// CHEMISTRY L5 — ELECTROCHEMISTRY
Build a battery from two metals.
Connect two different metals in their salt solutions and electrons flow — that's a cell. The bigger the gap in their electrode potentials, the bigger the voltage. Pick two and read the EMF.
// FIG. 01 — THE CELL BUILDER
{{ negTerm }} TERMINAL (−)
{{ negMetal }}
OXIDATION (loses e⁻)
{{ negHalf }}
{{ posTerm }} TERMINAL (+)
{{ posMetal }}
REDUCTION (gains e⁻)
{{ posHalf }}
electrons flow through the wire {{ negMetal }} → {{ posMetal }} · a salt bridge completes the circuit
E°cell = E°(+) − E°(−) = ({{ posE }}) − ({{ negE }})
CELL VOLTAGE E°cell =
{{ emf }} V
{{ feasible }}
// STANDARD ELECTRODE POTENTIALS
Every half-reaction has a standard electrode potential, E° — measured against a hydrogen electrode set at exactly 0.00 V, under standard conditions (298 K, 1 mol/dm³, 100 kPa). A more negative E° means the metal gives up electrons more readily — it gets oxidised and becomes the negative electrode. The more positive half-cell gains those electrons (reduction) and is the positive electrode.
// THE ELECTROCHEMICAL SERIES
The famous Daniell cell (zinc + copper) sits in this table: E°cell = +0.34 − (−0.76) = 1.10 V. This is also exactly why a reactive metal displaces a less reactive one — electrochemistry and the reactivity series are the same order.
// FROM CELLS TO BATTERIES & FUEL CELLS
A positive E°cell means the reaction is feasible — it will drive a current. Everyday batteries are cleverly packaged versions of this; rechargeable ones simply run the reaction backwards using an external supply. A hydrogen fuel cell combines hydrogen and oxygen to make electricity directly, with only water as the by-product — a leading candidate for clean transport.
// REFERENCES & FURTHER READING
Cross-check against your own exam board's specification (AQA, Edexcel, OCR, SQA).