AP Chemistry Unit 9: Applications of Thermodynamics

TLDR

• Apply thermodynamics to real systems using electrochemistry.

• Understand galvanic vs electrolytic cells and redox processes.

• Calculate cell potential and connect E°, ΔG°, and K.

• Use Faraday’s laws to relate charge to amount of substance produced.

• Recognize spontaneity through signs of Ecell and ΔG°.

Why This Unit Matters

Unit 9 brings everything together. It connects thermodynamics, equilibrium, and redox chemistry into one applied framework: electrochemistry.

On the AP Chemistry exam, this unit tests whether you can move fluidly between concepts — using cell potential to predict spontaneity, free energy to determine equilibrium, and charge to calculate chemical change. It is less about memorization and more about conceptual synthesis.

1. Oxidation–Reduction Review

Electrochemistry is built on redox reactions.

• Oxidation: loss of electrons

• Reduction: gain of electrons

Mnemonic: OIL RIG

Key locations:

• Oxidation occurs at the anode.

• Reduction occurs at the cathode.

This rule applies to all electrochemical cells.

2. Electrochemical Cells

Electrochemical cells convert between chemical and electrical energy.

Galvanic (Voltaic) Cells

• Reaction is spontaneous.

• Chemical energy → electrical energy.

• Ecell > 0 and ΔG < 0.

Electrolytic Cells

• Reaction is nonspontaneous.

• Electrical energy → chemical energy.

• Requires an external power source.

Mnemonic:
“Galvanic gives, electrolytic eats.”

3. Cell Potential (Ecell)

Cell potential measures the tendency for electrons to flow.

Key equation:
Ecell = Ecathode − Eanode

Rules:

• Always use standard reduction potentials.

• Do not multiply E° values by coefficients.

• The higher the E°, the greater the tendency to be reduced.

A positive Ecell indicates a spontaneous reaction.

4. Standard Reduction Potentials

Standard reduction potentials are measured under standard conditions.

Important ideas:

• More positive E° → stronger oxidizing agent.

• Reduction potentials are always written as reductions.

• Reversing a half-reaction flips the sign of E°.

AP exam questions often test relative strength using E° values rather than calculations.

5. Spontaneity and Electrochemistry

Electrochemistry mirrors thermodynamic behavior.

• Ecell > 0 → spontaneous

• Ecell = 0 → equilibrium

• Ecell < 0 → nonspontaneous

These signs correspond directly to ΔG values.

6. Free Energy and Cell Potential

Thermodynamics and electrochemistry are mathematically linked.

Key equation:
ΔG° = −nFE°cell

Where:

• n = number of moles of electrons transferred

• F = 96,485 C/mol e⁻

A positive E°cell produces a negative ΔG°, indicating spontaneity.

7. Cell Potential and Equilibrium

Electrochemistry also connects to equilibrium.

Thermodynamic relationship:
ΔG° = −RT ln K

Combined form:
ln K = (nFE°cell) / (RT)

Interpretation:

• Large E°cell → large K → product-favored reaction

• Small E°cell → small K → reactant-favored reaction

This relationship is frequently tested conceptually.

8. Electrochemical Cell Diagrams

Cell notation shows how a cell is constructed.

General format:
Anode | anode solution || cathode solution | cathode

• Single line indicates phase boundary.

• Double line represents the salt bridge.

Example:
Zn | Zn²⁺ || Cu²⁺ | Cu

Electrons always flow from anode to cathode.

9. Electron Flow and the Salt Bridge

• Electrons flow through the wire from anode to cathode.

• The salt bridge maintains charge balance by allowing ion movement.

• Conventional current flows opposite to electron flow.

Failure to mention charge balance often costs explanation points on FRQs.

10. Electrolysis

Electrolysis uses electrical energy to drive nonspontaneous reactions.

Applications:

• Electroplating

• Metal purification

• Water splitting

Key relationships:

• Charge passed depends on current and time.

• Amount of substance produced depends on number of electrons required.

Formulas:
q = It
moles of electrons = q / F

11. Faraday’s Laws of Electrolysis

• The amount of product formed is proportional to the charge passed.

• Reactions requiring more electrons produce less product for the same charge.

This concept frequently appears in calculation-based questions.

Common Pitfalls

• Mixing up anode and cathode definitions.

• Forgetting that E° values are reduction potentials.

• Multiplying E° by coefficients.

• Confusing galvanic and electrolytic cells.

• Forgetting to include n when calculating ΔG°.

Tutor Tip

When answering AP free-response questions, always link signs together:

Positive E°cell → negative ΔG° → large K → spontaneous reaction.

This logical chain demonstrates deep understanding and earns full conceptual credit.

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