Redox Reaction Chapter 6

Syllabus Outline
6.4 Redox

1 Use a Roman numeral to indicate the oxidation number of an element in a compound

2 Define redox reactions as involving simultaneous reduction and oxidation

3 Define oxidation in terms of:

(a) gain of oxygen

(b) loss of electrons

(c) an increase in oxidation number

4 Define reduction in terms of:

(a) loss of oxygen

(b) gain of electrons

(c) a decrease in oxidation number

5 Identify redox reactions as reactions involving gain and loss of oxygen, or gain and loss of
electrons

6 Identify redox reactions by changes in oxidation number using:

(a) the oxidation number of elements in their uncombined state is zero

(b) the oxidation number of a monatomic ion is the same as the charge on the ion

(c) the sum of the oxidation numbers in a compound is zero

(d) the sum of the oxidation numbers in an ion is equal to the charge on the ion

7 Identify redox reactions by the colour changes involved when using acidified aqueous
potassium

manganate (VII) or aqueous potassium iodide

8 Define an oxidising agent as a substance that oxidises another substance and is itself
reduced

9 Define a reducing agent as a substance that reduces another substance and is itself oxidised

10 Identify oxidation, oxidising agents, reduction and reducing agents in redox reactions.

1
Sadia Shahzad Redox Reaction
Redox: Reduction and Oxidation 6.4
6.4.1 Definitions
Reactions which involve both oxidation and reduction are called
redox reactions. In these reaction Oxidation and reduction take
place simultaneously.
It’s a combination of two terms:
REDOX

Oxidation Reduction

Oxidation Reduction
1. Addition of Oxygen 1. Removal of Oxygen
C + O2 → CO2 CuO + H2 → Cu + 2H2O
CH4 + 2O2 → CO2 + 2H2O Pb + Ag2O → PbO + 2Ag

2. Removal of Hydrogen. 2. Addition of Hydrogen.

H2S(g) + Cl2(g) → 2HCl(g) + S(g) H2S(g) + Cl2(g) → 2HCl(g) + S(g)

3. Loss of Electrons 3. Gain of Electrons
2Na(s)→ 2Na+ + 2e- Cl2(g) + 2e+→ 2Cl
-

4. Increase in Oxidation state 4. Decrease in oxidation state

Oxidation O I L
Oxidation is loss of electrons.
Reduction R I G
Reduction is gain of electrons

2
Sadia Shahzad Redox Reaction
Oxidation state:
Ability of an atom to gain or lose the electrons.
Valency:
Ability of an atom to make bonds. For example,
C can make four bonds so its valency is 4.
N makes 3 bonds so its valency is 3.
Oxygen makes 2 bonds so its valency is 2.
Valency can never be positive or negative.
Number of electrons lost or gained during a chemical reaction is
determined by number of valence electrons.
 Atoms whose valence shell has less than 4 electrons will
always lose electrons. Elements of Group I, II and III will lose
1,2 and 3 electrons respectively.
Example: Ionic bond formation involves oxidation and
reduction
Examples: lets understand with Diagram of Na and O.

Over all charge on sodium atom is increased from 0 to +1.

Na → Na+ + 1e-
0 → +1

1. Charge is increased from 0 to +1.

3
Sadia Shahzad Redox Reaction
 2. Na is losing one electron.
3. Hence two definitions of oxidation are fulfilled.

O+ 2e- → O2-
1. Charge is reduced from o to 2-.
2. O has gained two electrons.
3. Hence two definitions of Reduction are fulfilled.
Note: whenever we identify redox, last two definitions are
considered because apparently; oxygen and hydrogen are
not gained or lost in all reactions.
Rules to assign oxidation state.
1. Group I elements (Li, Na, K, Rb, Cs, Fr) always have +1
charge in compounds.
2. Group II elements always have +2 charge in compounds.
Example CaO
3. Group III elements always have +3 charge in compounds.
Example Al2O3
4. In Group VII F always show -1 charge.
5. In most of the cases Oxygen has -2 charge. Except in per
oxide where oxygen has -1 charge.
Oxide O2- Peroxide O1-

4
Sadia Shahzad Redox Reaction
 6. Hydrogen is +1 (except metal hydrides. Such as NaH).
Examples H2SO4 HCl, HNO3
7. When an element exists in free state; its oxidation state is
zero. Examples Fe, Na, He.
8. If an element makes a bond with its own atom; its oxidation
state is zero. Examples Cl-Cl, H-H,
Charges to Remember

CO₃²⁻ NO₃⁻ SO₄²⁻ OH⁻

Examples:
Work out the oxidation states of the underlined elements in
these compounds using the rules:

+2 -2 +4 -2*2 +1 +5 -2*3 +4 -2*2 +6 -2*3

NO
NO2 HNO3 SO2 SO3
+4 -2*2 +1*2 +6*2 -2*7 +1 +7 -2*4 +7 -2*2
+1*2 +4 -2*3
SO2 H2SO4 K2Cr2O7 KMnO4 MnO4
+1*2 -1*2 +1 +5 -2*3 +2 -2
+3*2 -2*3

Cr2O3 H 2 O2 HClO3 PbO

5
Sadia Shahzad Redox Reaction
 Note: Transition metals and some common elements may have
different oxidation states in different compounds. Examples of
elements with variable oxidation states.

Some compounds with possible variable oxidation states have

roman numeral as a guide about their oxidation state, e.g.
- Iron(II) chloride has formula FeCl2 and iron oxidation state +2.
- Potassium(VI) dichromate has formula K2Cr2O7 and potassium
oxidation state +6.
- Manganese(IV) oxide has formula MnO2 and manganese
oxidation state +4.
Note: Romen numeral aways represent a positive oxidation state.
Example of Redox Reactions
Elements are involved in these reactions.
A reaction is said to be redox if oxidation states of elements in
reactants and products are different.
1. Reaction of metals and dilute acids.

0 +2 -2 +2 -2 0
Zn(s)+2HCl(aq)→ZnCl2(aq)+H2(g)

6
Sadia Shahzad Redox Reaction
 It is a redox reaction because oxidation states of Zn and H
are changing.

Zn 0 →+2 (Oxidation) H +1 → 0 (Reduction)

Zinc is reducing agent HCl is oxidizing agent
Oxidizing Agent
Substance which oxidizes the other substance and itself is
reduced.
Reducing Agent
Substance which reduces the other substance and itself is
oxidized.

2. Displacement reaction
0 +1 -1 0 +1 -1

Cl2(g) + 2KI(aq) → I2(s) + 2KCl(aq)

It is a redox reaction because oxidation states of Zn and H
are changing.
I -1→ 0 (Oxidation) Cl 0 → -1 (Reduction)
KI is reducing agent Cl is oxidizing agent

3. Carbon + Oxygen

0 0 +4 -2*2

7
Sadia Shahzad Redox Reaction
C (S) +O2 (g) →CO2(g)

C 0→ +4 (Oxidation) O 0 → -2 (Reduction)
C is reducing agent O2 is oxidizing agent

4. Reaction of zinc and copper sulphate

0 +2 -2 0 +2 -2

Zn(s) + 2CuSO4(aq) → Cu(s) + ZnSO4 (aq)

Zn 0→ +2 (Oxidation) Cu +2 → 0 (Reduction)
Zn is reducing agent CuSO4 is oxidizing agent

Some common oxidizing agents: HNO3, H2O2, CO2

Name of compound Formula Application
Potassium manganate KMnO4 Test for reducing agent;
(VII) A reducing agent converts
purple KMnO4 to
colourless Mn2+
Chlorine and other Cl2 Oxidizes Br- to Br2
non-metals. and I- to I2

Some common reducing agents:

Name of compound Formula Application
Potassium Iodide KI Test for oxidizing
agent;
8
Sadia Shahzad Redox Reaction
 A colourless I- ions
oxidizes to brow/red
I2.

Testing oxidizing and reducing agent:

Testing reducing agent Testing oxidizing agent

X Z Aq. KI
colourless
Aq. KMnO4
(purple)
If colour change from purple to
colourless X will be reducing agent. If colour change from colourless to
red brown Z will be oxidising agent.

More examples of reducing agents:

Carbon monoxide CO Reduces metal oxide to metal in
heat
Fe2O3+3CO → 2Fe +3 CO2
Hydrogen H2 Reduces copper (II) oxide to
copper.
CuO + H2 → Cu + H2O
Sulfur dioxide SO2 Used to bleach wood pulp and
preservatives.
Metals (highly reactive) Na, Mg, etc. Displace less reactive metals
9
Sadia Shahzad Redox Reaction
Ethanol C2H5OH Convert orange K2Cr2O7
Some other reducing agents: C, FeCl2, FeSO4
6.4.11 Reactions which are not Redox!
If there is no change in the oxidation number of element during
chemical reaction then it is not a redox.

1. Decomposition of carbonates by heat:

heat

CaCO3(s) ) → CaO(s)+CO2(g)
Oxidation state of each element doesn’t change. This is not a
redox reaction.
2. Neutralization:
HCl+NaOH→NaCl+H2O
Oxidation state of each element doesn’t change. This is not a
redox reaction.
3. Precipitation reaction:
AgNO3(aq)+NaCl(aq)→AgCl(s)+NaNO3(aq)
Oxidation state of each element doesn’t change. This is not a
redox reaction.

10
Sadia Shahzad Redox Reaction