Group 7: The Halogens

If Group 1 is the group of the most reactive metals, Group 7 is the group of some of the most reactive non-metals. The halogens — fluorine, chlorine, bromine and iodine — are coloured, poisonous non-metals that exist as pairs of atoms bonded together, and they show a set of clear, learnable trends. In this lesson, part of Topic C4 of OCR Gateway Combined Science A, we look at their properties, at how their reactivity decreases down the group (the opposite direction to Group 1), and — most importantly for predicting reactions — at halogen displacement, where a more reactive halogen pushes a less reactive one out of its salt.

By the end of this lesson you should be able to describe the properties and trends of the halogens, explain the reactivity trend down Group 7 in terms of gaining an electron, predict and write equations for halogen displacement reactions, and use the trend to predict how an unfamiliar halogen will behave.

This lesson builds AO1 recall of the properties and trends of the halogens, AO2 application when you write balanced equations for halogen displacement reactions, and AO3 analysis when you explain the reactivity trend by electron gain and predict how an unfamiliar halogen will behave.

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Properties of the Halogens

The Group 7 elements have their own characteristic set of properties:

• they are non-metals with seven electrons in the outer shell;
• they exist as diatomic molecules — pairs of atoms joined by a covalent bond, written $\text{F}_2$F2​, $\text{Cl}_2$Cl2​, $\text{Br}_2$Br2​ and $\text{I}_2$I2​;
• they are coloured, and the colour gets darker as you go down the group;
• they react by gaining one electron to complete the outer shell, forming a −1 ion (a halide ion such as $\text{Cl}^-$Cl−), or by sharing electrons in covalent bonds.

Their appearance and physical state at room temperature change steadily down the group, which is itself one of the trends you should know:

Halogen	Formula	Colour and state at room temperature
Fluorine	$\text{F}_2$F2​	Pale yellow gas
Chlorine	$\text{Cl}_2$Cl2​	Green (yellow-green) gas
Bromine	$\text{Br}_2$Br2​	Red-brown liquid
Iodine	$\text{I}_2$I2​	Grey-black solid (purple vapour)

Exam Tip: The halogens are diatomic — always write $\text{Cl}_2$Cl2​, $\text{Br}_2$Br2​ and $\text{I}_2$I2​, never a lone $\text{Cl}$Cl. A common misconception is to write halogens as single atoms like the noble gases; only the noble gases are monatomic.

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Trends Down Group 7

Two trends run down Group 7, and they point in opposite directions:

• Melting and boiling points increase down the group. This is why the physical state shifts from gas (chlorine) to liquid (bromine) to solid (iodine): larger molecules have stronger forces between them, so more energy is needed to separate them.
• Reactivity decreases down the group — the reverse of Group 1.

Alongside these, the elements get darker and denser further down. The reactivity trend is the one that drives most predictions, so it is worth pinning down carefully.

Explaining the reactivity trend. A halogen reacts by gaining one electron to fill its outer shell and form a $-1$−1 ion. Going down the group, the outer shell is further from the nucleus and more shielded by inner shells of electrons, so the atom attracts an incoming electron less strongly. Because gaining that electron becomes harder, the halogens get less reactive down the group — fluorine and chlorine are the most reactive, iodine the least. It is worth pausing on how neat this is: a single structural change — an extra shell each step down — accounts for the whole trend, without your having to memorise each element separately.

Exam Tip: Group 1 and Group 7 trend in opposite directions but for the same underlying reason — the outer shell moving further from the nucleus and gaining more shielding. For Group 1 that makes losing an electron easier (more reactive); for Group 7 it makes gaining an electron harder (less reactive).

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Reaction of Halogens with Metals

Before looking at displacement, it helps to see the halogens reacting with metals, because this is the reaction that makes the halide salts in the first place. A halogen reacts with a metal to form a metal halide — an ionic compound in which the metal has formed a positive ion and each halogen atom has gained an electron to become a $-1$−1 halide ion. The general pattern is:

$\text{metal} + \text{halogen} \rightarrow \text{metal halide}$metal+halogen→metal halide

For example, sodium burns in chlorine to form sodium chloride (ordinary table salt):

$2\text{Na} + \text{Cl}_2 \rightarrow 2\text{NaCl}$2Na+Cl2​→2NaCl

and iron reacts with chlorine to form iron(III) chloride:

$2\text{Fe} + 3\text{Cl}_2 \rightarrow 2\text{FeCl}_3$2Fe+3Cl2​→2FeCl3​

Because the halogens differ in reactivity, the same metal reacts more vigorously with a more reactive halogen: iron reacts faster and more fiercely with chlorine than with bromine, and more fiercely with bromine than with iodine. So a metal-plus-halogen reaction is another way to see the reactivity order — chlorine, being highest of the common three, reacts most vigorously. The products of these reactions, the metal halides, dissolve in water to give the halide solutions used in the displacement reactions that follow.

Exam Tip: A halide is the compound or ion formed when a halogen gains an electron — chloride, bromide, iodide. Keep the two words straight: the element is the halogen (chlorine), the ion or salt is the halide (chloride).

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Halogen Displacement Reactions

The reactivity trend leads directly to a set of reactions you can predict: a more reactive halogen displaces a less reactive halogen from a solution of its salt (a halide). This is just like metal displacement, but with halogens.

For example, if you bubble chlorine (more reactive) into a solution of potassium bromide, the chlorine displaces the bromine:

$\text{chlorine} + \text{potassium bromide} \rightarrow \text{potassium chloride} + \text{bromine}$chlorine+potassium bromide→potassium chloride+bromine $\text{Cl}_2 + 2\text{KBr} \rightarrow 2\text{KCl} + \text{Br}_2$Cl2​+2KBr→2KCl+Br2​

You can see the reaction happen because the solution takes on the orange colour of the displaced bromine. Chlorine will also displace iodine from potassium iodide (turning the solution brown as iodine forms), and bromine will displace iodine. But the reverse never happens: iodine cannot displace chlorine or bromine, because iodine is the least reactive of the three.

The grid below summarises which displacements occur. Read a row as "the halogen added" and a column as "the halide already in solution"; a tick means the reaction happens.

Halogen added ↓ / Halide in solution →	Chloride	Bromide	Iodide
Chlorine	—	✓ displaces Br	✓ displaces I
Bromine	✗	—	✓ displaces I
Iodine	✗	✗	—

Notice the same diagonal pattern you saw for metals: each halogen can only displace those below it in the group. These displacement results are, in fact, a way of confirming the reactivity order of the halogens — the element that displaces the others is the most reactive.

Exam Tip: For halogen displacement, ask "is the halogen I am adding more reactive (higher in Group 7) than the halide in solution?" If yes, it displaces it and the colour of the newly released halogen appears; if no, there is no reaction.

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Higher Tier: Displacement as Redox

Higher tier only. Halogen displacement is a redox reaction, just like metal displacement. Take chlorine displacing bromide, written as an ionic equation (the potassium ions are spectators):

$\text{Cl}_2 + 2\text{Br}^- \rightarrow 2\text{Cl}^- + \text{Br}_2$Cl2​+2Br−→2Cl−+Br2​

Following the electrons:

• Each chlorine atom gains an electron to become a chloride ion, $\text{Cl}^-$Cl− — chlorine is reduced (gain of electrons).
• Each bromide ion loses its extra electron to become a bromine atom in $\text{Br}_2$Br2​ — bromide is oxidised (loss of electrons).