Chemical elements
  Nitrogen
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    Physical Properties
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      Nitrogen Chloride
      Nitrogen Iodide
      Monochloramine
      Nitrosyl Fluoride
      Nitrosyl Chloride
      Nitrosyl Bromide
      Nitryl Fluoride
      Nitryl Chloride
      Di-imide
      Nitramide
      Nitrohydroxylamine
      Hyponitrous acid
      Nitrous Oxide
      Nitric Oxide
      Nitrogen Trioxide
      Nitrogen Tetroxide
      Nitrogen Pentoxide
      Nitroso-nitrogen Trioxide
      Nitrous Acid
      Pernitric Acid
      Sulphur Nitride
      Pentasulphur Dinitride
    Ammonia
    Hydroxylamine
    Hydrazine
    Azoimide
    Nitric Acid

Nitrous Acid, HNO2





History

As early as 1777 it was realised that there was more than one "acid of nitre." Scheele distinguished " phlogisticated acid of nitre " from nitric acid as being a weaker volatile acid produced by the reduction of nitric acid. He also showed that nitre, when strongly heated, lost oxygen, and left a deliquescent salt which readily decomposed into a volatile acid when treated with acid. Priestley had previously described the brown fumes produced by oxidising nitric oxide as "nitrous acid vapour," while the terms "nitrous acid gas " and " nitrous acid " were used later by Davy and Gay-Lussac respectively. Confusion of terms, however, existed, due to the fact that both nitric and nitrous acids were present, and the means of distinguishing the two were not available at that time. It was understood clearly, however, that there were two distinct salts, nitrates and nitrites, and Cavendish showed that silver nitrite was precipitated when potassium nitrite was added to a solution of silver nitrate. Gay-Lussac was the first to prepare nitric and nitrous acids by the careful oxidation of nitric oxide with oxygen in the presence of water.

Nitrous acid is an unstable compound, and all the methods of preparation yield an aqueous solution of the acid.


Preparation

  1. Nitrogen trioxide, N2O3, which is nitrous anhydride, produces a weak solution of nitrous acid when treated with ice- cold water:

    N2O3 + H2O = 2HNO2.
  2. Dilute hydrochloric acid decomposes silver nitrite with the liberation of nitrous acid:

    AgNO2 + HCl = AgCl + HNO2.
  3. Alkali nitrites similarly decompose with dilute acids, and this affords a convenient laboratory method. Nitrites may be prepared by reducing nitrates with metals, or sulphites, or electrolytically; also by absorption of nitrogen trioxide by alkalies.
  4. Oxidation of ammonia with hydrogen peroxide produces nitrous acid, but there is always some ammonium nitrite present in the solution:

    NH3 + 3H2O2 = HNO2 + 4H2O.
  5. Nitric oxide passed into nitric acid reduces the latter to nitrous acid:

    HNO3 + 2NO + H2O = 3HNO2.

Physical Properties of Nitrous Acid

The aqueous solution of nitrous acid is blue in colour, which quickly fades with evolution of brown fumes, leaving a solution containing only nitric acid:

3HNO2HNO3 + 2NO + H2O.

The decomposition of nitrous acid is unimolecular, and the strongest solution which can be obtained at 0° C. is 0.185N, prepared from the decomposition of barium nitrite with dilute sulphuric acid.

The aqueous solution is more stable when kept at low temperatures and small concentration, and also when under pressure of nitric oxide. The decomposition of a cold dilute solution follows the reaction

3HNO2HNO3 + 2NO+H2O,

whereas stronger solutions at higher temperatures decompose according to the equations

2HNO2N2O3 + H2ONO + NO2 + H2O.

Other factors also influence the decomposition, such as agitation, surface area, and presence of nitric acid. The decomposition of nitrous acid in dilute acid solutions has been studied.

The heat of formation of nitrous acid is as follows:

N2O3, Aq. = -6,820 calories.
H, N, O2, Aq. = 30,770 calories
2NO, O, Aq. = 36,330
H, NO, O, Aq. = 52,345 calories

N2, 2H2O = 71,770 calories (product as NH4NO2) Heat of solution = 4,700 calories

The decomposition of nitrous acid in dilute solution is attended with the absorption of 18.4 Cals.

The heat of neutralisation with ammonia is 9.100 Cals., and with barium hydroxide 10.600 Cals.

The velocity constant of decomposition is 0.00014 at 0° C., 0.00022 at 21° C., and 0.00057 at 40° C.

The dissociation constant at 0° C. is 6×10-4 and the calculated mobility of the NO2 ion is 64.5, that determined from AgNO2 being 63, and from Ba(NO2)2 61.7.

The electrical conductivities at 25° C. for dilutions of 512, 1024, and 1536 are 150.7, 189.0, and 217.0 respectively.

Chemical Properties of Nitrous Acid

Nitrous acid functions both as a reducing and an oxidising agent. Thus all the ordinary oxidising agents, such as hydrogen peroxide, permanganates, chromates, ozone, bromine water, are reduced and nitric acid is the product:

HNO2 + O = HNO3.

On the other hand, many reducing agents are oxidised, the primary decomposition being:

2HNO2 = 2NO + O + H2O.

Stannous chloride is converted into stannic chloride, sulphuretted hydrogen into sulphur, sulphur dioxide into sulphur trioxide. Iodine is liberated from potassium iodide,

2KI + HNO2 ⇔ 2KOH + 2NO + I2,

or in acid solution,

2HI + HNO2 = 2H2O + 2NO + I2,

and this reaction is used in the detection of nitrous acid. Many organic colouring matters are bleached by a process of oxidation, e.g. indigo, litmus, and methyl orange.

A large number of secondary decomposition products of the reduction of nitrous acid include nitrous oxide, hyponitrous acid, hydroxylamine, nitrogen, and ammonia. Urea and nitrous acid react to give nitrogen and carbon dioxide:

CO(NH2)2 + 2HNO2 = 2N2 + CO2 + 3H2O.

The interaction of hydrazine and nitrous acid seems to be of a complex nature, and the following equations have been given to represent the reactions:

N2H4 + HNO2 = N2O + NH3 + 3H2O;
NOH4 + 2HNO2 = N2O + N2 + 3H2O.

Azoimide and nitrous acid yield nitrous oxide and nitrogen: N3H + HNO2 = N2O + N2 + H2O.

Concentrated sulphuric acid forms nitrosyl-sulphuric acid with nitrous acid (or the anhydride, N2O3):

HNO2 + H2SO4 + H2O.

Sulphurous acid shaken with nitrous acid yields hydroxylamine-disulphonic acid:

HNO2 + 2H2SO3OH.N.(SO2OH)2 + H2O.

There is no action between nitrites and normal sulphites, but with acid sulphites various salts of hydroxylamine-sulphonic acids are produced.

Constitution of Nitrous Acid. - Nitrous acid resembles nitric acid inasmuch as it shows two kinds of absorption spectra. There is probably an equilibrium between two tautomeric forms:

[NO2]H and .

Thus silver nitrite appears to exist in two forms, one an ionisable salt, [NO2]Ag, and the other a non-electrolyte, . This will account for the fact that the product of the reaction between AgNO2 and CH3I is methyl nitrile, , an ester the hydrolysis of which shows that the methyl group is attached to the oxygen atom; and also of nitro-methane, [NO2]CH3, a nitro-paraffin the reduction of which shows that the methyl group is attached to the nitrogen atom.

The modern valency theory represents the tautomerism of nitrous acid and nitrites as due to the shift of the hydrogen atom from oxygen to nitrogen:

NO'2 + HHNO2;

or, adopting the theory of the cubic arrangement of octets:

or

Detection and Estimation

The liberation of iodine from potassium iodide, which gives a blue colour in the presence of starch, is a delicate but not characteristic test for nitrous acid or nitrites in acid solution.

Metaphenylene diamine in hydrochloric acid produces a brown colour. A very sensitive reagent is an acetic-acid solution of sulphanilic acid and β-naphthylamine (Griess-Ilosvay reagent), which gives a pink colour with nitrites, sensitive to 1 part per million of water.
  1. The estimation of nitrites and nitrous acid may be carried out by direct titration with standard potassium permanganate. The solution is run from a burette into a standard solution of permanganate until the colour is just discharged.

    A modification of this method is to add a known excess of standard permanganate solution to the nitrite, acidify with sulphuric acid, liberate iodine by adding potassium iodide, and titrate with sodium thiosulphate, which gives the unused permanganate.
  2. The liberation of iodine from potassium iodide by nitrites may be used directly, and the iodine estimated with thiosulphate. The solution should be allowed to stand for two minutes before titration with the thiosulphate, otherwise high results are obtained.
  3. A nitrometer method of estimation may be used where presence of organic acids rules out the permanganate method. The nitric oxide evolved is directly measured, but the reaction requires the presence of potassium ferrocyanide, and also of one organic acid such as acetic, tartaric, citric, or oxalic:

    K4Fe(CN)6 + KNO2 + 2CH3COOH = K3Fe(CN)6 + 2CH3COOK + NO + H2O.

    Two factors likely to cause errors in this method are the solubility of the nitric oxide in the aqueous solution, and the vapour pressure of the acetic acid.
  4. Among other methods which have been proposed is the decomposition of nitrites by hydrazine sulphate and measuring the volume of nitrogen, two-thirds of which is due to nitrite; also the reaction with hydroxylamine hydrochloride and titration with sodium hydroxide before and after the reaction.


The reaction with azoimide, yielding nitrogen and nitrous oxide, is also quantitative in dilute solution, and is used in the estimation of nitrites. Nitrites may be determined gravimetrically by a reaction with an excess of silver bromate and acetic acid. The nitrous acid forms an equivalent amount of silver bromide by reduction.
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