Nitrogen(II) Oxide: Nitric Oxide

Nitrogen(II) oxide (Nitric oxide), NO, has a great affinity for oxygen as such it's difficult to find it in its pure form. That is, it reacts immediately with oxygen in air even at mild reaction conditions to form the reddish-brown nitrogen(IV) oxide gas.

A unique observation is the inability to tell the actual smell of the gas as it immediately become nitrogen(IV) oxide when exposed to air. Also, it is a colourless and poisonous gas that is neutral to moist litmus paper, insoluble in water and slightly denser than air.

Nitrogen(II) oxide is prepared from the reaction of 50% trioxonitrate(V) acid, HNO₃, with most metals in a process that forms undesirable brown fumes of nitrogen(IV) oxide. The nitrogen(IV) oxide is formed from reaction of the nitrogen(II) oxide product with the oxygen content of the preparation flask.

2NO(g) + O₂(g) → 2NO₂(g)

The nitrogen(IV) oxide is separated from the Nitrogen(II) oxide gas by bubbling the gas through water which dissolve the nitrogen(IV) oxide.

In laboratory preparation, 50% trioxonitrate(V) acid is reacted with copper turnings to liberate nitrogen(II) oxide that is collected by downward delivery.

3Cu(s) + 8HNO₃(aq) → 3Cu(NO₃)₂(aq) + 4H₂O(l) + 2NO(g)

Reaction with iron(II) tetraoxosulphate(VI)

The reaction with iron(II) tetraoxosulphate(VI) solution is essential for the brown ring test for trioxonitrate(V) and for test for the presence of nitrogen(II) oxide in an unknown sample. Here, the nitrogen(II) is dissolved in the tetraoxosulphate(VI) solution to form a dark solution which decomposes on heating. This reaction is essential for the purification and removal of nitrogen(II) oxide in a mixture of gasses.

FeSO₄(aq) + NO(g) ↔ FeSO₄.NO(aq)

As a reducing agent

Nitrogen(II) oxide acts as a reducing agent, but a weak reducing agent. The reduction property of the gas is demonstrated in the presence of potassium tetraoxomanganate(VII) where it slowly decolourizes the acidified solution reducing manganate(VII) to manganese(II).

3MnO₄^-(aq) + 4H⁺(aq) + 5NO(g) → 3Mn²⁺(aq) + 5NO₃^-(aq) + 2H₂O(l)

You should also observe the oxidative transformation where the nitrogen(II) oxide is oxidised to trioxonitrate(V).

Action of heat

Compared to nitrogen(I) oxide that rekindles a strongly glowing splint, nitrogen(II) oxide gas actually extinguishes both glowing splint and burning sulphur. But it is known to support the combustion of strongly burning substances such as burning lead or magnesium forming nitrogen gas and the oxides of the metals. Magnesium, for instance, will continue burning in the gas to produce nitrogen gas and magnesium oxide.

2Mg(s) + 2NO(g) → 2MgO(s) + N₂(g)

An important combustion reaction is the thermal decomposition of the nitrogen(II) oxide itself. Here, equal volume of nitrogen and oxygen gas is formed when the gas is heated to about 1000°C.

2NO(g) + N₂(g) + O₂(g)

Reduction in the presence of hot metals

Apart from being a weak reducing agent, nitrogen(II) oxide can be reduced to nitrogen in the presence of hot metals. Heated copper is reduces nitrogen(II) oxide to nitrogen gas with the formation of copper(II) oxide.

2Cu(s) + 2NO(g) → 2CuO(s) + N₂(g)

Reaction with Oxygen

Nitrogen(II) oxide also reacts directly with oxygen at room temperature to form brown fumes of nitrogen(IV) oxide.

2NO(g) + O₂(g) → 2NO₂(g)

Exposure to air

The first way to identify nitrogen(II) oxide is expose it to open air. A sample is suspected to be nitrogen(II) oxide if it turns reddish-brown when exposed to air. Usually, the unknown sample is initially trapped in a test-tube by a stopper. Once the test-tube is removed, nitrogen(II) oxide is immediately oxidised to nitrogen(IV) oxide which is reddish-brown in colour.

Test with iron(II) tetraoxosulphate(VI)

This is a confirmatory test for nitrogen(II) oxide. Here, a dark brown solution is form when acidified iron(II) tetraoxosulphate(VI) solution is poured into the test-tube containing the unknown sample. It is essential to state that the solution is acidified with dilute tetraoxosulphate(VI) acid.