Nitrogen makes up 78% of the air and it is a highly unreactive gas.  This is due to the fact that nitrogen has a triple bond, and the activation energy to break this bond is very high.

Nitrogen reacts with oxygen under extreme condition (thunderstorms, car engines) to form Nitrogen monoxide, which is then oxidised to nitrogen dioxide in the air.




The Haber process is used to manufacture ammonia in industry.

Nitrogen 1

Catalyst: Iron

Temperature: 400oC (Even though the reaction is exothermic, a high temperature is used in order to increase the rate of reaction.

Pressure: 200 atm


Reacting Ammonium salts with alkali. This gives off the salt, water and ammonia.

NH4Cl + NaOH –> NaCl + H2O + NH3



Ammonia has a lone pair on the nitrogen and therefore this can donate a pair of electrons making it a Lewis base. Ammonia can react with acids to produce ammonium salts while it can form bonds with B and Al since this both are electron deficient.

Nitrogen 2


Ammonia is very soluble in water to give a basic solution.

Complex formation

Concentrated ammonia can be used for identification of cations since this can react with a number of metals to produce complexes. One such complex is with Copper, where it can produce a deep blue solution.

Cu(OH)2 + xsNH3 –> [Cu(NH3)4]2+(aq)

Reducing properties

Nitrogen has a number of different oxidation states, and ammonia is the lowest of these, meaning that it can be oxidised to higher oxidations states.

NH3 –> N2

Cl–> Cl

CuO –> Cu

O2 –> H2O

Ammonium Salts

Most ammonium salts decompose to give ammonia and an acid, such as:

NH4Cl –> NH3 + HCl

(NH4)2CO3 –> NH3 + H2O + CO2

Nitrates, nitrites and chromates decompose differently, and these give a different form of nitrogen.

NH4NO3 –> N2O + H2O

NH4NO2 –> N2 + H2O

(NH4)2Cr2O7 –> N2 + H2O + CrO3

Oxides of Nitrogen

Dinitrogen oxide N2O

Nitrogen 3

N2O has a sweet smell, decomposes into elements on reacting with Oxygen and has anaesthetic properties.

This is prepared form the decomposition of ammonium nitrate.

NH4NO3 –> N2O + H2O

Nitrogen monoxide NO

Nitrogen 4

Nitrogen monoxide can be prepared by the oxidation of copper using cold nitric acid

Cu –> Cu2+

NO3 –> NO


NO is quite unstable due to the fact that it has an odd number of electron and therefore it readily oxidises in air to form NO2. It also dimerises to pair up the unpaired electron.

Nitrogen dioxide NO2

Nitrogen 5

Just like NO, NO2 has an unpaired electron and therefore dimerises to produce N2O4.

It is a brown gas and can be prepared by the oxidation of copper with hot nitric acid.

Cu  –> Cu2+

NO3–  –>NO2

It can also be prepared by the decomposition of nitrates bar group I nitrates.

2Pb(NO3)2 –> 2PbO + 2NO2 + O2

This can be collecting and purified by passing from a u-shaped tube over ice cold water. The NO2 will dimerise to produce a liquid while the O2 will remain gaseous.

Dissolving NO2 in water disproportionated the Nitrogen to give two acids:

2NO2 + H2O –> HNO3 + HNO2

Dinitrogen pentoxide N2O5


N2O5 can be either a gas or a solid, with both having different molecular structures.

Solid: NO2 NO3+                                                                        Gas:       Nitrogen 6

It can be prepared by the dehydration of nitric acid using P2O5:

HNO3 + P2O5 –> HPO3 + N2O5

Nitric (III) acid HNO2

Nitrogen 7

HNO2 is prepared at cold temperature via the reaction of a nitrite salt and acid. The cold temperature is important because it would otherwise disproportionate the products.

HCl + KNO2 –> HNO2 + KCl

On heating HNO2 disproportionates to give:

2HNO2 –> HNO3à NO

Only the alkali nitrate (III) salts are stable to heat, all the others would decompose in heating.

The NO2can be both  a reducing agent nad an oxidising agent:


NO2 –> NO3

Cr2O72- –> Cr3+


Oxidisng agent

NO2 –> NO

I –> I2


A test to distinguish between Nitrate (III) and Nitrate (V) is to add an acid. Nitrate (III)  will give a brown gas, NO2, while Nitrate (V) will not.

Nitric (IV) acid HNO3

Nitrogen 8

Nitriv (V) can be prepared from any nitrate salt reacting with an acid.

NaNO3+  HCl –> NaCl + HNO3

In industry, ammonia is used, first oxidising it to NO, then oxidation to NO which is, in turn, turned into NO3.

4NH3 + 5O2 –> 4NO + 6H2O (Pt catalyst at 800oC)

2NO + O2 –> 2NO2

NO2 + H2O + O2 –> HNO3

Nitric acid is a very good oxidising agent, and it has a number of reaction that it can undergo:

Cold copper

Cu –> Cu2+

NO3–> NO

Hot copper

 Cu –> Cu2+

NO3–> NO2

With zinc the product of the reaction is N2O

Zn –> Zn2+

NO3 –> N2O

Dilute HNO3 also oxidised non-metals to their highest oxidation state.

NO3 –> NO2

P –> PO43-

S –> SO42-



Nitrates are prepared with the reaction of nitric acid with a base.

Alkali nitrates break down to give nitrites while all other nitrates give nitrogen dioxide.

Test for the NO3

Brown ring test as describes in iron

Devarda’s alloy

The devarda’s alloy is a reaction with aluminium under basic conditions:

NO3–> NH3

Al –> Al3+