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XII Topic: P Block Elements- 16Th Group Subject: Chemistry Oxygen

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XII Topic: P Block Elements- 16Th Group Subject: Chemistry Oxygen Class: XII Topic: P block elements- 16th group Subject: Chemistry Oxygen family or Chalcogens Oxygen family The elements of group 16: oxygen (O), sulphur (S), selenium(Se), tellurium (Te) and polonium (Po) having general electronic configurationns2np4, 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 extra stability 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 (O) Oxygen is the first element of Group 16 with the electronic configuration of 2 2 4 1s 2s 2p in the ground state. Oxygen has two allotropes: dioxygen (O2) and trioxygen or ozone (O3). Dioxygen (O2) Oxygen usually exists in the form of dioxygen. Preparation: Dioxygen is prepared in the laboratory by thermal decompositions of oxygen rich compounds such as KClO3, Properties: (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. Examples are: 2Ca + O2 → 2CaO P4 + 5O2 → P4O10 Uses: (i) Dioxygen is used in making steel. (ii) Dioxygen is also used for sewage treatment, river revival and paper pulp bleaching. (iii) It is used as an oxidiser in underwater diving and in space shuttles. Oxides 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 3+ ‒ Al2O3 + 6HCl + 9H2O → 2 [Al(H2O)6] + 6 Cl Base Acid Al2O3 + 6NaOH + 3H2O → 2Na3[Al(OH)6] Acid Base Note-The above content has been absolutely prepared from home .
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