Le Chatelier's principle states that if a system at equilibrium is subjected to a change in concentration, pressure or temperature, the equilibrium shifts in the direction that tends to counteract (reduce) the effect of that change. Increasing the pressure shifts this equilibrium in the forward direction (towards NH3), because the forward reaction reduces the number of gas moles from 4 (1 + 3) to 2, thereby lowering the pressure.

According to the Bronsted-Lowry concept, an acid is a substance that donates a proton (H+) and a base is a substance that accepts a proton. A conjugate acid-base pair is a pair of species that differ by a single proton; for example, in HCl + H2O -> H3O+ + Cl-, HCl/Cl- and H2O/H3O+ are conjugate acid-base pairs.

acid <=> base + H+

HCl is a strong acid and dissociates completely, so [H+] = 0.001 M = 1 x 10^-3 M. pH = -log[H+] = -log(10^-3) = 3. The solution is acidic (pH < 7).

pH = -log[H+]

The ionic product of water Kw is the product of the molar concentrations of hydrogen and hydroxide ions in water: Kw = [H+][OH-]. At 298 K, Kw = 1.0 x 10^-14 mol^2 L^-2. Taking negative logarithm gives pKw = pH + pOH = 14 at 298 K. For pure (neutral) water, [H+] = [OH-] = 10^-7 M, so pH = 7.

Kw = [H+][OH-] = 1.0 x 10^-14

The common ion effect is the suppression of the degree of dissociation (ionisation) of a weak electrolyte by the addition of a strong electrolyte that provides an ion common to the weak electrolyte. For example, adding NH4Cl (which provides NH4+) to a solution of NH4OH suppresses the ionisation of NH4OH, decreasing the OH- concentration. This is used in qualitative analysis and to control pH.

Oxidation is the loss of electrons or an increase in oxidation number of an element; reduction is the gain of electrons or a decrease in oxidation number. Both occur simultaneously in a redox reaction. For example, in Zn -> Zn2+ + 2e-, zinc is oxidised (oxidation number 0 -> +2).

In KMnO4: K = +1, O = -2 (four O = -8). Let Mn = x. Then +1 + x + (-8) = 0, so x = +7. In K2Cr2O7: 2 K = +2, 7 O = -14. Let each Cr = y. Then +2 + 2y - 14 = 0, so 2y = 12, y = +6. Thus Mn is +7 and Cr is +6.

sum of oxidation numbers = charge on species

An oxidising agent is a substance that oxidises another substance by accepting electrons and is itself reduced; e.g. KMnO4, O2. A reducing agent is a substance that reduces another substance by donating electrons and is itself oxidised; e.g. H2, C (carbon). In a redox reaction the oxidising agent gains electrons and the reducing agent loses electrons.

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).