November 1950
INDUSTRIAL A N D E N G I N E E R I N G C H E M I S T R Y
(17) Davis, F. O., and Fettes, E. M., J . A n . C h a . Soc., 70, 2611 (1948). (18) Davis, F. O., and Patrick, J. C., unpublished experiments. (19) Dawson, I. M., Mathieson, A. M., and Robertson, J . M., J . C h a . Soc., 1948,322. (20)Daweon, I. M.. and Robertson, J. M., I W . , P. 1256. Donohue, J., and Schomaker, V., J : C h m . Phya., 16, 92 (1948). Duecker, W. W., Chem. Met. Eng., 41, 683 (1934). Farmer, E. H., and Shipley, F. W., J. Polymer Xn‘., 1,293 (1946).
Fettes, E. M., and Davis, F. O., unpublished experimenta. Flory, P. J., J. Am. Chem. Soe., 58, 1877 (1936). Fuller, C. S., Cha. Reus.,26, 143 (1940). Fueon, R. C., Price, C. C., Burness, D. M., Foster, R. E., Ratchard, W. R., and Lipscomb, R. D., J . Org. Chem., 11,487 (1946). Gee, G., Trans. Inst. Rubber Znd., 18, 266 (1943). Harragan, T. F., I n d k Rubber World, 95, 3 (1936). Jorcrak, J. S.,and Boswell, W. E., Ibid., 120,334 (1949). Kah, J. R., Tram. Far&y SOC.,32, 77 (1936). Koch,‘H. P., J. C h a . SOC.,1949, 394. Laurence. A., and Perrine, V. E., Rubber Age, 54, 139 (1943). Liska, J. W., IND.ENQ.CHEM.,36,40 (1944). Macallurn, A. D., J. Org. Chem., 1 3 , l M (1948). Macy, R., Jarman, 0.N., Morrison, A., and Reid, E. E., Scirnre, 106,355 (1947). Martin, 8. M., Jr., India Rubber Wor?d,99, 38 (1938); 104, 30 (1941). Martin, 8. M., Jr., and Laurence, A. E., IND.ENB.CHEM.,35, 986 (1943). Martin, S. M., Jr., and Patrick, J. C., IbiiE., 28, 1145 (1936).
Mason, J. P., and Manning, J. F., “The Technology of Plastics and Resins,” p. 388, New York, D. Van Nostrand Co., 1945. Mesrobian, R. B.. and Tobolsky, A. V., J. Polyner Sci., 2, 463 (1947).
Mochulsky. M., and Tobolksy, A. V., IND. ENQ.CHEM.,40,2155 (1948).
Moll, H. W., and Le Fevre, W. J., I W . , p. 2172. Morris, R. E., James, R. R., and Evans, E. R., India Rubber World, 110,529 (1944).
Morris, R. E., James, R. R., and Werkenthin, T. A,, IND.ENQ. CHN.,
35. 864 (1943).
2223
Mullins, L., Trans. Inst. Rubber Ind., 21,247 (1945). Patrick, J. C., Tram. Fara&y Soc., 32, 347 (1938). Patrick, J. C., U. 8. Patent 1,950,744 (March 13, 1934). Patrick, J. C., and Ferguson, H. R. (to Thiokol Corp.), Ibid., 2,466,963 (April 12, 1949). (50) Powers, P. O., “Synthetic Resins and Rubber,” p. 216, New York, John Wiley & Sons, 1943. (51) Powers, P. O., and Billrneyer. B. R., IND. ENG.CHEM.,37, 64 (46) (47) (48) (49)
(1945). (52) Sager, T. P., IND.ENQ.CHEM.,29, 747 (1937). (53) Sarbach, D. V., and Garvey, B. S., India Rubber World, 115, 798 (1947). (54) Sdker, M. L., Winspear, G. G., and Kemp., A. R., IND.ENQ. CHEM.,34,157 (1942). (85) Spielberger, O., Chem.-Ztg., 63, 29 (1938). (56) Spielberger,U.,Kautechuk, 13, 137 (1937). (57) Stern, M. D., and Tobolsky,A. V., J. Chem. Phys., 14,93 (1946). (58) Tertian, Robert, Reu. 01%. caoutchouc., 23, 245 (1936). (59) Tobolsky, A. V., and Andrews, R. D., J . Chem. Phgs., 13, 3 (1945). (60) Tobolsky, A. V., Leonard, F., and Roeser, G. P., J . Polynier Sci., 3, 604 (1948). (61) Tobolsky, A. V., Prettyman, I. B., and Dillon, J. H., J. Applied Phys., 15,380 (1944). (62) Trillat, J. J., and Tertian, R., C m p t . Tend., 219, 395 (1944). (63) Wakeman, R. L., “Chemistry of Commercial Plastics,” p. 600, New York, Reinhold Publishing Corp., 1947. (64)Werkenthin, T. A., Richardson, D., Thornley, R. F., and Morris, R. E., Rubber Age, 50, 103 (1941). (65) Westlake, H. E., and Dougherty, G., J. Am. C‘hsm. SOC.,64, 149 (1942). (66) Westlake, H. E., Mayberry, H. G., Whitlock, M. H., West, ,J, R., and Haddad, G. J., Ibid., 68, 748 (1946). (67) Wood, L. A., India Rubber World, 102, 33 (1940). (68) Yost, D. M., and Russell, H., “Systematic InorganioChemistry,” p. 162, New York, Prentice-Hall, Inc., 1944. (69) Zelinski, N. D., Denisenko. Y. I., Euentova, M. S.. and Khromov, S. I., J. Rubbcr I n d . (U.S.S.R.),11, 111 (1934). RBCETVED Mmch 27, 1950.
Inorganic Compounds Containing Sulfur and Fluorine HENRY C. MILLER AND J. F. GALL Research and Development Divinion, Pennsylvania Salt Manufacturing Company, Wyndmoor, Pa. Although few inorganic compounda containing oulfur and fluorine have found practical application, and in no case have they found large scale uoe, they have a wide range of properties and unusual combinationa of characteristics. The preparation, physical characteristics, and chemical reactivity of inorganic sulfur-fluorine compounds are reviewed here,
v
ERY few of the inorganic compounds containing sulfur and fluorine have found practical application, and in no case have these compounds found any very large scale use. Yet the sulfur-fluorine compounds show a great range of properties, and some have unusual combinations of characteristics. This article reviews the preparation, the physical characteristics, and the chemical reactivity of the inorganic sulfur-fluorine compounds. Although no valence states other than -1 have been clearly demonstrated for fluorine, the number of inorganic compounds containiig both sulfur and fluorine is large. This is the result of the possession of several valence states by sulfur, and of the strong tendency of fluorine, through its small atomic radius ana extreme electronegativity, to stabilize many moleaular structurea.
It is convenient to discuss the sulfur-fluorine conipounds in three groups: (1) compoun.ds in which sulfur is directly bonded to fluorine; (2) compounds in which sulfur is bonded to fluorine through a third element; and (3) compounds in which sulfur and fluorine are remote.
COMPOUNDS WITH SULFUR DIRECTLY BONDED TO FLUORIDE The interatomic distance between sulfur and fluorine has been measured for sulfur hexafluoride and sulfuryl fluoride. In both rases, the observed distance is 1.56 A. (88,8’). This is considerably less than the distance which may be calculated from the table of covalent atomic radii, according to which the sulfurfluorine distance for a covalent-type bond would be 1.68 A. (M).This may be regarded as confirmation of the expectation that the sulfur-fluorine bond would be of an ionic type, suggested by the extreme electronegativity of fluorine as compared with sulfur. It has been reported, however, that thionyl fluoride, SOFa, has a pyramidal structure, implying the preservation of the directional character of the valence bonds on sulfur and there fore at least a partial retention of covalent bonding (87). It has also been postulated that in the cme of sulfuryl fluoride the
INDUSTRIAL A N D ENGINEERING CHEMISTRY
2224
bond shortening results from manifold resonating structures imparting partial multiple bond characteristics (29). BINARY COMPOUNDS
As with chlorine, sulfur forms a series of binary compounds with fluorine. The group of sulfur fluorides, however, unlike the corresponding chlorides, contains compounds in which sulfur exhibits its maximum valence. Furthermore, these sulfur fluorides of higher valency are remarkably inert and unreactive. Thus all the chlorides are rapidly hydrolyzed by water, whereas sulfur hexafluoride is resistant even to hot aqueous alkalies. The physical properties of the fluorides differ from those of the chlorides in the anticipated direction-for example, the fluorides characteristically possess much higher vapor pressures than the corresponding chlorides. A list of all the known sulfur fluorides with some of their properties is given in Table I.
Table I. Sulfur Fluorides Name Sulfur monofluoride Sulfur difluonde sulfur tetraflu0rid.e Disulfur deoafluonde Sulfur hexafluoride
Formula &FI SFa SF^ SXFIO SF4
Color Colorless Colorlesa Colorless Colorless Colorless
Meltip Point, C.
- 105 5 -i24 -92
-50.8 at 16 Ib./sq. inch gage
Boiling Point, C.
-99
-.4d 29
Sublimation -63 8
sulfur Monofluoride, SzF,, was first identified in 1923 by Centnerszwer and Strenk ( 9 ) who piepared ii by heating a mixture of silver fluoride and sulfur. I t is a colorless gas with a disagreeable odor. The gas is thermally unstable. When exposed to moisture, it reacts immediately, disproportionating to sulfur dioxide and sulfur a8 well as hydrogen fluoride. It is rapidly absorbed by alkalies and by oxidizing solutions such as potassium permanganate. The gas attacks glass and other silicate materials at ordinary temperatures, forming sulfur dioxide thionyl fluoride, and silicon tetrafluoride. The dry gas does not attack steel, tin, or platinum. I t is reported to be toxic. Sulfur Difluoride, SF2, has been sought because of its formal analogy with the corresponding chloride of sulfur (22), but it has probably never been prepared in a pure state. Its formation by thermal decomposition of the monofluoride has been indicated by several investigators; it is evidently thermally unstable and highly reactive. Sulfur Tetrafluoride, SF,. Fischer and Jaenckner (19) in 1929 prepared sulfur tetrafluoride by reacting cobalt, trifluoride with sulfur in a quartz apparatus. Later work has shown Ohat this reaction produces a mixture of sulfur fluorides (16). The monoand Muorides of sulfur can be removed from the mixture by shaking with mercury, while the tetrafluoride may be separated from sulfur hexafluoride by fractional distillation. This gas is less reactive than sulfur monofluoride. It does not attack glass, but it is rapidly absorbed by caustic alkali solutions forming, among, other products, the sulfite and fluoride of the alkali metal. The vapor pressure of sulfur tetrafluoride has - 1132 7.746, corresponding to a been given as loglo pmm.= normal boiling point of about -40' C. Its melting point has been reported as 124' C. Sulfur Hexafluoride, SFB. Because of its unusual character as a stable gm of high molecular weight, and its useful electrical characteristics, sulfur hexafluoride has been the moat studied of all the sulfur fluorides. It is formed by direct combination of sulfur with elemental fluorine. From the mixture of sulfur fluorides produced by this combustion, the sulfur hexafluoride is purified by taking advantage of the lower stability or higher reactivity of the other fluorides. The gas mixture is first scrubbed with aqueous alkali (preferably potassium hydroxide because of tho
+
-
Vol. 42, No. 11
high solubility of potassium fluoride) to remove sulfur monofluoride and sulfur tetrafluoride. The gas is then passed through a heated zone which decomposes disulfur decafluoride to sulfur hexafluoride and sulfur tetrafluoride; and the latter impurity is removed by a second alkali scrubbing. Water may be removed from the gas by strong sulfuric acid or by the usual solid-type desiccants. At atmospheric temperature and pressure sulfur hexafluoride is a very heavy, colorless, odorless, nontoxic, and chemically inert gas. Its molecular weight of 146.1 gives it a density approximately five times that of air under equal conditions. Its degree of chemical reactivity can be compared with that of nitrogen; it is noncorrosive and nonreactive at atmospheric and moderately elevated temperatures. It will show characteristic reactions with some of the more active chemical Rubstances when strongly heated; thus it reacts with strong reducing agents such as hydrogen and some metals at red heat, forming the corresponding fluorides and sulfides. Glass and silicates are attacked by sulfur hexafluoride only at temperatures above red heat. Many of the physical properties of sulfur hexafluoride may be conveniently compared with those of carbon dioxide; both gases have no liquid range a t atmospheric temperature but condense at low temperature directly into the solid form. The vapor pressure curves are closely parallel. Under pressure, both compounds can be retained in the liquid form up to their critical temperatures, which lie a short distance above atmospheric temperatures. Sulfur hexafluoride is very sparingly soluble in water. Sulfur hexafluoride is characterized also by extremely high dielectric strength. This in combination with its other properties makes it unusually effective as a gaseous insulator. It is superior to the chlorofluorohydrocarbons, which, while very good insulating gases, have higher dew points at a given pressure The gas has found useful application in high voltage x-ray transformers and in high voltage generators, particularly of the electrostatic type (22). Disulfur Decafluoride, SzFlv. In 1934 Denbigh and WhytlawGray (IO), examining the mixture of gases resulting from the reaction between sulfur and elemental fluorine, discovered a high boiling sulfur fluoride which they proved to be disulfur decafluoride. This compound, upon structural examination, appears to be a homolog of sulfur hexafluoride in the same sense that ethane is a homolog of methane; thus it possesses a sulfur-tw sulfur bond. The compound is a colorless and nearly odorless liquid boiling at 29" C. and freezing at -92' C. The liquid has a specific gravity of 2.08. The compound is considerably more stable than the lower fluorides described previously. Thus, it is not immediately hydrolyzed by water or by dilute alkali solutions. It is decomposed by fused caustic or by heating to about 300" C. in the presence of metals. It does not attack glass or metals at low temperatures, but reacts with heated copper and iron, forming the metal sulfides and fluorides. Disulfur decafluoride is a toxic gas. SULFUR OXYHALIDES A N D DERIVATIVES
A rather large number of compounds contain sulfur directly linked to fluorine and oxygen. Type examples of oxyfluorides containing sulfur in the tetra and hexavalent states are, respectively, thionyl fluoride, SOFz, and sulfuryl fluoride, SOaF2. The latter compound is a mixed anhydride of fluosulfonic acid, HSOsF, and hydrogen fluoride. No fluosulfinic acid, corresponding to thionyl fluoride, is known. Some properties of the thionyl and sulfuryl fluorides and mixed halides are shown in Table 11. Thionyl Fluoride, 809, waa first prepared by Meslans (18) in 1890, by reaction of thionyl chloride with zina fluoride. The compound is better obtained by reaction of thionyl chloride with arsenia trifluoride or with antimony trifluoride, using antimony
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INDUSTRIAL AND ENGINEERING CHEMISTRY
Table 11. Thionyl and Sulfuryl Fluorides Name Thionyl fluoride Thionyl ahlorofluoride Sulfuryl fluoride Sulfuryl chlorofluoride Sulfuryl bromofluoride Thionvl tetrafluoride Pyrosulfuryl fluoride
Formula SOF, SOClF SOiFi SOd3F SOaBrF SOFI SzOaF:
Color Colorless Colorless Colorless Colorless
... ... ...
Melting Point, C.
Boili~p Point, C.
-129.5
-43.8 12;2 52
-139.6
- 120
-124.7 86
-- 107 ...
-
t7.1
+-48.6 40 ...
pentachloride as a catalyst (6). Thionyl fluoride is a relatively stable compound which hydrolyzes only slowly in water, and does not attack glass below about 400' C. I t is rapidly attacked by dilute alkali solutions. It reacts with sodium a t elevated temperatures. With ammonia it forms thionyl hemipentamminofluoride, 2SOFa.5NHs, then the hemiheptammino compound, 2SOFz.7NHI, and finally ammonium fluoride and thionyl amide, SO(NH& Thionyl Chlorofluoride, SOClF, was prepared by Booth and Maricola in 1940 (6), by fluorinating thionyl chloride with antimony trifluoride, using antimony pentachloride as a catalyst It is a colorless gas a t atmospheric temperature, and is more reactive than thionyl fluoride but less reactive than thionyl chloride. Thus thionyl chlorofluoride will react with mercury and, like the other thionyl compounds, is hydrolyzed slowly by moist air and readily in water. Sulfuryl Fluoride, S02Fz, was first described in 1901 (f7) by Moissan, who prepared it by reaction of fluorine with sulfur dioxide in the presence of a platinum catalyst. Subsequent work indicated that platinum is not necessary if a sufficiently high reaction temperature is maintained. This compound can also be prepared in fair yields by decomposing barium fluosulfonate, Ba(SOaF)%(S4). Sulfuryl fluoride is a very stable compound not decomposed by hot water. It does not attack glass even at high temperatures, and molten sodium has no effect upon it. It is soluble in liquids such as chloroform and alcohol. Like sulfur hexafluoride it has a high dielectric strength, but is somewhat less stable than that gas at high temperatures and under corona discharge. Sulfuryl Chlorofluoride, SOnClF, was first described by Davies and Dick in 1932 and was more carefully characterized by Booth and Herrmann in 1936 (6). The latter authors prepared it in good yields from sulfuryl chloride and antimony trifluoride in the presence of antimony pentachloride as a catalyst. This compound is intermediate in stability and in reactivity between the chloride and the fluoride. It hydrolyzes slowly in water, and very rapidly in dilute sodium hydroxide solution. It does not attack dry glass, mercury, or common metals a t atmospheric temperatures. Because of its susceptibility to hydrolysis it may be toxic. Sulfuryl Bromofluoride, SOsBrF, is reportedly obtained from the reaction of sulfur dioxide with a mixture of bromine trifluoride and bromine (8). No further information has been published. Thionyl Tetrafluoride, SOF4, has been reported, but details of its preparation and properties do not appear to have been published. It L said to be an excellent fluorinating agent for the preparation of organic fluorine compounds (8). Fluosulfonic Acid and Its Salts. Fluosulfonic acid, HSOsF, may be regarded as an intermediate mixed anhydride of hydrofluoric acid and sulfuric acid, with sulfuryl fluoride as the total anhydride of these two acids. A convenient laboratory preparation of fluosulfonic acid is the distillation of a mixture of potassium acid fluoride and oleum. This distillation can be carried out in glass apparatus. Fluosulfonic acid is made commercially by reaction of sulfur trioxide distilled from oleum and hydrogen fluoride in an iron column. Pure fluosulfonic acid is a colorless, somewhat viscous liquid. As might be expected, it is rapidly hydrolyzed by water, although
2225
less quickly than the corresponding chloro compound. This ready hydrolysis results in a tendency to fume strongly in air. The melting point of fluosulfonic acid is -87.3' and its boiling point is 162.6' C. In the absence of reactive substances it is reported to be remarkably stable up to 900' C. (SO). With a limited quantity of water the hydrolysis of fluosulfonic acid is incomplete, the equilibrium HSOsF H,O*HsSO, HF being set up. The increasing interest in fluosulfonic acid, revealed by the technical and patent literature, is suggestive of a wide range of potential utility in organic and inorganic chemistry. Thus fluosulfonic acid is proving to be a versatile catalyst and reagent for the production of increasingly important organic fluorine compounds. As a tool in preparative chemistry, it is similar to chlorosulfonic acid but is generally more stable. It has been used in several alkylation processes, in the preparation of organic fluorine compounds, and in the electropolishing of certain metals.
+
+
Table 111. Salts of Fluosulfonic Acid Salt NHdSOrF LiSOaF FiSOrF~3HzO NaSOrF KSOaF RbSOrF CsSOrF
Appearance Colorless needles White powder Needles Leaflets White prisms Colorless needles Colorless rhombio
Melting Point, C. 245 360 60-61
iii
304 292
Some of the salts of fluosulfonic acid are listed in Table 111. Several methods are reported for the preparation of these salts, and in most cases the yields of the pure compounds are rather low. The five methods most commonly mentioned are:
++ + + +
++
HSOsF M F +MSOsF H F M F +MSOIF HCI HSOaCl NH8OiF MOH +MSOiF NHiOH MSOiF KzS04 KiSzOl M F M F SO3 ----f MSOsF
+ +
The salts of fluosulfonic acid are relatively stable. They can be recrystallized from neutral solution, and the aqueous solutions decompose only slowly even on boiling. (The salta of this acid are usually very soluble in water. Thus the barium, lead, copper, and silver salts are soluble. The only highly insoluble fluosulfonate is formed with the organic base nitron.)
COMPOUNDS WITH SULIFUR AND FLUORINE INDIRECTLY CONNECTED In compounds of this type, exemplified by the phosphorus thiofluorides, the reactivity of sulfur is often enhanced by the presence of fluorine attached to the common element. Furthermore, the activity frequently increases with the number of fluorine atoms in the molecule, and the reactivity of the fluorine compound is sometimes greater than that of the corresponding chloro compound. This may be regarded as an effect of the electronegativity of fluorine acting inductively through the cvnmon element. PHOSPHORUS THIOFLUORIDES
The only known compounds containing fluorine and sulfur bonded directly to phosphorus contain phosphorus in the pentavalent state and sulfur in the divalent state. A similar series of compounds with chlorine and bromine is well known. The phosphorus thiofluorides range in activity from most active and spontaneously inflammable substances to very inert compounds. Some of these compounds have unusual solvent characteristics, resembling nonpolar compounds in this respect. Some of the properties of the phosphorus thiofluorides are listed in Table IV.
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Vol. 42, No. 11
T H I O C A R B O N n FLUORIDES
Table IV. Phosphorus Thiofluoridea Name Phosphorua thiofluoride
Formula PSFa
Phoephorus monoohlorothiofluoride Phos horus diokorothiofluoride Phosphorus monobromothiofluoride Phosphorus dibromothiofluoride Phosphorus diamidothiofluoride
PSClFz
Color Colorless, high refractive indev
Melting Point, C. -148 8 2
6 3
-96 0
64.7
-155
PSCl2F PSBr2Fn
Boiling Point, ' C. -52 3
-136
35.5
9
PSBrd.'
...
-75.2
125.3
P6( NH2)xF
...
...
...
The existence of thiocarbonyl fluorides has not been definitely established. It has been reported that Ruff prepared thiocarbonyl fluorides by reacting carbon disulfide with iodine pentafluoride and that this compound forms a dimer and a trimer (8). The melting and boiling points of the reported thiocarbonyl fluorides are given in Table V. The carbonyl fluorides are known to have fluorinating power and it is possible that this property is carried over to thiocarbonyl fluoride. SILICION THIOFLUORIDES
These compounds have not been reported, although the corresponding thiochlorides are known (1.9). It is probable that adequate study in this field has not yet been made, COMPOUNDS CONTAINING FLUORINE AND SULFUR REMOTELY CONNECTED
Phosphorus Thiofluoride, PSFa, is formed by reaction of phosphorus pentasulfide and lead difluoride at 200' C. in an atmosphere of nitrogen (14). The compound is a gas which is spontaneously flammable in moist air. It is slowly hydrolyzed by water in the following steps: PSFa +HSPOFz e HzSPOtF +HISPO,
&PO4
In an alkaline solution the end product is sodium thiophosphate, N&SPOs, while in strong acid solution the final product is phosphoric acid. On treating a solution of phosphorus thiofluoride with nitron acetate, the insoluble nitron salt of HSPOFn is formed. This is the only intermediate acid of this hydrolysis which has been separated. The reaction with moist air can be represented as lOPSF2 1502 ----f 6PFs 2PzOs 1OSO2. Phosphorus thiofluoride reacts with metallic sodium, forming a substance that reacts with water, giving off spontaneously flammable phosphene. Phosphorus thiofluoride has no action on mercury. It reacts with ammonia to form thiophosphoryldiamido fluoride, P(NH&SF, a white solid that hydrolyzes in moist air, forming hydrogen fluoride and thiophosphoryldiamido hydroxide [PS(NH& ]OH. Phosphorus Thiochlorofluorides, PSCIF2 and PSCIY, can be made by the fluorination of phosphorus thiochloride, using antimony trifluoride as the fluorinating agent in the presence of antimony pentachloride (4). The compounds have properties resembling both phosphorus thiofluoride and phosphorus thiochloride, showing an increase in reactivity as the fluorine content increases. They do not attac,k mercury but react with copper. They react completely but slowly with potassium hydroxide solution. They hydrolyze readily in moist air (distinction from phosphorus thiochloride). In contrast to phosphorus thiofluoride, the chlorofluorides are not spontaneously flammable. Phosphorus Thiobromofluorides, PSBrFz and PSBrzF, are formed by fluorinating phosphorus thiobromide using antimony trifluoride as the fluorinating agent (7). Pentavalent antimony is not required in this reaction. The two bromofluorides can be separated by distillation. The dibromomonofluoride, PSFBre, does not attack lead, tin, zinc, or sulfur. 1t.begins to react with mercury at 70" C. I t does not react with water or cold alkali solution. At 100" C. it react-9 dowly with hot alkali solution. It fumes slightly in air and has a peculiar but not sharp odor. It is miscible with acetone and carbon tetrachloride and dissolves some paraffin hydrocarbons. The monobromodifluoro compound is somewhat more reactive than the monofluoride, combining slowly with cold alkali and vigorously with hot alkali solutions, and attacks mercury at 35" C. No other mixed phosphorus thiohalides have been reported, although the chlorobromothiofluorides may exist. Corresponding compounds with iodine are probably not stable.
+
+
+
There are many compounds reported containing sulfur and fluorine in which these elements are not directly connected or are connected through a third element. Examples are the many addition products of boron trifluoride with sulfur compounds, and a number of double salts of sulfates and fluorides. As may be anticipated, these compounds have the expected combination of properties of the sulfur-containing and the fluorine-containing moieties. Phosphorus Difluoride Isothiocyanate, PFz(CNS), was reported by Anderson in 1947 as being formed when phosphorus triisothiocyanate reacts with antimony trifluoride (I). It is a liquid boiling a t 90.3' C. and freezing at -95' C. It is a rather stable compound and has unlimited solubility in carbon disulfide. BORON TRIFLUORIDES AND SULFATES
A number of addition products of boron trifluorides with metal sulfates have been formed by heating the sulfate in the presence of boron trifluoride (3). Examples are KzSO~.BFS,CszSOa.2BFa, NazS04.BFs,and TlZSO4.BF3. These compounds are generally stable in the solid state, but are readily decomposed by water, forming a solution of the metal sulfate and mixed hydrosyfluoboric acids.
Table V. Thiocarbonyl Fluorides Formula CSFP CnSnF4
CsSnFs
Melt i,ng Point, C.
Boiling Point, C.
-- 133 120
-26
--4033
- 136
A compound between calcium sulfate and boron trifluoride has been prepared and used as a means of preparing pure boron trifluoride by subsequent thermal decomposition. A compound related to the boron trifluoride complexes is made by direct addition of sulfur trioxide to potassium fluoborate (8). It has a formula KBF4.SOaand a melting point of 45' C. It is a colorless crystalline substance reacting vigorously with water to form pyrosulfate. It decomposes at 80" C. and prolonged heating at 120' to 130' C. yields potassium fluosulfonate. Boron Trifluoride-Hydrogen Sulfide Complex. Sabatier has shown the existence at low temperature of two compounds with the formulas, BFs.HaS, melting a t 137" C., and BFa.7H& having a transition point at -99" C. (21). These compounds are undoubtedly highly dissociated a t atmospheric temperatures.
-
SULFATE AND FLUORIDES
A number of double salts are known, such as the mineral schairerite, NatSO,.NaF, which is slowly but completely soluble
November 1950
INDUSTRIAL AND ENGINEERING CHEMISTRY
in water and forms colorless transparent easily fusiblc crystals. It is found in Searles Lake, Calif. A related compound is sulphohalite, 3NsnSO~.NaCINaF, occurring as pale greenish-yellow octahedrons which are very easily fusible, and have a refractive index of 1.455. This mineral is found in borax or Searles Lake. Double salts of aluminum sulfate and aluminum fluoride have een shown ,to exist, including A12(SO4)a.4AIF~.12HaOand Alp?h.AlF*.15H20 & (II,I8). One of these compounds, often referred to as aluminum fluorosulfate (AlF2)2S0,.12H20, is formed when calcium fluoride and aluminum sulfate are heated in boiling water. By rapid evaporation of the solution it forms a sirupy concentrate, and by slow evaporation a crystalline powder can be obtained. The compound is soluble in water, but when dried a t 300” C. it changes to an insoluble form. A solution of aluminum fluosulfate will react with hydrofluoric acid and sodium sulfate, forming chiolite or cryolite depending on the amount of sodium used. The compound K2BeF4. Ah( SOi)a,24Hz0 has been reported (12).
Arsenic trilluoride sulfur tetrachloride, 2AsFa.SCl4, was p r e pared by Ruff (IO), It forms yellow crystals which attack glass slowly and decomposes carbon tetrachloride, carbon disulfide, alcohol, ether, benzene, and petroleum ether. It is probably decomposed by water into sulfurous and thiosulfuric acids. Potaasium difluodithionate, &&Od?~.3Hpo, rubidium difluodithionate, Rb1S206F2.3H~0,and cesium hydroxyfluoditbionate, Cs&Os(OH)F.HeO, are formed by adding hydrofluoric acid to the saturated solution of the dithionate, and cooling (368).These salts are unstable, and when exposed to air quickly decompose with the evolution of water and hydrofluoric acid, leaving a residue of the dithionpte. When heated they give off water, hydrogen fluoride, and sulfur dioxide, leaving a residue of potassium sulfate.
(1) (2) (3) (4)
2227
LITERATURE CITED Anderson, H. H., J. Am. Chem. Soc., 69, 2496-7 (1947). Baumgarten, P., Ber., 73B, 1397-8 (1940). Baumgarten, P., and Hennig, H., Ibid,, 72B, 1743 (1939). Booth, H. S., and Cassidy, M. C., J.Am. Chem. Soc., 62,2369-72 (1940).
Booth, H. S., and Herrmann, C. V., Zbid., 58, 63 (1936). Booth, H. S., and Mericola, P.C., Zbid., 62, 640-2 (1940). Booth, H. S., and Seabright, C. A., Zbid., 65, 1834-6 (1943). British Intelligence Objectives Sub-committee, BIOS Final Rept. 1595, Item No. 22. (9) Centnersawer, M., and Strenk, C., Bw., 56B, 2249 (1923);
(5) (6) (7) (8)
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SULFUR IN FUNGICIDES M. M. BALDWIN Battelle Memorial Institute, Columbus, Ohio Although only a small percentage of the sulfur produced is used in the manufacture of agricultural fungicides, this element holds a key position in the fungicide industry. Varioum types of sulfur compounds are considered a0 fungicides. Besides being the most extensively used fungicide in ita elemental form, sulfur also is an active aonstituent of the dithiocarbamates and playa an indirect
role, through copper sulfate, in the preparation of copper fungicides. Sulfur and ita compounds are used to control dinewes of fruita, vegetables, and grasses. The mechanism of the fungicidal action of sulfur is discusaed brieny. Each of the various mechanisms proposed by different investigators may account for the toxicity of sulfur and its compeunds under diflerent operative conditions to fungi.
T
Comparisons are made in Table I from the standpoint of the sulfur industry. These figures show that only a relatively small percentage of the total amount of sulfur produced finds its way into fungicidal uses. This is not too surprising in view of the fact that almost three fourths of the sulfur in this country is converted to sulfuric acid, one of the most fundamental materials in the chemical industry. Even if only the nonacid uses of sulfur are considered aa a basis for comparison, the utilization of sulfur in fungicides amounts to only about 9% of that total. Nevertheless, around 81,500 long tons are an appreciable quantity of material. On the other hand, from the standpoint of the fungicide industry, sulfur overshadows other basic materials in this field. Estimated comparative figures are shown in Table 11. Not only is more elemental sulfur used than any other type of agricultural
HIS paper presents a summary of the position of sulfur and its compounds in the field of agricultural fungicides. Elemental sulfur is one of the oldest fungicides known, and sulfur compounds are among the most recently developed toxiaants to fungi. It can be said that sulfur is one of the fundamental elements of fungicidal materials. An indication of its prominent position during the development of fungicides is given by Horsfali in the frequent references to sulfur-containing fungicides in his list of landmarks in fungicide history (f7).
RELATIVE POSITION OF SULFUR F”0ICIDES It is difficult to obtain any very firm or clear-cut statistica on the production of fungicides. However, estimates can be made that give indications of the current position of sulfur fungicidw relative to the sulfur industry and to the fungicide industry.
