Chapter 20: Electrochemistry

Introduction to Electrochemistry

• Study of relationships between electricity and chemical reactions.
• Covers spontaneous and non-spontaneous processes.

Key Processes:

1. Electrolytic Cells: Use electrons to produce substances (e.g., electroplating, aluminum production).
2. Galvanic/Voltaic Cells: Generate electricity from chemical reactions (e.g., batteries).

Oxidation Numbers

• Essential for understanding where electrons come from and go.
• Rules for Assigning Oxidation Numbers:
  • Pure elements = 0
  • Monatomic ions = charge
  • Specific elements: F (-1), O (-2, peroxide -1), H (+1, metal hydride -1)
  • Sum equals overall charge.

REDOX Reactions

• Oxidation: increase in oxidation number (loss of electrons).
• Reduction: decrease in oxidation number (gain of electrons).
• Agents:
  • Oxidizing agent: causes oxidation in others, is reduced itself.
  • Reducing agent: causes reduction in others, is oxidized itself.

Balancing REDOX Equations: The Half-Reaction Method

1. Split into oxidation and reduction half-reactions.
2. Balance all atoms except O and H.
3. Balance O with H₂O and H with H⁺.
4. Balance charges with electrons.
5. Multiply half-reactions to equalize electrons.
6. Combine and simplify.

Example Reaction: MnO₄⁻ and C₂O₄²⁻

• Mn reduced from +7 to +2, C oxidized from +3 to +4.
• Balance the manganese and oxygen, then add H⁺ and e⁻ accordingly.

Electrochemical Cells

• Galvanic/Voltaic Cells: Separate half-reactions, allowing electron flow through a wire.
  • Anode: oxidation occurs.
  • Cathode: reduction occurs.
  • Bridge allows ion flow to maintain neutrality.
• Mnemonics: REDuction at CATode, OXidation at ANode.

Electromotive Force (EMF)

• Electrons flow spontaneously from high to low potential energy.
• EMF (E_cell): potential difference between anode and cathode (measured in volts).

Standard Cell Potentials

• Standard conditions: 1 M solutions, 1 atm gases, 25°C (298 K).
• Standard Hydrogen Electrode (SHE) as a reference (0 V).
• Standard potential for oxidation = negative of reduction potential.

Thermodynamics and Electrochemistry

• Relationship between standard free energy (G°) and equilibrium constant (K).

  [ G^° = -nFE^° ]

• Nernst Equation for non-standard conditions:

  [ E = E^° - \frac{RT}{nF} \ln Q ]

pH Meters

• Measure potential based on ion concentration differences across a membrane.
• Potential changes by 59.2 mV/pH unit change.

Electrolysis

• Electrolysis: Use electrical energy to drive nonspontaneous reactions.
• Faraday’s Law of Electrolysis: Amount of substance produced is proportional to charge passed.
  • Formula: [ \text{mass} = \frac{\text{charge} \times \text{molar mass}}{nF} ]

Michael Faraday: Developed foundational principles for electrolysis.

Important Equations and Concepts

• Standard EMF: Calculated as the difference in reduction potentials of cathode and anode.
• Faraday's constant: ( 96,485 \text{ C/mol} ) - number of coulombs per mole of electrons.
• Work and Free Energy: Related through redox potentials and thermodynamics.