Halogens – SS2 Chemistry Lesson Note (Second Term)

Halogens (salt formers) are the group VII elements: fluorine, chlorine, bromine, iodine and astatine.

Behavioural objectives

By the end of the lesson, students should be able to:

1. write the electronic configurations of fluorine, chlorine, bromine and iodine;
2. state the physical and chemical properties of the halogens;
3. describe the laboratory and industrial preparation of chlorine;
4. state the properties, test and uses of chlorine;
5. describe the preparation, properties and test for hydrogen chloride;
6. state the properties of chlorides and the test for a chloride.

What are the general features of the halogens?

The halogens are the most reactive non-metals, found in group VII of the periodic table.

• Valence electrons: each atom has seven electrons in its outermost shell.
• Molecules: they exist as diatomic molecules, such as F₂ and Cl₂.
• Colour: they are coloured elements.
• Ions: they ionise to form univalent negative ions, such as Cl^–.
• Compounds with metals: they form electrovalent (ionic) compounds with metals.

What is the electronic configuration of the halogens?

Each halogen has seven electrons in the outermost shell, one electron short of the noble gas structure.

• Octet: the chemistry of the group is dominated by the tendency to complete the octet by receiving one electron.
• Result: each atom gains one electron to form a univalent negative ion.

What are the physical properties of the halogens?

The halogens are coloured, volatile, non-polar diatomic molecules whose state changes from gas to solid down the group.

• Molecules: they exist in their natural states as non-polar diatomic molecules.
• State: fluorine and chlorine are gases, bromine is a liquid and iodine is a solid.
• Colour: fluorine is pale yellow, chlorine is greenish-yellow, bromine is red and iodine is violet. The colours deepen down the group.
• Odour: they have a typical penetrating smell.
• Volatility: they are volatile substances, and volatility decreases down the group.
• Solubility in water: all the halogens except fluorine dissolve to some extent in water.
• Fluorine and water: fluorine reacts with water to give oxygen and hydrogen fluoride.

What are the chemical properties of the halogens?

The halogens are very reactive, strongly electronegative elements that act as oxidising agents by gaining electrons.

Reactivity and electronegativity

• Reactivity: it decreases down the group. Fluorine is the most reactive halogen.
• Electronegativity: the halogens are strongly electronegative, and electronegativity decreases down the group.

Valency and bonding

• Ionic compounds: they are usually univalent and easily accept one electron from metals such as Na and K.
• Covalent compounds: they share electrons with themselves or with non-metals.

Halogens as oxidising agents

• Action: halogens accept electrons and form halide ions, especially in reactions with metals.
• Trend: oxidising power decreases down the group.
• Reaction with metals: halogens react with metals to form ionic compounds.
  2Na(s) + F₂(g) → 2NaF(s)

Reaction with hydrogen

The ease of reaction with hydrogen decreases down the group.

• Example: H₂(g) + Cl₂(g) → 2HCl(g)
• Stability: the stability of the hydrogen halides decreases down the group.
• State: hydrogen fluoride is a liquid with a boiling point of 19°C. The other hydrogen halides are gases.

Reaction with water

• Fluorine: it reacts vigorously with water and gives off oxygen gas.
• Chlorine: it reacts very slowly with water to give hydrochloric acid and oxochlorate (I) acid.
  H₂O(l) + Cl₂(g) → HCl(aq) + HOCl(aq)
• Decomposition: oxochlorate (I) acid later decomposes to give hydrochloric acid and oxygen gas.
• Bleaching: this oxygen is responsible for the bleaching action of moist chlorine gas and chlorine water.

What is chlorine?

Chlorine (Cl₂) is the most important element in the halogen family, a greenish-yellow poisonous gas.

• Occurrence: it does not occur as a free element in nature because it is too reactive.
• Combined state: it is usually found as chlorides, such as sodium chloride (common salt).

How is chlorine prepared in the laboratory?

Chlorine is prepared in the laboratory by oxidising concentrated hydrochloric acid, or a chloride with concentrated sulphuric acid, using a strong oxidising agent.

Method 1: oxidation of concentrated HCl

• Reagents: concentrated HCl and a strong oxidising agent such as MnO₂ or KMnO₄.
  MnO₂(s) + 4HCl(aq) → MnCl₂(aq) + 2H₂O(l) + Cl₂(g)

Method 2: heating NaCl and MnO₂ with concentrated H₂SO₄

• Reagents: a mixture of NaCl and MnO₂ heated with concentrated H₂SO₄.
  2NaCl(s) + MnO₂(s) + 2H₂SO₄(aq) → Na₂SO₄(aq) + MnSO₄(aq) + 2H₂O(l) + Cl₂(g)

Purifying and collecting the gas

1. Pass the gas through water to remove hydrogen chloride fumes.
2. Pass it through concentrated H₂SO₄ to dry it.
3. Collect it by downward delivery (upward displacement of air), because chlorine is denser than air.

How is chlorine prepared industrially?

Chlorine is manufactured industrially by the electrolysis of brine and molten NaCl, MgCl₂ or CaCl₂.

• Brine: concentrated sodium chloride solution.
• Molten chlorides: molten NaCl, MgCl₂ or CaCl₂.
• Where chlorine forms: it is released at the anode.

What are the physical properties of chlorine?

Chlorine is a greenish-yellow, poisonous gas with an unpleasant choking smell.

• Colour and smell: greenish-yellow gas with a choking smell.
• Toxicity: it is a poisonous gas.
• Density: it is about 2.5 times denser than air.
• Liquefaction: it can be liquefied under a pressure of about 6 atm.
• Solubility: it is moderately soluble in water.

What are the chemical properties of chlorine?

Chlorine is a very reactive element and a powerful oxidising agent that forms chlorides with most elements.

1. Bonding: it forms electrovalent compounds with metals and single covalent bond compounds with non-metals.
   2Na(s) + Cl₂(g) → 2NaCl(s)
   Cl₂(g) + H₂(g) → 2HCl(g)
2. Displacement: it displaces other halogens from solutions of their acids and salts.
   Cl₂(g) + 2NaI(aq) → 2NaCl(aq) + I₂(s)
3. Direct combination: it combines directly with other elements to form chlorides. The exceptions are oxygen, nitrogen, carbon and the noble gases.
   Ca(s) + Cl₂(g) → CaCl₂(s)
4. Affinity for hydrogen: it has a very strong affinity for hydrogen and removes hydrogen from its compounds.
   C₁₀H₁₆(l) + 8Cl₂(g) → 10C(s) + 16HCl(g)
5. Oxidising action: it oxidises green Fe²⁺ to yellow Fe³⁺.
   2FeCl₂(aq) + Cl₂(g) → 2FeCl₃(aq)
6. Bleaching action: in the presence of water, chlorine bleaches most dyes and inks except printer ink.
   H₂O(l) + Cl₂(g) → HCl(aq) + HOCl(aq)
   HOCl(aq) → HCl(aq) + [O]
   Dye (coloured) + [O] → [Dye + O] (colourless)
   The oxochlorate (I) acid decomposes to release oxygen, and the oxygen oxidises the dye to a colourless compound.
7. Reaction with slaked lime: it reacts with slaked lime solution to produce bleaching powder.
   Ca(OH)₂(aq) + Cl₂(g) → CaOCl₂.H₂O(s)
   CaOCl₂.H₂O is bleaching powder.

How does chlorine react with sodium hydroxide?

Chlorine reacts with sodium hydroxide to give different products, depending on the temperature and concentration of the alkali.

• Hot concentrated NaOH: it gives a mixture of trioxochlorate (V) and chloride of the metal.
• Cold dilute NaOH: it gives a pale yellowish mixture of oxochlorate (I) and chloride of the metal.

How do you test for chlorine?

Chlorine is identified by its effect on damp litmus paper and damp starch-iodide paper.

1. It turns damp blue litmus paper pink and then bleaches it. This shows it is an acidic gas.
2. It turns damp starch-iodide paper dark blue. Chlorine displaces iodine from the iodide, and the iodine liberated turns the starch blue.

What are the uses of chlorine?

Chlorine is used mainly as a germicide, a bleach and a raw material for making other chemicals.

• Germicide: it is a powerful germicide because of its oxidising nature.
• Bleaching agent: it bleaches cotton, wool, pulp and similar materials.
• Plastics and rubber: it is used to make polyvinyl chloride (PVC) and synthetic rubber.
• Organic compounds: it is used to make compounds such as CHCl₃ and CCl₄.
• Potassium chlorate: it is used to produce KClO₃, for making matches and fireworks.
• Weed killer: it is used to make NaClO₃, a weed killer.
• Antiseptics: it is used to make domestic antiseptics, such as acidified NaClO solution.

What is hydrogen chloride?

Hydrogen chloride (marine-acid gas) is a colourless gas at room temperature that dissolves in water to form hydrochloric acid.

• Occurrence: it occurs in traces in the air as an industrial by-product.
• Pollution: it is considered an air pollutant.
• Removal: it is very soluble in water, so it is easily washed down as acid rain.

How is hydrogen chloride prepared in the laboratory?

Hydrogen chloride is prepared by the action of hot concentrated H₂SO₄ on sodium chloride or another chloride.

2NaCl(s) + H₂SO₄(aq) → Na₂SO₄(aq) + 2HCl(g)

1. At a lower temperature, NaHSO₄ is formed first.
   NaCl(s) + H₂SO₄(aq) → NaHSO₄(aq) + HCl(g)
2. At a higher temperature, HCl gas is formed with Na₂SO₄.
   NaCl(s) + NaHSO₄(aq) → Na₂SO₄(aq) + HCl(g)
3. Pass the gas through concentrated H₂SO₄ in another flask to dry it.
4. Collect the gas by downward delivery, because it is denser than air and very soluble in water.

How is hydrogen chloride prepared industrially?

Pure HCl gas is produced on a large scale by direct combination of hydrogen and chlorine gas obtained from the electrolysis of brine.

H₂(g) + Cl₂(g) → 2HCl(g)

What are the physical properties of hydrogen chloride?

Hydrogen chloride is a colourless, sharp-smelling gas that is very soluble in water.

• Appearance and smell: pure HCl gas is colourless with a sharp irritating smell.
• Litmus: it turns damp blue litmus paper red.
• Density: it is about 1.25 times denser than air.
• Solubility in water: it is very soluble in water, forming aqueous HCl acid.
• Misty fumes: it forms misty fumes in moist air because it dissolves in the moisture to form tiny droplets of HCl acid.
• Combustion: it does not support combustion.

How does hydrogen chloride behave in different solvents?

• In non-polar solvents: it dissolves readily in chloroform and toluene, but the solution does not conduct electricity and has no acidic properties. HCl is a covalent molecule and does not ionise here.
• In water: it ionises. The ions formed are responsible for the acidic property and conductivity of the aqueous solution.

What are the chemical properties of hydrogen chloride?

1. With ammonia: it combines directly with NH₃ and produces white fumes of ammonium chloride.
   HCl(g) + NH₃(g) → NH₄Cl(s)
2. With metals: it reacts with various heated metals to form their chloride and hydrogen.
   Zn(s) + 2HCl(g) → ZnCl₂(s) + H₂(g)

How do you test for hydrogen chloride?

Hydrogen chloride is identified by dense white fumes with ammonia, or a white precipitate with silver trioxonitrate (V).

1. Dip a glass rod in ammonia solution and bring it near the gas jar. Dense white fumes on the glass rod show hydrogen chloride.
2. Add a few drops of silver trioxonitrate (V) solution to the gas jar and shake. A white precipitate of silver chloride shows hydrogen chloride.

What are chlorides?

Chlorides are normal salts formed when metallic ions replace the hydrogen ion in hydrochloric acid.

• Preparation: chlorides are prepared by a neutralisation reaction.
• Solubility: chlorides are soluble in water except a few.
• Soluble chlorides: NaCl, NH₄Cl, KCl, CuCl₂, CaCl₂.
• Insoluble chlorides: AgCl, PbCl₂, etc.

What are the properties of chlorides?

Chlorides are mostly stable to heat and release hydrogen chloride or chlorine when heated with concentrated sulphuric acid.

1. Effect of heat: most chlorides are not decomposed by heat. Ammonium chloride is an exception.
2. Recovery: chlorides are recovered from solution by evaporation to dryness or sometimes by crystallisation.
3. With hot concentrated tetraoxosulphate (VI) acid: they produce hydrogen chloride gas.
   2NaCl(s) + H₂SO₄(aq) → Na₂SO₄(aq) + 2HCl(g)
4. With the acid and a strong oxidising agent: on heating, chlorine is produced.
   2NaCl(s) + MnO₂(s) + 2H₂SO₄(aq) → Na₂SO₄(aq) + MnSO₄(aq) + 2H₂O(l) + Cl₂(g)

How do you test for chlorides?

A chloride is identified by a white precipitate of silver chloride that dissolves in excess ammonia solution.

1. Acidify the test solution with dilute trioxonitrate (V) acid to prevent precipitation of other salts.
2. Add a few drops of AgNO₃(aq) to the acidified solution in a test tube.
3. A white precipitate of AgCl forms.
4. Add excess NH₃(aq). The precipitate readily dissolves, which confirms a chloride.

Evaluation

1. Write the electronic configuration of chlorine and explain why halogens are univalent.
2. State the colour and physical state of fluorine, chlorine, bromine and iodine at room temperature.
3. Describe how chlorine is prepared in the laboratory from concentrated HCl, and write the equation.
4. Explain why moist chlorine bleaches dyes.
5. Describe the test for hydrogen chloride gas and the test for a chloride.

Assignment

1. Prepare a revision sheet on chlorine. Include four chemical properties with an equation for each, the difference between the products formed with hot concentrated NaOH and with cold dilute NaOH, and five uses of chlorine.