Nitrogen and Its Inorganic Compounds

Lessons

Senior 3 · Chemistry

Nitrogen and Its Inorganic Compounds

4 lessons

Lesson 1 of 4

Nitrogen: Preparation, Properties and Uses

25 min

Colder than anything in the kitchen

When a dairy farmer asks for artificial insemination, the technician arrives with semen straws that have been stored at −196 °C, the boiling point of liquid nitrogen. Producing liquid nitrogen and supplying it for artificial insemination is one of the jobs of the Rwanda Agriculture and Animal Resources Development Board (RAB). Nitrogen makes up 78% of the air you breathe, yet it takes part in almost none of the reactions around you. Why is nitrogen so unreactive, how is it obtained, and how do chemists persuade it to react at all?

Nitrogen from the air

In S2 you saw how oxygen is obtained from liquid air. The same process, fractional distillation of liquid air, is the industrial source of nitrogen. Air is cleaned, dried, cooled until it liquefies, and then allowed to warm up slowly in a tall column. Nitrogen boils at −196 °C, before argon (−186 °C) and oxygen (−183 °C), so it is collected as a gas from the top of the column.

In the laboratory, nitrogen can be obtained from air by removing the other gases chemically (Figure 1.2). The gas that is left is about 99% nitrogen; the rest is argon and other noble gases, which do not react with any of the reagents.

Figure 1.2: Obtaining nitrogen from air by removing the other gases.

Preparing nitrogen in the laboratory

Pure nitrogen is made by warming a solution containing ammonium chloride and sodium nitrite. Together they behave as ammonium nitrite, which decomposes to nitrogen and water:

$\mathrm{NH_4Cl(aq) + NaNO_2(aq) \rightarrow NaCl(aq) + N_2(g) + 2H_2O(l)}$

Nitrogen is collected over water (Figure 1.1): it is almost insoluble in water, and its density is so close to that of air ( $M_{r} = 28$ , against 29) that it cannot be collected by displacing air. The first jars are thrown away, because they contain the air that was in the flask.

Figure 1.1: Laboratory preparation of nitrogen.

Warning

Safety. Heat the mixture only gently. Once the reaction starts it gives out heat and can become violent, so remove the flame and heat again only if the bubbling slows. Sodium nitrite is toxic if swallowed: wear gloves and eye protection and wash your hands afterwards. Never heat solid ammonium nitrite or ammonium nitrate; they can decompose explosively.

Properties of nitrogen

Physical properties. Nitrogen is a colourless, odourless, tasteless gas. It is slightly less dense than air and almost insoluble in water. It is neutral to litmus. It liquefies at −196 °C.

Why it is unreactive. The two atoms of a nitrogen molecule share three pairs of electrons: a triple bond, $\mathrm{N{\equiv}N}$ . This is one of the strongest bonds known, and a great deal of energy is needed to break it before nitrogen can react. So nitrogen does not burn and does not support burning: a lighted splint goes out in it.

Reactions it does undergo, all at high temperature:

Burning magnesium continues to burn in nitrogen, forming magnesium nitride. Water decomposes the nitride and releases ammonia, which turns damp red litmus blue:
$\mathrm{3Mg(s) + N_2(g) \rightarrow Mg_3N_2(s)}$
$\mathrm{Mg_3N_2(s) + 6H_2O(l) \rightarrow 3Mg(OH)_2(s) + 2NH_3(g)}$
Oxygen combines with nitrogen only at very high temperatures, in lightning flashes and inside hot engines: $\mathrm{N_2(g) + O_2(g) \rightarrow 2NO(g)}$ .
Hydrogen combines with nitrogen to make ammonia, with a catalyst, high pressure and heat (the Haber process, Lesson 2).

Definition

Nitrogen fixation: any process that turns the unreactive nitrogen gas of the air into nitrogen compounds that plants can use, such as ammonium ions and nitrates. Example: bacteria in the root nodules of beans and other legumes fix nitrogen; lightning and the Haber process fix it too.

Table 1.1: Nitrogen and oxygen compared

Property	Nitrogen, $\mathrm{N_2}$	Oxygen, $\mathrm{O_2}$
Percentage of dry air (by volume)	78%	21%
Boiling point	−196 °C	−183 °C
Bond between the atoms	triple bond	double bond
Effect on a glowing splint	goes out	relights
Reactivity	very low; reacts only at high temperature	high; supports burning and rusting
Collected in the laboratory	over water	over water

Uses of nitrogen

Making ammonia (Lesson 2), and from it fertilisers, nitric acid and explosives: by far the biggest use.
Liquid nitrogen as a coolant: storing semen and embryos for animal breeding, quick-freezing food, preserving blood and tissue samples in hospitals.
An inert atmosphere: packets of crisps and some foods are filled with nitrogen so that the food does not react with oxygen and go stale; oil tanks and chemical reactors are flushed with nitrogen to prevent fires.

Worked example

Problem. What volume of nitrogen, at room temperature and pressure, is produced when 5.35 g of ammonium chloride reacts completely with excess sodium nitrite? (N = 14, H = 1, Cl = 35.5; molar volume $24\,\mathrm{dm^3\,mol^{-1}}$ )

Solution. $M(\mathrm{NH_4Cl}) = 14 + 4 + 35.5 = 53.5\,\mathrm{g\,mol^{-1}}$ .

$n(\mathrm{NH_4Cl}) = \frac{5.35}{53.5} = 0.100\,\mathrm{mol}$

From the equation, 1 mol $\mathrm{NH_4Cl}$ gives 1 mol $\mathrm{N_2}$ , so 0.100 mol of nitrogen forms. Volume $= 0.100 \times 24 = 2.40\,\mathrm{dm^3}$ .

Worked example

Problem. In 1892 Lord Rayleigh found that a litre of nitrogen obtained from air, as in Figure 1.2, was always about 0.5% heavier than a litre of nitrogen made from ammonium compounds. Use the fact that air contains 78.1% nitrogen and 0.93% argon to explain the difference. ( $M_{r}$ : $\mathrm{N_2} = 28$ , Ar = 40)

Solution. The "nitrogen from air" still contains the argon. Its fraction in that gas is $\frac{0.93}{78.1 + 0.93} = 0.0118$ , about 1.2%. The average $M_{r}$ of the mixture is

$0.988 \times 28 + 0.0118 \times 40 = 27.66 + 0.47 = 28.14$

Equal volumes of gases contain equal numbers of molecules, so density is proportional to $M_{r}$ : $\frac{28.14}{28} = 1.005$ , which is 0.5% heavier, just as Rayleigh found. Working with William Ramsay, he went on to discover argon in 1894.

Tip

Three bonds, three reasons. Nitrogen's triple bond explains why it is unreactive, why it is used as an inert atmosphere, and why the Haber process needs a catalyst, high pressure and heat.

Periodic Videos: Nitrogen (13 min), from the University of Nottingham. Watch the demonstrations with liquid nitrogen and list each property of nitrogen that they show.

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Practice

Name the two substances heated together to prepare nitrogen in the laboratory, and write the equation.
Give two reasons why nitrogen is collected over water.
Explain, in terms of bonding, why nitrogen is unreactive.
In Figure 1.2, what is removed by (a) sodium hydroxide solution, (b) heated copper? Why is the nitrogen obtained not pure?
Burning magnesium ribbon is lowered into a jar of nitrogen. It keeps burning and forms a yellowish powder. When water is added to the powder, a gas turns damp red litmus blue. Explain with equations.
Give three uses of nitrogen and the property each depends on.
What mass of ammonium chloride is needed to make $600\,\mathrm{cm^3}$ of nitrogen at room temperature and pressure? ( $\mathrm{NH_4Cl} = 53.5$ ; molar volume $24\,\mathrm{dm^3\,mol^{-1}}$ )

Show answersHide answers

Ammonium chloride and sodium nitrite (solution): $\mathrm{NH_4Cl(aq) + NaNO_2(aq) \rightarrow NaCl(aq) + N_2(g) + 2H_2O(l)}$ .
It is almost insoluble in water; its density is almost the same as that of air, so displacing air would not work.
The two nitrogen atoms are joined by a triple bond, which is very strong; a lot of energy is needed to break it before nitrogen can react.
(a) Carbon dioxide; (b) oxygen. It still contains about 1% of noble gases, mainly argon, which do not react with any of the reagents.
Magnesium burns in nitrogen to form magnesium nitride: $\mathrm{3Mg(s) + N_2(g) \rightarrow Mg_3N_2(s)}$ . Water decomposes the nitride to ammonia, an alkaline gas: $\mathrm{Mg_3N_2(s) + 6H_2O(l) \rightarrow 3Mg(OH)_2(s) + 2NH_3(g)}$ .
Making ammonia and fertilisers (it combines with hydrogen under the right conditions); liquid nitrogen for freezing (very low boiling point); inert atmosphere in food packets (unreactive).
$n(\mathrm{N_2}) = \frac{0.600}{24} = 0.0250\,\mathrm{mol}$ , so 0.0250 mol $\mathrm{NH_4Cl}$ ; mass $= 0.0250 \times 53.5 = 1.34\,\mathrm{g}$ .

Sources and credits

Chemistry 2e, 18.7: Occurrence, Preparation, and Properties of Nitrogen (obtaining nitrogen from liquid air, its triple bond and low reactivity, its reaction with magnesium, and its uses), Paul Flowers, Klaus Theopold, Richard Langley and William R. Robinson, OpenStax, via LibreTexts, licensed under CC BY 4.0. Adapted by Ganzaa.
RAB's work with liquid nitrogen: National Animal Genetic Improvement, Rwanda Agriculture and Animal Resources Development Board (read 30 September 2026).
Photo: Liquid nitrogen ŚFN 2025 (2).jpg, Adrian Tync, licensed under CC BY-SA 4.0, via Wikimedia Commons.
Video: Nitrogen – Periodic Table of Videos, Periodic Videos (University of Nottingham), embedded from YouTube.
The diagrams, table, worked examples and Practice items were written and drawn by Ganzaa.

1Colder than anything in the kitchen

2Nitrogen from the air

3Preparing nitrogen in the laboratory

4Properties of nitrogen

5Uses of nitrogen

Colder than anything in the kitchen

Nitrogen from the air

Preparing nitrogen in the laboratory

Properties of nitrogen

Uses of nitrogen