Nitrogen

Nitrogen

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.

Ammonia

Preparation

Industrial

The Haber process is used to manufacture ammonia in industry.

Catalyst: Iron

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

Pressure: 200 atm

Laboratory

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

NH₄Cl + NaOH –> NaCl + H₂O + NH₃

Reaction

Basic

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.

Solubility

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

Cl₂  –> 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

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 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 NO_2. It also dimerises to pair up the unpaired electron.

Nitrogen dioxide NO₂

Just like NO, NO₂ has an unpaired electron and therefore dimerises to produce N₂O_4.

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:       

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

HNO₃ + P₂O₅ –> HPO₃ + N₂O₅

Nitric (III) acid HNO₂

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 NO₂^– can be both  a reducing agent nad an oxidising agent:

Reducing

NO₂^– –> NO₃^–

Cr₂O₇^2- –> 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

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 800^oC)

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 –> PO₄^3-

S –> SO₄^2-

 

Nitrates

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⁺