Lesson on oxidation

1. OXIDATION – REDUCTION
REACTIONS
5C-I

3. What are
Redox
Reactions?
A redox reaction can be defined as
a chemical reaction in which
electrons are transferred between
two reactants participating in it. This
transfer of electrons can be identified
by observing the changes in the
oxidation states of the reacting
species

6. Electron-accepting species which tend to undergo a reduction in
redox reactions are called oxidizing agents. An electron-donating
species which tends to hand over electrons can be referred to as
a reducing agent. It can be noted that any redox reaction can be
broken down into two half-reactions, namely the oxidation half-
reaction and the reduction half-reaction.
When writing these half-reactions separately, each of them must
be balanced in a way that all the electrons are accounted for.

7. Types of Redox Reactions
Decomposition Reaction
Combination Reaction
Displacement Reaction
Disproportionation Reactions

9. Redox and
electron
transfer
Half-equations
•We can extend our definition of
redox to include reactions involving
ions.
Sodium reacts with chlorine to form
the ionic compound sodium
chloride.
We can divide this reaction into two
separate equations, one showing
oxidation and the other showing
reduction. We call these half-
equations.

12. Balancing
half-
equations
We can construct a balanced
ionic equation from two half
equations by balancing the
numbers of electrons lost and
gained and then adding the
two half-equations together.
The numbers of electrons lost
and gained in a redox
reaction must be equal.

15. When aluminium is heated in a
stream of chlorine gas
aluminium chloride is formed:

17. Oxidation number or oxidation state
An oxidation number can be assigned to each atom
or ion in a chemical substance. The oxidation
number indicates the number of electrons lost,
gained or shared as a result of chemical bonding.
Oxidation numbers are either positive, negative or
zero. Unless an oxidation number is zero, a plus or
minus sign is written in front of the number, and the
number 1 is always written, e.g. +1, -3.

18. Rules to follow when determining oxidation
numbers
1. The oxidation number of an element in the free state is zero.
2. A monoatomic ion will have an oxidation number that is equal to its charge.
3. In most compounds, the oxidation number of hydrogen is +1 (exception: when hydrogen is bonded to a
metal, the oxidation number -1.)
4. Oxygen, within a compound, will generally have an oxidation number of –2 (exception: in peroxides, the
oxidation number of oxygen is -1)
5. The oxidation numbers of combined elements from groups I and II are +1 and +2, respectively.
6. The algebraic sum of oxidation numbers of all elements in a neutral compound formula equals zero.
7. The algebraic sum of the oxidation numbers of all elements in a polyatomic ion is equal to the charge of
the ion

23. Examples
1. Determine the oxidation number of nitrogen in nitrogen dioxide
(NO2),
2. Determine the oxidation number of nitrogen in the nitrite ion (NO2
-)
3. Determine the oxidation number of chromium in the Cr2O7
-2 ion
and name the ion.
4. Determine the oxidation number of carbon in ethane (C2H6).

24. Oxidation-reduction reactions
in terms of oxidation number
In all redox reactions, the
oxidation number of one
element increases and the
oxidation number of
another element
decreases. Oxidation and
reduction can be defined
in terms of oxidation
number:

25. Using oxidation numbers to recognize redox
reactions
Any redox reaction can be recognized using the following steps:
• Write the balanced chemical equation for the reaction if it has not
been given.
• Write the oxidation number of each element below it in brackets. The
oxidation numbers of elements in polyatomic ions which remain
unchanged during a reaction need not be determined.
• Decide which element shows an increase in oxidation number. This
element has been oxidized.
• Decide which element shows a decrease in oxidation number. This
element has been reduced.
Note If the oxidation numbers of all elements remain unchanged, the
reaction is not a redox reaction.

27. Examples
Determine which reactant has been oxidized and
which has been reduced in the reaction between
iron(lIl) oxide and carbon monoxide.
Determine which reactant has been oxidized and
which has been reduced in the reaction between
magnesium and sulfuric acid.

28. Oxidizing and Reducing Agents
•The substance (atom, ion, and molecule) that gains
electrons and is thereby reduced to a low valency
state is called Oxidizing agent.
•The substance that loses electrons and is thereby
oxidized to a higher valency state is called a reducing
agent

29. Copper reacts with silver nitrate to
form silver and copper(II) nitrate.

30. oxidizing agent:
■ oxidizing agents increase the ox. no. of another atom
■an atom in the oxidizing agent decreases in ox. no.
■ the oxidizing agent is the substance which gets reduced – it gains
electrons
reducing agent:
■ reducing agents decrease the ox. no. of another atom
■ an atom in the reducing agent increases in ox. no.
■ the reducing agent is the substance that gets oxidized – it loses
electrons

34. Tests for the presence of an oxidizing agent
To test to see if a substance is an oxidizing agent, add it to a known
reducing agent, which gives a visible change when oxidized. The
reducing agents usually used are potassium iodide solution or an
aqueous solution of an iron(lI) salt.
1. To test to see if a substance is an oxidizing agent, add it to a known
reducing agent, which gives a visible change when oxidized. The
reducing agents usually used are potassium iodide solution or an
aqueous solution of an iron(lI) salt.
2. An oxidizing agent ca uses an aqueous solution of an iron(lI) salt,
e.g., iron(ll) sulfate, to change from pale green to yellow-brown
because it oxidizes the pale green Fe" ion to the yellow-brown Fe"
ion.

35. Tests for the presence of a reducing agent
To test to see if a substance is a reducing agent, add it to a
known oxidizing agent, which gives a visible change when it is
reduced. The oxidizing agents usually used are acidified
potassium manganate(VII) solution or acidified potassium
dichromate(VI) solution.
• A reducing agent causes acidified potassium manganate(VII)
solution to change from purple to colorless because it reduces
the purple MnO4
-; ion to the colorless Mn+2 ion.
• A reducing agent ca uses acidified potassium dichromate(VI)
solution to change from orange to green because it reduces
the orange Cr2O7
2- io n to the green Cr3+ ion.

37. REDOX REACTION IN OUR
EVERYDAY LIFE