Galvanic cells, such as the Daniell Cell, function based on spontaneous redox reactions to produce electrical energy.

Primary batteries are single-use cells that cannot be recharged once the reactants are consumed.

The overall reaction in an H₂–O₂ fuel cell is the reaction of hydrogen and oxygen to produce water, which is a major application advantage.

The Standard Hydrogen Electrode (SHE) is universally used as the primary reference electrode, assigned a potential of zero volts.

This specific condition (unity concentration) defines the Standard Electrode Potential.

Galvanization is the process of applying a zinc coating to steel or iron to prevent rusting, as zinc acts as a sacrificial anode.

Conductivity (κ) decreases with dilution because the number of ions per unit volume decreases.

The relative values of the Standard Reduction Potentials of competing species (water vs. ions) determine which substance is oxidized or reduced.

For weak electrolytes, dilution increases the degree of dissociation (α), leading to more ions and thus a sharp increase in Λₘ.

1 Faraday represents the charge on one mole of electrons (approx. 96,500 C).

Fuel cells convert the chemical energy of a fuel (like H₂ or CH₄) directly into electrical energy.

The salt bridge maintains electrical neutrality by allowing the migration of ions between the two half-cells, thereby completing the circuit.

Secondary batteries are rechargeable, and the Lead Storage Battery is the most common example mentioned.

The Nernst Equation is applied to single and complete cells to account for the effect of ion concentrations on the cell potential.

The first law states that m ∝ Q, where Q is the quantity of electricity passed.

Molar conductivity is typically measured in S cm² mol⁻¹ or S m² mol⁻¹.

A negative value of Gibbs free energy indicates a spontaneous process.

Conductance (G) is defined as 1/R and measures the ease with which current flows through the solution.

During discharge, the lead anode (Pb) is oxidized to form lead sulfate (PbSO₄).

Corrosion, such as rusting, involves redox reactions, where the metal acts as the anode and is oxidized.

An electrolytic cell uses electrical energy to drive a non-spontaneous chemical reaction.

The second law states that masses deposited are proportional to their equivalent weights (or electrochemical equivalents).

Kohlrausch’s Law is crucial for determining the molar conductivity at infinite dilution (Λ°ₘ) for weak electrolytes.

Nickel–Cadmium cells have a longer life span and better portability compared to the bulkier lead storage battery.

Based on the Nernst equation, increasing reactant concentration favors the forward reaction, increasing Q’s denominator and thus slightly increasing the cell potential.

The relationship between cell potential and the equilibrium constant is derived from the Nernst equation when E₍cell₎ = 0.

The relationship between cell potential and Gibbs Energy is given by ΔG = -nFE₍cell₎, where n is moles of electrons and F is Faraday’s constant.

Oxidation always occurs at the anode (negative terminal) in a Galvanic cell.

Oxygen acts as the oxidizing agent (electron acceptor) that gets reduced at the cathode, completing the circuit.

Fuel cells convert chemical energy directly to electrical energy, minimizing the energy lost as heat during combustion.