Oxidation Numbers

From MyMCAT

Introduction

In redox (oxidation/reduction) reactions, it is critical to determine how the electrons around molecules are changing. Is this compound being reduced, is this reactant an oxidizing agent? How much does the charge change in this atom? etc. Oxidation numbers answer these questions and can be calculated systematically using a few simple rules.

For all compounds, whether covalent, polar covalent, or ionic, we treat them as ionic when counting electrons. In other words, we can break all the complex molecules down into parts when assigning and calculating oxidation numbers.

Known Oxidation Numbers

While the rules themselves are simple, one must also know what the oxidation numbers are for some basic elements.

Atoms as Elements ( H₂,N₂,O₂,p₄, etc)
The oxidation number of these molecules is always equal to 0.

Monoatomic Ions
Cations ( Na⁺,Ca²⁺,Fe³⁺, etc The oxidation number of these ions are always equal to their charge. (+1, +2, +3 respectively). Or one can consider this as, their oxidation number is equal to their group number. (Na is in the first group, therefore +1, and so on.)

Anions( Cl^-, Br^-, etc)
The oxidation number of these ions are always equal to their charge. Or one can consider this as, their oxidation number is equal to their group number - 8. (Cl is in group 8, then subtract 7, and you get -1.)

Hydrogen Combined with Nonmetals (NH₃, H₂O, HCl)
Hydrogen has an oxidation number of +1.
Combined with Metals (NaH, CaH₂)
Hydrogen has an oxidation number of -1.

Oxygen
Oxygen has an oxidation number of -2, except when it exists in peroxides (O₂^-2).

Basic Rules

Rule 1: The sum of the oxidation numbers of all the atoms in a compound equals the charge on the compound. Neutral compounds: Sum of oxidation numbers = 0 Ionic species: Sum of oxidation numbers = charge of the ion

Rule 2: In Binary Compounds, the more Electronegative (EN) element is assigned the negative oxidation number. (Remember, electronegativity increases towards the top/right corner of the periodic table.)

Rule 3: Some oxidation numbers can not exist regardless of the compound the atom is in. The valid range is:

Maximum oxidation number possible = + Group number. Minimum oxidation number possible = (Group number - 8) (this number will be negative)

Examples

Basic Examples

Al₂O₃ This is a neutral compound so the sum of the oxidation states is zero. Oxygen has an oxidation number of -2 and their are three oxygens. So, the sum of the Al component must be equal and opposite to the oxygen component for the overall number to be zero. In other words 2x(the oxidation number of Al) + x(-2 for oxygen) = 0. If one works this out, Al must be +3. 2x(+3) + 3x(-2) = 0.

CO Again, the sum will equal 0 since it is a neutral molecule. O will have an oxidation number equal to -2. C, therefore, must be +2 for the sum to be zero.
1 C + 1 O = 0
(Oxy# of C) + (-2) = 0
C = +2

H₂O₂ H is found with a nonmetal in this case, thus its oxidation number must equal +1. The species is neutral, thus the overall oxidation number must equal zero.
2(H) + 2(O) = 0
2(+1) + 2(O) = 0
2(O) = -2
(O) = -1.

Notice, that as mentioned above, Oxygen was in a peroxide state and so the normal oxidation number of -2 was not observed.

Difficult Examples

K₂Cr₂O₇ The species is neutral, therefore the sum is zero. K from the first group, so we expect it to have an oxidation number of +1.
2(K) + (Cr₂O₇) = 0
2(+1) + (Cr₂O₇) = 0
(Cr₂O₇) = -2

So, Cr₂O₇, has an oxidation number of -2. This is expected as this ion has a charge of -2 on its own in aqueous conditions. Cr₂O₇ = -2 Oxygen generally has an oxidation number of -2, and there are 7 of them thus:
2(Cr) + 7(O) = -2
2(Cr) + 7(-2) = -2
2(Cr) = +12
Cr = +6.

Thus, Cr, must have an oxidation number of +6.