Standard reduction potentials

1. Understanding Reduction Potentials

• $E°$(Standard Electrode Potential) represents the potential difference of a half-reaction under standard conditions ($1M$ concentration, $25°C,$ $1$ atm pressure).
• Fluorine $(F₂)$ has the highest reduction potential, meaning it is the strongest oxidising agent.
• Lithium $(Li)$has the lowest reduction potential, meaning it is the strongest reducing agent.

2. Rules for Using the Standard Reduction Table

• A strong reducing agent will displace a weaker reducing agent from its compound.
• A redox reaction occurs when a reducing agent reacts with an oxidising agent.
• Always start with the reduction half-reaction (where electrons are gained).
• Oxidation and reduction reactions must be balanced by:
  • Ensuring the same number of electrons are transferred.
  • Multiplying the half-reactions by appropriate factors.

3. Combining Half-Reactions

• Remove common ions or molecules from both sides of the reaction.

• Add the E° values of the two half-reactions to get the total cell potential: $E°_{\text{cell}} = E°_{\text{cathode}} - E°_{\text{anode}}$

  • $E°$ (cathode): Reduction potential of the species being reduced.
  • $E°$ (anode): Reduction potential of the species being oxidised.

4. Predicting Spontaneity of a Reaction

• If $E°_{\text{cell}} > 0 →$ Reaction is spontaneous.
• If $E°_{\text{cell}} < 0 →$ Reaction is non-spontaneous.

5. Key Takeaways

• Higher $E°$ values indicate stronger oxidising agents.
• Lower $E°$ values indicate stronger reducing agents.
• Spontaneous reactions have a positive $E°$ value.
• Redox reactions must be balanced properly for accurate calculations.