Copper was oxidized because its electrons were accepted by an oxidizing agent, nitrogen (or nitrate ion).Īlthough they are useful and necessary for recognizing redox reactions, oxidation numbers are a highly artificial device. The nitrogen was reduced by electrons donated by copper, and so copper was the reducing agent. Gained one electron, so 2 e – were needed for the 2 NOģ –. The oxidation number of nitrogen went down from 5 to 4, and so the nitrogen (or nitrate ion) was reduced.

Since the oxidation number of copper increased from 0 to +2, we say that copper was oxidized and lost two negatively charged electrons. Applying the oxidation number rules to the following equation, we have: Oxidation corresponds to increasing the oxidation number of some atom.

As a general rule, reduction corresponds to a lowering of the oxidation number of some atom. This arbitrarily assigned gain of one electron corresponds to reduction of the nitrogen atom on going from NO 3 – to NO 2. In NO 2, on the other hand, the nitrogen has an oxidation number of + 4 and may be thought of as having one valence electron for itself, that is, one more electron than it had in NO 3 –. This arbitrary assignment corresponds to the nitrogen’s having lost its original five valence electrons to theĮlectronegative oxygens. This is done by assigning oxidation numbers to each atom before and after the reaction.įor example, in NO 3 – the nitrogen is assigned an oxidation number of +5 and each oxygen an oxidation number of –2. In order to be able to recognize redox reactions, we need a method for keeping a careful account of all the electrons.

Redox reactions may involve proton transfers and other bond-breaking and bond-making processes, as well as electron transfers, and therefore the equations involved are much more difficult to deal with than those describing acid-base reactions.

Assigning oxidation numbers