Group 7 — The Halogens

This lesson covers the Group 7 elements (the halogens) as required by the Edexcel GCSE Chemistry specification (1CH0), Topic 6: Groups in the Periodic Table. You need to know the physical properties of the halogens, their trend in reactivity, and how they react with metals. You must also be able to explain the reactivity trend using electron configuration.

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The Group 7 Elements

Group 7 is found on the right-hand side of the periodic table, one column to the left of the noble gases. The halogens are non-metals that exist as diatomic molecules (molecules made of two atoms bonded together).

Element	Symbol	Molecular Formula	Atomic Number	Electron Configuration
Fluorine	F	F₂	9	2, 7
Chlorine	Cl	Cl₂	17	2, 8, 7
Bromine	Br	Br₂	35	2, 8, 18, 7
Iodine	I	I₂	53	2, 8, 18, 18, 7
Astatine	At	At₂	85	— (radioactive, very rare)

Exam Tip: Remember that all halogens have seven electrons in their outer shell. This is why they are in Group 7 and why they all behave in a similar way chemically.

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Physical Properties of the Halogens

The halogens show a clear trend in physical properties as you go down the group.

Element	Colour	State at Room Temperature (25 °C)	Melting Point (°C)	Boiling Point (°C)
Fluorine	Pale yellow	Gas	−220	−188
Chlorine	Yellow-green	Gas	−101	−34
Bromine	Red-brown	Liquid	−7	59
Iodine	Dark grey (purple vapour)	Solid	114	184

Trends Down Group 7

As you go down Group 7:

Property	Trend	Explanation
Colour	Gets darker	From pale yellow to dark grey/black
State	Gas → Liquid → Solid	Due to increasing intermolecular forces
Melting point	Increases	Larger molecules have stronger intermolecular forces (London dispersion forces), requiring more energy to overcome
Boiling point	Increases	Same reason as melting point
Atomic radius	Increases	Each element has one more electron shell
Relative molecular mass	Increases	More protons and neutrons in the nucleus

Exam Tip: When explaining the increase in melting and boiling points down the group, refer to the increase in intermolecular forces between the molecules. The molecules become larger, so there are stronger London (dispersion) forces between them — more energy is needed to overcome these forces.

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Reactivity of the Halogens

The halogens are reactive non-metals. Unlike Group 1, reactivity decreases down Group 7.

Fluorine is the most reactive halogen. Iodine is the least reactive of the commonly studied halogens.

Explaining the Trend in Reactivity

All halogens have seven electrons in their outer shell. When they react, they need to gain one electron to form a negative ion (X⁻) with a stable noble gas electron configuration.

As you go down the group:

1. The atoms get larger (more electron shells).
2. The outer shell is further from the nucleus.
3. There is more shielding from inner electron shells.
4. The attraction between the nucleus and an incoming electron is weaker.
5. It is harder to gain an extra electron.
6. Therefore the element is less reactive.

Halogen	Shells	Distance of Outer Shell from Nucleus	Ease of Gaining an Electron	Reactivity
F	2	Closest	Easiest	Most reactive
Cl	3	Further	Less easy	Reactive
Br	4	Even further	Harder	Less reactive
I	5	Furthest	Hardest	Least reactive

Exam Tip: The key difference between Group 1 and Group 7 trends is: Group 1 metals lose electrons (reactivity increases down the group), while Group 7 non-metals gain electrons (reactivity decreases down the group). Make sure you explain the correct process for each group.

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Reactions of Halogens with Metals

Halogens react with metals to form metal halides — ionic compounds in which the halogen has gained one electron to form a halide ion (e.g. Cl⁻, Br⁻, I⁻).

The general word equation is:

metal + halogen → metal halide

Examples

Sodium and chlorine:

Word equation: sodium + chlorine → sodium chloride

Symbol equation: 2Na(s) + Cl₂(g) → 2NaCl(s)

Observations: Sodium burns with a bright yellow flame in chlorine gas. A white solid (sodium chloride) is formed.

Iron and chlorine:

Word equation: iron + chlorine → iron(III) chloride

Symbol equation: 2Fe(s) + 3Cl₂(g) → 2FeCl₃(s)

Observations: The heated iron wool glows brightly in chlorine gas. A brown solid (iron(III) chloride) is formed.

Iron and bromine:

Word equation: iron + bromine → iron(III) bromide

Symbol equation: 2Fe(s) + 3Br₂(g) → 2FeBr₃(s)

Aluminium and iodine:

Word equation: aluminium + iodine → aluminium iodide

Symbol equation: 2Al(s) + 3I₂(s) → 2AlI₃(s)

Observations: A vigorous, exothermic reaction producing clouds of purple iodine vapour and a white/yellow solid product.

Exam Tip: When writing symbol equations for halogen reactions, remember that halogens are diatomic (Cl₂, Br₂, I₂) — never write just Cl or Br on their own in an equation.

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Naming Metal Halides

The name of the salt formed depends on the halogen used:

Halogen	Halide Ion	Example Salt
Fluorine	Fluoride (F⁻)	Sodium fluoride (NaF)
Chlorine	Chloride (Cl⁻)	Sodium chloride (NaCl)
Bromine	Bromide (Br⁻)	Sodium bromide (NaBr)
Iodine	Iodide (I⁻)	Sodium iodide (NaI)

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Halogens as Diatomic Molecules

Halogens exist as diatomic molecules (two atoms bonded together by a single covalent bond). Each atom shares one electron with the other to achieve a full outer shell.

For example, in Cl₂, each chlorine atom has 7 outer electrons. They share one pair of electrons (one from each atom), giving each atom effectively 8 outer electrons — a stable arrangement.

This covalent bonding within the molecule is strong, but the intermolecular forces between the Cl₂ molecules are relatively weak — which is why chlorine is a gas at room temperature.

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Comparing Reactivity with Experiments

The relative reactivity of the halogens can be compared experimentally by:

1. Reacting each halogen with iron wool and comparing the vigour of the reaction.
2. Observing how quickly each halogen bleaches damp litmus paper (a more reactive halogen bleaches it faster).
3. Displacement reactions — adding halogen solutions to metal halide solutions (covered in the next lesson).

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Summary