Oxidation, Reduction, and Redox Reactions

Flashcards covering the definitions of oxidation, reduction, redox agents, oxidation state rules, and disproportionation based on the lecture notes.

Redox reaction

A chemical process that is short for reduction-oxidation, involving the movement of electrons from one species to another.

Oxidation (historical definition)

A chemical reaction in which oxygen is added to a substance or hydrogen is removed from it.

Reduction (historical definition)

A chemical reaction in which oxygen is removed from a substance or hydrogen is added to it.

Oxidation (electron transfer)

The loss of electrons by a chemical species during a reaction ($\text{OIL}$: Oxidation Is Loss).

Reduction (electron transfer)

The gain of electrons by a chemical species during a reaction ($\text{RIG}$: Reduction Is Gain).

Electron transfer reactions

Another name for redox reactions, identifying that they always involve the movement of electrons between species.

Half equation

A balanced equation that shows only the gain or loss of electrons by a single species, such as $\text{Cu} \rightarrow \text{Cu}^{2+} + 2\text{e}^-$.

Spectator ion

An ion that takes no part in a chemical reaction and whose oxidation state remains unchanged, such as the $\text{O}^{2-}$ ion in the reaction between copper oxide and magnesium.

Ionic equation

A simplified chemical equation for a redox reaction that shows only the species that are oxidised or reduced, omitting spectator ions.

Reducing agent

A species that gives away or donates electrons to another substance; it facilitates reduction and is itself oxidised.

Oxidising agent

A species that accepts electrons from another substance; it facilitates oxidation and is itself reduced.

Oxidation state

A number assigned to an element in a compound that represents the number of electrons lost or gained compared to its uncombined state, or the distribution of electrons in a molecule.

Oxidation state of uncombined elements

The oxidation number assigned to any element in its pure, uncombined form, which is always $0$.

Sum of oxidation states in a compound

The total value of all oxidation states of the elements in a neutral compound, which must always equal $0$.

Sum of oxidation states in a complex ion

The total value of all oxidation states of the elements in a complex ion, which must equal the charge on the ion (e.g., $-2$ for $\text{SO}_4^{2-}$).

Oxidation state of Group 1 metals

In compounds, these elements always have an oxidation state of $+1$.

Oxidation state of Group 2 metals

In compounds, these elements always have an oxidation state of $+2$.

Oxidation state of Aluminum

In compounds, this element always has an oxidation state of $+3$.

Oxidation state of Fluorine

As the most electronegative element, it always has an oxidation state of $-1$ in all its compounds.

Oxidation state of Hydrogen

Usually $+1$, except in metal hydrides such as $\text{NaH}$ where it is $-1$.

Oxidation state of Oxygen

Usually $-2$, except in peroxides (where it is $-1$), superoxides (where it is $-\frac{1}{2}$), or in $\text{OF}_2$ (where it is $+2$).

Copper(II) oxide

Black $\text{CuO}$, where the copper has an oxidation state of $+2$.

Copper(I) oxide

Red $\text{Cu}_2\text{O}$, where each copper atom has an oxidation state of $+1$.

Potassium superoxide

An orange-yellow compound with the formula $\text{KO}_2$ containing the $\text{O}_2^-$ ion, where each oxygen atom has an unusual oxidation state of $-\frac{1}{2}$.

Thermite reaction

A strongly exothermic redox reaction in which aluminium reacts with iron(III) oxide ($\text{Fe}_2\text{O}_3$) to produce molten iron ($\text{Fe}$).

Disproportionation

A reaction in which atoms of the same element in a single substance are simultaneously oxidised and reduced, such as in the decomposition of $\text{H}_2\text{O}_2$.