Chemical reactions that involve the transfer of electrons are called oxidation-reduction (or redox) reactions. In this post, we will understand what redox reactions are, we will define Oxidation, Reduction, and Redox. Then we will focus on oxidation numbers and the rules for assigning oxidation numbers.
Understanding Redox reactions
To understand redox reactions, let’s consider this chemical reaction:
2Na(s) + Cl2(g) → 2NaCl
The reactants are elements, and it is assumed that they are electrically neutral; they have the same number of electrons as protons.
The product, however, is ionic; it is composed of Na+ and Cl− ions.
Somehow, the individual sodium atoms as reactants had to lose an electron to make the Na+ ion, while the chlorine atoms as reactants had to each gain an electron to become the Cl− ion. This reaction involves the transfer of electrons between atoms.
In reality, electrons are lost by some atoms and gained by other atoms simultaneously. However, mentally we can separate the two processes.
Oxidation, Reduction & Redox – definitions
Oxidation is defined as the loss of one or more electrons by an atom. Reduction is defined as the gain of one or more electrons by an atom. In reality, oxidation and reduction always occur together; it is only mentally that we can separate them. Chemical reactions that involve the transfer of electrons are called oxidation-reduction (or redox) reactions.
Redox reactions require that we keep track of the electrons assigned to each atom in a chemical reaction. And to do that we use an artificial count called the oxidation number to keep track of electrons in atoms.
Oxidation numbers are assigned to atoms based on a series of rules. Oxidation numbers are not necessarily equal to the charge on the atom; we must keep the concepts of charge and oxidation numbers separate.
The rules for assigning oxidation numbers to atoms
The rules for assigning oxidation numbers to atoms are as follows:
- Atoms in their elemental state are assigned an oxidation number of 0.
- Atoms in monatomic (i.e., one-atom) ions are assigned an oxidation number equal to their charge. Oxidation numbers are usually written with the sign first, then the magnitude, which differentiates them from charges.
- In compounds, fluorine is assigned a −1 oxidation number; oxygen is usually assigned a −2 oxidation number (except in peroxide compounds [where it is −1] and in binary compounds with fluorine [where it is positive]); and hydrogen is usually assigned a +1 oxidation number (except when it exists as the hydride ion, H−, in which case rule 2 prevails).
- In compounds, all other atoms are assigned an oxidation number so that the sum of the oxidation numbers on all the atoms in the species equals the charge on the species (which is zero if the species is neutral).
Some examples of Oxidation Number assignment
Let us work through a few examples of oxidation number assignments for practice.
In H2, both hydrogen atoms have an oxidation number of 0, by rule 1.
In NaCl, sodium has an oxidation number of +1, while chlorine has an oxidation number of −1, by rule 2.
In H2O, the hydrogen atoms each have an oxidation number of +1, while the oxygen has an oxidation number of −2, even though hydrogen and oxygen do not exist as ions in this compound as per rule 3.
By contrast, by rule 3 in hydrogen peroxide (H2O2), each hydrogen atom has an oxidation number of +1, while each oxygen atom has an oxidation number of −1.
We can use rule 4 to determine oxidation numbers for the atoms in SO2. Each oxygen atom has an oxidation number of −2; for the sum of the oxidation numbers to equal the charge on the species (which is zero), the sulfur atom is assigned an oxidation number of +4. Does this mean that the sulfur atom has a 4+ charge on it? No, it only means that the sulfur atom is assigned a +4 oxidation number by our rules of apportioning electrons among the atoms in a compound.