IB syllabus > acids and bases (hl) > 18.4 

These notes were written for the old IB syllabus (2009). The new IB syllabus for first examinations 2016 can be accessed by clicking the link below.

IB syllabus for first examinations 2016

18.3 - Salt hydrolysis

18.3.1 - State and explain whether salts form acidic, alkaline or neutral aqueous solutions. Examples should include salts formed from the four possible combinations of strong and weak acids and bases. The effect of the charge density of the cations in groups 1, 2, 3 and d-block elements should also be considered, eg [Fe(H2O)6]3+ [Fe(OH)(H2O)5]2+ + H+ .

Salt hydrolysis

Salts are ionic, this means that they dissociate 100% in solution to give free aqueous ions

NaCl Na+ + Cl-

When both ions come from strong acid and bases they have no interactions with the ions formed by the dissociation of water (hydrogen and hydroxide ions), however if the ions come from weak acids and bases then they interact with the ions from water establishing equilibria.

Hence salts of ethanoic acid produce free ethanoate ions in solution that can interact with the hydrogen ions from the water.

Example: Sodium ethanoate solution

sodium ethanoate is 100% dissociated into ions:-


Sodium ions are from a strong base (sodium hydroxide) and do not interact with the water ions.

However, the ethanoate ions do interact with the hydrogen ions from the water equilibrium (H2O H+ + OH-)


We know that this last equilibrium lies to the side of the ethanoic acid (to the right), removing the hydrogen ions from the solution. As [H+] decreases the pH rises.

Hence a solution of sodium ethanoate has a pH greater than 7. We say that it is basic by hydrolysis.

Example: Ammonium chloride solution

Ammonium chloride dissociates 100% into ions in solution

NH4Cl NH4+ + Cl-

The ammonium ions interact with the hydroxide ions from the water removing them from the solution (equilibrium lies to the right)

NH4+ + OH- NH3 + H2O

This increases the concentration of hydrogen ions (as [H+] x [OH-] is constant) increasing the acidity of the solution (decrease pH)

We say that a solution of ammonium chloride is acidic by hydrolysis.

General rules

Salts involving ions with a high charge density


Ionic compounds dissociate 100% into ions in solution. These ions become solvated by the water molecules (the water molecules bond to the ions - this is one of the driving forces behind dissolution). The polar water molecules use the lone pairs on the oxygen of the water to coordinate to the positive metal ion. The ions are then enclosed by a 'cage' of water molecules usually in an octahedral arrangement.

Octahedral arrangement of water molecules around a positive ion (in this case a 3+ ion)

Charge density

This means the charge to size ratio of the ion.

charge density = ionic charge/ionic size

When the ion has a charge of 3+ or when it is very small this charge to size ratio is enough to polarise the water molecules surrounding the ion in solution. This results in a weakening of the O-H bonds within the water molecules allowing hydrogen ions to be released into the solution. Hence the solutions are acidic.

This effect is typified in aluminium salts (the aluminium ion has a charge of 3+) which are very acidic in solution

The aluminium hexaaqua ion

Aluminium ions are surrounded by six water molecules in an octahedral arrangement. This is called the aluminium hexaaqua ion. The high charge density of the aluminium ion polarises the water molecules and hydrogen ions are released into solution. The solution is so acidic that it releases carbon dioxide from sodium carbonate (this reaction is used in some fire extinguishers to produce foam in conjunction with detergent)

[Al(H2O)6]3+ [Al(OH)(H2O)5]2+ + H+

[Al(OH)(H2O)5]2+ [Al(OH)2(H2O)4]+ + H+

Transition metals

As the transition metals have variable oxidation states the ions that are formed with high charges (high oxidation state) also produce acidic solutions. A good examle of this is the Iron III ion. Salts such as iron III sulphate are acidic in solution.

[Fe(H2O)6]3+ [Fe(OH)(H2O)5]2+ + H+


Dissolution of ionic compounds in water