An acid is classified as either

• strong
• weak

Strong acids

A strong acid is defined as one which dissociates completely in water to form H₃O⁺ and A^–. Therefore, it needs to be a better proton donor than H₃O⁺. For example, HCl is a strong acid.

Calculating the pH

Ignoring the auto-ionisation of water, because a strong acid dissociates completely then the concentration of H₃O⁺ will equal the acid concentration. Therefore, a solution of 1.0M HCl will be 1.0M in terms of both H₃O⁺ ions and Cl^– ions (as long as the auto-ionisation of water is ignored):

HCl(aq) + H₂O(l) ? H₃O⁺(aq) + Cl^–(aq)

1.0M 1.0M

Therefore, the pH = 0

Theoretically, pH possesses no minimum value as the molarity of an acid is able to increase indefinitely. However, in practice, it is rare for the pH to be less than -0.5 because even strong acids are not able to dissociate completely at high concentrations. This is because there simply is not enough water. Therefore, a strong acid can only be said to have completely dissociated if the solution is quite dilute (less than 0.1M, for example).

If you dilute an acid tenfold the pH level will increase by one unit.

Weak acids

A weak acid is defined as one which only dissociates partially in water. Equilibrium is reached by:

HA(aq) + H₂O(l) ? H₃O⁺(aq) + A^–(aq)

C-x x x

Therefore, a weak acid needs to be a worse proton donor than H₃O⁺. For example, CO₂ is a weak acid.

For a weak acid the equilibrium expression is:

K_c = ([H₃O⁺][A^–]) / ([HA][H₂O])

Due to the fact that water is present in a large amount, this equilibrium does not significantly affect its concentration and, therefore, can be assumed as constant. Therefore, the equilibrium expression can be written instead as:

K_c [H₂O] = K_a = ([H₃O⁺][A^–]) / [HA]

K_a is called the acid dissociation constant for the acid and has the units of moldm^-3. A lot of the time it is quoted as:

pK_a = -log₁₀K_a

The larger the value of K_a, the more the acid dissociates with its ions and so the stronger the acid is.

Calculating the pH

H₃O⁺ concentration can be calculated using:

HA(aq) + H₂O(l) ? H₃O⁺(aq) + A^–(aq)

C-x x x

K_a = ([H₃O⁺][A^–]) / [HA]

If you ignore the H₃O⁺ due to the auto-ionisation of water, the concentrations of H₃O⁺ and A^– can be assumed as equal.

When the degree of dissociation is assumed small, the molarity of HA can be assumed to be equal to that of the solution: C-x = C

K_a = [H₃O⁺]² / [HA]

Therefore: [H₃O⁺]² = K_a [HA] and [H₃O⁺] = ? K_a [HA]

When a weak acid is diluted tenfold the pH does not increase by one unit; it only increase by 0.5 units. What this suggests is that a weak acid dissociates to a larger extent in order to compensate for the dilution.

Therefore, the larger the dilution of the solution the larger the extent to which the acid dissociates into ions. This follows the principle of Le Chatelier:

HA(aq) + H₂O(l) ? H₃O⁺(aq) + A^–(aq)

As the concentration of water increases so the equilibrium shifts to the right and the degree of dissociation increases. Strong acids on the other hand are fully dissociated even at relatively high concentrations and so do not respond like weak acids to dilution.

Acidic salts

The salts of weak bases are able to act like acids. For example, NH₄⁺(aq):

NH₄⁺(aq) + H₂O(l) ? NH₃ (aq) + H₃O⁺(aq)

Ka = ([NH₃][H₃O⁺]) / [NH₄⁺]

Dilute and concentrated solutions

The dilute solution of an acid has a low molarity. A concentrated solution of an acid, on the other hand, has a high molarity.

However, the concentration of a solution is not necessarily a good indicator of how acidic a solution is because the concentrated solution of a weak acid might be less acidic than the more dilute solution of a stronger acid.

Nevertheless, it is not always the case that a stronger acid will provide a more acidic solution. A lot of the time the concentrated solution of a weaker acid will be more acidic than the more dilute solution of a stronger acid.

Therefore, a solution’s acidity is dependent on two factors:

• the concentration of the solution
• the strength of the acid

Very dilute solutions

In the above examples the auto-ionisation has been ignored because water only ionises very slightly (in pure water [H₃O⁺] = 1 x 10^-7 M).

However, if the solution is very dilute then the auto-ionisation of water becomes significant and, therefore, must be taken into consideration. In general, this is the case if the solution’s concentration is below 1 x 10^-6 M.

Polybasic acids

An acid which can lose more than one proton is called polybasic. For example, H₂SO₄.

• An acid which can lose two protons in solution is called dibasic.
• An acid which can lose three protons is solution is called tribasic.

In aqueous solution a polybasic acid will dissociate more than once and will form more than one salt.

• A dibasic acid forms two salts.
• A tribasic acid forms three salts.

The number of protons an acid will lose can be calculated through its reaction with an alkali:

H_xA(aq) + xNaOH(aq) ? Na_xA(aq) + _xH₂O(l)

• Monoprotic acids react with sodium hydroxide in a 1:1 ratio.
• Diprotic acids react with sodium hydroxide in a 2:1 ratio.
• Triprotic acids react with sodium hydroxide in a 3:1 ratio.

With the majority of acids all the hydrogen available is not lost. Therefore, the amount of replaceable hydrogens needs to be experimentally determined.