`=>` Dinitrogen is produced commercially by the liquefaction and fractional distillation of air.

`=>` Liquid dinitrogen (b.p. `77.2 K`) distils out first leaving behind liquid oxygen (b.p. `90 K`).

`=>` In the laboratory, dinitrogen is prepared by treating an aqueous solution of ammonium chloride with sodium nitrite.

`color{red}(NH_4Cl(aq) +NaNO_2 (aq) → N_2 (g) +2H_2O (l) + NaCl (aq))`

● Small amounts of `color{red}(NO)` and `color{red}(HNO_3)` are also formed in this reaction; these impurities can be removed by passing the gas through aqueous sulphuric acid containing potassium dichromate. It can also be obtained by the thermal decomposition of ammonium dichromate.

`color{red}((NH_4)_2 Cr_2O_7 oversettext(Heat)→ N_2+4H_2O +Cr_2O_3)`

`=>` Very pure nitrogen can be obtained by the thermal decomposition of sodium or barium azide.

`color{red}(Ba(N_3)_2 → Ba +3N_2)`

`=>` Dinitrogen is a colourless, odourless, tasteless and non-toxic gas.

`=>` It has two stable isotopes : `color{red}(text()^14N)` and `color{red}(text()^15N)`.

`=>` It has a very low solubility in water (`23.2 cm^3` per litre of water at `273 K` and `1` bar pressure) and low freezing and boiling points (Table 7.1).

`=>` Dinitrogen is rather inert at room temperature because of the high bond enthalpy of `color{red}(N ≡ N)` bond. Reactivity, however, increases rapidly with rise in temperature.

`=>` At higher temperatures, it directly combines with some metals to form predominantly ionic nitrides and with non-metals, covalent nitrides. A few typical reactions are :

`color{red}(6Li +N_2 oversettext(Heat)→ 2Li_3N)`

`color{red}(3Mg +N_2 oversettext(Heat)→ Mg_3N_2)`

`=>` It combines with hydrogen at about `773 K` in the presence of a catalyst (Haber’s Process) to form ammonia.

`color{red}(N_2 (g) +3H_2 (g) overset(773 K) ⇌ 2NH_3 (g) ; \ \ \ \ \ \ Delta_fH^(⊖) = -46.1 k J mol^(-1))`

`=>` Dinitrogen combines with dioxygen only at very high temperature (at about 2000 K) to form nitric oxide, `NO.`

`color{red}(N_2+O_2 (g) oversettext(Heat) ⇌ 2NO (g))`

`=>` The main use of dinitrogen is in the manufacture of ammonia and other industrial chemicals containing nitrogen (e.g., calcium cyanamide).

`=>` It also finds use where an inert atmosphere is required (e.g., in iron and steel industry, inert diluent for reactive chemicals).

`=>` Liquid dinitrogen is used as a refrigerant to preserve biological materials, food items and in cryosurgery.

`=>` Ammonia is present in small quantities in air and soil where it is formed by the decay of nitrogenous organic matter e.g., urea.

`color{red}(NH_2CONH_2 +2H_2O →(NH_4)_2 CO_3 ⇌ 2NH_3 + H_2O +CO_2)`

`=>` On a small scale ammonia is obtained from ammonium salts which decompose when treated with caustic soda or lime.

`color{red}(2NH_4Cl+Ca(OH)_2 → 2NH_3+2H_2O +CaCl_2)`

`color{red}((NH_4)_2 SO_4+2NaOH → 2NH_3+2H_2O +Na_2SO_4)`

`=>` On a large scale, ammonia is manufactured by Haber’s process.

`color{red}(N_2(g) +3H_2(g) ⇌ 2NH_3 (g) ; \ \ \ \ \ \ \ \ \ Delta_fH^(⊖) = -46.1 kJ mol^(-1))`

● In accordance with Le Chatelier’s principle, high pressure would favour the formation of ammonia.

● The optimum conditions for the production of ammonia are a pressure of `color{red}(200 × 10^5)` Pa (about 200 atm), a temperature of `color{red}(~ 700 K)` and the use of a catalyst such as iron oxide with small amounts of `color{red}(K_2O)` and `color{red}(Al_2O_3)` to increase the rate of attainment of equilibrium. The flow chart for the production of ammonia is shown in Fig. 7.1.

`=>` Ammonia is a colourless gas with a pungent odour.

`=>` Its freezing and boiling points are `198.4` and `239.7 K` respectively.

`=>` In the solid and liquid states, it is associated through hydrogen bonds as in the case of water and that accounts for its higher melting and boiling points than expected on the basis of its molecular mass.

`=>` The ammonia molecule is trigonal pyramidal with the nitrogen atom at the apex. It has three bond pairs and one lone pair of electrons as shown in the structure.

`=>` Ammonia gas is highly soluble in water. Its aqueous solution is weakly basic due to the formation of `color{red}(OH^–)` ions.

`color{red}(NH_3 (g) +H_2O (l) ⇌ NH_4^(+) (aq) + OH^(-) (aq))`

`=>` It forms ammonium salts with acids, e.g., `color{red}(NH_4Cl, (NH_4)_2 SO_4)`, etc. As a weak base, it precipitates the hydroxides of many metals from their salt solutions. For example,

`color{red}(2FeCl_3(aq) +3NH_4OH (aq) → undersettext{(brown ppt)}(Fe_2O_3 . x H_2O (s)) +3NH_4Cl (aq))`

`color{red}(ZnSO_4 (aq) +2NH_4OH (aq) → undersettext{(white ppt)} (Zn (OH)_2 (s)) +(NH_4)_2SO_4 (aq))`

`=>` The presence of a lone pair of electrons on the nitrogen atom of the ammonia molecule makes it a Lewis base. It donates the electron pair and forms linkage with metal ions and the formation of such complex compounds finds applications in detection of metal ions such as `color{red}(Cu^(2+), Ag^(+))`.

`color{red}(undersettext{(blue)} (Cu^(2+)(aq))+4NH_3 (aq) ⇌ undersettext{(deep blue)} ([Cu(NH_3)_4]^(2+) (aq)))`

`color{red}(undersettext{(Colourless)} (Ag^(+) (aq)) +Cl^(-) (aq) →undersettext{(white ppt)}(AgCl (s)))`

`color{red}(undersettext{(white ppt)} (AgCl(s)) +2NH_3 (aq) → undersettext{(colourless)}([Ag(NH_3)_2])Cl (aq))`

`=>` Ammonia is used to produce various nitrogenous fertilisers (ammonium nitrate, urea, ammonium phosphate and ammonium sulphate).

`=>` In the manufacture of some inorganic nitrogen compounds, the most important one being nitric acid.

`=>` Liquid ammonia is also used as a refrigerant