Group 2

Dominated by +2 oxidation state chichis normally in the form of the M2+ ions, ionic group 2 compound stabilised by relatively low values of the first and second ionisation enthalpies of the metals and the high values of the lattice enthalpies for group 2 salts, beryllium is exceptional as it forms a number of covalent compounds

Elements are metallic solids at room temperature and pressure, higher melting points, boiling points and enthalpy changes of atomisation than their group 1 neighbours, this indicates stronger metallic bonding in the metal lattices due to the presence of two rather than one delocalised valence electrons per atom

Group2 metals much less reactive than group 1 metals and with the exception of barium do not need to be stored under oil. Most group 2 metals have a coating of oxide on their surfaces which makes them much less reactive than might be expected. They are sufficiently reactive however the free metals do not occur naturally

Obtained from sea water, in the first stage of the process aqueous calcium hydroxide solution is added to the sea water causing Mg(OH)2 to precipitate Mg2+ + Ca(OH)2 –> Mg(OH)2 + Ca2+ The solid magnesium hydride is treated with HCl to produce MgCl2 and the metal is liberated from molten MgCl2 using electrolysis Reduction at cathode Mg2+ + 2e- —> Mg Oxidation at anode 2Cl- —> Cl2 + 2e-

Calcium metal is obtained from minerals such as limestone CaCO3, firstly the carbonate is converted into the oxide by heating to 900-1100 degrees CaCO3 –> CaO + CO2 Calcium metal is obtained form the reaction of CaO (quicklime) with aluminium 6CaO + 2Al –> 3CaO.Al2O3 + 3Ca

Group 2 metals burn in air reacting with oxygen to give the oxides 2M + O2 —> 2MO Magnesium burns with a bright white flame with some of the nitride Mg3N2 produced alongside MgO The reaction becomes more exothermic down the group and barium can ignite with moist air

Barium can also form the peroxide with excess oxygen Ba + O2 —> BaO2

They are very basic and with the exception of BeO dissolve in water to give alkaline solutions of the hydroxides, this reaction is very exothermic fro CaO CaO + H2O –> Ca(OH)2 but less so for MgO which only reacts slowly

With the exception of beryllium the group 2 metals react with water to give the hydroxide liberating hydrogen gas in the process Ba + 2H2O –> Ba(OH)2 + H2

As with group 1 metals the reaction become more vigorous as the group is descended. Magnesium requires hot water or steam to react whereas calcium reacts with cold water. The hydroxides are generally basic in solution

The group 2 hydroxides decompose on heating to give oxides M(OH)2 –> MO + H2O

Be(OH)2 is amphoteric, as well as reacting with acids in the manner expected for a hydroxide to dissolves in excess alkali to form a tetrahedral complex anion Be(OH)2 + 2H3O+ –> Be2+ + 4H2O Be(OH)2 + 2OH- –> [Be(OH)4]2-

The group 2 metals react with the halogens to give the halide. Eg magnesium burns in chlorine gas to form magnesium chloride Mg + Cl2 —> MgCl2 Most of the group 2 halides are ionic but BeX2 compounds have polymeric structures with covalent bonds and bridging halides

Both BeCl2 and BeH2 have structures containing bridging atoms but they differ with respect to the bonding, BeH2 is electron deficient and the Be-H-Be bridges involve 3 centre 2 electron bonding, BeCl2 is not electron deficient because the bridging chloride donates three electrons

It is stable with respect to its elements but is unstable with respect to disproportionation

Group 2 metals form binary compounds with carbon, although sometimes called carbides these groups are really ethynides since they contain the C22- anion. These compound react with water to give ethyne CaC2 + 2H2O —> Ca(OH)2 + C2H2 Ethyne is burned to produce light

The group 2 metals all react with nitrogen on heating to give nitrides 3Mg + N2 –> Mg3N2 The nitrides decompose in water to form the metal hydroxide and ammonia Mg3N2 + 6H2O –> 3Mg(OH)2 + 2NH3

Group 2 carbonates are not very soluble in water and decompose on heating to the oxides and CO2 MCO3 –> MO + CO2 The decomposition temperature in degrees C of the carbonates increase as the group is descended

The decomposition temperature is related to the temperature at which a reaction becomes spontaneous that is when Gibbs energy becomes negative, for decomposition of a carbonate the entropy change is positive and largely independent of M this means that any changes in Gibbs energy are caused by changes in enthalpy. The enthalpy change is related to difference between the lattice enthalpies of MCO3 and MO

These also decompose on heating to give the oxides 2M(NO3)2 –> 2MO + 4NO + 3O2 MSO4 –> MO + SO3

Limestone CaCo3 is a very common minerals that is used directly as a building material, when its heated to 900 - 1100 degrees C it loses CO2 forming calcium oxide known as quicklime CaCO3 –> CaO + CO2

Tricalcium silicates (Ca3SiO5) is a typical component of comment. It hardens by absorbing water to form Ca3SiO5.2H2O long needle like crystals of this hydrated calcium silicate are formed which bon the cement particles together. Adding gypsum (CaSO4.2H2O) slows the solidification this mixture is called Portland cement Concretes typically contain a cement together with an aggregate concrete is the most common building material in the world by mass

The trend in solubility of group 2 compounds are similar to those for group 1 but as the lattice enthalpies are higher the compounds are generally less soluble. Compounds with large cations and large anions tend to be insoluble eg MgSO4 is very soluble, CaSO4 much less so and BaSO4 very poorly soluble.

Test for sulfate ions as their presence gives a dense white precipitate Ba2+ SO42- —> BaSO4