The Group 2 elements — beryllium, magnesium, calcium, strontium, barium and radioactive radium — are the alkaline earth metals. The name reflects that their oxides (the ‘earths’) are alkaline. Each has two electrons in its outermost shell, giving the valence configuration ns2, and they form +2 ions. Compared with the Group 1 metals in the same period, they are harder, denser, higher-melting and somewhat less reactive, because the +2 ion involves a higher nuclear charge holding two electrons.
General characteristics
They are silvery-white, fairly hard metals. Because two valence electrons take part in metallic bonding, the bonds are stronger than in alkali metals, so melting points and densities are higher and the metals are harder. They are still good reducing agents but less so than Group 1.
Trends down the group
Atomic and ionic radii increase down the group, though each Group 2 atom is smaller than the Group 1 atom of the same period (higher nuclear charge). Ionisation enthalpy decreases down the group; the first ionisation enthalpies are higher than those of the corresponding alkali metals, but the second is still moderate so the +2 state dominates. Hydration enthalpies are larger than for the +1 alkali ions (because of the +2 charge) and decrease down the group; this is why MgCl2 and CaCl2 are hydrated solids while NaCl is not. Reactivity increases down the group as ionisation enthalpy falls. Be and Mg do not impart colour to a flame, but Ca gives brick-red, Sr crimson and Ba apple-green flames because their electrons are excited at energies within the visible range.
Anomalous behaviour of beryllium
Beryllium, the first member, is anomalous due to its very small size, high charge density and high electronegativity. Its compounds are largely covalent (e.g. BeCl2 is covalent and polymeric in the solid state), it does not react with water even at high temperature, its oxide and hydroxide are amphoteric (dissolving in both acids and alkalis), and it forms complexes such as [BeF4]2−. Beryllium also shows a maximum covalency of 4.
Diagonal relationship: Be and Al
Beryllium resembles aluminium (Group 13, diagonally placed) because of similar charge-to-radius ratios. Both BeO and Al2O3 are amphoteric, both metals dissolve in alkali to evolve hydrogen, both chlorides (BeCl2, AlCl3) are covalent Lewis acids soluble in organic solvents and act as catalysts, and both form covalent carbides.
Important compounds
Quicklime, CaO is made by heating limestone: CaCO3 → CaO + CO2. It is a basic oxide used in cement, in steel-making and to make slaked lime. Slaked lime, Ca(OH)2 forms when quicklime reacts with water: CaO + H2O → Ca(OH)2; lime water (its solution) turns milky with CO2 and is used in whitewashing and to soften water. Limestone/marble, CaCO3 is used in cement, glass and as a building stone. Plaster of Paris, CaSO4·½H2O is made by heating gypsum (CaSO4·2H2O) to about 393 K: 2CaSO4·2H2O → 2CaSO4·½H2O + 3H2O. On mixing with water it sets to a hard mass of gypsum again, which is why it is used in plaster casts, moulds and statues.
Why are the ionisation enthalpies of alkaline earth metals higher than those of the alkali metals of the same period?
Solution- An alkaline earth metal has one more proton in its nucleus than the alkali metal of the same period.
- The greater nuclear charge pulls the electrons in more tightly, so the atom is smaller.
- A smaller atom holds its valence electrons more firmly.
- Hence more energy is needed to remove the first electron from a Group 2 atom than from the Group 1 atom of the same period.
Answer: The higher nuclear charge and smaller size of Group 2 atoms make their valence electrons more tightly held, so their ionisation enthalpies are higher than those of Group 1.
Give three ways in which beryllium behaves anomalously compared with the rest of Group 2.
Solution- Beryllium is very small with high charge density, so its compounds (e.g. BeCl2) are largely covalent rather than ionic.
- Beryllium oxide and hydroxide are amphoteric, dissolving in both acids and alkalis, unlike the basic oxides of Mg, Ca, etc.
- Beryllium does not react with water, whereas the other members react with water or steam.
Answer: Covalent compounds, amphoteric BeO/Be(OH)2, and no reaction with water are three anomalies of beryllium.
Write the equations for the preparation of plaster of Paris from gypsum and for its setting with water.
Solution- Gypsum is calcium sulphate dihydrate, CaSO4·2H2O.
- On heating to about 393 K it loses most of its water: 2CaSO4·2H2O → 2CaSO4·½H2O + 3H2O.
- The product, CaSO4·½H2O, is plaster of Paris.
- On mixing with water it sets by reabsorbing water: 2CaSO4·½H2O + 3H2O → 2CaSO4·2H2O (a hard mass of gypsum).
Answer: Preparation: 2CaSO4·2H2O → 2CaSO4·½H2O + 3H2O; setting: 2CaSO4·½H2O + 3H2O → 2CaSO4·2H2O.
How is quicklime obtained from limestone, and what happens when quicklime is treated with water?
Solution- Limestone is calcium carbonate, CaCO3.
- Heating it strongly (calcination) drives off CO2: CaCO3 → CaO + CO2.
- The residue CaO is quicklime, a basic oxide.
- Adding water to quicklime is highly exothermic and gives slaked lime: CaO + H2O → Ca(OH)2.
Answer: CaCO3 → CaO + CO2 gives quicklime; with water, CaO + H2O → Ca(OH)2 (slaked lime), a vigorous exothermic reaction.
List three properties shared by beryllium and aluminium that illustrate their diagonal relationship.
Solution- Be and Al have similar charge-to-radius ratios, the basis of the diagonal relationship.
- Both BeO and Al2O3 are amphoteric, reacting with both acids and alkalis.
- Both BeCl2 and AlCl3 are covalent, soluble in organic solvents, and act as Lewis acid catalysts.
- Both metals dissolve in alkali, liberating hydrogen.
Answer: Amphoteric oxides, covalent Lewis-acid chlorides, and dissolution in alkali with H2 evolution are shared properties showing the Be–Al diagonal relationship.
Why are Be and Mg unable to impart colour to a flame, whereas Ca, Sr and Ba do?
Solution- A flame colour appears only when an electron can be excited by the small energy available in a Bunsen flame and re-emit visible light.
- Be and Mg are small atoms whose electrons are tightly bound, so the excitation energy needed lies in the ultraviolet, not the visible region.
- Ca, Sr and Ba are larger with more loosely held electrons that can be excited by the flame.
- These excited electrons emit light in the visible range — brick-red (Ca), crimson (Sr) and apple-green (Ba).
Answer: In Be and Mg the electrons are too tightly held, so the emitted light falls in the UV; in Ca, Sr and Ba the looser electrons emit visible light, giving characteristic flame colours.