(i) a Brønsted–Lowry acid as a species that donates a proton and a Brønsted–Lowry base as a species that accepts a proton (see also 2.1.4 Acids)
(ii) use of the term conjugate acid–base pairs
(iii) monobasic, dibasic and tribasic acids
Learners should be able to identify acid–base pairs in equations for acid–base equilibria.
the role of H+ in the reactions of acids with metals and bases (including carbonates, metal oxides and alkalis), using ionic equations (see also 2.1.4c, 2.1.5e)
(i) the acid dissociation constant, Ka, for the extent of acid dissociation (see also 2.1.4b)
(ii) the relationship between Ka and pKa
use of the expression for pH as:
pH = –logH+
H+ = 10–pH
use of the expression for the ionic product of water, Kw
calculations of pH, or related quantities, for:
(i) strong monobasic acids (ii) strong bases, using Kw
calculations of pH, Ka or related quantities, for a weak monobasic acid using approximations
*Approximations for weak acid calculations:
HAequilibrium ≈ HAundissociated i.e. HA >> H+
H+equilibrium ≈ A–equilibrium i.e. negligible dissociation of H2O.
Learners will not be required to solve quadratic equations.*
limitations of using approximations to Ka related calculations for ‘stronger’ weak acids
Including reasons why HAequilibrium ≈ HAundissociated may no longer be valid.
a buffer solution as a system that minimises pH changes on addition of small amounts of an acid or a base
formation of a buffer solution from:
(i) a weak acid and a salt of the weak acid, e.g. CH3COOH/CH3COO–Na+ (ii) excess of a weak acid and a strong alkali, e.g. excess CH3COOH/NaOH
explanation of the role of the conjugate acid–base pair in an acid buffer solution, e.g. CH3COOH/CH3COO–, in the control of pH
calculation of the pH of a buffer solution, from the Ka value of a weak acid and the equilibrium concentrations of the conjugate acid–base pair; calculations of related quantities
explanation of the control of blood pH by the carbonic acid–hydrogencarbonate buffer system
The H2CO3/HCO3– buffer is present in blood plasma, maintaining a pH between 7.35 and 7.45.
pH titration curves for combinations of strong and weak acids with strong and weak bases, including:
(i) sketch and interpretation of their shapes (ii) explanation of the choice of suitable indicators, given the pH range of the indicator (iii) explanation of indicator colour changes in terms of equilibrium shift between the HA and A– forms of the indicator
No indicator is suitable for a weak acid/weak base titration. The indicator should be considered as a weak acid, HA.
the techniques and procedures used when measuring pH with a pH meter.
PAG11