Inorganic Synthesis

Nitric acid preparation (2 methods)


Pure nitric acid HNO3 is a colorless liquid with m.p. -41.6 °C. It is a strong mineral acid with pKa value of -1.64 [1]. In 0.1M water solution it is 93% ionized [2]. Pure nitric acid undergoes autoprotolysis characterized by equilibria

HNO3 + HNO3 ⇄ H2NO3+ + NO3
H2NO3+ + HNO3 ⇄ NO2+ + H2O3+ + NO3

Nitric acid dissolves almost all metals except gold, platinum, rhodium, iridium, niobium and tantalum. Many electronegative metals do not dissolve in concentrated nitric acid as well – they form a thin protective layer of an oxide (passivation process). When dissolving noble metals nitric acid is reduced to nitrogen dioxide or nitrogen monoxide depending on its concentration. Nitric acid can be also reduced to ammonium nitrate. In presence of stronger acids it acts as a base. A mixture of concentrated nitric and sulfuric acid is used in organic chemistry for nitration of aromatic compounds.

2 H2SO4 + HNO3 ⇄ NO2+ + 2 HSO4 + NO3

It is commercialy available as an azeotropic mixture composed of 68% HNO3 and 32% of H2O but higher concentrations can be prepared. 90% Fuming nitric acid has density of 1.50 g/mL. Red fuming nitric acid is composed of 84% nitric acid and 13% od dinitrogen tetraoxide. Due to higher nitrogen oxide content it is red colored.

Nitric acid is industrially prepared via Ostwald process. Ammonia gas is first converted into nitrogen monoxide by catalytic oxidation over platinum catalyst at temperatures above 700 °C. In the second step nitrogen monoxide is oxidized by oxygen to nitrogen dioxide gas that is further reacted with water at low temperatures.

4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g)
2 NO(g) + O2(g) → 2 NO2(g)
3 NO2(aq) + H2O(l) → 2 HNO3(aq) + NO(g)

Reaction of nitrogen dioxide at low temperatures yields a mixture of nitrous and nitric acid. Nitrous acid is however not stable in a solution and spontaneously break down into nitric acid and nitrogen monoxide gas. An easy way to produce nitrogen dioxide gas is a reaction described below. Potassium nitrate is used as an oxidant.

Cu(s) + KNO3(s) + 2HCl(aq) → Cu(NO3)2(aq) + NO2(g) + 2KCl(aq) + 2H2O/l)
N2O4(g) + H2O(l) → HNO3(aq) + HNO2(aq)
3 HNO2(aq) → HNO3(aq) + 2 NO(g) + H2O(l)

Procedure I:

Equimolar amounts of copper metal and potassium nitrate were added to a 500 mL side-arm flask. Large excess of concentrated hydrochloric acid was poured into a solid mixture and the flask was enclosed by a stopcock. The mixture was gently heated until evolution of nitrogen dioxide gas appeared. Nitrogen dioxide gas was  bubbled through 25 mL of distilled water (ice-cold) for an half of hour. A blue mixture of nitrous/nitric acid that resulted was heated until the coloration disappeared leaving a solution of nitric acid. VIDEO

Procedure II:

A round-bottom flask was filled with potassium nitrate (20 g, 198 mmol) and 30 % excess of sulfuric acid diluted with water (v/v ratio 1:1). Distillation apparatus was set up and the round-bottom flask was heated until distillate started to pass-through a condenser. Temperature was kept between 115-120 °C. The end of distillation was indicated by formation of nitrogen dioxide gas in the round-bottom flask. The apparatus was cooled down and the flask was detached. 17.5 mL of distillate was collected. Concentration of prepared acid was determined by a volumetric analysis and also by a measurement of density by pycnometer. The value of 45.25 % was found in both cases.


[1] Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. p. 171. ISBN 978-0130399137.
[2] Gažo J. a kol.: Všeobecná a anorganická chémia, ALFA Bratislava 1981.

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