Oxides of Nitrogen

Introduction to Oxides of Nitrogen

Nitrogen has a position in second period of group V in the modern periodic table. It has molecular formula N₂. It has atomic number 7 and atomic weight 14.08 and its electronic configuration of 2,5.

Besides combining with hydrogen and forming NH₃ , nitrogen combines with oxygen in different ratios and forms five different oxides.
The oxides of nitrogen and the details of the oxygen states of nitrogen and the N:O ratio can be presented in a tabular form as:

– All of these oxides of nitrogen are gases, excepting NO and N₂O the other oxides are brownish gases.
– Except the oxides of NO and N2O all the other oxides are acidic. NO and N₂O are neutral.
– The other oxides are prepared in the laboratory using different methods characteristic to each oxide.

Nitric Oxide – NO

Structure

Laboratory Preparation

The oxide is prepared in the laboratory by treating the metallic copper with a moderately concentrated nitric acid (1:1) at room temperature.
The reaction is given as :

3Cu + 8HNO₃ → 3Cu(NO₃)₂ + 2NO + 4H₂O

– The gas is collected by downward displacement of water.
– The apparatus used is Wolfe’s apparatus.
– The purification is done by absorbing the NO gas in freshly prepared ferrous sulphate solution.
– Ferrous sulphate absorbs all the NO gas and forms Fe(H₂O)₅NO and the solution becomes brown.
– On heating this solution pure Nitric Oxide is obtained.

Physical Properties

NO is not a combustible gas. At high temperature around 1000°C it decomposes into N₂ and O₂.

2NO → N₂ + O₂ (at high temperature)

From the equation above we can see that once the decomposition starts 50% O₂ gets evolved and this O₂ supports combustion thus making the reaction more violent.

Chemical properties

1) NO acts as an oxidising agent, oxidising SO₂ in presence of water to give H₂SO₄.

SO₂ + 2NO + H₂O → H₂SO₄ + N₂O

2) NO acts as a reducing agent

i) It can also reduce aqueous solution of I₂ to HI

3I₂ + 2NO + 4H₂O → 2HNO₃ + 6HI

3) With halogens NO can form addition compounds as:

2NO + Cl₂ → 2NOCl    (NOCl is nitrosyl chloride)

It reacts in the same way with fluorine and bromine.
4) With ferrous sulphate NO forms an addition compound as

FeSO₄ + 5H₂O + NO → [Fe(H₂O)₅NO]SO₄
penta aqua nitrosyl iron (II) sulphate

This is the famous brown ring test used to identify the nitrate radical or the NO radical.

Uses

Nitric Oxide is used to prepare nitric acid.

Nitrogen Dioxide – NO₂

Structure

Laboratory Preparation

In the laboratory NO₂ is prepared by thermal decomposition of Pb(NO₃)₂. Thus:

2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂

Please Note: Hydrated nitrate salts on heating react violently and explode.

Physical Properties

NO2 is a poisonous gas, main source being the exhaust of auto mobiles.

1. At room temperature it is a deep brown gas.
2. It does not support combustion.
3. It is not combustible

Chemical properties

1) With cold water NO₂ reacts to give a mixture of HNO₂ and HNO₃ acid.

2NO₂ + H₂O → HNO₂ + HNO₃

2) With hot water the reaction is

3NO₂ + H₂O → 2HNO₃ + NO

3) Being acidic it reacts with bases as

2NO₂ + 2KOH → KNO₃ + H₂O + KNO₂

4) It is also a strong oxidising agent.

H₂S + 3NO₂ → 3NO + H₂0 + S0₂

5) With excess oxygen and water NO₂ gives HNO₃.

4NO₂ + O₂ + 2H₂O → 4HNO₃

6) It reacts with concentrated H₂SO₄ to give nitrosyl hydrogen sulphate

2NO₂ + H₂SO₄ → SO₂(OH)ONO + HNO₃

Uses

NO₂ is used as a fuel in rockets besides being used to prepare HNO₃

Nitrous Oxide – N₂O (laughing Gas)

Structure

Laboratory Preparation

– N₂O can be prepared in the laboratory by heating NH₄NO₃ below 200°C to avoid explosion.
– Sometimes as a safety measure instead of directly using NH₄NO₃, a mixture of (NH₄)₂SO₄ and NaNO₃ are heated to give NH₄NO₃ which decomposes further to give N₂O.

NH₄NO₃ → N₂O + 2H₂O
(endothermic reaction)

Physical Properties

1) N₂O has a faint sweet smell and produces a tickling sensation on the neck when inhaled and makes people laugh hysterically. Excess of inhalation leads to unconsciousness.
2) Unlike other oxides of nitrogen, N₂O supports combustion though it does not burn itself.

Chemical properties

1) At very high temperature N2O decomposes to N₂ and O₂

2N₂O → 2N₂ + O₂

If a glowing piece of Mg, Cu, or P is introduced in such an environment, these pieces burn brightly due to the O₂ produced from decomposition of N₂O.
2) With Sodium and potassium N₂O reacts to give the corresponding peroxides liberating N₂ in the process.

2N₂O + 2Na → Na₂O + 2N₂
Na₂O is sodium peroxide

Uses

1) It is used as propellent gas.
2) Used in combination with oxygen in the ratio N₂O : O₂ = 1:10 as a mild anaesthetic