Standard electrode potentials

1. Definition

• Standard Electrode Potential $(E°)$ is the potential difference between a half-cell and a standard hydrogen electrode under standard conditions.
• It is measured in volts $(V)$ and determines whether a species is a strong oxidising agent or a strong reducing agent.

2. Standard Conditions

• Temperature: 25°C $(298 K)$
• Pressure: 1 atm
• Concentration: 1 mol·dm⁻³ of electrolyte solution

3. Uses of Standard Electrode Potentials

• Identify oxidising and reducing agents.
• Predict spontaneity of redox reactions.
• Calculate cell potential $E°_{\text{cell}}$ using:

$E°_{\text{cell}} = E°_{\text{cathode}} - E°_{\text{anode}}$

or

$E°_{\text{cell}} = E°_{\text{reduction}} + E°_{\text{oxidation}}$

4. Table of Standard Electrode Potentials

• The table lists half-reactions and their $E°$ values.
• Stronger reducing agents (metals) have more negative $E°$ values and are more easily oxidised.
• Stronger oxidising agents (non-metals) have more positive $E°$ values and are more easily reduced.

Example: Zinc and Copper Half-Reactions

Reaction	E° (V)
$Zn^{2+} + 2e^• \rightarrow Zn$	-0.76 V
$Cu^{2+} + 2e^• \rightarrow Cu$	+0.34 V

• Zinc $(Zn)$ is a stronger reducing agent (more negative $E°$).
• Copper (Cu²⁺) is a stronger oxidising agent (more positive $E°$).
• Since $E°_{\text{cell}} = 0.34 - (-0.76) = :highlight[1.1 V],$ the reaction is spontaneous.

5. Predicting Spontaneous Reactions

• A reaction is spontaneous if $E°_{\text{cell}} > 0$.
• Example: $Zn$ and $Cu$ cell
  • $Zn$ is oxidised (loses electrons).
  • $Cu²⁺$ is reduced (gains electrons).
  • Reaction occurs because Zn is a stronger reducing agent than $Cu.$
• A reaction is not spontaneous if $E°_{\text{cell}} < 0.$
• Example: $Cu$ and $Ag$ cell
  • $Cu$ cannot reduce $Ag⁺$ because $Cu$ is a weaker reducing agent than $Ag.$
  • No reaction occurs.

6. Applications

• Electrochemical cells: Used to design Galvanic cells and batteries.
• Metal corrosion protection: Predicting metal reactivity.
• Industrial electrolysis: Identifying best electrodes.

7. Key Takeaways

• More negative $E° →$ Stronger reducing agent → Easily oxidised.
• More positive $E° →$ Stronger oxidising agent → Easily reduced.
• $E°_{\text{cell}} > 0$ → Spontaneous reaction (produces electricity).
• $E°_{\text{cell}} < 0 →$ Non-spontaneous reaction (requires electricity).