WEAK COMPLEXES OF SULFUR COMPOUNDS WITH HALIDE LIGANDS Science and Technology, Special Review (2000) 31-53 © 2000 Sultan Qaboos University 31 Weak Complexes of Sulfur Compounds with Halide Ligands Salama B. Salama* and Saad Wasif** *Department of Chemistry, College of Science, Sultan Qaboos University, P. O. Box 36, Al Khod 123, Muscat, Sultanate of Oman. **University of Toronto, Canada دراسة للمعقدات بين مركبات الكبريت وأيونات الهالوجين سالمة بطرس سالمة و سعد واصف مع أيونات SOBr2, SO2Cl2, SOCl2, SO2ثالثين عاماً الماضية تمت دراسة تفاعالت مركبات الخالل في : خالصة SO2Cl2-X فأمكن الكشف عن وجود المعقدات -SCN-, I-, Br -, Clالهالوجين -, SOBr2-X -, SOCl2-X -, SO2-X - Hfوجميعها معقدات يستحيل فصلها من المحاليل إلنخفاض طاقة تكونها -X-= Cl-, Br-, I-, SCNحيث نجد أن o ∆ ولذا كان االعتماد على إثبات وجودها من طيف االمتصاص لكل منها في المدى kJ/mol 6.03- ,22.51–في المدى ويعتمد ثبات المعقدات آنفة الذكر على طبيعة المذيب المستخدم للدراسة، ورقم االستقبال 380nm – 290nmالطيفي لمركب الكبريت وعلى قدرة أيونات الهالوجين على المنح االلكتروني للكبريت وكذلك على درجة حرارة التفاعل موضع -:ما نجد في التفاعلالدراسة ولقد أمكن دراسة تفاعالت إحالل أيونات الهالوجين مثل SO2-I - + Cl- = SO2-Cl - + I- وقد تمت دراسة تفاعالت إحالل لجميع ) درست بواسطة أشعة الليزر ( 1010s-1وتحدث بسرعة عالية تصل إلى المدى .خرالمعقدات آنفة الذكر ويعتمد مدى اإلحالل األيوني على المذيب المستخدم وطبيعة األيون الذي يحل محل األيون اآل SO2-Xوتنتمي مركبات في وسط الهرم وتعتمد زاوية S في قمة الهرم وذرة -X هندسياً إلى العائلة الهرمية حيث تقع - . -Xالهرم على طبيعة األيون ABSTRACT: This review summarizes the study of the formation of SO2-X -, SOCl2-X -, SOBr2-X -, SO2Cl2-X - complexes (X- = Cl-, Br-, I-, and SCN-) in pure and mixed solvents of acetonitrile and dimethylsulfoxide over the past 30 years. Spectrometry (290nm-380nm) was the principal technique used for the investigation, since the enthalpies of formation ∆Hf o of the above complexes are low ( -6.03  -22.51 kJ/mol ). The stability of the complexes depend on the nature of the solvent, acceptor properties of the sulfur compounds, the donor properties of the halide ions and the temperature of the reaction. Also, it was found that the extent of the iodide ion replacement reactions by the other halide depends on the nature of the solvent and the halide ions. SO2-I - + X- → SO2-X - + I- The structures of SO2-X - compounds are pyramidal, with X- at the top of the pyramid, the S atom at the center, and the angle depends on the nature of the halide ion. SALAMA and WASIF CONTENTS 1. Introduction 32 2. Detection, Stoichiometry and Stability of SO2-X -, SOCl2-X -, and SO2Cl2-X - Complex Species 33 2.1 Detection of Sulfur Complex Species 33 2.2 Stoichiometry and Complex Species 34 2.3 Stability Constants of Complex Species 34 2.4 Thermodynamic Constants of SO2-X –, SOCl2-X – and SO2Cl2-X – 34 2.5 The Donor-Acceptor Nature of the Complex Species 36 2.6 Correlation of the Order of Stability Constants of Different Complexes with Donor-Acceptor Properties 37 3. Solvent Effects on Stability of SO2-X –, SOCl2-X – and SO2Cl2-X – 38 3.1 Pure Solvents 38 3.2 Mixed Solvents: (MeCN-dmso) 39 3.3 The SO2-X – Species 40 3.4 The SOCl2-X – and SO2Cl2-X – Series 40 3.5 Vertical Correlations in Table 9 41 3.5.1 The Iodide Complex Series 41 3.5.2 The Bromide and Chloride Complex Series 41 3.6 Evidence for Solvent-Solvent Interactions 41 3.6.1 Spectroscopic 41 3.6.2 Vapor-Pressure, Viscosity, and Excess Functions from Refractive Index, Dielectric Constant and Volume 42 3.7 A Thermodynamic View of Solvent Effects on The Stability of SO2-X -, SOCl2-X - and SO2Cl2-X - 44 3.7.1 The Significance of ∆Gf o of Complexes in Relation to Solute-Solvent Interactions 44 3.7.2 The Dependence of ∆Hf o and ∆Sf o of Complex Species on Solvent Composition 45 4. Ligand Replacement Reactions 46 4.1 Correlation of Stability Constants with Ligand Replacement 48 4.1.1 The Cl--I- Reaction 48 4.1.2 The Br––I– and SCN-–I- Reactions 49 4.2 The Role of Solvents in Replacement Reactions 49 4.3 Replacement Reactions in Mixed Solvents 50 5. Structure of SO2-X - 50 6. Conclusion 51 7. References 51 1. Introduction Throughout the past 30 years there has been a growing interest in the field of sulfur compounds-halide ligands chemistry. The motivation for this interest has undoubtedly arisen from various sources. Witeckowa and Witok (1955) investigated the reaction between SO2 and iodine in the gas phase and in solutions by spectrophotometric and kinetic techniques. They suggested that the interaction between HI and SO2 in aqueous solutions is due to dipole-dipole interaction. Burke and Smith (1959) studied the molecular complexes between HF and SO2 by infrared spectroscopy. Jander and Tuerk (1962; 1963) studied the adduct of iodine with H2S in dichloroethane at –95 oC. The low enthalpy of formation (∆Hf o = -31.8 kJ /mol) was taken as indication of the charge transfer nature of the adduct formation. Burow (1970) studied 32 WEAK COMPLEXES OF SULFUR COMPOUNDS the solvate formation between SO2 and Cl -, Br-, and I- ligands in liquid SO2. Gutmann (1956) isolated a number of adducts of SOCl2 and SO2Cl2 with halide ligands. Sandhu et al (1960; 1962) discussed the tendency of SO2Cl2 to form adducts with Lewis acids and Lewis bases. Salama and Wasif et al (1971; 1973; 1994) studied the interaction of some organic sulfur compounds and SO2, SOCl2, SO2Cl2, & SOBr2 with Cl-, Br-, I-, and SCN- ligands in acetonitrile (MeCN), dimethylsulfoxide (dmso) and water (for SO2 species only). The collated data are summarized in four parts: 2. Detection, stoichiometry and stability of SO2-X -, SOCl2-X -, and SO2Cl2-X - complex species. 3. Effects of solvents on the stability of complex species. 4. Ligand replacement reactions in complex species and factors which affect them. 5. Structure of SO2-X - species. 2. Detection, Stoichiometry and Stability of SO2-X-, SOCl2-X-, and SO2Cl2-X- Complex Species 2.1 Detection of Sulfur Complex Species Figure 1 shows the absorbance peaks of mixtures of SO2 with (A) tetramethylammonium iodide, (B) tetramethylammonium bromide and (C) tetramethylammonium chloride. Table 1 includes the absorbance peaks of sulfur compounds, SO2, SOCl2 and SO2Cl2 with tetramethylammonium halides in which the halide was in abundance of the sulfur compound and in acetonitrile (MeCN) solvent. Similar peaks were obtained in dmso and water confirming earlier studies (Jander et al., 1937; Seel et al., 1955). Figure 1. Absorbance peaks of SO2-X - Species. Table 1. λmax in MeCN at 298°K Sulfur compound S + X– (S) Cl– Br– I– SO2 (280)nm SOCl2 (280) SO2Cl2 (275) 292 292 293 320 322 322 380 382 375 33 SALAMA and WASIF 34 2.2 Stoichiometry and Complex Species Job’s (1928) and Asmus’s (1960) methods were used to determine the complex stoichiometry. The former gave the empirical formula while the latter gave its molecular formula. The two methods showed that all complex species were of 1:1 type irrespective of the solvent used (MeCN or dmso). 2.3 Stability Constants of Complex Species If the complex formation is represented by the equations SO2 + X – = SO2-X – (1) SOCl2 + X – = SOCl2-X – (2) SO2Cl2 + X – = SO2Cl2-X – (3) S compound + Halide Ligand = Complex Species then the equilibrium constants for reactions (1) – (3), which will be defined throughout this article as stability constants, may be represented by the equation Kc = [Complex]/[S compound] [X –] (4) where the parentheses represent molar concentration of each species and Kc is defined by the units dm 3 mol–1 . The magnitude (or value) of Kc is taken as a measure of the ability of the reacting species to associate in a complex. Spectrophotometry was the principal technique used to find the concentration of all the terms in equation (4). Two procedures (Salama et al., 1971) were adopted to evaluate Kc: (a) graphical and (b) by calculation from the molar concentrations of reactants and products. Both methods (a & b) depend on the absorbances of individual species before and after they are mixed together. Table 2 (Salama et al., 1971) includes Kc values for SO2-X –, SOCl2-X – and SO2Cl2-X – in MeCN at 293° or 298°K. In most cases the difference between Kc values by the graphical and calculation methods does not exceed 5% which may be taken as the limit of the experimental error. The constancy of the Kc data is further evidence to confirm that all the complex species were of a 1:1 type (Salama et al., 1971). Table 2. Stability Constants of SO2-X –, SOCl2-X – and SO2Cl2-X – in MeCN at 293° or 298° K S Compound SO2 SOCl2 SO2Cl2 Halide ligand I– Br– Cl– I– Br– Cl– I– Br– Cl– Kc(Graphical) 37.9 192 348 152 242 367 78 40.6 10.6 Kc (Calculation) 38.5 190 363 150 241 362 77 41 10.2 Temp.°K 298 293 293 298 298 298 298 298 298 2.4 Thermodynamic Constants of SO2-X –, SOCl2-X – and SO2Cl2-X – Table 3 includes Kc data for SO2-X –, SOCl2-X – and SO2Cl2-X – over a range of temperatures and their relevant thermodynamic constants. The data in Table 3 point to a weak association between the sulfur compounds (electron acceptors) and the halide ligands (electron donors) of a charge transfer nature (Salama WEAK COMPLEXES OF SULFUR COMPOUNDS 35 et al., 1971; Ketelaar et al., 1952; Benesi et al., 1949; Drago, 1959; Rossotti et al., 1969; Andrew et al., 1961). Since the solvents used are polar, they possess varying tendencies to solvate the species in solution ( ions, molecules and complex species) and although we are mainly concerned with complex species in acetonitrile, yet a simple interpretation of such enthalpy data will be complicated by solvation and/or dipole interaction. Comparison of the enthalpy data for the different complex species in Table 4 cannot lead to linear correlations. Table 3. Thermodynamic Constants of SO2-X –, SOCl2-X – and SO2Cl2-X – in MeCN Complex Species Kc dm 3mol–1 –∆Gf o kJ/mol –∆Hf o kJ/mol –∆Sf o J/K/mol SO2-I – 37.9(298),34.4(303),30.8 (308) 8.99 17.5 28.5 SO2-Br – 260(283), 192 (293), 137 (303) 12.7 22.4 32.8 SO2-Cl – 519(284), 400 (293), 348 (303) 14.7 15.1 1.34 SOCl2-I – 190 (288), 150 (298) 12.2 16.9 15.1 SOCl2-Br – 276 (288), 241 (298) 13.6 9.82 –12.6 SOCl2-Cl – 447(288), 362 (298), 264 (308) 14.6 15.4 2.72 SO2Cl2-I – 86 (288), 77 (298), 71 (308) 10.8 6.01 -15.9 SO2Cl2-Br – 45 (288), 39 (298), 37 (308) 9.20 10.7 4.98 SO2Cl2-Cl – 12(288), 10.5 (298), 9.3 (308) 5.76 7.82 6.90 Of the components taking part in the formation of these complexes only the acceptors (sulfur compounds) have UV absorption peaks, SO2 (280nm), SOCl2 (280nm) and SO2Cl2 (275nm) (Friedman, 1967). The appearance of new peaks due to the formation of the complex species arises from donor- acceptor interactions. These result in spectral shifts for the acceptor which must be a function of the donor character of each halide ligand. An attempt was made to correlate such spectral shifts with the reversible potential for: e- + X = X- Table 4. Stability Constants of SO2-X -, SOCl2-X -, and SO2Cl2-X - in MeCN at 298oK Cl– Br– I– SO2-X – 372 160 38 SOCl2-X – 362 241 150 SO2Cl2-X – 10.5 41.0 77 The linear plots of Figure 2 supplement this assumption for the different species. SALAMA and WASIF Figure 2. Correlation of Eo for the Reaction e- + X ⇔ X- with Acceptor Spectral Shifts: A, SO2-X -; B, SO2Cl2-X -; C, SOCl2-X - 2.5 The Donor-Acceptor Nature of the Complex Species In order to understand the nature of these complexes we shall try to rationalize the stability constants data of Table 4 with the nature of the halide ligands (donors) and the sulfur compounds (acceptors). Table 5 summarizes some important trends (Basolo et al., 1958 ; Gould, 1960). Table 5. Physical Constants of Halide Ions. Cl- Br- I- Ionisation Potential (kJ/mol) 1251 1136 1000 Electron Affinity (kJ/mol) 349 325 295 Hydration Energy of X– kJ g-ion-1 356 310 255 Polarisability/Å 2.3 3.3 5.1 Electronegativity 2.83 2.74 2.21 The iodide ion with high polarisability, low electronegativity and easy oxidation is considered to be a soft Lewis base (Pearson, 1963; Day et al., 1969). The chloride ion with low polarisability and high electronegativity is a hard Lewis base. The bromide ion is a borderline Lewis base. The acceptors include SO2 and SOCl2 in which the oxidation state of sulfur is four and SO2Cl2 in which it is six. SO2 is a borderline Lewis acid (Pearson, 1963; Day et al., 1969). It acts as a base toward BF3 to form the adduct SO2-BF3 and as acid towards water. Thionyl chloride SOCl2 is similar to SO2 in that it has a lone pair of electrons (3s2) but one of the double bonded oxygen atoms is replaced by two Cl atoms. The S-Cl bond is more polarized than the S-O bond owing to the higher electronegativity of Cl-, and would be expected to act as a stronger Lewis acid than SO2, or a better acceptor. Sulphonyl chloride SO2Cl2, may be related to SO3 (known as a hard Lewis acid) (Pearson, 1963; Day et al., 1969) in the same manner as SOCl2 is related to 36 WEAK COMPLEXES OF SULFUR COMPOUNDS 37 SO2, and the order of acid strength is SO2Cl2 > SOCl2 > SO2. The formation of the present complex species is the result of acid-base interactions between the acceptors and the donors and the order of stability given in Table 4 can be discussed on this basis. 2.6 Correlation of the Order of Stability Constants of Different Complexes with Donor-Acceptor Properties Table 4 shows that the stability of SO2-X – species falls in the order SO2-Cl – > SO2-Br – > SO2-I –. Sulfur (IV) forms coordination compounds owing to the electrofilic and nucleofilic nature of the sulfur atom. The former is due to the availability of the empty 3d electron orbitals and the latter to the presence of a lone pair of 3s2 electrons on the sulfur atom. Thus, in such compounds as SO2, sulfur acts as a σ-donor only or a Π-acceptor. However, the donor-acceptor properties of the sulfur atom are exhibited almost synonymously. If the donor (or ligand) contains d-orbitals of the appropriate symmetry (i.e. not diffuse) back-donation from the sulfur atom to the donor may occur, giving rise to the d-d multiple bonding which will strengthen the ligand acceptor bond. The Cl– and Br– ions may accept back-donation but this seems doubtful for the I– ion because the d-orbitals become progressively diffuse and less available for back- donation as we go down the halide group. The order of complex stability can be explained on this basis. This interpretation of the stability constants order is supplemented by the classification of the halide ions as hard (Cl-) borderline (Br–) and soft (I–) bases and of SO2 as a borderline Lewis acid. The order of stability of the SO2-X – complex species would follow the strength of the base and SO2-Cl – species would be the strongest and SO2-I – the weakest, as actually found. We may now consider the SOCl2-X – species. Table 6 includes the ratios of stability constants of SO2-X – and SOCl2-X – species. Table 6. Stability Constants Ratios for SO2-X – and SOCl2-X – in MeCN I– Br– Cl– Kc(SO2-X –)/Kc(SO2-I –) 1.0 4.0 10.0 Kc(SOCl2-X –)/Kc (SOCl2-I –) 1.0 1.6 2.4 Table 4 shows that the order of stability of SOCl2-X – species is similar to that of SO2-X – where Cl– > Br– > I–. Table 6 shows that the stability constants of SO2-Cl - and SOCl2-I – are nearly of the same order of magnitude but Kc (SOCl2-I -) is merely 4 times greater than Kc(SO2-I -) and Kc(SOCl2-Br –) and Kc(SOCl2- Cl–) are only 1.6 and 2.4 times greater than Kc (SOCl2-I –) which calls attention to new factors responsible for the observed change in ratios. The order of stability constants ratio of SOCl2-X – species shows that as in the case of SO2-X – the chloride species is the most stable and the iodide is the least stable. This order suggests that the nature of association between the Cl– and SO2 is much the same as with SOCl2. Back-donation may be considered to be the factor contributing to the stability of SOCl2-X – species. In SOCl2 the electrophilic nature is enhanced over SO2 by the replacement of one oxygen atom by two chlorine atoms and the d-orbitals of sulfur are more exposed for coordination because the electron cloud is removed by the electronegative chlorine atoms, and this makes SOCl2 a better Lewis acid than SO2. Although the halide ions were classified by Pearson (1963) as Lewis bases of varying strengths and such classification could account for the order of stabilities of SO2-X – species, the situation may be different with SOCl2. The increased acceptor character does not appear to have changed or to have affected dramatically the nature of association with the Cl– ligand. For the I– ligand this increased acid character appears to have increased the basicities of the I– and Br– ligands relative to that of the Cl– , so SOCl2 appears to be leveling up the basic character of the I– and Br– ligands towards that of the Cl– ligand. This leveling of relative basicities of the halide ligands appears to be another factor which determines the ratios of Table 6. SALAMA and WASIF 38 A third factor relevant to Table 6 is the increased ionic radii and polarisabilities of the donors. Table 5 shows that the I– ligand is the most polarisable of the halide ions. Other factors remaining equal, an increase in the polarisability of the donor would make the donor-acceptor interaction stronger. The dipole moments of SO2 and SOCl2 are 1.61 and 1.60D respectively. If polarisability was the only factor one would expect the iodide complex to be the most stable. This was not so, indicating that back-donation is a still more important factor in deciding the nature of association of the halide ligands with SOCl2. Sulphonyl chloride, SO2Cl2, is the strongest acceptor of this group of sulfur compounds; it has the highest dipole moment (1.86 D). The d-orbitals of sulfur here are the most exposed for coordination than in the other acceptors. The order of stability of its halide ligand complexes is: SO2Cl2-I – > SO2Cl2-Br – > SO2Cl2-Cl – which is the reverse of SO2-X – and SOCl2-X - species (Table 4) suggesting that back-donation cannot be strong in the formation of SO2Cl2-X – complexes. The increased acidity of SO2Cl2 seems now to be very important. In the presence of such a relatively strong Lewis acid the three Lewis bases appear to lose their identity and are of merely equal strength. Thus the leveling effect observed for SOCl2 is probably more strongly displayed. In protonic systems this leveling explains why benzoic acid and sulfuric acid are equally strong in liquid ammonia while water, alcohol, ketones…. etc are equally strong bases in pure sulfuric acid (Bell, 1965; Waddington, 1965). It appears that increased polarisability towards I–, the increased polarity towards SO2Cl2 and the increased leveling effect can account for the order of stability constants observed for SOCl2–X - complex species. The effect of polarisability of the halide ligands on the order of stability of some metal complexes has been reported (Gould, 1960). We may conclude that as the acceptor is changed from SO2 to SO2Cl2 the nature of association also changes. With SO2 dΠ - dΠ multiple bonding from back-donation makes its association with halide ligands quite strong but with SO2Cl2 the dipole-dipole interaction seems to be a weaker force of association, as shown from ∆Hf o values in Table 3 and the Kc values of Table 4. 3. Solvent Effects on Stability of SO2-X–, SOCl2-X– and SO2Cl2-X– 3.1 Pure Solvents Solvent molecules are not impartial in chemical processes and the extent to which they participate sometimes overshadows that of the other species in the reaction media. This is because the solvent represents the environment in which a chemical reaction takes place and in most cases plays the role of a donor or acceptor. The role of environment and solvent effects on chemical reactions has been discussed by a number of workers (Bell, 1965; Waddington, 1965; Frost et al., 1961; Benson, 1960; Amis, 1965; Gutmann, 1967, 1971; Grunwald, 1949) . The Kc values for SO2-X –, SOCl2-X – and SO2Cl2-X – in MeCN, dmso and water recorded in Table 7 illustrate the solvent effects. Table 7. Stability Constants for SO2 –X –, SOCl2-X – and SO2Cl2 –X – in MeCN, dmso and water at 298° K. (a = MeCN, b = dmso, and c = water) Cl– Br– I– X– a b a b c a b c SO2-X – 372 26 160 21 0.22 38 12 0.36 SOCl2-X – 362 18 241 21 150 35 SO2Cl2-X – 10 36 41 14 71 6 WEAK COMPLEXES OF SULFUR COMPOUNDS Comparison of Kc data is limited to the values in MeCN and dmso. With the exception of SO2Cl2-X - species, the stability constants of different complexes decreased by a factor of nearly 20 in dmso as compared to MeCN which is shown from the data in Table 8. Table 8. Comparison of Stability Constants in MeCN and dmso Kc(MeCN)/Kc(dmso) at 298 oK. X- Cl- Br- I- SO2-X - 15 8 3 SOCl2-X - 20 11 4 SO2Cl2-X - 0.3 3 13 The data in Tables 7 and 8 express significant changes in complex stabilities between MeCN and dmso, as the Kc values are lower in dmso than in MeCN. There are at least two possible roles that can be played by dmso (or water) in affecting the stabilities of the complex species. It may solvate the halide ligands, which prevents them from interacting with the acceptors, i.e. sulfur compounds, or it may act as a competing acceptor, i.e. competes against SO2, SOCl2, and SO2Cl2 and thus makes a complex species with the ligands and perhaps it may play the two roles depending on the environmental conditions. Table 8 shows some horizontal and vertical trends. The data in column (1) show maximum decrease in Kc for SO2-Cl- and also for SOCl2-Cl- in dmso. This can be attributed to solvation of the Cl- ligand and possibly ion-dipole interaction in view of the high dipole moment of dmso (4.3D). In these species, stabilization results from back-donation (Mines et al., 1972; Chadwick, 1973) from sulfur 3d-orbitals to the donor ligand. Minimum effect is shown for SO2Cl2-Cl - where Kc(MeCN)/Kc(dmso) equals 0.3, which shows the importance of polarization and not back-donation in stabilizing SO2Cl2-Cl - species. In column (2) ratios for SO2-Br – and SOCl2-Br – are nearly half those reported for the Cl– species, which is regarded as reflecting the lower tendency to solvation of the Br– ligand as compared to Cl– by dmso. In column (3) the ratios of SO2-I – and SOCl2-I – are nearly ¼ of the values for the Cl– species. This shows the lower tendency of I– ligands to solvate and this is not unexpected due to the larger ionic size of I–. The horizontal trends in SO2-X – and SOCl2-X – appear to agree with the conclusion that solvation of the donor ligands by dmso is important in decreasing the stability constants compared with those in MeCN and that this lowering is maximal with Cl– ligands and minimal with I– ligands. The Kc values in water appear to support this view. For SO2Cl2-X – the horizontal tend is reversed compared to that shown by SO2-X – and SOCl2-X –. This supports the view that the nature of association in SO2Cl2-X – is different from that in SO2-X – and SOCl2-X – (Salama et al., 1971), the former being mainly ion-dipole interaction and the latter back-donation as mentioned earlier. The lowering in stability constants of SO2Cl2-I - and SO2Cl2-Br - is evidence that ion- dipole interaction is particularly strong between dmso and iodide ion (Salama et al., 1991), which is understandable in view of the higher dipole moment of dmso compared with SO2Cl2. The data in Table 8, column (3) for the iodide species show that ion-dipole interaction outweighs solvation, while those in column (1) show that solvation has the greater effect. 3.2 Mixed Solvents: (MeCN-dmso) Few workers have reported on chemical processes in mixed solvents and also specified the role of solvent. The work to be outlined reports on the stability of SO2-X –, SOCl2-X - and SO2Cl2 -X – in MeCN- dmso mixed solvent, and Table 9 (Salama et al., 1978) includes the Kc data of these complex species in MeCN, dmso and mixtures of the two solvents at 298º K. The data in columns 2 and 6 were quoted from Table 2. One feature appears throughout Table 9. The stability constants of all complex species, at 298ºK, vary with solvent composition. In order to rationalize 39 SALAMA and WASIF 40 the Kc values we shall discuss the horizontal and vertical trends in Table 9. Table 9. Stability Constants for SO2-X -, SOCl2-X -, and SO2Cl2-X - in MeCN, dmso, and their mixtures at 298oK Complex dmso MeCN+dmso MeCN 1:3 1:1 3:1 SO2 -I – 12.1 14.1 24.7 36.7 37.9 SO2 -Br – 21.0 14.5 24.1 40.1 40.1 SO2 -Cl – 26.0 35.5 50.0 52.4 372.0 SOCl2 -I – 35.0 40.2 58.9 223.5 150.0 SOCl2 -Br – 21.0 22.9 34.2 71.7 241.0 SOCl2 -Cl – 18.0 66.7 73.8 113.0 362.0 SO2Cl2 -I – 6.0 30.1 50.1 142.2 76.0 SO2Cl2 -Br – 14.0 15.5 24.6 36.2 41.0 3.3 The SO2-X – Species The Kc values of SO2-I – change gradually between the limits set for MeCN and dmso as the solvent composition changes. In MeCN-dmso solvent (3:1) the Kc value is reduced by 2.7% compared with that in 25% dmso solvent. For SO2-Br – and SO2-Cl – the Kc values are reduced in the same solvent by 75% and 85% respectively. (The Kc value in MeCN is taken as a reference in each case). Solvation of Br – and Cl– ligands by dmso is clearly greater than that for the I– ligand. A similar conclusion was reached when Kc for the same species were determined in MeCN and dmso (Salama et al., 1975). To correlate the variations in Kc values with solvent composition it should be remembered that the molecules of different solvents can act as donors and/or acceptors. Even if both solvent molecules have donor or acceptor character a slight difference in the donor or acceptor properties between different solvent molecules will invite donor- acceptor interaction between them. The nature of such interactions in non-protic solvent mixtures such as MeCN-dmso has not been studied before (Waddington, 1965). Over a wide range of molecular ratios of MeCN and dmso one expects such interactions to exhibit different patterns which depend on the structural and geometrical characters of the molecules. We shall call this solvent-solvent interaction and may define it in terms of donor-acceptor interaction or perhaps association which vary with solvent composition. Such solvent-solvent interactions may occur at the expense of other interactions in solution. For example in MeCN-dmso containing SO2 molecules and Cl – ligands and SO2-Cl – complex species the following interactions are likely to occur (a) Cl–-dmso, (b) Cl–-MeCN, (c) SO2-dmso, (d) SO2-MeCN and (e) MeCN-dmso. The stability constant of the SO2-Cl – species is determined by the relative magnitudes of such interactions, a strong Cl–-solvent interaction (solvation) would reduce the stability of SO2-Cl – species, since this steric factor may prevent, to some extent, SO2 and Cl – from approaching each other for coordination. On the other hand a strong MeCN-dmso interaction would allow more SO2 to coordinate with Cl– ligands and Kc values become greater than the limits set for each solvent. The data for SO2- X – show that MeCN-dmso interaction accounts partly for the change in Kc with solvent changes. Ligand solvation is also important in determining Kc values. The Kc data show that while the I–-solvent interaction is relatively weak and reduces Kc by 2.7% (in 25% dmso relative to its value in MeCN) those of Cl–-solvent interactions and Br–-solvent interactions are much stronger and reduce Kc by 85% and 75% for SO2-Cl –- and SO2-Br – respectively. 3.4 The SOCl2-X – and SO2Cl2-X – Series The Kc values for SOCl2-Cl – and SOCl2-Br – fall between the limits set for MeCN and dmso but that for SOCl2-I – species exceeds the upper limit in MeCN by 30% (in 25% dmso solvent). A change in solvent WEAK COMPLEXES OF SULFUR COMPOUNDS 41 composition from 50 to 25% dmso is coupled with a sudden change in Kc for SOCl2-I –. It appears that over this critical range of solvent composition the solvent-solvent interaction reaches its maximum. The nature of this interaction is not yet clear but is detectable from vapor pressure measurements (Salama et al., 1985). Such interactions have freed sufficient I– ligands and SOCl2 molecules for coordination which makes Kc 30% greater than the value in MeCN. With SOCl2-Cl – and SOCl2Br – (in 25% dmso solvent) the Kc values are reduced by 69% and 71% from the value in MeCN. Thus, despite strong solvent-solvent interaction the Br--solvent and Cl--solvent interactions have outweighed solvent–solvent interactions. A similar argument applies to SO2Cl2-I - and SO2Cl2-Br - where Kc increases by 61% for the iodide species and decreases by 15% for the bromide species. 3.5 Vertical Correlations in Table 9 3.5.1 The Iodide Complex Series Although Kc values for SOCl2-I – and for SO2Cl2 -I – in 25% dmso are 30% and 64% greater than the value in MeCN we find that Kc for SO2-I – is 2.7% less than its value in MeCN. Such differences in behavior are probably due to differences in the manner in which the I– ligand is coordinated to the three acceptors. The I- ligand is polarisable with diffuse d-orbitals suitable for ion-dipole interaction. The dipole moment of the three acceptors are in the order SO2 (1.61D) SOCl2 (1.60D) and SO2Cl2 (1.86D). In 25% dmso solvent strong solvent-solvent interaction favors SO2Cl2-I – coordination so that the Kc value exceeds that in MeCN by 64%. The acceptor character of SOCl2 is enhanced by the replacement of one O atom in SO2 by 2 Cl atoms and coordination by ion-dipole interaction is relatively stronger with I – than with SO2. Solvent-solvent interaction helps coordination of SOCl2 to I - and results in a Kc value which exceeds that in MeCN by 30%. For SO2-I – the ion-dipole interaction is probably so weak that is nearly balanced by solvent- solvent interaction and Kc is slightly reduced by 2.7%. 3.5.2 The Bromide and Chloride Complex Series The decrease in the Kc values for SO2-Br –, SOCl2-Br – and SO2Cl2-Br – (in 25% dmso solvent) by 76, 71 and 15% respectively resulted from Br–- solvent interaction. The difference arises from differences in the modes of coordination of the three acceptors. For SO2 and SOCl2 back-donation is the principal mechanism by which coordination takes place to S(IV) . For SO2-Br – and SOCl2-Br – solvent-solvent interaction is outweighed by Br–- solvent interaction and Kc is reduced accordingly. For SO2Cl2-Br - coordination occurs by an ion-dipole mechanism involving S(VI). The Br– ligand, being a borderline Lewis base, responds favorably to this mechanism and by its solvation by dmso slightly outweighs solvent-solvent interaction and Kc is reduced to a smaller extent than in the other bromides. For SO2-Cl – and SOCl2-Cl – back-donation is the principal mechanism for coordination. In both species the Cl–-solvent interaction outweighs solvent-solvent interaction and Kc values are reduced by 85 and 69% respectively. This differing effect on Kc might be due to selective solvation in the presence of different acceptors which we may describe as chemical environmental factors. 3.6 Evidence for Solvent-Solvent Interactions 3.6.1 Spectroscopic Using the IR (infra red) techniques, it is found that the S=O vibration band of dmso appears at 1080 cm–1 and on addition of MeCN the band is shifted to 1070cm–1. The C ≡ N vibration band of MeCN appears at 2250cm–1 and adding dmso at x1 = 0.5 (x1 = mol fraction) the band is shifted to 2240cm –1 (Salama et al., 1985). Using Raman spectra it is found that the S = O vibration band appears at 1044cm–1 and on adding MeCN a peak appears at 1062cm–1 while that at 1044cm–1 disappears. For MeCN the C ≡ N vibration band appears at 2255cm–1 and on addition of dmso the band disappears gradually. The observed vibrational shifts SALAMA and WASIF may be taken as evidence for solvent-solvent interaction through adduct formation of the type represented by Figure 3. Figure 3. MeCN-dmso Adduct 3.6.2 Vapor-Pressure, Viscosity, and Excess Functions from Refractive Index, Dielectric Constant and Volume Further confirmation for solvent-solvent interaction was obtained from measurements of vapor pressure and viscosity of MeCN-dmso mixtures. Table 10 includes vapor pressure, viscosity, ∆Hvap, ∆Svap, and ∆Hvis, of the mixtures over the whole composition range. Table 10. Vapour pressures and viscosities of MeCN-dmso mixtures 10(a). Vapour Pressure Data 25% MeCN t/°C p/mmHg 61.6 58.0 69.0 72.0 75.6 84.0 80.7 99.0 85.7 115.5 50% MeCN t/°C p/mmHg 29.0 69.7 34.5 76.1 39.9 95.6 45.0 115.6 49.0 136.3 54.0 163.6 75% MeCN t/°C p/mmHg 28.4 78.9 33.1 95.4 38.5 119.4 43.7 147.9 49.4 186.9 80% MeCN t/°C p/mmHg 28.4 84.7 33.4 106.5 38.9 134.0 45.1 160.0 51.6 213.0 10(b). Viscosity Data t/°C 25.0 30.0 35.0 MeCN 3.55 3.38 3.30 x 10–4 Pas 75%MeCN 5.17 4.89 4.73 50%MeCN 7.44 7.04 6.68 25%MeCN 11.71 10.85 10.16 dmso 19.57 17.88 16.40 42 WEAK COMPLEXES OF SULFUR COMPOUNDS 10(c). Thermodynamic Data for Vaporization and Viscosity Solvent MeCN 80% MeCN 75% MeCN 50% MeCN 25% MeCN dmso ∆Hvap kJ/mol 33.2 31.8 33.3 29.1 28.2 52.9 ∆Svap J/K/mol 96.0 87.1 91.4 72.9 62.7 95.8 ∆Hvis kJ/mol 5.48 6.57 7.74 8.24 10.6 13.3 Figure 4 shows a plot of ∆Hvap and ∆Hvis against solvent composition. The plots are not linear and deviate over the composition range 50–70% (maximum deviation which may be taken to indicate solvent- solvent interaction). The ∆Svap data show a minimum value at 75% dmso suggesting maximum order for the system at this composition with probable formation of MeCN-dmso adduct. 0 20 40 60 80 100 %MeCN 20 30 40 50 60 E nt ha lp y of V ap or iz at io n/ kJ 0 4 8 12 E nt ha lp y of V is co si ty /k J Figure 4. Plot of Enthalpy of Vaporization/kJ (•) and Enthalpy of Viscosity /kJ (♦) against %MeCN The excess functions from refractive index nE, dielectric constant εE and densities VE of MeCN-dmso mixtures are given in Table 11 (Salama et al., 1985). The data in Table 11 show that the magnitude of any excess function reaches a maximum at approximately 1:1 composition. This is again taken to indicate solvent-solvent interaction. Table 11. Excess Functions of MeCN-dmso mixtures at 298oK 11(a). Refractive Index (nE) x1 0.1319 0.2676 0.4258 0.6899 0.7836 0.8697 nE 0.0030 0.0076 0.0087 0.0102 0.0065 0.0039 43 SALAMA and WASIF 44 11(b). Dielectric Constant (εE) x1 0.1319 0.2676 0.4258 0.6009 0.6899 0.7836 0.8697 εE 0.5710 0.9540 1.3180 1.4910 1.3680 1.1960 0.8630 11(c). Volume (VE). (x1 = mole fraction of dmso) x1 0.1006 0.3066 0.6037 0.7061 0.8083 0.8999 0.9455 VE –0.1300 –0.2523 –0.2519 –0.1773 –0.1275 –0.0876 –0.0486 3.7 A Thermodynamic View of Solvent Effects on The Stability of SO2-X-, SOCl2-X- and SO2Cl2-X- 3.7.1 The Significance of ∆Gf o of Complexes in Relation to Solute-Solvent Interactions: Table 12 includes the standard free energies of formation ∆Gf o of SO2-X - and SOCl2-X - in MeCN, dmso, and their mixtures. Table 12. –∆Gf o (kJ/mol) of complex species in different solvents at 298oK dmso: MeCN dmso 3:1 1:1 1:3 MeCN SO2-I – 6.19 6.57 8.03 9.00 9.08 SO2-Br – 7.61 6.74 7.91 9.20 8.49 SO2-Cl – 8.12 8.95 9.75 10.9 10.6 SOCl2-I – 8.87 9.20 10.2 13.5 12.5 SOCl2-Br – 7.61 7.82 8.79 10.7 13.7 SOCl2-Cl – 7.20 10.5 10.7 11.8 14.6 The data in Table 12 show that for every complex species there are several free energy minima, each corresponding to a different solvent composition. This situation is only possible if a change in the solvent composition affects continuously the coordinating ability of the halide ligand (X–) with the sulfur acceptor (SO2 or SOCl2) and partially hinders them from complex formation and which was described as solvation. Preliminary studies (Wasif, unpublished work) show that halide ion solvation in MeCN and dmso falls in the order Cl– > Br > I–,which agrees with their ionic radii and charge densities. Solvation of the sulfur acceptors was studied in the present work by UV spectroscopy. Figure 5 shows the absorbance of SOBr2 in (a) MeCN, (b) dmso and (c) 1:1 mixed solvent of MeCN-dmso. It shows three distinct species which obey Beer’s law. The intermediate absorbance of the 1:1 mixed solvent shows that SOBr2 forms an absorbing species of intermediate character between the species in dmso and MeCN. There are two possibilities in which this could happen: (1) The formation of a constant ratio of the adduct species SOBr2-MeCN and SOBr2-dmso, (2) that SOBr2 makes a species with a mixed solvent adduct e.g. (SOBr2-MeCN:dmso). Using CCl4 as solvent the species SOBr2-MeCN and SOBr2-dmso were detectable and their stability constants are given in Table 13. The data in Table 13 show that solvents MeCN and dmso play a competing role against halide ligands in their coordination with the sulfur acceptors. The data also show dmso to have a greater destabilizing role WEAK COMPLEXES OF SULFUR COMPOUNDS Table 13. Stability Constants of Adducts of MeCN and dmso with Sulfur Compounds at 298°K. A SO2 SOCl2 SOBr2 SO2Cl2 dmso-A 2.82 6.86 3.73 17.79 MeCN-A 0.34 0.07 0.11 0.68 for the complex species than does MeCN, by its strong ability to solvate the halide ligands and the sulfur compounds. Figure 5. Dependence of Absorbance of SOBr2 On Solvent Mole Fraction 3.7.2 The Dependence of ∆Hf o and ∆Sf o of Complex Species on Solvent Composition: A second thermodynamic aspect would be to consider the significance of ∆Hf o and ∆Sf o for complex species in mixed solvents. Table 14 includes the standard thermodynamic constants for the formation of the complex species SOBr2-Cl – and SOBr2-Br –. Table 14 shows that the standard enthalpy of formation varies as the solvent composition changes from MeCN to dmso. For SOBr2-Br – ∆Hf o is nearly 10 times greater in MeCN than in dmso but for SOBr2- Cl– it is nearly 5 times greater than the value in dmso. Such differences in ∆Hf o values suggest that the 45 SALAMA and WASIF 46 measured enthalpy of formation is a rather complex function. It does not probably represent the heat of formation of the complex species but other heat terms are possibly embodied in this term such as heats of solvation of the halide ligands and sulfur compounds in MeCN and dmso. If ∆Hf o be taken as a rough measure for the complex stability, then data in Table 14 would show that the complex species SOBr2-X – (X- = Cl, Br) are more stable in MeCN than they are in dmso. A thorough discussion of the significance of ∆Hf o data requires a knowledge of the heats of solvation of the different species in both solvents which are not at present available. This situation permits a qualitative discussion of ∆Hf o data. Since solute-solvent interactions are a dynamically changing process we may expect the magnitude of ∆Hsolvation values to change over the solvent concentration range, which is confirmed directly from Table 14. With this situation in mixed polar solvents a discussion of ∆Sf o values would be difficult to interpret. Table 14. Thermodynamic constants of SOBr2-Cl – and SOBr2-Br – in MeCN, dmso and their Mixtures at 298oK. dmso: MeCN dmso 3:1 1:1 1:3 MeCN SOBr2-Br – Kc dm 3 /mol –∆Gf o kJ/mol –∆Hf o kJ/mol ∆Sf o J/K/mol 12 5.86 1.88 13.4 19 7.36 4.60 9.20 25 8.03 4.60 11.3 32 8.66 20.1 -38.5 203 13.3 19.2 -20.1 SOBr2-Cl – Kc dm 3.mol–1 –∆Gf o kJ/mol –∆Hf o kJ/mol ∆Sf o J/K/mol 19 7.36 1.88 18.4 36 8.95 10.0 -2.93 40 9.20 9.62 -1.26 79 10.9 16.3 -18.0 100 11.5 10.0 5.02 4. Ligand Replacement Reactions The coordination of different ligands (Cl–, Br– I– or SCN–) with the same sulfur acceptors giving varying stability constants suggested that they are differently coordinated and could accordingly be able to replace each other with the same acceptor. Equation (5) shows a general replacement reaction between SO2- I– and X– (X- =Cl, Br, SCN) and Figure 6 illustrates a spectrophotometric scan when SCN– is added to SO2- I- in MeCN (Salama et al., 1978). SO2-I - + X- = SO2-X - + I- (5) WEAK COMPLEXES OF SULFUR COMPOUNDS Figure 6. Replacement of I– by SCN– in MeCN a) No SCN–, b) [SCN–] / [I–] = 0.4, c) [SCN–] / [I–] = 1, d) [SCN–] / [I–] = 2 In Figure 6 the gradual addition of SCN– solution causes a gradual disappearance of the SO2-I – peak at 378 nm and appearance of a new peak at 322 nm for the SO2-SCN – species. Table 15 includes the results of the above replacement reactions. Table 15. Replacement Reactions in MeCN, dmso and Water at 298o K* (a) In MeCN Reaction 1) Cl– + SO2-I – 2) Cl– + SOCl2-I – 3) Cl– + SO2Cl2-I – [Cl–] : [I–] 0.05(5), 0.10(8), 0.15(12), 0.20(15), 0.30(19), 0.40(23) 0.05(4), 0.10(7), 0.15(11), 0.20(15), 0.30(18), 0.40(22) 0.05(2), 0.10(3.5), 0.15(5.5), 0.20(8), 0.30(10), 0.40(12.5) 4) Br– + SO2-I – 5) Br– + SOCl2-I – 6) Br– + SO2Cl2-I – [Br–] : [I–] 0.05(1.5), 0.10(3), 0.20(5), 0.30(8), 0.40(10) 0.05(2), 0.10(3.6), 0.20(5.8), 0.30(7.7), 0.40(9.3) 0.05(2.5), 0.10(4), 2.20(6), 0.30(8), 0.40(9.4) 7) [SCN]– + SO2-I – 8) [SCN]– + SO2Cl2-I – [SCN–] : [I–] 0.20(–), 0.50(2), 1.0(3), 2.5(4), 4.0(6) 0.20(–), 0.50(1.8), 1.0(2.9), 2.5(4), 4.0(5.9) (b) In dmso Reaction 9) Cl– + SO2-I – 10) Cl– + SOCl2-I – 11) Cl– + SO2Cl2-I – [Cl–] : [I–] 0.05(3), 1.0(5), 1.5(7), 2.0(10), 3.0(13), 4.0(17) 0.50(3), 1.0(4), 1.5(6), 2.0(9), 3.0(12), 4.0(15) 2.0(8), 4.0(16) 12) Br– + SO2-I – 13) Br– + SOCl2-I – [Br–] : [I–] 0.50(4), 1.0(6), 2.0(11), 4.0(13) 0.50(4), 1.0(5), 1.5(7), 2.0(10) (c) In water 14) Cl– + SO2-I – 4.0(3), 8.0(5), 12.0(7) *No replacement was observed by [SCN]– in dmso. Percentage replacements of the iodide species are given in parentheses. In all the above reactions the extent (or magnitude) of replacement depends on the stability constant for the reactant and product complex species. Table 16 includes the Kc data for reactants and products in the 47 SALAMA and WASIF replacement reactions. Another factor which determines the magnitude of the replacement is the Lewis basic character of the ligands and a third factor appears to be related to the donor and acceptor number of the solvent used (Salama et al., 1971; Pearson, 1963; Day et al., 1969). Table 16. Stability Constants for Reactants and Products in Replacement Reactions at 298oK X- Cl- Br- I- SCN- SO2-X - 372 160 38 65 SOCl2-X - 362 240 150 77 SO2Cl2-X - 10 41 77 298 4.1 Correlation of Stability Constants with Ligand Replacement 4.1.1 The Cl--I- Reaction The data in Table 15 show that in the presence of two halide ligands, Cl- and I- and acceptors such as SO2, SOCl2 and SO2Cl2 in MeCN solvent the thermodynamics would be more favorable for SO2-Cl - and SOCl2-Cl - than for SO2-I - and SOCl2-I -. We may add that Kc(SO2-Cl -) is nearly equal to Kc(SOCl2-Cl -) and both are much higher than Kc(SO2Cl2-Cl -)in the ratio 37: 36: 1. Such a large difference in stability constants makes replacement of I- by Cl- much easier for SO2-I - and SOCl2-I - than it is for SO2Cl2-I -. The data in Table 15 and the plot of Figure 7 illustrate this observation and we note that the percentage replacement in SO2-I - and SOCl2-I - fall at a low [Cl-]/[I-] on the same line and are much higher for SO2Cl2-I -. The replacement reaction, equation (6), seems anomalous in view of the Kc(SO2Cl2-I -) / Kc(SO2Cl2-Cl -) ratio = 8. That this reaction, equation (6), can take place despite the reversed order of stability constants for reactant and product can only be due to the relative abundance of the Cl– ligand which seems to outweigh the difference in stability constants. Table 15 shows that a 20% disappearance of SO2-I – requires a ligand ratio of 0.35:1 while in the case of SO2Cl2-I – ligand ratio [Cl–] / [I–] = 1 : 1 was necessary for the same percentage replacement. By changing the ligand ratio we are merely increasing the chances of effective collisions leading to a replacement by the more abundant ligand. SO2Cl2-I - + Cl- = SO2Cl2-Cl - + I- (6) Figure 7. Comparison of Replacement Reactions by added Cl– in MeCN at 298°K (•) SO2-I - and SOCl2-I - ; (o) SO2Cl2-I - 48 WEAK COMPLEXES OF SULFUR COMPOUNDS 4.1.2 The Br––I– and SCN-–I- Reactions The magnitudes of replacement for the ligands Br– and SCN– are not widely different (Table 15). This is not unexpected since the stability constants for SO2-Br –, SOCl2-Br – and SO2Cl2-Br – are in the ratio 4:6:1 and for SO2-SCN – and SOCl2-SCN – and SO2Cl2-SCN – are in the ratio 1: 1.2 : 2.4. Figure 8 shows a comparison of the replacing ability of the SCN– ligand with that of Cl– and Br– ligands. The replacing ability of ligands fall in order Cl– > Br– > SCN–, which parallels their Lewis base character as hard, borderline and soft respectively (Pearson, 1963; Day et al., 1969). Figure 8. Comparison of Amounts of Replacements of I– in SOCl2-I – 4.2 The Role of Solvents in Replacement Reactions Table 15 and Figure 9 show the effect of solvents on the reaction SO2-I – + X– where X- = Cl–. For a ligand ratio [Cl–] /[I– ] = 4 the extent of replacement is 70% in MeCN but only 16% in dmso and less than 2% in water. Clearly a change of solvent greatly affects the degree of replacement. In dmso two factors are operative: a) dmso may solvate the Cl– ligand and b) dmso may act as a potential acceptor at its S atom and compete with SO2 in solution. The conclusion that dmso acts as a potential acceptor was reached using the observation of the spectrum of a solution containing Cl– and I– ligands. This solution shows slow growth of two peaks at 292 nm and 365 nm and the rate of peaks growth depends on the ligand concentration and a rate constant was found to be 2 x 10–2 min–1 at 298°K. The replacement reaction SO2-I – + Cl– = SO2-Cl – + I– in dmso is ionic with a rate constant of 1010 sec-1 and therefore the addition of Cl– to SO2-I – in dmso is accompanied by the disappearance of the peak at 378nm (SO2-I –) and the emergence of the new peak at 292 nm (SO2-Cl –). If this solution is left for some time the peaks at 292nm and 365nm will slowly appear for the complex species of dmso with Cl– and I– ligands respectively. The sulfur atom in dmso will be the acceptor center with the following structures: OMe2S-Cl – and OMe2S-I – respectively. Recent work on photoelectron spectroscopy of sulfur compounds favors this view (Guest et al., 1972; Buncel et al., 1975). Water appears to be a more drastic solvating agent towards the Cl– ligand which results in a low degree of replacement with high ligand ratio of [Cl–] / [I–] = 12. 49 SALAMA and WASIF Figure 9. Effect of Solvents on the Reaction SO2-I – + Cl– → SO2-Cl – + I– in a) MeCN, b) dmso and c) water at 298°K. 4.3 Replacement Reactions in Mixed Solvents Table 17 includes the replacement reaction data for the SO2-I – + Cl– reaction in MeCN-dmso mixed solvent. For all runs the ligand ratio was [Cl–]/[I–] = 4. The replacement percentage of iodide is given in parentheses. As the percentage of dmso in the solvent increases the ability of Cl– to replace I– diminishes. Here it is possible that the solvating power of dmso towards Cl– makes it sterically difficult to replace the I– in SO2-I –. Table 17. Replacement Reaction of SO2-I – + Cl– in MeCN-dmso Mixture at 298°K dmso% 0.00(70%) 25(53%) 50(32%) 75(20%) 100(17%) 5. Structure of SO2-X- The structures of SO2-X - (X- = F, Cl, Br & I) (Latajka et al., 1995) were investigated using ab initio method at the electron correlation level with effective core potential double zeta valence basis set with polarization functions. It has been found that the minimum on the potential energy surface corresponds to the CS structure of the complexes whereas the planar C2V structure is the transition state for the inversion process. The stability of SO2-X - complexes fall in the order SO2-F -> SO2-Cl -> SO2-Br -> SO2-I -. Careful studies of the potential energy surfaces of the complexes clearly indicates that only one structure corresponds to the minimum. The X- is in close contact with the S atom having a positive total charge. The intermolecular distance R(S …..X) increases in this series with the increase of the atomic number of X-. The shortest intermolecular distance is noted for SO2-F - complex. Replacement of F- by Cl- increases the distance by 2.61Å and successive replacements by Br- and I- ligands cause an increase in R(S….O) value by about 0.3Å. The complex formation slightly distorts the SO2 subunit. The SO bond is stretched while the OSO bond angle slightly decreases in comparison with the values for the isolated SO2 subunit. Perturbation of the geometrical structure of the SO2 subunit is essentially pronounced with the F - ligands. 50 WEAK COMPLEXES OF SULFUR COMPOUNDS 51 SO2-X - complexes are not planar and have CS symmetry. The nonplanarity of the complex is denoted by α which measures the angle between the C2 axis of the SO2 subunit and the S….X axis. In the series of SO2-X - complexes the angle decreases from 71o for the SO2-F - complex to 61o for the SO2-I - complex which makes this complex less pyramidal. Since in this series the intermolecular distance has increased from 1.93Å for the SO2-F – complex to 3.25 Å for the SO2-I – complex, the dipole-charge interaction becomes most important and in consequence the SO2-I – is less pyramidal. In conclusion we may report that SO2-F – has a rigid pyramidal structure while SO2-I – is the least pyramidal in the above series of SO2-X – complex species. 6. Conclusion The formation of the complex species SO2-X -, SOCl2-X -, SO2Cl2-X - and SOBr2-X - where (X- = Cl-, Br-, I- and SCN-) was reported during the last thirty years. The above complexes are weak of charge transfer nature and can not be isolated from solution but can be detected by spectrophotometry. 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Chem., 2(4): 149-154. SALAMA, S. B. and WASIF, S. 1994. Formation of iodide ions with some arenesulfonyl chloride, Asian Journal of Chemistry, 6(2): 381-388. SANDHU, S. S. and SINGH, A. 1960. Complexes of solvacids and ansolvo bases, J. Indian Chem. Soc. 37: 329-333. SANDHU, S. and SINGH, A. 1962. Mechanism of the reactions of solvo-bases and solvo-acids, J. Indian Chem. Soc . 39: 589-594. WEAK COMPLEXES OF SULFUR COMPOUNDS SEEL, F. and RIEHL, L. 1955a. Fluorosulfinates, Z. anorg. Chem., 282: 293-306. SEEL, F. and RIEHL, L. 1955b. Fluorosulfinates, Angew. Chem., 67: 32-33. WADDINGTON, T. C. 1965. Non-Aqueous Solvent Systems, Academic Press, London. WASIF, S. Unpublished Work. WITECKOWA, S. and WITOK, T. 1955. Reaction between iodine and sulfur dioxide, Zeszyty nauk, Politech lodz.Chem.Spoz., 1: 73-86. Received 12 December 1999 Accepted 24 June 2000 53 Salama B. Salama* and Saad Wasif** ÏÑÇÓÉ ááãÚÞÏÇÊ Èíä ãÑßÈÇÊ ÇáßÈÑí� ÓáÇãÉ ÈØÑÓ ÓáÇãÉ æ ÓÚÏ æÇÕÝ CONTENTS 1. Introduction 32 4.1 Correlation of Stability Constants with Ligand Replacement 48 Conclusion 51 1. Introduction Figure 3. MeCN-dmso Adduct Table 13. Stability Constants of Adducts of MeCN and dmso with Sulfur Compounds at 298(K. 4.1 Correlation of Stability Constants with Ligand Replacement Conclusion