Oxidation Number

Uncombined Elements: Any element in its pure, uncombined state (e.g., $S_8$, $O_2$, $Fe$) always has an oxidation number of $0$.

Monoatomic Ions: The oxidation number is equal to the charge of the ion; for example, $Mg^{2+}$ is $+2$ and $Cl^-$ is $-1$.

Fixed Values in Compounds: Group 1 metals are always $+1$, Group 2 metals are always $+2$, and Fluorine is always $-1$ because it is the most electronegative element.

Hydrogen and Oxygen: Hydrogen is usually $+1$ (except in metal hydrides where it is $-1$), and Oxygen is usually $-2$ (except in peroxides where it is $-1$ or in $F_2O$ where it is $+2$).

A Disproportionation Reaction is a specific type of redox reaction where a single element in a single species is simultaneously oxidised and reduced.

To identify these, one must calculate the oxidation number of the element in the reactant and compare it to the two different oxidation numbers it possesses in the products.

This process results in the element 'splitting' its electronic state into one higher and one lower oxidation level.

Always include signs: Oxidation numbers must be written with a $+$ or $-$ sign (e.g., $+2$, not just $2$) to distinguish them from mathematical counts.

Per-Atom Basis: When calculating, remember that the oxidation number refers to a single atom of that element, even if the formula contains multiple atoms (e.g., in $O_2$, each $O$ is $0$).

Verification: After assigning numbers, always sum them up to ensure they match the total charge of the molecule or ion; this is the most effective way to catch calculation errors.

Roman Numerals: In naming compounds (e.g., Iron(III) sulfate), the Roman numeral specifically denotes the oxidation state of the metal, which is vital for elements with variable states.