AMMONOLYSIS

Ammonolysis

$\begin{array}{l}{\left[\text{M}{\left({\text{H}}_2\text{O}\right)}_n\right]}^{x+}\stackrel{{\text{H}}_{\text{2}}\text{O}}{\rightleftharpoons }{\left[\text{M}{\left({\text{H}}_{\text{2}}\text{O}\right)}_{n-1}\left(\text{OH}\right)\right]}^{(x-1)+}+{\text{H}}_{\text{3}}{\text{O}}^{\text{+}}\hfill \\ {\text{BX}}_{\text{3}}{\text{ + H}}_2\text{O }\stackrel{{\text{H}}_{\text{2}}\text{O}}{\rightleftharpoons }{\text{ BX}}_2\left(\text{OH}\right)+{\text{H}}_{\text{3}}{\text{O}}^{\text{+}}+{\text{X}}^{-}\hfill \end{array}$

Similar equations can be written for ammonolytic reactions in liquid ammonia:

$\begin{array}{l}{\left[\text{M}{\left({\text{NH}}_3\right)}_n\right]}^{x+}\stackrel{{\text{NH}}_3}{\rightleftharpoons }{\left[\text{M}{\left({\text{NH}}_3\right)}_{n-1}\left({\text{NH}}_2\right)\right]}^{(x-1)+}+{\text{NH}}_4^{+}\hfill \\ {\text{BX}}_{\text{3}}+{\text{NH}}_{\text{3}}\stackrel{{\text{NH}}_3}{\rightleftharpoons }{\text{BX}}_{\text{2}}\left({\text{NH}}_{\text{2}}\right)+{\text{NH}}_{\text{4}}^{\text{+}}+{\text{X}}^{-}\hfill \end{array}$

It should be noted that the first equation of each of these pairs really represents only an acid–base reaction (in the Lowry–Brønsted sense), the solvated ion (acid) losing a proton to the solvent (base). In the second equation, the BX3 molecule loses an X and gains an OH or NH2 group from the solvent, by what may be termed a metathetic reaction involving the solvent. However, it is convenient to use solvolysis to describe both kinds of reaction. Two general considerations can be given before ammonolysis is discussed in detail. First, N—H bonds of ammonia are less easily broken than O—H bonds of water and hence ammonolysis is, in general, less likely to occur than hydrolysis. This has important consequences since it means that liquid ammonia can often be used in the preparation of substances which would undergo hydrolysis if the synthesis were attempted in aqueous solution; examples of this will be given in the next chapter. The metal–ammonia solutions to be discussed later in this chapter also owe their stability to the fact that the reaction (M = alkali metal):

${\text{M+NH}}_3\to {\text{MNH}}_{\text{2}}+\frac{1}{2}{\text{H}}_{\text{2}}$

occurs much less readily than the hydrolysis

$\text{M}+{\text{ H}}_2\text{O}\to \text{MOH + }\frac{1}{2}{\text{H}}_2$

Secondly, even if ammonolysis does occur, the rate at which it does so will depend upon temperatuie. In many cases, since the reactions in liquid ammonia can be carried out at temperatures as low as — 78°C, the rate of ammonolysis can be made slow enough not to matter especially where, for example, metathetic ionic reactions (which are still rapid at low temperatures) are carried out.

As with hydrolytic reactions, ammonolytic reactions will be favoured if an ammono-base is present such as potassium amide (cf. KOH); but here we shall assume that only solvent is involved.

Ionic compounds — salts of many of the metals — do not in general undergo ammonolysis, although the cations are ammonated in solution. Relatively little is known of equilibria such as

$\begin{array}{ccc}{\left[\text{M}{\left({\text{NH}}_3\right)}_6\right]}^{3+}& \rightleftharpoons & {\left[\text{M}{\left({\text{NH}}_3\right)}_5{\text{NH}}_2\right]}^{2+}+{\text{H}}^{+}\\ & & \Updownarrow \\ & & {\left[\text{M}{\left({\text{NH}}_3\right)}_4{\left({\text{NH}}_2\right)}_2\right]}^{+}+{\text{H}}^{+}\end{array}$

etc., where M is a transition metal.

With wholly or partly covalent compounds, ammonolysis is more common, the extent to which it occurs being determined in many cases by the same factors as affect hydrolysis. Thus to take an example from Group III of the Periodic Classification, the covalent BX3 compounds ammonolyse and hydrolyse readily when, for example, X = halogen or some similar electronegative group but not when X = alkyl; the same applies to the compounds of silicon (SiX4) in Group IV and phosphorus (PX3) in Group V. In these reactions, intermediate formation of an ammonia adduct precedes ammonolysis; for this reason the carbon tetrahalides do not undergo ammonolysis at low temperatures since carbon has no vacant orbital of suitable energy to accept electrons from an ammonia molecule. This behaviour is parallel to that observed with water; carbon tetrahalides are not easily hydrolysed. The compound Cl4.2NH3 formed when carbon tetraiodide reacts with ammonia at — 33°C is thought to be held together by purely van der Waals forces, the ammonia being lost when the solid is warmed up to room temperature. Rather similar behaviour to that just described is observed with many covalent halides of the transition metals such as titanium tetrachloride TiCl4, molybdenum pentachloride MoCl5 and the halides of tin(IV). In these ammonolyses, reaction may be complete

$\begin{array}{cc}{\text{BCl}}_3+{\text{NH}}_3\to & \left[{\text{Cl}}_3\overline{B}-\stackrel{+}{\text{N}}{\text{H}}_3\right]\\ & \downarrow {\text{NH}}_3\\ \text{C}1\text{B}{\left({\text{NH}}_2\right)}_2\stackrel{{\text{NH}}_3}{\leftarrow }& {\text{Cl}}_2{\text{BNH}}_2+{\text{NH}}_4\text{Cl}\\ \downarrow {\text{NH}}_3& \\ B{\left({\text{NH}}_2\right)}_3& \end{array}$

or incomplete as in

${\text{ZrCl}}_{\text{4}}+2{\text{NH}}_3\to {\text{ZrCl}}_3{\text{NH}}_2+{\text{NH}}_4\text{Cl}$

where no further ammonolysis takes place. In certain cases decomposition of the amide to the imide takes place as happens, for example, in the ammonolysis of germanium tetrachloride where the final product of the reaction is the imide

${\text{GeCl}}_4\stackrel{{\text{NH}}_3}{\to }\text{Ge}{\left({\text{NH}}_{\text{2}}\right)}_4\to \text{Ge}{\left(\text{NH}\right)}_2+{\text{2NH}}_3$

This step can be taken further since, on heating, several amides and imides can evolve ammonia until finally the nitride of the element remains.

In the presence of an excess of an ammono-base (i.e. an ionic amide) some of the ammonolytic products will dissolve to form soluble complexes. Thus in the presence of potassium amide, iridium(III) bromide in liquid ammonia gives first the amide Ir(NH2)3 and this then dissolves in excess of the potassium amide

${\text{Ir(NH}}_2\text{)}{}_3{\text{ + 3KNH}}_2\to {\text{ K}}_3{\text{Ir(NH}}_2\text{)}{}_6$

Here again there is a parallel with the amphoteric behaviour of some metal hydroxides in presence of excess hydroxyl ion.

$\begin{array}{l}{\left[\text{M}{\left({\text{NH}}_3\right)}_n\right]}^{x+}\stackrel{{\text{NH}}_3}{\rightleftharpoons }{\left[\text{M}{\left({\text{NH}}_3\right)}_{n-1}\left({\text{NH}}_2\right)\right]}^{(x-1)+}+{\text{NH}}_4^{+}\hfill \\ {\text{BX}}_{\text{3}}+{\text{NH}}_{\text{3}}\stackrel{{\text{NH}}_3}{\rightleftharpoons }{\text{BX}}_{\text{2}}\left({\text{NH}}_{\text{2}}\right)+{\text{NH}}_{\text{4}}^{\text{+}}+{\text{X}}^{-}\hfill \end{array}$