Electrode potentials and electrochemical cells - AL only (3.1.11)
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Standard electrode potential () is the potential difference between a half-cell and the standard hydrogen electrode (SHE) under standard conditions.
These conditions are:
Under non-standard conditions, the measured electrode potential will be different.
Changes to concentration, pressure and temperature will alter the actual electrode potential of a half-cell. Standard conditions must be applied for a reliable comparison.
When the concentration of the solution in a half-cell is altered from (standard conditions) there is a shift in equilibria to maintain the equilibrium constant. This causes a change in the total amount of electrons moving through the cell.
The total cell potential is proportional to the current, and hence the number of electrons moving through the cell.
If there is an increased concentration of the species being reduced the total number of electrons in the cell will drop and the cell potential will also reduce.
Increasing the concentration of above will cause the equilibrium to shift to the right, the number of electrons in the cell to drop, and the cell potential to decrease.
Increasing the concentration of above will cause the equilibrium to shift to the left, the number of electrons in the cell to increase, and the cell potential to increase.
The standard electrode potential indicates how easily a species gains electrons (is reduced) compared to hydrogen.
A more positive value suggests a greater tendency to be reduced, making the substance a stronger oxidising agent.
Conversely, a more negative value indicates a greater tendency to lose electrons, making the substance a stronger reducing agent.
The electrochemical series is a list of standard electrode potentials () for various half-cells, arranged in order of their ability to gain electrons and therefore be reduced.
The electrochemical series is a key tool in predicting the redox reactions and understanding the reactivity of different species.
Standard cell potential () is calculated by combining the standard electrode potentials of the two half-cells involved in the reaction.
The must be positive for a reaction that is thermodynamically feasible.
Where:
- (anode) is the standard electrode potential at the positive electrode; that of the oxidation half-reaction.
- (cathode) is the standard electrode potential at the negative electrode; that of the reduction half-reaction.
The most positive half-cell from the electrochemical series will always be the reduction half-reaction.
Redox Reactions can be predicted using the reactivity series. The overall must be positive for the reaction to be thermodynamically favourable.
A species higher in the series (with a more positive reduction potential ) can be most easily reduced; it is an oxidising agent to a species lower in the series.
For example, copper () has an of and zinc () has an of . Copper ions can reduce to copper metal in order to oxidise zinc metal to zinc ions.
Question walkthrough
Predicting redox reactions
Use given electrode data to deduce the reactions occurring in cells
Cell notation is used to quickly convey the setup of an electrochemical cell.
When conveying electrochemical cells using cell notation:
- The anode is shown on the left and the cathode on the right.
- Single vertical lines | represent phase boundaries.
- Double vertical lines || represent the salt bridge.
- Commas represent separate species in the same half-cell and the same phase.
- Write from left to right, reactant followed by product.
- Include electrodes, even if they are inert.
Question walkthrough
Cell notation
Describe a cell using conventional representation
Understanding of the electrochemical series is important for:
- Battery design: Selection of metals for anodes and cathodes based on their relative values.
- Corrosion prediction: Metals lower in the series are more prone to oxidation (e.g., iron can be corroded by oxygen and water).
- Electroplating and extraction: Metals with specific values are chosen for plating or extraction processes.
The standard hydrogen electrode (SHE) is used as a reference electrode to measure the standard electrode potentials of other half-cells. The SHE is assigned a potential of by definition, under standard conditions.
- Hydrogen gas () at a pressure of .
- Platinum electrode coated with platinum black, which acts as an inert conductor.
- Acidic solution with ions at a concentration of
- Temperature
When a cell is set up to calculate the standard electrode potential:
- The half-cell of interest is connected to the standard hydrogen electrode (SHE).
- The two half-cells are connected by a salt bridge allowing ion transfer.
- An external cable with a voltmeter completes the circuit and measures the potential difference between the two half-cells.
- The potential difference recorded is the standard electrode potential () of the half-cell under investigation.
Measuring standard electrode potential of a zinc half-cell.
In this setup, the cell potential is measured for a solid electrode in contact with a solution containing its ions.
Electrode: A strip of zinc metal ().
Electrolyte: A solution of zinc sulfate () containing ions at concentration.
Procedure:
- Prepare the Half-Cell:
- Place the zinc strip in the solution.
- Ensure the solution concentration is to meet standard conditions.
- Connect to a standard hydrogen electrode (SHE):
- Connect the zinc half-cell to the SHE using a wire between the electrodes and a salt bridge between the electrolytes to maintain the ionic balance.
- Measure the cell potential:
- Use a voltmeter to measure the potential difference between the zinc electrode and the SHE.
The observed voltage is the standard electrode potential () of the zinc half-cell, relative to the SHE.
Measuring standard electrode potential of a chlorine half-cell.
In this setup, the cell potential is measured for a gaseous substance in contact with a solution containing its ions in the presence of an inert electrode.
Electrode: Platinum electrode (inert).
Electrolyte: A solution containing ions at concentration.
Procedure:
- Prepare the Half-Cell:
- Bubble chlorine gas into the aqueous solution containing ions. Ensure the solution concentration is to meet standard conditions.
- Immerse the platinum electrode in the solution.
- Connect to the SHE:
- Connect the chlorine half-cell to the SHE using a wire between the electrodes and a salt bridge between the electrolytes to maintain the ionic balance
- Measure the Cell Potential:
- Record the potential difference between the chlorine half-cell and the SHE.
The observed voltage is the standard electrode potential, , of the chlorine half-cell, relative to the SHE.
In this setup, the cell potential is measured for a redox system in which both species are in solution (homogenous half-cell), using an inert solid electrode to allow electron transfer.
Electrode: A platinum electrode is used because iron is not present as a solid metal, it provides a conductive surface for the redox reaction.
Electrolyte: An equimolar solution containing of and of ions.
Note that the total concentration of iron ions is .
- Prepare the half-cell
- Place a clean platinum electrode into the solution.
- Ensure an equimolar solution of and of ions to meet standard conditions.
- Connect to the standard hydrogen electrode (SHE)
- Connect the half-cell to the SHE using a wire between the electrodes and a salt bridge between the two electrolytes to maintain ionic balance.
- Measure the cell potential
- Use a high-resistance voltmeter to measure the potential difference between the half-cell and the SHE.
The observed voltage is the standard electrode potential of the half-cell relative to the SHE.
When designing an experimental set-up to measure half-cell potential, consider the state of your reacting species.
The half-cell to be quantified should be connected to a standard hydrogen electrode via a conductive wire and a salt bridge.
Electrochemical cells convert chemical potential energy into electrical energy through redox reactions. They are a source of electrical energy.
A spontaneous redox reaction causes electrons to flow from the anode (oxidation) to the cathode (reduction).
Electrochemical cells can be rechargeable or non-rechargeable.
The chemical reactions in rechargeable cells can be reversed by applying an external current, restoring the original reactants. Lithium-ion batteries are rechargeable.
When the reactants in non-rechargeable cells are used up, the cell stops producing electricity. The reactants cannot be regenerated by an external current. Alkaline batteries are an example of non-rechargeable cells.
You need to know the simplified electrode reactions for a Lithium-ion cell.
Overall cell potential for a Lithium-ion cell is high.
Lithium-ion batteries have a high energy density, meaning they can store a large amount of energy in a small volume. This makes them portable.
These batteries can be recharged multiple times, making them cost-effective over the long term.
There are also some risks associated with Lithium ion batteries:
- Lithium and other materials used in these batteries, like cobalt, are toxic and can pose environmental hazards if not disposed of properly.
- Lithium-ion batteries are prone to overheating which can lead to fires or explosions.
Where input of electricity causes a redox reaction to proceed in reverse, an electrochemical cell can be made rechargeable.
Rechargeable cells can be used multiple times, reducing the number of units required, and therefore the environmental impact of their manufacture and disposal.
Fuel cells, like other electrochemical cells, generate a voltage.
This voltage drives the flow of electrons (current), which can then power electrical devices.
Fuel cells are innovative devices that use chemical energy from a fuel, and oxygen, to produce a voltage.
Fuel cells are assembled with:
- an anode where the fuel (e.g. hydrogen) is oxidised.
- a cathode where oxygen is reduced.
- an electrolyte allowing ions to move between the anode and cathode, while keeping the electrodes separate.
- a wire allowing the flow of electrons.
These are not combustion reactions. There is no direct contact between the fuel and oxygen.
The electrode reactions in a fuel cell differ with of the electrolyte.
When the electrolyte is alkaline:
Anode (fuel electrode)
- Fuel, such as hydrogen gas, is oxidised, releasing electrons.
- Reaction:
Cathode (oxygen electrode)
- Oxygen gas is reduced, gaining electrons from the external circuit.
- Reaction:
Overall cell reaction
The overall reaction combines both half-equations to form water as the main product.
The electrode reactions in a fuel cell differ with of the electrolyte.
When the electrolyte is acidic:
Anode (fuel electrode)
- Fuel, such as hydrogen gas, is oxidised, releasing protons ( ions) and electrons ().
- Reaction:
Cathode (oxygen electrode)
- Oxygen is reduced by gaining electrons and reacting with protons to form water.
- Reaction:
Overall cell reaction
The overall reaction combines both half-equations to form water as the main product.
Portable electricity is a key component of modern society.
Factors to consider when selecting an appropriate energy source include:
- Recharability – Reversible electrochemical cells can be recharged and fuel cells and combustion engines can be topped up. Time taken to recharge/top up should also be considered.
- Waste: Irreversible electrochemical cells are single use.
- Type of fuel – methanol is easier to store and transport than hydrogen.
- Byproducts – typical combustion reactions are associated with the production of carbon dioxide, a greenhouse gas.
- Controllability – The reaction between hydrogen and oxygen is easier to control and therefore safer in a fuel cell than in a combustion reaction.