Chemical elements
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      Nitramide
      Nitrohydroxylamine
      Hyponitrous acid
      Nitrous Oxide
      Nitric Oxide
      Nitrogen Trioxide
      Nitrogen Tetroxide
      Nitrogen Pentoxide
      Nitroso-nitrogen Trioxide
      Nitrous Acid
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      Sulphur Nitride
      Pentasulphur Dinitride
    Ammonia
    Hydroxylamine
    Hydrazine
    Azoimide
    Nitric Acid

Hyponitrous acid, H2N2O2





Historical

It was shown by Divers in 1871, that the reduction of sodium nitrate with sodium amalgam gave a salt having the empirical composition NaNO, the acid being an isomer of nitramide. This salt can also be prepared by electrolytic reduction of a nitrite. The reactions of the ethyl ester showed that this is constituted as a diazo-compound:

C2H5-O-N=N-O-C2H5.

The ester reacts with water, giving C2H5OH, CH3.CHO, and N2.

A bibliography of hyponitrous acid has been compiled by Divers, to whose researches so much of our knowledge of this interesting compound is due.


Methods of Preparation of Hyponitrites

Four general methods of preparation are known, under which numerous special methods may be classified. There are also a few other methods of subsidiary importance.
  1. By the reduction of nitrites.
  2. By the action of nitrous acid on hydroxylamine (or oxidation of hydroxylamine).
  3. By the alkaline hydrolysis of hydroxylamine sulphonates.
  4. By the hydrolysis of some organic compounds containing the diazo-group.
  1. The reduction of sodium or potassium nitrite by sodium amalgam proceeds according to the equation

    2NaNO2 + 4H = Na2N2O2 + 2H2O.

    The amalgam is added to the nitrite in the ratio 4Na to 1KNO2. The solution is kept cold. When the gas evolution ceases, the solution is neutralised by acetic acid and silver nitrate is added. The precipitate of yellow Ag2N2O2 is washed in the dark with cold water, dissolved in cold dilute nitric acid, and reprecipitated with sodium carbonate solution. The purified Ag2N2O2 is again washed and dried in vacuo over concentrated sulphuric acid. The yield is not good.

    Or, Ba(NO2)2 may be reduced with sodium amalgam, and the Ag2N2O2 precipitated as above.
  2. Ferrous hydroxide will also reduce nitrates to hyponitrites. Pure ferrous sulphate is precipitated with milk of lime, and to the cooled

    mixture is added 1 mol. of NaNO3 for each 10 mols. of FeSO4. The hyponitrite is precipitated with silver nitrate.

    A solution of NaNO2 is added to a solution of NH2OH.H2SO4. The mixture is heated quickly to 60° C. and AgNO3 is added at once. This method does not give good yields, but it proves that hyponitrous acid is a dioxime. It proceeds as follows: -

    HO-NH2 + O=N-OHHO-N=N-OH + H2O.

    Hyponitrous acid has also been prepared by the oxidation of hydroxylamine with CuO, HgO, Ag2, and by the oxidation of hydroxylamine with N2O3 in methyl-alcoholic solution.
  3. The most serviceable method of preparation is that which proceeds from the interaction of sulphites and nitrites. The potassium salt of hydroxylamine ββ-disulphonic acid is partly hydrolysed by hot water, giving potassium bisulphate and the potassium salt of the β-mono-sulphonic acid, i.e. HO-NH-SO3K. When this salt is fused with alkalies and the melt dissolved in water, the hyponitrite may be precipitated by silver nitrate as the silver salt, or, by the addition of a large excess of alcohol, the alkali hyponitrite may be separated from the excess of alkali and hydroxylamine. The yield may be 60 per cent, of the theoretical:

    2HO-NH-SO3Na + 4KOH = K2N2O2 + 2KNaSO3 + 4H2O.
  4. The alkaline hydrolysis of organic substitution products of nitrosohydroxylamine gives, by intramolecular change, salts of the tautomeric hyponitrous acid:

    (CH3)2=N-CO-N(NO)-OH + KOH = NH(CH3)2 + CO2 + H-O-N=N-OK.

    Further directions and improvements in these methods of preparation have been described.

Properties of Hyponitrites

The normal alkali salts of the dibasic acid are soluble in water, and are hydrolysed, giving alkaline solutions. Normal salts of other bases are very slightly soluble. The acid salts are very unstable, like the free acid. Silver hyponitrite is a yellow amorphous substance, non-hygroscopic, which may be boiled in water without decomposition. In the dry state it decomposes at 100° C., giving AgNO3, and it explodes at 150° C. It dissolves in nitric acid, and is reprecipitated by alkalies. Acetic acid and hydrogen sulphide set free hyponitrous acid. It is decomposed by hot alkalies. It reacts with alkyl iodides, giving alkyl hyponitrites.

Preparation and Properties of the Free Acid

Solutions of the acid are obtained when the silver salt is treated with hydrochloric acid, nitric acid, hydrogen sulphide, or phosphoric acid. A large excess of the silver salt is rubbed in a mortar with dilute cold HCl and filtered quickly.

Silver hyponitrite is added in portions to an ethereal solution of hydrochloric acid until there is no more free hydrochloric acid. The filtered solution when evaporated in a desiccator leaves the free acid in the form of white leaflets, which easily explode when rubbed, especially in the presence of acid vapour or solid potassium hydroxide. The free acid is very soluble in water, and also dissolves in alcohol, ether, chloroform, and in benzene, but not in petroleum ether.

The molar weight, determined by the cryoscopic method in aqueous solution, is 59 (H2N2O2 requires 62 ).

The acid does not expel CO2 from carbonates. When titrated it behaves much like carbonic acid, one-half of the hydrogen being neutralised at the phenolphthalein end point. In fact, as an acid it appears to be about as strong as carbonic acid (see Conductivity below). On standing it slowly changes into nitrous and nitric acids, and when an aqueous solution is boiled it gives N2O. This hydrolysis is rapidly effected by sulphuric acid; nitrous oxide is the anhydride. The reaction is, however, not reversible:

H2N2O2H2O + N2O.

In the presence of nitric acid hyponitrous acid can be boiled without decomposition. When freshly prepared, it should give only a faint yellow colour with potassium iodide; this effect is probably due to the nitrous acid formed by its decomposition. It is very stable towards reducing agents, although sodium bisulphite, followed by zinc and acetic acid, gives hydrazine.

It is easily oxidised; potassium permanganate in acid solution converts it into nitric acid, in alkaline solution into a nitrite. Hyponitrites can be quantitatively titrated with permanganate.

The heat of formation is negative, and has been determined by oxidising the calcium salt:

2N + O + H2O aq. = H2N2O2 aq. -57,400 calories.

The conductivity of the free acid is low, approximating to that of carbon dioxide; that of the salts is high on account of considerable dissociation and some hydrolysis.

The following values of the equivalent conductivities refer to 0° C.:

Hyponitrous Acid.

V6.2212.1112.4422.3724.2244.7473.4789.48146.94
λ1.081.491.522.122.393.103.854.215.43


For the sodium and calcium salts the following values have been obtained:-

Vλ
Na2N2O21685.3268.32
CaN2O2120073.12


The maximum equivalent conductivity of the sodium salt is about 68.

By introducing the Kohlrausch values for the conductivities of the sodium and calcium ion at 0° C., and allowing for hydrolysis, the conductivity of the NO' ion is found to be 38.0 and 38.7 from the acid and its sodium salt.

Structure

That this compound contains the diazo-group is indicated by many of the reactions used in its preparation, especially by the hydrolysis of organic compounds known to contain this group.

It is in fact a dioxime, as is shown more especially by its preparation from nitrous acid and hydroxylamine.

The molar weight is decided by the cryoscopic method, and the dibasic character of the acid by the neutralisation experiments and conductivities and by the existence of dialkyl derivatives and esters.

It has been suggested that hyponitrous acid is the antidioxime



the isomeric nitramide being perhaps the syn-compound.
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