Oxidation Numbers Overview

Overview

This lecture explains the concept of oxidation numbers, rules for assigning them, and their importance in chemical nomenclature and reactions.

The oxidation number is the hypothetical charge an atom would have if the compound was made of ions.
They help track electron transfer during chemical reactions.

Atoms in elements (e.g., O₂, S₈, Al) have oxidation number 0.
Simple ions' oxidation number equals their charge (Na+ is +1, Cl- is -1).
Hydrogen is +1 with nonmetals (e.g., H₂O, NH₃) and -1 with metals (e.g., NaH, LiAlH₄).
Group IA metals have oxidation number +1 in compounds.
Group IIA metals have oxidation number +2 in compounds.
Oxygen is usually -2, except in peroxides (e.g., H₂O₂, O₂²⁻) and elemental forms.
Group VIIA nonmetals (halogens) often have oxidation number -1 in compounds.
The sum of oxidation numbers in a neutral compound is zero.
The sum of oxidation numbers in a polyatomic ion equals the ion’s charge.
Elements in the lower left of the periodic table are more likely to have positive oxidation numbers.

H₂O: 2(+1) + (-2) = 0 (neutral compound sum is zero)
SO₄²⁻: (+6) + 4(-2) = -2 (sulfur’s oxidation number is +6 to balance the ion charge)
SO₂: (+4) + 2(-2) = 0 (sulfur’s oxidation number is +4 to balance oxygen)

Oxidation number — the hypothetical charge an atom would have if all bonds were ionic.
Ion — atom or group with a net electric charge.
Group IA/IIA/VIIA — columns of the periodic table (alkali metals, alkaline earth metals, halogens).
Peroxide — compounds featuring an O-O single bond, often changing the typical oxidation state of oxygen.