1. Acid-Base Equilibria L.O.: To understand the difference between strong and weak acids. To be able to carry out calculations on strong and weak acids.

2. STRONG ACIDS completely dissociate (split up) into ions in aqueous solution e.g. HCl H + (aq) + Cl¯(aq) MONOPROTIC1 replaceable H HNO 3 H + (aq) + NO 3 ¯(aq) H 2 SO 4 2H + (aq) + SO 4 2- (aq) DIPROTIC 2 replaceable H’s STRONG BASES completely dissociate into ions in aqueous solution e.g. NaOH Na + (aq) + OH¯(aq) STRONG ACIDS AND BASES

3. Weak acids partially dissociate into ions in aqueous solution e.g. ethanoic acid CH 3 COOH(aq) CH 3 COO¯(aq) + H + (aq) When a weak acid dissolves in water an equilibrium is set up HA(aq) + H 2 O(l) A¯(aq) + H 3 O + (aq) The water stabilises the ions To make calculations easier the dissociation can be written... HA(aq) A¯(aq) + H + (aq) The weaker the acid the less it dissociates and the more the equilibrium lies to the left. The relative strengths of acids can be expressed as K a or pK a values The dissociation constant for the weak acid HA is K a = [H + (aq)] [A¯(aq)] mol dm -3 [HA(aq)] WEAK ACIDS

4. Partially react with water to give ions in aqueous solution e.g. ammonia When a weak base dissolves in water an equilibrium is set up NH 3 (aq) + H 2 O (l) NH 4 + (aq) + OH¯ (aq) as in the case of acids it is more simply written NH 3 (aq) + H + (aq) NH 4 + (aq) The weaker the base the less it dissociates and the more the equilibrium lies to the left The relative strengths of bases can be expressed as K b or pK b values. WEAK BASES

5. Hydrogen ion concentration [H + (aq) ] Hydrogen ion concentration determines the acidity of a solution Hydroxide ion concentration determines the alkalinity For strong acids and bases the concentration of ions is very much larger than their weaker counterparts which only partially dissociate.

6. Hydrogen ion concentration [H + (aq) ] pH: hydrogen ion concentration can be converted to pH pH = - log 10 [H + (aq)] To convert pH into hydrogen ion concentration [H + (aq)] = antilog (-pH) pOH: an equivalent calculation for bases converts the hydroxide ion concentration to pOH pOH = - log 10 [OH¯(aq)] N.B: [ ] represents the concentration in mol dm -3 STRONGLY ACIDIC 10 0 10 -1 10 -2 10 -3 10 -4 10 -5 10 -6 10 -7 10 -8 10 -9 10 -10 10 -11 10 -12 10 -13 10 -14 10 -13 10 -12 10 -11 10 -10 10 -9 10 -8 10 -7 10 -6 10 -5 10 -4 10 -3 10 -2 10 -1 10 -0 01234567891011121314pH OH¯ [H + ] WEAKLY ACIDIC NEUTRAL STRONGLY ALKALINE WEAKLY ALKALINE Hydrogen ion concentration determines the acidity of a solution Hydroxide ion concentration determines the alkalinity For strong acids and bases the concentration of ions is very much larger than their weaker counterparts which only partially dissociate.

7. Ionic product of water - K w Despite being covalent, water conducts electricity to a very small extent. This is due to the slight ionisation...H 2 O(l) + H 2 O(l) H 3 O + (aq) + OH¯(aq) or, more simply H 2 O(l) H + (aq) + OH¯(aq) Applying the equilibrium law to the second equation gives K c = [H + (aq)] [OH¯(aq)] [ ] is the equilibrium concentration in mol dm -3 [H 2 O(l)] As the dissociation is small, the water concentration is very large compared with the dissociated ions. Therefore, any changes to its value are insignificant. So, its concentration can be regarded as constant. This “constant” is combined with (K c ) to get a new constant (K w ). K w = [H + (aq)] [OH¯(aq)] mol 2 dm -6 = 1 x 10 -14 mol 2 dm -6 (at 25°C) N.B: The constant is based on an equilibrium, therefore K w varies with temperature

8. Ionic product of water - K w The value of K w varies with temperature because it is based on an equilibrium. Temperature / °C 0 20 25 3060 K w / 1 x 10 -14 mol 2 dm -6 0.110.681.01.475.6 H + / x 10 -7 mol dm -3 0.330.821.01.272.37 pH7.487.08 76.926.63 What does this tell you about the equation H 2 O(l) H + (aq) + OH¯(aq) ? K w gets larger as the temperature increases so, the concentration of H + and OH¯ ions gets greater and so, the equilibrium moves to the right If the concentration of H + increases then pH must decrease, so pH decreases as the temperature increases Because the equation moves to the right as the temperature goes up, it must be an ENDOTHERMIC process (heat taken out of the system by going in the forward direction as temperature increases)

9. Relationship between pH and pOH Because H + and OH¯ ions are produced in equal amounts when water dissociates[H + ] = [OH¯] = 1 x 10 -7 mol dm -3 their concentrations will be the same. K w = [H + ] [OH¯] = 1 x 10 -14 mol 2 dm -6 take logs of both sideslog[H + ] + log[OH¯] = -14 multiply by minus- log[H + ] - log[OH¯] = 14 change to pH and pOHpH + pOH = 14 (at 25°C) N.B.As they are based on the position of equilibrium and that varies with temperature, the above values are only true if the temperature is 25°C (298K) Neutral solutions may be regarded as those where [H + ] = [OH¯]. Therefore a neutral solution is pH 7 only at a temperature of 25°C (298K) K w is constant for any aqueous solution at the stated temperature