Nitrogens – SS2 Chemistry Lesson Note (Second Term)

Nitrogen is a colourless, odourless gas that makes up about 78% of the air and heads Group VA (Group 15) of the periodic table.

Behavioural objectives

1. State the members of Group VA and describe their general properties.
2. Describe the laboratory and industrial preparation of nitrogen and state its properties and uses.
3. Explain the stages of the nitrogen cycle.
4. Prepare, test for and state the properties of the oxides of nitrogen.
5. Describe the preparation, properties, tests and uses of ammonia and trioxonitrate (V) acid.

What are the elements of the nitrogen family (Group VA)?

The nitrogen family has five elements: nitrogen, phosphorus, arsenic, antimony and bismuth.

• Differences: there are more differences than similarities between the elements.
• Nitrogen: a diatomic gas that does not show allotropy and is the most electronegative member.
• Phosphorus: a typical non-metal that shows allotropy, for example white and red phosphorus.
• Arsenic and antimony: metalloids.
• Bismuth: a metal.

What are the electronic configurations of Group VA elements?

Each Group VA element has five electrons in its outer shell, written ns² np³.

• Nitrogen (7): 1s² 2s² 2p³
• Phosphorus (15): 1s² 2s² 2p⁶ 3s² 3p³
• Arsenic (33): 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p³
• Antimony (51): 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ 5s² 5p³

What is nitrogen and where does it occur?

Nitrogen is a Group VA gas that occurs chiefly as a free element in the air.

• Free state: it makes up 78% of the volume of the atmosphere.
• Combined state: it occurs in ammonia, urea, proteins and many other compounds.

How is nitrogen prepared in the laboratory?

Nitrogen is prepared in the laboratory either from air or by chemical methods.

From air (not pure)

1. Pass air through caustic soda solution to remove carbon (IV) oxide.
2. Pass the gas over heated copper turnings to remove oxygen.
3. Collect the gas left. It contains about 1% by volume of rare gases and is denser than pure nitrogen.

Chemical methods (pure nitrogen)

• Thermal decomposition of ammonium dioxonitrate (III): the salt is made by mixing sodium dioxonitrate (III) and ammonium chloride, then heated.
  
  NaNO₂(aq) + NH₄Cl(aq) → NH₄NO₂(aq) + NaCl(aq)
  
  NH₄NO₂(aq) → 2H₂O(l) + N₂(g)
• Thermal decomposition of ammonium heptaoxodichromate (VI):
  
  (NH₄)₂Cr₂O₇(s) → Cr₂O₃(s) + 4H₂O(l) + N₂(g)
• Oxidation of ammonia by hot copper (II) oxide:
  
  2NH₃(g) + 3CuO(s) → 3Cu(s) + 3H₂O(g) + N₂(g)
• Reduction of dinitrogen (I) oxide by red-hot copper:
  
  N₂O(g) + Cu(s) → CuO(s) + N₂(g)

How is nitrogen prepared industrially?

Nitrogen is obtained industrially by the fractional distillation of liquid air.

What are the physical properties of nitrogen?

• Appearance: colourless, odourless and tasteless gas.
• Density: pure nitrogen is slightly lighter than air.
• Solubility: slightly soluble in water.
• Melting point: -210°C.
• Boiling point: -196°C.

What are the chemical properties of nitrogen?

• With very electropositive metals: it forms nitrides.
  
  3Mg(s) + N₂(g) → Mg₃N₂(s)
• With hydrogen: it forms ammonia under the conditions of the Haber process.
  
  N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
• With oxygen: it forms nitrogen (II) oxide at very high temperature, such as in lightning or an electric arc.
  
  N₂(g) + O₂(g) → 2NO(g)

What are the uses of nitrogen?

• Ammonia manufacture: it is the raw material for making ammonia industrially.
• Cooling agent: liquid nitrogen is used to cool and freeze things.
• Food preservation: it fills packaged foods to prevent rancidity.

What is the nitrogen cycle?

The nitrogen cycle is the series of changes by which atmospheric nitrogen becomes soil nitrogen and returns to the air as free nitrogen.

1. Oxidation of atmospheric nitrogen

• Lightning: electrical discharge makes nitrogen and oxygen combine to form oxides of nitrogen.
• Rainwater: the oxides dissolve to form dioxonitrate (III) and trioxonitrate (V) acids.
• Soil: the acids react with mineral salts to form trioxonitrate (V) salts, which plants absorb.

N₂(g) + O₂(g) → 2NO(g)

2NO(g) + O₂(g) → 2NO₂(g)

4NO(g) + O₂(g) + 2H₂O(l) → 4HNO₂(aq)

4NO₂(g) + O₂(g) + 2H₂O(l) → 4HNO₃(aq)

2. Action of nitrogen-fixing bacteria

Nitrogen fixation is the conversion of atmospheric nitrogen into nitrogen compounds by soil microorganisms.

• Rhizobium: lives in root nodules of leguminous plants and makes organic nitrogen compounds for the host plant.
• Release: the compounds enter the soil when the plants die.
• Azotobacter and Clostridium: free-living soil bacteria that also fix atmospheric nitrogen into compounds plants can absorb.

3. Decay of organic matter

Nitrification is the conversion of ammonia in the soil into trioxonitrate (V) by bacteria.

1. Plants and animals die.
2. Putrefying bacteria and fungi change the organic nitrogen compounds into ammonia.
3. Nitrifying bacteria (Nitrosomonas and Nitrobacter) change the ammonia into trioxonitrate (V).
4. Plants absorb the trioxonitrate (V) again.

4. Denitrification

Denitrification is the conversion of trioxonitrate (V) salts in the soil into free nitrogen by denitrifying bacteria.

• Result: nitrogen escapes into the atmosphere.
• Effect on soil: it reduces the amount of trioxonitrate (V) in the soil.

What are the oxides of nitrogen?

The three common oxides of nitrogen are nitrogen (I) oxide (N₂O), nitrogen (II) oxide (NO) and nitrogen (IV) oxide (NO₂).

What is nitrogen (I) oxide, N₂O?

Nitrogen (I) oxide is a colourless gas called laughing gas because it causes uncontrollable laughter when inhaled.

Laboratory preparation

• Method: thermal decomposition of ammonium trioxonitrate (V).
• Safety: the salt is not heated directly, since the reaction is exothermic and may become uncontrollable and explode.

KNO₃(s) + NH₄Cl(s) → KCl(s) + NH₄NO₃(s)

NH₄NO₃(s) → 2H₂O(g) + N₂O(g)

Physical properties

• Appearance: colourless gas with a faint pleasant smell and a sweetish taste.
• Solubility: fairly soluble in cold water.
• Density: about 1.5 times denser than air.
• Litmus: neutral to moist litmus paper.

Chemical properties

• Heating: it decomposes on strong heating (about 600°C) to nitrogen and oxygen.
  
  2N₂O(g) → O₂(g) + 2N₂(g)
• Combustion: it supports the burning of any substance hot enough to decompose it.
  
  Mg(s) + N₂O(g) → MgO(s) + N₂(g)
• Reduction: heated copper or iron reduces it to nitrogen.
  
  Cu(s) + N₂O(g) → N₂(g) + CuO(s)

Test for N₂O

1. Insert a glowing splint into the gas jar. If it rekindles, the gas is oxygen or nitrogen (I) oxide.
2. Smell the gas and add nitrogen (II) oxide. Oxygen gives brown fumes.
3. A gas with a pleasant smell that gives no brown fumes with nitrogen (II) oxide is nitrogen (I) oxide.

Use

• Anaesthetic: used for minor surgical operations.

What is nitrogen (II) oxide, NO?

Nitrogen (II) oxide is a colourless, poisonous gas that is almost insoluble in water.

Laboratory preparation

• Method: react copper with 50% (dilute) trioxonitrate (V) acid.
• Brown fumes: some NO reacts with oxygen in the flask to form brown nitrogen (IV) oxide, which dissolves in the water the gas is passed through.

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

Physical properties

• Appearance: colourless and poisonous gas.
• Solubility: almost insoluble in water.
• Density: slightly denser than air.
• Litmus: neutral.

Chemical properties

• With oxygen: it reacts readily to form brown fumes of nitrogen (IV) oxide.
  
  2NO(g) + O₂(g) → 2NO₂(g)
• Heating: at high temperatures it decomposes into equal volumes of nitrogen and oxygen.
  
  2NO(g) → N₂(g) + O₂(g)
• Reduction by hot metals:
  
  2Cu(s) + 2NO(g) → 2CuO(s) + N₂(g)
• Reducing agent: it slowly decolourises acidified potassium tetraoxomanganate (VII).
  
  3MnO₄^–(aq) + 4H⁺(aq) + 5NO(g) → 3Mn²⁺(aq) + 5NO₃^–(aq) + 2H₂O(l)

Test for NO

• Using air: open the gas jar. If the gas turns reddish-brown, it is NO.
• Using iron (II) tetraoxosulphate (VI): pour FeSO₄ solution, acidified with a little dilute H₂SO₄, into the gas jar. If the solution turns dark brown, the gas is NO.

What is nitrogen (IV) oxide, NO₂?

Nitrogen (IV) oxide is a reddish-brown, poisonous gas with an irritating smell.

Laboratory preparation

• Method: thermal decomposition of lead (II) trioxonitrate (V).
• Reason for the salt: it has no water of crystallisation to interfere with the preparation.
• Collection: the gas mixture passes through a U-tube in a freezing mixture. NO₂ liquefies and oxygen escapes.

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

Physical properties

• Appearance: reddish-brown gas.
• Smell: irritating and poisonous.
• Litmus: turns damp blue litmus paper red, and dissolves in water to give an acidic solution.
• Liquefaction: it liquefies to a yellow liquid at 21°C.
• Density: much heavier than air.

Chemical properties

• Dimer: at low temperatures it exists mainly as dinitrogen tetraoxide, N₂O₄. On heating it breaks down in stages.
  
  N₂O₄(g) ⇌ 2NO₂(g) ⇌ 2NO(g) + O₂(g)
• Combustion: it supports combustion because it releases oxygen on heating.
  
  2NO₂(g) → 2NO(g) + O₂(g)
• Reduction: a reducing agent such as hot copper reduces it to nitrogen.
  
  4Cu(s) + 2NO₂(g) → 4CuO(s) + N₂(g)
• With water: it forms a mixture of dioxonitrate (III) and trioxonitrate (V) acids, so it is a mixed acid anhydride.
  
  H₂O(l) + 2NO₂(g) → HNO₂(aq) + HNO₃(aq)
• With alkalis: it forms a mixture of dioxonitrate (III) and trioxonitrate (V) salts.
  
  2KOH(aq) + 2NO₂(g) → KNO₃(aq) + KNO₂(aq) + H₂O(l)

How do the three oxides of nitrogen compare?

What is ammonia?

Ammonia (NH₃) is a hydride of nitrogen and the only common alkaline gas.

• Natural source: it forms when nitrogenous matter decays in the absence of air.
• In the air: only traces occur, because it is very soluble and rainwater washes it into the soil.

How is ammonia prepared in the laboratory?

1. Heat a mixture of calcium hydroxide (slaked lime) and ammonium chloride.
2. Dry the gas by passing it through calcium oxide (quicklime).
3. Do not use concentrated H₂SO₄ or fused CaCl₂ to dry it, because ammonia is alkaline and reacts with them.

Ca(OH)₂(s) + 2NH₄Cl(s) → CaCl₂(s) + 2H₂O(l) + 2NH₃(g)

How is ammonia made industrially?

Ammonia is manufactured from nitrogen and hydrogen by the Haber process.

• Mixture: nitrogen and hydrogen in the ratio 1:3 by volume.
• Catalyst: finely divided iron.
• Temperature: about 450°C.
• Pressure: about 200 atm.
• Yield: about 15% under these conditions, because the reaction is reversible.

N₂(g) + 3H₂(g) ⇌ 2NH₃(g) + heat

What are the physical properties of ammonia?

• Appearance: colourless gas with a characteristic choking smell.
• Toxicity: poisonous in large quantities because it affects the respiratory muscles.
• Nature: the only known alkaline gas.
• Density: about 1.7 times less dense than air.
• Melting point: about -78°C (-77.7°C).
• Boiling point: about -33°C.

What are the chemical properties of ammonia?

• Burning in oxygen: it forms water vapour and nitrogen.
  
  4NH₃(g) + 3O₂(g) → 6H₂O(g) + 2N₂(g)
• Reducing agent, with copper (II) oxide:
  
  3CuO(s) + 2NH₃(g) → 3Cu(s) + 3H₂O(l) + N₂(g)
• Reducing agent, with chlorine:
  
  3Cl₂(g) + 8NH₃(g) → 6NH₄Cl(s) + N₂(g)
• With carbon (IV) oxide: it forms urea and water.
  
  2NH₃(g) + CO₂(g) → (NH₂)₂CO(s) + H₂O(l)
• With acids: it forms ammonium salts.
  
  2NH₃(g) + H₂SO₄(aq) → (NH₄)₂SO₄(aq)

How do you test for ammonia?

Ammonia is recognised by its choking smell and confirmed by two tests.

1. Litmus paper: damp red litmus turns blue in the gas.
2. Hydrochloric acid: a glass rod dipped in concentrated HCl gives white fumes in the gas.

What are the uses of ammonia?

• Acid and soda manufacture: making trioxonitrate (V) acid, and sodium trioxocarbonate (IV) by the Solvay process.
• Refrigerant: liquid ammonia is used for cooling.
• Water softening: aqueous ammonia softens temporary hard water.
• Laundry: aqueous ammonia removes grease and oil stains.
• Fertilisers: it is used to make urea and ammonium salts.

What is trioxonitrate (V) acid, HNO₃?

Trioxonitrate (V) acid (nitric acid) is a volatile, corrosive acid that is also a strong oxidising agent.

How is trioxonitrate (V) acid prepared in the laboratory?

• Method: concentrated H₂SO₄, which is less volatile, displaces the acid from a trioxonitrate (V) salt.
• Salt used: potassium or sodium trioxonitrate (V), because they are cheap.
• Apparatus: all-glass, because the acid vapour attacks cork and rubber.

KNO₃(s) + H₂SO₄(conc) → KHSO₄(s) + HNO₃(g)

How is trioxonitrate (V) acid made industrially?

It is made by the catalytic oxidation of ammonia in three steps.

1. Ammonia and excess air pass over a platinum-rhodium catalyst at about 700°C to form nitrogen (II) oxide (about 96% conversion).
   
   4NH₃(g) + 5O₂(g) → 4NO(g) + 6H₂O(g)
2. Nitrogen (II) oxide is cooled and mixed with excess air to form nitrogen (IV) oxide.
   
   2NO(g) + O₂(g) → 2NO₂(g)
3. Nitrogen (IV) oxide and excess air dissolve in hot water to give an acid solution of up to 50% concentration.
   
   4NO₂(g) + 2H₂O(l) + O₂(g) → 4HNO₃(aq)

What are the physical properties of trioxonitrate (V) acid?

• Appearance: the pure acid is a colourless fuming liquid with a sharp choking smell.
• Yellow colour: it turns yellow because it decomposes to nitrogen (IV) oxide, which redissolves.
• Boiling point: 86°C.
• Melting point: -42°C.
• Density: 1.52 g cm^-3.
• Solubility: miscible with water in all proportions. It forms a constant boiling mixture at 121°C.
• Corrosiveness: the concentrated acid is corrosive.
• Litmus: the dilute acid turns blue litmus red.

What are the chemical properties of trioxonitrate (V) acid?

As an acid

• With bases and alkalis: it forms a trioxonitrate (V) salt and water only.
  
  NaOH(aq) + HNO₃(aq) → NaNO₃(aq) + H₂O(l)
• With trioxocarbonates (IV): it liberates carbon (IV) oxide.
  
  CaCO₃(s) + 2HNO₃(aq) → Ca(NO₃)₂(aq) + H₂O(l) + CO₂(g)
• With metals: unlike other acids, it rarely gives hydrogen. Only a very dilute solution does so, with Ca, Mg or Mn.

As an oxidising agent

• With non-metals: it oxidises them to their oxides or acids.
  
  S(s) + 6HNO₃(aq) → H₂SO₄(aq) + 2H₂O(l) + 6NO₂(g)
• With Cu, Pb, Hg and Ag: it forms the trioxonitrate (V) salt and nitrogen (IV) oxide if concentrated, or nitrogen (II) oxide if moderately concentrated.
• With aluminium and iron: concentrated acid does not attack them, because a passive oxide coating forms. Aluminium or iron containers can therefore carry it.
• With hydrogen sulphide: it oxidises it to sulphur.
  
  H₂S(g) + 2HNO₃(aq) → S(s) + 2H₂O(l) + 2NO₂(g)
• With iron (II) salts: it oxidises them to iron (III) salts.
  
  6Fe²⁺(aq) + 8H⁺(aq) + 2NO₃^–(aq) → 6Fe³⁺(aq) + 4H₂O(l) + 2NO(g)

What are the uses of trioxonitrate (V) acid?

• Laboratory: used as an acid, an oxidising agent and a nitrating agent.
• Fibres: used in making nylon and Terylene.
• Rocket fuel: used as an oxidiser in rocket fuel.
• Manufacture: used to make fertilisers, dyes, drugs and explosives.

Evaluation

1. Name the five elements of Group VA and state which is a metal.
2. Write the equation for the laboratory preparation of nitrogen from ammonium dioxonitrate (III).
3. Explain how lightning adds trioxonitrate (V) to the soil.
4. State the tests that distinguish nitrogen (II) oxide from ammonia.
5. Describe the three steps in the industrial preparation of trioxonitrate (V) acid.

Assignment

Draw a labelled flow chart of the nitrogen cycle, then write the conditions used in the Haber process and give two uses of ammonia.