IB Chemistry HLTopic 6 — Electron SharingPaper 1 & 2Core idea~9 min read
Free Radicals
Up to now electrons have moved around in pairs. A radical breaks that habit: it has one electron sitting on its own, with no partner. That single lonely electron is why radicals are so desperate to react, and why the chemistry on the next two pages looks nothing like anything you have met so far.
📚 What you need to know
A radical is any species with an unpaired electron. That is the whole definition.
The unpaired electron is shown as a dot, placed on the atom that carries it.
Radicals can be atomic, molecular, anionic or cationic — charge is irrelevant.
Radicals are very reactive and short-lived, because the unpaired electron makes them high in energy.
They react by taking an atom or electron from something else (making a new radical), or by joining with another radical.
An equation shows what reacts; a mechanism shows how, step by step.
Paired, or not paired
Every stable molecule you have drawn so far has its electrons in pairs: bonding pairs holding atoms together and lone pairs sitting on atoms. A radical has one electron left over with nowhere to go.
The red electron is the whole story. Give the chlorine atom one more electron and it becomes a stable chloride ion — which is exactly why chlorine radicals grab electrons so eagerly.
A shortcut worth knowing: add up the outer-shell electrons of the whole species, adjusting for charge. If the total is odd, one electron cannot possibly be paired, so it must be a radical. NO has 5 + 6 = 11 electrons, so it is a radical without you having to draw anything.
The four types
Notice the superoxide ion O2−• is both an anion and a radical at the same time. The two labels are describing completely different things.
Where to put the dot
In a species made of several atoms, the dot belongs on the atom that actually holds the unpaired electron — not floating somewhere near the formula.
•CH3 — the dot goes on the carbon, because that is where the electron sits after a C−H bond has broken.
OH• — the dot goes on the oxygen.
In a displayed formula, draw the dot right next to the atom it belongs to.
Do not confuse the radical dot with a lone pair or with the dots in a Lewis diagram. The radical dot is a single electron. If you can see two dots together, that is a pair, and pairs do not make radicals.
Why radicals are so reactive
An unpaired electron is an unstable arrangement. The species is high in energy, and it can lower that energy by getting the electron paired up somehow. There are only two ways to do that:
🧩 The two things a radical can do
Steal from something else. It takes an atom or electron from a neighbouring molecule. That molecule is now short of an electron, so it becomes a radical too — the problem has been passed on, not solved.
Pair up with another radical. Two unpaired electrons join to form a covalent bond, and both radicals disappear at once. This is the only way the chain actually stops.
Either way it happens fast. Radicals typically survive for a tiny fraction of a second before reacting.
Why this matters beyond the exam: radicals in the upper atmosphere are what destroy ozone, and radicals in your body are what antioxidants are sold to mop up. Step 1 above — passing the problem on — is exactly why one chlorine radical can wreck thousands of ozone molecules.
Worked examples
WORKED EXAMPLE
Which of these are radicals? H, HCl, NO, Br−, •CH2CH3, CO2
Step 1: Count outer electrons, adjusting for chargeH = 1 HCl = 1 + 7 = 8 NO = 5 + 6 = 11Br− = 7 + 1 = 8 C2H5 = 8 + 5 = 13 CO2 = 4 + 12 = 16Step 2: Odd totals cannot pair up1, 11 and 13 are odd, so those three must have an unpaired electron.H, NO and •CH2CH3 are radicalsBr− is charged but not a radical — charge and radical are separate ideas
WORKED EXAMPLE
Classify each of these radicals as atomic, molecular, anionic or cationic: Br•, •CH3, O2−•
Step 1: Ask how many atoms, then ask about chargeOne atom means atomic. Several atoms with no charge means molecular.Step 2: Work through themBr• — one atom, no charge → atomic•CH3 — four atoms, no charge → molecularO2−• — two atoms, one negative charge → anionicAtomic, molecular and anionicall three are still radicals — the type is just extra description
WORKED EXAMPLE
Explain why radicals are described as short-lived, and why one radical often leads to many reactions.
Step 1: Start from the unpaired electronIt makes the species high in energy and therefore unstable.Step 2: Say what it does about thatIt reacts almost immediately with the first thing it meets, so it does not exist for long.Step 3: Follow what happens nextTaking an atom from another molecule turns that molecule into a radical.Short-lived because they react on contact; the problem is passed along, creating a chain“it makes a new radical” is the phrase that earns the second mark
💡 Exam tip
Define a radical as a species with an unpaired electron. Do not say “an unpaired pair” or “a free electron”.
Use the odd electron count shortcut to spot radicals quickly, then check by drawing if you have time.
Put the dot on the correct atom. Examiners mark •CH3 and CH3• as the same, but a dot on a hydrogen is wrong.
Remember that charge is independent. A radical can carry a charge, or none at all.
When explaining reactivity, link it back to the unpaired electron and to high energy.
Learn the two ways radicals react — they become initiation, propagation and termination two pages from now.
⚠ Common mix-up
Thinking all ions are radicals. Br− has a full outer shell of paired electrons and is not a radical.
Thinking radicals must be neutral. Anionic and cationic radicals both exist.
Confusing the radical dot with a lone pair. One dot is one electron; a lone pair is two.
Writing the dot in the wrong place in a polyatomic radical.
Saying radicals are unstable “because they are charged”. Most are not charged at all.
Forgetting to count the charge when totalling outer electrons. An anion has gained one.
Up next: Homolytic Fission — radicals have to come from somewhere. They are made by splitting a covalent bond straight down the middle, giving one electron to each atom.
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