Chemistry
25 flashcards
~13 min
Electrochemistry
Understand the chemistry of electricity with this comprehensive electrochemistry flashcard deck. Electrochemistry is the study of how chemical reactions produce electricity and how electricity drives chemical reactions — it underpins batteries, fuel cells, met...
About this deck
Understand the chemistry of electricity with this comprehensive electrochemistry flashcard deck. Electrochemistry is the study of how chemical reactions produce electricity and how electricity drives chemical reactions — it underpins batteries, fuel cells, metal plating, corrosion, and industrial processes like aluminium smelting. This deck covers oxidation states, redox reactions, galvanic (voltaic) cells, standard electrode potentials, the electrochemical series, electrolysis, Faraday's laws of electrolysis, the Nernst equation, and real-world applications including batteries and the Hall-Héroult process. Essential for AP Chemistry, A-Level, and university physical and inorganic chemistry.
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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.
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).
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.
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).
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.
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.
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.
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).
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.
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.
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).
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.
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.
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.
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.
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.
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).
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.
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.
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.
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.
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.
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.
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.
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.