Oxidation Number: The Electron Accounting System
The Fundamental Rules of Electron Bookkeeping
To assign oxidation numbers consistently, chemists follow a set of rules. Think of these as the guidelines for our electron accounting system. The rules are hierarchical, meaning if two rules conflict, the higher-numbered rule takes precedence.
| Rule # | Description | Example |
|---|---|---|
| 1 | The oxidation number of an atom in its elemental form is always 0. | Na (sodium metal), O$_2$ (oxygen gas), Cl$_2$ (chlorine gas) all have an oxidation number of 0. |
| 2 | The oxidation number of a monatomic[2] ion is equal to the charge of the ion. | In Na$^+$, the oxidation number is +1. In Cl$^-$, it is -1. |
| 3 | Fluorine always has an oxidation number of -1 in its compounds. | In NaF and CF$_4$, fluorine is -1. |
| 4 | Oxygen usually has an oxidation number of -2. Exceptions are peroxides (like H$_2$O$_2$) where it is -1, and when bonded to fluorine. | In H$_2$O and CO$_2$, oxygen is -2. In H$_2$O$_2$, it is -1. |
| 5 | Hydrogen is usually +1 when bonded to non-metals and -1 when bonded to metals. | In H$_2$O, hydrogen is +1. In NaH, it is -1. |
| 6 | Group 1 metals (Alkali metals) are always +1 and Group 2 metals (Alkaline earth metals) are always +2 in their compounds. | In NaCl, sodium is +1. In MgO, magnesium is +2. |
| 7 | The sum of oxidation numbers in a neutral compound is zero. In a polyatomic[3] ion, the sum equals the ion's charge. | For H$_2$O, the sum is 0. For the sulfate ion SO$_4^{2-}$, the sum is -2. |
Putting the Rules into Practice: A Step-by-Step Guide
Let's use the rules to find the oxidation number of sulfur in the sulfate ion, SO$_4^{2-}$.
Step 2: Set up an equation. Let the oxidation number of sulfur (S) be x.
Step 3: Apply Rule 7: The sum of all oxidation numbers equals the ion's charge.
(S) + 4*(O) = -2
x + 4*(-2) = -2
x - 8 = -2
Step 4: Solve for x.
x = -2 + 8
x = +6
Therefore, the oxidation number of sulfur in SO$_4^{2-}$ is +6.
Oxidation Numbers in Action: From Rust to Respiration
Oxidation numbers are not just for homework; they help us understand the world around us. The most important application is in identifying redox reactions.
What is a Redox Reaction? It's a chemical reaction where electrons are transferred between species. This transfer is tracked by changes in oxidation numbers.
- Oxidation is an increase in oxidation number (loss of electrons).
- Reduction is a decrease in oxidation number (gain of electrons).
A simple mnemonic is "OIL RIG": Oxidation Is Loss, Reduction Is Gain.
Example 1: The Formation of Rust
When iron rusts, it reacts with oxygen and water. The key change is the oxidation of iron:
Elemental iron (Fe) has an oxidation number of 0 (Rule 1). In rust (Fe$_2$O$_3$), we can calculate its oxidation number. Let y be the oxidation number of Fe. Oxygen is -2.
2y + 3*(-2) = 0 (Rule 7 for a neutral compound)
2y - 6 = 0
2y = 6
y = +3
The oxidation number of iron increased from 0 to +3. Iron was oxidized (it lost electrons).
Example 2: Cellular Respiration
The process that powers our cells involves the oxidation of glucose (C$_6$H$_{12}$O$_6$). The carbon in glucose is mostly in a zero or low oxidation state. During respiration, it is converted to carbon dioxide (CO$_2$), where carbon has an oxidation number of +4. This increase in oxidation number shows that carbon is being oxidized, releasing energy that our bodies use.
Important Questions
Yes, but only if it is in its elemental form within the compound. This is rare. A more common scenario is when the average oxidation number of an element in a compound is zero, which happens in molecules where the same element is bonded to itself, like in hydrogen peroxide (H$_2$O$_2$). The two oxygen atoms are bonded to each other, so the oxidation number for each oxygen is -1, not zero. True zero oxidation states in compounds are found in coordination complexes[4] where a metal atom is surrounded by neutral molecules like carbon monoxide (CO).
This is an advanced but important distinction. Both are bookkeeping tools, but they are calculated differently. The oxidation number assumes all bonds are 100% ionic. The formal charge assumes all bonds are 100% covalent (shared equally). Formal charge is calculated as: (Number of valence electrons in a free atom) - (Number of lone pair electrons + 1/2 the number of bonding electrons). Oxidation number is more useful for tracking electron transfer in redox reactions, while formal charge is better for predicting the most stable structure of a molecule.
Many elements, especially transition metals like iron (Fe), copper (Cu), and manganese (Mn), can have multiple oxidation numbers. This is because the energy difference between their inner electron shells is small, allowing them to lose different numbers of electrons depending on the chemical environment. For example, iron can be +2 (as in FeO) or +3 (as in Fe$_2$O$_3$). This versatility is what makes them crucial in biological systems and industrial catalysts.
Footnote
[1] Redox Reaction: A shorthand for reduction-oxidation reaction, a chemical process in which the oxidation states of atoms are changed through the transfer of electrons.
[2] Monatomic: Consisting of a single atom (e.g., Na$^+$, Cl$^-$).
[3] Polyatomic: Consisting of multiple atoms (e.g., SO$_4^{2-}$, NH$_4^+$).
[4] Coordination Complex: A structure consisting of a central metal atom or ion bonded to a surrounding array of molecules or anions, known as ligands.