Chambers, Holliday. Modern inorganic chemistry

126 GROUPS I AND II

The alkali metals have the interesting property of dissolving in some non-aqueous solvents, notably liquid ammonia, to give clear coloured solutions which are excellent reducing agents and are often used as such in organic chemistry. Sodium (for example) forms an intensely blue solution in liquid ammonia and here the outer (3s) electron of each sodium atom is believed to become associated with the solvent ammonia in some way, i.e. the system is Na+(solvent) + e"(solvent).

The solution is energetically unstable (Chapter 3); the sodium slowly reacts with the ammonia solvent thus:

2Na + NH3 -* 2NaNH2 + H2t

sodium amide (sodamide)

(a reaction which can be written 2e~ + 2NH3 -» 2NH^ + H2|). This reaction is catalysed by such ions as iron(III) and should be compared to the reaction with water

2Na + 2H2O -> 2NaOH + H2|

COMPOUNDS OF GROUP I AND II ELEMENTS

GENERAL

For the most part it is true to say that the chemistry of the alkali and alkaline earth metal compounds is not that of the metal ion but rather that of the anion with which the ion is associated. Where appropriate, therefore, the chemistry of these compounds will be discussed in other sections, for example nitrates with Group V compounds, sulphates with Group VI compounds, and only a few compounds will be discussed here.

THE HYDRIDES

All Group I and II elements, except beryllium, form hydrides by direct combination with hydrogen. The hydrides of the metals except those of beryllium and magnesium, are white mainly ionic solids, all Group I hydrides having the sodium chloride lattice structure. All the hydrides are stable in dry air but react with water, the vigour of the reaction increasing with the molecular weight of the hydride for any particular group.

MH + H2O -> MOH -f H2 T

MH2 + 2H2O -> M(OH)2 4- H2|

Hydrolysed by water

THE OXIDES AND HYDROXIDES

The white solid oxides M^O and M"O are formed by direct union of the elements. The oxides M!2O and the oxides MUO of calcium down to radium have ionic lattices and are all highly basic; they react exothermically with water to give the hydroxides, with acids to give salts, and with carbon dioxide to give carbonates. For example

Na2O + H2O -» 2NaOH

BaO + CO2 -> BaCO3

Magnesium oxide is almost inert towards water, but dissolves in

+ CO2-*HCOJ

The reaction between Ca(OH)2 + CO2 to produce sparingly soluble CaCO3 is the common test for carbon dioxide.

Beryllium hydroxide is obtained as a white gelatinous precipitate when OH~ ions are added to a solution of a beryllium salt. It is only sparingly soluble in water, and is weakly basic, dissolving in strong acids to give the hydrated beryllium ion [Be(H2O)4]2+ , but also dissolving in solutions containing the hydroxide ion to give the tetrahydroxoberyllateill) ion [Be(OH)4]2" ; addition of acid first reprecipitates the hydroxide Be(OH)2 (as a white gelatinous hydrated precipitate) and then re-dissolves it to give the hydrated ion ; hence we have the sequence*

This behaviour distinguishes beryllium hydroxide from the other hydroxides of Group II which are not amphoteric; this amphoterism is also shown by aluminium hydroxide in Group III, and it has been discussed more fully in Chapter 2, where we saw it as characteristic of small ions of high charge, i.e. Be2+ and A13+ .

The other Group II hydroxides are sparingly soluble in water, the solubility increasing down the group ; magnesium hydroxide is

* The species involved are more complicated than this sequence indicates, see note on p. 46; the simplified representation is. however, quite adequate.

132 GROUPS) AND I!

THE CARBONATES AND HYDROGENCARBONATES

As with the hydroxides, we find that whilst the carbonates of most metals are insoluble, those of alkali metals are soluble, so that they provide a good source of the carbonate ion COf ~ in solution; the alkali metal carbonates, except that of lithium, are stable to heat. Group II carbonates are generally insoluble in water and less stable to heat, losing carbon dioxide reversibly at high temperatures.

Table 6.4

DECOMPOSITION TEMPERATURES* ( K ) OF SOME CARBONATES

* The temperature at which the pressure of CO2 reaches 1 atmosphere.

A further peculiarity of the Group I and II carbonates is the ability to form the hydrogencarbonate or bicarbonate ion HCOa:

CO?" + H3O+ ^ HCOJ + H2O

This ion is produced by the prolonged passage of carbon dioxide through neutral or alkaline solutions containing Group I or II ions (except lithium or beryllium which do not form a hydrogencarbonate). The hydrogencarbonates of Group 1 elements can be isolated as solids but these solids readily decompose when heated to form the carbonate with the evolution of carbon dioxide and water, for example

2NaHCO3 -> Na2CO3 + H2O + CO2

Group II hydrogencarbonates have insufficient thermal stability for them to be isolated as solids. However, in areas where natural deposits of calcium and magnesium carbonates are found a reaction between the carbonate, water and carbon dioxide occurs:

M"CO3 + CO2 + H2O -> M2 + + 2H(X>3

Insoluble In solution

This produces sufficient concentrations of magnesium and calcium ions to render the water hard. The above reaction is readilyreversed by boiling the water when the magnesium and calcium ionsresponsible for the hardness are removed as the insoluble carbonate.

GROUPS I AND II 133

Some carbonates are important industrial chemicals. Calcium carbonate occurs naturally in several forms, including limestone, and is used in the production of quicklime, calcium oxide CaO, slaked (or hydrated) lime, calcium hydroxide Ca(OH)2 and cement.

Several million tons of sodium carbonate are used every year, almost one third of this being used in glass making and the rest being used for a variety of purposes including paper manufacture, chemicals, and as a water softener in soap powder. Sodium sesquicarbonate, Na2CO3 . NaHCO3 . 2H2O, occurs naturally in theUS and approximately 1000 000 tons of sodium carbonate are produced from this annually. Until recently almost all the sodium carbonate required commercially in the UK (5000 000 tons annually) was manufactured by the soda-ammonia process but some is now produced by carbonation of sodium hydroxide, surplus to requirements, made during the electrolysis of brine :

2NaOH + CO2 -» Na2CO3 + H2O

The soda-ammonia process occurs in two main stages. First, brine is saturated with ammonia gas and this "ammoniacal brine' is then treated with carbon dioxide. The equilibrium

CO2 4- 2H2O ^HCO3~ + H3O+

is moved to the right by the competition of the ammonia for protons. i.e. NH3 + H3O+ ?± NH + 4- H2O. The ions then present are NH^. HCOa, Cl~ and Na+ and the least soluble salt sodium hydrogen carbonate, is precipitated when ionic concentrations increase, and is removed by vacuum filtration.

When heated, sodium hydrogencarbonate readily decomposes evolving carbon dioxide, a reaction which leads to its use as baking powder when the carbon dioxide evolved 'aerates' the dough. In the soda-ammonia process the carbon dioxide evolved is used to supplement the main carbon dioxide supply obtained by heating calcium carbonate :

CaCO3 -* CaO 4- CO2

The calcium oxide so produced is slaked to give a suspension of calcium hydroxide and this is heated with the filtrate from the carbonator which contains ammonium chloride:

2NH4C1 + Ca(OH)2 -> CaCl2 + 2NH3f + 2H2O

The ammonia gas is used again and the only by-product, calcium chloride, is used to melt snow, prevent freezing of coal in transit and as an antidust treatment since it is hygroscopic and forms a solution of low freezing point.

134 GROUPS I AND II

ABNORMAL PROPERTIES OF LITHIUM AND

BERYLLIUM

As any group is descended the size of the atom and number of electrons shielding the outer electrons from the nucleus increases and the ionisation energy falls (see Table 6.2.)

Shielding of the outer electrons is least for the small lithium and beryllium atoms and their ionisation energies are consequently higher than other members of their respective groups. In the case of beryllium the higher ionisation energy results in the bonding in many beryllium compounds being covalent rather than ionic. (This tendency is shown to a much lesser extent by magnesium which forms some covalent compounds.)

The small lithium Li+ and beryllium Be2+ ions have high chargeradius ratios and consequently exert particularly strong attractions on other ions and on polar molecules. These attractions result in both high lattice and hydration energies and it is these high energies which account for many of the abnormal properties of the ionic compounds of lithium and beryllium.

In view of the ionisation energies the electrode potentials for

Ionisation energy refers to the process Li(g) -* Li+ (g) -f e~ .whereas the electrode potential measured in aqueous solution also includes

lithium compensates for the high ionisation energy. The value of the second ionisation energy of beryllium (the energy to remove the second electron) is so great that even the large hydration energy of the Be2+ cannot compensate for it, and E^~ is less negative.

The hydroxide of lithium, although soluble in water, is a weak base owing to the great attraction between the Li^ and OH~ ions (p. 74); the hydroxide of beryllium is really a neutral, insoluble beryllium complex [Be(OH)2] (p.45).

L (H20)J

When considering the fluorides, the high hydration energy of the small fluoride ion, F", must also be considered (p. 78). The lattice energy of beryllium fluoride is high but the combined hydration energies of the Be2+ and F~ ions are sufficient for the BeF2 to dissolve, whilst the other fluorides of Group II elements having lower M 2 + hydration energy are insoluble in spite of lower lattice

energies. The insolubility of lithium fluoride results from the high lattice energy which in this case is not exceeded by the combined hydration energies. Other Group I fluoridesdissolve since the lattice energies are smaller and are exceeded by the combined hydration energies.

In this discussion, entropy factors have been ignored and in certain cases where the difference between lattice energy and hydration energy is small it is the entropy changes which determine whether a substance will or will not dissolve. Each case must be considered individually and the relevant data obtained (see Chapter 3), when irregular behaviour will often be found to have a logical explanation.

The abnormal properties of lithium and beryllium are summarised in Tables 6.5 and 6.6.

Table 6.6