Nitrogen and its compounds - Dr D. Mizzi Page 1 of 15

Watch this video about the fractional distillation of liquid air.

10.1 GROUP 5 – Nitrogen and its Compounds

Nitrogen

I can describe the properties, test and uses of nitrogen.

Nitrogen is an essential element necessary for the wellbeing of animals and plants. It is present in proteins found in all living things. It is essential that nitrates are present in the soil for it to remain fertile. Farmers add either farmyard manure or artificial fertilisers e.g. ammonium nitrate to replace the lack of nitrates in the soil.

It is a colourless, odourless gas, which makes up 78% of the Earth’s atmosphere.

It is only slightly soluble in water.

It is a very unreactive gas.

It is diatomic. N2 contains a triple covalent bond N N.

Nitrogen is neutral gas and has no effect on litmus.

Nitrogen gas is obtained industrially by fractional distillation of liquid air. This is carried out as follows:

First air is filtered to remove dust. Then it is cooled in stages. As air is cooled down water vapour

condenses and is removed followed by carbon dioxide, which solidifies at -78°C.

Air is cooled further by compressing it and then allowed to expand suddenly. After many cycles some air liquefies at - 200°C.

Liquid air is pumped into the fractionating column and it is slowly warmed up. Since the gases boil off at different temperatures (nitrogen -196◦C, argon -186◦C, oxygen -183◦C) they will condense in the different trays in the fractionating column. Oxygen, nitrogen and argon are redistilled. They are stored under pressure in strong steel cylinders.

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Nitrogen will combine directly with metals to form nitrides. 3Mg (s) + N2 (g) Mg3N2 (s)

Nitrogen combines with oxygen at high temperatures to form first nitrogen monoxide (NO)

which rapidly combines with O2 to form nitrogen dioxide (NO2). N2 (g) + O2 (g) 2NO (g) 2 NO (g) + O2 (g) 2 NO2 (g)

This happens in car engines, where toxic oxides of nitrogen (NOx) are emitted as car exhaust fumes. Oxides of nitrogen will dissolve in water vapour to form acid rain.

4NO2(g) + O2(g) + 2H2O (g) → 4HNO3 (aq) Catalytic converters prevent emission of oxides of nitrogen by converting them back to N2 gas before they leave the exhaust pipe.

CO (g) + O2 (g) 2 CO2 (g)

2 NO (g) + 2 CO (g) N2 (g) + 2 CO2 (g)

Test for Nitrogen Since the gas is inert it has no positive tests and it is identified by elimination of other tests, i.e. it does not support combustion and it does turn lime water milky. Uses of Nitrogen

to make ammonia, nitric acid, & fertilisers; to fill crisps bags. Nitrogen will stop the oxidation of food in packed food bags; liquid nitrogen is used as a refrigerant for e.g. to preserve biological specimens.

Take this short quiz about nitrogen. Some questions have one correct response, whilst others have more than one.

Watch this video about the uses of nitrogen

Nitrogen and its compounds - Dr D. Mizzi Page 3 of 15

This experiment can be seen on the following video: https://www.youtube.com/watch?v=16gfUwFAjlU

Ammonia

I can describe the properties, test and the lab preparation of ammonia.

Ammonia is a colourless gas with a sharp or pungent smell.

It is a toxic gas which is less dense than air.

It is easily liquefied by cooling to –33C or by compressing it, so it is easy to transport in tanks or cylinders.

It is very soluble in water forming ammonia solution which is weakly alkaline (since few OH- are present as the reaction is reversible).

NH3(g) + H2O(l) ⇌ NH4+(aq) + OH-(aq)

Test for Ammonia Ammonia turns damp red litmus paper blue.

Ammonia produces white fumes with hydrogen chloride gas. NH3 (g) + HCl (g) → NH4Cl (s)

Chemical Properties of Ammonia 1. It reacts with hydrogen chloride gas to form white smoke of ammonium chloride (NH4Cl).

NH3 (g) + HCl (g) → NH4Cl (s)

cotton wool soaked in concentrated hydrochloric acid

cotton wool soaked in concentrated ammonia solution

white ring of ammonium chloride

This can be demonstrated by the fountain experiment as seen in this video:

Nitrogen and its compounds - Dr D. Mizzi Page 4 of 15

2. Since ammonia is alkaline, it reacts with acids to form salts. Examples: Ammonia reacts with nitric acid to form ammonium nitrate, which is an important

fertiliser. NH3 (g) + HNO3 (aq) → NH4NO3 (aq)

Ammonia reacts with sulfuric acid to form ammonium sulfate. 2NH3 (g) + H2SO4 (aq) → (NH4)2SO4 (aq)

3. Ammonia will burn in atmosphere of air slightly enriched in oxygen but not air alone.

In the presence of a platinum catalyst it produces nitrogen monoxide and steam. 4NH3 (g) + 5O2 (g) ⇌ 4NO (g) + 6H2O (g)

4. Ammonia acts as a reducing agent. It reacts with copper (II) oxide to reduce it to copper.

Ammonia is oxidised to nitrogen and water. 3CuO (s) + 2NH3 (g) → 3Cu (s) + 3H2O (l) + N2 (g)

Using oxidation numbers: CuO is reduced to Cu since the oxidation numbers decrease from +2 to 0 NH3 is oxidised to N2 since the oxidation numbers increase from -3 to 0

5. Ammonia solution precipitates insoluble metallic hydroxides,

e.g. NH4OH (aq) + FeSO4 (aq) Fe(OH)2 (s) + (NH4)2SO4 (aq) green gelatinous precipitate

Uses of ammonia

Most of the ammonia from the Haber process is used to make fertilisers, such as ammonium nitrate, NH4NO3 and ammonium sulfate, (NH4)2SO4.

It is used to make household cleaners, dyes, explosives, nylon and nitric acid.

ammonia

copper (II)

oxide

liquid cold water

heat

gas

water

Nitrogen and its compounds - Dr D. Mizzi Page 5 of 15

Preparation of Ammonia in the Laboratory Ammonia is produced by heating an ammonium salt with an alkali.

alkali + ammonium salt salt + water + ammonia Ca(OH)2 (s) + 2NH4Cl(s) CaCl2 (s) + 2H2O(g) + 2NH3 (g)

Grind the mixture and put it in a flask. The neck of the flask should slope downwards towards A so that water will not run back.

If allowed to run back into the hot flask, it might break. Heat the flask. Water vapour is removed from the ammonia gas by passing it through a drying tower

containing calcium oxide as the drying agent. (CaO does not react with ammonia). Since ammonia is less dense than air and very soluble in water, it is collected by upward

delivery.

Preparation of ammonia solution Heat gently the flask. The rim of the inverted funnel should just touch the surface of the water. This is a device to prevent water from ‘sucking back’ into the flask. After some time, the water will be found to have acquired the smell of ammonia gas, which has dissolved in it. The solution is called ammonia solution.

NH3 (g) + H2O (l) ⇌ NH4+ (aq) + OH- (aq)

calcium hydroxide &

ammonium chloride

calcium oxide

to dry

ammonia

NH3

card cover

heat

ammonium chloride & calcium hydroxide

heat

H2O

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Industrial Manufacture of Ammonia – The Haber Process

I can describe the steps in the industrial manufacture of ammonia.

Until the beginning of the 20th century deposits of nitrogen fertilisers such as sodium nitrate, bird

droppings and animal manure were sufficient to put the nitrogen back into the soil to maintain soil

fertility. With the demands for nitrates for explosives, a method of making ammonia synthetically

was necessary. In 1908 Fritz Haber developed a process to combine nitrogen and hydrogen to

produce ammonia. In 1913 the first ammonia factory was making 30 tonnes a day. The Haber

process is still used today and about 85% of ammonia goes to make fertilisers to help feed the world.

Raw materials Nitrogen obtained from fractional distillation of liquid air. Hydrogen obtained from the reaction between methane and steam.

CH4 (g) + 2H2O (g) ⇌ CO2 (g) + 4H2 (g) Procedure The reaction is reversible so the percentage of ammonia produced depends on the conditions.

N2 (g) + 3 H2 (g) ⇌ 2NH3 (g) H = - 92KJ/mol 1. The two gases are mixed. The mixture is cleaned or scrubbed to get rid of any

impurities. 2. Nitrogen and hydrogen in the correct proportions (1:3) are then pressurised to

approximately 200 atmospheres at the compressor. 3. The gases pass through the converter over a catalyst of freshly produced, finely divided

iron at a temperature of between 350C and 500C. 4. Under these conditions, the gas mixture leaving the reaction vessel contains about 15%

of ammonia gas which is removed by cooling and condensing it as a liquid. The unreacted N2 and H2 are recirculated into the reaction vessel to react together to produce further quantities of ammonia.

5. Ammonia is run into tanks and stored as a liquid, under pressure.

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Nitrogen

N2

Hydrogen

H2

Converter

450°C temperature

200 atm pressure

Fe catalyst

Ammonia

NH3

air

natural gas

15%

unreacted, recycled

Watch this video about the Haber process: (reference to equilibrium will be explained in Year 11)

Summary: The chemical industry is geared up to make new substances in a cost effective way. The conditions for any industrial chemical process are often chosen for financial rather than chemical reasons. The constant flow of nitrogen and hydrogen gases will displace the mixture before it has

the opportunity to establish equilibrium.

The highest yields of ammonia are obtained at high pressures and low temperatures.

High pressures are expensive as the apparatus has to have reinforced pipe work etc. to avoid explosions and the cost of doing this restricts the pressure used. 200 atmospheres are chosen to save money.

At low temperature the yield is high but the rate of the reaction is slower. As a

compromise the temperature is raised to 450C and the use of a catalyst of finely divided iron speeds up the reaction.

Nitrogen and its compounds - Dr D. Mizzi Page 8 of 15

solid

ammonium

chloride

white sublimate

of ammonium

chloride

heat

Ammonium salts 1. They are all soluble in water.

2. Ammonium salts give off NH3 when heated with an alkali, this property is used to test for the NH4

+ ion. NH4Cl (s) + NaOH (aq) NaCl (s) + NH3 (g) + H2O(l)

3. Some ammonium salts show thermal dissociation when heated.

These gases recombine on cooling. Thus solid sublimes. NH4Cl (s) ⇌ NH3 (g) + HCl (g)

Ammonium carbonate decomposes on heating but products do not recombine on cooling.

(NH4)2CO3 (s) 2NH3 (g) + CO2 (g) + H2O (l)

Take this short quiz about ammonia. Some questions have one correct response, whilst others have more than one.

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Watch this video which shows the acidic nature of nitrogen dioxide: https://www.youtube.com/watch?v=e6LF7m1zEUM

Nitrogen dioxide

I can describe the properties and the lab preparation of nitrogen dioxide.

It is a brown gas.

It is toxic and denser than air.

It is an acidic oxide. Nitrogen dioxide reacts with water to form a mixture of

nitrous (HNO2) and nitric acid (HNO3).

2NO2 (g) + H2O (l) HNO2 (aq) + HNO3 (aq)

Test for nitrogen dioxide

It is brown gas which turns damp blue litmus red (being an acidic gas).

Nitrogen dioxide is known to be an atmospheric pollutant, it contributes to acid rain.

Watch this video which shows how the emissions of NO2 have decreased during the Covid-19 pandemic over China.

Watch this video showing the test for nitrogen dioxide.

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Laboratory preparation of nitrogen dioxide It is produced by the action of heat on the fairly reactive metal nitrates. Lead(II) nitrate makes a crackling sound and melts on heating. It effervesces giving a brown gas (NO2) and oxygen. NO2 is liquefied and condenses as liquid dinitrogen tetroxide [N2O4] in the freezing mixture and collects in the boiling tube. It may appear greenish due to dissolved water if the lead(II) nitrate was not absolutely dry. O2 escapes as a gas. Lead(II) oxide remains in the test tube as a yellow solid fused into the glass.

2 Pb(NO3)2 (s) 2PbO (s) + 4NO2 (g) + O2 (g)

Thermal dissociation of dinitrogen tetraoxide, N2O4 Nitrogen dioxide, NO2 and dinitrogen tetraoxide, N2O4 occur together in equilibrium. The position of this equilibrium and colour of the system varies with temperature and pressure.

cool 2NO2 (g) ⇌ N2O4 (g) ΔH -ve brown gas heat colourless gas

The intensity of the brown colour of NO2 decreases as the temperature decreases since the equilibrium shifts to the N2O4 side upon a decrease in temperature.

On increasing pressure the equilibrium shifts to the N2O4 side that is to the side of a lower number of molecules. (You will study the topic of equilibria in Year 11).

Nitrogen dioxide dimerizes (i.e. 2 molecules of NO2 bond together) to form dinitrogen tetraoxide

Below

21.5°C

Above

21.5°C

heat

lead(II)nitrate

boiling tube

liquid N2O4

freezing mixture

(ice salt mixture)

glass tube

rubber tube

glass tube

oxygen escapes from here

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Nitrogen dioxide can also be produced by the reaction of copper with concentrated nitric acid. Cu (s) + 4HNO3 (aq) Cu(NO3)2 (aq) + 2H2O (l) + 2NO2 (g)

Nitrogen monoxide

I can describe the properties and test of nitrogen monoxide.

a colourless gas almost insoluble in water it is a toxic gas slightly denser than air neutral to litmus readily combines with oxygen to form NO2

2 NO (g) + O2 (g) 2 NO2 (g) Test for Nitrogen monoxide Expose nitrogen monoxide to air and brown fumes are formed at once by oxidation of the gas in air. Summary: The oxides of nitrogen NOx include nitrogen dioxide and nitrogen monoxide. The following table compare the properties of the two oxides: Nitrogen dioxide, NO2 Nitrogen monoxide, NO

colour Reddish-brown gas Colourless gas

solubility in water Very soluble in water Almost insoluble

effect on litmus Turns litmus red Neutral to litmus

density Denser than air Very slightly denser than air

effect on splint Extinguishes a burning splint and candle

Extinguishes a burning splint and candle

test for gas Brown gas and turns damp blue litmus paper red Turns brown on exposure to air

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Take this short quiz about oxides of nitrogen and nitric acid. Some questions have one correct response, whilst others have more than one.

Nitric Acid I can describe properties and uses of nitric acid.

Nitric acid is a very strong acid. As a dilute acid it shows most properties of the acids i.e. Turns pH paper red. Reacts with alkalis or bases to form nitrates and water.

e.g. NaOH (aq) + HNO3 (aq) → NaNO3(aq) + H2O (l) CuO (s) + 2HNO3 (aq) → Cu(NO3)2 (aq) + H2O (l)

Reacts with carbonates to give salt (nitrates), water and CO2 e.g. MgCO3 (s) + 2HNO3 (aq) → Mg(NO3)2 (aq) + H2O (l) + CO2(g) The nitrates produced from the reaction of nitric acid are highly soluble in water.

Cold very dilute HNO3 reacts with Mg or Ca to give the salt (nitrate) and hydrogen.

Nitric acid is a powerful oxidizing agent. When concentrated nitric acid reacts with copper any hydrogen produced initially is oxidised by more acid to form water and the reduction products of the acid are like NO2 , NO or even NH3.

Cu (s) + 4HNO3 (aq) Cu(NO3)2 (aq) + 2H2O (l) + 2NO2 (g)

When the concentrated acid is diluted with its own volume of water, it would react with copper to give:

3Cu (s) + 8HNO3 (aq) 3Cu(NO3)2 (aq) + 4H2O (l) + 2NO (g)

Uses of nitric acid

Manufacture of artificial fertilisers e.g. NH4NO3 Explosives TNT (trinitrotoluene) Dyes and artificial fibres such as nylon.

Watch this video of the reaction of copper and concentrated nitric acid

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End of Unit Questions 1. Ammonia, NH3, is easily identified as it changes damp red litmus paper to blue.

a) Suggest a further confirmatory test for ammonia.

b) Some dry ammonia is prepared in the laboratory by heating an ammonium salt with a metal hydroxide as shown in the diagram.

i) Complete and balance the reaction: Ca(OH)2(s) + 2NH4Cl(s) → CaCl2(s) + ____________(l) + ____________(g)

ii) Explain why the two reagents are in the solid state rather than in the aqueous phase. iii) Calcium oxide, CaO, is used to dry the ammonia gas produced. Suggest a reason why

concentrated sulfuric acid cannot be used instead of calcium oxide.

c) Ammonia solution can also be produced in the laboratory. i) Draw a labelled diagram of the set up you would use to produce ammonia solution. ii) Ammonia solution is weakly alkaline. With the help of a balanced chemical equation

explain why this is so. Adapted from State School Exam Chemistry – Year 10 – Track 3 – 2019

2. Ammonia is produced on an industrial scale by the Haber process.

a) Name the starting materials and also state the source of each one. b) Write a balanced chemical equation for the formation of ammonia. c) State three conditions required for the reaction in (b) to reach equilibrium in a short

time. d) Ammonia is an important product because it is used to make fertilisers. One such

example is ammonium nitrate. i) Name the two compounds required to prepare ammonium nitrate. ii) Write a balanced chemical equation for the preparation of ammonium nitrate using the

compounds mentioned in d (i). iii) Describe a chemical test to show the presence of the ammonium ion (NH4

+) in ammonium nitrate.

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3. A student decides to prepare nitrogen dioxide in the laboratory by heating lead(II) nitrate.

a) Write a balanced chemical equation for the reaction to produce nitrogen dioxide. b) Give two observations noted whilst heating lead(II) nitrate. c) Predict what would be observed when a piece of damp blue litmus paper is inserted in

nitrogen dioxide gas. d) Nitrogen dioxide is very soluble in water. Give the names of the two products formed

when nitrogen dioxide dissolves in water. e) Oxides of nitrogen are polluting and cause harmful effects. Explain how oxides of

nitrogen are formed from the emission of vehicles. Support your answer with balanced chemical equations.

4. Nitric acid is a very strong acid. a) Explain what the term ‘strong acid’ means. b) As a strong acid it can react with many different substances.

i) Write a balanced chemical equation for the reaction of nitric acid with magnesium oxide.

ii) Write a balanced chemical equation for the reaction of nitric acid with zinc carbonate. c) Besides behaving as an acid nitric acid is also an oxidising agent.

i) Complete the reaction below: Cu + 4HNO3 → _________ + ________ + _________

ii) Give the oxidation state of copper before and after the reaction: oxidation state of copper before the reaction: ___________ oxidation state of copper after the reaction: ___________

iii) Using oxidation numbers, explain what happens to copper in this reaction.

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Answers to End of Unit Questions 1a) Ammonia forms dense white cloud if exposed to hydrogen chloride gas. b i) Ca(OH)2(s) + 2NH4Cl(s) → CaCl2(s) + 2H2O(l) + 2NH3(g) b ii) The reagents must be in the solid state otherwise the ammonia formed would dissolve in the

round bottom flask rather than be collected by upward delivery. b iii) Ammonia being alkaline would react and neutralise the sulfuric acid.

ci) cii) NH3 (g) + H2O (l) ⇌ NH4

+ (aq) + OH- (aq) As ammonia dissolves in water OH- ions are formed which make it turn alkaline.

2a) Nitrogen is obtained from fractional distillation of liquid air. Hydrogen is produced from steam. b) N2 (g) + 3H2 (g) ⇌ 2NH3 (g) c) Finely divided iron as a catalyst; a temperature of 450◦C and a pressure of 200atm. di) ammonia solution and nitric acid. dii) NH3 (aq) + HNO3 (aq) → NH4NO3 (aq) diii) A solution of ammonium nitrate is prepared. Sodium hydroxide solution is added and the mixture

is heated. A pungent smelling gas is evolved which turns damp red litmus paper blue. 3a) 2Pb(NO3)2 (s) → 2PbO (s) + 4NO2 (g) + O2(g)

b) A brown gas is evolved and a yellow solid remains in the boiling tube. c) The damp blue litmus paper turns red. d) Nitric acid and nitrous acid. e) At high temperature in vehicles nitrogen from air combines with oxygen to form nitrogen

monoxide. This gas is then oxidised to form nitrogen dioxide. N2 (g) + O2 (g) ⇌ 2NO (g) 2NO (g) + O2(g) → 2NO2 (g)

4a) Nitric acid is a very strong acid which means that it completely ionises in solution. bi) 2HNO3 (aq) + MgO (s) → Mg(NO3)2 (aq) + H2O (l) bii) 2HNO3 (aq) + ZnCO3 (s) → Zn(NO3)2 (aq) + H2O (l) + CO2 (g) ci) Cu + 4HNO3 → Cu(NO3)2 (aq) + 2H2O (l) + 2NO2 (g) cii) oxidation state of copper before: 0 oxidation state of copper after: +2 ciii) due to an increase in oxidation numbers oxidation takes place.