Electrochemistry: Oxidation-Reduction and Galvanic Cells

Electrochemistry: The study of the interchange of chemical and electrical energy.
Redox reaction: A chemical reaction that involves the transfer of electrons between two species.
Galvanic (Voltaic) Cells: Devices that convert chemical energy into electrical energy through spontaneous redox reactions.
Electrolytic Cells: Devices that use electrical energy to drive non-spontaneous redox reactions.
Anode: The electrode where oxidation occurs (loss of electrons).
Cathode: The electrode where reduction occurs (gain of electrons).
Salt Bridge: A device that maintains electrical neutrality by allowing ions to flow between the two half-cells of a galvanic cell.
E° cell: Standard cell potential, measured in volts (V).

OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons).
LEO the lion says GER: Lose Electrons in Oxidation; Gain Electrons in Reduction.

Split the overall reaction into oxidation and reduction half-reactions.
Balance atoms for each half-reaction (excluding H and O).
Balance O by adding H₂O molecules.
Balance H by adding H+ ions.
Balance charges by adding electrons.
Ensure equal number of electrons in both half-reactions, adjusting coefficients if necessary.
Add the half-reactions together and simplify.
In basic solutions, add OH- to balance H+ ions, forming H₂O where applicable.

Galvanic Cells:
- Spontaneous reactions.
- Convert chemical energy to electrical energy.
- Use inert electrodes in certain configurations.
Electrolytic Cells:
- Non-spontaneous reactions.
- Use external energy sources to drive reactions.

Anode (−): Where oxidation occurs; mass decreases over time.
Cathode (+): Where reduction occurs; mass increases over time.
Electron Flow: From anode to cathode.
Salt Bridge: Contains neutral salts; facilitates movement of ions to maintain charge balance.

Oxidation half-reaction: Zn(s) → Zn²⁺(aq) + 2e⁻.
Reduction half-reaction: Cu²⁺(aq) + 2e⁻ → Cu(s).
Overall reaction: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s).
Standard cell potential (E° cell): Determined by the difference in reduction potentials of the half-reactions.

Standard Reduction Potentials: Can be found in tables, indicating a substance's tendency to gain electrons (be reduced).
The more positive the reduction potential, the more favorable the reduction.
E° cell Calculation:
1. Write reduction reactions and associated potentials.
2. Identify which half-reaction will be oxidized (reverse its equation and change its sign).
3. Sum the potentials: E° cell = E° oxidation + E° reduction.

Cell potentials may vary due to changes in concentration, temperature, or pressure.
The Nernst Equation can adjust for these variables, detailing how non-standard conditions favor different reaction directions.

Related to E° cell by the equation: ΔG° = -nFE° cell, where n is the number of moles of electrons transferred and F is Faraday's constant (96485 C/mol).
Positive ΔG° indicates non-spontaneous reactions, while negative ΔG° indicates spontaneous reactions.

Equilibrium is reached in a cell when ΔG°= 0, meaning the forward and reverse reactions are balanced.
Factors influencing shift in equilibrium can be described using Le Chatelier's principle.

Always track electron flow and the identity of reducing and oxidizing agents.
Maintain clarity in distinguishing between the roles of different components within electrochemical cells.
Familiarize yourself with the reactions and how they indicate the favorability of spontaneous processes, particularly in galvanic vs. electrolytic contexts.