Batteries, Corrosion & Electrolysis – CHM2046 Session 8

Vocabulary flashcards summarizing essential terms and definitions on batteries, corrosion, and electrolysis from the lecture notes.

Battery

A self-contained group of voltaic cells in series whose individual voltages add together.

Primary Battery

A non-rechargeable battery that irreversibly converts chemical energy to electrical energy.

Secondary Battery

A rechargeable battery whose chemical reactions can be reversed by supplying electrical energy.

Fuel Cell

An electrochemical device that continuously converts external fuel and oxidant into electricity with separated half-reactions.

Alkaline Battery

Primary cell with a zinc anode case, MnO₂/KOH paste cathode mix, graphite rod, ~1.5 V output.

Mercury Battery

Primary dry cell using Zn anode and HgO cathode in basic KOH paste; compact button form.

Silver Battery

Primary dry cell using Zn anode and Ag₂O cathode in basic medium; common in watch batteries.

Lead-Acid Battery

Secondary 12 V car battery of six 2.1 V cells with Pb anode, PbO₂ cathode, 4.5 M H₂SO₄ electrolyte.

Lithium-Ion Battery

Rechargeable cell with graphite-hosted Li anode, Li metal oxide cathode, Li⁺ conducting electrolyte.

Nickel-Cadmium (Ni-Cd) Battery

Secondary battery where Cd(s) is oxidized and NiOOH is reduced during discharge, E° ≈ 1.2 V.

Proton Exchange Membrane (PEM) Fuel Cell

Hydrogen-oxygen fuel cell in which H₂ is oxidized at anode and O₂ reduced at cathode to form H₂O.

Galvanic (Voltaic) Cell

Spontaneous electrochemical cell that produces electrical energy during discharge.

Electrolytic Cell

Non-spontaneous cell driven by external electricity; operates during battery recharge and electrolysis.

Anode (Electrolytic)

Positive electrode where oxidation occurs and electrons leave the cell.

Cathode (Electrolytic)

Negative electrode where reduction occurs and electrons enter the cell.

Corrosion

Natural redox process that oxidizes metals to oxides or sulfides, e.g., rusting of iron.

Rust

Hydrated iron(III) oxide, Fe₂O₃·nH₂O, formed when Fe reacts with O₂ and moisture.

Cathodic Protection

Prevention of corrosion by attaching a more active (sacrificial) metal that is preferentially oxidized.

Sacrificial Anode

More active metal (e.g., Zn) that corrodes instead of the protected iron structure.

Faraday Constant (F)

96 500 C mol⁻¹ of electrons; converts between charge and moles of e⁻.

Faraday’s Law of Electrolysis

Mass deposited: m = Q·MM / (F·n), where Q is charge, MM molar mass, n electrons.

Overvoltage

Extra potential required above E° to drive certain electrode reactions, significant for O₂ evolution.

Molten Electrolysis Rule – Cations

Species with highest ionization energy is reduced at the cathode in molten mixture.

Molten Electrolysis Rule – Anions

Species with lowest electron affinity is oxidized at the anode in molten mixture.

Aqueous Electrolysis – Cathode

Species with highest E° is reduced; active metal cations (Group 1, 2, Al) are bypassed in favor of water.

Aqueous Electrolysis – Anode

Species with lowest E° is oxidized; halides (except F⁻) oxidize easier than water at high [Cl⁻], etc.

Hydrogen Fuel Cell Half-Reactions

Anode: H₂ → 2H⁺ + 2e⁻; Cathode: O₂ + 4H⁺ + 4e⁻ → 2H₂O.

Relationship ΔG°, E°cell, K

ΔG° = –nFE°, E° > 0 gives K > 1 (spontaneous); electrolytic cells have E° < 0, K < 1.

Current (I)

Rate of charge flow, 1 A = 1 C s⁻¹; used with time to calculate Q for electrolysis.

Coulomb (C)

Unit of electric charge; 1 C = charge carried by ~6.242×10¹⁸ electrons.