Acid-Base equilibria

Bronsted-Lowry definition of acid and base

• Acid: a proton donor.HCl (aq) → H⁺ (aq) + Cl^- (aq)

• Base: a proton acceptor OH^- (aq) + H⁺ (aq) → H₂O (l)

what are conjugate acid-base pairs

• Conjugate acid-base pairs are a pair of reactants and products that are linked to each other by the transfer of a proton

Difference between strong and weak acids?

• Strong acids fully dissociate in water.The position of the equilibrium is so far over to the right that you can represent the reaction as an irreversible reaction

• Weak acids partially dissociate in water. The position of the equilibrium is more over to the left and an equilibrium is established. Reversible arrow.

pH definition

pH = -log[H⁺]

[H⁺] = 10^-pH

what are Ka expressions

• For weak acids as there is an equilibrium we can write an equilibrium constant expression for the reaction

  

what is pKa

• The range of values of K_a is very large and for weak acids, the values themselves are very small numbers

• For this reason it is easier to work with another term called pK_a

• The pK_a  is the negative log of the K_a value, so the concept is analogous to converting [H⁺] into pH values

pK_a = -logK_a

assumptions made when calculating Ka

• (H+) = (A-)

• (HA) doesnt change on dissociation

assumptions made when writing the equilibrium expression for weak acids

• The concentration of hydrogen ions due to the ionisation of water is negligible

• The value of K_a indicates the extent of dissociation

  • The higher the value of K_a the more dissociated the acid and the stronger it is

  • The lower the value of K_a the weaker the acid

pH calculations involving strong acids

• Strong acids are completely ionised in solution

HA (aq) → H⁺ (aq) + A^- (aq)

• Therefore, the concentration of hydrogen ions, H⁺, is equal to the concentration of acid, HA

• The number of hydrogen ions formed from the ionisation of water is very small relative to the [H⁺] due to ionisation of the strong acid and can therefore be neglected

• The total [H⁺] is therefore the same as the [HA]

• so conc of acid= conc of H+

how to calculate pH of weak acids?

• the equilibrium concentration of [H⁺] and [A^-] will be the same since one molecule of HA dissociates into one of each ion

• This means you can simplify and re-arrange the expression to

K_a x [HA] = [H⁺]²

[H⁺]² = K_a x [HA] 

• Taking the square roots of each side

[H⁺] = √(K_a x [HA])

• Then take the negative logs

pH = -log[H⁺] = -log√(K_a x [HA]

ionisation of water

water ionises slightly- the equilibrium lies to the left

water behaves as both an acid and a base

equilibrium constant/ionic product of water

• the concentration of water is constant so:

• Kw= (H+)(OH-)

• at 298K, Kw=  1 x 10^-14 mol² dm^-6

how to calculate pKw

pK_w = -logK_w

what is Kw dependent on

• Kw is temeprature dependent

• the ionisation of water is endothermic. An increase in temperature shifts the equilibrium to the right

how to calcualte pH of strong base

• a strong base fully dissociates in water

• BOH (aq) → B⁺ (aq) + OH^- (aq)

  • Therefore, the concentration of hydroxide ions [OH^-] is equal to the concentration of base [BOH]

    • Even strong alkalis have small amounts of H⁺ in solution which is due to the ionisation of water

  • The concentration of OH^- in solution can be used to calculate the pH using the ionic product of water

  • Once the [H⁺] has been determined, the pH of the strong alkali can be founding using pH = -log[H⁺]

  • 

finding Ka for a weak acid, experimentally

• weak acids only partially dissociate, in water

• when a weak acid reacts with an alkali it all reacts

what determines pH of salt

• not all salts form at pH 7

• it depends on the relative strengths of the acid and base used to form the salt

• eg with equivalent amounts reacting (equal conc and volume-all monoprotic (monobasic))

• Strong acid+ strong base: pH of resultant solution if pH7

• Strong acid+ weak base: pH of the solution is <pH 7

• Weak acid + strong base: pH of the solution is >pH 7

• Weak acid+ weak base: the pH of the resultant solution depends on the extent of dissociation of acid and base. Reversible reaction

• If Ka=Kb salt =pH7

• If Ka> Kb <pH7

• If Ka < Kb >pH7

what are serial dilutions

this means successive dilution by the same factor eg by a factor of 10 (10x)

effect of dilution on the pH of aqueous solutions of acids

• For strong acids a decrease in concentration by a facor of 10 leads to an increase in pH by 1

• For weak acids, a decrease in concentration of a factor of 10 leads to an increase in pH of 0.5

• as dilution increases, dissociation increases