Trends in solid compounds: the importance of lattice enthalpy
Solid compounds of Group 1 and Group 2 elements:
All these elements form stable simple binary hydrides, halides, oxides and hydroxides with the group oxidation state. The +1 oxidation state of Group 2 is unstable with respect to disproportionation.
The stability of compounds with small anions (F–, H–, OH–,O2-) decrease down the groups, while the stability of those with large anions increase. The exception is Mg in Group 2.
For halides of the same element, the stability decreases from F– to I–, but the decrease is less for large cations.
Only Li in group 1 forms a stable nitride. The other Group 1 metals instead form azides MN3.
Structural chemistry of the binary compounds:
All hydrides, halides and hydroxides of Group 1 metals have the rock salt structure except CsCl, CsBr and CsI, which have the cubic CsCl structure. All the oxides M2O have the anti-fluorite structure, except Cs2O. The superoxides MO2 have the pyrites structure.
The Group 2 oxides except BeO have the rock salt structure. Of the difluorides, BeF2 has BeF4 tetrahedra sharing vertices, MgF2 has the rutile structure, and the other MF2 have the fluorite structure.
The other dihalides show increasing tendency to adopt more distorted and layered structures, and the coordination numbers exceed 8 for some barium compounds. This is very difficult to predict on ionic grounds.
As the cation increases in size, the thermal stability of compounds with large complex ions increases.
There are many examples of this, such as the oxides formed by heating the alkali metals in air: lithium forms the oxide, sodium forms the peroxide, while potassium, rubidium and cesium form the superoxide.( Note that rubidium and cesium also form a series of suboxides with interesting structures based on the face- and edge-sharing of OM6 octahedra.)
Heating Lithium nitrate gives the oxide, but the other Group 1 nitrates form the nitrite on heating. The stability of the carbonates, sulphates, etc. increase down the group. The large Group 1 metal ions can be used to stabilize complex interhalogen ions such as ClF4–. All of this behaviour can be accounted for using the ionic model, and is discussed in the Groups 1 and 2 section.
The ionic model can often successfully predict the feasibility of double decomposition reactions. Here, the combined lattice enthalpies of the products is greater than the combined lattice enthalpies of the reactants.
There are notable exceptions to this, however, especially those involving complex ion salts. Both of the reactions below are favourable even though the radius of the chloride ion is smaller than the radius of the nitrate ion.
Summary of trends in solid ionic compounds
The ionic radius increases down the group, so the lattice enthalpies decrease. Li+ and Mg2+ are very small, and hence have an anomalously high polarizing power, with a resultant high degree of covalent character.In Groups 1 and 2, which are the standard bearers of the ionic model, the valence electrons are outside the noble gas core, and so there are low ionization energies and relatively large cations.
The enthalpy of atomization and the ionization energy decrease down the group, so the enthalpy of formation of the Mn+(g) ion also deceases.
As the lattice enthalpy depends inversely on the sum of the anionic and cationic radii, it decreases with increasing ionic radius. The lattice enthalpy decreases with increasing cation radius faster for as small anion than for a large anion.
Trends in aqueous compounds: the importance of hydration enthalpy
In general, ions with a high value of z2/r, where z is the effective nuclear charge and r is the ionic radius, have higher hydration enthalpies. The enthalpy of hydration is expected to decrease down a group, as it is inversely proportional to the ionic radius.In solution metal ions are heavily hydrated, and it is probable that there are more than one structural forms in equilibrium. The primary (first) hydration sphere of most alkali metal and alkaline earth metal ions most likely has six water molecules, except the small Li+ ion which has four. However, since ion-dipole interactions decrease slowly with distance, the influence of Li+ extends well beyond the primary coordination sphere.
Ionic mobility increases down the group, and this reflects the decreasing size of the hydrated ion. Note that while the ionic radius of the naked ion increases down the group, as z2/r decreases and so the attraction with the solvent water molecules decreases and the size of the hydrated ion therefore also decreases.
Due to the great significance of electrostatic bonding in Group 1 and Group 2 metal ions, they do not form stable complexes with many ligands in aqueous solution.
Complexes with neutral ligands such as ethers, thioethers, ammonia and amines, pyridines, and large polarizable ions such as chloride, bromide and iodide are not very stable.
Complexes with small anions, such as F– and OH– are quite strong, as are those with O as donor ligands, especially for chelating ligands such as P2O74- and EDTA4-. This is a result of the Hard and Soft Acid/Base theory.
In general, the stability of a complex with a Group 2 metal ion is higher than the corresponding complex with a Group 1 metal ion.
The stability of complexes with small donor ligands usually follow the order predicted by the ionic model, ie. Li>Na>K and Mg>Ca>Sr>Ba.
These trends are also discussed in the sections: trends in the Group 1 and 2 metals, and solutions of the Group 1 and 2 metals.