1. Write the half-reactions. 2. Find the standard reduction potentials for each half-reaction. 3. Flip the sign of the oxidation half-reaction. 4. Calculate E°: E° = E°(reduction) + E°(oxidation).

1. Identify the half-reactions (reduction and oxidation). 2. Determine the standard reduction potentials for each half-reaction. 3. Calculate the overall cell potential (E°cell). 4. Apply a voltage greater than or equal to the absolute value of E°cell to drive the non-spontaneous reaction.

A reaction involving the transfer of electrons between chemical species.

A substance that gains electrons and is reduced in a redox reaction.

A substance that loses electrons and is oxidized in a redox reaction.

The force that pushes electrons in a redox reaction, measured in volts (V).

The voltage associated with a reduction half-reaction under standard conditions.

An electrochemical cell that uses spontaneous redox reactions to generate electricity.

An electrochemical cell that uses electrical energy to drive non-spontaneous redox reactions.

The electrode where oxidation occurs.

The electrode where reduction occurs.

Anode: Where oxidation occurs; Cathode: Where reduction occurs; Electron flow: From anode to cathode through the wire.

Anode: Where oxidation occurs; Cathode: Where reduction occurs; Electron flow: From the external power source to the cathode, and from the anode back to the power source.

Salt bridge: A tube containing an electrolyte (e.g., $NaNO_3$) that connects the two half-cells; Function: Maintains charge neutrality by allowing ions to flow between the half-cells.