Class 11 Chemistry — Chapter 303: Elements of Group 1 and 2

Key concepts: Elements of Group 1 and 2 (Class 11 Chemistry)

The s-block holds the two most reactive families of metals: the alkali metals of group 1 (Li, Na, K, Rb, Cs) with an ns¹ configuration, and the alkaline earth metals of group 2 (Be, Mg, Ca, Sr, Ba) with ns². Because the valence electrons are few and loosely held, almost everything in this chapter follows from a single chain of reasoning — size increases down a group, ionisation enthalpy and hydration enthalpy fall, and reactivity therefore rises. Layered on top of that smooth trend are two exceptions worth memorising, the anomalous behaviour of lithium and beryllium and their diagonal relationships with magnesium and aluminium, followed by the industrial chemistry of NaOH, Na₂CO₃, NaHCO₃, CaO, Ca(OH)₂, CaCO₃ and plaster of Paris, and the biological roles of Na, K, Mg and Ca.

Electronic configuration of the s-block

Alkali metals have the valence configuration ns¹ and alkaline earth metals ns², both sitting just after a noble gas core. Losing these one or two electrons gives M⁺ or M²⁺ with a noble gas arrangement, which is why the +1 and +2 states are the only ones normally seen.

Atomic and ionic radii

Radii increase down each group as a fresh shell is added, so Li < Na < K < Rb < Cs and Be < Mg < Ca < Sr < Ba. Across a period, a group 2 atom is smaller than its group 1 neighbour because the extra nuclear charge pulls the same shell inwards; the cation is always much smaller than the parent atom.

Ionisation enthalpy

Alkali metals have the lowest first ionisation enthalpy in their periods and the value falls further down the group as the valence electron moves away from the nucleus and is better shielded. The second ionisation enthalpy of an alkali metal is enormous because it would break into the noble gas core, which is why M²⁺ is never formed.

Hydration enthalpy

Smaller ions attract water molecules more strongly, so hydration enthalpy follows Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺ and Be²⁺ > Mg²⁺ > Ca²⁺ > Sr²⁺ > Ba²⁺. Li⁺ therefore carries the thickest water sheath, giving it the largest hydrated radius and the lowest ionic mobility in solution even though it is the smallest bare ion.

Reducing character and electrode potential

All s-block metals are strong reducing agents. Lithium is the strongest in aqueous solution, with E° = −3.04 V, in spite of having the highest ionisation enthalpy in its group, because its exceptionally large hydration enthalpy more than pays back the energy spent on atomising and ionising the metal.

Physical properties

Alkali metals are soft, silvery, light and low melting, and lithium is the least dense of all metals at about 0.53 g cm⁻³. Alkaline earth metals are harder, denser and higher melting than the alkali metal of the same period because each atom contributes two electrons to a stronger metallic bond.

Reaction with water

The general reaction 2M + 2H₂O → 2MOH + H₂ becomes more violent down group 1, from gentle fizzing with lithium to an explosion with caesium. In group 2, beryllium does not react even with steam, magnesium reacts with hot water or steam, and calcium, strontium and barium react with cold water.

Reaction with oxygen

The product depends on the size of the cation: lithium gives the normal oxide Li₂O, sodium the peroxide Na₂O₂, and the larger K, Rb and Cs give superoxides such as KO₂. Group 2 metals burn to normal oxides MO, while barium can also form BaO₂.

Hydrides and halides

Alkali metals give ionic hydrides MH containing the H⁻ ion, whereas BeH₂ and MgH₂ are covalent and polymeric. All the halides MX are ionic high melting solids, except that lithium halides show growing covalent character in the order LiF < LiCl < LiBr < LiI, in line with Fajans' rules.

Flame colours

The loosely held valence electron is excited in a flame and emits a characteristic colour on returning: Li crimson, Na golden yellow, K lilac, Rb reddish violet, Cs blue; Ca brick red, Sr crimson and Ba apple green. Beryllium and magnesium give no colour because their ionisation enthalpies are too high for the flame to excite them.

Solutions in liquid ammonia

Alkali metals dissolve in liquid ammonia to give deep blue, conducting, paramagnetic solutions in which the metal has separated into ammoniated cations and ammoniated electrons. On standing or on becoming concentrated, the solution turns bronze coloured and slowly liberates hydrogen with the amide.

Anomalous behaviour of lithium

Lithium is set apart by its very small size, high charge density, high ionisation enthalpy and absence of d orbitals. Consequently Li₂CO₃, LiOH and LiNO₃ decompose on heating unlike their sodium analogues, LiF and Li₂CO₃ are only sparingly soluble, LiCl is deliquescent and soluble in ethanol, and lithium alone combines directly with nitrogen to give Li₃N.

Anomalous behaviour of beryllium

Beryllium forms mainly covalent compounds because of its small size and high polarising power, and its maximum covalency is four since it has no d orbitals. BeO and Be(OH)₂ are amphoteric, solid BeCl₂ is a polymeric chain that dissolves in organic solvents, and beryllium does not react with water at all.

Diagonal relationship

Li resembles Mg and Be resembles Al because a diagonal pair has almost the same charge-to-radius ratio. Both Li and Mg form nitrides and give oxides on heating their nitrates and carbonates; both Be and Al have amphoteric oxides, dissolve in alkali to liberate hydrogen, are rendered passive by concentrated nitric acid, and form chlorides that act as Lewis acid catalysts.

Sodium hydroxide and sodium carbonate

NaOH is made by electrolysing brine in the Castner-Kellner cell with a mercury cathode, giving sodium amalgam that is then decomposed by water. Na₂CO₃ is made by the Solvay process, in which sparingly soluble NaHCO₃ is precipitated from ammoniacal brine and then heated; the method fails for K₂CO₃ because KHCO₃ is far too soluble to settle out.

Lime, limestone and plaster of Paris

Heating limestone CaCO₃ gives quicklime CaO, and slaking quicklime with water gives Ca(OH)₂, whose clear solution is lime water. Gypsum CaSO₄·2H₂O heated to about 393 K loses three-fourths of its water to give plaster of Paris CaSO₄·½H₂O, which sets hard on remixing with water; overheating gives dead-burnt plaster that will not set.

Biological importance

A 70 kg adult holds roughly 90 g of sodium and 170 g of potassium. Na⁺ is the chief cation outside the cells and K⁺ inside, and the Na⁺-K⁺ pump that maintains this difference drives nerve impulses; Mg²⁺ sits at the centre of chlorophyll and activates ATP-handling enzymes, while Ca²⁺ builds bone and teeth as calcium phosphate and is essential for blood clotting.