1.1.4 Redox
Answers
Oxidation numbers
We use the concept of oxidation numbers to help us analyse redox chemistry at AS/A2. The oxidation number for a species is not the same as ionic charge; the numbers are assigned to species within covalent compounds, and ionic compounds. Oxidation numbers are given Roman Numerals (or regular numbers with the charge before the number, for example +2).
Here are some simple rules to follow when asked to assign oxidation numbers:
1. The oxidation number of an uncombined element is always zero;
2. The oxidation number of a simple ion is the same as the charge on the ion;
3. The oxidation numbers of all the atoms in a molecule or complex ion must add up to the total charge on that species;
4. For non metallic elements in an ion or molecule, we pretend that the most electronegative element ‘takes’ the electrons to give it the minus charge;
5. Hydrogen usually has an oxidation number of +1 in compounds (except in metal hydrides when it is –1);
6. Oxygen usually has an oxidation number of –2 in compounds (except in peroxides when it is –1);
7. Fluorine is -1 in a compound, but the other elements in group 7 have variable oxidation numbers.
Student activity
Work out the oxidation numbers of the elements shown in the given formulae:
Species / Oxidation number / Oxidation number / Oxidation numbera) Cl2 / Cl=0
b) NaCl / Na=+1 / Cl=-1
c) MgCl2 / Mg=+2 / Cl=-1
d) Mg / Mg=0
e) H2O / H=+1 / O=-2
f) H2O2 / H=+1 / O=-1
g) NaH / Na=+1 / H=-1
h) OH- / H=+1 / O=-2
i) MnO4- / Mn=+7 / O=-2
j) MnO2 / Mn=+4 / O=-2
k) CO2 / C=+4 / O=-2
l) CO32- / C=+4 / O=-2
m) SO42- / S=+6 / O=-2
n) HSO4- / S=+6 / O=-2 / H=+1
o) NH4+ / N=-3 / H=+1
p) NO3- / N=+5 / O=-2
q) SO32- / S=+6 / O=-2
Identifying oxidation and reduction within a reaction
There are many ways of thinking about reduction and oxidation, but remember they are opposites and they take place together. It is not possible to oxidise one species without reducing another.
For most inorganic chemistry at this level, we will use the following definitions:
oxidation = loss of electrons or increase in oxidation number
reduction = gain of electrons or decrease in oxidation number
A good way to remember it is the mnemonic OILRIG (oxidation is loss, reduction is gain). If a species loses electrons, then it will become less negative (or more positive), hence the oxidation number increases.
When analysing organic reactions it may still be useful to use the following definitions:
oxidation = loss of hydrogen or gain of oxygen
reduction = gain of hydrogen or loss of oxygen
Student activity
In the following equations, write the oxidation number of each species under the equations; note that for CuO you need a number underneath the Cu and another number underneath the O. Your teacher will show you how to do the first one.
Fe2O3 + 3C → 2Fe + 3CO
In the exam always write oxidation number on the appropriate line.
Never write the multiple
1. 2Mg + O2 → 2MgO
oxidation number 0 0 +2 -2
2. CuO + H2 → H2O + Cu
+2 -2 0 +1 -2 0
3. Zn + H2SO4 → ZnSO4 + H2
0 +1+6-2 +2+6-2 0
4. CO + ½ O2 → CO2
+2 -2 0 +4 -2
5. 2NO + 2CO → N2 + 2CO2
+2 -2 +2 -2 0 +4 -2
Student activity
a) For the same equations, draw arrows showing which change is the oxidation and which is the reduction. Your teacher will show you how to do the first one.
Fe2O3 + 3C → 2Fe + 3CO
1. 2Mg + O2 → 2MgO
Mg oxidised O2 reduced
2. CuO + H2 → H2O + Cu
H2 oxidised Cu reduced
3. Zn + H2SO4 → ZnSO4 + H2
Zn oxidised H reduced
4. CO + ½ O2 → CO2
C oxidised O2 reduced
5. 2NO + 2CO → N2 + 2CO2
C oxidised N reduced
b) In a redox reaction, there is a reducing agent (which becomes oxidised) and an oxidising agent (which becomes reduced). Sometimes we call these the reductant and oxidant. Identify the oxidising and reducing agents, by writing R under the reducing agent and O under the oxidising agent (in the examples above).
Student activity
Complete the following table.
What happens in terms of electron transfer?(does it lose or gain electrons?) / What happens to the oxidation number of this agent?
(does the oxidation number increase or decrease)
Oxidising agent / The oxidising agent will
gain electrons / The oxidation number will
decrease
Reducing agent / The reducing agent will
lose electrons / The oxidation number will
increase
Using Roman numerals to indicate the magnitude of the oxidation state within a species
Some compounds have the same name but contain species with different oxidation states. For example:
sulphuric (IV) acid H2SO3
sulphuric (VI) acid H2SO4
Here the sulphur has an oxidation state of +4 (IV) in H2SO3 and +6 (VI) in H2SO4.
Student activity
Name the following species.
a) MnO2
Manganese (iv) oxide
b) MnO4-
Manganate (vii)
c) ClO3-
Chlorate (v)
d) SO32-
Sulphate (iv)
e) HSO4-
Hydrogen sulphate (vi)
How do metal and non-metal elements behave when taking part in reactions?
Look at the equations below.
a) Zn + H2SO4 → ZnSO4 + H2
b) 2Fe + 3Cl2 → 2FeCl3
c) Mg + 2HCl → MgCl2 + H2
d) 3Mg + N2 → Mg3N2
e) Mg + I2 → MgI2
1. a) What has happened to the oxidation numbers of the metal elements? they always increase
b) Have they lost or gained electrons? lost electrons
2. a) What has happened to the oxidation numbers of the non-metal elements? they always decrease
b) Have they lost or gained electrons? gained
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