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+]

NaOH is a strong base and dissociates completely, so [OH-] = 0.01 M = 10^-2 M. pOH = -log[OH-] = -log(10^-2) = 2. Since pH + pOH = 14, pH = 14 - 2 = 12. The solution is basic (pH > 7).

pH + pOH = 14

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

A buffer solution is one that resists a change in its pH on the addition of a small amount of acid or base. An acidic buffer is made from a weak acid and its salt with a strong base, e.g. acetic acid + sodium acetate (CH3COOH + CH3COONa). The Henderson-Hasselbalch equation is pH = pKa + log([salt]/[acid]).

pH = pKa + log([salt]/[acid])

AgCl dissociates as AgCl <=> Ag+ + Cl-. If solubility = s mol/L, then [Ag+] = [Cl-] = s. Ksp = [Ag+][Cl-] = s x s = s^2. So s = sqrt(Ksp) = sqrt(1.8 x 10^-10) = 1.34 x 10^-5 mol/L.

Ksp = s^2 (for AB type salt)

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

Reduction half: MnO4- + 8H+ + 5e- -> Mn2+ + 4H2O. Oxidation half: Fe2+ -> Fe3+ + e-. To balance electrons, multiply the oxidation half by 5: 5Fe2+ -> 5Fe3+ + 5e-. Add the two halves: MnO4- + 8H+ + 5Fe2+ -> Mn2+ + 4H2O + 5Fe3+. This is the balanced equation.

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.

Hard water is water that does not give lather easily with soap because it contains dissolved calcium and magnesium salts. Temporary hardness is due to bicarbonates of Ca and Mg (Ca(HCO3)2, Mg(HCO3)2) and can be removed by boiling or by Clark's method (adding calculated slaked lime). Permanent hardness is due to chlorides and sulphates of Ca and Mg and is removed by adding washing soda (Na2CO3) or by the ion-exchange (permutit/resin) method.

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.

Solubility of hydroxides increases down the group (Mg(OH)2 is sparingly soluble while Ba(OH)2 is fairly soluble) because lattice energy decreases faster than hydration energy. Solubility of sulphates decreases down the group (MgSO4 is soluble but BaSO4 is almost insoluble) because hydration energy decreases faster for the larger ions. Flame colours: calcium gives brick-red, strontium gives crimson-red and barium gives apple-green.