Who are the halogens?
The halogens sit in group 17 (old name group 7) of the periodic table: fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At). The word halogen means salt maker, because they join with metals to make salts like sodium chloride.
Each atom has 7 electrons in its outer shell. It needs just 1 more to have a full shell. So a halogen atom easily gains one electron and becomes a halide ion with charge −1 (F⁻, Cl⁻, Br⁻, I⁻).
The atoms do not stay alone. Two atoms share a pair of electrons and make a diatomic molecule: F₂, Cl₂, Br₂, I₂.
Physical properties and trends
At room temperature (about 25 °C):
- F₂: pale yellow gas
- Cl₂: yellow-green gas
- Br₂: red-brown liquid (gives orange vapour)
- I₂: grey-black shiny solid (gives purple vapour when warmed)
Down the group:
- Atoms get bigger (more shells).
- Molecules get bigger, so the forces between them get stronger. Melting and boiling points go up.
- Colour gets darker.
- Electronegativity (how strongly an atom pulls shared electrons) goes down. Fluorine is the most electronegative element of all.
Chemical properties and reactivity
Reactivity goes down the group. Why? A halogen reacts by pulling in one electron. In a small atom (F) the outer shell is close to the positive nucleus, so the pull is strong. In a big atom (I) the outer shell is far away and shielded by inner shells, so the pull is weak.
With metals: halogens make ionic salts. 2Na + Cl₂ → 2NaCl. Iron wool burns in chlorine to make iron(III) chloride.
With hydrogen: they make hydrogen halides. H₂ + Cl₂ → 2HCl. HCl gas dissolves in water to make hydrochloric acid.
Displacement: a more reactive halogen takes the place of a less reactive one in a salt. Cl₂ + 2KBr → 2KCl + Br₂. Here chlorine gains electrons, so it is an oxidising agent. Oxidising power falls down the group.
Testing for halide ions and uses
Test: add dilute nitric acid to the solution, then a few drops of silver nitrate. A precipitate (solid) forms:
- chloride: white (AgCl)
- bromide: cream (AgBr)
- iodide: yellow (AgI)
Uses: chlorine kills germs in drinking water and pools and is used to make bleach and PVC plastic. Fluoride in toothpaste and some water makes teeth stronger. Iodine is used as an antiseptic, and iodised salt prevents goitre. Bromine compounds were used in old film photography.
Safety: halogen gases are toxic. Chlorine was used as a poison gas in World War I. Always work in a fume cupboard.
Key formulas and definitions
- Halogen atom + 1 e⁻ → halide ion: Cl + e⁻ → Cl⁻
- Metal + halogen → salt: 2Na + Cl₂ → 2NaCl
- Hydrogen + halogen → hydrogen halide: H₂ + Cl₂ → 2HCl
- Displacement: Cl₂ + 2KBr → 2KCl + Br₂
- Test: Ag⁺ + Cl⁻ → AgCl (white)
- Reactivity: F₂ > Cl₂ > Br₂ > I₂
Worked examples
1. Chlorine water is added to potassium iodide solution. What do you see? Write the equation.
The colourless solution turns brown because iodine is formed. Chlorine is more reactive than iodine. Cl₂ + 2KI → 2KCl + I₂.
2. Astatine is below iodine. Predict its state at room temperature and its reactivity compared with iodine.
Following the trend, astatine is a dark solid with a higher melting point than iodine, and it is less reactive than iodine.
3. A solution gives a cream precipitate with acidified silver nitrate. Which ion is present?
Bromide ion (Br⁻). The cream solid is silver bromide, AgBr.
Common mistakes
- Writing chlorine as Cl in an equation. As an element it is a molecule, Cl₂.
- Thinking reactivity goes up down the group, as with group 1 metals. For halogens it goes down.
- Mixing up chlorine (the element, a toxic gas) and chloride (the ion in salt, safe to eat in NaCl).
- Forgetting to add nitric acid first in the silver nitrate test. It removes other ions like carbonate that also make a precipitate.