What are the rules for assigning oxidation numbers?

The rules for assigning oxidation numbers help us keep track of electron transfer in chemical reactions. Here are the main rules written simply:

Rules for Assigning Oxidation Numbers

1. Pure elements → The oxidation number of an atom in its natural form is 0.
   • Example: O₂, H₂, Na, Cl₂.
2. Monatomic ions → The oxidation number is the same as the ion’s charge.
   • Example: Na⁺ is +1, Cl⁻ is –1, Mg²⁺ is +2.
3. Oxygen → Usually –2 in compounds.
   • Exception: In peroxides (–1), superoxides (–½), and OF₂ (+2).
4. Hydrogen → Usually +1 when bonded to nonmetals, but –1 when bonded to metals in hydrides.
5. Fluorine → Always –1 in compounds (most electronegative element).
   • Other halogens (Cl, Br, I) are usually –1, unless bonded to oxygen or fluorine.
6. Alkali metals (Group 1) → Always +1 in compounds.
   • Example: NaCl, K₂O.
7. Alkaline earth metals (Group 2) → Always +2 in compounds.
   • Example: CaCl₂, MgO.
8. Aluminum → Always +3 in compounds.
9. Neutral compounds → The sum of all oxidation numbers is 0.
10. Polyatomic ions → The sum of all oxidation numbers equals the ion’s overall charge.

• Example: In sulfate (SO₄²⁻), the total is –2.

In short: Oxidation numbers are assigned based on element type, usual charges, and must balance out to the charge of the compound or ion.