Group 7: The Halogens

This lesson covers Group 7 of the periodic table — the halogens — as required by AQA GCSE Chemistry specification (5.1.2). The halogens are reactive non-metals with distinctive properties. Understanding their physical properties, reactions, displacement reactions, and the trend in reactivity down the group is essential for GCSE Chemistry.

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The Halogens

Group 7 contains the following elements:

Element	Symbol	Atomic Number	Electronic Configuration	State at Room Temp	Colour
Fluorine	F	9	2, 7	Gas	Pale yellow
Chlorine	Cl	17	2, 8, 7	Gas	Green-yellow
Bromine	Br	35	2, 8, 18, 7	Liquid	Red-brown
Iodine	I	53	2, 8, 18, 18, 7	Solid	Dark grey/purple vapour
Astatine	At	85	(radioactive, very rare)	Solid	Dark (predicted)

Key Features of Halogens

• All halogens have 7 electrons in their outer shell.
• They exist as diatomic molecules — F2, Cl2, Br2, I2.
• They are non-metals.
• They react by gaining 1 electron to form -1 ions (called halide ions: F-, Cl-, Br-, I-).

Exam Tip: Remember that halogens are diatomic molecules (they exist in pairs). In equations, always write Cl2, Br2, I2 — never just Cl, Br, or I. This is one of the most common mistakes students make when writing equations.

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Physical Properties of Halogens

Trends Down Group 7

Property	Trend Down the Group	Explanation
Melting and boiling points	Increase	Intermolecular forces (van der Waals forces) get stronger as the molecules get larger with more electrons
Colour	Gets darker	F2 is pale yellow; Cl2 is green-yellow; Br2 is red-brown; I2 is dark grey
State at room temperature	Gas → Liquid → Solid	Due to increasing boiling points
Density	Increases	Molecules are heavier

This explains the change in state: fluorine and chlorine are gases because they have weak intermolecular forces and low boiling points; bromine is a liquid; iodine is a solid with relatively stronger intermolecular forces.

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Reactivity of Halogens

Unlike Group 1 metals (where reactivity increases down the group), halogen reactivity decreases down the group.

Element	Reactivity	Explanation
Fluorine	Most reactive	Smallest atom; outer shell closest to nucleus; strongest attraction for incoming electron
Chlorine	Very reactive	Small atom; strong attraction for incoming electron
Bromine	Moderately reactive	Larger atom; outer shell further from nucleus; weaker attraction
Iodine	Least reactive	Largest common halogen; outer shell furthest from nucleus; weakest attraction

Why Does Reactivity Decrease?

graph TD
    A["Going Down Group 7"] --> B["More electron shells"]
    B --> C["Outer shell is further from the nucleus"]
    B --> D["Greater electron shielding"]
    C --> E["Weaker attraction for incoming electron"]
    D --> E
    E --> F["Harder to gain an extra electron"]
    F --> G["Less reactive"]

    style A fill:#2c3e50,color:#fff
    style E fill:#e74c3c,color:#fff
    style G fill:#e67e22,color:#fff

All halogens react by gaining 1 electron to achieve a full outer shell (noble gas configuration). Going down the group:

• Atoms have more electron shells, so the outer shell is further from the nucleus.
• There is greater electron shielding from inner shells.
• The attractive force between the nucleus and an incoming electron is weaker.
• It is therefore harder for the atom to gain an electron.
• This makes the element less reactive.

Exam Tip: The reactivity trend for halogens is the OPPOSITE of Group 1 metals. Group 1: reactivity increases down (easier to LOSE electrons). Group 7: reactivity decreases down (harder to GAIN electrons). Make sure you explain the correct mechanism for each group.

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Reactions of Halogens

Reaction with Metals

Halogens react with metals to form metal halides (ionic compounds). The halogen atoms gain 1 electron to form halide ions with a -1 charge.

Reaction	Word Equation	Symbol Equation
Sodium + Chlorine	Sodium + Chlorine → Sodium chloride	2Na + Cl2 → 2NaCl
Iron + Chlorine	Iron + Chlorine → Iron(III) chloride	2Fe + 3Cl2 → 2FeCl3
Iron + Bromine	Iron + Bromine → Iron(III) bromide	2Fe + 3Br2 → 2FeBr3

Reaction with Hydrogen

Halogens react with hydrogen to form hydrogen halides, which dissolve in water to form acidic solutions.

Reaction	Product	Acid in Water
H2 + F2 → 2HF	Hydrogen fluoride	Hydrofluoric acid
H2 + Cl2 → 2HCl	Hydrogen chloride	Hydrochloric acid
H2 + Br2 → 2HBr	Hydrogen bromide	Hydrobromic acid
H2 + I2 → 2HI	Hydrogen iodide	Hydroiodic acid

The vigour of these reactions decreases down the group (fluorine reacts explosively; iodine reacts slowly and requires a catalyst or high temperature).

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Halogen Displacement Reactions

A more reactive halogen can displace (push out) a less reactive halogen from a solution of its salt. This is used to demonstrate the reactivity trend.

The Rules

• A halogen higher in Group 7 (more reactive) will displace a halogen lower in the group (less reactive).
• A halogen cannot displace a more reactive halogen.

Displacement Reaction Results

Halogen Added	Potassium Chloride (KCl)	Potassium Bromide (KBr)	Potassium Iodide (KI)
Chlorine (Cl2)	No reaction (same element)	Displacement occurs — solution turns orange (bromine produced)	Displacement occurs — solution turns brown (iodine produced)
Bromine (Br2)	No reaction (Br is less reactive than Cl)	No reaction (same element)	Displacement occurs — solution turns brown (iodine produced)
Iodine (I2)	No reaction	No reaction	No reaction (same element)

Example equations:

• Cl2 + 2KBr → 2KCl + Br2 (chlorine displaces bromine)
• Cl2 + 2KI → 2KCl + I2 (chlorine displaces iodine)
• Br2 + 2KI → 2KBr + I2 (bromine displaces iodine)

Exam Tip: Displacement reactions prove the reactivity order. If a halogen CAN displace another, it is MORE reactive. Remember: chlorine displaces both bromine and iodine; bromine displaces only iodine; iodine displaces neither. Colour changes are key observations: Br2 is orange, I2 is brown.

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Halide Ions

When halogens gain an electron, they form halide ions: