Redox reactions involve the transfer of electrons from one reactant to another, resulting in oxidation and reduction. Oxidation is the loss of electrons and reduction is the gain of electrons. Common redox reactions include photosynthesis, respiration, combustion, and production and use of fertilizers.

1. Mr.Harangule Y. R.
M.Sc Organic chemistry Bed

2. What is it?
Redox Reactions are chemical reactions that involve
oxidation and reduction.
Oxidation can be defined as a loss of electrons to another
substance. Reduction can be defined as an acceptance of
electrons from another substance.
Redox reactions are those in which electrons are transferred
from one reactant to another.
Everyday redox reactions include:
 Photosynthesis
 Respiration
 Combustion of coal
 Production and use of fertilisers

3. Key Terms
If electrons are transferred, it is a redox
reaction.
1) A loss of electrons is called oxidation. A gain in
electrons is called reduction.
2) Reduction and oxidation happen simultaneously,
hence the name “redox”
3) An oxidising agent (oxidant) accepts electrons and
thus gets reduced
4) A reducing agent donates electrons and thus gets
oxidised
OIL RIG
Oxidation Is Loss of electrons Reduction Is Gain of
electrons

4. The oxidant is the species which causes oxidation and is itself reduced
The reductant is the species which causes reduction and is itself
oxidised

5. From this equation you can see that Na goes from an
oxidation state of 0 to +1, it has donated an electron
and has been oxidised. We can say that Na is the
reducing agent (or reductant) as it has reduced the Cl
Cl2 goes from 0 to -1, it has been reduced as it has gained
an electron. It can also be called the oxidant (or
oxidising agent) as it has oxidised the Na.

6. Oxidation/Reduction/Oxidant/Reduct
ant
Mg(s) + S (s) MgS (s)
Fe(s) + 2Ag+ 2Ag(s) + Fe2+
(aq)
Cl2 (aq) + 2Br-
(aq) Br2 (aq) + 2Cl-
(aq)
Fe3+
(aq) + 3OH-
(aq) Fe(OH)3 (s)

8. Oxidation Numbers - Rules
There are a lot of rules:
1) All atoms are treated as ions for this, even if they are
covalently bonded
2) Uncombined elements have an oxidation number of
0
3) Elements just bonded to identical atoms, eg O2 or H2
also have an oxidation number of 0.
4) The oxidation number of a simple monatomic ion, eg
Na+, is the same as its charge.

9. Oxidation Numbers - Rules
5) In compounds or compound ions, the overall
oxidation number is just the ion charge
SO4
2- - overall oxidation number is -2
Oxidation number of O = -2 (total -8)
Oxidation number of S = +6
***Within an ion, the most electronegative element has a
negative oxidation number equal to its ionic
charge***

10. Oxidation Numbers - Rules
6) The sum of the oxidation numbers in a neutral
compound is 0
Fe2O3 – overall oxidation number is 0.
oxidation number of O = -2 (total = -6)
so oxidation of Fe= +3
7) If you see roman numerals, this is an oxidation
number
Copper (II) Sulphate:
Copper has oxidation number of +2

11. Oxidation Numbers - Rules
8) The oxidation number of Hydrogen is +1 in its
compounds with non-metals (eg HCl)
The oxidation of Hydrogen is -1 in metal hybrides (eg
NaH)
9) The oxidation number of Oxygen is usually -2
Exceptions:
- peroxide compounds where O is -1 (eg H2O2)
- compounds where it is bonded to Fluorine where O
is +2

12. Assigning oxidation numbers to
the atoms in the following
substances
***Assign as many oxidation numbers as possible
and then find the oxidation number of the
unknown***
a) HBr
b) Na2O
c) CH4
d) Al2O3

13. Determine the oxidation
number of S in each of the
following:
a) H2SO4
b) SO3
c) SO2
d) H2S

15. Has a redox reaction taken
place??
Oxidation numbers are used to determine whether a
REDOX reaction has taken place
Oxidation is an INCREASE in the Oxidation Number of an
ATOM
Reduction is a DECREASE in the Oxidation Number of an
atom
***Keep in mind that oxidation cannot happen without
reduction***

16. Has a redox reaction taken
place??
1. Assign oxidation numbers to all species present
2. Determine whether a change in oxidation numbers has
occurred
3. Has oxidation and reduction both taken place?

17. Identify the following equations
as redox, state the substances that have been oxidised
and reduced
1) 2Fe(s) + 3Cl2 (g) 2FeCl3(s)
2) NH3(g) + HCl(g) NH4(s)
3) 2NO(g) + O2(g) 2NO2(g)
4) P4O10(s) + 6H2O(l) 4H3PO4(aq)

18. Half Equations
Half equations are a useful way of understanding the
processes involved in a redox reaction.
Although reduction and oxidation reactions occur
simultaneously, it is possible to consider the two reactions
separately.
To do this we separate the conjugate pair of oxidant and
reductant and we balance the equations by showing the
electrons.
Combining these half equations make up the ionic equation.

19. Half Equations
When an iron nail is placed in a blue copper sulfate solution,
the nail becomes coated with metallic copper and the blue
colour of the solution fades.
The full equation is:
Fe(s) + CuSO4(aq) FeSO4(aq) + Cu(s)
Fe(s) + Cu2+ + SO4
2-
(aq) Fe2+ + SO4
2-(aq) + Cu(s)
***No change in SO4
2- oxidation number so can be disregarded ad spectator ions
Fe(s) + Cu2+ Fe2+ + Cu(s)

20. Half Equations
Fe(s) + Cu2+ Fe2+ + Cu(s)
Oxidation Half-Equation:
Fe(s) Fe2+ + 2e-
Reduction Half-Equation:
Cu2+
(aq) + 2e- Cu(s)

21. Half Equations
The balanced ionic equation for the displacement of
silver from an aqueous silver nitrate solution by metallic
lead is:
2Ag+
(aq) + Pb(s) 2Ag(s)+ Pb2+
(aq)
a) Write balance oxidation and reduction half-equations
b) Which reactants accept electrons?
c) Which reactant is oxidised?
d) Which reactant is the reductant?

22. Chapter Review Questions
22, 23, 26, 27, 41