Halogen and Interhalogen Chemistry

Core Idea

The group 17 elements form diatomic molecules X₂ whose oxidizing power falls steadily from fluorine to iodine. Fluorine, the most electronegative element, shows only the −1 oxidation state in its compounds, while chlorine, bromine and iodine also reach +1, +3, +5 and +7 in oxoacids, oxoanions and interhalogens. Interhalogen compounds such as ClF₃, BrF₅ and IF₇ always have the heavier halogen at the center, and their shapes follow VSEPR.

How It's Best Learned

Build a group table of state at room temperature, X–X bond energy and standard reduction potential, and look for the one anomaly (the weak F–F bond). Use the reduction potentials to predict displacement reactions such as Cl₂ + 2Br⁻. Then draw Lewis structures for ClF₃, BrF₅ and IF₇ and predict their shapes before checking them.

Common Misconceptions

• Bond energy falls smoothly down the group, so F₂ has the strongest X–X bond. (The F–F bond is weaker than the Cl–Cl bond because lone pairs on the two small fluorine atoms repel each other.)
• All halogens can take positive oxidation states. (Fluorine is the most electronegative element, so it is always −1 in its compounds.)
• In an interhalogen, the more electronegative halogen sits at the center. (The larger, less electronegative halogen is central and carries the positive oxidation state; fluorine atoms are terminal.)

Explainer

The halogens are fluorine, chlorine, bromine, iodine and the radioactive astatine. Each atom needs one electron to complete its octet, so each element exists as X₂ and acts as an oxidizing agent. At room temperature, fluorine is a pale yellow gas, chlorine a yellow-green gas, bromine a red-brown liquid and iodine a dark solid that gives a violet vapor. The change from gas to solid reflects stronger London forces between larger molecules.

Oxidizing strength falls down the group. Standard reduction potentials are about +2.87 V for F₂, +1.36 V for Cl₂, +1.07 V for Br₂ and +0.54 V for I₂. This order predicts displacement reactions: a halogen oxidizes the halide ions of any halogen below it. Chlorine water added to bromide gives bromine, and added to iodide gives iodine.

Fluorine is unusual in several ways. Its F–F bond, about 158 kJ/mol, is weaker than the Cl–Cl bond, about 242 kJ/mol, because the small atoms bring their lone pairs close together. This weak bond, together with the strong hydration of the small fluoride ion, explains fluorine's extreme reactivity. Fluorine is also the most electronegative element, so it never takes a positive oxidation state. Its acid, HF, is weak in water, while HCl, HBr and HI are strong.

The heavier halogens use positive oxidation states when bonded to oxygen or fluorine. Chlorine forms four oxoacids: hypochlorous acid HOCl (+1), chlorous acid HClO₂ (+3), chloric acid HClO₃ (+5) and perchloric acid HClO₄ (+7). Their acid strength rises with the number of oxygens. Hypochlorite, made from chlorine and base, is the active ingredient in household bleach.

Halogens also combine with each other to form interhalogens. Their general formulas are XY, XY₃, XY₅ and XY₇, where X is the larger halogen and Y is usually fluorine. Larger central atoms fit more fluorines around them, which is why IF₇ exists but ClF₇ does not. VSEPR predicts the shapes: ClF₃ is T-shaped, BrF₅ is square pyramidal and IF₇ is pentagonal bipyramidal. ClF₃ is among the most aggressive fluorinating agents known.

Finally, halide ions can bind halogen molecules to form polyhalide ions. The best known is triiodide, I₃⁻, which forms when iodine dissolves in iodide solution. It is linear. Polyiodide chains held inside the helix of starch's amylose give the familiar blue-black starch-iodine color.