The terms "oxidation number" and "oxidation state" are used interchangeably to indicate the apparent charge of an atom within a species, as if all atoms are ionic.
The IUPAC does differentiate between the two terms, representing oxidation numbers using Roman numerals and oxidation states with numbers preceded by a sign.
Syllabus reference S3.1.6Structure 3.1.6 - The oxidation state is a number assigned to an atom to show the number of electrons transferred in forming a bond. It is the charge that atom would have if the compound were composed of ions.
- Deduce the oxidation states of an atom in an ion or a compound.
- Explain why the oxidation state of an element is zero.
Guidance
- Oxidation states are shown with a + or – sign followed by the Arabic symbol for the number, e.g. +2, –1. Examples should include hydrogen in metal hydrides (–1) and oxygen in peroxides (–1).
- The terms “oxidation number” and “oxidation state” are often used interchangeably, and either term is acceptable in assessment.
- Naming conventions for oxyanions use oxidation numbers shown with Roman numerals, but generic names persist and are acceptable. Examples include NO3– nitrate, NO2– nitrite, SO42– sulfate, SO32– sulfite.
Tools and links
- Reactivity 3.2 - How can oxidation states be used to analyse redox reactions?
Valence electrons
Atoms on their own are not stable under normal circumstances (apart from the noble gases) and react with other atoms forming molecules or ions, depending on the difference in electronegativity between the reacting atoms.
They do this by using their outer shell electrons in such a way as to produce molecular shells with stable configurations of electrons, usually in shells of eight (an octet). Hence, the reactivity of an atom is determined mainly by its outer shell, or valence, electrons.
The vertical groups in the periodic table have the same number of valence electrons. It follows then, that vertical groups have similar chemical properties and, where dependent on the outer shell electrons, similiar physical properties.
Group trends
The outer electrons behaviour is modified by the attraction from the nucleus. As the nucleus increases in charge descending a group, it follows that there will be a gradual change in chemical characteristics, and any physical characteristics dependent on outer electrons, within a group as the group is descended.
These group properties are particularly noticeable in groups 1, 2 and 17, and for this reason they are usually the groups chosen to exemplify the concept of periodicity.
Oxidation state
The oxidation state is a number assigned to an atom to show the number of electrons transferred in forming a bond. It is the charge that atom would have if the compound were composed of ions.
If an atom accepts electrons, it is assigned a negative oxidation state. If it donates electrons, it is assigned a positive oxidation state. The term "oxidation number" is used interchangebly with oxidation state.
The redox rules
- The sum of all the oxidation states in a compound must add up to the charge on the species.
- The oxidation state of an uncombined element is always zero (0).
- The most electronegative atom takes a negative oxidation number.
- By convention, when naming compounds the oxidation number is written as a Roman numeral in the compound name immediately after the atom to which it refers.
Examples
- The oxidation state of iron in iron(II) sulfate is +2
- The oxidation state of phosphorus in phosphorus(V) oxide is +5
- The oxidation state of chromium in potassium dichromate(VI) is +6
Calculating the oxidation state
Oxidation state in compounds
There are some elements that nearly always have the same oxidation state. These can be used to calculate the oxidation states of the atoms to which they are bonded.
Hydrogen, for example has an oxidation state of -1 when bonded to a metal (more electropositive element) and +1 when bonded to a more electronegative element (non-metal). Oxygen is nearly always -2 (the exception is when it is in the form of the peroxide ion, (O-O)2-, where each oxygen atom has an oxidation state of -1).
Group 1 and 2 metals have an oxidation state of +1 and +2 respectively.
Example: Calculate the oxidation state of sulfur in sulfuric acid H2SO4
- Hydrogen = +1 oxidation state
- Oxygen = -2 oxidation state
(2 x H) + S + (4 x O) = 0
2 + S - 8 = 0
S = +6
Example: Determine the oxidation state of nitrogen in calcium nitrate Ca(NO3)2
- Calcium is in group 2 = +2 oxidation state
- Oxygen = -2 oxidation state
Therefore:
(+2) + [(2 x N) + (6 x -2)] = 0
+2 + 2N -12 = 0
2N = 10
N = +5
Oxidation state in ions
When you are dealing with an ion, the process for working out the oxidation state of the atoms within the ion is the same, with the important exception that the sum of the oxidation states must add up to the charge on the ion.
Example: Calculate the oxidation state of the chlorine atom in the chlorate ion, ClO3-
Each of the oxygen atoms has an oxidation state of -2
Therefore as Cl + (3 x -2) = -1
Then Cl = -1 + 6 = +5
Beware! IB examiners are very particular about differentiating between oxidation number/state and ionic charge. The oxidation number must be expressed as such, +2, -1 etc., whereas the ionic charge shows the number of charges in that order, 2+, 1- etc..