Compare oxidising and reducing agents using E°
Use standard reduction potentials to compare the electron-accepting and electron-donating tendencies of redox couples.
A Level; Cambridge International Chemistry 9701, 2028–2030 syllabus, Topic 24. Treat all tabulated electrode potentials as reduction potentials; state conditions, phases, signs and units. Written practice is self-checked
Before you begin
- Use AS redox half-equations, oxidation states, mole ratios, current and charge.
Work on paper. Try each question before revealing help, and explain why your method works.
Quick readiness check
Which species is the oxidising agent in a reduction half-equation Ox + e⁻ → Red?
What you will learn
- Identify the oxidising agent from the species reduced
- Identify the reducing agent from the species oxidised
- Relate relative agent strength to the signs and sizes of E° values
More positive reduction potential means stronger oxidising form
A couple with a more positive E° has an oxidised form with a greater tendency to gain electrons. That oxidised form is a stronger oxidising agent because it is reduced while causing another species to be oxidised.
The reduced form of a couple with a more negative E° is generally a stronger reducing agent: it more readily donates electrons and is oxidised. Compare species using their relevant tabulated couples, not by the ion charge alone.
Name the form that accepts or donates electrons
In Ox + ne⁻ ⇌ Red, Ox is the oxidising form and Red is the reducing form. A positive shift in E° strengthens the reduction tendency of Ox; a more negative E° makes Red a stronger electron donor in a comparison.
Use observed or predicted electron transfer to confirm the labels. The oxidising agent is itself reduced; the reducing agent is itself oxidised.
Pause and explain
Which oxidised form is generally the stronger oxidising agent?
Worked example
Which is the stronger oxidising agent, Cu²⁺ or Zn²⁺, given E°(Cu²⁺/Cu)=+0.34 V and E°(Zn²⁺/Zn)=−0.76 V?
Show the worked solution
- Both oxidising forms are the ions written on the left of reduction equations.
- Cu²⁺/Cu has the more positive E°.
- Cu²⁺ has the greater tendency to gain electrons.
Answer Cu²⁺ is the stronger oxidising agent because the Cu²⁺/Cu reduction potential is more positive.
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Test this topic →From guided practice to a new situation
Write your answer and reasoning first. Use a hint only if you are stuck. The model solution is for self-checking; your written work is not automatically marked.
Given E°(Cl₂/Cl⁻)=+1.36 V and E°(I₂/I⁻)=+0.54 V, identify the stronger oxidising agent.
Give me a hint
Compare the reduction potentials of the oxidised forms.
Compare my reasoning
- Both equations are reductions to halide ions.
- Cl₂/Cl⁻ is more positive.
- Cl₂ has the greater tendency to accept electrons.
Cl₂ is the stronger oxidising agent because its reduction potential is more positive.
Look for these in your work
- Couples compared
- Potential ordering correct
- Oxidising form identified
Given E°(Mg²⁺/Mg)=−2.37 V and E°(Ag⁺/Ag)=+0.80 V, which reduced form is the stronger reducing agent?
Give me a hint
Compare which reduced form is more readily oxidised in a cell comparison.
Compare my reasoning
- Mg²⁺/Mg has the more negative reduction potential.
- The reduced form Mg is more ready to donate electrons.
- Mg is the stronger reducing agent.
Mg(s) is the stronger reducing agent because the Mg²⁺/Mg reduction potential is much more negative than Ag⁺/Ag.
Look for these in your work
- More negative potential identified
- Reduced species selected
- Reducing-agent role correct
In Zn + Cu²⁺ → Zn²⁺ + Cu, identify the oxidising and reducing agents and state which one is reduced.
Give me a hint
The oxidising agent accepts electrons; the reducing agent donates them.
Compare my reasoning
- Cu²⁺ gains electrons to form Cu.
- Zn loses electrons to form Zn²⁺.
- Cu²⁺ is oxidising agent; Zn is reducing agent.
Cu²⁺ is the oxidising agent and is reduced. Zn is the reducing agent and is oxidised.
Look for these in your work
- Electron directions correct
- Agent roles correct
- Own changes stated
Common mistakes
- Calling the species oxidised the oxidising agent
- Comparing reduced forms when asked for oxidising strength
- Ignoring which side of a redox couple is the oxidised form
- Equating a more positive E° with a stronger reducing agent
What can you explain now?
Which species is the oxidising agent in a reduction half-equation Ox + e⁻ → Red?
Compare with the explanation
Ox
Ox accepts electrons and is reduced.
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Make it stick
Which form in Ox/Red accepts electrons, and how does E° compare oxidising and reducing strength?
If you needed the solution, close it and solve the problem again from a blank page. A correct answer is stronger when you can explain the reason for each step.