Interhalogen Compounds

What Interhalogen Compounds Are

The halogens of Group 17 — fluorine, chlorine, bromine and iodine — react not only with other elements but also with one another, producing a series of mixed binary compounds known as interhalogen compounds. Each molecule contains atoms of exactly two different halogens, written in the general form $\ce{XX'_n}$, where X is the larger, more electropositive halogen and X′ is the smaller, more electronegative halogen, and $n = 1, 3, 5$ or $7$.

The reason the larger halogen sits at the centre is geometric. A central atom can only accommodate as many surrounding atoms as its size permits. As the ratio of the radius of X to the radius of X′ increases, more X′ atoms can pack around X, so the number of atoms per molecule rises. This single principle explains the whole family: small differences in size give the 1:1 type $\ce{XX'}$, while the extreme case — large iodine surrounded by the smallest halogen, fluorine — gives $\ce{IF7}$, the only species with seven peripheral atoms.

Because the bonding is between two non-metals of similar but unequal electronegativity, every interhalogen is a covalent, diamagnetic molecule. Their behaviour is broadly intermediate between that of the two parent halogens, with one important twist in reactivity that NEET examiners love to probe.

The Four Types: XX′ to XX′7

Interhalogens are classified purely by stoichiometry into four types. The 1:1 type is known for every halogen pair; the higher types appear only where the size difference is large enough, and fluorine — being smallest — is the peripheral atom in almost all of the higher members.

Note that fluorine never sits at the centre of an interhalogen molecule. As the most electronegative and smallest halogen, it cannot expand its valence shell or act as the larger, electropositive partner — so it is always the peripheral X′ atom, never X. This is also why iodine, the largest halogen, is the central atom in the highest-coordination species $\ce{IF5}$ and $\ce{IF7}$.

Preparation

Two general routes give interhalogen compounds: direct combination of the two halogens, or the action of a halogen on a lower interhalogen. Which product forms depends on the proportions of reactants and the temperature, so controlling conditions is the whole game.

The pattern is consistent: a larger excess of the smaller, more reactive halogen (almost always fluorine) drives the equilibrium toward the higher interhalogen. Lowering the fluorine ratio stops the reaction at the lower type. This dependence on stoichiometry is the key to selecting between $\ce{ClF}$ and $\ce{ClF3}$, or between $\ce{ICl}$ and $\ce{ICl3}$.

General Physical Properties

All interhalogen compounds are covalent and diamagnetic. They are volatile solids or liquids at 298 K, with the single exception of $\ce{ClF}$, which is a gas. Their physical properties — colour, melting and boiling points — are intermediate between those of the two constituent halogens, except that their melting and boiling points run a little higher than a naive average would predict, owing to the polarity of the $\ce{X-X'}$ bond.

Why They Out-React the Halogens

The headline chemical fact — and the one most often tested — is that interhalogen compounds are, in general, more reactive than the halogens, the lone exception being fluorine. The cause is bond strength: the $\ce{X-X'}$ bond in an interhalogen is weaker than the $\ce{X-X}$ bond in the parent halogens, except for the unusually strong $\ce{F-F}$ bond. A weaker bond breaks more readily, so the molecule is a more aggressive reagent.

Hydrolysis Pattern

Every interhalogen undergoes hydrolysis with water, and the products follow a clean rule. The halide ion comes from the smaller halogen (X′), while the oxoanion comes from the larger halogen (X), and the type of oxoanion rises step-by-step with the formula. The 1:1 case is the template:

$\ce{XX' + H2O -> HX' + HOX}$

The oxidation state of the larger halogen in the oxoanion climbs from +1 (hypohalite) to +7 (perhalate) exactly as the number of fluorine atoms in the formula climbs, which is a useful internal consistency check when reconstructing a hydrolysis product under exam pressure.

VSEPR Structures & Hybridisation

The geometries of the higher interhalogens are a direct application of VSEPR theory to the central atom. Count the central atom's seven valence electrons, subtract the bonding electrons, and the remainder gives the lone pairs that decide the shape. The four types map onto four distinct geometries.

XX′3 — the bent T-shape of $\ce{BrF3}$

In $\ce{BrF3}$ the central Br has seven valence electrons. Three of these form electron-pair bonds to three fluorine atoms, leaving four electrons as two lone pairs. With three bond pairs and two lone pairs, the electron domains adopt a trigonal bipyramidal arrangement ($sp^3d$). The two lone pairs occupy the equatorial positions — minimising the strong lone-pair–lone-pair and lone-pair–bond-pair repulsions — and the two axial fluorines bend slightly toward the equatorial fluorine. The result is a slightly bent T-shape.

XX′5 and XX′7 — $\ce{IF5}$ square pyramidal, $\ce{IF7}$ pentagonal bipyramidal

In $\ce{IF5}$ the central iodine forms five bonds and retains one lone pair (five bond pairs, one lone pair, $sp^3d^2$). The six electron domains sit at the vertices of an octahedron; the single lone pair takes one vertex, pushing the five fluorines into a square-pyramidal shape. In $\ce{IF7}$ the iodine forms seven bonds with no lone pairs ($sp^3d^3$), so the seven fluorines spread to the vertices of a pentagonal bipyramid.

Uses as Fluorinating Agents

Interhalogen compounds serve as non-aqueous solvents and, more importantly, as very powerful fluorinating agents. The weak, polar $\ce{X-F}$ bonds release fluorine readily, making compounds such as $\ce{ClF3}$ and $\ce{BrF3}$ ideal for converting metals and oxides into fluorides.

The textbook industrial application is the production of uranium hexafluoride, the volatile compound used to enrich uranium-235:

$\ce{U(s) + 3ClF3(l) -> UF6(g) + 3ClF(g)}$

Here $\ce{ClF3}$ fluorinates elemental uranium straight to gaseous $\ce{UF6}$, which is then separated by diffusion or centrifugation. This single reaction is the most frequently cited use of interhalogens, so tie "fluorinating agent + uranium enrichment" together in your memory.