Aqueous solution chemistry

In solution, the metal ions are heavily hydrated. The primary hydration sphere is six water molecules (except Lithium which is four, due to its small size), but the ion-dipole coulombic interactions extend beyond the first sphere, ie. attraction between the ionic charge and the dipole on the water molecule. The enthalpy of hydration increases as z²/r, therefore the enthalpy of hydration decreases down the group. Also, the size of the hydrated ion decreases down the group (this is counter to the trend in size of the dehydrated ion, which increases down the group), and hence the ionic mobility increases down the group.

Complexes with neutral ligands, eg. R₂O, R₂S, R₃N, and halide ions are generally unstable, except those with Be²⁺ and Mg²⁺. In general, the smaller the metal ion, the greater the stability of the complex.

Complexes with small anions, eg. F^- and OH^-, are generally quite strong. Other strong complexes are formed with O donor ligands:

This preference for forming compounds with O, F (and N) ligands reflects the hard acid nature of the aqueous cations, and the interaction is the hard acid metal cation with the hard base O, F, or N donor atom.

The stability of complexes with small anions follows the order predicted by the ionic model. This says that small cations favour small anions, and large cations favour large anions, and this is reflects in the order of stability: Li>Na>K and Mg>Ca>Sr>Ba.

In general, the stability of a compound with a given ligand is higher for a group 2 metal than for a group1 metal: this reflects the higher charge of the group 2 metal, and hence the increased coulombic interaction with the ligands.

Non-aqueous solution chemistry

Group 1 metals dissolve in liquid ammonia and undergo ionization in the process.

This forms a dilute solution of electrons: these electrons polarize the NH₃ molecules to form a cavity in which the electron becomes trapped. The energy levels of the electrons within these cavities can be predicted using the particle in a box method, and transitions between these levels give rise to the strong blue colour of the dissolved electron solution. Other metals also dissolve in ammonia and give a colour which does not depend on the metal; the colour is due to the energy levels of the trapped electron not the dissolved metal ion.

More concentrated solutions become metallic, with the electrons becoming delocalized over the solution, and the solution takes on a bronze appearance.

These solutions can survive for long periods, but their decomposition is catalyzed by the presence of d-block compounds.

In effect, the dissolved electron reduces the H⁺ ion in the ammonia to hydrogen gas. Other reduction reactions may take place in the solutions, and unusual complexes, such as [Ni₂(CN)₆]^4-, may be formed (in the absence of air).