 50-cc. aliquot. Chloride was determined by titration with standard mercuric nitrate and diphenyl carbazide indicator (18). Sulfate was determined by the standard gravimetric method (6).
HOURS AT 60%. AFTER HEATING FROM ROOM TEMPERATURE
100
BO
40
60
20
4
4.0
3.6
8 3.2
3 3
2.8
2.4
eo
I
I
I
I
i
i
0 20 40 60 8 9 1 0 0 HOURS AT ROOM TEMPERATURE AFTER HEATING TO 60°C.
FIGURE 2. BY
SORPTION OF SULFURTRIOXIDE SODIUMCHLORIDE, RUN 1
The products were checked for the possible presence of hypochlorite or sulfite ions with starch-iodide solution or by attempting to decolorize a pale blue starch-iodine solution. No evidence of oxidizing or reducing materials was detected in these tests. DATA AND OBSERVATIONS
I n brief, the procedure was to react either sulfur trioxide or chlorosulfonic acid vapors with sodium chloride on glass yarn a t temperatures varying from 25' to 100' C., and then to decompose the reacted material by heat and evacuation. A detailed account of the method and observations on the several runs follows:
Run 1. 6.613 milliequivalents of sodium chloride were deposited on reel X by the method outlined under "Manipulations". After drying the apparatus, measuring its volume (906.9 cc.), and evacuating to a pressure of less than 0.5 mm. mercury, the sulfur trioxide ampoule in tube C was broken by flaming the outside of the tube. This procedure vaporized 17.794 millimoles of sulfur trioxide into the system and permitted its diffusion into contact with the sodium chloride at 27' C. The absolute pressure in the system rose immediately on breaking the ampoule but by no means to the value which would have resulted if none of the sulfur trioxide had condensed. The temperature of the portion of the apparatus near the salt rose, without external heating, until it was distinctly warm to the touch. After 96 hours a t room temperature, the pressure in the apparatus (50.4 mm.) showed that there were 2.46 millimoles of gas in it. Some solid material, apparently &SOs was noted on the walls of tube A . The salt attached to the glass yarn had changed from its original small sparkling crystal form to a matte white, crusty material somewhat wet in appearance. Heating the pyrolysis tube at 60' C. for 23 hours resulted in the disappearance of what was just called P-SOS and in the solid becoming dry. At this time the pressure in the apparatus was 97.8 mm., corresponding to a free gas quantity of 4.69 millimoles.
DISSOLVED GASEOUS PRODUCTS AND RESIDUES
TABLE I. ANALYSESOF
THE
Run
Milliequivalents Sod-Acid
No. 1
Product No.
... ...
16.20 10.32 4.61 2.60 0.98
25- 60 30-133 125-264 240-500
245
.....
...
1.48 0.34 2.40 1.48 0.96 6.74
9.62 6.28 2.41 2.20 6.54 27.34
11.18 6.68 4.86 3.68 0.06
25- 30 30-160 98-278 236-394
275 380
0.21 2.83 1.43 5.88 10.40
7.45 2.96 1.54 4.5s 16.69
7.72 5.81 2.95 0.54
25-142 132-312 209-391
258 375
....
....
1 2 3
8.15 3.74 1.87 4.31
4.90 3.89 2.09
13.07 1.67 3.87 0.005
30-112 110-271 258-418
258 390
Charge' NaCl HSOsCl
10.14 8.19
..... .....
... ...
1 2 3 4
F
Chargea 1 2 3
F
Charge' 4
Temperature, C. Range Major
13.63 10.11 2.14 1.48 7.57 35.59
F
3
C11.82 0.17 2.40 1.12 0.002 5.61
1
2 3 4
Chargea 2
Vol. 33, No. 12
INDUSTRIAL AND ENGINEERING CHEMISTRY
1476
F
a By calculation b By differenoe.
5.51b 16:38
.... ....
.... ....
...
Furnace residue
...
...
Furnace residue ..... ...
paratus was then evacuated, and product gases were trapped out while maintaining the furnace temperature a t 132" C. (Product 2 was a yellowish-green and white soIid a t liquid nitrogen temperature.) Continuation of heating for 128 hours a t 125" to 166.5" C. and gradually raising the temperature to 184' C. over a period of a n hour showed negligible gas evolution (0.05 millimole). Lowering the temperature to 165" and reheating to 200" C. again showed no appreciable gas evolution. Raising the temperature above this point showed rapid gas formation starting a t 225" C. Figure 3 is a graph of the gases evolved against temperature variation up to 240" C.
...
Furnace residue
...
Furnace residue
2.0
p 1.9 Y
L
1.8
I3
5
1.7
Y
from quantities charged to the systems.
0
40
80
120
160
200
240
280 320
TIME IN MINUTES
The apparatus was allowed to cool to room temperature (25-29" C.) and to remain there for 92 hours. The pyrolysis tube was then heated to GO" C. and held there (59-62" C.) for 98 hours. The progress of gas condensation and reappearance as vapor during this cooling and heating period is shown in Figure 2. The apparatus was then evacuated to less than 0.5 mm., and the gases were frozen out in the trap. (Product 1 was a greenish-white solid a t liquid nitrogen temperature; the analysis is given in Table I.) S l o d y raising the temperature of the material remaining in the furnace to 100" C. over 8 hours caused gradual evolution of 0.96 millimole of gaseous products with a pressure rise to 20 mm. Holding the temperature around 100" C. for another 8 hours resulted in no appreciable change in pressure. Following this period a t 100" C., the temperature was raised to 132" C. for 1 hour. Rapid evolution of gaseous products as measured by pressure rise was noted (Figure 3). Evolution of product gases between 100" and 132" C. was 2.85 millimoles. Holding the temperature a t 132' C. for 4 hours resulted in no further evolution of gaseous material. The ap-
TEMPERATURE, 'C.
FIGURE 3. EVOLUTION OF GASESON HZATING SULFUR TRIOXIDE-SODIUX CHLORIDE ADDITION COMPLEX, RUN1
OF
FIGURE 4. DECOMPOSITION REACTION VELOCITYAT 245' C., RUN1
Holding the temperature of the solid'material at 244" C. for about 50 hours showed evolution of gases as a first-order reaction (Figure 4). At the close of this period the system was re-evacuated, and the products were trapped out. (Product 3 was a yellowish-green solid a t liquid nitrogen temperature; and a white solid condensed in the warmer portion of the tube-i. e., nearer the entrance from the system to the immersed portion of the trap.) Five days more a t about 240" C. showed no further gas evolution. Heating the solid was then continued a t rapidly rising temperatures to 464' C. (Figure 5) with the first appreciable evolution of gas occurring between 350" and 370" C. The evolution of gas ceased a t 464" C.; the pressure in the system actually decreased as the temperature rose to 500" C. and continued to decrease as the temperature was again lowered to 370" C. The temperature vs. gas evolution curve for this portion of the run is shown in Figure 5 . Analyses of all the products are shown in Table I.
330
340
380 420 TEMPERATURE, 'C
460
500
FIGURE 5. EVOLUTION OF GASESIN SECOND DECOMPOSITION O F SODIUM CHLOROSULFONATE RUE 1
December, 1941
INDUSTRIAL AND ENGINEERING CHEMISTRY
1411
After removing product 3, heating was continued with gradual temperature rise to 380' C. in 6 hours; the first significant evolution of gases occurred between 330' and 350' C. (Figure 8). With the temperature maintained a t 380' C. for 70 hours, a second-order reaction was found which produced 1.47 millimoles of gases. Figure 9 is the reaction velocity graph. At the end of this 70-hour period, the system was evacuated through the trap. (Product 4 was a green and white solid.) The furnace residue appeared to have partially melted at this temperature. Further heating of the furnace to 473' C. gave evolution of a small amount of gas (0.17 millimole) which recondensed on cooling. TEMPERATURE, *C.
FIGURE6. EVOLUTION OF GASESON HEATING OF SULFUR TRIOXIDE-SODIUM CHLORIDD ADDITION COMPLEX, RUN2
R u n 2. The procedure in this run was somewhat similar to that in run 1. A larger excess of sulfur trioxide (27.337 millimoles) over sodium chloride (6.744 millimoles) was used. The time of contact of sulfur trioxide vapors a t room temperature with the salt was 100 hours. At the end of this time the pressure in the apparatus was 31.0 mm., showing 1.46 millimoles of uncondensed gases in a volume of 895 cc. The uncondensed gases were removed by evacuation through the trap. (Product 1 was a greenish solid.) During the evacuation the solid salt on the glass reel changed from a wet to a dry appearance. Heating of the solid from room temperature to 154' C, was effected gradually (Figure 6) in 5 hours with a similar gradual evolution of gases to that obtained in run 1. After holding the temperature at about 154" C. for 24 hours without any appearance of further gas evolution (total gases evolved during heating, 2.81 millimoles) the product gases were removed by evacuation through the trap. (Product 2 was a greenish-white solid.) The temperature of the solid in the furnace was lowered to 98" and then gradually raised to 275' C. A gas evolution ws. temperature graph (Figure 6) shows the results in this period of heating. It was again evident that the first ap-
TEMPERATURE,
'C.
FIGURE8. EVOLUTION OF GASES IN SECOND DECOMPOSITION OF SODIUM CHLOROSULFONATE, RUN 2
I30 I
2
4
O
j
120
3n 110 $ 3 P g
100
90
80 70 40
I20
200 280 TIME IN MINUTES
360
440
FIGURE9. SECONDDECOMPOSITION REACTION VELOCITY AT 380' C., RUN2
20
80
I20
160
2W
TIME IN MINUTES
FIGURE7. DECOMPOSITION REACTION VELOCITY AT 275" C., RUN2
preciable evolution of gas started at 225' C. The first-order reaction velocity graph for this reaction at 275" C. is shown in Figure 7. After heating a t 275' for 20 hours the product gases were removed. (Product 3 was a green and white solid.) Total gases evolved before evacuation were 2.44 millimoles.
R u n 3. I n this experiment the charge conditions were reversed; that is, less sulfur trioxide (8.345 millimoles) than sodium chloride (10.399 millimoles) was used. The volume of the system was 894.1 cc. The sulfur trioxide was reacted with the sodium chloride a t room temperature for 12 hours, a t which time the uncondensed gases showed a pressure of 47.5 mm., equivalent to 2.28 millimoles. On heating, in 6 hours, to 94' C., 0.10 millimole condensed and a pressure of 46.3 mm. remained. The temperature was held at 94' C. for 20 hours, during which time 0.08 millimole condensed. Heating further to 142" C. and maintaining this temperature for 5 hours caused an increase in pressure to 65.8 mm., showing an added evolution of 0.88 millimole of gas. At this point the system was evacuated and the gaseous products frozen out (Figure 10). (Product 1 was a greenish-white solid.) Little further gas evolution a t 0.1 to 0.6 mm. was noted on further heating until 225' C. was reached (Figure lo), when rapid generation commenced. A temperature of 258" C. was then maintained for 20 hours. The total pressure (63.4
INDUSTRIAL AND ENGINEERING CHEMISTRY
1478 I
I
1
CONTINUED ON FIG II
Vol. 33, No. 12
2.0 ~
18
p
16 1.4
3
:: 12 Y g IO
4 TEMPERATURE, ' C
FIGURE 10. EVOLUTION OF GASXS ON HEATINGOF SULFURTRIOXIDESODIUM CHLORIDEADDITIONCOMPLEX, RUN3
OB 0.6
TIME IN MINUTES
VELOCITY A%' 11. DECO&fPOSITIONREACTION 258' C., RUN3
FIG-
I10
105
190
2 100 +
3 z
t3 TEMPERATURE, " C
FIGURE 12. EVOLUTION OF GASESIN SECOND
OF SODIUM CHLOROSULFONATE, DECOMPOSITION RUN 3
3
h
B
80 75 ~
340
4x1
500
TIME IN MINUTES
mm.) indicated the evolution of 2.83 millimoles of gas during this period. Figure 11shows the reaction velocity a t 258" C. The system was then evacuated through the trap. (Product 2 was a green and white solid.) Continuation of the temperature rise to 330" C. showed little gas evolution a t about 0.2 111111. pressure. The first appearance of gas evolution occurred between 330" and 340" C. (Figure 12). The reaction was continued a t 375" C. for 48 hours. The velocity graph of this reaction is given in Figure 13. The pressure change indicated a production of 1.48 millimoles of gas. Evacuation of the system through the trap completed the run. (Product 3 was a green and white solid.)
FIGURE13. SECOND DECOMPOSITIox REACTION VELOCITY AT 375" C., RUN3
Run 4. To determine whether the reaction product of the interaction of sulfur trioxide and sodium chloride was sodium chlorosulfonate or not, the decomposition of sodium chlorosulfonate was studied in this run. Sodium chloride (10.137 millimoles) was spread on glass reel X , and an ampoule containing chlorosulfonic acid was suspended in the top of pyrolysis tube A . After the usual drying, volume measurement (904.4 cc.), and evacuation, the sodium chloride was heated to 45' C. and the ampoule caused to fall into the glass reel. Heating was continued until the ampoule broke by internal pressure. This occurred when the salt temperature was 118" C. The furnace was immediately lowered and the reaction of chlorosulfonic acid with sodium chloride proceeded as the salt returned to room temperature. After 22 hours a t room temperature the noncondensed gases in the system amounted to 6.49 millimoles a t a pressure of 132.9 mm. Evacuation of the evolved gases was continued for 21 hours. (Product 1 was a white solid.) During the early part of this period, the temperature of the furnace was raised to 112" C. and maintained for the balance of the evacuation.
TIME IN
FIGURD14.
MINUTES
DECOMPOSITION REACTIONVE258" C., RUN4
LOCITY AT
Heating of the solid in the furnace was continued with a rise in temperature to 220" C. in 2 hours. No evolution of gas was noted, a t less than 1 mm. pressure, to this point. Evolution commenced between 225" and 229" C. The temperature was raised t o 258" C. and held there for 24 hours. During this time 3.93 millimoles of gas were evolved with a rise in pressure to 86.0 mm. The reaction velocity graph (Figure 14) shows the progress of the reaction. At the termination of the 24-hour period, the system was evacuated. (Product 2 was a green and white solid.) Again raising the temperature of the solid residue caused no gas evolution at less than 0.5 mm. pressure until a temperature of 340" to 360" C. was reached (Figure 15). This reaction was continued a t 390" C. for 41 hours. Figure 16
December, 1941
INDUSTRIAL A N D ENGINEERING CHEMISTRY
indicates the rate of the reaction; 1.89 millimoles of gas giving a h a 1 pressure of 41.7 mm. were evolved. The apparatus was then evacuated. (Product 3 was a greenish-white solid.)
0.4
"
o.o o
lis
'
le5
/
I
L53.55&
0
65
1
I
1
Run 5. The purpose of this run was to determine whether the second reaction a t 350' C. and above, found in all previous runs, could be accounted for by decomposition of sodium pyrosulfate. The latter was prepared on the glass reel in situ by dehydration of sodium acid sulfate. The method of preparation follows the procedure of Ishikawa, Masuda, and Hagisawa (8) who reported that pure sodium pyrosulfate could be obtained by heating sodium acid sulfate a t 240-250' C. A solution containing 11.82 millimoles of sulfuric acid and 11.83 millimoles of sodium hydroxide was placed in the bottom of pyrolysis tube A , and glass reel X was inserted. Careful evaporation of this solution in a water bath resulted in the precipitation of the major part of the dissolved sodium acid sulfate on the glass reel. The pyrolysis tube was then attached to the system and the drying completed by evacuation. After measuring the volume of the system (889.2 cc.) and drying by evacuation and admission of dry air several times, the temperature was raised to 250' C. and maintained there for 8 hours. During this period the system was evacuated through a liquid nitrogen trap. To ensure dry-
1479
ness, the system was purged with dry air and re-evacuated twice. After thorough evacuation the temperature was raised slowly to 378" from 179' C. A plot of gas evolved (total 0.193 millimole) against time in this range is shown in Figure 17. Holding the temperature a t 378-382' C. for 12 hours resulted in further evolution of gas (0.231 millimole). Raising the temperature again slowly to 418' C. caused further evolution (0.226 millimole), also shown on Figure 17. Continuation of the heating for 25 hours a t 410-425' C. showed some further gas evolution (0.10 millimole). Raising the temperature 27' more to 452' C. in 5 hours caused the evolution of 0.307 millimole; and 15.5 hours later when the temperature had subsided to 446' C., 0.045 millimole of gas had recondensed. The maximum amount of gas evolved (1.044 out of a possible 5.91 millimoles), the shape of the gas evolution vs. time curves, and the recondensation of gas with small temperature decreases point toward the conclusion that sulfur trioxide in appreciable quantity is in equilibrium with sodium sulfate and pyrosulfate in the temperature range 330450' C.
Analytical Data. The results of the analyses of the several products in each of the runs are indicated in Table I. For reference these data are converted back to millimoles of gas evolved, as calculated from the analyses, and compared in Table I1 with the quantities of gases evolved a t the corresponding times from pressure-volume-temperature values.
TABLE11. GASEOUS PRODUCTS OF REACTIONS Product
No.
No.
1
1 2 3 4
2
1 2 3
.
Millimoles of Gases Calcd.: From analyses From pressure Total 2 X Ch measurement 1.82 7.73 4.69 5.14 0.17 4.17 2.27 2.40 2.40 1.30 1.12 1.21
.
7
Run
Temp. Range, C. 25- 60 30-133 125-164 240-500
4
5.55 3.31 2.41 1.84
1.48 0.34 2.40 1.48
1.46 2.81 2.44 1.47
25- 30 30-160 98-278 236-394
3
1 2 3
3.83 2.90 1.49
0.21 2.83 1.43
3.02 2.83 1.48
25-142 132-312 209-391
4
1 2 3
0.64 3.91 1.98
3:?4 1.87
6.49 3.93 1.89
30-112 110-271 258-418
THERMODYNAMIC CONSIDERATIONS
I n any proposed engineering process of a chemical nature, knowledge of the heats and free energies of the reactions involved is highly desirable. From the best information now available in the literature, the following thermodynamic quantities may be calculated for:
+ SO, + C11 2N&&Or + 2NaC1+ 3N&S04 + SO2 + Cln 3NaC1+ 3SOS NatS,O, + NaCl + Clr + SO1 NazSaO, Na&04 + SO8 2S01
+ 2NaC1+
N&SO1
(3) (8)
(11)
---c
(12)
+
Equation No. 3
8 11 12
TIME IN MINUTES
FIGUFCE 16. SECONDDECOMPOSITION REACTION AT 390" C., RUN4 VELOCITY
A&m
-23 16,470 520 -36:465 20,000
AHi~8.i
- 21,900
-39,210 52,455 30,000
AFo -21,900 -I- 18.2" 39,210 52.65" -52,455 f 53.6" 30,600 36.4T
-
-
These quantities were calculated from the values given by Kelley (9) for the enthalpies and free energies of formation of sodium chloride, sulfur trioxide, sodium sulfate, sulfur
1480
INDUSTRIAL AND ENGINEERING CHEMISTRY
dioxide, and sodium pyrosulfate. The equations for the standard free energies were calculated, disregarding changes in ACp. They are probably sufficiently accurate for temperatures near 25' C. If these equations are used t o calculate the free energies of these reactions a t the working temperatures of this study, it is apparent that, within the validity of the assumption that changes in AC, will be small, all of the above reactions are thermodynamically possible with the exception of the reactions of sodium pyrosulfate.
980
1060
1140
1220
1300
TIME IN MINUTES
FIGURE 17.
DECOMPOSITION O F S O D I U M PYROSULFATE
Since run 5 showed that the decomposition of sodium pyrosulfate is appreciable a t 350" C. and since Ishikawa, Masuda, and Hagisawa (8) noted a partial decomposition of sodium pyrosulfate a t temperatures above 300" C., it is evident that the use of the simple expression for A F " derived above is not justified. This casts serious doubt on the validity of the other similarly calculated expressions for A F O when extended to temperatures very far from 25" C. Unfortunately high-temperature specific heat data are lacking for sodium sulfate and pyrosulfate, nor is published information regarding the thermodynamic properties of sodium chlorosulfonate available. Using known values for the specific heats of the other materials, together with the information listed in this section, does not materially alter the values calculated for the free energies of these reactions at 275" and 375" C. INTERPRETATION OF RESULTS
Original Data. It was felt impractical to present voluminous tables of original data in the form of pressure instrument readings. Every value of recorded original data, with the slight exception noted below, has been shown on the curves, in the tables, or in the description. The conversion of pressure measurements to millimoles of gases is simple; it involves only correction of mercury manometer readings to 0" C. by means of a large nomograph constructed for the instrument used, conversion of the differential manometer readings to pressure difference in millimeters of mercury by the calibration factor 0.870, addition, subtraction, and application of the simple equation given in the discussion of reaction velocity measurements. Therefore it seems unnecessary to report the original data in tabular form. Should it be desired to ascertain what the absolute pressure was in the apparatus a t any time, a simple gas law calculation using the corresponding point from some curve, and the volume of the apparatus (reported in the description of each run) will suffice. No gases other than those evolved by reaction were contained in the system during the experiments.
Vol. 33, No. 12
The only data omitted from these graphs are a few points a t temperatures near 225" C. on Figures 3, 6, and 10, which fell so close to the points shown that they could not be properly plotted.
Precision and Accuracy. Wherever possible the analyses of dissolved products were made on duplicate aliquots. These duplicate results in no case deviated from one another by more than 2.0 per cent of the absolute value and in most cases agreed within 0.5 per cent. The average of these duplicate determinations is reported in Table I. The analyses on the several gaseous products show close agreement between the sum of the C1- and Soh-- and the acidity value. In all cases where adsorbed sulfur trioxide was not likely to interfere, the quantities of gases appearing in the products as calculated from analysis and from pressure-volume-temperature relations show good agreement (Table 11). Assuming a probable inaccuracy in the measurement of furnace temperature as *2O C. and that of the gases in the system a t the ambient temperature as * 1O C., the probable error in the recorded figures for gas quantities in the system was calculated to be =t0.008 millimole a t 55 mm. pressure, At pressures above 5 mm. the error in gas temperature in the apparatus a t room temperature is the controlling factor. Since this error is proportional to the pressure, the measured values of gas quantity greater than about 0.25 millimole have a probable error of not more than 1 per cent. Below 0.25 millimole of gases in the system the pressure measurement contributes the greatest source of error. This amounts to =!=0.002millimole, assuming an accuracy of h 0 . 2 mm. in the pressure reading. The controlling factor for the error in the calculated values of the reaction velocity constants for the reaction measured a t 245", 258", and 275" C. is the temperature variation of the reactants as measured on TI. Manual control of this temperature held this variation to *2" C. Using this value for the temperature variation and the calculated value of 42,000 calories for the experimental energy of activation of the reaction indicates that the reported values of k are probably good to within * 15 per cent. Similarly for the velocity constants a t 380" and 390" C. for the second reaction, the probable error is approximately *7 per cent. I n the case of the 375" C. constant the curvature of the graph makes the value highly uncertain. The reported value is that calculated from the mean slope for the first 30 per aent of the reaction. Sodium Chloride in the Pyrolysis Residues. The figures in Table I indicate a high degree of probability that the sodium chloride remaining at the end of a run never took part in any of the reactions except possibly as a diluent or support. This conclusion is based on (a) the similarity of the course of the reaction in all runs in spite of the fact that the amount of sodium chloride remaining in the residue varied from 0.002 to 5.88 millimoles; ( b ) the observation noted in the preliminary experiments that the penetration of the reaction of sulfur trioside with sodium chloride into the interior of masses of salt is very slow (corroborated by the same observation of Deacon, 4) ; and (c) the apparent independence of the "order of the reaction" on the quantity of this "inactive" sodium chloride. (For brevity, the residual sodium chloride a t the close of a run is hereafter termed "inactive" and the balance "active". No inference that an activated state is involved is intended.) Most of the interpretation in this study is predicated on the validity of this observation.
Reaction of Shallfur Trioxide with Sodium Chloride. The quantit.yof sulfur trioxide condensed on the active sodium chloride a t room temperature was greater than 2
December, 1941
INDUSTRIAL AND ENGINEERING CHEMISTRY
moles of sulfur trioxide to 1 of sodium chloride. This was ascertained very simply by a material balance, using sulfur trioxide condensed] the amount charged, and the amount present in the gaseous state as measured by pressure measurement a t the appropriate times. I n runs 1and 2 where the excess of sulfur trioxide was sufficient and little inactive sodium chloride remained, calculation shows that the quantity of sulfur trioxide condensed was in the neighborhood of 3 moles to 1mole of active sodium chloride. This sulfur trioxide in excess of 2 moles to 1, at temperatures between 25' and 60' and perhaps up to 90' or 100' C., is in major part adsorbed on the solid sodium chloride or a lower sulfur trioxidecontent addition compound. Evidence for this is shown in Figure 2, by the gradual evolution of gases as the temperature of the material was raised from room temperature to 100' C. (Figures 3 and 9), and by the excess of sulfur trioxide found in the analysis of product 1 in runs 1and 2 over that calculated from pressure measurement. I n the case of run 3 (Figure lo), because insufficient sulfur trioxide was present to provide an excess over 2:1, no evolution of gases below 100' C. took place. This evidence supports the claim made by Traube (18), whose method of removing "adsorbed sulfur trioxide" was probably good (absorption from the gases over the addition compound by 95 per cent sulfuric acid), that the compound formed at room temperature is a 2 to 1 mole ratio of sulfur trioxide to sodium chloride. It also substantiates his criticism of Schultz-Sellack (16) that the compound reported by the latter was due to adsorbed excess sulfur trioxide. On Figures 3, 6, and 10 there is a horizontal section of the gas evolution temperature curve terminating a t approximately 110-120' C. From the analytical data and the amount of evolved gases up to this temperature, it may be shown that the quantity in millimoles of sulfur trioxide condensed at just below 100" C. is a good approximation of twice the amount of active sodium chloride present in runs 1, 2, and 3. The evolution of gas, then, a t about 115' C. corresponds to the decomposition of NaC1.2SOa to give NaC1.SOa. Since in runo 1,2, and 3 little chlorine was evolved after removal of the free and adsorbed gases a t room temperature up to 225' C., it seems probable that this compound is in major part NaS03C1.SOs rather than the NaSOsSOaCl as proposed by Traube. The presence of chlorine in the gases evolved a t room temperature warrants explanation. If we accept the customary formulation of sodium chlorosulfonate as 0
I I
Na--OS-Cl,
account for the presence of chlorine in the gases evolved a t room temperature and for the comparative absence of chlorine in the products formed between room temperature and 150' C. On this assumption the sulfur trioxide and chlorine formed a t room temperature would have resulted from the Na(S03),SOsCI (where z is probably 1 but might be higher) on "spontaneous" decomposition; the jog in the temperaturegas evolution curve would have been caused by the decomposition of the NaSO&I.SOs which did not have time to rearrange to the NaSOsSOaCl form. The data are not adequate to prove or disprove the suggestion.
Nature of the 1 :1 S0s:NaCIAddition Compound. The sum of the furnace residue acidity and the Sod-- in the last two gaseous products from each run would represent the available sulfur trioxide a t 225' C., all of which we know to be condensed. I n each run this quantity in millimoles is always larger than or nearly equal to the remaining active sodium chloride at the same point. This suggests the possibility that a compound NaCLSOa may exist. That this compound does exist and is identical with NaS03C1 is believed established b y its similarity in behavior with the sodium chlorosulfonate of run 4. The points of similarity are: (a) decomposition with the liberation of sulfur dioxide and chlorine at temperatures between 225' and 275' C.; (b) liberation of a third of the sulfur content as SO2 and twothirds of the chlorine content as Cla in the first decomposition; ( c ) commencement of appreciable decomposition at a temperature just above 225' C.; (d) decomposition as a first-order reaction; (e) velocity constants of decomposition a t 258' C. within the limits of experimental error being identical (Table I11 and Figure 18) ; (f) liberation of the balance of the chlorine as Clz and one-fourth of the remaining sulfur as SO2 above approximately 330" C.; (y) the second decomposition probably being second order in both cases; and (h) the final solid decomposition product being sodium sulfate (shown by absence of acidity in the furnace residue in runs 3 and 4).
Decomposition of Sodium Chlorosulfonate. The argument already presented for the identity of the 1:l S0a:NaCl addition compound with sodium chlorosulfonate supported by the data of run 4,in which sodium chlorosulfonate was formed direct by reaction of sodium chloride with chlorosulfonic acid, shows that the following reaction represents the stoichiometry of decomposition above 225" C.: 3NaSOsCl-
and if we accept that the 1:l addition com-
0 pound of sulfur trioxide and sodium chloride is in fact sodium chlorosulfonate, an extension of the mechanism necessary for the formation of the sodium chlorosulfonate will serve to account for the evolution of chlorine a t the lower temperatures. It is evident that in the customary formulation of sodium chlorosulfonate the chlorine atom is attached to the sulfur. Since it has been demonstrated that an additional molecule of sulfur trioxide adds to sodium chlorosulfonate, it seems plausible to assume that a portion of the chlorine may shift to the second added sulfur trioxide as proposed by Traube (18). I n runs 1, 2, and 3, respectively, the addition compound remained for successively shorter times below looo C. The amount of chlorine evolved from the salt below 225' C. (Table I) decreases in the same order. If it is assumed that the formation of the NaSOaS03Cl requires time and that this compound decomposes spontaneously at room temperature to give sulfur dioxide and chlorine, we can
1481
NsSnO,
+ NaCl + 802 + Clz
(7)
This is substantiated by these observations: (a) SO2 and Clz to the extent of one-half and two-thirds the content of the material, respectively, are liberated in equivalent proportions (Table I, product 3 of runs 1 and 2, and product 3 of runs 3 and 4); also, oxidizing or reducing substances for I- or Iz were absent in the products after solution in water; (b) a further reaction takes place at higher temperatures and liberates the balance of the C1 and one-fourth the remaining s. No evidence of this further reaction was observed after evacuating product 3 of runs 1 and 2, and product 2 of runs 3 and 4 until temperatures above 300" C. were reached. The reaction of decomposition of sodium chlorosulfonate by heat does not occur below 225" C. (Figures 3, 6, and 10 and discussion of run 4). Despite the lack of exact equivalence in the analyses for Sod-- and C1- in the several dissolved gaseous products, the absence of oxidizing or reducing substances in these products is a definite indication of the equivalence of SO2 and C12. The excess sulfate then necessarily indicates SO*.
Decomposition of Solid Reaction Product of First Decomposition of Sodium Chlorosulfonate. It has been established that the first reaction of decomposition on pyrolysis of sodium chlorosulfonate a t 245", 258", and 275" C. produces a solid material which may be represented by NaCl Na2Sz07or the equivalent in some other molecular arrangement. On further heating, the second decomposition is then a reaction of these products which may be represented by:
+
+
or
Vol. 33, No. 12
INDUSTRIAL AND ENGINEERING CHEMISTRY
1482
+
+
2Na&07 2NaC1+ 3NazS04 SO2 C11 ZNar&Ol 4 2SOS 2NazSOk 2508 2NaC1+ NalSO, SO0 Clz
+
+
+
+
(8) (12)
(3)
The evidence is absence of C1 in the residue from run 1, absence of acidity in the residue of run 4,liberation of SO2 and Clz in equivalent quantities, and evolution of half as much SO2 and Clz as in the first decomposition of NaSOaC1. The question of the molecular configuration of the 3 sodium, 2 sulfur, 7 oxygen, and 1 chlorine atoms in the solid product of the first decomposition of 3NaSOaCl is still in doubt. It seems probable, however, that Equations 12 and 3 represent the course of the reaction since the reaction does not proceed a t an appreciable rate below 330" C., the temperature at which, under similar conditions, the dissociation of Na2SaOr becomes evident (Figure 17).
Reaction Velocities, First Decomposition. Since these reaction velocities were measured by pressure measurements converted to moles of gases evolved, it is pertinent to outline the method used for that conversion. The system containing the gases was essentially at two temperatures, that of the furnace as measured by Tl,and that of the air in the laboratory as measured by T2, with a zone in tube A where intermediate temperatures existed. It was assumed for purposes of this calculation that all the gas in that portion of the system enclosed in the furnace was a t T1and that portion outside of the furnace was a t Te. The volume in the furnace (94 cc.) was obtained by determining the amount of water required to fill tube A , containing glass reel X and thermometer TI, to the level of the top of the furnace. The total volume V' of the system was measured for each run by pressure-volume-temperature relations as described. The amount of gas in millimoles contained in the system a t any pressure, P mm., was obtained from the expression: - 94) p millimoles = 0.01603 ( TI 94 273 -k TV' m
+
where T I , Ta = temperatures read from thermometers Tl and
Too.
The concordance of the values given in Table I1 for gas quantities as calculated from analyses and by this equation justifies the assumption. I n cases where the values do not correspond closely (products 1 and 2 of runs 1 and 2, and product 1 of run 3), the discrepancy has already been explained by evidence of sulfur trioxide adsorption, which would make the pressure-volume-temperature measurement doubtful. The fraction of unreacted sodium chlorosulfonate a t any time t was calculated by taking the difference between the total quantity of gas in millimoles evolved during the reaction and the amount evolved up to time t , and dividing this difference by the total quantity of gas evolved on completion. Plotting the logarithm of this quantity against time gives a straight line over a large part of the decomposition reaction (Figures 4, 7, 11, 14). The first-order reaction velocity constants a t these temperatures, calculated from the slopes of these lines, are given in Table 111.
TABLE111. FIRST-ORDER REACTION VELOCITYCONSTANTS OF FIRST DECOMPOSITION OF SODIUM CHLOROSULFONATB Run No. 1 3 4 2
T o C.
k X 108, M h - 1
245
4.49
14.7 11.8 41.7
258 258 275
f
=k
0.70 2.6
* 1.8 * 5.8
Fig. No. 4 11 14 7
Since these velocity curves are straight lines over a considerable portion of the probable reaction, 3NaSOsC1 -L NaCl NazSz07 SO2 Clg (7) and since gas partial pressures varied from ten to fifty fold, it seems probable that effects such as diffusion or interfacial area had little control over the rate of the reaction. The long straight portion of the curves is also a fairly good indication that the reaction was taking place for the most part under conditions well removed from equilibrium.
+
+
+
Reaction Velocities, Second Decomposition. To obtain useful velocity constants from the data, some assump tions were necessary. The concentration of the equimolar mixture of sodium chloride and sodium pyrosulfate as obtained in each case from the decomposition of sodium chlorosulfonate was taken as unity a t the start of the second reaction. Since both materials, sodium chloride and pyrosulfate, were formed simultaneously by decomposition of crystalline sodium chlorosulfonate, these two solid products (or any other molecular combination of the 3 Na, 2 S, 7 0, and 1 C1 atoms) must have been laid down in almost perfect molecular juxtaposition, regardless of the original quantity of sodium chlorosulfonate per unit volume. This having been the case, significant dilution of the sodium pyrosulfatechloride mixture during the reaction was caused only by the reaction itself (Equation 8) and is assumed to have been linear with the progress of the reaction. The expression for a second-order reaction where the reactants are present in stoichiometric proportions may be written as - '(cz) = k(cx)l dt
where the initial concentration of either reactant is unity, x is the number of moles of either reactant present at time t , and c is a proportionality constant based on linear dilution as reaction proceeds. Then cz represents the concentration of either reactant a t any time t referred to an initial concentration of unity. The slope of the velocity curve as plotted on Figures 9, 13, and 16, converted to molar units, is then ck. Since this slope was found constant a t 380" and 390" C. over more than one third of the reaction, the assumptions made above seem justified. This velocity coefficient (ck) is just as useful in predicting reaction times as k alone would be. Coefficient ck has the units of reciprocal time and is independent of the actual concentration obtaining at any time, provided the sodium pyrosulfate and chloride (or other equivalent configuration) were formed by decomposition of sodium chlorosulfonate (Equation 7). Table IV shows the calculated values of velocity coefficient ck. TABLEIv. VELOCITY COEFFICIENTS FOR REACTION (EQUATION 8) Run No.
T o C.
ck X loa, M h - 1
* 0.48
Fig. No.
16 380 3.19 * 0.2: 9 375 (2.8 * 7) 13 3 a Based on the mean slope of the curve on Figure 13 from 6 t o 38% oompletion of t h e reaction. 4
2
390
G.41
December, 1941
INDUSTRIAL A N D ENGINEERING CHEMISTRY
There are a number of possible mechanisms for this overall reaction (Equation 8) ; each probably involves a series of consecutive reactions. Little can be said from present knowledge as to the most probable course of the changes except that at temperatures below 380" C. a different one of the series becomes the slow or controlling step. The curvature of the graph (Figure 13) is an indication that such may be the case. If the dissociation of sodium pyrosulfate is the initial step, which seems plausible, the shift to the left of this dissociation equilibrium as temperature is decreased may account for the curvature of the graph of the reaction at 375" C. (Figure 13).
Influence of Temperature on Reaction Rates. Figure 18 is the customary log k us. 1/T plot of the simple Arrhenius equation d lnk E -adt RTa for the first decomposition of sodium chlorosulfonate (Equation 7). The points on this curve were obtained from the data of Table 111. The size of the circles representing individual points show the probable precision of the values. On Figure 18 the excellence of the agreement in the rate of decomposition of the sodium chloridesulfur trioxide addition aompound prepared from SO2 and NaCl (solid circles) and that of the sodium chlorosulfonate prepared from HSOsCl and NaCl (dashed circle) is an indication that the 1:1addition compound is in fact sodium chlorosulfonate. The experimental energy of activation for Equation 7 is calculated from this plot to be 42,000 calories. Substitution of ck for k in t h e Arrhenius -12 equation does not alter - 1.4 I I I I I the form, and the experimental energy of activation may still be - 1.8 c a l c u l a t e d from the slope of a log ck us. D -2.0 1/T graph. Using the 4 -2.2 values from Table IV at 380' and 390" C., -2.4 the experimental -2.6 energy of activation for 1.81 183 1.85 1.87 1.89 1.91 1.93 1% Equation 8 was found IOOO/T' K t o be 61,000 calories. FIQURB 18. DECOMPOSITION OF On the graph of these ck SODIUM CHLOROSULFONA~ values the point corresponding t o 375" C. (oalculated from the slope of the curve of Figure 13) falls slightly below the straight line through the points for 380' and 390" C. if the mean slope of Figure 13 is used and slightly above it if the initial slope is used. CHEMICAL AND KINETIC CONCLUSIONS
From the foregoing discussion, the following conclusions seem justifiable: 1. The reaction of gaseous sulfur trioxide with sodium chloride a t room temperature results in the formation of an addition compound containing 2 moles of sulfur trioxide to 1 of sodium chloride on which is adsorbed a variable quantity of sulfur trioxide, depending on the temperature and pressure. 2. This compound decomposes at low preRsures (around 20 mm.) above 110-120° C.; it yields essentially sulfur trioxide and sodium chlorosulfonate. 3. Sodium chlorosulfonate begins t o decompose a t pressures near 1 mm. on heating to above about 225' C. 4. The decomposition of 3 moles of sodium chlorosulfonate a t tem eratures between 245' and 275" C.results in the formation o P 1 mole of C12 and l mole of sulfur dioxide and a solid residue (Equation 7).
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5. This decomposition is a first-order react.ion whose velocity constants have been measured as: T o C.
k X 108, M h - 1
245 258 276
4.49 13.3
41.7
t
0.70
* 2.0 * 5.8
6. The experimental energy of activation for the decomposition of sodium chlorosulfonate as calculated from these values is about 42,000 calories. 7. The solid residue from this decomposition is, most plansibly, a mixture of sodium pyrosulfate and sodium chloride in equimolar pro ortions. 8. This soEd residue from the first decomposition of sodium chlorosulfonate at pressures near 1 mm. begins to decompose appreciably at temperatures somewhat above 330' C. 9. This reaction yields '/e mole of CL, l/g mole of SO*, and mole of NatSOt per mole of original NaSOoC1. 10. The reaction formin SO2 and C12 from the solid residue of the first decomposition NaSOaCl (Equation 8) is apparently a second-order reaction. 11. The velocity coefficients for this reaction (Equation 8) have been measured as:
02
T oC . 380 390
ck X 108, M k - 1 3.19 if 0.22 6.41 * 0.45
12. The experimental energy of activation for the second decomposition as calculated from these values is of the order of 60,000 calories. SUMMARY OF INDUSTRIAL IMPLICATIONS
I n the present incomplete state of our knowledge of the engineering variables involved in the use of the chemical reactions studied here, it is difficult to make any definite prediction as to the ultimate success of a process based upon them. We may, however, suggest the probable shape Qf such an industrial process. Moreover, it is possible to state definitely some facts which indicate that further investigation may lead to a workable process: 1. The production of chlorine and sodium sulfate from sulfur and sodium chloride is chemically possible. 2. From a ractical standpoint the process is probably industrially workabk because: a. The operating temperatures of the several reactions are within the range normally encountered in industry. b. The several reactions may be carried substantially to completion in reasonable time. c. The over-all process may be carried out in steps, each of which is amenable to close control. d. Assuming that SO1 and Clo may be successfully separated and the SOUrecycled, it seems probable that the yields of both products may be made to ap roach 100 per cent on both the sulfur and the sodium chloride%ases. e. The better the purity of the products obtained, the higher will be the yields.
Some idea as to the probable industrial setup to carry out these reactions may be indicated by interpretation of the results of this study to predict optimum conditions: 1. To obtain ood yields of sodium chlorosulfonate from sodium chloride an! sulfur trioxide, the reaction should be carried out in apparatus which provides means for abrading the solid material. The formation of sodium chlorosulfonate is rapid on fresh salt surface but is physically inhibited by a coating of product. This suggests the use of a sigma bladed incorporator, a baffled screw conveyor, or a pebble mill as possible reactora for this stage of the process. 2. Sulfur trioxide should be added t o sodium chloride between 115' and 220' C. to avoid the formation of material of a higher molecular ratio of SOo to NaCl than 1 to 1. 3. If the sodium chlorosulfonate so formed is removed from this reaction to some sort of rabble furnace or rotary oven, it may be desirable to calcine the material in two zones or stages, one at 275300" C. and the other at 380' C. upward. 4. In the first zone two thirds of the ultimate yield of chlorine will be obtained in less than 2 hours at 275' C. and considerably more rapidly at higher temperatures (Figure 7). 5. In the second heating sta e, substantially complete liberation of chlorine gas will be o%tained at temperatures above
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INDUSTRIAL AND ENGINEERING CHEMISTRY
390” C. in less than 24 hours (Figure 16). Higher temperatures
will give much more rapid reaction. 6. It is imperative that water in any form be excluded from the process.
The similarity of this process to two others abandoned as uneconomical makes desirable a discussion of its potential advantages over them. Comparing our method with the Deacon process for chlorine, the following advantages seem evident : 1. The temperatures required are somewhat lower. 2. The reactions involved are substantially irreversible under the conditions for rapid completion. 3. No catalyst seems necessary. 4. A potentially important intermediate product, sodiuni chlorosulfonate, is formed. 5. The process may involve at least one less intermediate step. 6. The materials handled are inherently less corrosive. 7. Chlorine is produced in the absence of nonoondensable gases. Comparing our method with the Hargreaves process for sodium sulfate, the following advantages seem apparent: 1. Control of the process is simpler and a higher purity of product may be expected. 2. The second product of the process, chlorine, is more valuable than hydrochloric acid. 3. The materials handled are less corrosive. 4. An intermediate product of potential value is obtained.
The possible advantages of our process over the present widely used electrolytic method for the manufacture of chlorine seem t o be as follows: 1. Proximity t o cheap electric power is unnecessary. 2. There is a rapidly increasing demand for both salt cake and chlorine, while the demand for chlorine has outrun the demand for camtic soda. 3 Since the over-all process is exothermic, it is theoretically possible to eliminate power and fuel costs almost entirely. 4. For customers who can use chlorine for bleaching, sodium sulfate for sulfate pulp, or sodium sulfite for sulfite pulp, shipment
Vol. 33, No. 12
of the intermediate sodium chlorosulfonate offers interesting possibilities as a means of economical transport of chlorine and sodium sulfate. 5. The intermediate product, sodium chlorosulfonate, is potentially valuable as a sulfonating agent. ACKNOWLEDGMENT The exploratory work of Ralph Miller brought to the authors’ attention the interesting possibilities of the process. His interested criticism during the progress of the study was a valuable asset. The financial support of the Chemical Foundation made the study possible. Patents covering the use of the process will be administered by the Chemical Foundation, Inc. LITERATURE CITED Bodenstein, M., and Fink, C. G., 2. physilo. Chem., 60, 1 (1907). Brown, E. H., J . Chem. Education, 10, 119 (1933). Clemm, Brit. Patent 15,152 (1899). Deacon, Ibid., 1908 (1871). Fales, I-I., and Kenny, F., “Inorganic Quantitative Analysis”, p. 278, New York, D. Appleton-Century Co., 1939. I. G. Farbenindustrie, German Patent 644,222 (1937). Iler, U. 9. Patent 2,219,103 (1940). Ishikawa, Masuda, and Hagisawa, Scieltce Repts. T6hoku Imp. Uniu., 23, 164 (1934). Kelley, X. K., U. S. Bur. Mines, Bull. 406 (1937). Laury, N. A. (to American Cyanamid Co.), U. S. Patent 2,254,014 (Aug. 26, 1941). Patosz, T., and Rabek, T. J., Przemysl Chem., 14, 529 (1930). Roberts, I., IND. ENQ.CHEM.,ANAL.ED., 8, 365 (1936). Rose, H., Ann. Physik, 28, 120 (1833). Salley, D. J., J . Am. Chem. SOC.,61,834 (1939). Schmidt, Brit. Patent 249,474 (1926). Schulta-Sellack, C., Ber., 4 , 109 (1871). Stephens, H., J . Am. Chem. Soc., 52, 636 (1930). Traube, W., Ber., 46, 2513 (1913). U. S. Tariff Commission, Rept. 124, 4 (1937).
PRESENTED before t h e Division of Industrial a n d Engineering Chemistry a t t h e 102nd Meeting of t h e American Chemical Society, Atlantic City, N. J. Based upon a dissertation presented by A. H. Tenney in partial fulfillment of t h e requirements for t h e degree of doctor of philosophy in t h e Faculty of Pure Science, Columbia University.
Settling and Thickening of Aqueous Suspensions KARL KAMMERMEYER Drexel Institute of Technology. Philadelphia, Penna.
HE factors involved in the settling of suspensions can roughly be divided into two groups. One group will comprise the factors actually coming into play during the settling of the particles from the original suspension and thus primarily determines the rate of sedimentation and therefore the area of the settling tank. The second group will include the factors which characterize the ultimate condition of a settled sludge-that is, the sedimentation volume-and therefore will largely determine the height of the tank. The object of this paper is t o discuss the effects of some of the factors of the second group upon the sedimentation volumes of the aqueous suspensions of calcium carbonate, barium sulfate, silica, and sludges from alum manufacture. More specifically, these factors are settling tube diameter, initial
T
weight concentration, slow agitation during settling (e. g., thickening), and, in connection with the latter, the effect of stirrer height. Some of the more recent investigations on the settling behavior of suspensions are those of Work and Kohler (7) and Comings (2), dealing with the settling of sludges with slow stirring, and of Egolf and McCabe (3) and Ward and Kammermeyer (6) on the quiescent settling of suspensions. It is to be expected that the ultimate settling height of a suspension will be different for the two types of settling, and a comparison of the behavior under these different settling conditions will be made in a subsequent section. It was felt, however, that there are two factors having a general bearing on the problem: (a) the effect of the tube
