Grade 12 Chemistry Exam Review Unit 5: Electrochemistry

Flashcards covering the fundamentals of electrochemistry, including oxidation-reduction rules, half-reactions, agents, galvanic cell components, and standard reduction potentials.

Oxidation (LEO)

The loss of electrons from a substance; results in an increase in the oxidation number.

Reduction (GER)

The gain of electrons by a substance; results in a decrease in the oxidation number.

Oxidation Number Rule 1

The oxidation number for any element by itself (e.g., $\text{Na}$, $\text{O}_2$, $\text{Cl}_2$, $\text{Mg}$) is $0$.

Oxidation Number Rule 2

For a neutral compound, the sum of the oxidation numbers must equal $0$.

Oxidation Number Rule 3

For a polyatomic ion, the sum of the oxidation numbers must equal the ion charge.

Oxidation Number Rule 5

The oxidation number of Hydrogen in a metal hydride (e.g., $\text{LiH}$) is $-1$.

Oxidation Number Rule 7

The oxidation number of Oxygen in peroxides (e.g., $\text{H}_2\text{O}_2$) is $-1$.

Group 1 Metals

Metals that always have an oxidation number of $+1$ (e.g., $\text{Na}$, $\text{K}$, $\text{Li}$).

Group 2 Metals

Metals that always have an oxidation number of $+2$ (e.g., $\text{Mg}$, $\text{Ca}$, $\text{Ba}$).

Redox Reaction

A chemical reaction in which the oxidation numbers of atoms change due to the transfer of electrons.

Oxidizing Half-Reaction

A half-reaction where electrons are written on the product side (e.g., $\text{Mg} \rightarrow \text{Mg}^{2+} + 2\text{e}^-$).

Reducing Half-Reaction

A half-reaction where electrons are written on the reactant side (e.g., $\text{Cu}^{2+} + 2\text{e}^- \rightarrow \text{Cu}$).

Oxidizing Agent (OA)

A substance that causes another substance to be oxidized; the oxidizing agent itself gains electrons and is reduced.

Reducing Agent (RA)

A substance that causes another substance to be reduced; the reducing agent itself loses electrons and is oxidized.

Basic Solution Rules

Converted from a balanced acidic solution by adding $\text{OH}^-$ to both sides equal to the number of $\text{H}^+$ present, then combining $\text{H}^+$ and $\text{OH}^-$ into $\text{H}_2\text{O}$.

Galvanic Cell

A voltage cell that converts chemical energy from a spontaneous redox reaction into electrical energy.

Anode

The negative electrode where oxidation occurs (AN OX); it produces electrons and its mass typically decreases.

Cathode

The positive electrode where reduction occurs (RED CAT); it consumes electrons and its mass typically increases.

Electron Flow

The path of electrons in a galvanic cell, moving from the Anode to the Cathode (or Oxidation to Reduction).

Electrolyte

A solution containing ions, such as $\text{Zn(NO}_3\text{)}_2(aq)$ or $\text{Cu(NO}_3\text{)}_2(aq)$, used in electrochemical cells.

Salt Bridge

A component that completes the circuit and maintains electrical neutrality by allowing ions to move between half-cells.

Standard Conditions for Half Cells

Solution concentration of $1.0\,\text{M}$, temperature of $25^{\circ}\text{C}$, and gas pressure of $100\,\text{kPa}$.

Standard Hydrogen Electrode (SHE)

The reference half-cell used to determine all reduction potentials, representing the reaction $2\text{H}^+(aq) + 2\text{e}^- \rightarrow \text{H}_2(g)$ with $E^\circ = 0.00\,\text{V}$.

Standard Reduction Potential ($E^\circ$)

A measure of the tendency of a species to gain electrons and be reduced; a more positive value indicates a stronger oxidizing agent.

Cell Potential Formula

$E^\circ_{\text{cell}} = E^\circ_{\text{red}}(\text{cathode}) - E^\circ_{\text{red}}(\text{anode})$.

Spontaneity

A redox reaction is spontaneous if the calculated $E^\circ_{\text{cell}} > 0$.

Cell Notation General Formation

Anode | Anode solution || Cathode solution | Cathode.

Phase Boundary symbol ($|$)

In cell notation, this symbol separates different phases such as a solid electrode and an aqueous solution.

Salt Bridge symbol ($||$)

In cell notation, this symbol separates the two half-cells.