What a radical is

A radical is a molecular entity with at least one unpaired electron. The definition is broad on purpose: a radical may be a single atom, a molecule, a cation or an anion, as long as the species contains an unpaired electron.

We show radicals with a dot, •, placed beside the atom carrying the unpaired electron. So Cl• is a chlorine atom with an unpaired electron, while •CH₃ is a methyl radical with the unpaired electron on carbon. In a larger species, that dot matters. It shows where the radical character sits.

A molecular entity is a distinct atom, molecule, ion or radical that can be treated as an individual chemical species. So the word “radical” does not automatically mean “neutral”. Examples include:

• atom radical: Cl•
• molecular radical: •CH₃
• radical cation: NH₄•⁺
• radical anion: O₂•⁻

Because the electron is unpaired, radicals are highly reactive: pairing that electron in a new covalent bond is usually favourable. In many organic reactions, radicals appear as short-lived intermediates, which are species formed during a reaction mechanism but not usually isolated as final products.

Why electron sharing is the theme

In ordinary covalent bonding, two atoms share a pair of electrons. Radical chemistry works differently, because single electrons move around. A radical can collide with a stable molecule, take part in bond breaking and bond making, then leave a new radical behind. That is the basis of a chain reaction: one reactive radical generates another reactive radical, so the process keeps going.

Two radicals can also meet and form a normal covalent bond. This removes unpaired electrons from the system and stops that particular chain. Later, we’ll call that termination. For now, the main idea is straightforward: radicals are reactive because they have unpaired electrons, and their formulae must show those unpaired electrons clearly.

Homolytic fission

Homolytic fission is bond breaking where each bonded atom takes one electron from the shared pair, producing two radicals. It differs from heterolytic fission, where both electrons move to one atom and ions are formed.

For a halogen molecule, the general equation is:

X₂ → 2X•

For chlorine:

Cl₂ → 2Cl•

This occurs when enough energy is supplied, usually from ultraviolet light or heat. UV light matters especially for halogens, as it can provide the energy needed to break the X–X bond.

Fish-hook arrows

A single-barbed arrow, often called a fish-hook arrow, is a curved arrow showing the movement of one electron. Use it in radical mechanisms. A normal double-barbed curly arrow shows the movement of an electron pair, so it isn’t the right arrow for homolytic fission.

In homolytic fission of Cl₂, draw one fish-hook arrow from the Cl–Cl bond to the left chlorine atom, and another fish-hook arrow from the same bond to the right chlorine atom. The two electrons in the covalent bond split evenly.

Be precise when drawing these arrows. The tail starts at the electron source, such as the covalent bond or an unpaired electron. The head finishes where that single electron goes. In radical chemistry, the arrowhead is single-barbed because only one electron is moving.

Initiation in a chain reaction

An initiation step is the first step in a radical chain reaction, producing radicals from non-radical reactants. In halogenation reactions, initiation is usually homolytic fission of the halogen molecule:

Cl₂ → 2Cl•

or, for bromine:

Br₂ → 2Br•

A chain reaction is a reaction sequence where a reactive intermediate made in one step causes further steps of the same type to happen. The initiation step gets the chain going by making the first radicals.

Reverse process and bond strength links

Homolytic fission can run in reverse: two radicals join to form a covalent bond. In that process, two single electrons pair up and make a bond, for example:

Cl• + Cl• → Cl₂

This also counts as a termination reaction when it happens during a radical chain mechanism.

CFCs in the atmosphere can release chlorine radicals more readily than fluorine radicals because the relevant C–Cl bonds are weaker than C–F bonds. A weaker bond needs less energy for homolytic fission, so UV radiation in the atmosphere is more able to break C–Cl bonds than C–F bonds.

Chlorine radicals can break down ozone, O₃, in the stratosphere, but typically do not break down oxygen, O₂, in the same way. That tells us the bonding in O₂ is stronger and less easily disrupted than the bonding arrangement in O₃. Put simply, ozone is more vulnerable to radical attack than oxygen.