Oxidation States - Complete Explanation
Oxidation state (or oxidation number) is a concept used in chemistry to describe the degree of
oxidation of an atom in a chemical compound. It represents the hypothetical charge an atom would
have if all bonds were purely ionic.
Rules for Assigning Oxidation States:
1. The oxidation state of a free element (uncombined element) is always 0.
2. The oxidation state of a monoatomic ion is equal to its charge.
3. Hydrogen is usually +1, except in metal hydrides where it is -1.
4. Oxygen is usually -2, except in peroxides (-1) and superoxides (-1/2).
5. Fluorine is always -1 in compounds.
6. The sum of oxidation states in a neutral compound is 0.
7. The sum of oxidation states in a polyatomic ion equals its charge.
Common Oxidation States of Elements:
Hydrogen: +1 (except in hydrides, -1)
Oxygen: -2 (except peroxides, -1)
Alkali metals (Group 1): +1
Alkaline earth metals (Group 2): +2
Halogens: Usually -1 (except when bonded with oxygen or fluorine)
Transition metals: Variable oxidation states (e.g., Fe2+, Fe3+, Cu+, Cu2+)
Redox Reactions and Balancing:
A redox reaction involves both oxidation (loss of electrons) and reduction (gain of electrons). To
balance redox reactions:
1. Identify oxidation states of elements.
2. Determine which elements are oxidized and reduced.
3. Balance elements and charges using electrons.
4. Use the half-reaction method if necessary.
Oxidation state (or oxidation number) is a concept used in chemistry to describe the degree of
oxidation of an atom in a chemical compound. It represents the hypothetical charge an atom would
have if all bonds were purely ionic.
Rules for Assigning Oxidation States:
1. The oxidation state of a free element (uncombined element) is always 0.
2. The oxidation state of a monoatomic ion is equal to its charge.
3. Hydrogen is usually +1, except in metal hydrides where it is -1.
4. Oxygen is usually -2, except in peroxides (-1) and superoxides (-1/2).
5. Fluorine is always -1 in compounds.
6. The sum of oxidation states in a neutral compound is 0.
7. The sum of oxidation states in a polyatomic ion equals its charge.
Common Oxidation States of Elements:
Hydrogen: +1 (except in hydrides, -1)
Oxygen: -2 (except peroxides, -1)
Alkali metals (Group 1): +1
Alkaline earth metals (Group 2): +2
Halogens: Usually -1 (except when bonded with oxygen or fluorine)
Transition metals: Variable oxidation states (e.g., Fe2+, Fe3+, Cu+, Cu2+)
Redox Reactions and Balancing:
A redox reaction involves both oxidation (loss of electrons) and reduction (gain of electrons). To
balance redox reactions:
1. Identify oxidation states of elements.
2. Determine which elements are oxidized and reduced.
3. Balance elements and charges using electrons.
4. Use the half-reaction method if necessary.
