Electrons can be lost or gained when some atoms interact with other atoms and bond or react with them. Why are oxidation numbers important in this context?
Oxidation numbers are used by chemists to deduce and keep track of the number of electrons transferred or shared during chemical reactions. Oxidation numbers are also useful to chemists when it comes to naming inorganic compounds.
Firstly, we will define the term oxidation number.
Then, we will look at the oxidation number rules, as well as their exceptions.
After that, we will explore how oxidation numbers relate to naming compounds.
Finally, we'll have a go at oxidation number calculations for various compounds and ions.
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 numbers are numbers assigned to ions that show how many electrons the ion has lost or gained, compared to the element in its uncombined state. A positive oxidation number shows that the element lost electrons, whilst a negative oxidation number shows that it gained electrons. They can also be referred to as oxidation states.
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.
- 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 CO32-, 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:
Look at the charge of the ion or compound, if any. This will help you know what you are aiming for.
Identify any atoms with fixed oxidation states.
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.
What are the oxidation numbers of sulphur in the following compounds and ions?
- S8
- H2S
- SO32-
- H2SO4
a. Because this is an uncombined element, the oxidation number of sulphur in S8 is 0.
b. H2S is a neutral compound, and so the overall sum of all the oxidation numbers is zero. Each hydrogen ion has an oxidation number of +1. Therefore, sulphur must have the oxidation number -2, as 2(1) + (-2) = 0.
c. The overall charge on the SO32- ion is -2. Therefore, the sum of the oxidation numbers must equal -2. Each oxygen has an oxidation number of -2, and so their combined total is 3(-2) = -6. This means that the oxidation number of sulphur must be +4, as (-6) + 4 = -2
d. Once again, H2SO4 is a neutral compound and so the sum of all the oxidation numbers must equal zero. There are four oxygens, each with an oxidation number of -2, and so their combined total is 4(-2) = -8. There are two hydrogens, each with an oxidation number of +1, and so their combined total is 2(1) = 2. Therefore, the oxidation number of sulfur must be +6, as (-8) + 2 + (+6) = 0.
Oxidation Number - Key takeaways
- Oxidation numbers are numbers assigned to ions that show how many electrons the ion has lost or gained, compared to the element in its uncombined state.
- There are certain rules to follow when assigning oxidation numbers:
- The oxidation number of all uncombined elements is zero.
- The sum of the oxidation numbers in an ion is equal to the ionic charge.
- The oxidation number of a neutral compound is zero.
- In an ion or a compound, the more electronegative element is given the more negative oxidation number.
- Some elements always take certain oxidation states, although there are exceptions to the general rules.
- Roman numerals and compound prefixes and suffixes give us clues about the oxidation numbers of the elements involved.
- We can work out oxidation numbers using chemical formulae and the rules listed above.
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