The s-metals consist of the Alkali Metals (Group 1) and the Alkaline Earth Metals (Group 2). They generally occur in compounds with oxidation states +1 and +2 respectively, though in the absence of air and water, some compounds with the metals in lower oxidation states may be prepared.
Some of the Group 1 and 2 metals are amongst the most abundant: calcium, sodium, magnesium and potassium are the 5th to 8th most abundant metals respectively, though others like Lithium and Beryllium have very low abundances.
Most of the metals are isolated by electrolysis of their molten salts; because the metals are so reducing (see the standard reduction potentials in the table below), electronic reduction of their cations is generally the only way of their isolation. This can often be very expensive.
|Group 1||E (V)||Group 2||E (V)|
The standard reduction potentials of these metals mean that their oxidation by water proceeds rapidly: the evolution of hydrogen gas means that the reaction can be explosive. This is because the heat evolved from the burning hydrogen can melt the metals, which have low melting points, thus greatly increasing the surface area of metal available to react with the water, and so increasing the rate of reaction. The reactivity increases down the groups: in fact, Beryllium and Magnesium are stable in water and air due to the presence of a thin oxide layer formed by reaction with the air which prevents reaction with the water.
|n = 1 for group 1, or n = 2 for group 2|
The similarity in the standard reduction potentials of the Group 1 metals is due to the balancing of various terms in the Born-Haber cycle for the process. As the group is descended, the enthalpies of ionization and sublimation both decrease, which favours oxidation, but this is balanced by the less exothermicenthalpy of solvation, which disfavours oxidation. The result is that the thermodynamics of the process are broadly similar for each of the Group 1 metals, and this reflected by the similar reduction potentials.
The general chemistry of Groups 1 and 2
Group 1 metals most clearly show the effect of increasing size and mass on the decent of a group. For example,
The enthalpy of sublimation and melting point.
The lattice energies.
The effective hydrated ionic radii.
The ease of thermal decomposition on carbonates and nitrates (see table)
the strength of covalent bonds in M2
All of these decrease down the group.
All form simple binary hydrides, halides, oxides and hydroxides with the metal in the group oxidation state. The +1 oxidation state is unstable with respect to disproportionation in group 2.
The stability of the compounds with small anions increases and the stability with large anions decreases down the group.
For MX, the stability decreases from F– to I–, but the decrease in stability is less for large cations.
All group 2 metals form stable nitrides, but only Lithium in group 1. The other group 1 metals form Azides [M+(N3)–]. The structure of Lithium Nitride is as shown, based on hexagonal layers of Li+ ions.
Li3N: Hexagonal layers of Li with N at the centers of the hexagons, forms (Li2N)–. The other Li+ ions bridge N ions in adjacent layers.
As the cation increases in size down the group, the thermal stability of compounds with large complex ions increases. For example, in group 1 oxides, the energetically favoured forms are (Li+)2O2-, (Na+)2O22-, and Rb+O2–.
The Stability of Group 1 carbonates, M2CO3. The thermal stability with respect to loss of CO2 decreases down the group. Why?Consider the thermochemical cycle for the loss of CO2 from the carbonate. ΔHr is the enthalpy of reaction for the conversion of the carbonate ion into the oxide ion and CO2. It does not depend on M.
The total enthalpy change for the process depends on three terms. One of these is a constant with M, and the other two depend inversely on the size on the metal cation. The size of the cation therefore determines where crossover from stable carbonate to stable oxide comes.
As the cation gets bigger, the carbonate gets more stable relative to the oxide.
Simply, large cations are more stable with large anions, and small cations are more stable with small anions.
This fact also explains the trend in stability of the Group 1 oxides, nitrides/azides, and halides, as discussed above.
All the MH, MX, and MOH have the rock salt, NaCl, structure (with the exceptions of CsCl, CsBr and CsI, which have the Caesium Chloride, CsCl, structure).
All M2O have the antifluorite structure (except Cs2O).
All MIIO have the NaCl structue (except BeO, which has the wurtzite structure).
All MIIF2 have the fluorite structure (except BeF2, which has the quartz structure, made up of vertex shared BeF4 tetrahedra, and MgF2 which has the rutile structure). This reflects the increasing size of the cations down the group.
Other MX2 have an increasing tendency to form distorted and layered structures, eg. the coordination number of Ba2+ is greater than 8 in some compounds.
The structures of Be2+ often contain the cation in a tetrahedral environment: it is small and highly charged, and so has a high polarizing power and tends to form bonds with a high degree of directionality, ie. it displays a high covalency