Oxidation Numbers

Oxidation numbers describe the balance between electrons and protons on an atom, whether that atom is happily floating around all by itself or part of a massive molecule. Oxidation numbers can be determined two different ways: by using rules, or by looking at structure. Let's look at the rules first.

Oxidation Numbers by the Rules:
1. For neutral, uncombined elements, Ox# = 0. Examples: Fe(s), H₂(g), Hg(l), Ne(g)
2. For monoatomic ions, Ox# = charge. Examples: Fe²⁺(aq) {Ox# = +2}, P^3-(g) {Ox# = -3}
3. Oxygen is almost always Ox# = -2, except in O₂ {Ox# = 0, Rule #1} and peroxides {Ox# = -1}
4. Hydrogen is almost always Ox# = +1, except in H₂ {Ox# = 0, Rule #1} and hydrides {Ox# = -1}
5. The sum of the Ox#s on all the atoms in a polyatomic molecule or ion is equal to the charge on the whole polyatomic molecule or ion.

Let's look at a redox reaction and assign Ox#s by the rules:

Cl₂(g) + 2 O₂(g) ↔ 2 ClO₂(g)

Cl₂(g) : Rule #1, Ox# = 0

O₂(g) : Rule #1, Ox# = 0

ClO₂(g) : Rule #3, oxygen is Ox# = -2.

ClO₂(g) : Rule #5, (Ox# Cl) + 2(Ox# O) = 0 (the charge on a neutral molecule)

(Ox# Cl) + 2(-2) = 0

(Ox# Cl) = +4

So in this redox reaction, each Cl is going from 0 to +4, losing 4 electrons, Losing Electrons is Oxidation; and each O is going from 0 to -2, gaining 2 electrons, Gaining Electrons is Reduction.

For many substances, it's actually easier to assign Ox#s by looking at the structure. The process is very similar to finding Formal Charge, the electrons are just assigned a little differently. Formal Charge assigns electrons as if all bonds are purely covalent, meaning that all bonding pairs of electrons are split with one electron given to each atom in the bond. Oxidation Number assigns electrons as if all bonds are purely ionic, meaning that all of the bonding electrons go to the more electronegative element in the bond.

Oxidation Number by the Structure:
1. Draw a good Lewis Structure
2. Assign all bonding electrons to the more electronegative element in the bond
3. Compare the electrons assigned to each atom to the valence electrons of the neutral element

Again, let's look at an example, in fact, let's look at the same example as above, ClO₂(g). Drawing a good Lewis Structure:

This is a very interesting molecule, it violates the octet rule and it has an unpaired electron. Looks like it would be pretty reactive. Looking at the electronegativities, oxygen is more electronegative than chlorine, so all of the bonding electrons will be assigned to oxygen, giving each oxygen 8 assigned electrons and the chlorine 3 assigned electrons.

Neutral oxygen has 6 valence electrons, we've assigned 8 electrons to oxygen, so the oxidation number for oxygen in this molecule is -2, just like we predicted using the rules. Neutral chlorine has 7 valence electrons, we've assigned 3 electrons to chlorine, so the oxidation number for chlorine in this molecule is +4, again, just like we predicted using the rules. Both methods work. Why would you ever use structures when the rules work? Try looking at hydrogen peroxide, H-O-O-H, or a more complex molecule like glucose. The rules don't always give the best picture of what's happening in a molecule, and this can ultimately make it harder to predict reactivity or other behaviors.