Oxidation States
Oxidation States
Oxidation states, also known as oxidation numbers, are a concept in chemistry that provides a way to keep track of electrons in atoms as they form compounds. Understanding oxidation states is crucial for predicting the reactivity of elements, especially in redox (reduction-oxidation) reactions.
Definition
The oxidation state of an atom is a measure of the degree of oxidation of an atom. It is defined as the charge an atom would have if all bonds to atoms of different elements were 100% ionic.
Rules for Assigning Oxidation States
Several rules are used to assign oxidation states to each atom in a compound:
The oxidation state of an element in its standard state is zero.
- Examples: ( \text{H}_2, \text{O}_2, \text{Fe} )
For monoatomic ions, the oxidation state is equal to the charge of the ion.
- Examples: ( \text{Na}^+ ) has an oxidation state of +1, ( \text{Cl}^- ) has an oxidation state of -1.
Oxygen usually has an oxidation state of -2 in most of its compounds.
- Exceptions include peroxides (e.g., ( \text{H}_2\text{O}_2 )) where it is -1, and in compounds with fluorine (e.g., ( \text{OF}_2 )) where it is +2.
Hydrogen usually has an oxidation state of +1 when bonded to nonmetals and -1 when bonded to metals.
- Examples: ( \text{HCl} ) (H is +1), ( \text{NaH} ) (H is -1).
The sum of oxidation states in a neutral compound is zero.
- Example: In ( \text{H}_2\text{O} ), the oxidation state of O is -2, and H is +1, so ( 2(+1) + (-2) = 0 ).
The sum of oxidation states in a polyatomic ion is equal to the charge of the ion.
- Example: In ( \text{SO}_4^{2-} ), the oxidation state of O is -2, and S is +6, so ( 4(-2) + (+6) = -8 + 6 = -2 ), which is the charge of the sulfate ion.
Table of Common Oxidation States
| Element | Common Oxidation States | Example Compounds |
|---|---|---|
| Hydrogen | +1, -1 | ( \text{H}_2\text{O} ) (+1), ( \text{NaH} ) (-1) |
| Oxygen | -2, -1, +2 | ( \text{H}_2\text{O} ) (-2), ( \text{H}_2\text{O}_2 ) (-1), ( \text{OF}_2 ) (+2) |
| Nitrogen | -3, -2, -1, 0, +1, +2, +3, +4, +5 | ( \text{NH}_3 ) (-3), ( \text{NO} ) (+2), ( \text{NO}_2 ) (+4), ( \text{NO}_3^- ) (+5) |
| Carbon | -4, -3, -2, -1, 0, +1, +2, +3, +4 | ( \text{CH}_4 ) (-4), ( \text{CO} ) (+2), ( \text{CO}_2 ) (+4) |
| Iron | +2, +3 | ( \text{Fe}^{2+} ), ( \text{Fe}^{3+} ) |
Examples and Explanation
Example 1: ( \text{H}_2\text{O} )
In water (( \text{H}_2\text{O} )), oxygen has an oxidation state of -2, and hydrogen has an oxidation state of +1. Since there are two hydrogen atoms, the total positive oxidation state is +2, which balances the -2 from oxygen, resulting in a neutral molecule.
Example 2: ( \text{KMnO}_4 )
Potassium permanganate (( \text{KMnO}_4 )) is an ionic compound composed of ( \text{K}^+ ) and ( \text{MnO}_4^- ) ions. Potassium has an oxidation state of +1. Oxygen has an oxidation state of -2, and since there are four oxygen atoms, the total negative oxidation state from oxygen is -8. To balance the charge of the ( \text{MnO}_4^- ) ion, manganese must have an oxidation state of +7.
Example 3: ( \text{Fe}_2\text{O}_3 )
In iron(III) oxide (( \text{Fe}_2\text{O}_3 )), iron has an oxidation state of +3, and oxygen has an oxidation state of -2. There are two iron atoms, so the total positive oxidation state is +6. There are three oxygen atoms, so the total negative oxidation state is -6. The sum of the oxidation states is zero, which is consistent with a neutral compound.
Conclusion
Oxidation states are an essential tool in understanding the electronic structure of compounds and the changes that occur during chemical reactions. They are particularly useful in balancing redox reactions and in determining the reactivity of different elements within a compound. By following the rules for assigning oxidation states and practicing with examples, one can gain a solid understanding of this fundamental concept in chemistry.