2.5.4

# Weak Acids & Bases (A2 Only)

Test yourself

## Weak Acids and Bases

Not all acids and bases fully dissociate in solution.

### Weak acids

• Weak acids do not fully ionise in solution.
• For a strong acid, like HCl, the equilibrium for the following equation lies very far to the right:
• HCl ⇋ H+ + Cl-
• For a weak acid (e.g. ethanoic acid), the equilibrium lies closer to the middle.

### Weak bases

• Weak bases behave similarly. They do not fully ionise in solution.
• So, for weak acids and bases, the concentration of H+ or OH- ions is not just the concentration of the acid or base.
• This makes it harder to calculate the concentrations in solution.

### Ka

• We define a new constant as Ka = $\frac{[H^+][A^-]}{[HA]}$
• HA refers to the un-ionised acid molecule, and A- refers to the anion remainder after loss of a proton.
• A large Ka value means that a lot of the acid ionises in solution.
• A small Ka value means that very little acid ionises in solution.

### pKa

• Just like pH, we use pKa because Ka values vary a lot.
• The definition of pKa is:
• pKa = −log10(Ka)

## Calculating the Concentration of Ions

We can use pKa along with pH to calculate the concentrations of varying ions in solution.

### Converting pKa to Ka

• The definition of pKa is:
• pKa = -log10(Ka)
• This can be rearranged into:
• Ka = 10-pKa

### Calculating [H+] from Ka

• When you dissolve a weak monoprotic acid in solution, [H+] is always equal to [A-].
• This means you can simplify Ka into:
• Ka = $\frac{[H^+]^2}{[HA]}$
• But, [HA] is the concentration of the acid added ([HA]0, minus the concentration of the ionised acid), so we can simplify further to:
• Ka = $\frac{[H^+]^2}{[HA]_0 - [H^+]}$
• If you know the values of Ka and [HA]0, you can calculate the concentration of H+ ions.

### Calculating [HA] from Ka and pH

• First, use the definition of pH to calculate the [H+].
• Next, use [H+] = [A-] to simplify the equation for Ka.
• Then, rearrange Ka to give [HA]:
• [HA] = $\frac{[H^+]^2}{K_a}$