Department of Environmental Sciences

Jahangirnagar University
Batch #52
Course # Env. 1117
Session: 2022–2023 (Year I Semester I Examination 2023)
Course Title # Foundations of Chemistry
File #4-redox reactions

1
 Oxidation-reduction
1. The Classical Concept
According to classical concept, oxidation simply means the addition of oxygen or
electronegative element to a compound, and the reduction is the removal of oxygen or
electronegative element from a compound.
Oxidation reactions:
2Mg + O2 = 2MgO
2CO + O2 = 2CO2
FeCl2 + Cl2 (+ heat) → FeCl3
Cu + S (+ Heat) → CuS
Reduction reactions:
2H2S (+ heat) → 2S + 2H2
2HCl → Cl2 + H2
2KI + H2O2 → 2KOH + I2
2N2 + 3H2 → 2NH3
2
2. Valence State Concept
Oxidation is a chemical reaction which involves the change of valence state in the positive
direction.
Reduction is a chemical reaction which involves the change of valence state in the negative
direction.
Oxidation reaction: Cu + S (+ Heat) → CuS
Reduction reaction: 2CuO (+ Heat) → 2Cu + O2

3
3. The Electronic Concept
Oxidation is a process which involves loss of electrons by an atom or ion.
Similarly, reduction can be attributed due to the gain of electrons by an atom.
An oxidizing agent is one that accepts electrons and is, thereby, reduced.
A reducing agent is one that gives electrons and is, thereby, oxidized.
Therefore, in oxidation-reduction reactions, electrons are transferred from the reducing
agent to the oxidizing agent. The oxidation-reduction reactions (may also known redox
reactions) are complementary and take place simultaneously.

4
In oxidation (the loss of electrons) and reduction (the gain of electrons) must always
accompany each other in a reaction. In other words, if a given reaction involves an oxidation,
it must also involve a reduction.
2Mgo + O2o = 2Mg2+O2– or 2MgO
Here, Mgo is oxidized to Mg2+ by losing of two electrons and oxygen O2o is reduced to
oxide O2– ion by gaining two electrons. In this reaction, Mg is the reducing agent and O2 is
acting as an oxidizing agent.

5
In the reaction between Zn and CuSO4 solution, metallic copper is deposited and some of
the zinc rapidly dissolves:
Zno + Cu2+SO42– = Cuo + Zn2+SO42–
Notice that SO42– ion does not undergo any change in this oxidation-reduction reaction. The
zinc metal is oxidized from Zno to Zn2+ ion by losing two electrons. The copper ion, Cu2+, is
reduced to Cuo by gaining of these two electrons. Here, Zno is oxidized and acts as the
reducing agent and Cu2+ is reduced and acts as the oxidizing agent.

6
 Oxidation Number and Oxidation State
When the electrons in the outermost energy level (or valence electrons) are removed from
an atom during a chemical reaction, the atom is said to be in a positive oxidation state.
Similarly, when an atom takes up electron during a chemical reaction, the atom is said to be in
a negative oxidation state.
The number of electrical charges on an atom is known as the oxidation number. Thus, in the
reaction
Mgo + Cl2o = Mg2+Cl–Cl– (or, MgCl2)
the oxidation number of Mg undergoes a change from zero in the free state to 2+ and the
oxidation number of chlorine changes from zero in the free state to –1. In general, the
oxidation-reduction reaction, the oxidation number of an element is increased and during
reduction the oxidation number is decreased.
7
8
9
10
In the oxidation-reduction reaction, during oxidation, the oxidation number of an element is
increased and during reduction the oxidation number is decreased.
A substance that causes an increase in the oxidation number of another substance is called
an oxidizing agent (or, an oxidant).
A substance that causes a decrease in the oxidation number of another substance is called a
reducing agent (or, a redundant).

11
 H2SO4
2(+1) + x + 4(–2) = 0
2 + x – 8=0
x = +6 (the oxidation number of S in H2SO4)
KNO3
+1 + x + 3(–2) = 0
+1 + x – 6 = 0
x = +5 (the oxidation number of N in KNO3)

12
Some elements may exist in several oxidation states and can have several oxidation
numbers.Thus, nitrogen can have oxidation states:
–3, –2, –1, 0, +1, +2, +3, +4, +5 in the following compounds:
–3 –2 –1 0 +1 +2 +3 +4 +5
NH3 N2H4 NH2OH N2 N2O NO N2O3 NO2 N2O5

Similarly, chlorine has oxidation numbers +7, +5, +3, +1, 0, and –1, as in

+7 +5 +3 +1 0 –1
HClO4 HClO3 HClO2 HClO Cl2 HCl

13
Obtain the oxidation number for the element
noted in each of the following:
a) Ga in Ga2O3
b) Mn in K2MnO4
c) Br in KBrO4
d) Cr in K2Cr2O7
e) N in NH2–
f) I in IO3–
g) Al in Al(OH)4–

14
 Rules for Assigning Oxidation Numbers
Rule Applies to Statement
1 Elements The oxidation number of an atom in an element is zero. [Na has an oxidation number
of zero]
2 Monatomic The oxidation number of an atom in a monatomic ion equals the charge on the ion.
ions [The oxidation number of Cu2+ is 2]
3 Oxygen The oxidation number of oxygen is –2 in most of its compounds. [An exception is O
in H2O2 and other peroxides, where the oxidation number is –1. The oxidation
number of O in KO2 is –½]
4 Hydrogen The oxidation number of hydrogen is +1 in most of the compounds. [The oxidation
number of hydrogen is –1 in binary compounds with a metal, such as CaH2]
5 Halogens The oxidation number of fluorine is –1 in all its compounds. Each of the other
halogens (Cl, Br, I) has an oxidation number of –1 in binary compounds, except when
the other element is another halogen above it in the periodic table or the other
element is oxygen.
6 Compounds The sum of the oxidation numbers of the atoms in a compound is zero.
and ions The sum of the oxidation numbers of the atoms in a polyatomic ion equals the charge
on the ion.
15
Now, look:
▪ What species is being oxidized? (or, what is the reducing agent)? What species is being reduced?
(or, what is the oxidizing agent)?
What are the oxidation numbers of nitrogen in each of the following compounds?
N2O, NO, N2O3, N2O4, HNO2, NH3, NH4Cl, N2H4, NH2OH, and AlN

16
Identification of Oxidizing and Reducing Agents
▪ If an element is in its higher possible oxidation state in a compound, it can function as an
oxidising agent, e.g., KMnO4, K2Cr2O7, HNO3, H2SO4, HClO4, etc.
▪ If an element is in its possible lower oxidation state in a compound, it can function as a
reducing agent, e.g, H2S, H2C2O4, FeSO4, SnCl2, etc.
▪ The compound will act as an oxidising agent if a highly electronegative element is in its
highest oxidation state.
▪ The compound acts as a reducing agent if a highly electronegative element is in its
lowest oxidation state.

17
18
Types of Redox Reactions
The different types of redox reactions are:
• Decomposition Reaction
• Combination Reaction
• Displacement Reaction
• Disproportionation Reactions
Decomposition Reaction
This kind of reaction involves the breakdown of a compound into different compounds.
Examples of these types of reactions are
• 2NaH → 2Na + H2
• 2H2O → 2H2 + O2
• Na2CO3 → Na2O + CO2
All the above reactions result in the breakdown of smaller chemical compounds in the form of
AB → A + B
But, there is a special case that confirms that all the decomposition reactions are not redox
reactions.
For example: CaCO3 → CaO + CO2
19
Combination Reaction
These reactions are the opposite of decomposition reactions and hence, involve the
combination of two compounds to form a single compound in the form of A + B → AB.
For example:
• H2 + Cl2 → 2HCl C + O2 → CO2
• 4Fe + 3O2 → 2Fe2O3
Displacement Reaction
In this kind of reaction, an atom or an ion in a compound is replaced by an atom or an ion of
another element. It can be represented in the form of X + YZ → XZ + Y.
For example: CuSO4 + Zn → Cu + ZnSO4
Disproportionation Reactions
Disproportionation reactions are known as reactions in which a single reactant is oxidized
and reduced.
For example: P4 + 3NaOH + 3H2O → 3NaH2PO2 + PH3

20
 Oxidation Number and Oxidation State
When the electrons in the outermost energy level (or valence electrons) are removed from
an atom during a chemical reaction, the atom is said to be in a positive oxidation state.
Similarly, when an atom takes up electron during a chemical reaction, the atom is said to be in
a negative oxidation state.
The number of electrical charges on an atom is known as the oxidation number. Thus, in the
reaction
Mgo + Cl2o = Mg2+Cl–Cl– (or, MgCl2)
the oxidation number of Mg undergoes a change from zero in the free state to 2+ and the
oxidation number of chlorine changes from zero in the free state to –1. In general, the
oxidation-reduction reaction, the oxidation number of an element is increased and during
reduction the oxidation number is decreased.
21
In the oxidation-reduction reaction, during oxidation, the oxidation number of an element is
increased and during reduction the oxidation number is decreased.
A substance that causes an increase in the oxidation number of another substance is called
an oxidizing agent (or, an oxidant).
A substance that causes a decrease in the oxidation number of another substance is called a
reducing agent (or, a redundant).

22
Iron(II) ion may be oxidized easily by a number of oxidizing agents to yield the iron(III) ion.
Permanganate ion in acidic, aqueous solution is a strong oxidizing agent and can oxidize Fe2+
to Fe3+. In the process of oxidizing the iron(II) ion, the permanganate ion (MnO4−) is
reduced to manganese(II) ion (Mn2+).
We can express these facts in the following skeleton equation:

We have written oxidation numbers over the appropriate atoms. Note that the skeleton
equation is not balanced. Nor is it complete; in an acidic, aqueous solution, H+(aq) and H2O(l)
may be possible reactants or products in the equation. (If this were a basic solution, OH− ion
would replace the H+ ion.)
23
Example. Zinc metal reacts with nitric acid, HNO3, to produce a number of products,
depending on how dilute the acid solution is. In a concentrated solution, zinc reduces nitrate
ion to ammonium ion; zinc is oxidized to zinc ion, Zn2+.
Write the net ionic equation for this reaction.
Note that nitric acid is a strong acid, so it exists in solution as H+ and NO3− ions.
For the skeleton equation, write just the NO3− ion.
The skeleton equation, with oxidation numbers for those atoms that
change values, is:

24
 Half reactions
The reaction of iron with copper(II):
A half-reaction is one of two parts of an oxidation-reduction reaction, one part of which
involves a loss of electrons (or increase of oxidation number) and the other a gain of
electrons (or decrease of oxidation number). The half reaction for the preceding equation
are:
Fe(s) → Fe2+ (aq) + 2e– (electrons lost by Fe)
Cu2+ (aq) + 2e– → Cu(s) (electrons gained by Cu2+)

25
Oxidation is the half-reaction in which there is a loss of electrons by a species (or an
increase of oxidation number of an atom).
Reduction is the half-reaction in which there is a gain of electrons by a species (or a
decrease in the oxidation number of an atom).
Thus,
the equation Fe(s) → Fe2+ (aq) + 2e– represents the oxidation half-reaction, and
the equation Cu2+ (aq) + 2e– → Cu(s) represents the reduction half-reaction.

26
Recall that a species that is oxidized loses electrons (or contains an atom that increases in
oxidation number) and a species that is reduced gains electrons (or contains an atom that
decreases in oxidation number).
An oxidizing agent is a species that oxidizes another species; it is itself reduced.
Similarly, a reducing agent is a species that reduces another species; it is itself oxidized. In the
above example, the copper(II) ion is the oxidizing agent, whereas iron metal is the reducing
agent.

27
Fe2+ is oxidized to Fe3+ by hydrogen peroxide when an acid is present. This reaction is
provided below.
2Fe2+ + H2O2 + 2H+ → 2Fe3+ + 2H2O
Oxidation half-reaction: Fe2+ → Fe3+ + e–
Reduction half-reaction: H2O2 + 2e– → 2 OH–
Thus, the hydroxide ion formed from the reduction of hydrogen peroxide combines with
the proton donated by the acidic medium to form water.

28
Application of Redox Reaction in Photosynthesis
Green plants convert water and carbon dioxide into carbohydrates, defined as
photosynthesis. The reaction is given as 6CO2 + 6H2O → C6H12O6 + 6O2

In the above reaction, we can see that carbon dioxide is reduced to carbohydrates, while the
water gets oxidized to oxygen; hence, it is a redox reaction. The energy is provided by the
sunlight for this reaction. This reaction is a source of food for animals and plants.
29
What are oxidation-reduction reactions?
Oxidation-reduction reactions are chemical reactions that involve the transfer of electrons
between the reacting species. These electron transfers are accompanied by a change in the
oxidation state of the reactants.
What are oxidizing agents?
An oxidizing agent is an electron-accepting species that is readily reduced in an oxidation-
reduction reaction. The oxidation numbers of these species tend to decrease in redox
reactions. Examples: nitric acid (HNO3) and hydrogen peroxide (H2O2).
What are reducing agents?
Reducing agents are the electron-donating species that readily undergo oxidation in
oxidation-reduction reactions. These species tend to lose electrons in redox reactions, and
their oxidation number increases. Examples: zinc (Zn) and lithium (Li).

30
Why is the reaction between hydrogen and fluorine a redox reaction?
The H2 molecule loses its electrons in this reaction, yielding two protons. The F2 molecule
accepts the electrons, leading to the formation of two fluoride ions. A polar covalent bond is
formed between the fluoride ion and the proton. Thus, this reaction is redox.
H2 + F2 → 2HF
Is every chemical reaction a redox reaction?
No, not every chemical reaction is a redox reaction. Reactions, like double decompositions,
acid-base neutralisation, and precipitation reactions are non-redox reactions.
• Na2SO4 + BaCl2 = 2NaCl + BaSO4
• 2KI + Pb(NO3)2 = 2KNO3 + PbI2
• 2HCl + CaCO3 = H2O + CaCl2 + CO2
• HNO3 + NaOH = H2O + NaNO3
Some decomposition reactions are also non-redox reactions.
• NH4Cl = HCl + NH3
• CaCO3 = CO2 + CaO
• 2Al(OH)3 = 3H2O + Al2O3
31
A cell potential is a measure of the driving force of the cell reaction. This reaction occurs in
the cell as separate half-reactions: an oxidation-half reaction and a reduction-half
reaction.The general forms of these half-reactions are:
Reduced species → oxidized species + ne– (oxidation/anode)
Oxidized species + ne– → reduced species (reduction/cathode)
The cell potential is the sum of the potentials for the reduction and oxidation half-
reactions.

32
33
Electrochemical Cells
If we place a strip of zinc metal in a beaker containing a solution of CuSO4, Zn is oxidized to
Zn2+ ions, while Cu2+ ions are reduced to metallic copper
Zn(s) + Cu2+(aq) ==> Zn2+(aq) + Cu(s)
The electrons are transferred directly from the reducing agent, Zn, to the oxidizing agent,
Cu2+ in solution. We have already seen, when we balanced redox equations, that we can
separate a redox reaction like the one above into two half-reactions. One of the half-
reactions being oxidation and the other being a reduction half-reaction.
Zn(s) ==> Zn2+(aq) + 2 e- oxidation half- reaction
Cu2+(aq) + 2 e- ==> Cu(s) reduction half- reaction

34
35
36
Example. A voltaic cell is constructed from a half-cell in which a cadmium rod dips into a
solution of cadmium nitrate, Cd(NO3)2, and another half-cell in which a silver rod dips into a
solution of silver nitrate, AgNO3. The two half-cells are connected by a salt bridge. Silver ion
is reduced during operation of the voltaic cell.
i) Draw a sketch of the cell.
ii) Label the anode and cathode, showing the corresponding half reactions at these
electrodes.
iii) Indicate the electron flow in the external circuit (with a lightbulb), the signs of the
electrodes, and the direction of cation migration in the half-cells.

37
38
39
40
The electrochemical process involved in the rusting of iron
Here a single drop of water containing ions forms a voltaic cell in which iron is oxidized to
iron(II) ion at the center of the drop (this is the anode). Oxygen gas from air is reduced to
hydroxide ion at the periphery of the drop (the cathode). Hydroxide ions and iron(II) ions
migrate together and react to form iron(II) hydroxide. This is oxidized to iron(III)
hydroxide by more O2 that dissolves at the surface of the drop. Iron(III) hydroxide
precipitates, and this settles to form rust on the surface of the iron.

41