Oxidation Number Rules
Assign oxidation states systematically
Understanding Oxidation Numbers
Oxidation numbers (also called oxidation states) are a bookkeeping tool that tracks electron distribution in chemical compounds and helps identify which atoms are oxidized or reduced during chemical reactions. While oxidation numbers don't represent actual charges on atoms in covalent compounds, they provide a systematic method for analyzing electron transfer in redox reactions. The concept dates to the early 20th century and has become indispensable for balancing redox equations, predicting reaction products, understanding electrochemistry, and organizing chemical knowledge about element reactivity patterns.
The rules for assigning oxidation numbers follow a priority system that resolves conflicts when multiple rules could apply. For example, in most compounds oxygen has oxidation number -2, but this rule is overruled by the fluorine rule (fluorine always -1) in OF₂, where oxygen becomes +2. Understanding this priority hierarchy is crucial for correctly assigning oxidation states. The most fundamental rule is that oxidation numbers must sum to the overall charge: zero for neutral molecules, or the ion charge for polyatomic ions. This constraint allows calculation of unknown oxidation states through algebraic equations.
Oxidation numbers reveal patterns in periodic table chemistry: Group 1 metals always +1, Group 2 always +2, while transition metals exhibit variable oxidation states (Fe can be +2 or +3, Mn ranges from -1 to +7). These patterns help predict compound formulas and reaction outcomes. In organic chemistry, tracking carbon oxidation states helps classify reactions: oxidation increases oxygen/decreases hydrogen content (alcohols → aldehydes → acids), while reduction does the opposite (aldehydes → alcohols). Mastering oxidation numbers is essential for understanding redox chemistry, electrochemistry, corrosion, metabolism, and industrial chemical processes.
Priority Rules (Apply in Order)
- Free elements: Oxidation number = 0 (e.g., Na, O₂, Cl₂)
- Monatomic ions: Oxidation number = ion charge (e.g., Na⁺ = +1, Cl⁻ = -1)
- Fluorine: Always -1 in compounds
- Oxygen: Usually -2 (exceptions: peroxides -1, OF₂ +2)
- Hydrogen: +1 with nonmetals, -1 with metals (hydrides)
- Group 1 metals: Always +1
- Group 2 metals: Always +2
- Sum rule: Sum of all oxidation numbers = overall charge
Example: H₂SO₄
Compound is neutral, so sum = 0.
H: +1 (rule 5) → 2 H = +2 total
O: -2 (rule 4) → 4 O = -8 total
Let S = x: 2(+1) + x + 4(-2) = 0
2 + x - 8 = 0 → x = +6
Answer: S has oxidation number +6
Additional Worked Example: Dichromate Ion (Cr₂O₇²⁻)
Step 1: Identify the overall charge: -2
Step 2: Oxygen usually -2 (rule 4), so 7 O = 7(-2) = -14 total
Step 3: Let Cr oxidation state = x for each Cr atom
Step 4: Sum rule: 2(x) + 7(-2) = -2
Step 5: Solve: 2x - 14 = -2 → 2x = 12 → x = +6
Answer: Each Cr has oxidation number +6
Cr₂O₇²⁻ is a powerful oxidizing agent, as Cr⁺⁶ readily accepts electrons to become Cr⁺³.
Key Applications
Balancing Redox Equations
Identify oxidation/reduction by tracking oxidation number changes. Oxidation = increase in oxidation number; reduction = decrease.
Predicting Compound Formulas
Use common oxidation states to predict formulas: Fe⁺³ + O⁻² → Fe₂O₃ (charges must balance).
Identifying Redox Reactions
If oxidation numbers change, it's a redox reaction. If no change, it's not redox (e.g., precipitation, acid-base).
Common Mistakes
- Forgetting peroxide exception (H₂O₂: O is -1, not -2).
- Not accounting for overall charge in polyatomic ions.
- Applying rules out of priority order.