Redox Reactions and Electrode Potentials
Redox Reactions and Electrode Potentials
Lesson duration is: 50 minItems in this lesson
Redox Reactions and Electrode Potentials
Open with the big idea that redox is about electron transfer and that electrode potentials help us predict whether a redox change is feasible.
Learning objective
I can balance redox equations and use standard electrode potentials to predict whether a reaction is feasible.
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What words do you link with redox reactions?
Use this to surface prior knowledge such as electrons, oxidation, reduction, OIL RIG, oxidation numbers, anode and cathode. Invite students to notice which ideas feel secure and which need refreshing.
Lesson overview
This KS5 chemistry lesson explores Redox Reactions and Electrode Potentials. Students balance redox equations and use standard electrode potentials to predict feasibility and construct electrochemical cells. It is suitable for A level classroom teaching, independent study or exam revision.
This text is included verbatim as requested. Keep it on its own near the beginning of the lesson.
What makes a reaction redox?
A redox reaction is a reaction in which electrons are transferred.
Oxidation = loss of electrons
Reduction = gain of electrons
Oxidation and reduction always happen together
One species is oxidised while another is reduced
Reinforce OIL RIG if helpful, but keep the formal definitions as electron loss and gain.
Oxidation numbers track electron change
Oxidation numbers help us see which species are oxidised or reduced.
Increase in oxidation number → oxidation
Decrease in oxidation number → reduction
Useful when electrons are not shown explicitly
You may want to remind students of common rules: elements as 0, Group 1 as +1, oxygen usually -2, hydrogen usually +1.
Spotting redox in equations
Example: Zn + Cu²⁺ → Zn²⁺ + Cu
Zn: 0 → +2, so zinc is oxidised
Cu: +2 → 0, so copper ions are reduced
Zn is the reducing agent
Cu²⁺ is the oxidising agent
Students often confuse the agent names. Emphasise that the reducing agent causes reduction by being oxidised itself.
Writing half-equations
Half-equations show electron transfer clearly.
Oxidation half-equation: electrons on the right
Reduction half-equation: electrons on the left
Atoms and charge must both balance
Example:
Zn → Zn²⁺ + 2e⁻
Cu²⁺ + 2e⁻ → Cu
If needed, model how the electrons account for charge balance, not atom balance alone.
Balancing redox equations
A standard method is:
Split the equation into half-equations
Balance atoms other than O and H
Balance O with H₂O and H with H⁺ when needed
Balance charge with electrons
Multiply half-equations so electron numbers match
Add and cancel common species
This is the A level method for acidic conditions. If alkaline conditions come up, note that extra steps with OH⁻ may be needed, but do not overload unless required by the course context.
Worked redox balance
For Fe²⁺ + MnO₄⁻ → Fe³⁺ + Mn²⁺ in acidic solution:
Fe²⁺ → Fe³⁺ + e⁻
MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O
Multiply the iron half-equation by 5, then add:
5Fe²⁺ + MnO₄⁻ + 8H⁺ → 5Fe³⁺ + Mn²⁺ + 4H₂O
Point out why 8H+ and 4H2O appear: they balance hydrogen and oxygen in acidic solution.
In Zn + Cu²⁺ → Zn²⁺ + Cu, which species is the oxidising agent?
Zn
Cu
Zn²⁺
Cu²⁺
Cu²⁺ is reduced, so it acts as the oxidising agent. Distractors target confusion between the reacting metal, the product metal and the ion formed.
Complete the redox summary
A redox reaction involves transfer of electrons. In a half-equation, oxidation has electrons on the right and reduction has electrons on the left. When balancing in acidic solution, oxygen is balanced with H2O.
This checks the core method from slides 3 to 8: electron transfer, half-equation structure and balancing oxygen in acidic conditions.
Why multiply one half-equation before adding?
Expected elements: oxidation is loss of electrons or rise in oxidation number; reduction is gain of electrons or fall in oxidation number; electrons must cancel after multiplying the half-equations.
Standard electrode potentials
A standard electrode potential, E°, measures the tendency of a species to be reduced.
Measured under standard conditions
Written as a reduction half-equation
Units are volts (V)
More positive E° means reduction is more favourable
Standard conditions at A level are usually 298 K, solutions at 1 mol dm⁻³ and gases at 100 kPa. Mention only if useful for context.
Why values are measured relative to hydrogen
Electrode potentials cannot be measured on their own.
They are measured relative to the standard hydrogen electrode
The standard hydrogen electrode is defined as 0.00 V
Other half-cells are compared against this reference
Students do not need every construction detail unless your specification requires it. The key point is that electrode potentials are relative measurements.
Reading E° values
Compare these reduction half-equations:
Cu²⁺ + 2e⁻ → Cu E° = +0.34 V
Zn²⁺ + 2e⁻ → Zn E° = -0.76 V
Copper ions are more easily reduced than zinc ions because +0.34 V is more positive than -0.76 V.
Connect this back to the earlier Zn/Cu reaction so students see the same chemistry described in two ways.
Why is the standard hydrogen electrode used?
It gives 1.00 V
It is most reactive
It is a zero reference
It speeds reactions
Electrode potentials are relative measurements, so the standard hydrogen electrode provides the agreed 0.00 V reference point.
Complete the E° paragraph
A standard electrode potential, E°, is a measure of the tendency of a species to be reduced. Values are written as reduction half-equations and measured relative to the standard hydrogen electrode, which is assigned 0.00 V.
This anchors the definition of E° and the role of the standard hydrogen electrode from slides 9 to 11.
Building an electrochemical cell
An electrochemical cell turns a feasible redox reaction into electrical energy.
Oxidation happens at the anode
Reduction happens at the cathode
Electrons flow from anode to cathode in the external circuit
A salt bridge completes the circuit by allowing ion movement
Students often mix up electron flow and conventional current. Focus here on electron flow, which is most useful for redox reasoning.
The Zn/Cu cell
For a cell made from zinc and copper half-cells:
Zn is oxidised at the anode
Cu²⁺ is reduced at the cathode
Electrons flow from zinc to copper
Zinc is the more reactive metal
This is the classic Daniell cell. If useful, mention that zinc dissolves and copper is deposited.
Calculating E°cell
Use reduction potentials from the data table:
E°cell = E°(reduction at cathode) − E°(reduction at anode)
For Zn/Cu:
E°cell = (+0.34) − (−0.76) = +1.10 V
A positive E°cell suggests the reaction is feasible under standard conditions.
Stress that students should use the tabulated reduction potentials directly, rather than changing signs and then adding unless they are very secure.
Using E° to predict feasibility
A more positive E°cell means a greater tendency for the reaction to occur.
Positive E°cell: reaction is feasible under standard conditions
Negative E°cell: reaction is not feasible under standard conditions
This is a prediction, not a guarantee of reaction rate
Important distinction: thermodynamic feasibility is not the same as fast reaction rate. Students often blur these ideas.
Example feasibility prediction
Can chlorine oxidise bromide ions?
Cl₂ + 2e⁻ → 2Cl⁻ E° = +1.36 V
Br₂ + 2e⁻ → 2Br⁻ E° = +1.07 V
If chlorine is reduced and bromide is oxidised:
E°cell = +1.36 − +1.07 = +0.29 V
So chlorine can oxidise bromide ions under standard conditions.
This links electrode potentials to halogen displacement, a familiar UK A level context.
In a Zn/Cu cell, electrons flow from the:
cathode to anode
copper to zinc
salt bridge to zinc
zinc to copper
Zinc is oxidised at the anode, so electrons leave zinc and travel to copper, where Cu²⁺ is reduced.
Complete the cell calculation
In an electrochemical cell, oxidation happens at the anode and reduction happens at the cathode. Cell potential is calculated using E°cell = E° at the cathode minus E° at the anode. A positive E°cell suggests the reaction is feasible under standard conditions.
This brings together structure of the cell, the E°cell formula and the feasibility rule from slides 12 to 16.
Why can chlorine oxidise bromide ions?
Expected elements: chlorine has the more positive reduction potential; chlorine is reduced while bromide is oxidised; E°cell is +0.29 V so the reaction is feasible under standard conditions.
Limits of E° predictions
Electrode potentials are useful, but conditions matter.
Values are for standard conditions only
Concentration changes can alter cell potential
Reactions may be feasible but still very slow
Some reactions need a catalyst or different conditions
Keep this concise. The aim is not to teach Nernst equation work here, only to stop overgeneralisation.
Applied A level task
To analyse an unfamiliar cell, students need to:
identify the two half-equations from data
choose which half-equation is reduction
reverse the other to show oxidation
calculate E°cell
state the direction of electron flow
decide whether the overall reaction is feasible
This slide sets up application without becoming an activity slide. You can turn it into a worked exam-style example in delivery.
Differentiation: support and stretch
Support
use oxidation numbers before writing half-equations
label anode, cathode, oxidation and reduction on every example
use the E°cell formula exactly as written
Stretch
justify feasibility in full chemical language
explain why a feasible reaction may still be slow
compare predictions with reactivity series ideas
This preserves the requested differentiation in a student-facing way without turning the slide into instructions for an activity.
A reaction has E°cell = +0.45 V. What can you conclude?
It is feasible
It is very fast
It is impossible
It needs no ions
This checks the limit made on slides 17 to 19: positive E°cell predicts feasibility under standard conditions, not reaction rate.
Which part still needs the most practice?
Half-equations
Balancing
E°cell
Electron flow
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What is one idea you will remember best?
Use responses to identify what students found memorable or useful, such as oxidation numbers, the E°cell formula or anode/cathode logic.
Which part of this lesson felt hardest?
Invite honest reflection. Likely themes: balancing in acidic conditions, choosing which half-equation is reversed, or separating feasibility from rate.
Glossary
redox reaction - a reaction in which oxidation and reduction happen together through electron transfer
oxidation - loss of electrons
reduction - gain of electrons
oxidising agent - a species that causes oxidation by being reduced
reducing agent - a species that causes reduction by being oxidised
half-equation - an equation showing electron loss or gain for one part of a redox reaction
oxidation number - a number showing the apparent charge used to track electron transfer
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Glossary
standard electrode potential (E°) - the potential of a half-cell measured relative to the standard hydrogen electrode under standard conditions
standard hydrogen electrode - the reference electrode assigned a potential of 0.00 V
electrochemical cell - a system that produces electrical energy from a feasible redox reaction
anode - the electrode where oxidation happens
cathode - the electrode where reduction happens
E°cell - the cell potential found from electrode potential data to predict feasibility
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