Oxidation Number

What are oxidation numbers?

In "Redox", you learned that many reactions involve a movement of electrons. One species loses electrons and is oxidised, whilst another gains electrons and is reduced. Overall, we call these processes redox reactions. Oxidation numbers help us keep track of which species is oxidised and which species is reduced in such a reaction.

Oxidation number rules

There are a few rules that can help and simplify the way we work out oxidation numbers.

• The oxidation number of all uncombined elements is 0. The reason behind this is that the element has neither lost any electrons, nor gained any, and is therefore neutral.
  • e.g. Zn, H, and Cl.

• The sum of the oxidation numbers of all the atoms or ions in a neutral compound equals 0.
  • e.g. In NaCl, the oxidation number of Na is +1 and the oxidation number of Cl is -1. These add up to make 0.

• The sum of the oxidation numbers in an ion equals the charge on the ion. This applies to monatomic ions as well as complex ions.
  • e.g. The oxidation number of the monatomic ion F^- is -1.
  • e.g. In the ion CO₃^2-, C has an oxidation number of +4 and the three O each have an oxidation number of -2. 4 + 3(-2) = -2, which is the charge on the ion.

• In an ion or a compound, the more electronegative element generally has the more negative oxidation number. Remember that electronegativity decreases down a group and increases across a period.
  • e.g. In F₂O, F is more electronegative than oxygen, and so takes the more negative oxidation number. Here, F has an oxidation number of -1 and O has an oxidation number of +2.

Lots of elements have the same oxidation number in all of their compounds:

• Group 1 elements all have the oxidation number +1.
• Group 2 elements all have the oxidation number +2.
• Aluminium always has the oxidation number +3.
• Fluorine always has the oxidation number -1.
• Hydrogen usually has the oxidation number +1, except in metal hydrides.
• Oxygen usually has the oxidation number -2, except in peroxides and in compounds with fluorine.
• Chlorine usually has the oxidation number -1, except in compounds with oxygen and fluorine.

Periodic table with oxidation numbers

To help with working out the oxidation numbers of different compounds, here is an image of the periodic table with the common oxidation numbers per group.

A periodic table with the oxidation numbers of the elements within their groups - StudySmarter Originals

However, you must always remember the exceptions to the oxidation number rules. We'll look at these in more detail next.

Oxidation number exceptions

As we've learned, there are a few exceptions to the oxidation numbers of elements within compounds.

Oxidation number exceptions: Hydrogen

Hydrogen usually has an oxidation number of +1. But in metal hydrides, such as NaH or KH, it has an oxidation number of -1. This is because we know that the sum of the oxidation numbers in a neutral compound is always 0, and that group 1 metals always have an oxidation number of +1. This means that in a metal hydride, hydrogen must have an oxidation state of -1, as 1 + (-1) = 0. For example, in NaH, Na has an oxidation state of +1 and H has an oxidation state of -1.

Oxidation number exceptions: Oxygen

Oxygen usually has an oxidation number of -2. But in peroxides, such as H₂O₂, it has an oxidation number of -1. Once again, this is a neutral compound, and therefore the sum of the oxidation numbers must be zero. For example, in the case of H₂O₂, each hydrogen atom has the oxidation number +1, so each oxygen atom must have the oxidation number -1.

Oxygen also deviates from its usual oxidation number in compounds with fluorine. This is because we know that the more electronegative element takes the more negative oxidation number, and fluorine is more electronegative than oxygen. For example, in F₂O, the more electronegative element is fluorine, so it gains the negative oxidation number -1. We have two fluorines for every oxygen, and so the oxidation number of oxygen is +2.

Oxidation number exceptions: Chlorine

Likewise, chlorine takes variable oxygen numbers in compounds with oxygen or fluorine. Once again, this is because oxygen and fluorine are more electronegative than chlorine. For example, in HClO, O is the most electronegative element and so takes the most negative oxidation number. Here, it has the oxidation number of -2. H isn't in a metal hydride and so has an oxidation number of +1. This means that Cl must also have an oxidation number of +1, as 1 + 1 + (-2) = 0.

Oxidation numbers and naming compounds

Although we've just learned some rules for assigning oxidation numbers, they don't cover every element. In fact, many elements can take numerous possible oxidation numbers, which can cause confusion in many compounds. Here are some tips to help you.

Oxidation numbers and naming compounds: Roman numerals

If there is any risk of ambiguity, the specific oxidation number of an element in a given compound is shown using Roman numerals. However, this only applies to positive oxidation states. For example, iron (II) sulphate (FeSO₄) contains iron ions with an oxidation number of +2, whilst iron (III) sulphate (Fe₂(SO₄)₃) contains iron ions with an oxidation number of +3.

Oxidation numbers and naming compounds: Prefixes and suffixes

We can also use prefixes and suffixes to give information about the formula of a compound, which helps us work out each element's oxidation state:

• Compounds containing oxygen end in -ate or -ite. There's a difference between the two: the -ate compound always has one more oxygen than the -ite compound. If we encounter a compound with one more oxygen than the -ate compound, we add the prefix per-. If we encounter a compound with one fewer oxygen than the -ite compound, we add the prefix hypo-.
  • e.g. The perchlorate ion (HClO₄⁻) has 4 oxygens, the chlorate ion (ClO₃⁻) has three, the chlorite ion (ClO₂⁻) has two and the hypochlorite ion (ClO⁻) has just one.

• Inorganic acids containing oxygen end in -ic.
  • e.g. Sulphuric acid (H₂SO₄).

Oxidation number calculation examples

The sum of all the oxidation states in a neutral compound must add up to zero, and the sum of all the oxidation numbers in a complex ion must add up to the charge of the ion - we know this from our rules for assigning oxidation numbers. But how do we work out the oxidation numbers of the individual elements within the compound or ion? For this, we can apply our knowledge of fixed oxidation numbers and work out the unknown oxidation numbers by deduction.

It can help to follow this process:

1. Look at the charge of the ion or compound, if any. This will help you know what you are aiming for.

2. Identify any atoms with fixed oxidation states.

3. Deduce the oxidation states of the remaining atoms, making sure the sum of all the oxidation states adds up to the charge of the ion or compound.

It's now your turn: Have a go at working out the oxidation numbers of some elements using the rules we covered above. If you get stuck, we'll work through the solutions together.