What you'll learn
Group 7, the halogens, are a family of reactive non-metals with clear trends down the group. In this guide you will learn the properties of the halogens, how they exist as diatomic molecules, the trends in their physical properties and reactivity, how they react with metals and hydrogen, how displacement reactions show their order of reactivity, and how electronic structure explains these trends. This topic shows how the periodic table predicts patterns of behaviour.
Key terms and definitions
Group 7 / halogens — the non-metals in the seventh column (fluorine, chlorine, bromine, iodine).
Diatomic molecule — a molecule made of two atoms (e.g. Cl₂); halogens exist as diatomic molecules.
Halide ion — the 1− ion formed when a halogen gains an electron (e.g. Cl⁻).
Displacement reaction — a more reactive halogen taking the place of a less reactive one in a compound.
Reactivity — how readily an element reacts; for halogens it decreases down the group.
Shielding — inner electron shells reducing the nucleus's pull on the outer shell.
Core concepts
Properties of the halogens
The halogens are reactive non-metals that exist as diatomic molecules (F₂, Cl₂, Br₂, I₂). They have coloured vapours and show a clear trend in state at room temperature: fluorine is a pale yellow gas, chlorine a green gas, bromine a red-brown liquid, and iodine a grey-black solid (that sublimes to a purple vapour). Each halogen has seven electrons in its outer shell.
Trends in physical properties
Going down Group 7:
- Melting and boiling points increase (so the state changes from gas to liquid to solid), because the molecules get larger and intermolecular forces increase.
- The colour gets darker and the elements become denser.
Reactions with metals and hydrogen
Halogens react with metals to form metal halides (ionic salts), gaining one electron to form 1− halide ions (e.g. sodium + chlorine → sodium chloride). They react with hydrogen to form hydrogen halides (e.g. hydrogen chloride, HCl), which dissolve in water to form acidic solutions.
Trend in reactivity down the group
Reactivity decreases down Group 7 — the opposite of Group 1. Halogens react by gaining one electron to complete their outer shell. Down the group, the outer shell is further from the nucleus and has more shielding, so the nucleus attracts an incoming electron less strongly. This makes it harder to gain an electron, so reactivity decreases. Fluorine (at the top) is the most reactive halogen; iodine is the least.
Displacement reactions
A more reactive halogen displaces a less reactive halogen from a solution of its salt. For example, chlorine added to potassium bromide solution displaces bromine: chlorine + potassium bromide → potassium chloride + bromine (the solution turns orange as bromine forms). These reactions confirm the order of reactivity: chlorine displaces bromine and iodine; bromine displaces iodine; but iodine cannot displace chlorine or bromine.
Worked examples
Example 1: Predicting a displacement
Will chlorine displace iodine from potassium iodide?
Yes. Chlorine is more reactive than iodine, so it displaces it: chlorine + potassium iodide → potassium chloride + iodine (the solution darkens as iodine forms).
Example 2: Explaining the trend
Why is chlorine more reactive than bromine?
Chlorine's outer shell is closer to the nucleus with less shielding, so it attracts an incoming electron more strongly and gains an electron more easily, making it more reactive.
Example 3: Forming an ion
What happens to a chlorine atom when it reacts with a metal?
It gains one electron to complete its outer shell, forming a chloride ion, Cl⁻ (a 1− halide ion).
Common mistakes and how to avoid them
Reversing the trend. Reactivity decreases down Group 7 (the opposite of Group 1).
Forgetting halogens are diatomic. Write them as Cl₂, Br₂, etc.
Explaining the trend by electron loss. Halogens gain an electron; reactivity falls because gaining gets harder down the group.
Predicting impossible displacements. A less reactive halogen cannot displace a more reactive one (iodine cannot displace chlorine).
Wrong ion charge. Halogens form 1− ions (halide ions).
Exam technique for Group 7
Describe properties and states of the halogens, noting the colour and state trend.
State the physical trends (melting/boiling points increase down the group) and explain by molecule size.
Explain the reactivity trend using distance, shielding and ease of gaining an electron.
Predict displacement reactions and describe the colour changes.
Show halide ion formation by gaining one electron.
Quick revision summary
Group 7, the halogens (fluorine, chlorine, bromine, iodine), are reactive non-metals that exist as diatomic molecules (Cl₂, Br₂…) with seven outer electrons. Down the group the state changes from gas (F₂, Cl₂) to liquid (Br₂) to solid (I₂) as melting and boiling points increase (larger molecules, stronger intermolecular forces), and the colours darken. Halogens react with metals to form ionic metal halides and with hydrogen to form hydrogen halides (acidic in water), gaining one electron to form 1− halide ions. Reactivity decreases down the group (opposite to Group 1): because the outer shell is further from the nucleus with more shielding, the nucleus attracts an incoming electron less strongly, so gaining an electron gets harder. A more reactive halogen displaces a less reactive one from its salt (chlorine displaces bromine and iodine), confirming the order, with characteristic colour changes. State the properties and trends, explain reactivity by the ease of gaining an electron, and predict displacement reactions correctly.