Rucete ✏ AP Chemistry In a Nutshell
12. Oxidation-Reduction Reactions and Electrochemistry
This chapter explores oxidation-reduction (redox) reactions, focusing on electron transfer, oxidation states, balancing techniques, electrochemical cells, and the connection between redox chemistry and electricity.
Oxidation and Reduction
• Oxidation: loss of electrons.
• Reduction: gain of electrons.
• These processes always occur together and can be represented as half-reactions.
• Example: Ba loses 2e– (oxidized), S gains 2e– (reduced).
Oxidation Numbers
• Oxidation number (oxidation state) helps track electron movement in reactions.
• A change in oxidation number indicates a redox process.
Oxidation Number Rules
• Rule 1: The sum of all oxidation numbers equals the total charge.
• Rule 2: Alkali metals = +1, alkaline earth = +2, Group 3 = +3.
• Rule 3: H = +1; F = –1.
• Rule 4: O = –2.
• Rule 5: Halogens = –1.
• Rule 6: Group 6 nonmetals = –2.
• Apply rules hierarchically; higher rules override lower ones if conflicts occur.
Using Oxidation Numbers
• Redox identified by oxidation number changes between reactants and products.
• Example: MnO₄⁻ (Mn = +7) → Mn²⁺ (Mn = +2) indicates reduction (gain of 5e–).
Balancing Redox Reactions: Ion-Electron Method (in Acid)
1. Write separate oxidation and reduction half-reactions.
2. Balance atoms other than O and H.
3. Balance O by adding H₂O.
4. Balance H by adding H⁺.
5. Balance charge using electrons (e⁻).
6. Multiply reactions to equalize electrons and combine them.
7. If in basic solution, neutralize H⁺ with OH⁻ to form H₂O, then simplify.
Common Redox Reactions
• Single-replacement: metals displace ions or hydrogen (e.g., Zn + HCl → ZnCl₂ + H₂).
• Very active metals react with water (e.g., Na, K), active metals react with acids (e.g., Mg, Zn), and inactive metals do not displace hydrogen (e.g., Ag, Au).
Permanganate Reactions
• In acid: MnO₄⁻ → Mn²⁺ (5e⁻ reduced).
• In neutral: MnO₄⁻ → MnO₂ (3e⁻ reduced).
• In base: MnO₄⁻ → MnO₄²⁻ (1e⁻ reduced).
Electrolysis
• Non-spontaneous redox driven by electricity.
• Cathode: reduction; Anode: oxidation.
• In molten salts: cation reduced at cathode, anion oxidized at anode.
• In aqueous solution, water or solutes may be oxidized/reduced depending on reactivity.
Electrochemical Cells (Voltaic/Galvanic)
• Spontaneous redox reactions produce electricity.
• Oxidation occurs at the anode (electrons flow out).
• Reduction occurs at the cathode (electrons flow in).
• Salt bridge allows ion flow to maintain electrical neutrality.
• Electrons always flow from anode to cathode.
Cell Notation
• Format: Anode | Anode solution || Cathode solution | Cathode
• Example: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)
Standard Electrode Potentials (E°)
• Reduction potentials are measured relative to the standard hydrogen electrode (SHE), which is 0.00 V.
• Positive E° → greater tendency to be reduced.
• The more positive E° is favored in redox reactions.
• E°cell = E°cathode – E°anode
Using E° Values
• Do not multiply E° values by coefficients from balanced equations.
• A positive E°cell indicates a spontaneous reaction.
• Reverse sign of E° when writing oxidation half-reactions.
Free Energy and Cell Potential
• ΔG° = –nFE°cell – n = number of electrons transferred – F = Faraday’s constant = 96,485 C/mol e⁻
• Negative ΔG° indicates a spontaneous process.
Nernst Equation
• E = E° – (RT/nF) × lnQ – Q = reaction quotient – Used to calculate cell potential under nonstandard conditions.
• Simplified at 25°C: E = E° – (0.0591/n) × logQ
Electrolytic Cells
• Require an external power source to drive nonspontaneous reactions.
• Anode is positive (connected to battery’s + terminal), cathode is negative.
• Electrolysis used in electroplating, metal refining, and decomposition of compounds.
In a Nutshell
Redox reactions involve electron transfer, tracked by oxidation numbers. Electrochemical cells harness spontaneous redox reactions to generate electricity, while electrolytic cells use energy to force nonspontaneous reactions. Cell potentials, calculated from standard reduction potentials and adjusted by the Nernst equation, determine feasibility and direction of redox processes.