L3 - Electrolysis of Aqueous Solutions

objectives

Understandings

• When aqueous solutions are electrolysed, water can be oxidized to oxygen at the anode and reduced to hydrogen at the cathode.

Applications and Skills

• Explanation of the products formed during the electrolysis of aqueous solutions.

Guidance

• Electrolytic processes to be covered in theory should include the electrolysis of aqueous solutions (eg sodium chloride, copper(II) sulfate etc) and water using both inert platinum or graphite electrodes and copper electrodes. Explanations should refer to Eº values, nature of the electrode and concentration of the electrolyte.

Electrolysis of Water (also dilute Sulphuric Acid)

• It is possible to split water into hydrogen and oxygen gas by electrolysis. Acids are typically added to increase the concentration of hydrogen ion in solution (Figure 2). The reactions are

• Note that the sulfuric acid is not consumed and that the volume of hydrogen gas produced is twice the volume of oxygen gas produced.
• The minimum applied voltage is 1.229 V.

Electrolysis of Aqueous Sodium Chloride

• The electrolysis of aqueous sodium chloride is the more common example of electrolysis because more than one species can be oxidized and reduced.
• Considering the anode first, the possible reactions are

• These values suggest that water should be oxidized at the anode because a smaller potential would be needed—using reaction (ii) for the oxidation would give a less-negative cell potential.
• When the experiment is run, it turns out chlorine, not oxygen, is produced at the anode.
• The unexpected process is so common in electrochemistry that it has been given the name overpotential.
• The overpotential is the difference between the theoretical cell voltage and the actual voltage that is necessary to cause electrolysis.
• It turns out that the overpotential for oxygen is rather high and effectively makes the reduction potential more positive.
• As a result, under normal conditions, chlorine gas is what actually forms at the anode.
• Now consider the cathode. Three reductions could occur:

• Reaction (v) is ruled out because it has such a negative reduction potential.
• Under standard state conditions, reaction (iii) would be preferred to reaction (iv).
• However, the pH of a sodium chloride solution is 7, so the concentration of hydrogen ions is only 1× 10−7M.
• At such low concentrations, reaction (iii) is unlikely and reaction (iv) occurs. The overall reaction is then

• As the reaction proceeds, hydroxide ions replace chloride ions in solution.
• Thus, sodium hydroxide can be obtained by evaporating the water after the electrolysis is complete.
• Sodium hydroxide is valuable in its own right and is used for things like oven cleaner, drain opener, and in the production of paper, fabrics, and soap.

Electrolysis of Aqueous Copper Sulphate

• The ions present in the solution are: copper ions, Cu2+; sulfate ions,SO42-; hydrogen ions, H+; and hydroxide ions, OH-.

At the cathode

The positive ions are attracted to the negative cathode. There is competition between the copper ions and the hydrogen ions. As the hydrogen ion | hydrogen redox equilibrium appears higher in the electrochemical series than the copper ion | copper equilibrium, then the copper ions are preferentially reduced and copper metal is deposited at the electrode (a pink layer is observed)

Cu2+ + 2e --> Cu

At the anode

There is competition between the negative ions at the positive anode. The sulfate ions compete with the hydroxide ions to release their electrons to the anode. The hydroxide ions are much better reducing agents and are preferentially released AS OXYGEN GAS and water

4OH- - 4e --> 2H2O + O2

Ions remaining in solution

The ions that are removed from the solution, then, are the copper ions and the hydroxide ions. This means that the hydrogen ions and the sulfate ions remain in the solution - i.e sulfuric acid is also produced.