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Properties of the Halogens

You might know chlorine as the chemical added to swimming pools, but its pure, natural form looks very different.
Group 7 elements are known as the halogens, and they all have 7 electrons in their outer shell.
They naturally exist as a diatomic molecule (e.g., $F_2$ , $Cl_2$ , , ), meaning two atoms share a pair of electrons in a single covalent bond.

As you go down Group 7, the relative molecular mass increases. This leads to stronger intermolecular forces, meaning higher melting and boiling points. At AQA standard, you must know their physical states and appearances at room temperature:

Fluorine: Pale yellow gas.
Chlorine: Pale green gas (forms a pale green aqueous solution in water).
Bromine: Dark red-brown liquid (forms an orange or yellow-orange aqueous solution).
Iodine: Dark grey/grey-black solid (forms a brown aqueous solution, but sublimates into a purple vapor).

If you need to test for chlorine gas, place damp blue litmus paper in the gas. The paper will turn red because the gas is acidic, and it will then bleach white.

Explaining the Trend in Reactivity

Why does fluorine react aggressively with almost everything, while iodine barely reacts at all?
In Group 7, reactivity decreases as you move down the group from fluorine to astatine.
To achieve a full outer shell and become stable, halogens need to gain one electron to form a negative halide ion with a -1 charge (e.g., $Cl^-$ ).

To secure full marks when explaining this trend, you must describe the causal mechanism in step-by-step order. As you go down the group, the atoms have more electron shells, which increases the atomic radius (the distance from the nucleus to the outer shell). These extra inner shells also increase electron shielding.

The combination of increased distance and increased shielding results in a weaker nuclear attraction between the positive nucleus and an incoming electron. Because this electrostatic attraction is weaker, it becomes harder for the atom to gain the one electron it needs, making the element less reactive. Even though astatine is radioactive and rarely tested in school labs, we can use this trend to predict it is the least reactive element in the group.

Halogen Displacement Reactions

In chemistry, an aggressive element can "steal" a place in a compound, forcing a weaker element out.
This process is called a displacement reaction, where a more reactive halogen displaces a less reactive halogen from an aqueous solution of its metal halide salt.
Metal halide salts (like potassium bromide, $KBr$ ) do NOT have a colour in aqueous solution; they are perfectly colourless.

To perform this reaction, follow these steps: First, add an aqueous solution of a halogen to a colourless metal halide solution. Then, observe the mixture. If the added halogen is higher up Group 7 (more reactive), it displaces the less reactive halide ion from the salt. Finally, the newly formed halogen molecule causes the solution to change colour. If the added halogen is lower in the group (less reactive), there is no reaction and no colour change.

Typical AQA displacement observations include:

Chlorine water + KBr: Solution turns orange (bromine is formed).
Chlorine water + KI: Solution turns brown (iodine is formed).
Bromine water + KI: Solution turns brown (iodine is formed).

For Higher Tier students, an organic solvent like cyclohexane can be added to clearly distinguish the products. Bromine in cyclohexane appears orange/red-brown, whereas iodine in cyclohexane appears purple/violet/lilac.

Redox and Ionic Half-Equations (Higher Tier)

Tracking the movement of electrons reveals the true mechanism behind displacement reactions.
Displacement reactions are redox reactions because both oxidation and reduction happen simultaneously.
The halogen molecule acts as an oxidising agent because it gains electrons (it is reduced), while the halide ion acts as a reducing agent because it loses electrons (it is oxidised).

During these reactions, the metal ion (like potassium, $K^+$ ) acts as a spectator ion. This means it does NOT change its state or its electronic configuration during the reaction, so we can ignore it when writing ionic equations.

Worked Example: Chlorine Displacing Bromine

Step 1: Write the full balanced symbol equation. Remember that halogens are diatomic.

Cl_2(aq) + 2KBr(aq) \rightarrow 2KCl(aq) + Br_2(aq)

Step 2: Write the ionic equation. Remove the potassium spectator ions.

Cl_2(aq) + 2Br^-(aq) \rightarrow 2Cl^-(aq) + Br_2(aq)

Step 3: Split into oxidation and reduction half-equations.

Reduction half-equation: $Cl_2(aq) + 2e^- \rightarrow 2Cl^-(aq)$ (Chlorine gains electrons).

Students often confuse Group 7 with Group 1. In Group 1, reactivity increases down the group because an outer electron is lost more easily, but in Group 7, reactivity decreases because an electron is gained less easily.

In 3-mark 'Explain the trend in reactivity' questions, examiners expect you to explicitly link the increase in atomic radius and shielding to a weaker nuclear attraction, resulting in an electron being gained less easily.

Always use the term 'halide' for the negative ion (e.g., Cl⁻) and 'halogen' for the neutral element (e.g., Cl₂); mixing these up in your descriptions will result in lost marks.

When writing balanced symbol equations for displacement reactions, do not forget that halogens are diatomic, meaning you need two halide ions to form one halogen molecule.

For 4-6 mark explanations on Higher Tier redox questions, clearly state which species gains electrons (the halogen) and which species loses electrons (the halide ion), linking this to their half-equations.

The outermost energy level of an atom that contains electrons, which determines the element's chemical properties.

A molecule consisting of exactly two atoms chemically bonded together, such as Cl2.

A mixture where a substance is dissolved in water.

A halogen atom that has gained one electron to form a negative ion with a -1 charge.

The distance from the center of the nucleus to the outermost shell of electrons.

The process by which inner electron shells reduce the effective nuclear charge felt by outer electrons or incoming electrons.

The electrostatic force of attraction between the positive protons in the nucleus and the negative electrons.

A chemical reaction in which a more reactive element takes the place of a less reactive element in a compound.

A type of chemical reaction where oxidation and reduction occur simultaneously.

The loss of electrons from a substance during a chemical reaction.

The gain of electrons by a substance during a chemical reaction.

An ion that does not change state or electronic configuration during a reaction and is omitted from ionic equations.