Thermal decomposition or reduction of gypsum СaSО4.2H2O

Oxidative combustion of sulfides of metals (pyrites)

Oxide of sulfur (IV). Preparation.

In industry:

4FeS₂ + 11О₂ = 2Fe₂О₃ + 8SО₂

2. Incineration of native sulfur (in countries with rich deposits of elementary sulfur):

S + О₂ = SО₂; DН⁰₂₉₈ = -296.8 kJ/mol

2СaSО₄ 2СаО + 2SО₂ + О₂

СaSО₄ + С СаО + SО₂ + СО₂ (in cylinder furnaces)

In laboratory: 1. Using sulfites:

Na₂SО₃ + 2H₂SО₄ = 2NaHSО₄ + SО₂­ + H₂O,

2. Reduction of concentrated H₂SО₄ by low active metals (e.g. with copper):

Cu + 2H₂SО_4(conc.) = CuSО₄ + SО₂­ + 2H₂O

Structure. SО₂ has a bent structure (S–O bond length 143.2 pm, O–S–O valence angle 119.5◦ in the gas phase),

that allows to ascribe for the sulfur atom sp²- hybridization (theoretical angle 120^о). Bonds of the central sulfur atom with every oxygen atom have equal energy (497.5 kJ/mol) and length. According to the Valence Bond method, the atoms in SО₂ are bound by two s-bonds and by a delocalized three-centered p-bond.

The state of hybridization of sulfur orbitals in SО₂ corresponds to the structure draft:

The formation of s-bonds in SО₂ molecule is a result of overlap of two sulfur monoelectronic hybrid orbitals with 2р_х monoelectronic orbitals of two oxygen atoms. To achieve an octet electronic configuration 2р_z -orbital of the oxygen atom accepts one electron of unhybridized 3р_z -orbital of sulfur (electron transition is shown on the chart below with an arrow):

This arrangement results in the formation of a p-bond between sulfur and the second atom of oxygen. The resonance structure of this overlapping state of atomic orbitals is shown near the chart. However, this structure has no similarity to actual SО₂ structure since S-О bonds are identical in the real molecule. The same is the description of validity of the opposite resonance structure of overlapping orbitals.

According to the resonance theory a real molecule is an intermediate state between the limiting cases of resonance structures, and its structure is the result of imposition (superposition) of all variants of resonance structures. Therefore, the real structure of SО₂ corresponds to the first structure chart, where together with s-bonds delocalised three-centred p-bond is present.

Properties. SO₂ is a poisonous, colourless gas with a sharp smell; it condenses (−10 ◦C) to give a colourless liquid and ultimately (−76 ◦C) white crystals. SO₂ is an acid oxide, its molecule is polar (m = 1.59 D = 0.53•10^-29 C•m).

SO₂ demonstrates reducing properties (that are more characteristic):

SO₂ + Br₂ + 2H₂O = H₂SO₄ + 2HBr,

and oxidizing properties:

SO₂ + 2H₂S = 3S + 2H₂O.

It is well soluble in water (36 volumes SO₂ in one volume of water at 20 ^oC) with the formation of hydrates of variable composition SO₂∙nH₂O. Some dissolved molecules react with water and form sulfurous acid, Н₂SO₃.

SULFURUOUS ACID, H₂SO₃.

SO₂ + H₂O Н₂SO₃

Н₂SO₃ exists only in diluted solutions as a result of reversibility of this reaction. Its equilibrium is significantly shifted to the left hand side. On being heated, the solubility of SO₂ diminishes and the equilibrium is even more displaced toward reagents. A reaction of acids with salts of sulfurous acid causes evolving SO₂ because of Н₂SO₃ instability; therefore, Н₂SO₃ cannot be isolated in the free state.

Structure of sulfite-ion, SO₃^2-. The structure is pyramidal with the atom of sulfur on the top.

The value of valence angles (Ð 105⁰) is the evidence of sp³-hybridization. All bonds here are equivalent, and their length (151 pm) is somewhat longer of SO₂ bonds (143 pm). At bonds formation, every atom surrounds electron octet that provides SO₃^2- ion stability. The identity of S—O bonds suggests equivalent distribution of two negative charges between three atoms of oxygen.

Diminishment of valence angle in the sulfite ion (105^o) in comparison with tetrahedral (109^o28') is caused by the electrostatic repulsion between the unshared electron pair and the bond electrons.

Since the unshared electron pair of sulfur is located on spatially directed hybrid orbitals, the sulfite ion is an active donor of electron pair and its transformations into tetrahedral ions НSO₃^- and SO₄^2- can be realised easily.

Properties. Sulfurous acid is a dibasic acid of intermediate strength:

Н₂SO₃ Н⁺ + НSO₃^-; K₁ = 2∙10^-2;

Н₂SO₃^- Н⁺ + НSO₃^2-; K₂ = 6∙10^-8.

It forms two types of salts: sulfites and hydrogen sulfites:

NaOH + SO₂ = Na₂SO₃

Na₂SO₃ + SO_2(excess) + H₂O = 2NaНSO₃

The hydrogen sulfite-ion, НSO₃^-, has two tautomeric forms:

The first structure corresponds to generally neutral salts, the second structure is close to the salts of some low active metals.

Only sulfites of alkali metals and hydrogen sulfites are well soluble. Only solid sulfites can be prepared. Solid hydrogen sulfites are known for alkali metals only.

Salts of Н₂SO₃ are partially hydrolysed by anion, their solutions have alkaline reaction.

The sulfite and bisulfite anions are both moderately strong reducing agents.

· Н₂SO₃ and its salts are gradually oxidized even by air oxygen in aqueous solutions:

2Na₂SO₃ + O₂ = 2Na₂SO₄

· S trong oxidants oxidize SO₃^2- practically instantly:

Н₂S⁺⁴O₃ + Hal₂⁰ + Н₂О = Н₂S⁺⁶O₄ + 2НHal^1- (Hal₂ = Cl₂, Br₂, I₂)

5Na₂SO₃ + 2KMnO₄ + 3H₂SO₄= 5Na₂SO₄ + 2MnSO₄ + K₂SO₄ + 3H₂O

Na₂SO₃ + H₂O₂ = Na₂SO₄ + H₂O

Н₂SO₃ demonstrates oxidizing properties at the attack of strong reductants:

Н₂SO₃ + 2H₂S = 3S + 3H₂O

Being heated, sulfites disproportionate:

4Nа₂S⁺⁴O₃ = Na₂S^-2 + 3Na₂S⁺⁶O₄.

POLYSULFURUOUS ACIDS. Zinc dithionite is formed at SO₂ reduction with zinc dust in water:

Zn + 2SO₂ = ZnS₂O₄.

Dithionate is also obtained by the reaction:

2NаHSO₃ + Zn + H₂SO₃ = Na₂S₂O₄ + ZnSO₃ + 2H₂O.

Dithionites are salts of dibasic dithionic acid, Н₂S₂O₄. It is unstable and simultaneously strong acid (K₁ = 5∙10^-1, K₂ = 4∙10^-3) in aqueous solutions. It has a symmetric structure and forms tautomers:

Н₂S₂O₄ is one of the strongest reductants. In alkaline medium, it reduces even perchlorates. Salts are quickly converted into hydrogen sulfites in the presence of air oxygen:

2Na₂S₂O₄ + O₂ + 2H₂O = 4NаHSO₃.

Hydrogen sulfites lose water and transform into disulfite (pyrosulfite) which are salts of the disulfurous (pyrosulfurous) acid Н₂S₂О₅ at moderate heating, for example, during crystallisation from solutions, unknown in a free state:

2NaHSO₃ = Na₂S₂О₅ + Н₂O.

A more convenient method to prepare disulfites is presented below:

2SO₂ + 2NaHCO₃ = Na₂S₂О₅ + 2CO₂ + H₂O.

This acid has the following structure:

Dissolution of disulfites in water results in reverse reaction. Na₂S₂О₅ salt is a conservator of moist corn and vine.

Oxygen - element(IV) compounds. EO₂ oxides can be prepared from simple substances at combustion in air or oxygen atmosphere:

E + O₂ = EO₂.

Unlike SO₂, EO₂ oxides are polymer solids at STP: SeO₂ (t_subl. = 315 °), TeO₂ (m.p. = 733 °), PoO₂ (m.p. = 885 °). Actually, SeO₂ has a chain structure:

This is a tetrahedron where Se is in the state of sp³- hybridisation.

In the series SeO₂—TeO₂—PoO₂ acid properties are weakened and basic properties are strengthened. The reason for this phenomenon is the growth of metallic properties of elements and also increase of ionic mechanism contribution in E—O bond. Certainly, SeO₂, like SO₂, is dissolved in water and alkalis easily with the formation of selenous acid, H₂SeO₃, which is similar to sulfurous acid:

SeO₂ +H₂O H₂SeO₃

SeO₂ + 2NaON = Na₂SeO₃ +H₂O

TeO₂ already shows amphoteric properties. It is not soluble in water, but like SeO₂, is soluble in alkali solutions, forming a colourless telluride:

TeO₂ + 2NaON = Na₂TeO₃ +H₂O (TeO₂ has acidic properties),

and in solutions of strong acids to form salts:

TeO₂ + HI = TeI₄ + 2H₂O (TeO₂ has basic behaviour),

TeI₄ + 2HI H₂[TeI₆].

PoO₂ is a basic oxide. That is why PoO₂ has no interaction with solutions of alkali (the reaction takes place only with solid alkalis in the molten state), although gives salts with acids:

PoO₂ + 2H₂SO₄ = Po(SO₄) ₂ + 2H₂O.

Due to the action of strong acids on selenites and tellurites selenous and tellurious acids are formed. Unlike H₂SO₃, they can be obtained in the free state. H₂SeO₃ is a white hygroscopic solid, which easily loses water:

H₂SeO₃ SeO₂ + H₂O

H₂TeO₃ has the ability to form polymers (its composition is TeO₂∙ h H₂O). Low soluble TeO₂∙ h H₂O can be deposited from concentrated solutions.

H₂SeO₃ and H₂TeO₃ are easily obtained during elemental Se and Te oxidation by concentrated HNO₃.

In the series of acids H₂TeO₃ — H₂SeO₃ — H₂SO₃ the acid strength increases (K₁ = 3^.10^-6, 2^.10^-3, 2^.10^-2). First of all, this behaviour is defined by the growth of chalcogen’s EN in this direction that leads to increased displacement of electrons in OH-groups to the oxygen atom and increases its polarity:

The stronger polarisation of H—O bond facilitates dissociation of acids under the influence of polar water molecules.

Elements display intermediate oxidation state in compounds Se⁴⁺ and Te⁺⁴, and these substances can be reductants and also oxidants in redox reactions, but unlike S⁴⁺ derivatives, oxidising properties are more typical of them, as evidenced by comparing E^o values:

H₂SO₃ + 4H⁺ + 4e^- = S + 3H₂O; E ° = 0.45 V

H₂SeO₃ + 4H⁺ + 4e^- = Se + 3H₂O; E ° = 0.74 V (a consequence of secondary periodicity)

H₂TeO₃ + 4H⁺ + 4e^- = Te + 3H₂O; E ° = 0.59 V

Therefore,

H₂SO₃ + HI ® no reaction

H₂SeO₃ + 4HI = 2I₂ + Se + 3H₂O

but as mentioned above: TeO₂ + 4HI = TeI₄ + 2H₂O

Selenous acid oxidises SO₂:

H₂SeO₃ + 2SO₂ + H₂O = Se + 2H₂SO₄

These acids can also have reducing properties (with very strong oxidants):

if H₂SO₃ + Br₂ + H₂O = H₂SO₄ + 2HBr

then

5H₂SeO₃ + 2HClO₃ = 5H₂SeO₄ + Cl₂ + H₂O

Na₂SeO₃ + Cl₂ + NaOH = Na₂SeO₄ + NaCl + H₂O

5H₂SeO₃ + 2KMnO₄ + 3H₂SO₄ = 5H₂SeO₄ + 2MnSO₄ + K₂SO₄ + 3H₂O

and

TeO₂ + H₂O₂ + 2H₂O = H₆TeO₆ ortho-telluric acid

Sulfur (VI) oxide

Production in industry. Catalytic oxidation of SO₂ by air oxygen takes place in the presence of catalyst, V₂O₅ (440 ^oC):

2SO₂ + O₂ 2SO₃; DН⁰₂₉₈ =-184.4 kJ/mol

In the laboratory it is commonly prepared by the reaction between sulfur dioxide and oxygen in the presence of platinum catalyst and high temperature:

2SO₂ + O₂⁰ 2SO₃

Initially, sulfur trioxide was prepared by heating iron(III) sulfate:

Fe₂(SO₄)₃ = Fe₂O₃ + 3SO₃

Itis also obtained by the dehydration of concentrated sulfuric acid with phosphorus(V) oxide:

2H₂SO₄ + P₄O₁₀ = 4HPO₃ + 2SO₃

Structure. SO₃ is a nonpolar molecule that has a plane structure with the valence angle 120^o (sp²- hybridisation of valence orbitals of sulfur) and equivalent bonds due to delocalised fourcentred p-bond (like SO₂):

Polymorphic modifications of SO₃. Molecules SO₃ exist only in the gaseous phase. In the solid state, it is a polymer with (SO₃)_n chains:

This solid polymorph is referred to as a-SO₃. It looks like transparent mass similar to ice; m.p. = 16.9 ^oC.

b - SO₃. It forms when a- SO₃ absorbs water. This polymer has lower molar mass than a-SO₃. It does not melt, but sublime at 50^oC. Externally, it is similar to asbestos.

g - SO₃ is the product of condensation of SO₃ vapours consisting of cyclic trimer (SO₃)₃ where S has sp³ -hybridisation state:

Properties. This acid oxide reacts with basic oxides, and hydroxides. Moreover, what is most important: SO₃ is an anhydride of sulfuric acid. It reacts with water rapidly:

SO₃ + Н₂O = Н₂SO₄; DН⁰₂₉₈ = -89.2 kJ/mol

Sulfuric acid

Sulfuric acid is probably the most important chemical substance which is not found in Nature. The total world consumption is about 25 000 000 tons per year.

Production in industry:

1. Contact method. Catalytic oxidation of SO₂ by air oxygen (440 ^oC, in the presence of V₂O₅ catalyst):

2SO₂ + O₂ 2SO₃; DН⁰₂₉₈ =-184.4 kJ/mol

SO₃ is absorbed by concentrated H₂SO₄. Usage of water is not efficient because gaseous SO₃ reacts first of all with water steam and then predominately H₂SO₄. Fog is the product of this interaction. The solution of SO₃ in H₂SO₄ has a technical name oleum that emphasizes its high viscosity (oleum - Lat. Oil).

H₂SO₄ dissolves SO₃ in any proportions, so the composition of oleum is H₂SO₄·nSO₃. Oleum contains several polysulfuric acids:

The content of SO₃ in oleum is 20-65%. To produce acid, oleum is mixed with H₂SO₄ having the necessary amount of water. Under the influence of water S—O—S bonds of polysulfuric acid are destroyed and converted into H₂SO₄ acid of required concentration.

H₂SO₄ produced by this method has high purity and can be of any concentration.

2. Lead Chamber process. It was called so because the chamber is lined with lead, which resists the action of cold sulfuric acid; the homogeneous catalyst is nitrogen oxide. The method was applied for the first time in the middle of 18^th century.

It was the most important method of industrial production of H₂SO₄ before the modern contact method was developed. Its essence is SO₂ oxidation by nitrogen oxide, NO_2, in the presence of water:

SO₂ + NO₂ + H₂O = Н₂SO₄ + NO

Nitrogen(II) oxide (NO) is oxidised again into the initial NO₂ and returned to the reaction:
2NO + О₂ = 2NO₂.

The process is carried out in special towers (see scheme below). This method gives about 76% H₂SO₄ used in the manufacture of fertilizers:

Structure. Н₂SO₄ and НSO₄^- -ion are the distorted tetrahedrons, with sulfur atom in the centre of a polyhedron. SO₄^2--ion has the shape of regular tetrahedron where all distances and angles are equivalent. Orbitals of sulfur have the state of sp³ -hybridisation:

S—O bonds are very strong because of the additional p-bonding: displacement of unshared electron pairs of O atoms to the vacant 3d-orbital of S. The equal S—O bonds lengths is a consequence of delocalisation of p-electron density and the negative ion charge.

Properties. Anhydrous (100%) Н₂SO₄ under the normal conditions is a colourless oily liquid (m.p. = 10.4 ^oC). Н₂SO₄ molecules have intermolecular H-bonds with one another in the liquid and solid state either.

Н₂SO₄ azeotrope [i] contains Н₂SO₄ 98.3 % wt. and 1.7% by weight of water, b.p. = 338.8 ^oC.

Н₂SO₄ is a strong dibasic acid in aqueous solutions:

Н₂SO₄ Н⁺ + НSO₄^-, K₁= 1·10³

НSO₄^- Н⁺ + SO₄^2-, K₂= 1,2·10^-2

Due to the large differences in dissociation constants, Н₂SO₄ forms sulfates and hydrogen sulfates. The scheme of their mutual conversion is shown below:

H₂SO₄

K₂SO₄ KHSO₄

KOH

Being dissolved Н₂SO₄ in water, a lot of heat liberates owing to hydrates formation. Therefore, concentrated Н₂SO₄ and water must be mixed with caution. To prevent spray of liquid, pour Н₂SO₄ (as a heavier component) into water, not vice versa. Formation of very stable hydrates explains Н₂SO₄ water absorption properties.

Sulfuric acid is a strong oxidising agent. Its reaction with metals depends on the concentration and activity of metal:

1. Diluted acid (oxidant is hydrogen ion H⁺). Therefore, it dissolves (oxidises) only metals that are situated before H in the electromotive series of metals:

Zn⁰ + Н₂SO_4(dil) = Zn⁺²SO₄ + Н₂⁰­

Zn + 2H⁺ = Zn²⁺ + H₂

2. Concentrated Н₂SO₄ (oxidant is sulfur(VI)):

· Sulfur is reduced to SO₂ at the interaction with nonactive metals (Cu, Ag, Hg):

Cu⁰ + 2Н₂SO_4(conc.) = CuSO₄ + SO₂­ + 2H₂O

· Sulfur is reduced to the elemental S or H₂S (a mixture of reduction products often appears)at the interaction with active metals:

3Zn + 4Н₂SO_4(conc.) = 3ZnSO₄ + S¯ + 4H₂O

4Zn + 5Н₂SO_4(conc.) = 4ZnSO₄ + H₂S­ + 2H₂O

It can oxidise other simple substances and compounds:

2HBr + Н₂SO_4(conc.) = Br₂ + SO₂­ + 2H₂O

8HI + Н₂SO_4(conc.) = 4I₂ + H₂S­ + 4H₂O

C + 2Н₂SO_4(conc.) = CO₂­ + 2SO₂­ + 2H₂O

S + 2Н₂SO_4(conc.) = 3SO₂­ + 2H₂O

Н₂SO₄ is a very acidic solvent. CH₃COOH and HNO₃ behave as bases in Н₂SO₄. HClO₄ is one of the strongest acids, but it is a weak one in Н₂SO₄ medium:

HClO₄ + Н₂SO₄ = Н₃SO₄⁺ + ClO₄^-

The only acid that is stronger than Н₂SO₄ in its medium is hydrogen tetra(hydrogen sulfato) borate, H[B(HSO₄)₄]. It has not been obtained in the free state, but can be isolated from its solutions in Н₂SO₄:

B(OH)₃ + 6Н₂SO₄ = [B(HSO₄)₄]^- + 3H₃O⁺ + 2HSO₄^-

Н₂SO₄ salts. Sulfates are very abundant. Most of them are well-soluble in water. The low soluble sulfates are CaSO₄, SrSO₄, BaSO₄, PbSO₄.

Solid hydrogen sulfates are obtained only for the most active metals (NaHSO₄, KHSO₄ etc.). Disulfates (pyrosulfates) are salts of disulfuric (pyrosulfuric) acid obtained at heating of hydrogen sulfates:

2NaНSO₄ = Na₂S₂O₇ + Н₂O,

in turn, when heated stronger, the following reaction proceeds:

Na₂S₂O₇ = Na₂SO₄+ SO₃

Under the influence of water, disulfates are transformed again into hydrogen sulfates.

Chemical properties of sulfuric acid can be summarised in the scheme:

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