1The idea in one paragraph

Almost every stable molecule has its electrons in pairs. A radical is a species with one electron left unpaired, and that lone electron makes it hungry: it will pull an atom off almost any molecule it hits in order to pair up. Radicals are made by homolytic fission, where a covalent bond splits evenly and each atom walks away with one electron of the pair; UV light or heat supplies the energy. Alkanes are so unreactive that they ignore acids, alkalis and most other reagents, but a halogen radical can take a hydrogen atom from an alkane. The new radical it leaves behind reacts with a halogen molecule and makes another halogen radical, so the process runs as a chain reaction: initiation makes radicals, propagation passes the unpaired electron on while making product, and termination ends the chain when two radicals meet. Nothing controls which hydrogen is swapped or when the swapping stops, so the product is always a mixture.

2What a radical is

A radical is a molecular entity that has an unpaired electron. "Molecular entity" is the guide's deliberately wide term: a radical can be an atom, a molecule, or an ion. What makes it a radical is only the odd electron.

The quickest test is to count. Add up the valence electrons of every atom, subtract one for each positive charge, add one for each negative charge. An odd total means at least one electron is unpaired, so the species is a radical. An even total usually means every electron is paired. Figure 1 shows four cases.

• An atom. A chlorine atom has 7 valence electrons: three lone pairs and one left over. Written Cl•.
• A molecule. The methyl radical, CH₃, has 4 + 3 = 7 valence electrons. Three pairs form the three C–H bonds and the seventh sits on the carbon. Written •CH₃. The hydroxyl radical, OH, has 6 + 1 = 7: one bonding pair, two lone pairs on the oxygen and one odd electron. Written HO• or •OH. Nitrogen dioxide, NO₂, with 5 + 12 = 17 valence electrons, is a radical too: the brown gas in polluted city air, stable enough to last for hours.
• An ion. Take one electron away from methane (8 valence electrons) and you get [CH₄]•⁺, with 7: a radical cation. This is exactly what a mass spectrometer makes when it ionizes a molecule, which is why the molecular ion of S3.2 is sometimes written M⁺•. Add one electron to an oxygen molecule (12 valence electrons) and you get O₂•⁻, the superoxide ion, with 13: a radical anion.

A radical has an unpaired electron. Write it with a dot, and put the dot on the atom that carries the electron: •CH₃, not CH₃ with a dot floating somewhere.

Why are radicals so highly reactive? Pairing an electron into a bond releases energy, so a radical gains a great deal by finding a partner for its odd electron, and the reactions that let it do so need very little activation energy. Almost any collision with a molecule it can take an atom from is enough. So radicals exist only briefly, at very low concentrations: you make them in the flask, where they are used.

The other side of this is the reverse process. When two radicals meet, their two unpaired electrons pair up to form a covalent bond, and energy equal to the bond enthalpy is released. Bond formation from two radicals is the reverse of homolytic fission, and it is the step that ends a chain reaction (section 5).

3Making radicals: homolytic fission

A covalent bond is a pair of electrons shared between two atoms. It can break in two ways. Homolytic fission splits the pair evenly: each atom takes one electron, so both fragments are radicals. (The other way, heterolytic fission, gives both electrons to one atom and makes ions; it belongs to R3.4.)

A single electron moving is shown by a fish-hook arrow, a curved arrow with a single barb. Two fish-hooks start on the bond, one curling to each atom. Figure 2 draws the homolytic fission of chlorine.

Cl₂ → 2Cl•

The energy comes from ultraviolet (UV) light or from heat: an absorbed UV photon carries enough energy to break a Cl–Cl bond, and strong heating does the same job through collisions. In an exam, "UV light" is the condition to write.

Why does the Cl–Cl bond break, and not the C–H bonds of the methane it is mixed with? Because it is the weakest bond in the mixture. Figure 3 compares typical bond enthalpies from the data booklet. (These are typical values; use the figures in your own booklet if they differ.)

The halogen–halogen bonds sit at 150 to 250 kJ mol⁻¹, against about 414 for C–H and 346 for C–C. Light with enough energy to split Cl₂ does not have enough to split C–H, so the chlorine molecules break and the methane molecules do not. This is the step called initiation: the step that creates radicals from molecules that had none.

Linking question: why do CFCs release chlorine radicals but not fluorine radicals? A chlorofluorocarbon such as CCl₂F₂ contains both C–Cl bonds (about 324 kJ mol⁻¹) and C–F bonds (about 492 kJ mol⁻¹). High in the stratosphere, UV radiation has enough energy to break the weaker C–Cl bond, so Cl• is released. The C–F bond is far stronger and survives. The chlorine radical then attacks ozone in a chain of its own:

The chlorine radical is regenerated, so one radical can destroy a very large number of ozone molecules. The same Cl• does not break O₂. That tells you the bonds in O₃ are weaker than the O=O double bond in O₂: in ozone each oxygen–oxygen bond is somewhere between a single and a double bond (S2.2, resonance), and weaker bonds are easier to break.

4Why alkanes react only with radicals

Alkanes are the least reactive organic family: hexane ignores acid, alkali and acidified potassium manganate(VII) at room temperature. The guide asks for two reasons.

The bonds are strong. Every bond in an alkane is either C–C (about 346 kJ mol⁻¹) or C–H (about 414 kJ mol⁻¹). Breaking one needs a lot of energy.

The bonds are essentially non-polar. Carbon and hydrogen have similar electronegativities, 2.6 and 2.2, so the electrons in a C–H bond are shared almost evenly. There is no δ+ carbon to attract a species with a lone pair, and no region of high electron density to attract a positive species. An alkane also has no lone pairs and no double bonds. Reagents that work by attracting charge, the acids, bases, nucleophiles and electrophiles of R3.1 and R3.4, have nothing to grip.

A radical does not need a charged target. It takes a hydrogen atom, and the new bond it makes pays for most of the one broken. So radicals are among the few things that react with alkanes at all.

Linking question: why are alkanes kinetically stable but thermodynamically unstable? Methane burns with ΔH of about −890 kJ mol⁻¹ (a typical data-booklet value), so the products are far lower in energy than the reactants: methane is thermodynamically unstable with respect to carbon dioxide and water. Yet methane mixed with air does not react at room temperature, because the activation energy is very high and almost no collisions succeed: methane is kinetically stable. A spark or flame supplies the activation energy, and the energy released keeps the reaction going. Figure 4 draws both ideas on one energy profile.

5The chain reaction: methane and chlorine

Mix methane and chlorine in the dark and nothing happens. Shine UV light on the mixture and chloromethane and hydrogen chloride form. The reaction is a substitution: a chlorine atom replaces a hydrogen atom. Because the reacting species are radicals, it is called radical substitution, and it takes place in three kinds of step. Figure 5 shows how they fit together.

Initiation. UV light splits a chlorine molecule. Radicals are created.

Propagation. A radical reacts with a molecule to make a product molecule and a new radical. There are two propagation steps, and they run in a loop.

Termination. Two radicals meet and pair their electrons into a bond. Radicals are removed and the chain stops.

The marks sit in the propagation steps. In step 1 the chlorine radical takes a hydrogen atom, giving HCl and a methyl radical; writing Cl• + CH₄ → CH₃Cl + H• loses the mark, because no hydrogen radicals form. In step 2 the methyl radical takes a chlorine atom from Cl₂, making CH₃Cl and handing back a Cl•, which goes straight back into step 1.

Figure 6 shows the same two steps with fish-hook arrows, one electron at a time. You are not required to draw arrows for propagation, but seeing them makes clear why a radical goes in and a radical comes out.

Add the two propagation steps and the radicals cancel: CH₄ + Cl₂ → CH₃Cl + HCl. That is the overall equation, and it is why propagation is the step that makes the product. Each propagation step uses up one radical and makes one, so the number of radicals never changes. A single chlorine radical from one photon can go round the loop thousands of times. That is what "chain reaction" means, and it is why a little light produces a lot of product.

Termination is rare by comparison, because radicals are so dilute that two seldom meet. One termination product is the best evidence for the mechanism: small amounts of ethane are found in the product, and ethane can only come from two methyl radicals joining.

Why the propagation steps are fast. Use bond enthalpies to find ΔH for each one (typical data-booklet values: C–H 414, H–Cl 431, Cl–Cl 242, C–Cl 324 kJ mol⁻¹).

Both steps are exothermic: each breaks one bond and makes a stronger one, so the activation energies are small. Try it Q4 runs the same sums for iodine and shows why iodine does not react with methane.

6Why you get a mixture of products

The guide's words are "producing a mixture of products", and you must be able to say why. Figure 7 shows three separate reasons.

It keeps going. Chloromethane still has C–H bonds, so a chlorine radical can take a hydrogen from it just as easily as from methane. That gives •CH₂Cl, which reacts with Cl₂ to give dichloromethane, CH₂Cl₂. The same happens again to give trichloromethane, CHCl₃, and tetrachloromethane, CCl₄. The propagation steps for each are the same pattern:

Termination makes new alkanes. Two methyl radicals give ethane, and ethane is then substituted in its turn to give chloroethane and more.

Bigger alkanes have more than one kind of hydrogen. Propane, CH₃CH₂CH₃, has six hydrogens on the end carbons and two on the middle one. Take an end hydrogen and you make 1-chloropropane; take a middle one and you make 2-chloropropane. These are position isomers (S3.2), and both form.

A chemist can shift the balance but not remove it. Excess alkane makes a chlorine radical far more likely to hit an alkane than a chloroalkane, so the monosubstituted product dominates; excess chlorine pushes towards CCl₄. Either way the products must be separated afterwards, which is why radical substitution is a poor way to make one pure chloroalkane.

7Writing the mechanism for any alkane

Paper 2 will not ask for methane every time. The pattern transfers. Here is ethane with bromine, worked the way you should write it.

Step 1. Name the reaction and the condition. Radical substitution, UV light.

Step 2. Initiation. Split the halogen.

Step 3. Propagation. The halogen radical takes an H atom and leaves an alkyl radical; the alkyl radical takes a halogen atom and hands back a halogen radical.

Step 4. Termination. Pair up any two radicals present.

Step 5. Check by counting radicals: one on each side of every propagation step, two on the left and none on the right of every termination step. Bromine reacts more slowly than chlorine, because its first propagation step is endothermic, about 414 − 366 = +48 kJ mol⁻¹ with typical values.

Linking question: why is bromine water decolourized in the dark by alkenes but not by alkanes? This is the standard test for unsaturation, and Figure 8 shows it.

Shake hexane with orange bromine water in the dark and the colour stays: hexane reacts with bromine only through radicals, and in the dark nothing splits Br₂. In UV light the colour slowly fades as radical substitution runs. Shake hex-1-ene with bromine water in the dark and the colour disappears at once, because the electron-rich C=C double bond attacks Br₂ directly, by the electrophilic addition of R3.4. The test works because the two families react by different mechanisms.