Standard reduction potentials are reference values that show how strongly a chemical species tends to gain electrons under standard conditions. They let chemists compare oxidizing agents, reducing agents, and possible redox reactions using one organized table. This matters because batteries, corrosion, electroplating, and many industrial reactions all depend on electron transfer.
A standard reduction potential scale acts like an electrochemical ladder, with stronger oxidizing agents near the top and stronger reducing agents near the bottom.
Key Facts
- Standard conditions are 1 M solutes, 1 atm gases, pure solids or liquids, and 25 degrees C.
- All listed half-reactions are written as reductions, meaning electrons appear on the reactant side.
- E°cell = E°cathode - E°anode when both values are taken from a standard reduction potential table.
- A more positive E°red means a stronger tendency to be reduced and a stronger oxidizing agent.
- A more negative E°red means the reduced form is a stronger reducing agent and is more easily oxidized.
- A redox reaction is spontaneous under standard conditions if E°cell > 0.
Vocabulary
- Standard reduction potential
- The voltage for a reduction half-reaction measured relative to the standard hydrogen electrode under standard conditions.
- Cathode
- The electrode where reduction occurs in an electrochemical cell.
- Anode
- The electrode where oxidation occurs in an electrochemical cell.
- Oxidizing agent
- A substance that causes another substance to lose electrons by accepting electrons itself.
- Reducing agent
- A substance that causes another substance to gain electrons by donating electrons itself.
Common Mistakes to Avoid
- Adding the two half-cell potentials directly is wrong because the anode reduction potential must be subtracted when using E°cell = E°cathode - E°anode.
- Multiplying E° values when balancing electrons is wrong because electric potential is an intensive property and does not scale with the number of electrons.
- Choosing the lower E° value as the cathode is wrong for a spontaneous galvanic cell because the more positive reduction potential is reduced at the cathode.
- Forgetting to reverse the anode half-reaction is wrong because oxidation occurs at the anode even though the table lists all half-reactions as reductions.
Practice Questions
- 1 Given Cu2+ + 2e- -> Cu with E° = +0.34 V and Zn2+ + 2e- -> Zn with E° = -0.76 V, calculate E°cell for a Zn-Cu galvanic cell.
- 2 Given Ag+ + e- -> Ag with E° = +0.80 V and Fe2+ + 2e- -> Fe with E° = -0.44 V, identify the cathode and anode, then calculate E°cell.
- 3 Using a standard reduction potential table, explain why fluorine is a strong oxidizing agent while lithium metal is a strong reducing agent.