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Oxidation

Loss of electrons

Reduction

Gain of electrons

Redox reaction

Reaction involving electron transfer

Oxidizing agent

Gains electrons and is reduced

Reducing agent

Loses electrons and is oxidized

Anode

Where oxidation occurs

Cathode

Where reduction occurs

Electron flow

From anode to cathode

Sign of anode

Negative

Sign of cathode

Positive

Cell potential (E°)

Voltage of electrochemical cell

Standard conditions

1 M, 1 atm, 25°C

E°cell

E°cathode - E°anode

Spontaneous reaction

E°cell > 0

Non-spontaneous reaction

E°cell < 0

n

Number of electrons transferred

ΔG° equation

ΔG° = -nFE°

F (Faraday constant)

96485 C/mol e⁻

Relationship between ΔG and E°

Negative ΔG corresponds to positive E°

Nernst equation

E = E° - (RT/nF) ln Q

Simplified Nernst (25°C)

E = E° - (0.0592/n) log Q

Reaction quotient (Q)

Ratio of products to reactants

At equilibrium

E = 0 and Q = K

Relationship between E° and K

E° = (0.0592/n) log K

Half-reaction

Shows oxidation or reduction separately

Balancing electrons

Multiply half-reactions so electrons cancel

Salt bridge

Allows ions to flow and maintain charge balance

Purpose of salt bridge

Prevents charge buildup in the cell

Voltaic (galvanic) cell

Spontaneous electrochemical cell

Electrolytic cell

Non-spontaneous cell requiring energy

Strong oxidizing agent

Large positive reduction potential

Strong reducing agent

Large negative reduction potential

How to identify cathode

Higher (more positive) E° value

How to identify anode

Lower (more negative) E° value

Common mistake

Forgetting to flip sign of anode

Common mistake

Not balancing electrons when finding n

Common mistake

Mixing up electron flow direction