Group VII: the halogens

What you'll learn

Group VII elements—the halogens—appear regularly in CIE IGCSE Chemistry papers, testing your understanding of periodic trends, displacement reactions, and bonding. This topic accounts for both short-answer and extended-response questions worth 8-12 marks across typical papers. Mastering halogen reactivity patterns and being able to predict and explain displacement reactions is essential for achieving top grades.

Key terms and definitions

Halogen — a Group VII element with seven electrons in its outer shell, existing as diatomic molecules (F₂, Cl₂, Br₂, I₂).

Displacement reaction — a reaction where a more reactive halogen removes (displaces) a less reactive halogen from its compound.

Diatomic molecule — a molecule consisting of two atoms covalently bonded together, such as Cl₂ or Br₂.

Halide ion — a negatively charged ion formed when a halogen atom gains one electron (F⁻, Cl⁻, Br⁻, I⁻).

Oxidising agent — a substance that accepts electrons from another substance during a redox reaction; halogens act as oxidising agents.

Reactivity series — the order of elements arranged by their reactivity; for halogens, reactivity decreases down the group.

Reduction — the gain of electrons; halide ions are formed when halogens are reduced.

Covalent bonding — the sharing of electron pairs between atoms; halogen molecules contain single covalent bonds.

Core concepts

Physical properties and trends down Group VII

The halogens show clear trends in physical properties as you move down the group from fluorine to iodine:

State at room temperature:

• Fluorine (F₂) — pale yellow gas
• Chlorine (Cl₂) — green gas
• Bromine (Br₂) — red-brown liquid (volatile, producing brown fumes)
• Iodine (I₂) — grey-black solid (sublimes to form purple vapour)

Melting and boiling points increase down the group because:

• The molecules become larger with more electrons
• Stronger intermolecular forces (van der Waals forces) exist between larger molecules
• More energy is required to separate the molecules

Atomic radius increases down the group as each successive element has an additional electron shell.

Colour intensity increases down the group, from pale yellow through green and brown to dark grey-black.

All halogens exist as diatomic molecules with single covalent bonds between atoms (X—X). This is because each halogen atom has seven outer electrons and needs one more to complete its outer shell. Sharing one electron pair satisfies both atoms.

Chemical reactivity and the reactivity trend

Halogens are highly reactive non-metals that readily form ionic compounds with metals and covalent compounds with non-metals. The key trend is:

Reactivity decreases down Group VII: F > Cl > Br > I

This trend exists because:

1. All halogens gain one electron to form 1− ions when reacting
2. The attractive force between the nucleus and incoming electron decreases down the group
3. As atomic radius increases, the outer shell is further from the nucleus
4. Increased electron shielding from inner shells reduces nuclear attraction
5. Therefore, fluorine attracts electrons most strongly and is most reactive

Reactions with metals

Halogens react with metals to form ionic halides (metal halides). The general equation:

Metal + Halogen → Metal halide

Specific examples:

Sodium + chlorine: 2Na(s) + Cl₂(g) → 2NaCl(s)

This reaction is vigorous, producing bright yellow flames and white sodium chloride.

Iron + chlorine: 2Fe(s) + 3Cl₂(g) → 2FeCl₃(s)

Iron glows red-hot and brown iron(III) chloride forms.

Aluminium + bromine: 2Al(s) + 3Br₂(l) → 2AlBr₃(s)

A more reactive halogen produces a more vigorous reaction with the same metal. Chlorine reacts more vigorously with iron than bromine does.

Displacement reactions of halogens

Displacement reactions are the most commonly tested aspect of halogen chemistry in CIE IGCSE papers. A more reactive halogen will displace a less reactive halogen from its compound (aqueous halide solution).

General pattern:

• Chlorine displaces bromide and iodide ions
• Bromine displaces iodide ions but not chloride ions
• Iodine cannot displace chloride or bromide ions

Chlorine + potassium bromide: Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq)

Observation: The colourless solution turns orange-brown as bromine forms.

Ionic equation: Cl₂(aq) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(aq)

Chlorine + potassium iodide: Cl₂(aq) + 2KI(aq) → 2KCl(aq) + I₂(aq)

Observation: The colourless solution turns brown (or black if concentrated) as iodine forms.

Ionic equation: Cl₂(aq) + 2I⁻(aq) → 2Cl⁻(aq) + I₂(aq)

Bromine + potassium iodide: Br₂(aq) + 2KI(aq) → 2KBr(aq) + I₂(aq)

Observation: The orange-brown solution turns darker brown/black as iodine forms.

No reaction examples:

Bromine + potassium chloride — no reaction (bromine is less reactive than chlorine) Iodine + potassium chloride — no reaction Iodine + potassium bromide — no reaction

Understanding displacement in terms of electron transfer

Displacement reactions are redox reactions:

When chlorine displaces bromide:

• Chlorine molecules are reduced: Cl₂ + 2e⁻ → 2Cl⁻
• Bromide ions are oxidised: 2Br⁻ → Br₂ + 2e⁻
• Chlorine acts as an oxidising agent (accepts electrons)

More reactive halogens are stronger oxidising agents because they attract electrons more readily.

Uses of halogens and their compounds

Understanding practical applications demonstrates the relevance of halogen chemistry:

Chlorine:

• Sterilising water supplies (kills bacteria)
• Manufacturing bleach (sodium chlorate(I), NaClO)
• Producing hydrochloric acid (HCl)
• Manufacturing PVC (polyvinyl chloride) plastic

Bromine:

• Flame retardants in furniture and electronics
• Agricultural pesticides
• Pharmaceuticals

Iodine:

• Antiseptic solutions (iodine in alcohol)
• Treatment of thyroid disorders
• Photographic chemicals (silver iodide)

Fluorine compounds:

• Toothpaste (sodium fluoride strengthens tooth enamel)
• Non-stick coatings (PTFE/Teflon)
• Refrigerants (though CFCs are now banned due to ozone depletion)

Testing for halide ions

Silver nitrate solution tests for halide ions:

1. Add dilute nitric acid to the test solution (removes carbonates that would interfere)
2. Add silver nitrate solution
3. Observe the precipitate colour:

• Chloride (Cl⁻) — white precipitate of silver chloride (AgCl)
• Bromide (Br⁻) — cream precipitate of silver bromide (AgBr)
• Iodide (I⁻) — yellow precipitate of silver iodide (AgI)

The ionic equation for chloride: Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

Worked examples

Example 1: Predicting displacement reactions

Question: A student adds chlorine water to separate solutions of sodium bromide and sodium iodide.

(a) State what the student would observe in each case. [2] (b) Write a balanced chemical equation for the reaction between chlorine and sodium iodide. [2] (c) Explain why chlorine can displace both bromine and iodine from their compounds. [2]

Solution:

(a)

• With sodium bromide: the solution turns orange-brown/brown [1]
• With sodium iodide: the solution turns brown/dark brown [1]

(b) Cl₂ + 2NaI → 2NaCl + I₂ [1 for correct formulae, 1 for balancing]

Alternative (ionic equation): Cl₂ + 2I⁻ → 2Cl⁻ + I₂

(c) Chlorine is more reactive than both bromine and iodine [1] because chlorine atoms attract electrons more strongly/have smaller atomic radius/less electron shielding [1]

Example 2: Explaining trends

Question: The table shows information about three halogens.

Halogen	Boiling point (°C)
Chlorine	−34
Bromine	59
Iodine	184

(a) Describe the trend in boiling points. [1] (b) Explain this trend in terms of molecular structure and forces. [3]

Solution:

(a) Boiling point increases down the group [1]

(b)

• The molecules become larger/have more electrons [1]
• The intermolecular forces/van der Waals forces become stronger [1]
• More energy is needed to separate the molecules [1]

Example 3: Practical observation question

Question: A student adds bromine water to a solution of potassium iodide in a test tube.

(a) State what the student would observe. [1] (b) Write a balanced symbol equation for the reaction. [2] (c) Explain what would happen if the student instead added bromine water to potassium chloride solution. [1]

Solution:

(a) The solution turns darker brown/brown-black [1]

(b) Br₂ + 2KI → 2KBr + I₂ [2] (1 mark for correct formulae, 1 mark for balancing)

(c) No reaction/no change observed [1] (because bromine is less reactive than chlorine and cannot displace it)

Common mistakes and how to avoid them

• Mistake: Writing halogen elements as single atoms (Cl, Br, I) instead of diatomic molecules (Cl₂, Br₂, I₂) in equations. Correction: Always remember halogens exist as diatomic molecules under normal conditions. Check every equation includes the subscript 2.

• Mistake: Stating that reactivity increases down Group VII, confusing halogens with Group I metals. Correction: Halogen reactivity decreases down the group (opposite to metals). Link this to electron gain—atoms at the top attract electrons more strongly.

• Mistake: Predicting that iodine will displace chlorine from chloride compounds. Correction: Only a more reactive halogen displaces a less reactive one. Use the reactivity order F > Cl > Br > I to predict displacement. Iodine is least reactive so displaces nothing.

• Mistake: Describing displacement observations vaguely as "colour change" without specifying colours. Correction: Learn specific colour observations: chlorine is green, bromine is orange-brown, iodine is brown. State both initial and final colours for full marks.

• Mistake: Confusing melting/boiling point trends with reactivity trends. Correction: Physical properties (melting/boiling points) increase down the group due to stronger intermolecular forces. Chemical reactivity decreases down the group due to nuclear attraction. These are separate trends with different explanations.

• Mistake: Writing incorrect formulae for halides, such as NaCl₂ or KBr₂. Correction: Halide ions have a 1− charge (Cl⁻, Br⁻, I⁻). When combined with Group I metals (1+ charge), the ratio is always 1:1 (NaCl, KBr, KI).

Exam technique for Group VII: the halogens

• Displacement reaction questions typically ask you to predict observations and write equations. Always state specific colours (not just "colour change") and include state symbols for full marks. Practice writing both full and ionic equations as either may be required.

• "Explain" questions about reactivity trends require you to link atomic structure to electron gain. Mention atomic radius increasing, electron shielding increasing, and nuclear attraction to incoming electrons decreasing. Two or three developed points usually earn full marks.

• Practical observation questions test whether you know halogen colours and can predict displacement. Create a table memorising: chlorine (green), bromine (orange-brown/red-brown), iodine (brown/grey-black solid, purple vapour). Know which displaces which.

• Comparison questions may ask you to compare halogens with Group I metals or explain why halogens form covalent molecules but ionic compounds with metals. Address both bonding types clearly, referencing electron transfer (ionic) versus electron sharing (covalent).

Quick revision summary

Group VII halogens exist as diatomic molecules with reactivity decreasing down the group (F > Cl > Br > I) due to increasing atomic radius and electron shielding. Physical properties like boiling point increase down the group because of stronger intermolecular forces. More reactive halogens displace less reactive ones from aqueous halide solutions in redox reactions—chlorine displaces both bromide and iodide; bromine displaces only iodide. Halogens react vigorously with metals to form ionic halides. Key colours: chlorine is green, bromine is orange-brown, iodine is brown-black. Silver nitrate tests produce white (chloride), cream (bromide), or yellow (iodide) precipitates.