Electrochemistry

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Oxidation: Loss of electrons (OIL). Reduction: Gain of electrons (RIG). Memory: OIL RIG. In any redox reaction, one species is oxidised and another is reduced simultaneously.

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A measure of the degree of oxidation of an atom. Rules: Pure element = 0. Simple ion = its charge. O = −2 (except in peroxides −1 and F₂O +2). H = +1 (except in metal hydrides −1). Sum of oxidation states in a compound = 0 (or ion charge).

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Reducing agent: The species that gets oxidised (loses electrons) — it causes reduction in another species. Oxidising agent: The species that gets reduced (gains electrons) — it causes oxidation in another species.

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An electrochemical cell that converts chemical energy into electrical energy through a spontaneous redox reaction. Two half-cells connected by a wire (electrons flow through) and a salt bridge (ions flow through).

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Anode: Electrode where oxidation occurs (negative in galvanic cell, positive in electrolytic cell). Cathode: Electrode where reduction occurs (positive in galvanic cell, negative in electrolytic cell). Memory: AN OX, RED CAT.

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Maintains electrical neutrality in both half-cells by allowing ions to flow between them, completing the circuit. Without it, charge builds up and the cell stops working. Usually contains KCl or KNO₃ in agar gel.

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The voltage (EMF) of a half-cell measured under standard conditions (25°C, 1 mol/L, 1 atm) relative to the standard hydrogen electrode (SHE, E° = 0.00 V). More positive E° = stronger oxidising agent.

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E°cell = E°cathode − E°anode (reduction potential of cathode minus reduction potential of anode). A positive E°cell means the reaction is spontaneous (ΔG < 0).

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A list of half-reactions ranked by standard electrode potential from most negative (strongest reducing agents: Li, Na, K) to most positive (strongest oxidising agents: F₂, Au). Species above hydrogen on the table can displace H₂ from acids.

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Using an external electrical current to drive a non-spontaneous redox reaction. The opposite of a galvanic cell. Used to decompose compounds, electroplate metals, and extract reactive metals.

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First Law: Mass of substance deposited is proportional to charge passed. Second Law: For the same charge, masses deposited are proportional to molar mass / number of electrons transferred. Faraday's constant F = 96,485 C/mol (charge per mole of electrons).

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m = (M × I × t) / (n × F) where m = mass (g), M = molar mass, I = current (A), t = time (s), n = number of electrons transferred, F = 96,485 C/mol.

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Industrial electrolysis of aluminium oxide (Al₂O₃) dissolved in molten cryolite (Na₃AlF₆) to extract aluminium. Cathode: Al³⁺ + 3e⁻ → Al. Anode: 2O²⁻ → O₂ + 4e⁻. Very energy-intensive.

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Cathode: 2H⁺ + 2e⁻ → H₂ (hydrogen gas). Anode: 2Cl⁻ → Cl₂ + 2e⁻ (chlorine gas). Remaining solution: NaOH (sodium hydroxide). Products used in PVC, bleach, and soap manufacturing.

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Coating an object with a thin layer of metal using electrolysis. The object = cathode (metal deposits on it). The plating metal = anode (dissolves to replenish solution). Example: silver-plating cutlery. Used for decoration, corrosion protection, and conductivity.

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E = E° − (RT/nF) × ln Q or at 25°C: E = E° − (0.0592/n) × log Q. Calculates cell potential under non-standard conditions. As Q increases (products accumulate), E decreases — the battery runs down.

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ΔG° = −nFE°cell. If E°cell > 0 → ΔG° < 0 → spontaneous reaction. Links electrochemistry to thermodynamics. n = moles of electrons transferred, F = Faraday's constant (96,485 C/mol).

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An electrochemical cell that converts the chemical energy of a fuel (usually H₂) directly into electricity. Hydrogen fuel cell: H₂ + ½O₂ → H₂O + electricity. More efficient than combustion engines. Used in electric vehicles and spacecraft.

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Corrosion is the oxidation of a metal by its environment (air, water). Iron rusting: Fe → Fe²⁺ + 2e⁻ (oxidation at anode). O₂ + 2H₂O + 4e⁻ → 4OH⁻ (reduction at cathode). Prevention: galvanizing (zinc coating), painting, sacrificial anode.

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A more reactive metal attached to a less reactive metal to prevent corrosion. The more reactive metal (e.g., zinc, magnesium) preferentially oxidises (sacrifices itself), protecting the less reactive metal. Used on ships, pipelines, and oil rigs.

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The reference electrode with E° = 0.00 V by definition. Consists of H₂ gas (1 atm) bubbled over a platinum electrode in 1 M H⁺ solution at 25°C. All standard electrode potentials are measured relative to it.

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A ranking of metals by reactivity (ability to displace other metals from solutions). Corresponds to the electrochemical series: more reactive metals = more negative E°. A metal higher in the series displaces ions of metals below it.

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Cathode: 2H⁺ + 2e⁻ → H₂ (hydrogen gas — reduction). Anode: 2H₂O → O₂ + 4H⁺ + 4e⁻ (oxygen gas — oxidation). Ratio of H₂ : O₂ = 2:1 by volume. Overall: 2H₂O → 2H₂ + O₂.

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A galvanic cell where both half-cells contain the same species but at different concentrations. The cell potential depends on the concentration difference. E = 0 when concentrations are equal. Described by the Nernst equation.

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Many half-reactions involve H⁺ or OH⁻. By the Nernst equation, changing [H⁺] (pH) changes the electrode potential. This is why the oxygen electrode potential varies with pH and why corrosion rates depend on acidity.