Introduction
Oxidation numbers (also called oxidation states) are used to track electron transfer in redox reactions. They help us determine which atoms are oxidized and which are reduced, even in complex compounds.
Key Concepts
- Definition:
An oxidation number is the hypothetical charge an atom would have if all bonds were 100% ionic.
- General Rules for Assigning Oxidation Numbers:
|
Rule |
Explanation |
Example |
|
Elements in free state |
Always 0 |
O₂, H₂, Na → 0 |
|
Monoatomic ion |
Same as the charge |
Na⁺ = +1, Cl⁻ = -1 |
|
Group 1 elements |
Always +1 |
Na, K |
|
Group 2 elements |
Always +2 |
Mg, Ca |
|
Hydrogen (H) |
+1 with non-metals, -1 with metals |
H₂O = +1, NaH = -1 |
|
Oxygen (O) |
Usually -2, but -1 in peroxides |
H₂O = -2, H₂O₂ = -1 |
|
Fluorine (F) |
Always -1 |
F₂, HF |
|
Sum of oxidation numbers in a compound |
Always 0 |
H₂O → 2(+1) + (-2) = 0 |
|
Sum in a polyatomic ion |
Equals the ion charge |
SO₄²⁻ → total = -2 |
- Identifying Redox Changes with Oxidation Numbers
A substance is oxidized if its oxidation number increases
A substance is reduced if its oxidation number decreases
Worked Example 1:
Reaction:
Fe + Cl₂ → FeCl₃
Assign oxidation numbers:
- Fe: 0 → +3 (oxidized)
- Cl: 0 → -1 (reduced)
Fe is the reducing agent, Cl₂ is the oxidizing agent
Worked Example 2 (NECTA-style):
Question:
Assign oxidation numbers and identify what is oxidized and what is reduced in:
2H₂S + SO₂ → 3S + 2H₂O
Solution:
- H in H₂S: +1
- S in H₂S: -2
- S in SO₂: +4
- S (elemental): 0
- O in H₂O: -2
Changes:
- S in H₂S: -2 → 0 → oxidized
- S in SO₂: +4 → 0 → reduced
Common NECTA Tip:
Always show oxidation numbers above atoms in redox equations to explain your answer clearly.
Applications of Oxidation Numbers
- Balancing redox equations
- Predicting reactions
- Identifying agents in complex reactions
Summary
- Oxidation numbers help track electron movement
- Use rules to assign correct values
- Changes in oxidation number indicate redox activity