Halogens – SS2 Chemistry Lesson Note (Second Term)

Halogens (salt formers) are found in group VII of the periodic table. They are the most reactive non-metals. They have seven valence electrons, exist as diatomic molecules, are coloured, and ionise to form univalent negative ions. They form electrovalent compounds with metals. The group contains fluorine, chlorine, bromine, iodine and astatine.

Behavioural objectives

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

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

What is the electronic configuration of the halogens?

The halogens, meaning “salt makers”, are found in group VII of the periodic table. Their electronic configuration is one electron short of the noble gas structure (that is, they have seven electrons in the outermost shell). The chemistry of the group is dominated by the tendency to complete the octet arrangement by receiving an electron. The electronic configurations of the halogens are shown below:

Fluorine = 9: 1s² 2s² 2p⁵

Chlorine = 17: 1s² 2s² 2p⁶ 3s² 3p⁵

Bromine = 35: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁵

Iodine = 53: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ 5s² 5p⁵

What are the physical properties of the halogens?

1. They are usually univalent, and easily accept one electron from other atoms to form ionic compounds (especially from metals, e.g. Na and K). They also share electrons with themselves or with non-metals to form covalent compounds.
2. They exist in their natural states as non-polar diatomic molecules.
3. Fluorine and chlorine are gases, bromine is a liquid and iodine is a solid.
4. The halogens are coloured, with a typical penetrating odour. The colours deepen down the group.
5. Fluorine is pale yellow, chlorine is greenish-yellow, bromine is red and iodine is violet.
6. They are volatile substances. Their volatility decreases down the group.
7. All the halogens except fluorine dissolve to some extent in water. Fluorine reacts with water to give oxygen and hydrogen fluoride.

What are the chemical properties of the halogens?

1. The halogens are very reactive elements. Their reactivity decreases down the group. Fluorine is the most reactive halogen. They are also strongly electronegative, and their electronegativity decreases down the group.
2. As oxidising agents: halogens are strong oxidising agents. They do so by accepting electrons and forming halide ions, especially in the reaction with metals. The oxidising power decreases down the group.
   Reaction with metals: halogens react with metals to form ionic compounds.
   2Na(s) + F₂(g) → 2NaF(s)
3. Reaction with hydrogen: fluorine explodes with hydrogen even in the dark. Chlorine reacts slowly in the dark but explodes in bright sunlight. Bromine reacts with hydrogen in the presence of a platinum catalyst, while iodine reacts partially with hydrogen on heating. Example:
   H₂(g) + Cl₂(g) → 2HCl(g)
4. The stability of the hydrogen halides decreases down the group. Hydrogen fluoride is a liquid with a boiling point of 19°C. The other hydrogen halides are gases.
5. Reaction with water: fluorine reacts vigorously with water to give off oxygen gas. Chlorine reacts very slowly with water to give a mixture of hydrochloric acid and oxochlorate (I) acid, which later decomposes to give hydrochloric acid and oxygen gas. The oxygen gas given off by the oxochlorate (I) acid is responsible for the bleaching action of moist chlorine gas and chlorine water.
   H₂O(g) + Cl₂(g) → HCl(aq) + HOCl(aq)

What is chlorine?

Chlorine is the most important element in the halogen family. It does not occur as a free element in nature because it is too reactive. It is usually found in a combined state as chlorides.

Laboratory preparation of chlorine

1. By the oxidation of concentrated HCl with a strong oxidising agent such as MnO₂ or KMnO₄.
   MnO₂(s) + 4HCl(aq) → MnCl₂(aq) + 2H₂O(l) + Cl₂(g)

2. By heating concentrated H₂SO₄ with a mixture of NaCl and MnO₂.
   2NaCl(s) + MnO₂(s) + 2H₂SO₄(aq) → Na₂SO₄(aq) + MnSO₄(aq) + 2H₂O(l) + Cl₂(g)

Industrial preparation of chlorine

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

Physical properties of chlorine

1. Chlorine is a greenish-yellow gas with an unpleasant choking smell.
2. It is a poisonous gas.
3. It is about 2.5 times denser than air.
4. It can be liquefied under a pressure of about 6 atm.
5. It is moderately soluble in water.

Chemical properties of chlorine

1. It is very reactive and tends to attain stability by forming 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. It displaces other halogens from solutions of their acids and salts.
   Cl₂(g) + 2NaI(aq) → 2NaCl(aq) + I₂(g)
3. It combines directly with other elements, except oxygen, nitrogen, carbon and the noble gases, to form chlorides.
   Ca(s) + Cl₂(g) → CaCl₂(s)
4. It has a very strong affinity for hydrogen. It removes hydrogen from its compounds.
   C₁₀H₁₆(l) + 8Cl₂(g) → 10C(s) + 16HCl(g)
5. It is a powerful oxidising agent. It oxidises green Fe²⁺ to yellow Fe³⁺.
   2FeCl₂(aq) + Cl₂ → 2FeCl₃(aq)
6. It has a bleaching action. In the presence of water, chlorine bleaches most dyes and inks except printer ink. The bleaching action is due to its ability to react with water to form oxochlorate (I) acid, which decomposes to release oxygen. The oxygen oxidises the dye to form a colourless compound.
   H₂O(l) + Cl₂(g) → HCl(aq) + HOCl(aq)
   HOCl(aq) → HCl(aq) + [O]
   Dye (coloured) + [O] → [Dye + O] (colourless)
7. It reacts with a hot concentrated NaOH solution to give a mixture of trioxochlorate (V) and chloride of the metal.
   6NaOH + 3Cl₂(g) → NaClO₃(aq) + 5NaCl(aq) + 3H₂O(l)
   NaClO₃ is sodium trioxochlorate (V).
8. With a cold dilute solution of NaOH, a pale yellowish mixture of oxochlorate (I) and chloride of the metal is formed.
   2NaOH(aq) + Cl₂(g) → NaOCl(aq) + NaCl(aq) + H₂O(l)
9. 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 do you test for chlorine?

1. It turns damp blue litmus paper pink and then bleaches it. 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?

1. It is a powerful germicide (due to its oxidising nature).
2. It is used as a bleaching agent for cotton, wool, pulp, etc.
3. It is used in the manufacture of polyvinyl chloride (PVC) and synthetic rubber.
4. It is used in the manufacture of organic compounds, e.g. CHCl₃ and CCl₄.
5. It is used in producing KClO₃, for making matches and fireworks.
6. It is used for making NaClO₃, a weed killer.
7. It is used for making domestic antiseptics, e.g. acidified NaClO solution.

What is hydrogen chloride?

Hydrogen chloride (marine-acid gas) exists as a gas at room temperature. It dissolves in water to form hydrochloric acid. It occurs in traces in the air as an industrial by-product and is considered an air pollutant, but it can easily be washed down as acid rain since it is very soluble in water.

Laboratory preparation of hydrogen chloride

The gas is prepared by the action of hot concentrated H₂SO₄ on any soluble chloride. Example:

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

Note: NaHSO₄ is first formed at a lower temperature and later at a higher temperature HCl gas is formed. The gas is dried by passing it through concentrated H₂SO₄ in another flask and collected.

Industrial preparation of hydrogen chloride

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

H₂(g) + Cl₂(g) → 2HCl(g)

Physical properties of hydrogen chloride

1. Pure HCl gas is colourless and has a sharp irritating smell.
2. It turns damp blue litmus paper red.
3. It is about 1.25 times denser than air.
4. It is very soluble in water, forming aqueous HCl acid.
5. It dissolves readily in non-polar solvents like chloroform and toluene, but the solution does not conduct electricity and has no acidic properties. Hydrogen chloride is a covalent molecule and does not ionise when it dissolves in non-polar solvents. It does ionise when it dissolves in water, and the ions formed are responsible for the acidic property and conductivity of the aqueous solution.
6. It forms misty fumes in moist air because it dissolves in the moisture to form tiny droplets of HCl acid.
7. It does not support combustion.

Chemical properties of hydrogen chloride

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

How do you test for hydrogen chloride?

1. A glass rod that has been dipped in ammonia solution is brought near the gas jar containing the unknown gas. If dense white fumes form on the glass rod, the gas is hydrogen chloride.
2. A few drops of silver trioxonitrate (V) solution are added to the gas jar containing the unknown gas and shaken. If a white precipitate of silver chloride is observed, the gas is hydrogen chloride.

What are chlorides?

Chlorides are normal salts formed when metallic ions replace the hydrogen ion in hydrochloric acid. Chlorides are prepared by a neutralisation reaction. Chlorides are soluble in water except a few.

Soluble chlorides: NaCl, NH₄Cl, KCl, CuCl₂, CaCl₂

Insoluble chlorides: AgCl, PbCl₂, etc.

Properties of chlorides

1. Chlorides are not decomposed by heat. They can only be recovered from solution by evaporation to dryness or sometimes by crystallisation.
2. They react with hot concentrated tetraoxosulphate (VI) acid to produce hydrogen chloride gas.
   2NaCl(s) + H₂SO₄(aq) → Na₂SO₄(aq) + 2HCl(g)
3. On heating a chloride with concentrated tetraoxosulphate (VI) acid in the presence of a strong oxidising agent, chlorine is produced.
   ZnCl₂(s) + KMnO₄(s) + 2H₂SO₄(aq) → ZnSO₄(aq) + K₂SO₄(aq) + 2MnO₂(aq) + 2H₂O(l) + Cl₂(g)

How do you test for chlorides?

Acidify the test solution with dilute trioxonitrate (V) acid to prevent precipitation of other salts. Add a few drops of AgNO₃(aq) to the acidified solution in a test tube. A white precipitate of AgCl that readily dissolves in excess NH₃(aq) solution indicates the presence of 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. List four chemical properties of chlorine and write an equation for each.
2. State the difference between the products formed when chlorine reacts with hot concentrated NaOH and with cold dilute NaOH.
3. List five uses of chlorine.