Overview
This lecture explains the setup and functioning of electrochemical cells, details how to calculate cell potentials, and explores examples like lithium-ion and hydrogen fuel cells.
Electrochemical Cell Basics
- An electrochemical cell consists of two different metals (electrodes) dipped in solutions containing their own ions.
- A salt bridge (often filter paper soaked in potassium nitrate) completes the circuit by allowing ion flow between half-cells.
- Electromotive force (emf) or cell voltage drives electron flow through the connecting wire.
- The more reactive metal is oxidized (loses electrons) and the less reactive is reduced (gains electrons).
- The setup can also use inert electrodes (e.g., platinum) with different ion solutions (e.g., Fe²⁺/Fe³⁺).
Half-Equations and Electrode Potentials
- Both oxidation and reduction can occur at each electrode; reactions are reversible.
- Write half-equations with the reduction reaction going forward (electrons on the left).
- The more negative electrode potential indicates the metal will be oxidized.
- Combine half-equations to form the overall redox reaction in the cell.
- Electrode potentials are measured relative to the standard hydrogen electrode (SHE), which is defined as 0 volts.
Standard Conditions and Calculating emf
- Standard conditions: solutions at 1 mol/dm³, temperature at 298 K, and pressure at 100 kPa (1 atm).
- The cell emf is calculated as the difference between the electrode potentials of the two metals (subtract one from the other, accounting for sign).
Electrochemical Cell Notation
- Shorthand notation: On the left, reduced form | oxidized form || oxidized form | reduced form on the right.
- A double vertical line (||) represents the salt bridge.
Battery and Cell Examples
- Lithium-ion batteries use a graphite and lithium cobalt oxide electrode with a lithium salt electrolyte.
- Lithium is oxidized at one electrode; lithium ions travel and are reduced at the other, forming a complex ion.
Hydrogen Fuel Cells
- Hydrogen fuel cells use hydrogen and oxygen gases at separate platinum electrodes with potassium hydroxide solution and an anion exchange membrane.
- The membrane allows hydroxide ions (OH⁻) to travel and combine with hydrogen to form water (H₂O).
- Electrons flow through the external circuit, creating usable current.
- Main advantage: only water is produced, with no release of pollutants.
Key Terms & Definitions
- Electromotive force (emf) — The voltage generated by an electrochemical cell.
- Salt bridge — A device allowing ion flow to complete the electrochemical circuit.
- Half-cell — One compartment of an electrochemical cell containing an electrode and ion solution.
- Electrode potential — The voltage of a half-cell compared to the standard hydrogen electrode.
- Redox reaction — A reaction involving simultaneous oxidation and reduction.
Action Items / Next Steps
- Review the table of standard electrode potentials.
- Practice writing half-equations and overall redox equations for given cell setups.