H₂S H⁺ + HS^-; K₁ = 8.7^.10^-8

HS^- H⁺ + S^2-; K₂ = 10^-14

The neutral salts of H₂S are sulfides, acidic salts – hydrogen sulfides.

A great majority of sulfides are insoluble in water. The soluble include only the sulfides of alkali metals and NH₄⁺, and also hydrogen sulfides. Most of them exist only in solution.

Preparation of sulfides. Sulfides can be obtained in a number of different ways:

1. Reactions between simple substances (sulfur reacts with metals vigorously at t^o).

2. Action of H₂S on soluble salts of metals. This method sometimes is unsuitable. The reason is a huge difference in solubility products of sulfides, e.g.: SP_HgS = 4∙10^-53, SP_FeS = 3.7∙10^-19.

Such reactions are reversible in general, since one of initial substances - H₂S is a weak acid, and among products of reaction is MS - sparingly soluble salt.

The displacement of equilibrium of the double replacement reaction:

M²⁺ + H₂S MS + 2H⁺

depends on which product is the weakest electrolyte: H₂S (K_H2S = K₁∙K₂ = 8.7∙10^–22) or the sparingly soluble sulfide, MS.

Two variants are possible:

1). In case of virtually insoluble sulfides (SP_MS << K_H2S) reactions proceed to the end, for example:

Cu²⁺ + H₂S = CuS + 2H⁺ (SP_CuS = 4∙10^-38)

Hg²⁺ + H₂S = HgS + 2H⁺ (SP_HgS = 4∙10^-53)

These sulfides are insoluble in water and acids, their preparation is possible by H₂S action on the solutions of salts of such metals.

2. In case of SP_MS << K_H2S, formation of sulfides does not occur (the equilibrium of reaction is displaced to the left). Such sulfides (FeS, MnS, BaS) are easily dissolved in acids, for example:

FeS + 2HCl = FeCl₂ + H₂S

Sulfides of both types are insoluble in water and their preparation using soluble sulfides instead of weak electrolyte H₂S is a better alternative, for example:

FeSO₄ + Na₂S = FeS + Na₂SO₄

Under such conditions, the deepest proceeding of direct reaction is realised up to concentration of ions defined by SP_MS.

Different solubility of metal sulfides in water and in acidic solutions make up a basis of qualitative analysis of cations in analytical chemistry. Some of them (Na⁺, K⁺, NH₄⁺ and etc.) form soluble sulfides; the second group contains insoluble in water metal sulfides (Fe²⁺, Mn²⁺, Zn²⁺and etc.) but well soluble in acids; the third group metal sulfides of Cu²⁺, Pb²⁺, Hg²⁺ and etc. are insoluble neither in water nor in diluted acids.

3. Reduction of sulfates by carbon (for the sulfides of active metals)

Na₂SO₄ + 4C Na₂S + 4CO

Hydrolysis. In the solutions sulfides, as the salts of a weak acid, H₂S, demonstrate strong hydrolysis by anіon. The hydrolysis of soluble sulfides proceeds step by step and reversibly:

S^2- + HOH HS^- + OH^-

HS^- + HOH H₂S + OH^-

As at hydrolysis forms a surplus of OH^- ions, solutions of sulfides have an alkaline reaction (pH > 7).

The hydrolysis of sulfides of multicharged ions which are salts of very weak bases and acids proceeds completely to the end:

Al₂S₃ + 6H₂O = 2Al(OH)₃ + 3H₂S

Cr₂S₃ + 6H₂O = 2Cr(OH)₃ + 3H₂S

Therefore, their preparation by any exchange reaction is impossible since the product hydrolyses in aqueous solutions:

2AlCl₃ + 3Na₂S + 6H₂O = 2Al(OH)₃ + 3H₂S + 6NaCl

Cr₂(SO₄)₃ + 3Na₂S + 6H₂O = 2Cr(OH)₃ + 3H₂S + 3Na₂SO₄

Sulfides, as well as oxides, have basic, amphoteric or acid character. Interaction between them gives salts of thioacids, therefore the sulfides of non-metals are called thioanhydrides of corresponding acids, for example:

Na₂S + CS₂ = Na₂CS₃

Basic acid thiocarbonate is a thioanhydride of thiocarbonic acid (Н₂CS₃).

Acids which correspond to thiosalts are most frequently unstable and readily decompose:

2Na₃AsS₄ + 6HCl = As₂S₅ + 3H₂S + 6NaCl

H₂S and sulfides have the lowest oxidation state of sulfur (-2) and demonstrate only reduction properties, they belong to strong reductants.

At ignition H₂S burns in air:

2H₂S^-2 + 3О₂⁰ = 2S⁺⁴O₂^-2 + 2H₂O^-2

In solution it is slowly oxidized by the oxygen of air to elementary sulfur:

H₂S^-2 + O₂⁰ = 2S⁰ + 2H₂O^-2

In aqueous solutions, H₂S and sulfides are readily oxidized by halogens, KМnO₄, K₂Cr₂O₇ and other oxidants, for example:

Na₂S^-2 + I₂⁰ = 2S⁰ + 2NaI^-1

H₂S^-2 + 4Br₂⁰ + 4H₂O = H₂SO₄ + 8HBr

3H₂S^-2 + K₂Cr₂O₇ + 4H₂SO₄ = 3S⁰ + Cr₂⁺³(SO₄)₃ + K₂SO₄ + 2H₂O

5H₂S^-2 + 2KMnO₄ + 3H₂SO₄ = 5S⁰ + MnSO₄ + K₂SO₄ + 8H₂O

H₂S^-2 + 2Fe⁺³ = S⁰ + 2Fe⁺² + 2H⁺

Depending on conditions and oxidants the products of oxidation can be sulfur, SO₂ or H₂SO₄ . It can be represented by the chart:

TESTS FOR HYDROGEN SULFIDE

1. It’s unpleasant smell of rotten eggs.

2. The blackening of filter paper moistened with a soluble lead(II)salt (e.g. nitrate) due to the formation of lead(II) sulfide.

PolysulFides.

The sulfur ability for chains formation is observed in numerous polysulfides. They are prepared by interaction of sulfides of alkaline metals or NH₄⁺ with sulfur in concentrated solutions (water solutions of sulfides simply dissolve sulfur):

Na₂S^-2 + S⁰ = Na₂S₂^-1

Na₂S + (n-1)S = Na₂S_n

These reactions can be considered as an oxidation of sulfides by sulfur. They usually form the mixtures of polysulfides: n = 2-8 (n=9 only in ((NH₄)₂S₉), more frequently n=2). At n growth the colour will change from yellow through orange to red. The richest in sulfur (NH₄)₂S₉ has intensive red colour.

Under the action of strong acids, polysulfides form sulfanes Н₂S_n:

Н₂S_n + 2НСl = Н₂S_n + 2NaCl

A heavy yellow oil-like liquid prepared is a mixture of sulfanes Н₂S_n. In the individual state sulfanes up to Н₂S₈ are isolated: Н₂S₂ (m.p. = -88⁰, b.p. = 75⁰), Н₂S₃ (m.p. = -52⁰), Н₂S₄ (m.p. = -85⁰), Н₂S₅ (m.p. = -50⁰). They are relatively stable only in a strongly acidic medium, under other conditions they decompose with elimination of sulfur.

Polysulfide and sulfane molecules are zig-zag chains made by atoms of sulfur, valence orbitals of which are in the state of sp³-hybridization:

The ends of these chains are the atoms of Н (H₂S_n) or atoms of alkali metals in polysulfides. The dissociation energy of S—S bonds is equal to 226 kJ/mol and is similar enough for all sulfanes.

Sulfanes are weak acids in water solution. However, they considerably exceed H₂S strength. Their strength grows at the increase of sulfur atoms number in the molecule, for example:

Н₂S Н₂S₄ Н₂S₅

K₁ = 8.7∙10^-8 2∙10^-4 3∙10^-4

K₂ = 1∙10^-14 5∙10^-7 2∙10^-6

At the increase of sulfur atoms number H—S bonds become more polarized, therefore, its dissociation is facilitated under the action of polar water molecules. In accordance with the increase of acid strength, hydrolysis of their salts appropriately diminishes:

n in Н₂S_n 2 3 4 5

Degree of hydrolysis, % 64.6 37.6 11.8 5.7

Like peroxides, polysulfides display reduction properties:

They can display oxidizing properties under certain conditions:

Sn⁺²S + Na₂S₂^-1 = Sn⁺⁴S₂ + Na₂S^-2.

oxidant

when heated, they disproportionate:

Na₂S₂^-1 = Na₂S^-2 + S.

selenium, telurium, and polonium

Compounds of elements (-2)

Apart from the simple hydrogen chalcogenides of Se and Te (which are more acidic and thermally less stable than the corresponding sulfur analogs) the most stable inorganic compounds of selenium (and its higher homologs) are:

1. Chalcogenides and polychalcogenides with strongly electropositive ions such as alkali or alkaline-earth metals, and those that are more covalently bonded with transition metals;

2. Large variety of Se(II), Se(IV), and Se(VI) compounds with electronegative elements O, F, Cl, and Br;

3. Homologs belonging to the interesting class of sulfur–nitrogen compounds;

4. Organoselenium compounds.

When heated, selenium and tellurium react with metals forming selenides and tellurides:

2K + E = K₂E^-2; 2Al + 3E = Al₂E₃^-2

Thus active metals form salt-like selenides and stoichiometric tellurides, which dissolve in water readily.

Reactions of Se and Те with transitional metals give selenides and tellurides of various types МЕ, М₂Е₃, М₃Е₄, МЕ₂ and etc., with many of them having a varying composition. At low content of non-metallic elements they have mainly metallic nature, at high content of chalcogens — semiconductors.

Hydrogen compounds. Along with H₂S hydrogen selenide, H₂Se, and hydrogen telluride, H₂Тe, are known. Hydrogen polonide, H₂Ро, is extraordinarily unstable. Radioactivity of a gas formed after HCl treatment of polonium coated magnesium can serve as a proof. Polonides (Na₂Po) are also known.

Preparation. The action of diluted HCl or Н₂О on selenides and tellurides is used in a laboratory:

Al₂E₃ +6HCl = 3H₂E + 2AlCl₃

Al₂E₃ +6H₂O = 3H₂E + 2Al(OH)₃

Н₂Se can be prepared if gaseous Н₂ passes over heated to 400C elemental selenium.

Properties. Under ordinary conditions H₂Se and H₂Тe are colourless toxic gases with a very unpleasant smell. Their structure and properties resemble those of H₂S:

	H2О	H2S	H2Se	H2Te
m.p., С	0	-85.6	-65.7	-51
b.p., С	100	-60.4	-41.4	-2
Bond length, nm	0.096	0.133	0.146	0.169
Bond energy, J/mol	463	347	276	238
Valence angle H-E-H	104.5	92.2	91.0	90
Нf,298, kJ/mol	-285.8	-20.6	33	99.7
Gf,298, kJ/mol	-237.2	-33.8	19.7	85.2
K1(aqueous solution)	1.8•10-16	1•10-7	1.9•10-4	2.3•10-3

Covalent radii of chalkogens grow in the series H₂О—H₂S—H₂Se—H₂Te diminishing Н-Е bonds energy and H₂Е molecules stability. That is why unlike H₂О and H₂S, selenium and tellurium compounds with hydrogen are endothermic compounds (Н_f>0), they readily decompose when heated: thermal decomposition of H₂Se takes place with a noticeable rate at t^o > 300, and H₂Тe is gradually decomposed already under ordinary conditions.

Similarly to molecule H₂S the molecule of H₂Se and H₂Te have bent shape with valence angles 91 and 90, which signifies that hybridization of selenium and tellurium orbitals is absent.

Solubility of H₂Se and H₂Te is better as compared to H₂S. In aqueous solutions they display properties of weak hydroselenic and hydrotelluric acids. The strength of these hydracids gradually grows in the series H₂О-H₂S-H₂Se-H₂Te. It is explained by diminishing of Н — Е bond energy and growth of their ability to polarization. Dissociation of compounds is hereupon facilitated under the action of polar molecules of water:

Н₂Е Н⁺ + НЕ^-

HЕ^- Н⁺ + Е^2-

H₂Se can form two types of salts: neutral and acidic. H₂Te has no reactions when acidic salts are the products. Exposed to air oxygen their solutions quickly acquire a red color owing to the formation of polyselenides and polytellurides of metals, similar to polysulfides:

12K₂Te + 5O₂ + 10H₂O = 2K₂Te₆ + 20KOH

Neverthless, as a result of large size and low electronegativity of elements, a reaction between selenides of different chemical nature, as well as between tellurides with the formation of analogues of thiosalts can not be realised.

In the series H₂S-H₂Se-H₂Te the reducing activity grows appropriately. So, sulfur is able to oxidize H₂Se аnd H₂Te:

H₂Se^-2 + S = Se + H₂S^-2,

and oxygen oxidizes hydrogen compounds and their derivatives of other elements of the subgroup:

2H₂Е + O₂ = 2E + 2H₂O (in solution)

2H₂Е + 3O₂ = 2EО₂ + 2H₂O (burning)

Oxide of sulfur (IV). Preparation.

In industry:

1. Oxidative combustion of sulfides of metals (pyrites)

4FeS₂ + 11О₂ = 2Fe₂О₃ + 8SО₂

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

S + О₂ = SО₂ ; Н⁰₂₉₈ = -296.8 kJ/mol

3. Thermal decomposition or reduction of gypsum СaSо4.2h2o

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 -bonds and by a delocalized three-centered -bond.

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

The formation of -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 -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.

A ccording 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 -bonds delocalised three-centred -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 ( = 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 ( 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

• Strong 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∙hH₂O). Low soluble TeO₂∙hH₂O 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; Н⁰₂₉₈ =-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

It is 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 -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 -SO₃. It looks like transparent mass similar to ice; m.p. = 16.9 ^oC.

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

 - 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; Н⁰₂₉₈ = -89.2 kJ/mol