Oxidation number Rules

The oxidation number rules for identifying the oxidant and reductant in a redox reaction seem complex at first but become quite simple with a small amount of practice and are an excellent tool for understanding the chemical reactions that occur all around us. Oxidation numbers are not charges even though many times they can be the same as ionic charges. They are an invented system for keeping track of electron loss and gain. We do not need to know anything about the bonding within the chemicals in order to be able to use these numbers effectively.

For each element or molecule that is involved in a reaction we need to follow these simple rules.

Rule 1. All pure substances have an oxidation number of zero. This applies to any pure substance whether it is a diatomic gas like O2 or a piece of pure metal like Iron (Fe). Examples of

Rule 2. In compounds, elements that usually have an ionic charge imparted by their position in a particular group have that same oxidation number. An example is Cl which is usually in the form Cl- in compounds; this will have an oxidation number of -1 in compounds.

Rule 3. When two or more usually negatively charged ions are involved in a compound, the one with the highest electronegativity value is given its ionic charge as the oxidation number; the others are worked out normally. An example is OF2. F is more electronegative, and so it is assigned the value of -1. Here is a good table of electronegativity values.

Rule 4. Oxygen in a compound always has an oxidation number of -2

Rule 5. Hydrogen in compounds always has an oxidation number of +1 except in the rare case of Metal Hydrides where it has a value of -1.

Rule 6. The oxidation numbers in the compound or molecule must total to the overall charge of that compound or molecule. For example CO2 has no overall charge and so the oxidation numbers must tally to zero. The sulfate ion, SO42-, has an overall charge of -2, so the oxidation numbers must tally to -2

Examples of Applying Oxidation Number Rules

These oxidation number rules only make sense when we start to use them. It is easiest to get the correct values if we break up molecules into their component atoms and line them up next to each other.

Example 1: Carbon Dioxide

Let's start with the very familiar Carbon Dioxide, CO2.

Carbon Dioxide contains one carbon atom and two oxygen atoms bound together in a discreet molecule. It is an uncharged molecule, meaning its overall charge value is zero.

First we draw each atom involved. Since there are two oxygen atoms in CO2, we put two of them in the drawing. We then make that collection of atoms equal to the overall charge which in this case is zero.

Next we go through the rules from 1 to 6, skipping any that are not relevant.

Rule 1: not relevant: CO2 is a compound and not a pure substance.

Rule 2: Oxygen is the most electronegative element and so it gets its standard charge of -2 for each oxygen atom.

Rule 3: not relevant: Carbon never forms negatively charged ions and even if it did it it is less electronegative than oxygen so what we have done is still good.

Rule 4: this has already been taken care of in Rule 2.

Rule 5: not relevant: there are no Hydrogen atoms involved.

Rule 6: In order to tally to zero, we MUST assign Carbon the oxidation number of +4. Therefore we can confidently state that after following the oxidation number rules, Carbon has an oxidation number of +4 in this compound.