The s-block elements are characterised by their larger atomic sizes, lower ionisation enthalpies, invariable +1 oxidation state and solubilities of their oxosalts. In the light of these features describe the nature of their oxides, halides and oxosalts.
The s-block elements contain two groups on the periodic table, the alkali metals (Group 1) and the alkaline earth metals (Group 2). The nature of their oxides, halides and oxosalts is given as follows:
1) Alkali metals:
(i) Oxides: Alkali metals form oxides on combustion in excess air. Lithium forms oxide Li2O and peroxide Li2O2 in minor quantities. Sodium forms the peroxide, Na2O2 (and superoxide NaO2 in minor quantities) whilst potassium, rubidium and caesium form the superoxides, MO2. Going down the group, the size of the atoms increases and the ability to form peroxides and superoxides also increases. This happens due to the stabilisation of large anions by larger cations through lattice energy effects.
(ii) Halides: Alkali metal halides can be produced by reaction with the appropriate carbonate, oxide or hydroxide and hydrohalic acid (HX). The melting and boiling points follow the trend flouride > chloride > bromide > iodide. All alkali metal fluorides have high negative enthalpy of formation. The values for fluorides become less negative down the group and for chlorides, bromides and iodides the values become more negative down the group. All of these halides are highly soluble in water except LiF which has high lattice energy.
(iii) Oxosalts: All alkali metals react with oxo-acids like carbonic acid H2CO3 and sulphuric acid H2SO4 to form oxo-salts like carbonate and sulphate respectively. They are soluble in water and thermally stable. The carbonates and hydrogencarbonates, M2CO3 and MHCO3 respectively are thermally stable. The stability of these compounds increases with increasing electropositive character down the group. The exception is Li2CO3 which is not thermally stable due to the small size of the Li atom, which polarizes the large CO23- ion. This leads to the decomposition and formation of LiO and CO2. The hydrogencarbonate does not exist as a solid.
2) Alkaline Earth Metals
(i) Oxides: Alkaline earth metals burn in oxygen to form metal oxides in the form of MO. The oxides are ionic in nature except BeO which is covalent in nature. The enthalpies of formation of oxides is high. The oxides are basic in nature except BeO which is amphoteric.
(ii) Halides: Halides of alkaline earth metals are ionic in nature with the exception of BeCl2 which is covalent and is soluble in organic solvents. The fluorides are relatively less soluble than the chlorides owing to their high lattice energies.
(iii) Oxosalts: Alkaline earth metals form oxosalts with oxo-acids. They form carbonates with carbonic acid, sulphates with sulphuric acid and nitrates with nitric acid. Carbonates are insoluble in water and they decompose on heating to form the oxides and carbon dioxide. Beryllium carbonate is unstable and has to be placed in a CO2 atmosphere to stop it from decomposing. The sulphates are stable in heat. BeSO4 and MgSO4 are readily soluble in water because of hydration enthalpies of Be2+ and Mg2+ overcoming lattice enthalpy, and the solubility reduces from CaSO4 to BaSO4. Nitrates show decreasing tendency to form hydrates with increasing size and decreasing hydration enthalpy in the group. All of them decompose to give metal oxide MO and NO2 and O2.
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