This article is a continuation of the revision notes on Class 12 Chemistry : Chapter- The p-Block Elements, Part-I.
In Part-I we have studied about the Group 15 elements (Nitrogen family), their important compounds like Dinitrogen, Ammonia, Oxides of Nitrogen, Nitric acid, Phosphine, Phosphorus Halides, etc.
In Part-II, we will get acquainted with Group 16 elements. All the these quick notes are based on the latest CBSE syllabus for Class 12th Chemistry.
The main topics covered in these quick notes are:
• Oxygen family or Chalcogens
• General properties of oxygen family
• Anomalous properties of oxygen
• Reactivity of group 16 elements towards:
• Preparation, properties and uses of dioxygen (O2)
• Simple Oxides
• Preparation, properties and uses of ozone (O3)
• Depletion of ozone layer
• Preparation, properties and uses of sulphur dioxide (SO2)
• Oxoacids of sulphur
• Preparation, properties and uses of sulphuric acid (H2SO4)
The notes of the chapter are as follow:
The elements of group 15: oxygen (O), sulphur (S), selenium(Se), tellurium (Te) and polonium (Po) having general electronic configuration ns2np4, are known as the oxygen family. All these elements collectively are also known as chalcogens. Polonium is a radioactive element.
General properties of oxygen family
Atomic and ionic radii: Due to increase in the number of shells, atomic and ionic radii increase from top to bottom in the group.
Ionisation enthalpy: Due to the increase in size of the atoms the ionisation enthalpy decreases down the group. IE1of group 16 elements is less than the IE1 of group 15. This is because group 15 elements have extrastability due to half-filled p-orbitals.
Electron gain enthalpy: Due the compact nature of oxygen, it has less electron gain enthalpy than sulphur. After sulphur, the electron gain enthalpy decreases down the group.
Electronegativity: The electronegativity decreases down the group. This implies that the metallic character increases down the group from oxygen to polonium.
Melting and boiling point: The melting and boiling point increases with increase in atomic number down the group.
Oxidation states: Group 16 elements show ‒2, +2, +4, +6 oxidation states. The stability of ‒2 oxidation state decreases down the group due to increase in atomic size and decrease in electronegativity. O shows only ‒2 oxidation state except when it combines with the most electronegative F with which it shows positive oxidation states. S shows + 6 only with O and F.
Anomalous behaviour of oxygen
Oxygen forms strong hydrogen bonding in H2O which is not found in H2S. Also, the maximum covalency of oxygen is four, whereas in a case of other elements of the group, the valence shells can be expanded and covalency exceeds four.
Reasons for the anomalous behaviour of oxygen are:
• Small size and high electronegativity
• Absence of d-orbitals
Reactivity towards hydrogen: All the elements of Group 16 form hydrides of the type H2E (E = S, Se, Te, Po).
Thermal stability: Thermal stability of group 16 elements decreases down the group.
H2O > H2S > H2Se> H2Te > H2Po
This is because the H-E bond length increases down the group, hence the bond dissociation enthalpy decreases down the group.
Acidic nature: Due to the decreasing bond dissociation enthalpy, acidic character of group 16 elements increases down the group.
H2O < H2S < H2Se < H2Te
Reducing character: The reducing character also decreases down the group due to the decreasing bond dissociation enthalpy.
H2O < H2S < H2Se < H2Te < H2Po
Reactivity towards oxygen: All group 16 elements form oxides of the type EO2 and EO3 Reducing character of dioxides decreases down the group. Acidity also decreases down the group. Besides EO2 type, sulphur, selenium and tellurium also form EO3 type oxides. Both types of oxides are acidic in nature.
Reactivity with halogens: Elements of Group 16 form a large number of halides of the type, EX2 EX4 and EX6, where X is a halogen. The stability of halides decreases in the order F− > Cl− > Br− > I− . This is because E-X bond length increases with increase in size. Among hexa halides, hexafluorides are the most stable because of steric reasons. Dihalides are sp3 hybridised and have tetrahedral geometry. H2O is a liquid while H2S is a gas. Because in water due to the small size and high electronegativity of O, strong hydrogen bonding is present there.
Oxygen is the first element of Group 16 with the electronic configuration of 1s2 2s2 2p4 in the ground state. Oxygen has two allotropes: dioxygen (O2) and trioxygen or ozone (O3).
Oxygen usually exists in the form of dioxygen.
Dioxygen is prepared in the laboratory by thermal decompositions of oxygen rich compounds such as KClO3,
(i) Oxygen is a colourless, odourless and is a highly reactive tasteless gas.
(ii) Due to the presence of pπ‒ pπ bonding, O2 is a discrete molecule and intermolecular forces are weak van der Waals forces, hence, O2 is a gas.
(iii) Dioxygen combines with metals and non-metals to form binary compounds called oxides.
2Ca + O2 → 2CaO
P4 + 5O2 → P4O10
(i) Dioxygen is used in making steel.
(ii) It is used in the production of oxygen containing organic chemicals.
(iii) Dioxygen is also used for sewage treatment, river revival and paper pulp bleaching.
(iv) It is used as an oxidiser in underwater diving and in space shuttles.
Oxygen combines with majority of the elements of the periodic table to forms oxides (O2‒). There are three types of oxides:
(i) Acidic oxides: Oxides of non- metals are usually acidic in nature. For example, SO2 combines with water to give H2SO3, an acid.
SO2 + H2O → H2SO3
(ii) Basic oxides: Metallic oxides are mostly basic in nature. Basic oxides dissolve in water to give basic solution. For example, CaO combines with water to give Ca(OH)2, a base.
CaO + H2O → Ca(OH)2
(iii) Amphoteric oxides: Some metallic oxides show characteristics of both acidic as well as basic oxides. Such oxides are known as amphoteric oxides. For example, Al2O3 reacts with acids as well as alkalies
Al2O3 + 6HCl + 9H2O → 2 [Al(H2O)6]3+ + 6 Cl‒
Al2O3 + 6NaOH + 3H2O → 2Na3[Al(OH)6]
Ozone is an allotropic form of oxygen. Ozone has angular structure with a bond angle of about 117o. Both O = O bonds are of equal bond length due to resonance.
It is formed when dioxygen is irradiated with UV light or silent electric discharge.
3O2 (g) → 2O3 (g) ΔHo = +142 KJ/mol
(i) Ozone is a pale blue gas with a characteristic pungent odour.
(ii) Ozone is diamagnetic in nature.
(iii) It is the second most powerful oxidising agent after fluorine. It liberates oxygen gas when acting as an oxidising agent.
2Fe2+ + O3 + 2H+ → 2Fe3+ + H2O + O2
PbS + 4O2 → PbSO4 + 4O2
Depletion of ozone layer:
Thinning of the ozone layer is termed as depletion of ozone layer. The depletion of ozone layer in the stratosphere is caused by the presence of chlorofluoro carbons. CFCs decomposed by UV radiation to produce chlorine which reacts with ozone and this causes a de-crease in the concentration of ozone at a rate faster than its formation from dioxygen. Another cause of depletion of ozone layer is the release of nitrogen oxides into the stratosphere by supersonic jet aeroplanes.
NO + O3 → NO2 + O2
Sulphur exhibits allotropy, two important allotropes of which are:
(a) Yellow Rhombic (α - sulphur)
(b) Monoclinic (β- sulphur)
At 369 K both forms are stable. S8 puckered shape in both forms and has crown shape.
Sulphur Dioxide (SO2)
The molecule of SO2 is angular. It is a resonance hybrid of the two canonical forms:
Sulphur dioxide is formed together with a little (6-8%) sulphur trioxide when sulphur is burnt in air or oxygen:
S (g) + O2(g) → SO2(g)
In the laboratory it is readily generated by treating a sulphite with dilute sulphuric acid.
Na2SO3 + H2SO4 → Na2SO4 + H2O + SO2
(i) SO2 is used in refining petroleum and sugar
(ii) It is used in bleaching wool and silk
(iii) It is also used as a disinfectant and preservative.
(iv) It is used in the manufacture of sulphuric acid, sodium hydrogen sulphite and calcium hydrogen sulphite.
Oxoacids of sulphur
Sulphur forms variety of oxoacids. All oxoacids of sulphur are dibasic.
Sulphur forms a number of oxoacids such as H2SO3, H2S2O3, H2S2O4, H2S2O5, H2SxO6 (x = 2 to 5), H2SO4, H2S2O7, H2SO5, H2S2O8 . Some of these acids are unstable and cannot be isolated. They are known in aqueous solution or in the form of their salts. Structures of some important oxoacids are shown in figure given below
Image Source: NCERT Books
Sulphuric acid (H2SO4)
Sulphuric acid is manufactured by the Contact Process as follows:
(i) Sulphuric acid is a colourless, dense oily liquid.
(ii) It is dibasic acid or diprotic acid.
(iii) It is a strong dehydrating agent.
(iv) It is a moderately strong oxidizing agent.
(i) H2SO4 is used in manufacture of fertilisers
(ii) It is also used in petroleum refining, manufacture of pigments, detergents, metallurgical application and storage batteries.