Electrochemical cells and redox

In a nutshell

Redox reactions are where reduction and oxidation occurs. Reduction is the gain of electrons whereas oxidation is the loss of electrons. Redox occurs in electrochemical cells at the anode and cathode, producing electricity. A voltmeter measures voltage difference of electrodes to calculate cell potential.

Redox reactions

A redox reaction is a reaction where both reduction and oxidation occurs, transferring electrons.

Oxidation is the loss of electrons and the oxidation number of an element will increase when it has been oxidised. Reduction is the gain of electrons and the oxidation number of an element will decrease when it has been reduced. You can determine if an element has been oxidised or reduced by half reactions.

Tip: You can remember the difference between these two terms by the acronym 'OILRIG' which stands for Oxidation Is Loss and Reduction Is Gain.

Example

$Na^++\frac{1}{2}Cl_2\rightarrow NaCl$

Half reactions:

$Na\rightarrow Na^+ +e^-$

$\frac{1}{2}Cl_2+e^-\rightarrow Cl^-$

$Na$ has been oxidised as there is a loss of electrons whereas $Cl_2$ has been reduced as there is a gain of electrons. The oxidation number of $Na$ has increased from $0$ to $+1$ and the oxidation number of $Cl_2$ has decreased from $0$ to $-1$ .

Element type	redox	Oxidation number	Example
$s$ -block	Oxidation	Oxidation number = group number	$Mg\rightarrow Mg^{2+}+2e^-$
$p$ -block elements (groups $3-4$ )	Oxidation	Oxidation number = group number	$Al\rightarrow Al^{3+}+ 3e^-$
$p$ -block elements (groups $5-8$ )	Reduction	Oxidation number = group number $-8$	$\frac{1}{2}O_2+2e^-\rightarrow O^{2-}$
$d$ -block	Oxidation	Variable oxidation numbers	$Cu\rightarrow Cu^{+}+ e^-$ $Cu\rightarrow Cu^{2+}+ 2e^-$ $Cu\rightarrow Cu^{3+}+ 3e^-$

Note: These are general rules for predicting oxidation numbers and there are exceptions. Groups $3-4$ in the $p$ -block can be referred to as metals and groups $5-7$ can be referred to as non-metals.

Electrochemical cells

Electrochemical cells exhibit redox reactions as there are two different metal electrodes (positive and negative) immersed in a solution of ions and connected by an electrical circuit. This allows electrons to flow between the electrodes which produces electricity.

At the anode

This is the negative electrode where oxidation (loss of $e^-$ ) will take place and electrons will flow from this electrode. The more reactive metal will typically be the anode as it loses electrons so has become oxidised.

At the cathode

This is the positive electrode where reduction (gain of $e^-$ ) will take place and electrons flow to this electrode. The less reactive metal will typically be the cathode as it loses electrons less easily so will gain electrons and be reduced.

Example

Chemistry; Redox II; KS5 Year 12; Electrochemical cells and redox

Zinc is labelled as the anode as it is more reactive compared to copper so will undergo oxidation. Copper is labelled as the cathode as it is less reactive than zinc so undergoes reduction. You can determine whether the metals are anodes and cathodes by half reactions that are occurring in the electrochemical cell.

Half reactions:

Anode:

$Zn\rightarrow Zn^{2+}+2e^-$

Cathode:

$Cu^{2+}+2e^-\rightarrow Cu$

Note: Half reactions can be shown with reversible arrows, however as one of the metals will oxidise more readily, they are drawn like this to show if the metal is oxidised or reduced.

Each electrode is immersed in an aqueous salt solution and connected to a wire to allow the flow of electrodes from the anode to the cathode. A salt bridge made out of filter paper soaked in a salt solution such as $KCl$ allows flow of ions to counteract the flow of electrons in the circuit.

A voltmeter is used within the electrical circuit to measure the voltage $V$ between the half cells (electrodes). The calculated difference in voltage between the half cells is referred to as the cell potential or the electromotive force (EMF) and shows the flow of electrons in an electrochemical cell.

Electrolysis in an electrochemical cell

procedure

1.	The reactive anode will oxidise to form positive ions and lose electrons
2.	The electrons will travel around the circuit which generates electricity
3.	When the electrons reach the cathode, positive ions in the solution will gain electrons
4.	As positive ions both build up and are reduced in the solutions, a salt bridge is used so that positive and negative ions can flow to either solution to maintain charge neutrality