Hydroxylamine, NH₂OH

Hydroxylamine hydrochloride was discovered in 1865 by Lossen. This salt was prepared by passing nitric oxide into a solution of hydrochloric acid which was acting on tin with evolution of hydrogen:

2NO + 6H = 2NH₂OH.

Later the salts were prepared in the solid state, and subsequently the free base was isolated by heating the addition compound with zinc chloride, ZnCl₂.2NH₂OH, or better, by adding sodium methoxide to hydroxylamine hydrochloride in methyl alcohol, and distilling off first the alcohol and then the hydroxylamine at reduced pressure:

NH₂OH.HCl + CH₃ONa = NaCl + CH₃OH + NH₂OH.

The preparations of hydroxylamine generally depend directly or indirectly upon the reduction of more oxidised compounds.

The reduction of nitric oxide can also be effected by hydrogen in platinum sponge over 100° C., as well as by nascent hydrogen (vide supra). Nitric acid, nitrates, and nitric esters can also be reduced to hydroxylamine by nascent hydrogen formed during the solution of Mg, Zn, etc., in acids, and by the action of water on sodium amalgam, as also by sulphites, sulphides, and kathodic hydrogen. The reduction of nitrites by the same reagents gives a most serviceable method of preparing the salts (vide infra). Another method depends upon the reduction of nitro-compounds, such as nitro-methane and nitroform.

The Reduction of Nitrites by Sulphites

Crystalline salts formed by the interaction of nitrates and sulphites were prepared and analysed by Fremy. These were later shown to be alkali salts of sulphonic acid derivatives. The direct addition of sulphurous to nitrous acid could give dihydroxylamine sulphonic acid, (HO)₂N.SO₃H, or its anhydride nitroso-sulphonic acid, ON.SO₃H. Actually, when nitrous acid is added to an acid sulphite, or sulphurous acid is added to a nitrite, hydroxylamine disulphonic acid is produced:

HO.NO + 2H₂SO₃ = HO.N(SO₃H)₂ + H₂O.

The acid is unstable, but the potassium or sodium salts may be crystallised from the solutions obtained by the interaction of potassium nitrite and metabisulphite at 0° C.:

2KNO₂ + 3K₂S₂O₅ + H₂O = 2HO.N(SO₃K)₂ + 2K₂SO₃.

The solution, at first neutral, becomes alkaline, and the potassium salt crystallises in rhombic prisms. On acid hydrolysis this gives a salt of hydroxylamine:

HO.N(SO₃H)₂ + 2H₂O = HO.NH₂.H₂SO₄ + H₂SO₄.

A concentrated solution of sodium nitrite (2 mols.) is mixed with one of Na₂CO₃ (1 mol.), and sulphur dioxide is passed into the mixture until it is slightly acid. The solution containing HO.N(SO₃Na)₂ is warmed with a little sulphuric acid in order to effect the hydrolysis, which is completed over 90° C. The resulting solution is neutralised with sodium carbonate, evaporated to small volume, when sodium sulphate (the decahydrate) crystallises. From the mother-liquor a hydroxylamine sulphate crystallises on cooling.

Electrolytic Preparation

Nitric acid may be reduced to hydroxylamine at a kathode of mercury or amalgamated lead in strongly acid solution, containing say 50 per cent, sulphuric acid and 25 per cent, hydrochloric acid. The nitric acid (50 per cent.) is run slowly into this cooled solution. The yield is over 80 per cent.

The base may be prepared in methyl-alcoholic solution as described above. Sodium ethoxide and a solution of ethyl alcohol may also be used, the yield being over 40 per cent. By using an alcohol of higher molecular weight the yield is increased. Thus the fraction distilled from commercial butyl alcohol at 115.5° to 117.5° C. is converted into sodium butylate. This is added to hydroxylamine hydrochloride in more of the butyl-alcohol fraction, using phenolphthalein as an indicator. The sodium chloride is filtered off and the filtrate cooled to 0° C., when hydroxylamine crystallises in large white flakes. It may be purified by distillation under reduced pressure. The base may also be prepared by distillation of the tribasic phosphate, (NH₃O)₃PO₄, in a vacuum. The yield is about 40 per cent.

Hydroxylamine is a colourless solid, crystallising in leaves or needles and having a density at 0° C. of 1.2255 (Bruhl). It melts at 32° to 33° C. and boils at 56° to 58° C. under 22 mm. pressure. The density of the vapour under these conditions shows a normal molar weight. The refractive indices for several wave-lengths have been determined; for the sodium line, n = 1.44047. The molecular refraction for the sodium line is 7.228.

Pure hydroxylamine is deliquescent in air, very soluble in water and in methyl and ethyl alcohols, but scarcely dissolves in typical organic solvents such as ether and chloroform. It may be recrystallised, however, from boiling ether. The solid is very unstable and begins to decompose slowly even at ordinary temperatures, explosively over 100° C. It also decomposes in solution, especially in the presence of alkali. The products are NH₃, N₂, and, to a smaller extent, N₂O.

3NH₂OH = NH₃ + N₂ + 3H₂O (1)
4NH₂OH = 2NH₃ + N₂O + 3H₂O (2)

If more combined oxygen is present, as in the nitrate and nitrite, the ammonia is completely oxidised, and the products are nitric and nitrous oxides and water. A comparison of these with the decomposition products of ammonium nitrate and nitrite shows clearly that hydroxylamine is oxyammonia:

NH₂OH.HNO₃ = 2H₂O + 2NO,
NH₂OH.HNO₂ = 2H₂O + N₂O.

Decomposition of Salts by Heat

The hydrochloride decomposes in a somewhat similar manner at about 150° C. In the presence of alkali only a little nitrous oxide is formed, in accordance with equation (1) above. In the presence of acids the reaction is displaced in the direction of equation (2). The change may occur through intermediate compounds: first two molecules of hydroxylamine give hydroxyhydrazine, HO.HN.NH₂; this is converted into the diamide of nitrous acid, HO.N(NH₂), which is hydrolysed to ammonia and nitrous acid; the latter reacts with more hydroxylamine to give nitrous oxide. In the presence of hydrazine hydrochloride, hydroxylamine hydrochloride decomposes at about 150° C. The products are nitrogen and ammonium chloride:

2NH₂OH + N₂H₄.HCl = 2NH₄Cl + N₂ + 2H₂O.

Reactions of Hydroxylamine Salts

Hydroxylamine and its salts are powerful reducing agents, the products being oxides of nitrogen, nitrogen, or nitrous acid, according to circumstances. At the same time it is capable of acting as an oxidising agent, especially in alkaline solution, when it is all present as free base. The product of the reaction is then ammonia.

Atmospheric oxygen, especially in the presence of alkali, oxidises it rapidly to oxides of nitrogen. Strong oxidising agents, such as chlorine, potassium permanganate, etc., may cause inflammation. Salts of mercury, silver, gold, and platinum are reduced with precipitation of the metal. Cupric salts are reduced to cuprous oxide in alkaline solution:

4CuSO₄ + 2NH₂OH + 8NaOH = 2Cu₂O + N₂O + 4Na₂SO₄ + 7H₂O.

This affords a delicate qualitative test, and also a method of estimating hydroxylamine by means of Fehling's solution.1 As a qualitative test it is sensitive to 1 in 100,000 parts of solution. Ferrous iron in alkaline solution is oxidised to the ferric state.2 Thus:

NH₂OH + 2Fe(OH)₂+ H₂O = 2Fe(OH)₃ + NH₃.

In weakly acid solution ferric chloride is reduced, giving ferrous chloride and nitrous oxide. When boiled with ferric chloride and a moderate amount of acid it is oxidised quantitatively to N₂O. The titration of the ferrous salt with permanganate allows a quantitative determination of the hydroxylamine,

2NH₂OH + O₂ = 3H₂O + 4N₂O.

If there is only a small excess of ferric salt the oxidation leads to NO:

2NH₂OH + 30 = 3H₂O + 2NO.

In concentrated solutions of sulphuric or phosphoric acids it reacts with ferrous sulphate, giving ferric and ammonium sulphates. It may also be directly oxidised by means of potassium permanganate.

The heat of formation of hydroxylamine from its elements is positive and about twice as great as that of ammonia. The value deduced from the heat of combustion of hydroxylamine nitrate is 23,700 cals. per mol. That deduced from the heat of reduction of the hydrochloride by silver nitrate is 24,290 cals.

The heat of neutralisation with strong acids, e.g. HCl, in dilute solution is much less - 9260 cals. - than that of ammonia - 12,300 cals. This is connected with the fact that hydroxylamine is a weaker base than ammonia (vide infra). The solution only reacts slightly alkaline.

The heat of formation of the hydrochloride, NH₂OH.HCl, from its elements is 75,500 cals. or 76,510 cals., about the same as that of solid ammonium chloride. The greater heat of formation of the free base is almost exactly compensated by the smaller heat of combination with hydrochloric acid.

In addition to the qualitative tests already mentioned, the following have been described. Sodium nitroprusside gives a magenta colour with a neutralised solution. When treated with yellow ammonium sulphide and ammonia it gives an evanescent purple colour. The appearance of this is accelerated by the addition of one or two drops of N/10 manganous sulphate. This test will show 0.00005 per cent, of hydroxylamine. Another test consists in the addition of an ammoniacal solution of diacetyl monoxime, which is converted into the dioxime by the hydroxylamine, and this is then identified by the formation of a scarlet precipitate with a solution of nickel salt.

In the presence of sodium carbonate or sodium phosphate hydroxylamine is quantitatively oxidised by iodine, and in the absence of other reducing agents it may be determined by means of this reaction. It is also oxidised by iodine in acetic acid solution, but in the presence of hydrochloric acid the reaction is a balanced one; hydroxylamine will liberate iodine from hydrogen iodide.

Hydroxylamine forms crystallised salts with acids; these are extensively hydrolysed and react acid in aqueous solution. Hydroxylamine is a monacid base, but these salts are only formed in presence of excess of acid. Many characteristic salts, with two or three molecules of the base attached to a monobasic acid, have been prepared. The salts are in most cases moderately soluble in water, very sparingly in ethyl alcohol, insoluble in ether, etc.

The monohydrochloride, NH₂OH.HCl, is formed in the presence of an excess of hydrochloric acid. It may be crystallised in pointed crystals or leaflets from hot alcohol. The density of the solid is 1.67. The solubility (grams in 100 grams of water) is 45.57 at 17° C. It melts with decomposition at about 150° C. When a solution of 20 grams of the free base in 260 grams of ethyl alcohol is added to one containing 42 grams of the hydrochloride in 42 c.c. of water, a crystalline precipitate of dihydroxylamine hydrochloride, (NH₂OH)₂.HCl, is thrown down.

The crystals after washing with alcohol and ether melt at 85° C. They are easily soluble in water. A tribasic salt, (NH₂OH)₃HCl, can be prepared by similar methods. Hydroxylamine hydrobromide is made by the double decomposition of the sulphate and barium bromide. It is rather soluble in water. A dihydroxylamine hydrobromide can be made by adding more of the free base. Hydroxylamine hydriodide can be crystallised by evaporation in vacuo. It explodes when heated to 83° or 84° C. The dibasic and tribasic hydriodides are more stable.

Hydroxylamine sulphate is well crystallised, either in the monoclinic or in the triclinic system. The solubility in water is high; according to Adams, 100 parts of water dissolve the following weights: -

t °C	-8	0	10	20	30	40	50	60	90
w	0.307	0.329	0.366	0.413	0.441	0.482	0.522	0.560	0.685

The salt can replace ammonium sulphate in alums, such as aluminium, ferric, and chromic alums. The bisulphate crystallises in prisms from a solution containing the theoretical proportions of the base and acid. The nitrate, NH₂OH.HNO₃, is prepared in solution by the action of silver nitrate on the hydrochloride or barium nitrate on the sulphate. It is very soluble in water, which only yields the salt in a hygroscopic mass after evaporation and strong cooling. It is decomposed on heating.

The phosphites, phosphates, sub-phosphates, arsenates, vanadates, the dithionate, and the salts of some organic acids, are known, as well as the ammonium or potassium hydroxylamine salts of some of these acids. Trihydroxylamine phosphate is more sparingly soluble in water than the salts considered above, the solubility at 20° C. being only 1.9 grams in 100 grams of water. The solution reacts acid. It is noteworthy that the tribasic ammonium phosphate is more difficult to prepare. This fact, taken in conjunction with the easy decomposition of the salt, has led to the belief that it should be formulated as hydroxylamine-dihydroxylamine-hydrogen phosphate,

NH₂OH (NH₃OH)₂HPO₄.

Hydroxylamine forms addition compounds, analogous to the ammines, with salts of the alkaline earths (including magnesium) and with those of zinc, cadmium, mercury, manganese, cobalt, and nickel. Hydroxylammines of silver and copper, if formed, cannot be isolated on account of the reduction to metals by the hydroxylamine. The derivatives of trivalent cobalt, e.g. Co(NH₂OH)₆Cl₃, are quite analogous to the ammines, Co(NH₃)₆Cl₃ [Co(NH₃)₅X]X₂, etc.

In the platinum series both platinic and platinous derivatives are known. The compound represented by the co-ordinated formula [Pt(NH₂OH)₄]Cl₂ dissociates electrolytically; but [Pt(NH₂OH)₄](OH)₂ is not an electrolyte, while the corresponding ammine is one, and reacts alkaline. Also the plato-compound [Pt(NH₂OH)₂X₂] is not an electrolyte.