Halogen Bonding: The Hydrogen Bond's Lesser-Known Cousin

Here is a puzzle. Halogen atoms in molecules are usually described as electron-rich: chlorine, bromine and iodine are electronegative, they carry a partial negative charge in a C–X bond, and they have three lone pairs each. So why would a halogen atom in one molecule be attracted to the lone pair of a nitrogen or oxygen atom in another?

Yet crystallographers kept finding exactly that: iodine or bromine atoms pressed up against nitrogen and oxygen atoms at distances shorter than the atoms’ van der Waals radii allow for a mere touching contact, and almost always in a straight line. The interaction is now called the halogen bond, and it was given a formal definition in an IUPAC recommendation in 2013. It turns out to be a close relative of the hydrogen bond, with a twist that makes it especially useful to designers of crystals and drugs.

The key idea: the σ-hole

The answer to the puzzle is that the electron density around a bonded halogen is not spread evenly.

Take a molecule R–X, where X is chlorine, bromine or iodine bonded to some group R. The halogen uses one of its valence electrons to form the σ bond to R. Its other electrons, including the three lone pairs, crowd into a belt around the “waist” of the atom, perpendicular to the R–X bond. The region on the far side of the halogen, pointing directly away from R along the bond axis, ends up comparatively short of electron density.

If you map the electrostatic potential on the surface of the molecule, that far tip often comes out positive, even though the halogen as a whole carries a small negative charge. This positive cap is called the σ-hole: a “hole” in the electron density, sitting on the extension of the σ bond.

So a bonded halogen is two-faced:

• Around its equator it is electron-rich (δ−) and can act as a Lewis base, accepting hydrogen bonds or binding to metal ions.
• At its tip it is electron-poor (δ+) and can act as a Lewis acid, attracting lone pairs, anions or π electrons.

A halogen bond is the attraction between that σ-hole and an electron-rich partner B. It is written R–X···B, with the dotted line representing the halogen bond, just as O–H···O represents a hydrogen bond.

An analogy

Picture a planet with a ring of thick cloud around its equator and clear skies at both poles. Viewed from the side, the cloud belt dominates. But fly directly over one pole and you look straight down onto the bare surface. The halogen’s σ-hole is that clear pole. An approaching lone pair only “feels” the positive region if it comes in exactly along the axis; approach from the side and it meets the cloud belt and is repelled.

Why halogen bonds are so directional

That analogy explains the most striking feature of halogen bonds: they are highly linear. The R–X···B angle is typically close to 180°. The σ-hole is a small, focused region on the bond axis, and moving the partner even modestly off-axis brings it into the negative belt.

Hydrogen bonds are directional too, but they tolerate more bending. A hydrogen atom has no lone pairs and no electron belt; its positive region is spread more broadly. That difference makes halogen bonds attractive to chemists who want to control geometry precisely.

The distance is also telling. In a halogen bond, the X···B separation is less than the sum of the van der Waals radii of X and B. The two atoms are closer than they would be if they were simply resting against each other, which is a sign of real attraction.

What controls the strength

Two factors dominate.

1. Which halogen

The σ-hole grows as the halogen gets larger and more polarisable, and as its electronegativity falls. So halogen-bond strength generally follows:

I > Br > Cl, with F rarely taking part.

Iodine is the star performer. Its outer electrons are held loosely and are easy to push around, and the positive σ-hole is large. Fluorine is the opposite case: it is so electronegative and so tightly held that its electron density usually covers the whole atom, leaving no positive tip. Fluorine only forms halogen bonds in unusual cases where it is attached to very strongly electron-withdrawing groups. You can review the group trends in the halogens.

2. What the halogen is attached to

Electron-withdrawing groups on R drain yet more electron density from the halogen and deepen the σ-hole. Iodine on a benzene ring whose other positions carry fluorine atoms (as in C₆F₅I) is a much better halogen-bond donor than iodine on plain benzene. Iodine in the diatomic molecule I₂ is also a respectable donor, because each iodine has only another iodine as its partner.

The acceptor matters too. Good partners include nitrogen atoms in amines and pyridines, carbonyl and ether oxygens, sulfur atoms, halide anions and π systems.

Strengths span a wide range, from weak contacts comparable to dispersion forces up to interactions stronger than many ordinary hydrogen bonds, especially with iodine donors and anionic acceptors.

A colour you have probably seen

The simplest everyday sign of halogen bonding is the colour of iodine solutions. Dissolve I₂ in a solvent with no lone pairs available, such as hexane, and you get a violet solution, close to the colour of iodine vapour. Dissolve it in a solvent with donor atoms, such as ethanol, and the solution turns brown. In the second case, solvent lone pairs interact with the σ-hole on iodine, forming loose complexes that shift how the molecule absorbs light. The same kind of interaction between I₂ and amines was known for well over a century before anyone called it halogen bonding.

Halogen bonds compared with hydrogen bonds

Feature	Hydrogen bond	Halogen bond
Electron-poor partner	H on N, O or F	Cl, Br or I (σ-hole)
Written as	X–H···B	R–X···B
Angle	Prefers near 180°, tolerates bending	Very close to 180°, bends little
Tuning	Mostly by changing the donor atom	By choosing the halogen and the groups on R
Size of donor	Tiny hydrogen atom	Large, polarisable atom
Behaviour in water	Competes with water’s own hydrogen bonds	Less affected, as halogens are fairly hydrophobic

The resemblance is why the name was chosen. In both cases an electron-poor region on one atom meets a lone pair on another, and in both cases the attraction is mostly electrostatic, with contributions from polarisation, charge transfer and dispersion. For a refresher on the parent interaction, see hydrogen bonding.

Halogen–halogen contacts

Halogen atoms in crystals often touch each other, and two geometries appear:

• Type I: the two C–X···X angles are roughly equal. This is a symmetric contact, often just the result of close packing, and is not a halogen bond.
• Type II: one C–X···X angle is near 180° and the other near 90°. Here the σ-hole of one halogen points at the electron-rich belt of the other. This is a genuine halogen bond.

Looking at the angles is a quick way to decide whether a short contact in a crystal structure is a halogen bond.

Where halogen bonding is used

• Crystal engineering. Because halogen bonds are strong enough to matter and linear enough to be predictable, chemists use them to assemble molecules into chains, sheets and networks by design, for example pairing iodoperfluorocarbons with nitrogen-containing bases.
• Drug design. Many drug molecules carry chlorine, bromine or iodine. In some cases the halogen forms a halogen bond with a backbone carbonyl oxygen or another Lewis base in the target protein, improving binding. Medicinal chemists now consider this when choosing where to put a halogen. More on the wider field in drug design.
• Anion recognition and sensing. Receptors that present iodine atoms towards a central cavity can bind halide ions selectively.
• Materials. Halogen bonds have been used to organise liquid crystals and other soft materials.

Common mistakes

1. “Halogens are negative, so they can’t attract lone pairs.” The halogen’s charge is anisotropic. The belt is negative; the tip can be positive. Both matter.
2. Assuming fluorine forms the strongest halogen bonds. It is the reverse of what electronegativity might suggest: iodine is best and fluorine rarely participates.
3. Treating every short X···X contact as a halogen bond. Check the geometry. Only type II contacts, with one angle near 180°, qualify.
4. Confusing a halogen bond with a hydrogen bond to a halogen. In O–H···Cl, chlorine is the acceptor of a hydrogen bond, using its lone pairs. In C–I···N, iodine is the donor of a halogen bond, using its σ-hole.
5. Ignoring angle. A halogen bond approached off-axis weakens rapidly. Drawing it bent is a sign something is wrong.

Key takeaways

• A halogen bond R–X···B is an attraction between an electron-poor σ-hole on a bonded halogen and an electron-rich partner such as N, O, S, an anion or a π system.
• The σ-hole sits on the extension of the R–X bond, which makes halogen bonds very close to linear (near 180°).
• Strength generally follows I > Br > Cl, with fluorine rarely involved, and increases when R is electron-withdrawing.
• Halogen bonds are close cousins of hydrogen bonds, formally defined in an IUPAC recommendation in 2013, and are used in crystal engineering, drug design and anion sensing.