Industrial Information

For industrial use Nitrogen is obtained by fractional distillation of the air. It is used for the manufacture of ammonia by the Haber-Bosch Process:

N₂(g) + 3H₂(g) ↔ 2NH₃(g)

A catalyst of finely-divided Iron is required for this equilibrium reaction. The mixture is heated to 450°C at 250 atm of pressure. Working at high pressure increases the yield of ammonia, but this adds to the cost of the plant, and a compromise between cost and yield is needed. The reaction is exothermic so low temperatures increase the yield of ammonia, but this slows up the reaction so again a compromise is needed. At the temperatures and pressures used in practice about 15% conversion is attained. The ammonia is condensed and removed from the plant and the unreacted gases are recycled.

Ammonia has numerous uses - approximately 100 megatonnes are produced worldwide each year. It is used as a fertilizer both directly and after conversion to other fertilizers such as Ammonium Nitrate. It is also a raw material for nitric acid manufacture and for the production of nylon. Gaseous Nitrogen is used to provide an inert atmosphere for reactions, which cannot be carried out in Oxygen. It is also used as a carrier gas in liquid-gas chromatography. Phosphorus is used in match-heads and on safety.

2. Nitrogen

Its electronic formula is 1s²2s²2p³. On the basis of graphic drawing of the external energy level (2s²2p³) the following valency state and oxidation numbers are possible for Nitrogen:

				2p
		2s		↑	↑			↑
		↑↓
1s
↑↓

Valency is III (owing to three unpaired electrons of 2p-sublevel). Valency is IV (three - owing to unpaired 2p-electrons, the forth bond - after donor-acceptor mechanism using 2s-electrons). Oxidation numbers are 3-, 2-, 1-, 0, 1+, 2+, 3+, 4+, 5+.

Nitrogen is the typical non-metal, only Fluorine and Oxygen have stronger electronegativity (3,0). The most widespread Nitrogen compounds of different oxidation levels (from 3- to 5+) are shown in Table 17.

Ammonia (NH₃) is colorless gas with purgen odor. In the laboratory heating of NH₄Cl with Calcium hydroxide or any strong alkalis can produce it:

2 NH₄Cl + Ca(OH)₂ 2 NH₃↑+ CaCl₂ + 2 H₂O.

Ammonia can react with water and acids, showing the main properties in the reactions. Practically it does not make NH₄OH hydroxide well dissolved in water (31% at 20⁰C), but reacts with H₂O by means of Hydrogen bond:

Table 17. General characteristic of Nitrogen compounds

Oxidation numbers	3-	2-	1-	0	1+	2+	3+	4+	5+
Nitrogen compounds	NH3 (NH4+)	N2H4 - hydrazine	NH2OH - hydroxyl-amine	N2	N2O -laughing gas	NO (N2O2)	N2O3	NO2 (N2O4)	N2O5
Properties in reactions without oxidation number	Weak basic properties; Proton acceptor	Basic properties; Proton acceptor	Weak base	Low chemical activity	Indifferent oxides		Acid oxides
							HNO2	HNO2 + HNO3	HNO3
Properties in reactions with oxidation number change	Reducing agent only	Reduction properties prevail over oxidation ones		Oxidizing agent; reducing agent relative to F2 and O2	Oxidizing and reducing agents (depending on reaction conditions)				Oxidizing agent only

NH₃ + H₂O ↔ NH₃·H₂O;

NH₃·H₂O ↔ NH₄OH.

The following important fertilizers are produced by ammonia interaction with acids:

NH₃ + HNO₃ = NH₄NO₃ - Ammonia Nitrate;

2NH₃ + H₂SO₄ = (NH₄)₂SO₄ - Ammonia Sulfate;

NH₃ + H₃PO₄ = NH₄H₂PO₄ - Ammonia dihydrophosphate (ammophos);

2NH₃ + H₃PO₄ = (NH₄)₂HPO₄ - Ammonia hydrophosphate (diammophos).

Nitric oxide (NO) is get by oxidation of Ammonia and it is used as the intermediate product in HNO₃ production:

4 NH₃ + 5 O₂ 4 NO + 6 H₂O

O20 + 4  2 O2-	5 (reduction);
N3- - 5  N2+	4 (oxidation).

Nitrogen oxide (III) and nitrous acid

Nitrogen oxide (III) is nitrous acid anhydride:

N₂O₃ + H₂O ↔ 2HNO₂.

As acid oxide it reacts with alkalis:

N₂O₃ + 2NaOH = 2NaNO₂ + H₂O.

Nitrous acid (HNO₂) dissociates with Hydrogen ions formation:

HNO₂ ↔ H⁺ + NO₂^- .

At it’s heating and under the influence of strong acids the process of disproportion takes place:

3HNO₂ ↔ HNO₃ + 2NO + H₂O

N3+ +  N2+	2 (reduction);
N3+ - 2  N5+	1 (oxidation).

Nitrous acid and its salts can be both reducing and oxidizing agents:

2NaNO₂ + 2KI + 2H₂SO₄ = I₂ + 2NO + K₂SO₄ + Na₂SO₄ + 2H₂O

N3+ + 1  N2+	2 (reduction);
2I- - 2  I20	1 (oxidation).

5KNO₂ + 2KMnO₄ + 3H₂SO₄ = 5KNO₃ + 2MnSO₄+ K₂SO₄ + 3H₂O

Mn7+ + 5  Mn2+	2 (reduction);
N3+- - 2  N5+	5 (oxidation).

Nitric oxide (V) and nitric acid

Nitric acid and its salts (Nitrates) have the oxidizing properties:

_{3KNO3 + 8Al + 5KOH + 2H2O = 3NH3 + 8KAlO}2

N5+ + 8  N3-	3 (reduction);
Al0 - 3  Al3+	8 (oxidation).

Particular feature of nitric acid is in interaction almost with all metals and non-metals, at the same time it oxidizes them. The reduction of N(5+) but not H(1+) as with acids-oxidizing agents always takes place. Generally NO and NO₂ prevail among reduction products. Active metals (Mg, Zn, Ca and others) reduce diluted HNO₃ up to N₂ and NH₄NO₃.

At the same time some N(5+) reduction products can be isolated. However, the equations of such reactions are relative and only one compound (NO₂, NO or N₂, NH₃), formed in quantity, is indicated in the products:

4 Mg + 10HNO₃ (diluted) = 4Mg(NO₃)₂ + NH₄NO₃ + 3H₂O;

Cu + 4HNO₃ (conc) = Cu(NO₃)₂ + 2NO₂ + 2H₂O;

S + 6HNO₃ (conc) = H₂SO₄ + 6NO₂ + 2H₂O;

3P + 5HNO₃ (diluted) + 2H₂O = 3H₃PO₄ + 5NO.

Nitric acid salts - Sodium, Potassium, ammonia, and calcium Nitrates are of great practical value. They are called saltpeters and are used in great quantity as fertililizers.