Redox Reactions - Oxidation Number

Oxidation Number (O.N.) is defined as the formal charge that an atom would carry in a molecule or ion if the shared electrons in a covalent bond were assigned completely to the more electronegative element.

Rules for assigning O.N.: (1) In free or elementary state (e.g., $H_2$, $O_2$, $P_4$, $S_8$, $Na$), the O.N. of each atom is $0$. (2) For monatomic ions, the O.N. is equal to the charge of the ion (e.g., $Na^+$ is $+1$, $Mg^{2+}$ is $+2$, $Cl^-$ is $-1$).

Oxygen Rule: In most compounds, the O.N. of oxygen is $-2$. Exceptions include peroxides (e.g., $H_2O_2$, $Na_2O_2$) where it is $-1$, superoxides (e.g., $KO_2$) where it is $-\frac{1}{2}$, and in $OF_2$ where it is $+2$.

Hydrogen Rule: The O.N. of hydrogen is $+1$ when bonded to non-metals. However, in metal hydrides (e.g., $LiH$, $CaH_2$), it is $-1$.

Halogen Rule: Fluorine, being the most electronegative element, always has an O.N. of $-1$. Other halogens ($Cl$, $Br$, $I$) are typically $-1$ unless combined with oxygen or fluorine.

In a neutral molecule, the algebraic sum of the oxidation numbers of all constituent atoms must be $0$. In a polyatomic ion, the sum must equal the net charge of the ion.

Oxidation is defined as an increase in the oxidation number of an element, whereas Reduction is defined as a decrease in the oxidation number of an element.

Stock Notation: The oxidation state of a metal in a compound is represented by a Roman numeral in parentheses after the symbol of the metal (e.g., $FeO$ is Iron(II) oxide, $Fe_2O_3$ is Iron(III) oxide).