A disproportionation reaction is a redox reaction in which the same element in a single oxidation state is simultaneously oxidised and reduced (to a higher and a lower oxidation state). Example: 2H2O2 -> 2H2O + O2, where oxygen in the -1 state is both oxidised (to 0 in O2) and reduced (to -2 in H2O). Another example is Cl2 + 2NaOH -> NaCl + NaOCl + H2O.

Zinc goes from 0 to +2 (loses electrons), so Zn is oxidised and acts as the reducing agent. Copper goes from +2 (in CuSO4) to 0 (in Cu) by gaining electrons, so Cu2+ is reduced and CuSO4 acts as the oxidising agent. Thus Zn is the reducing agent and CuSO4 is the oxidising agent.

Zn + Cu2+ -> Zn2+ + Cu

Hydrogen resembles both alkali metals (group 1) and halogens (group 17), so its position is anomalous. Like alkali metals, it has one valence electron (1s1), forms H+ and shows +1 oxidation state. Like halogens, it is one electron short of a noble gas configuration, is diatomic (H2), and can gain an electron to form the hydride ion (H-). Because it fits neither group perfectly, its placement is debated.

In H2O2 the oxidation number of oxygen is -1, which is intermediate between 0 (in O2) and -2 (in H2O). Therefore it can be reduced to -2 (acting as an oxidising agent) or oxidised to 0 (acting as a reducing agent), depending on the other reactant. For example, it oxidises PbS to PbSO4 (oxidising agent) and reduces acidified KMnO4 (reducing agent).

Alkali metals have low ionization enthalpies because of their large size and a single loosely held valence electron. They readily lose this electron to form unipositive ions, i.e. they are easily oxidised. Since a substance that is easily oxidised is a good reducing agent, alkali metals are strong reducing agents. Their reducing power generally increases down the group.

M -> M+ + e-

A diagonal relationship is the similarity in properties between an element and the element placed diagonally to its lower right in the periodic table (e.g. Li and Mg, Be and Al, B and Si). Lithium and magnesium resemble each other because they have similar atomic and ionic sizes and nearly the same charge-to-size (polarising power) ratio. For example, both form nitrides directly with nitrogen and both form covalent, water-soluble compounds unlike the rest of their groups.

Lithium differs from the other alkali metals because of its very small atomic and ionic size, high charge density (high polarising power) and the absence of d-orbitals in its valence shell. As a result, its compounds have appreciable covalent character; for example, LiCl is soluble in organic solvents, Li forms a nitride (Li3N) and its carbonate and hydroxide decompose on heating, unlike those of Na and K.

The inert pair effect is the reluctance of the two outermost s-electrons (ns2) to take part in bonding, which becomes more pronounced down a group. As a result the lower oxidation state (2 less than the group valency) becomes more stable for heavier elements. For example, in group 14 the +2 state becomes more stable than +4 down the group, so Pb2+ is more stable than Pb4+, while carbon and silicon prefer +4.

In BF3, boron has only six electrons in its valence shell after forming three B-F bonds, so it has an incomplete octet and an empty p-orbital. This makes BF3 electron-deficient, so it can accept a lone pair of electrons from a donor (Lewis base) such as NH3 to complete its octet (forming F3B<-NH3). A species that accepts an electron pair is a Lewis acid, hence BF3 is a Lewis acid.

BF3 + :NH3 -> F3B-NH3

Catenation is the self-linking of atoms of the same element to form long chains or rings. Carbon shows catenation to the maximum extent because the C-C bond is very strong (high bond energy) due to the small size of the carbon atom, allowing effective overlap. Three crystalline allotropes of carbon are diamond, graphite and fullerene (C60).

In diamond each carbon is sp3 hybridised and bonded to four others in a rigid three-dimensional tetrahedral network, so it is extremely hard and a non-conductor. In graphite each carbon is sp2 hybridised and bonded to three others in flat hexagonal layers held together by weak van der Waals forces, so the layers slide easily (soft/lubricating). The fourth (unhybridised) electron of each carbon is delocalised over the layer, allowing graphite to conduct electricity.

(CH3)2CHCH2CH3 has a four-carbon main chain (butane) with a methyl group on the second carbon, so its IUPAC name is 2-methylbutane. CH3COOH is a two-carbon carboxylic acid, so its IUPAC name is ethanoic acid (common name acetic acid).

The inductive effect is the permanent displacement of the shared sigma-bond electron pair towards the more electronegative atom, transmitted through a chain of carbon atoms and decreasing with distance. Groups that push electrons away from themselves (electron-releasing, e.g. alkyl groups -CH3) show the +I effect, while groups that pull electrons towards themselves (electron-withdrawing, e.g. -NO2, -Cl) show the -I effect.

Resonance is the representation of a molecule or ion by two or more structures (canonical/contributing structures) that differ only in the position of electrons, none of which alone describes it fully; the actual structure is a resonance hybrid. Resonance energy is the difference in energy between the resonance hybrid and the most stable contributing structure; the greater the resonance energy, the more stable the molecule. Example: benzene, which is a hybrid of two Kekule structures and is more stable than expected.

An electrophile (electron-loving) is an electron-deficient species that accepts a pair of electrons; examples: H+, NO2+ (also positively charged or neutral electron-deficient species like BF3). A nucleophile (nucleus-loving) is an electron-rich species that donates a pair of electrons; examples: OH-, CN- (also neutral species with lone pairs like NH3 and H2O).

In homolytic fission, a covalent bond breaks so that each bonded atom takes one of the shared electrons, producing neutral free radicals (species with an unpaired electron); it usually occurs in the presence of heat or light in non-polar bonds. In heterolytic fission, the bond breaks so that one atom takes both shared electrons, producing oppositely charged ions (a cation and an anion, e.g. a carbocation and an anion).

Wurtz reaction: two molecules of an alkyl halide react with sodium metal in dry ether to give a symmetrical alkane with double the number of carbon atoms, e.g. 2CH3Cl + 2Na -> CH3-CH3 + 2NaCl (ethane). Decarboxylation: the sodium salt of a carboxylic acid is heated with soda lime (NaOH + CaO) to give an alkane with one carbon less, e.g. CH3COONa + NaOH -> CH4 + Na2CO3 (methane).

2CH3Cl + 2Na -> C2H6 + 2NaCl

Markovnikov's rule states that when an unsymmetrical reagent (HX) adds to an unsymmetrical alkene, the negative part of the reagent (X) attaches to the carbon bearing fewer hydrogen atoms, while hydrogen adds to the carbon bearing more hydrogen atoms. So HBr adds to propene (CH3-CH=CH2) to give mainly 2-bromopropane (CH3-CHBr-CH3), because the more stable secondary carbocation is formed. In the presence of peroxide, addition is anti-Markovnikov (peroxide/Kharasch effect), giving 1-bromopropane.

CH3-CH=CH2 + HBr -> CH3-CHBr-CH3

Ozonolysis is the reaction of an alkene with ozone (O3) to form an unstable ozonide, which on reductive cleavage (with Zn and water) breaks the carbon-carbon double bond and gives carbonyl compounds (aldehydes and/or ketones). It is used to locate the position of the double bond. Propene (CH3-CH=CH2) on ozonolysis gives ethanal (CH3CHO) and methanal (HCHO).

CH3-CH=CH2 -> CH3CHO + HCHO

Benzene has a stable, delocalised aromatic pi-electron cloud (6 pi electrons). Addition reactions would destroy this aromatic stability, so benzene prefers substitution, which preserves the aromatic ring. The pi cloud attracts electrophiles, so it undergoes electrophilic substitution. Examples: nitration (with conc. HNO3 + conc. H2SO4, electrophile NO2+) and halogenation (with Cl2/FeCl3); others are sulphonation and Friedel-Crafts alkylation/acylation.