Halogens

Fluorine, chlorine, bromine, iodine and astatine sit together in Group 7 of the Periodic Table, and the group has a name of its own: the halogens. At first sight the members look almost nothing like one another — one is a pale gas, one a dense coloured liquid, one a shiny grey solid — and part of this chapter is learning to see the family underneath the differences.

Hold on to one idea from the start: the outer shell. Every Group 7 atom carries 7 electrons in its outer shell, and the outer shell is where all chemistry happens. That single shared fact is why these five elements react in similar ways, form similar compounds and show neat, predictable trends as you descend the group — exactly the pattern you already saw for Group 1 in Chapter 11 (u2:c11), now played out on the non-metal side of the table.

The name halogen means "salt-forming", and it is earned: react any halogen with a metal and you get a salt. Calcium fluoride in tap water chemistry, sodium chloride on your chips, silver bromide in photographic film, potassium iodide in the lab — all of them are halogen salts, and every one of them contains the same unit: the halide ion, X⁻ (where X stands for any halogen). One letter, one charge, one family.

Two pieces of machinery before the detail. First, the halogens are non-metallic and exist as diatomic molecules — F₂, Cl₂, Br₂, I₂ — two atoms bonded into one molecule, never lone atoms in the element. Second, this chapter's evidence engine is the displacement reaction: pitting one halogen against another's salt solution and watching which one wins. Those reactions are the experimental proof that reactivity falls down the group, and they are where the chapter earns its exam marks.

By the end of this chapter you will be able to state the colours and physical states of the halogens at room temperature and explain the physical-property trends; use the group trends to predict the properties of halogens you have never met; carry out and interpret halogen displacement reactions, including their ionic equations; and — for Paper 2 chemistry-only content — explain the reactivity trend in terms of electronic configurations. Each of those builds on the last, so the order above is the order to learn in.

The atoms of fluorine, chlorine, bromine, iodine and astatine all have 7 electrons in their outer shell, and that is the whole reason they are grouped together. The number of outer-shell electrons determines how an atom reacts, so every Group 7 element reacts in a similar way — and the reactions all carry the family signature: forming compounds that contain the X⁻ ion, the halogen having gained one electron.

The elements themselves are non-metallic and diatomic — their molecules are pairs of atoms, F₂, Cl₂, Br₂, I₂. Memorise the physical profile the family walks through as the molecules get heavier:

Read the table as a trend, not a list. Down the group the melting points and boiling points increase — gas → liquid → solid — and the colour becomes darker. Both trends have one cause. The halogens are all covalent molecular substances, and as the relative molecular mass increases down the group the intermolecular-forces of attraction between molecules become stronger, so more energy is needed to overcome them. That is what melting and boiling actually are: separating molecules from each other, not breaking the covalent bonds inside them. No covalent bond is broken when iodine melts — the I–I bonds within each I₂ molecule survive intact.

Being non-metals, the halogens are poor conductors of heat and electricity, in both their solid and liquid states — there are no delocalised electrons to carry charge, unlike the metallic bonding of Chapter 9 (u1:c09).

One member needs special handling: astatine. It is radioactive, formed during the radioactive decay of other elements such as uranium and thorium, and most of its isotopes are so unstable that a sample decays away almost as fast as it forms — which is why no one has ever assembled a visible quantity of it. Everything you can say about astatine in an exam is a prediction from the group trend, which is exactly what the closing section of this chapter trains you to do.

Before any halogen reaction, the hazards — because this family is the most dangerous set of elements you will meet. Fluorine is so dangerously reactive that you would never expect to find it in a school lab at all; the reactions it drives are violent beyond anything a teaching room is equipped to control. The school-taught members — chlorine, bromine, iodine — are manageable, but only inside a discipline: all the halogens have extremely poisonous vapours, so every experiment with them is carried out in a fume-cupboard, which pulls the toxic vapours away from the air you breathe.

Liquid bromine adds a second hazard: it is very corrosive, and great care is taken to keep it off the skin. A drop of bromine on bare skin burns chemically — that is what corrosive means — so bromine handling is gloves, goggles and a fume cupboard, every time. Chlorine is a gas with a choking, poisonous vapour; iodine is the gentlest of the three but still stains and irritates, and its vapour is not something to breathe by choice.

The safety picture is itself part of the chemistry. The same high reactivity that makes halogens superb salt-formers makes their vapours destructive to living tissue, and reactivity (like the boiling points) tracks the group trend — fluorine at the top is the worst of the family, which is precisely why it has no place in a school lab.

Two reaction families carry the halogens' chemistry, and both are formula-patterns you can write for any member of the group. Learn the pattern once, and every individual equation is a substitution exercise.

Reaction 1: with hydrogen. Every halogen combines with hydrogen to form a hydrogen-halide, HX — hydrogen fluoride, hydrogen chloride, hydrogen bromide, hydrogen iodide:

H₂(g) + Br₂(g) → 2HBr(g)

The hydrogen halides are all acidic, poisonous gases. In common with every compound formed between two non-metals, they are covalent molecular substances — and, like all acids, they dissolve in water to give acidic solutions. Hydrochloric acid is hydrogen chloride dissolved in water; the same gas-to-acid story gives hydrofluoric, hydrobromic and hydroiodic acids. The gas and the solution are different substances with different structures (molecules in the gas, ions in the solution) — the distinction the previous section's key point pinned down.

Reaction 2: with metals. React a halogen with a metal and you form an ionic salt. Sodium burning in chlorine is the demonstration to picture: it burns with its typical yellow flame and the product is white, solid sodium chloride,

2Na(s) + Cl₂(g) → 2NaCl(s)

an ionic solid built from Na⁺ and Cl⁻. The pattern generalises across the Group 1 and Group 2 metals: when halogens react with these metals, the metal atoms lose electrons and the halogen atoms gain them, so the products are ionic — M⁺ and X⁻ packed into a lattice. With non-metals, by contrast, halogens form covalent molecular compounds (the hydrogen halides are the standing example). Metal + halogen = ionic; non-metal + halogen = molecular. That single sentence sorts every halogen product on the course.

One caution about scope. Most of the group chemistry you meet — the reactions above, the displacement work coming next — is taught through the three commonly used halogens, chlorine, bromine and iodine. The patterns do not stop there: the trend runs on through the rest of the group, which is what lets you make confident statements about fluorine above them and astatine below them. Fluorine's extra violence is real but unprintable in a school context; astatine's behaviour is invisible but predictable. The three lab halogens are the family's public face.

Put a more reactive halogen into a solution of a less reactive halogen's salt and the newcomer takes over: it pushes the original halogen out of solution and takes its place. These are displacement-reactions, and they are the experimental engine of this chapter — a way of ranking the halogens by reactivity using nothing more than colour changes.

Demo 1: chlorine water + potassium bromide solution. Add orange-free — colourless — potassium bromide solution to chlorine water and the mixture turns orange: bromine has been formed.

Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq)

Chlorine is more reactive than bromine, and it has displaced the bromine from solution. The way to think about it: a more reactive element has a greater tendency to react and form a compound; a less reactive element is more likely to fall back to being the free, unreacted element. Chlorine wins the compound-forming contest, so bromine is pushed back out as the element.

Demo 2: chlorine water + potassium iodide solution. The same experiment with potassium iodide solution gives a brown solution (orange if very dilute) — iodine has been displaced:

Cl₂(aq) + 2KI(aq) → 2KCl(aq) + I₂(aq)

And the ladder keeps its order one rung down: bromine water added to potassium iodide solution also gives brown iodine,

Br₂(aq) + 2KI(aq) → 2KBr(aq) + I₂(aq)

because bromine, though weaker than chlorine, is still stronger than iodine. Reactivity falls down the group, and every one of these colour changes is a measurement of that fall.

Colour is your observation data in these experiments, so the shades need to be exact. Pure liquid bromine is red-brown — but bromine in solution is orange, and a very dilute solution can look yellow. Iodine in solution is brown (yellow-brown when dilute). Learn the in-solution colours, not the neat-element colours, because exams show you solutions.

The demos in the previous section generalise into a rule, and the rule has three parts. First, the rule itself: a more reactive halogen displaces a less reactive halogen from its salt solution — and the reverse never happens. Test the impossible directions and nothing occurs:

• bromine + potassium chloride solution: no reaction — the less reactive bromine cannot displace the more reactive chlorine;
• iodine + potassium chloride solution: no reaction;
• iodine + potassium bromide solution: no reaction.

Every "no visible change" here is as informative as a colour change: the displacement ladder Cl₂ > Br₂ > I₂ is confirmed by the reactions that fail as much as by the ones that work.

Second, what the metal ion is doing: nothing. Remember that when an ionic compound dissolves in water, it splits into its ions. So potassium bromide solution is really a soup of K⁺(aq) and Br⁻(aq) ions, and when chlorine reacts with it, the chlorine attacks the bromide ion — the potassium ion never changes. It would make no difference if we used sodium bromide or magnesium bromide instead: the reaction happens between the chlorine molecule and the bromide ion, and the metal ion is a spectator-ion. Write the chlorine + potassium iodide reaction showing every charge and you can see it:

2K⁺(aq) + 2I⁻(aq) + Cl₂(aq) → 2K⁺(aq) + 2Cl⁻(aq) + I₂(aq)

The K⁺ ions stand identical on both sides — untouched. Delete them and you get the equation that matters, the ionic-equation:

Cl₂(aq) + 2I⁻(aq) → 2Cl⁻(aq) + I₂(aq)

Same surgery on the bromide reaction gives:

Cl₂(aq) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(aq)

and for bromine against iodide:

Br₂(aq) + 2I⁻(aq) → 2Br⁻(aq) + I₂(aq)

Third, these are redox reactions. Remind yourself of the two definitions: oxidation is loss of electrons, reduction is gain. In the chlorine + bromide reaction the chlorine takes electrons away from the bromide ions — so it oxidises them, while the Cl₂ molecules gain electrons and are reduced to Cl⁻. The electron-thief is called an oxidising-agent: it oxidises something else by taking its electrons, and because it accepts those electrons it is itself reduced. The robbed party — the bromide ion, which gave electrons away and was oxidised — played the matching role of reducing-agent. Every halogen displacement reaction is a redox-reaction with the halogen molecule as oxidising agent and the halide ion as reducing agent, and the oxidising strength runs down the group exactly as reactivity does: chlorine is a stronger oxidising agent than bromine, which is stronger than iodine. Chlorine can oxidise bromide and iodide ions; bromine can oxidise iodide ions but not chloride ions; iodine can oxidise neither.

One asymmetry to carry carefully: this "most reactive = strongest oxidising agent" logic works for non-metals only. For metals the ordering flips purpose — the most reactive metal is the best reducing agent. The difference traces to the ions each family forms: non-metals react to form negative ions (they take electrons), metals react to form positive ions (they give them away).

And fluorine gets one honest footnote. The trend explanation coming in the next section would predict, correctly, that fluorine out-reacts chlorine — but the true microscopic reason does not fit the family pattern: fluorine atoms do not actually accept electrons more easily than chlorine atoms do. Its exceptional violence traces to the unusually weak bond between the F₂ atoms. At International GCSE you will not be asked to explain why fluorine beats chlorine; the standard group-trend argument is the accepted answer at this level if the question ever appears.

Here is the Paper-2 crown of the chapter: why does reactivity fall down Group 7? The answer runs through how a halogen reacts. In the reactions of this chapter, a halogen reacts by forming a 1− ion — it takes electrons away from something else (a metal atom, or a halide-ion in solution). So reactivity measures how eagerly the atom can gain an electron, and that eagerness depends on one geometric fact: the size of the atom.

A chlorine atom is smaller than a bromine atom. Add an electron to each and the chlorine's new electron enters a shell closer to the nucleus — chlorine has fewer filled shells in the way. An electron in a shell closer to the nucleus is more strongly attracted to it. So chlorine pulls the incoming electron in more strongly, forms its 1− ion more readily, and therefore has the stronger tendency to react: chlorine out-reacts bromine, and will take an electron away from a bromide ion. Down the group the atoms grow shell by shell, the incoming electron sits further from the nucleus, the attraction weakens, electron-gaining gets harder — and reactivity falls: Cl₂ > Br₂ > I₂.

This is the mirror image of the Group 1 story from Chapter 11 (u2:c11). There, reactivity rose down the group because losing an electron gets easier as the outer electron sits further out. Here, reactivity falls down the group because gaining one gets harder. Same physics — outer-shell distance controls electrostatics — opposite directions, because the families do opposite jobs with the electron.

Everything the chapter taught, compressed onto one card. The Group 7 elements:

• have diatomic molecules, X₂ (F₂, Cl₂, Br₂, I₂);
• go from gas to liquid to solid as you descend the group — melting points and boiling points increase;
• have coloured, poisonous vapours — the colours darken down the group;
• form compounds of formula HX with hydrogen (HF, HCl, HBr, HI), which dissolve in water to form acids;
• form ionic salts with metals and covalent compounds with non-metals;
• form X⁻ ions (F⁻, Cl⁻, Br⁻) in their ionic compounds;
• become less reactive towards the bottom of the group;
• displace elements lower down the group from their salts.

Each line is a lever for prediction, and prediction is the last skill this chapter hands you. Take astatine — the element below iodine, so rare and radioactive that no one has ever seen a sample of it. You have never observed astatine, and you do not need to: every property above extrapolates down one more rung. Astatine will be:

• a darker colour than iodine — very dark grey or black;
• a solid at room temperature, with a higher melting point than iodine's;
• diatomic, existing as At₂ molecules;
• reactive with hydrogen to form HAt, which dissolves in water to form an acid;
• a salt-former with metals — NaAt with sodium, KAt with potassium;
• built around the astatide ion, At⁻, in its ionic salts;
• less reactive than iodine;
• displaced from solution by iodine.

Notice what happened there: eight properties, zero experiments. The trend in physical properties predicted the state, the colour, the melting point; the reaction families predicted the formulas (HAt, NaAt, KAt, At⁻); the displacement ladder predicted who beats whom in solution. That is the real power of a chemical family — knowledge of a few members plus the trend is knowledge of all of them, and this exact move (trend + application) is what prediction questions on Paper 2 are built from.