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Comparison of H2S and H2O, calculated using the Restricted Hartree-Fock

wave function (RHF) using the 6-31G** basis set.

 

Comparison of H2O and H2S

Aqueous H2S and SO2

link; The water molecule

link; Aqueous alcohol

link; Aqueous ammonia

V Hydrogen sulfide

V Aqueous H2S

V Sulfur dioxide

V Aqueous SO2

V Sulfuric acid

 

Hydrogen sulfide

Hydrogen sulfide a (H2S) is a colorless, dense, flammable, reactive, and corrosive gas (boiling point, -60 °C; melting point, -85.5 °C, triple point, -85.49 °C, 23.1 kPa; critical temperature 100.25 °C, critical pressure 8.96 MPa, critical density 347.62 kg ˣ m-3; liquid density -60 °C, 949.0 kg ˣ m-3; surface tension -60.2 °C, 25. 43 kN ˣ m-1) r with a characteristic unpleasant and highly toxic odor of rotten eggs even at very low concentrations (>~0.5 ppb). Dangerously, at very high concentrations, its smell is undetectable. It is naturally produced and utilized in living metabolism, f in the putrefaction of organic matter in water, and from volcanic gases. It is found in amounts of about 5 μg ˣ m-3 (5 ppb) in the air. It is almost non-polar with dipole moment 0.97 D in the gas phase and 1.25 D in the aqueous phase [3846]. In aqueous solution, it acts as a weak acid. Its structure (see above) and properties are very different from H2O, in contrast to the apparent similarity, as an analog, at first glance. H2S is very soluble in common organic solvents.

Aqueous H2S

H2S is only slightly soluble in water with a minimum solubility at about 200 °C (1 atm, mole fraction (H2S + SH- + S2-) = 0.0005; mole fraction at 0 °C and 25 °C being 0.00375 and 0.0019 respectively); the solubility behavior being similar to other non-polar gases. It behaves as a weak acid. g The (disputed c) high pKa2 means that the concentration of S2- in aqueous solution is generally negligible.

 

                                 H2S + H2O equilibrium arrows  SH- + H3O+              pKa1 = 6.98 (25 °C) b

                                  SH- + H2O equilibrium arrows  S2- + H3O+              pKa2 = ~17 (25 °C) c

 

The pKa1 drops to a minimum of about 6.5 at about 100 °C. e H2S is a weak reducing agent. It forms colorless solutions that turn yellow with the formation of insoluble sulfur.

 

                             S0 + 2H+ + 2e- equilibrium arrows H2S                                 E°' = +0.17 V d              

                        S0 + H2O + 2e- equilibrium arrows HS- + OH-                      E°' = -0.478 V [70]

                             S0 + 2e- equilibrium arrows S2-                                  E°' = -0.508 V [70]

                 2 H2S + O2 -> 2 S0 + 2 H2O                          slow (pH <~ 6)

        2 HS-  + O2 + 2H+ -> 2 S0 + 2 H2O                       slow (pH >~ 6)

                          H2S + 2 O2 -> HSO4-    + H+                          very slow                

 

Comparison of H2S and H2O dimers, calculated using the Restricted

Hartree-Fock wave function (RHF) using the 6-31G** basis set.

Comparison of H2S and H2O dimers, calculated using the Restricted 

Hartree-Fock wave function (RHF) using the 6-31G** basis set.

 

 

The low atomic charges (see top right), large molecular dimensions, and low polarity of H2S mean that H2S - H2O interactions, including their very weak hydrogen bonds, are far smaller than H2O - H2O interactions. This is due to the increased dispersive interactions being more than compensated by the very decreased electrostatic interactions. The in vacuo dimers are given on the right with interactions:

 

H2O···H2O ≫ HSH···OH2 > HOH···SH2 ≫ H2S···H2S

 

In H2S solutions the water···water interactions are far stronger than the interactions involving H2S so that the mixed clusters are negligible with H2S molecules generally excluded from the water hydrogen bonding network and with the H2S preferring a dodecahedral water clathrate environment [3846].

 

Even if single hydrogen bonds transiently form between H2S and H2O, the H2S molecule cannot sustain any further hydrogen bonds whereas the H2O molecules prefer having four hydrogen bonds.

 

Dissolved H2S acts as a powerful surfactant lying at the surface with the sulfur atom pointed away from the bulk water [3846].

 


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Sulfur dioxide

 

Sulfur dioxide

Sulfur dioxide

Sulfur dioxide (SO2) is a toxic pollutant gas (boiling point, 101.3 kPa, -10 °C; melting point, -75.5 °C, critical temperature 157.5 °C, critical pressure 7.88 MPa, critical density 525 kg ˣ m-3) s released mostly from the combustion of fuels, particularly coal and coke (~20 Gkg ˣ yr-1), and oil and gas (~9 Gkg ˣ yr-1), but also naturally from volcanoes (~20 Gkg ˣ yr-1). The greatest polluting countries are India, Russia, and China. It is a colorless, dense, non-flammable, and reactive gas with a nasty irritating smell (boiling point, -10 °C; melting point, -72 °C, critical temperature 430.7 K, critical pressure 7.78 MPa). Its atmospheric concentration varies between good (1 ppb) and very poor (0.2 ppm) levels and is about four times that of hydrogen sulfide (the other significant atmospheric S-compound). It is responsible for 'acid rain'. Sulfur dioxide is primarily manufactured for sulfuric acid production.

 

Sulfur dioxide is formed by the burning of sulfur or materials that contain sulfur,

 

                                                   S2 (g) + 2 O2 (g) → 2 SO2 (g)                               ΔH = -723 kJ ˣ mol-1 at 298 °C,

2 H2S  (g) + 3 O2  (g) → 2 H2O + 2 SO2 (g)              

 

In the atmosphere, the following reactions have been proposed [3985],

 

1SO2 (g, singlet) + UV → 3SO2 (g, triplet)

3SO2 (g, triplet) + H2O (g) → OH• (g) + HOSO• (g)

1SO2(g) + OH• (g) → HOSO2• (g)
HOSO2• (g) + O2 (g) → SO3 (g) + HOO• (g)
SO3 (g) + H2O (g) → H2SO4 (g)

 

Sulfite results by the action of an aqueous base on sulfur dioxide,

 

SO2 + NaOH → NaHSO3

SO2 + 2 NaOH → Na2SO3 + H2O

 

Sulfuric acid is made industrially by further oxidation by the 'contact process' utilizing  vanadium(V) oxide (V2O5) catalyst.

 

2 SO2 + 2 H2O + O2 → 2 H2SO4

 

The ice snow pack has been shown to be a highly efficient sink for atmospheric SO3 depletion, especially in winter [3996].

Aqueous sulfur dioxide

 

The bisulfite ion has an equilibrium structure, Keq = 0.2 o

The aqueous chemistry of the oxides of sulfur and related acids is very complex, involves at least eight sulfur oxidation states, and is not restricted by the 'octet rule'. Sulfur dioxide (SO2) and sulfur trioxide (SO3) are the acid anhydrides of sulfurous acid (H2SO3) and sulfuric acid (H2SO4) respectively and are both very soluble in water. Sulfurous acid cannot be isolated from aqueous solution but does form bisulfites and is stable in the gas phase. p

 

SO2(g) -> SO2(aq) 

SO2(aq) + H2O ⇌ SO2.H2O

SO2.H2O + H2O ⇌ HSO3 + H3O+

 

H2SO3 is a strong acid (pKa = 1.86). HSO3 is weakly acidic (pKa  = 6.97) with the equilibrium above lying mostly towards hydrates SO2 (SO2.H2O, Keq ≪ 10−9). Raman spectra of solutions of  sulfur dioxide in water show only signals due to the SO2 molecule and the bisulfite ion. Because of its low concentration, the bisulfite dissociates to the  sulfite ion, SO32−.

 

SO2.H2O, calculated using the Restricted

Hartree-Fock wave function (RHF) using the 6-31G** basis set

SO2.H2O, calculated using the Restricted 

Hartree-Fock wave function (RHF) using the 6-31G** basis set

 

HSO3 + H2O -> SO32− + H3O+

 

Many bisulfites form metabisulfites (disulfites) on drying 

 

2 HSO3 ⇌ S2O52− + H2O

 

The photochemistry of SO2 at the air-water interface of water droplets leads to the formation of the strongly acidic (pKa  = -1) HOSO radicals that may play an important role in acid rain formation [3978].

 

Putative disulfurous acid; like sulfurous acid

is another phantom acid

 

Putative disulfurous acid

Many oxy-sulfur anions contain expandable S-S bonds (1.8 - 3.0 Å) that are flexible with bond angles between 90° and 180° and dihedral angles between 0° and 180°. n

 

The metabisulfite anions consists of an SO2 group linked to an SO3 group, with an S-S bond (2.22 Å m) and the negative charges localized on both ends. Such structures can exist as both [2OSOSO2]2− (see above right l) and [3OS-SO2]2− forms, with [3OS-SO2]2− found in crystals but with inequivalent ions. j The thermal decompositions of the sodium sulfur oxo-salts have been examined. k

 

Redox chemistry of sulfur showing some of the oxidation states h

The redox chemistry of sulfur

 

Peroxodisulfate (S2O82−, persulfate) is a very strong oxidizing agent, i used as a polymerization initiator. It can be serially reduced, producing the intermediates shown above. It decomposes to give sulfate radical anion or hydrogen peroxide,

 

S2O82−-> 2 SO4•−

S2O82−+ 2H2O → 2HSO4 + H2O2

 

The strongest reducing agent is dithionate followed by sufite.

 

Sulfurous (left) and sulfuric acids, calculated using the Restricted Hartree-Fock wave function (RHF) using the 6-31G** basis set

 

Sulfurous and sulfuric acids, calculated using the Restricted Hartree-Fock 

wave function (RHF) using the 6-31G** basis set

 

There are two possible forms of sulfurous acid (far left and middle, above). Modeling does give a stable (if more energetic) structure with an S-H bond similar to that in the bisulfite ion equilibrium (see middle above). Sulphuric acid (right above), but not sulfurous acid, is stable in aqueous solutions.

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Sulfuric acid

Sulfuric acid (see above right, (boiling point, 279.6 °C where it slowly decomposes; melting point, -10.4 °C; density 1835.6 kg ˣ L-1, 15 °C) t is a clear, polar (dielectric constant = 101 25 °C q), colorless, odorless, highly corrosive, and slightly viscous  liquid. It is very hygroscopic, miscible with water (highly exothermic) and strongly oxidizing and dehydrating, for example, removing water from sugars (charring) converting them to carbon and steam and producing a low density, solid, black substance including gas bubbles. Sulfuric acid is widely used in the chemical industry and in lead-acid batteries. Concentrated H2SO4 is usually 98.3 % by mass (18.4 mol ˣ L-1) and lead battery acid is about 30 % by mass (4.5 mol ˣ L-1). Pure sulfuric acid is ionized to a small extent,

 

H2SO4 + H2SO4 equilibrium arrows HSO4 + H3SO4+  

H2SO4 + H2SO4 equilibrium arrows H2S2O7 + H3O+    

 

Its acid dissociation constants in water are, (see also for Hammett description)

 

                                          H2SO4 + H2O -> HSO4 + H3O+                                            pKa1  = -6.4

                                           HSO4 + H2O -> SO42− + H3O+                                                            pKa2  = +1.98

 

Sulfuric acid has high electrical conductivity due to its autoprotolysis,

 

                                             H2SO4 + H2SO4 -> HSO4 + H3SO4+                                                          pK  = 3.6

 

This is about 1010 greater than the ionization of water (Kw = 10-14)

 

Disulfuric acid

 

Disulfuric acid

 

 

Pure sulfuric acid reacts with sulfur trioxide to form oleum (SO3)y.H2SO4; y = 0 - 1, also called fuming sulfuric acid, with the fumes mostly SO3; a mixture of sulfuric and disulfuric acids. If completely converted to H2S2O7 (i.e, y = 1), it is also called disulfuric acid or pyrosulfuric acid; melting point 36 °C) that reacts with water to reform H2SO4, This is the key final step in the 'Contact Process' for the production of sulfuric acid.

 

H2SO4 (anhydrous, pure) + SO3 -> H2S2O7                                                        

H2S2O7 + H2O -> 2 H2SO4                                                        

 

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Footnotes

a The spellings 'hydrogen sulphide', sulphur dioxide' and 'sulphur' (etc.) are not recommended [IUPAC] but still acceptable. Nowadays the 'ph' in 'sulphur' compounds should always be replaced by 'f'. [Back]

 

b F. J Millero, The thermodynamics and kinetics of the hydrogen sulfide system in natural waters, Marine Chemistry, 18 (1986) 121-147. [Back]

 

c W. Giggenbach, Optical spectra of highly alkaline sulfide solutions and the second dissociation constant of hydrogen sulfide. Inorganic Chemistry, 10 (1971) 1333-1338; this high value is disputed with others claiming a value around pKa2 = 14, R. S., Ramachandra and L. G. Hepler, Equilibrium constants and thermodynamics of ionization of aqueous hydrogen sulfide. Hydrometallurgy, 2 (1977) 293-299. [Back]

 

d O. Kabil and R. Banerjee, Redox biochemistry of hydrogen sulfide, Journal of Biological Chemistry, 285 (2010) 21903-21907. [Back]

 

e J. A. Barbero, K. G. Mccurdy and P. R. Tremaine, Apparent molal heat capacities and volumes of aqueous hydrogen sulfide and sodium hydrogen sulfide near 25°C: the temperature dependence of H,S ionization, Canadian Journal of Chemistry, 60 (1982) 1872-1880. [Back]

 

f E. Cuevasanta, M. N. Möller and B. Alvarez, Biological chemistry of hydrogen sulfide and persulfides, Archives of Biochemistry and Biophysics, 617 (2017) 9-25. [Back]

 

g The pKas of H2Se and H2Te are 3.7 and 2.7 respectively [3888b], with that for H2Te thirteen orders of magnitude greater than that for H2O. [Back]

 

h N. Wiborg, Inorganic Chemistry, Academic Press, San Diego, 2001, ISBN 0-12-352651-5. [Back]

 

i S. A. Shafiee, J. Aarons and H. H. Hamzah. , Review—Electroreduction of peroxodisulfate: A review of a complicated reaction, Journal of The Electrochemical Society, 165 (2018) H785-H798 [Back]

 

i K. L. Carter, T. A. Siddiquee, K. L. Murphy and D. W. Bennett, The surprisingly elusive crystal structure of sodium metabisulfite, Acta Crystallographica Structural Science, B60 (2004) 155-162. [Back]

 

k B. Jaszczak-Figiel and Z. Gontarz, Stages of thermal decomposition of sodium oxo-salts of sulfur, Journal of Thermal Analysis and Calorimetry, 96 (2009) 147-154. [Back]

 

l M. Abedi, M. Vahedpour, S. Farnia and H. Farrokhpour, Theoretical study on the structures, stabilities and electronic properties of S2O52− isomers in the gas and solution phases, Molecular Physics: An International Journal at the Interface Between Chemistry and Physics, 111 (2013) 581-588. [Back]

 

m B. I-C. Chen and Y. Wang, Reinvestigation of the structure of potassium pyrosulfite, K2S2O5, Acta Crystallographica, C40 (1984), 1780-1781. [Back]

 

n R. Steudel, Properties of sulfur-sulfur bonds, Angewandte Chemie, 14 (1975) 655-720. [Back]

 

o D. A. Hornert and R. E. Connick, Equilibrium quotient for the isomerization of bisulfite ion from HS03to S03H, Inorganic Chemistry, 25 (1986) 2414-2417. [Back]

 

p D. Siilzle, M. Verhoeven, J. K. Terlouw and H. Schwarz, Generation and characterization of sulfurous acid (H2SO3) and of its radical cation as stable species in the gas phase, Angewandte Chemie International Edition, 27(1988) 1533-1534; E. Bishenden and D. J. Donaldson, Ab initio Study of SO2+H2O, Journal of Physical Chemistry A, 102 (1998) 4638-4642. [Back]

 

q R. J. Gillespie and R. H. Cole, The dielectric constant of sulphuric acid, Transactions of the Faraday Society, 52 (1956) 1325-1331. [Back]

 

r F. Pouliquen, C. Blanc, E. Arretz, I. Labat, J. Nougayrede, G. Savin. R. Ivaldi, M. Nicolas, J. Fialaire, R. Millischer, C. Azema, L. Espagno, H. Hemmer, and J. Perrot, Hydroigen sulfide, Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim, Vol. 18 (2012) pp 429-449, doi: 10.1002/14356007.a13_467. [Back]

 

s H. Müller, Hydroigen sulfide, Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim, Vol. 35 (2012) pp 73-118, doi:10.1002/14356007.a25_569. [Back]

 

t H. Müller, Sulfuric acid and sulfur trioxide, Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim, Vol. 35 (2012) pp 141-211, doi:10.1002/14356007.a25_635. [Back]

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