JEE Chemistry The P Block Elements Notes Set 23

Group 16: Occurrence and Trends

The group 16 (chalcogen) elements are O, S, Se, Te and radioactive Po. Oxygen is the most abundant element in the earth's crust (about 46 percent, in oxides, silicates and water) and forms 21 percent of air by volume; sulphur occurs free in volcanic regions and as sulphides (galena, zinc blende, iron pyrites) and sulphates (gypsum, Epsom salt, barytes). The valence shell is \( \ce{ns^2 np^4} \).
Trends: size and metallic character increase; ionisation enthalpy and electronegativity fall; oxygen has a less negative electron gain enthalpy than sulphur because of repulsion in its small 2p shell. Oxidation states are -2, +2, +4 and +6; the +4 state becomes more stable than +6 down the group, and oxygen shows only -2 (+2 in \( \ce{OF2} \), -1 in peroxides), as it lacks d-orbitals. The hydrides \( \ce{H2E} \) show decreasing thermal stability and increasing acidity and reducing power from \( \ce{H2O} \) to \( \ce{H2Te} \); \( \ce{H2O} \) has the highest boiling point because of hydrogen bonding. The single O-O bond is weaker than S-S, which is the strongest of the group, so sulphur catenates most.

Allotropes of Sulphur

On heating rhombic sulphur: it melts to a pale yellow mobile liquid of \( \ce{S8} \) rings, near 430 K the rings open and polymerise so the viscosity rises, then at higher temperatures the chains break and the liquid thins again; the vapour is mostly \( \ce{S8} \) at about 473 K and mostly the paramagnetic \( \ce{S2} \) at about 1000 K. The standard state of sulphur is the rhombic form.

Oxygen

Preparation: \( \ce{2KClO3 ->[MnO2] 2KCl + 3O2} \), \( \ce{2HgO -> 2Hg + O2} \), \( \ce{2H2O2 -> 2H2O + O2} \), electrolysis of water, and industrially by fractional distillation of liquid air. Oxygen is paramagnetic (\( \ce{O2} \) has two unpaired electrons in antibonding pi* orbitals, bond order 2), supports combustion and is absorbed by alkaline pyrogallol. Oxides are classified as acidic (non-metals), basic, amphoteric (\( \ce{Al2O3} \), \( \ce{ZnO} \), \( \ce{Sb4O6} \)) and neutral (CO, NO, \( \ce{N2O} \)).

Ozone and Hydrogen Peroxide

Ozone \( \ce{O3} \) is made by passing a silent electric discharge through dry oxygen (about 10 percent ozone); the reaction \( \ce{3O2 -> 2O3} \) is endothermic, so ozone is less stable than oxygen. It is a bent, diamagnetic molecule (central atom sp2, two equal O-O bonds of order 1.5), a pale blue, pungent-smelling toxic gas, and a very strong oxidant: \( \ce{PbS + 4O3 -> PbSO4 + 4O2} \), \( \ce{2KI + H2O + O3 -> 2KOH + I2 + O2} \), \( \ce{NO + O3 -> NO2 + O2} \). It turns mercury sticky (tailing), is absorbed by oil of turpentine and oil of cinnamon, and in the stratosphere it screens ultraviolet light; chlorine atoms from CFCs and NO destroy it catalytically.
Hydrogen peroxide is a pale blue liquid, a weak acid with a non-planar (open-book) structure. Preparation: \( \ce{BaO2 + H2SO4 -> BaSO4 + H2O2} \) or from 2-ethylanthraquinol with air. Oxidising and reducing: in acid \( \ce{2Fe^2+ + H2O2 + 2H+ -> 2Fe^3+ + 2H2O} \), \( \ce{PbS + 4H2O2 -> PbSO4 + 4H2O} \); as a reductant it reduces \( \ce{MnO4-} \), \( \ce{Cl2} \) and \( \ce{O3} \) with release of oxygen. It decomposes in light, with dust and with alkali, so it is stored in dark plastic or wax-lined bottles, often with a stabiliser.

Sulphur Compounds

Hydrogen sulphide: \( \ce{FeS + H2SO4 -> FeSO4 + H2S} \) (Kipp's apparatus), a colourless gas with a rotten-egg smell, a weak dibasic acid and a good reducing agent: \( \ce{H2S + Cl2 -> 2HCl + S} \), and it burns to \( \ce{SO2} \) in excess air. With acidified \( \ce{KMnO4} \) or nitric acid it first gives colloidal sulphur.
Sulphur dioxide: \( \ce{S + O2 -> SO2} \) or roasting of sulphides; a bent, colourless, suffocating gas. Its aqueous solution is acidic and reducing, and bleaches by reduction (the colour returns on standing in air). It reduces \( \ce{MnO4-} \), \( \ce{Cr2O7^2-} \) (to green \( \ce{Cr^3+} \)) and \( \ce{Fe^3+} \). Sulphur trioxide is made by the contact process: \( \ce{2SO2 + O2 <=> 2SO3} \) with \( \ce{V2O5} \) at about 720 K and 2 bar; it is absorbed in 98 percent \( \ce{H2SO4} \) to give oleum \( \ce{H2S2O7} \) and diluted, since direct dissolution in water forms a mist.
Sulphuric acid is dibasic, a strong dehydrating agent (sugar, formic acid), a sulphonating agent and, when hot and concentrated, an oxidant: \( \ce{Cu + 2H2SO4 -> CuSO4 + SO2 + 2H2O} \), \( \ce{2HBr + H2SO4 -> Br2 + SO2 + 2H2O} \). Sodium thiosulphate is made as \( \ce{Na2SO3 + S -> Na2S2O3} \); it reacts with iodine \( \ce{2S2O3^2- + I2 -> S4O6^2- + 2I-} \), is used as a fixer in photography (\( \ce{AgBr + 2S2O3^2- -> [Ag(S2O3)2]^3- + Br-} \)) and as an antichlor.

Worked Examples

Example 1. How many sulphur atoms in sodium tetrathionate are in the +5 state, and what is the oxidation state of the others?
Answer:
In \( \ce{S4O6^2-} \), the two end sulphur atoms (each in an \( \ce{SO3} \) group) are +5 and the two middle sulphur atoms are 0; check: 2(5) + 2(0) + 6(-2) = -2.

Example 2. Why is the oxidation state of oxygen never above +2 and why does sulphur show +4 and +6?
Answer:
Oxygen is the second most electronegative element and has no d-orbitals, so it cannot expand its octet; the best it does is +2 with fluorine. Sulphur has vacant 3d orbitals, so electrons can be promoted to give 4 or 6 unpaired electrons.

Example 3. A mixture of white and red phosphorus is to be separated. Suggest one method.
Answer:
Treat the mixture with \( \ce{CS2} \): white phosphorus dissolves and red phosphorus does not. (Hot alkali would also destroy the white form by giving phosphine.)