Nitrogen(IV) Oxide: Nitrogen Dioxide

Nitrogen(IV) oxide is a poisonous reddish-brown gas with irritating smell. It is acidic when dissolved in water as evident in how it turns moist blue litmus paper red. It is heavier than air and will become a yellow liquid at 21°C. But it is the unique mode of preparation and its chemical and physical properties that are most interesting.

Nitrogen(IV) oxide (nitrogen dioxide), NO₂, is prepared by the thermal decomposition of the trioxonitrate(V) salt of heavy metals. But it is lead(II) trioxonitrate(V) that is most suitable for the preparation. The reason is that, water of crystallization that can interfere with the preparation, is absent in the lead(II) trioxonitrate(V).

Lead(II) trioxonitrate(V) is heated strongly so that the crystals crackles, melt and decompose to lead(II) oxide, nitrogen(IV) oxide and oxygen.

2Pb(NO₃)₂(s) → 2PbO(s) + O₂(g) + 4NO₂(g)

An essential part of the preparation process is how the nitrogen(IV) oxide is liquefied in a U-tube immersed in a freezing mixture. The freezing process allows the nitrogen(IV) oxide to form as a green liquid in the U-tube, the lead(II) oxide as the reddish-brown left over in the heating test tube that turns yellow on cooling and the escapes as gas from the U-tube.

Reaction with water

Nitrogen(IV) oxide is regarded as mixed acid anhydride as it dissolves in water to form a mixture of two acids: dioxonitrate(III) acid, HNO₂, and trioxonitrate(V) acid.

2NO₂(g) + H₂O(l) → HNO₂(aq) + HNO₃(aq)

The dioxonitrate(III) acid decomposes at room temperature to form more trioxonitrate(V) acid, nitrogen(II) oxide, NO, and water.

3HNO₂(aq) → HNO₃(aq) + 2NO(g) + H₂O(l)

The entire nitrogen(IV) oxide can be converted to only trioxonitrate(V) acid if the product of its reaction with water is further reacted with excess oxygen. Here, the dioxonitrate(III) (and not the dioxonitrate(V) acid) is oxidized to trioxonitrate(V) acid.

2HNO₂(aq) + O₂(g) → 2HNO₃(aq)

Reaction with alkalis

The mixed acid anhydride nature of nitrogen(IV) oxide is also evident in its reaction with alkalis where corresponding mixture of dioxonitrate(III) and trioxonitrate(V) salts are formed with water. This is expressed ionically as,

2OH^-(aq) + 2NO₂(g) → NO₃^-(aq) + NO₂^-(aq) + H₂O(l)

Potassium hydroxide forms potassium trioxonitrate(V) salt and potassium dioxonitrate(III) salt.

2KOH(aq) + 2NO₂(g) → KNO₃(aq) + KNO₂(aq) + H₂O(l)

Sodium hydroxide also form its own mixture of sodium trioxonitrate(V) and sodium dioxonitrate(III) salt.

2NaOH(aq) + 2NO₂(g) → NaNO₃(aq) + NaNO₂(aq) + H₂O(l)

Action of heat

Nitrogen(IV) oxide is actually in transition between two molecular forms that allows it to transit between three colours at varying temperature. It exist as dinitrogen(IV) oxide, N₂O₄, at low temperature and nitrogen(IV) oxide dissociation product on warming.

Dinitrogen(IV) oxide is pale yellow in colour. This colour gradually transform to the reddish-brown colour of nitrogen(IV) oxide with increase in temperature. Dinitrogen(IV) oxide is actually in equillibrium with nitrogen(IV) oxide at temperature below 140°C, and the entire dinitrogen(IV) oxide must have been converted to nitrogen(IV) oxide at 140°C. The nitrogen(IV) oxide will in turn dissociate to nitrogen(II) oxide and oxygen at temperature above 140°C. This form a transition from pale yellow, N₂O₄,through reddish-brown, NO₂, to colourless, NO.

(pale yellow) N₂O₄(g) ↔ (reddish-brown) 2NO₂(g) ↔ (colourless) 2NO(g) + O₂(g)

Reduction by reducing agent

Nitrogen(IV) can also be reduced to lower oxide or nitrogen by reducing agents. Heated copper or iron can reduce nitrogen(IV) oxide to nitrogen gas. Copper will form copper(II) oxide, CuO, and nitrogen gas.

4Cu(s) + 2NO₂(g) → 4CuO(s) + N₂(g)

Hydrogen sulphide, H₂S, is another reducing agent that can reduce the nitrogen(IV) oxide to nitrogen(II) oxide as the sulphur(IV) oxide is oxidized to sulphur, S.

H₂S(g) + NO₂(g) → H₂O(l) + NO(g) + S(s)

Carbon(II) oxide can also act as a reducing agent where it reduces nitrogen(IV) oxide to nitrogen(II) oxide as carbon(IV) oxide, CO₂, is formed in the process.

CO(g) + NO₂(g) → CO₂(g) + NO(g)

Support of combustion

Nitrogen(IV) oxide will support the combustion of burning substances when heated enough to release its oxygen for combustion.

2NO₂(g) → N₂(g) + 2O₂(g)

It allows burning carbon to continue burning in it forming carbon(IV) oxide and nitrogen.

2C(s) + 2NO₂ → 2CO₂(g) + N₂(g)

It also allow burning magnesium to continue burning forming magnesium oxide.

4Mg(s) + 2NO₂(g) → 4MgO(s) + N₂(g)

Finally, nitrogen(IV) oxide will also allow vigorously burning sulphur and phosphorus to continue burning, but will not rekindle a glowing splint.