Oxidation numbers
An oxidation number is a number given to each atom or ion in a compound that shows us its
degree of oxidation.
Oxidation numbers can be positive, negative or zero. The + or − sign must always be included.
Higher positive oxidation numbers mean that an atom or ion is more oxidised. Higher negative
oxidation numbers mean that an atom or ion is more reduced.
Oxidation number rules
1. The oxidation number of any uncombined element is zero. For example, the oxidation
numbers of each atom in S8, Cl2 and Zn is zero.
2. In compounds many atoms or ions have fixed oxidation numbers:
• Group 1 elements are always +1
• Group 2 elements are always +2
• fluorine is always −1
• hydrogen is +1 (except in metal hydrides such as NaH, where it is −1)
• oxygen is −2 (except in peroxides, where it is −1, and in F2O, where it is +2).
3. The oxidation number of an element in a monatomic ion is always the same as the charge. For
example, Cl− is −1, Al3+ is +3.
4. The sum of the oxidation numbers in a compound is zero.
5. The sum of the oxidation numbers in an ion is equal to the charge on the ion.
6. In either a compound or an ion, the more electronegative element is given the negative
oxidation number.
Compounds of a metal with a non-metal
The metal always has the positive ox. no. and the non-metal has the negative ox. no. For
example, in sodium oxide, Na2O, Na = +1 and O = −2.
If we do not know the ox. no. of one of the atoms, we can often work it out using the invariable
ox. nos. in rule 2. For example in sodium sulfide:
ox. no. of each Na atom = +1
for two sodium atoms = +2
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Na2S has no overall charge, so the total ox. no. is zero (rule 4)
ox. no. of S = −2
Compounds of a non-metal with a non-metal
In compounds containing two different non-metals, the sign of the ox. no. depends on the
electronegativity of each atom. The most electronegative element is given the negative sign (rule
6).
1. Sulfur dioxide, SO2
ox. no. of each O atom = −2
for two oxygen atoms = 2 × (−2) = −4
SO2 has no charge, so the total ox. no. is zero (rule 4)
ox. no. of S = +4
2. Iodine trichloride, ICl3
chlorine is more electronegative than iodine, so chlorine is − and iodine is +
ox. no. of each Cl atom = −1
for three chlorine atoms = 3 × (−1) = −3
ICl3 has no charge, so the total ox. no. is zero (rule 4)
ox. no. of I = +3
3. Hydrazine, N2H4
nitrogen is more electronegative than hydrogen, so nitrogen is − and hydrogen is +
ox. no. of each H atom = +1 (rule 2)
for four hydrogen atoms = 4 × (+1) = +4
N2H4 has no charge, so the total ox. no. is zero (rule 4)
ox. no. of two N atoms = −4
ox. no. of each N atom = −2
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Question
State the ox. no. of the bold atoms in these compounds or ions:
1. P2O5
2. SO42−
3. H2S
4. Al2Cl6
5. NH3
6. ClO2−
7. CaCO3
Redox and oxidation number
We can define oxidation and reduction in terms of the oxidation number changes of particular
atoms during a reaction:
Oxidation is an increase in oxidation number.
Reduction is a decrease in oxidation number.
When tin reacts with nitric acid, the oxidation numbers of each atom of tin and nitrogen change
as shown below.
Each tin atom (Sn) has increased in ox. no. by +4: tin has been oxidised. Each nitrogen atom has
decreased in ox. no. by −1: nitrogen has been reduced. The ox. no. of each oxygen atom is
unchanged at −2. The ox. no. of each hydrogen atom is unchanged at +1. Oxygen and hydrogen
are neither oxidised nor reduced.
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In this reaction, nitric acid is acting as an oxidising agent because it increases the oxidation
number of another atom.
In this reaction, tin is acting as a reducing agent because it decreases the oxidation number of
another atom.
Oxidising agents and reducing agents
An oxidising agent (oxidant) is a substance which brings about oxidation by removing electrons
from another atom or ion.
An oxidising agent increases the oxidation number of another atom or ion.
When this happens, the oxidation number of the oxidising agent decreases.
Typical oxidising agents are oxygen, chlorine and potassium manganate (VII).
A reducing agent (reductant) is a substance that brings about reduction by donating (giving)
electrons to another atom or ion.
A reducing agent decreases the oxidation number of another atom or ion.
When this happens, the oxidation number of the reducing agent increases.
Typical reducing agents are hydrogen, potassium iodide and reactive metals such as aluminium.
Since oxidation and reduction occur together, in every redox reaction there must be an oxidising
agent and a reducing agent. For example, in the reaction given by the equation:
MnO4− + 5Fe2+ + 8H+ → Mn2+ + 5Fe3+ + 4H2O
MnO4− is the oxidising agent because the oxidation number of Mn decreases from +7 to +2
MnO4− has increased the oxidation state of iron from +2 to +3
Fe2+ is the reducing agent because the oxidation number of Fe increases from +2 to +3
Fe2+ has decreased the oxidation state of Mn in MnO4− from +7 to +2
Question
(a) Deduce the change in ox. no. for the bold atoms or ions in each of the following equations. In
each case, state whether oxidation or reduction has taken place.
(i) 2I− + Br2 → I2 + 2Br−
(ii) (NH4)2Cr2O7 → N2 + 4H2O + Cr2O3
(iii) As2O3 + 2I2 + 2H2O → As2O5 + 4H+ + 4I−
(b) Identify the oxidising agent and the reducing agent in parts a i and a iii.
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