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T. K. WIEWIOROWSKI AND F. J. TOURO
ion and a chlorine atom is released. The overall picture is in fact closely similar to dye-photosensitized photolysis of silver halides. However, whereas in dye photosensitization it is often uncertain4 whether the photolysis results from electron or energy transfer,
Molten Sulfur Chemistry.1 I.
there is little doubt that in this case electron transfer occurs. Acknowledgments. The author wishes to thank Dr. P. P. Williams and Dr. R. 19. Golding for several helpful discussions concerning this work.
Chemical Equilibria in Pure Liquid Sulfur
by T. K. Wiewiorowski and F. J. Touro Freeport Sulphur Company, Belle Chasse, Louisiana
(Received May 16, 1966)
A theory of Lewis acid-base equilibria in pure liquid sulfur is advanced. Sulfur molecules in the form of short chains have a higher 7r-electron density and can behave as bases with respect to sulfur molecules in which the ratio of u bonds to sulfur atoms is equal to 1 ( o c b atomic sulfur rings and polymeric sulfur molecules). The theory rationalizes the experimentally established existence of short chain molecules in molten sulfur.
Nomenclature
Introduction
Sulfur is known to exist in a variety of molecular forms. Present symbolism of these forms is in many cases confusing and misleading. For instance, liquid sulfur just above the melting point is frequently referred to as Sh. Since sulfur at this temperature is considered by some workers to consist almost exclusively of octaatomic sulfur rings, there is a tendency to assume that Sh is synonymous with octaatomic sulfur ring molecules in the liquid state. I n view of the fact that the system consists of more than one molecular species, the use of a Greek letter is improper in this case. The IUPAC provides rules for naming the various molecular forms of sulfur, but no convenient unambiguous symbolism is available. The following symbols are therefore recommended : SsR, cyclooct* catenapolysulfur; SgCH, catenaoctasulfur; and SflCH, sulfur. Since branching of sulfur chains has not been observed, only two superscripts, R and CH, are required to differentiate between ring and chain sulfur molecules.
Chemical equilibria in pure molten sulfur have been the subject of extensive study. Several theoretical treatments of the sulfur system have been a d ~ a n c e d . ~ - ~ However, the theories proposed so far have certain shortcomings; each is capable of explaining some aspects of the sulfur system but fails to explain others. A theory of chemical equilibria in sulfur should be capable of explaining, or at least should be compatible with, the following experimentally observed properties of the element. 1. The melting point of monoclinic sulfur is higher than the freezing point of liquid ~ u l f u r . ~
The Journal of Physical Chemistry
(1) This is the first of a series of papers dealing with the chemistry of systems in which molten sulfur is the sole or major component. (2) (a) G. Gee, Trans. Faraday SOC.,48, 515 (1952); (b) F. Fairbrother, G. Gee, and G. T. Merrall, J . Polyner Sci., 16, 459 (1955). (3) R. E. Powell and H. Eyring, J . Am. Che-m. SOC.,65, 648 (1943). (4) A. V. Tobolsky and E. Eisenberg, ibid., 81, 780 (1959). (5) A. Smith, 2. Physik. Chem., 42, 469 (1903).
MOLTEN SULFUR CHEMISTRY
3529
2. A t least two molecular species, SsR and SgCH,are present, in molten sulfur just above its freezing point. These two species differ in their solubility in carbon disulfide and are found in sulfur quenched in CS2.6 The formation of SeCHis believed to be responsible for the freezing point depression of liquid sulfur. 3. Sulfur rapidly cooled from the liquid state contains a fraction of polymeric sulfur which does not dissolve readily in carbon disulfide. The size of this fraction depends on the temperature from which the sulfur is chilled, but even low-temperature liquid sulfur (115159") yields significant quantities of this fraction upon ~hilling.',~ These properties of sulfur cannot be rationalized in terms of presently known theories of chemical equilibria in pure liquid sulfur. One of these theories, developed by Gee,2*treats the system as consisting essentially of only two species, octaatomic sulfur rings, SsR, and polymeric chains SnCH. Neglect of the existence of octaatomic sulfur chains, SsCH,in liquid sulfur is perhaps the most serious weakness of Gee's theory. It is contrary to experimental observations and leads to the development of two distinct treatments, one applicable below and the other above the transition temperature (159"). A refinement of this theoryzb has several related shortcomings which are in fact recognized by the authors: (a) the theory does not explain the existence of the experimentally observed carbon disulfide-insoluble fraction in sulfur quenched from the liquid below 159"; (b) the theory is incompatible with heat capacity data just below this temperature. Another thermodynamic treatment of chemical equilibria in liquid sulfur was developed by Powell and E ~ r i n g . Their ~ treatment, however, also fails to account for the existence of appreciable concentrations of short sulfur chains near the freezing point of liquid sulfur. Tobolsky and Eisenberg4 developed a unified theory based on the assumption that the equilibrium of the system is governed by two types of reactions
Initiation:
R
s 6
+
+ +s s C H
SsR SBCH_T slaCH Propagation: These authors have derived a series of equations which mathematically describe the ring-chain equilibrium and are valid over the entire liquid range of sulfur. Their theory accounts very well for the pronounced viscosity change in the vicinity of 159" and explains the formation of polymeric sulfur above this temperature. However, this treatment also fails to explain the existence of considerable concentrations of SsCHin liquid sulfur. The concentration of this molecular species calculated from the theory is negligible compared to the experimentally determined values.
Discussion The discussion of chemical equilibria which is presented here provides a qualitative explanation of numerous facets in the behavior of elemental sulfur. The existence of octaatomic sulfur chains has been theorized by Schenk and Thummler,6 but is not accounted for in thermodynamic treatments of chemical equilibria in liquid sulfur. We are proposing that this phenomenon can be explained in terms of Lewis acid-base equilibria in this system. Sulfur molecules in the form of short chains can be regarded as the Lewis bases. Since the ratio of u-S-S bonds to sulfur atoms in these molecules is relatively low, they can have a higher electron density in their T clouds and behave as Lewis bases. In contrast, sulfur species in which the ratio of u-sulfursulfur bonds to sulfur atoms is equal to 1, that is octaatomic sulfur rings and polymeric sulfur chains, assume the role of Lewis acids in the system. Under these circumstances, the relatively high energy involved in breaking an S-S bond in an octaatomic sulfur ring is to a large extent compensated for by the solvation energy resulting from the interaction between the formed base and the acidic medium. Consequently, the formation of small amounts of complexed SBCHin molten sulfur becomes thermodynamically favorable, even at rei% tively low temperatures. A description of this system can be formulated in terms of the chemical equilibria
Ka
+ nSsR
S8CH'nS8R)r SsCH
(11)
An exact mathematical interpretation of these equilibria presents problems which are difficult to resolve unambiguously. Consequently, only a qualitative discussion is given here and no attempt is made to calculate the concentration of complexed octaatomic sulfur chains over the entire temperature range of liquid sulfur. The first equilibrium refers to the formation of complexed short sulfur chains. Presumably, this complex formation represents a significant contribution to the thermodynamic properties of liquid sulfur and is responsible for the freezing-point depression in pure sulfur. The second equilibrium describes the formation of free short chains which correspond to the species defined by Tobolsky and Eisenberg as being formed in the initiation reaction. In this sense, the theory pre(6) P. W. Schenk and U. Thammler, 2. EZektrochem., 6 3 , 1002 (1959). (7) A. Smith and W. B. Holmes, J. Am. Chem. Soc., 27, 979 (1905). ( 8 ) D. L. Hammick, W. Cousins, and E. Langford, J . Chem. Soc., 797 (1928).
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T. K. WIEWIOROWSKI AND F. J. TOURO
3530
sented in this paper is intended to supplement, but not to replace, the treatment developed by Tobolsky and Eisenberg. The value of n in these equilibria represents the average ratio of Lewis acid to Lewis base molecules in the complex. This ratio is temperature dependent and decreases with increasing temperature. If it is assumed that n = 2 at the freezing point of sulfur and that the concentration of complexed octaatomic sulfur chains is 5.S%,6 the free energy change involved in the formation of these species at the freezing point of sulfur can be estimated. A F 3 8 8 0 ~=
-RTr hl K1
=
1820 cal/mOk
Furthermore, if n is larger than 2, AF would be even smaller. This calculated value is, of course, considerably smaller than the free-energy change involved in the opening of a cyclooctasulfur molecule resulting in the formation of a free octaatomic sulfur chain (SgR e ~ ~ 3 8 8 0= K 23,900 cal/mole).4 Since the formation of short sulfur chains is an endothermic process, an increase in temperature will shift chemical equilibrium I toward the right. This implies that while near the melting point of sulfur, cyclooctasulfur predominates and the system behaves as a Lewis acid, as temperature increases ring molecules open to form short sulfur chains, and the basicity of the system gradually increases. The extent to which this process takes place cannot be quantitatively determined at the present time. It is believed, however, that this shift in the chemical equilibria is not very pronounced because formation of complexed chains can be expected to become increasingly difficult (endothermic) as the concentration of the Lewis acid (necessary to complex the base) decreases. This explanation of the thermodynamic feasibility of short-chain formation facilitates the interpretation of various aspects of sulfur chemistry. For example, the fact that sulfur molecules have a tendency to enter Lewis acid-base interactions suggests that the solubility of a compound in liquid sulfur would to a certain extent depend on its electron-donating capability. More specifically, since the octaatomic sulfur ring molecules (Lewis acid) predominate in the system, at least at relatively low temperatures, it should be expected that basicity of a compound would favor its solubility in liquid sulfur. There is ample experimental evidence available to verify this line of reasoning. Naphthalene (a Lewis base) is more soluble in liquid sulfur than either cis or trans-decalin,g aniline is more soluble than either benzene or toluene, which in turn are more soluble than cyclohexane, etc. Lewis acids are generally weakly soluble in liquid sulfur and, moreover, it should be expected that molecules of the dissolved Lewis acid would tend to associate with the chain species in liquid sulfur. The Journal of Physical Chemistry
Experimental verification for this conclusion can be found in the work of Smith and Holmes,’ who provide evidence that sulfur dioxide or sulfuric acid dissolved in liquid sulfur enhance the concentration of sulfur chains in the liquid. Lewis acid-base interactions should give rise to charge-transfer bands in the absorption spectrum of pure liquid sulfur. The exact location of these bands cannot be reliably assessed on the basis of presently available spectroscopic data. However, it is quite possible that the shift of the absorption edge of liquid sulfur toward longer wavelengths with increasing temperature’O is related at least partially to charge-transfer phenomena. The displacement of the absorption edge is dependent upon the rate of heating of the sample and consequently cannot be due exclusively to the thermal excitation of vibrational energy states of the sulfur molecules. This dependence upon the rate of sample heating can, however, be a result of the slow kinetics involved in the formation of short sulfur chains which enter the charge-transfer complexes in liquid sulfur. It is reasonable to assume that 3d orbitals are involved in this complex formation. I n fact, if the possibility of 3d orbital participation in sulfur bonding were excluded, the short sulfur chains would have to be paramagnetic. Electron spin resonance” and magnetic susceptibility measurements,12 however, indicate the absence of unpaired electrons in low-temperature liquid sulfur. It can be concluded, therefore, that an electron pair which forms a a-sulfur-sulfur bond in the ring molecule, upon bond breaking becomes delocalized and occupies a molecular orbital in the short sulfur chain. General aspects concerning the participation of d orbitals in sulfur bonding have been previously discussed.1 3 9 l 4 I n summary, the qualitative aspects of the theory we are proposing are restated. The presence of short chains in liquid sulfur is interpreted in terms of Lewis acid-base interactions which drastically reduce the net free energy change involved in the formation of the chain species. These interactions result from the higher n electron density in the short chains as com-
(9) R. L. Scott, “Elemental Sulfur,” B. hfeyer, Ed., Interscience Publishers, Inc., New York, N. Y., 1965, p 337. (10) A. M. Bass, J . Chem. Phy8., 21, 80 (1953). (11) D. M. Gardner and G. K. Fraenkel, J . Am. Chem. SOC.,78, 3279 (1956). (12) J. A. Poulis, C. H. Massen, and P. Leeden, Trans. Faraday SOC.,5 8 , 474 (1962). (13) C. C. Price and S. Oae, “Sulfur Bonding,” The Ronald Press Co., New York, N. Y., 1962. (14) M. Schmidt, “Elemental Sulfur,” B. Meyer, Ed., Interscience Publishers, Inc., New York, N. Y., 1965, p 301.
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MOLTENSULFURCHEMISTRY
pared to sulfur molecules in which the ratio of u bonds to sulfur atoms is equal to 1. The development of the concept of acid-base equilibria in liquid sulfur is extremely useful in the inter-
pretation of various previously unexplained aspects of sulfur chemistry. This will be demonstrated in the forthcoming series of publications dealing with the chemistry of molten sulfur.
Molten Sulfur Chemistry. 11. The Solubility of Sulfur Dioxide in Molten Sulfur
by F. J. Touro and T. K. Wiewiorowski Freeport Sulphur Company, Belle Chasse, Louisiana
(Received May 16, 1966)
The solubility of sulfur dioxide in molten sulfur has been determined. Infrared absorptivity coefficients of gaseous sulfur dioxide as well as dissolved sulfur dioxide have been obtained at 125 and 150”. An infrared cell designed and built to obtain vapor-liquid equilibrium data has been employed in this study. As a result of this investigation, a van’t Hoff equation relating the solubility of sulfur dioxide in molten sulfur to absolute temperature and the sulfur dioxide partial pressure has been derived. The heat of solution for dissolving sulfur dioxide in molten sulfur is -1850 cal/mole.
Introduction A previous investigation‘ reported on the “unusual” solubility behavior of hydrogen sulfide in liquid sulfur. The increase in the solubility of the gas with rising temperature was interpreted in terms of hydrogen polysulfide formation and correlated with the well-known viscosity reducing effect of hydrogen sulfide on sulfur.2 Bacon and Fanell? found that the presence of sulfur dioxide did not reduce the viscosity of pure sulfur. Consequently, it could be expected that sulfur dioxide should exhibit a “normal” solubility pattern in molten sulfur. The results of this study confirm this expectation.
used as reagents. Molten sulfur was saturated with sulfur dioxide a t 1 atm of sulfur dioxide overpressure for 2 hr. Nitrogen was then bubbled through the sulfur, and the effluent was passed through an iodine solution to collect the sulfur dioxide. Thiosulfate back-titration followed to complete the determination. Solubility data obtained by the chemical method were utilized for determining the infrared absorptivity coefficient for dissolved sulfur dioxide in molten sulfur. An infrared cell equipped with a small set of cartridge heaters, as previously de~cribed,~ was filled with molten sulfur and saturated with sulfur dioxide. A PerkinElmer 221G spectrophotometer scanned the molten
Experimental Section A chemical method was employed for determining the solubility of sulfur dioxide in molten sulfur a t temperatures from 121 to 140’. Sulfur purified by the method of Bacon and Fanelli4 and sulfur dioxide supplied in lecture bottles by the Matheson Co. were
(1) T.K. Wiewiorowski and F. J. Touro, J. Phys. Chem., 7 0 , 234 (1966). (2) F. J. Touro and T. K. Wiewiorowski, ibid., 70, 239 (1966). (3) R. F . Bacon and R. Fanelli, J . Am. Chem. Soc., 65, 639 (1943). (4) R. F. Bacon and R.Fanelli, Ind. Eng. Chem., 34, 1043 (1942). (5) T. K. Wiewiorowski, R. F. Matson, and C. T. Hodges, Anal. Chem., 37, 1080 (1965).
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Volume 70.Number 11
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