OXIDATION AND REDUCTION

IMPORTANT DEFINATIONS

Many chemical reactions involve the addition of oxygen or hydrogen to the reactants. The

reaction in which oxygen is added is called oxidation whereas the reaction which involves

the addition hydrogen is called reduction. The two reactions always occur together, these

reactions are called Redox reactions.

1.1 ELECTRONIC INTERPRETATION OF OXIDATION AND REDUCTION:

Oxidation and reduction can be predicted in terms of the transference of electrons in the

reactants. This is called as electronic concept of oxidation and reduction. According to this

concept oxidation and Reduction are defined as follows.

a) Oxidation is a process in which an atom or ion loses one or more electrons. It is also

called de-electronation.

Example: Na ------------ Na+ + e-

S2- ------------ S + 2e-

During oxidation there is an increase in positive valency or decrease in negative valency.

The substances which can lose electrons during redox reactions are called reducing agents.

b) Reduction is a process in which an atom or ion gains one or more electrons. It is also

called electronation.

Example: Cl + e- ---------------- Cl-

Sn4+ + 2e- ---------------- Sn2+

During oxidation there is an increase in negative valency or decrease in positive valency. The

substance which can gain the electrons during redox reactions are called oxidizing agents.

Consider the following reaction occurring in aqueous medium.

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

IMPORTANT FACTS:

In water medium it occurs as follows

Zn(8) + 2H+ (aq) + 2C1-

(aq) ------------- Zn2+ (aq) + 2C1-(aq) + H2(g)

In this reaction Zn metal gets oxidized and hydrogen gets reduced.

Lose Two electrons

Zn(s) + 2H+(aq) Zn2+

(aq) + H2(g)

Gain Two Electrons

The two electrons from Zn(8) is transferred 2H(8) is transferred to 2H+ ions, hence it is a redox

reaction. (Oxidation is also addition of electronegative element and reduction is addition of

electropositive element.

1. The Oxidation State or oxidation number is the Charge an atom of an element would

have if it existed as an ion in a compound .

IN TERM OF

OXIDATION

REDUCTION

ELECTRONS

The Loss of electrons

The Gain of electrons

HYDROGEN

The LOSS of hydrogen

The GAIN of hydrogen

OXYGEN

The Gain of oxygen

The Loss of oxygen

OXIDATION STATE

The INCREASE in the oxidation

The DECREASE in the oxidation state of

3. An Oxidizing Agent is a substance that cause oxidation on another substance.

4. A Reducing Agent is a substance that cause reduction on another substance.

5. A Redox Reaction is a chemical reaction in which oxidation and reduction both takes place.

TEST YOURSELF:

Define oxidation and reduction ( redox ) in terms of gain and loss of oxygen hydrogen, electron transfer and charges in oxidation state

OXIDATION AND REDUCTION RECTIONS:

1. A substance is Oxidised if it there is a Gain of Oxygens in a chemical reaction.

EXAMPLE1:

Oxidation

( Gain of Oxygen )

4Na (s) + O2 (g) 2 Na2O (s)

Sodium oxygen sodium oxide

2. A substance is Oxidised if it there is Loss of Hydrogens in a reaction .

state of the element the element

1.1

Sodium Gains

Oxygen and is

Oxidised to

sodium oxide .

EXAMPLE 2:

Oxidation

( Loss of Hydrogen )

H2S (g) + CI2 (g) 2 HCI (g) + S(s)

Hydrogen sulfide chlorine hydrogen chloride sulfur

3. A substance is Oxidised if there is loss of Electrons in a chemical reaction .

EXAMPLE 3:

Oxidation

( Loss of Electrons )

Mg (s) + CI2 (g) MgCI2(s)

Magnesium chlorine Magnesium chloride

Electrons have been transferred during this reaction, shown by the following half equations .

Mg (s) CI2 (g)

4. A substance is Reduced if there is Loss of Oxygens in a chemical reaction .

H2S Looses

Hydrogen and is

Oxidised to sulfur

EXAMPLE4 :

Reduction

( Loss of Oxygen )

CuO (s) + H2 (g) Cu(s) + H2O(g)

Copper (II) oxide hydrogen copper water

5. A substance is Reduced if it Gains Hydrogen in a reaction .

EXAMPLE5 :

Reduction

( Gain of Hydrogen )

H2(g) + CI2 (g) 2HCI(g)

Hydrogen chlorine hydrogen chloride

6. A substance is Reduced if there is Gain of Electrons in a chemical reaction .

The copper (II)

oxide loses oxygen

and is reduced to

copper metal

Chlorine Gains

Hydrogen and is

reduced to

hydrogen chloride

EXAMPLE6 :

Reduction

( Gain of Electrons )

2FeCI3(aq) + H2S (g) 2FeCI2(aq) + 2HCI (aq) + S (s)

Iron (III) hydrogen Iron (II) hydrogen sulfur

Chloride sulfide chloride chloride

The ionic equation for the reaction:

2Fe3+ (aq) + H2S (g) 2Fe2+ (aq) + 2H+ (aq) + S (s)

EXAMPLE7 :

Which compound is (a) oxidized and (b) reduced in the following reaction ?

2NH3 (g) + 3CuO( s) N2 (g) + 3CuO (s) + 3H2O(I)

SOLUTION:

> CI- ions spectator ions.

> Iron (III) ion gains an

electron and is

reduced to iron (II) ion.

Oxidation

( Loss of Hydrogen )

2NH3 (g) + 3CuO( s) N2 (g) + 3CuO (s) + 3H2O(I)

Reduction

( Loss of Oxygen )

7. OXIDATION STATE:

The Oxidation State (oxidation number) is the charge an atom of an element would have

if it existed as an ion in a compound .

Oxidation state is also called Oxidation Number.

The oxidation state can be a positive number (e.g.+3), a negative number (e.g.-3) or

zero.

The oxidation state of an element can be determined by using the following rules.

RULE

EXAMPLE

OXIDATION

The oxidation state of an element in the

uncombined state (free state) is zero.

Na C

Mg O2

0 0 0 0

The oxidant state of a simple ion is equal

to the charge on the ion.

Na+ Fe3+ O2- P3-

+1 +3 -2 -3

The oxidation states of Groups I and II

elements in their compounds are fixed (refer to Periodic Table)

Group I elements Group II elements

+1 +2

The oxidation states of hydrogen and

oxygen in their compounds are fixed

H ( in H2O ) O ( iN CuO )

+1 -2

The oxidation states of the atoms present

in the formula of the compound add up to zero.

AI2O3

2AI = 2 X (+3) = +6 3O = 3 X (2-) = -6

Total = (+6 ) + (-6) = 0

The oxidation states of all the atoms in a

polyatomic ion is equal to the charge on the ion.

SO4

2-

S = + 6

40 = 4 X (-2) = - 8 Total = (+6) + (-8) = -2

COMMON ERROR ACTUAL FACTS

The reaction as shown below is

a neutralization reaction :

Acid + Base Salt and Water

4HI(aq) + MnO2 (s) MnCI2 (aq) + H2O (l)

+ Cl2 (g)

The reaction is not a neutralization

reaction as chlorine is also produced

other than salt and water .

( Comment : It is a redox reaction . MnO2 oxidises

HCI to CI2 and itself reduced to MnCI2 . )

1.3 OXIDATION NUMBER:

Oxidation number is defined as the charge present on the atom of an element which is present in

the combined state. It may have a positive or negative value. The oxidation number may be real

or apparent charge.

Difference between valency and oxidation number:

Valency Oxidation number

1. It is the combining capacity of an element. 1. It is a charge assigned to an atom

or ion ion in a molecule by using

arbitrary rules.

2. It has no negative or positive sign 2. It carries a sign – ve or + ve.

3. Valency is a fixed value 3. It is not fixed for an element. It

Depends on the compound.

4. It is a whole number but never zero 4. Is is a whole number, fraction or

Even zero.

Table 4.1 : Differences between Valency and Oxidation Number

1.4 RULES FOR COMPUTING OXIDATION NUMBER:

The following arbitrary rules have been adopted, to calculate oxidation number or

Element on the basis of periodic properties of elements.

1. In an uncombined state or free state the ON in zero.

2. In the combined state of elements the ON’s are

a) F = -1

b) O = -2 (peroxides – O – O - =1) In F2O it is +2

c) H = +1

d) Metals always + Ve

e) Alkali and Alkaline earth metals = +1 and + 2.

f) Halogen = -1

g) Sulphur = - 2

3. The algebraic sum of all oxidation numbers of elements in a compound is zero

Example in K2MnO4

4. The algebraic sum of all oxidation numbers of elements in a radical is equal to net

charge on that radical

Example C2O4 2-

5. Maximum oxidation number (except O and F) = Group number.

Minimum oxidation number (except metals) = group number – 8.

(Note group number is an Mendeleev’s modern periodic table).

1.4 CALCULATION OF OXIDATION NUMBERS:

Examples:

1. KMnO4 Let on O.N. of Mn be ‘x’

1 + x + 4 (-2) = 0

1 + x – 8 = 0

1 + x =8

∴ x = 8-1

∴ x = +7

∴ The ON of Mn = + 7

2. H2SO4 Let the O.N. of s be ‘x’

2 + x + 4 (-2) = 0

2 + x -8 = 0

∴ x = +6

3 K2Cr2O7 Let the O.N. of Cr be ‘x’

2+2x+7(-2)=0

2 + 2x – 14 = 0

∴ x = +6

3. H3POa Let the O.N. of P be ‘x’

3 + x + 4 (-2) = 0

3 + x + (-8) = 0

∴ x = -5

4. H3PO4 Let the O.N. of P be ‘x’

3 + x + 4(-2) =0

3 +x+(-8) =0

∴ x= -5

5. HNO3 Let the O.N. of N be ‘x’

1 + x-6=0

X=6-1

∴ X = +5

6. (Fe(CN)6)3- Let the O.N. of Fc be ‘x’

X + 6(-1) = -3

x-6 = -3

x = -3+6

∴ x = +3

7. MnO4- Let the O.N. of Mn be ‘x’

X + 4 (-2) = 0

X +(-8) = 0

X – 8 = 0

∴ x = +8

8. Cr(H2O)3+ Let the O.N. of Mn be ‘x’

The ON of H2O is zero as it is neutral molecule

X + 6 (0) = +3

∴ x = +3

9. Na2S4O6 Let the O.N. of S be ‘x’

2 + 4x +6 (-2) = 0

2+4x-12=0

2+4x=12

4x=12-2

4x=10

10 5

X = --- = --

4 2

∴ x = + 5

----

2

10. FeSO4 (NH4) SO4 6H2O

Let O.N. of Fe be ‘x’

X +(-2) +0+0=0

∴ x = +2

1.5 BALANCING EQUATIONS:

The chemical equations involving redox reactions can be balanced by following two

Methods.

a) Oxidation number method

b) Ion electron Method (Half reaction method)

We shall learn the balancing of equations involving redox reactions by oxidation

number method.

1.6a) Oxidation number method:

The following steps must be applied in balancing a redox equation by oxidation

number method.

1. Write the skeleton equation.

2. Write the oxidation numbers of all elements on their symbols.

3. Identify the elements that undergo change in oxidation number.

4. Calculate the increase or decrease in ON per gram atom, with respect to the

reactants. If more than one atom is present, then mulply the number of atoms

undergoing change to calculate the total change in oxidation number.

5. Equate the increase and decrease in ON on the reactant side by multiplying the

oxidizing and reducing agents suitably.

6. Balance the equations except hydrogen and oxygen. Later balance hydrogen and

oxygen also.

7. In reactions occurring in acid medium, balance Oxygen by adding H2O where

Oxygen is less. Then balance Hydrogen atoms by adding H+ where Hydrogen is

less.

8. In Basic medium, balance negative charges by adding OH where negative charge

is less. Then and H2O molecules to other side and balance the equation.

EXAMPLE8:

Calculate the oxidation states of nitrogen in the following compounds:

a) NH3 b) N2O4 c) NO3-

SOULATION:

a) Let the oxidation state of N = x . x 3 X (+ 1)

x + 3 X (+1) = 0 (Rule 5 )

x = - 3 NH3

The oxidation state of nitrogen in NH3 is – 3 .

b) Let the oxidation state of N = y. 2y 4 X (-2)

2y + 4 X (-2) = 0 (Rule 5 )

Y = + 4 N2O4

The oxidation state of nitrogen in N2O4 is +4 .

c) Let oxidation state of N = z . z 3 X (-2)

z + 3 X (-2) = -1 ( Rule 6 )

z = + 5 NO3-

The oxidation state of nitrogen in NO3- is + 5,

Oxidation state is not written in the same way as the charge on an ion .

For example, in PbCl2, oxidation state of Pb is +2, but the charge on Pb2+ is 2+.

Similarly, the oxidation state of phosphorus in Na3P is -3, but the charge on P3- is 3- .

8. To check whether oxidation has taken place in any reaction, follow the three steps :

a) Write the balanced equation for the reaction .

b) Write the oxidation states of all the substances in the reaction .

c) Compare the oxidation states to check which reactant has been oxidised .

9. Oxidation occurs when the oxidation state of an element increases .

Reduction occurs when the oxidation state of an element decreases .

EXAMPLE9 :

TIPS FOR STUDENT

Oxidation

( Increase in Oxidation State )

+2 +3

2FeCl2(s) + Cl(g) 2FeCl3(s)

0 -1

Reduction

( Decrease in Oxidation State )

In this reaction, the oxidation number of

Iron increases from +2 in FeCl2 to +3 in FeCl3. This is an oxidation process .

Chlorine decreases from O in Cl2 to -1 in Cl- . This is a reduction process .

10. Oxidation and reduction always occur together. If one reactant is oxidised, the other

reactant must be reduced. We call the combined process the redox reaction .

TIP FOR STUDENT

OXIDISING AND REDUCING AGENTS:

1. A substance that oxidizes other substances is called an Oxidizing Agent .

An oxidising agent is reduced when it oxidises another substance.

2. A substance that reduces other substances is called an reducing Agent .

An reducing agent is oxidised when it reduces another substance.

EXAMPLE10 :

COMMON ERROR ACTUAL FACTS

The oxidation number of hydrogen is 1

and the oxidation number of Oxygen is 2.

The “ + ” and “ – ”signs must be

shown . Thus, the oxidation number of

hydrogen +1 the oxidation number of

oxygen is -2 .

10.2

In the extraction of iron from iron ( III ) oxide, the following reaction occurs between iron ( III )

oxide and carbon monoxide .

Reduced

( Loss of Oxygen )

+3 0

Fe2O2 (s) + 3CO(g) 2Fe( l ) + 3CO2 (g)

oxidising agent +2 +4

reducing agent

Oxidised

( Gain Of Oxygen )

a) Iron ( III ) oxide is an oxidizing agent . It oxidises carbon monoxide to carbon dioxide and is

itself reduced to iron.

b) Carbon monoxide is a reducing agent . It reduces iron (III) oxide to iron and is itself oxidised

to carbon dioxide.

3. In terms of electron transfer ,

a) An oxidising agent is an acceptor of electrons,

b) A reducing agent is a donor of electrons.

EXAMPLE 11:

Reaction between chlorine gas and potassium iodide solution.

Cl2 (g) + 2KI (aq) 2KCl (aq) + I2 (aq)

Chlorine Potassium iodide Potassium chloride iodine

Chlorine gas is the acceptor of electron, and is thus an oxidising agent .

Cl2 (g) + 2e- 2CI- (aq) … reducation

The iodide ion is the donor of electrons, and is thus a reducing agent .

2I- (aq) I 2 (aq) + 2e- … oxidation

EXAMPLE 12:

Identify the oxidising and reducing agents in the following reaction .

a) PbO(s) + H2 (g) Pb(s) + H2 O (l)

b) Zn(s) + CuSO4 ( aq) ZnSO4 (aq) + Cu (s)

SOLUTION:

An reducing agent

undergoes

oxidation

An oxidising agent

undergoes

reduction.

a) Reduction

( loss of oxygen )

PbO(s) + H2(g) Pb (s) + H2O(l)

Oxidation

( gain of oxygen)

Lead (ll) oxide, PbO, is the oxidising agent . It oxidizes hydrogen to water.

Hydrogen is the reducing agent . It reduces leas (ll) oxide to lead.

b) Oxidation

( donor of electrons )

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

Reduction

( acceptor of electrons )

In this reaction,

Zn(s) Zn2+(aq) + 2e-

( electron acceptor )

Cu2+ (aq) + 2e- Cu(s)

( electron acceptor )

Zinc metal is the reducing agent . It reduces Cu2+ to Cu by donating electrons to Cu2+ .

Cu2+ ion is the oxidising agent. It oxidizes Zn to Zn2+ by accepting electrons from Zn .

4. Some substances can act as both an oxidising agent and a reducing agent.

EXAMPLE:

Sulfur dioxide, SO2’ acts as an oxidising agent in reaction (1) but as a reducing agent in reaction

(2).

-2 0

SO2 (g) + 2H2S (g) 3S (s) + 2H2O (I) …… ( 1 )

oxidising agent

0 -2

2SO2 (g) + O2 (g) 2SO3 (g) …… ( 2 )

Reducing agent

5. The table below shows some common oxidising and reducing agents.

OXIDISING AGENT

REDUCING AGENT

Chlorine, CI2

Potassium iodide KI

Bromine, Br2

Carbon C

Nitric acid HNO3

Carbon monoxide, CO

Hydrogen Peroxide, H2O2

Ammonia, NH3

Potassium manganate (VII), KMnO4

Sulfur dioxide, SO2

Potassium dichromate(VI), K2 Cr2 O7

Hydrogen, H2

Oxygen, O2

Hydrogen sulfide, H2S

Concentrated sulfuric acid, H2SO4

Metals

6.

TEST

DIAGRAM

OBSERVATION AND

INFERENCE

EXPLANATION

Add potassium iodide solution (colourless).

Potassium Iodide

Unknown Brown

If the solution turns brown, The substance is an oxidising agent.

An oxidising agent oxidises iodide ion to iodine. 2I-(aq) I2(aq) +2e-

(Colourless) (Brown)

Substance Solution

Test with starch-iodide paper.

Moist starch-iodide paper

An oxidising agent changes the colour of moist starch-iodide paper from white to blue,

The iodine reacts with the starch to give a blue colour.

7.

TEST

DIAGRAM

OBSERVATION AND

INFERENCE

EXPLANATION

Add acidified potassium dichromate (VI) solution. Dilute sulfuric acid is always used to acidify Potassium dichromate (VI) solution.

Acidified potassium dichromate (VI)solution

Orange Green solution solution

If the orange solution turns green, The substance is a reducing agent.

The reducing agent reduces the dichromate (VI) ion, Cr2 O7

2-, to chromium(III) ion, Cr3+. Cr2 O7

2- (aq) + 14H+ (aq) + 6e- (orange) 2Cr3+(aq) + H2 O(I) (green)

Add acidified potassium managnate (VII) solution.

Acidified potassium Managnate (VI)solution

Purple Colourless solution solution

If the purple solution turns colourless (decolourisation), the substance is a reducing agent.

The reducing agent reduces the managnate (VII) ion to managanese (II) ion. MnO4

- (aq) + 8H+ (aq) + 5e- (Purple) Mn2+(aq) + H2 O(I) (colourless)

SUMMARY AND KEY POINTS

1.)The reaction in which oxygen is added is called oxidation whereas the reaction which

involves the addition hydrogen is called reduction. The two reactions always occur

together, these reactions are called Redox reactions.

2.) Oxidation is a process in which an atom or ion loses one or more electrons. It is also

called de-electronation.

COMMON ERROR ACTUAL FACTS

A substance is always an oxidising

agent or a reducing agent in all

reactions.

Some substance, such as sulfur dioxide, SO2,

hydrogen peroxide, H2O2, and sodium

nitrite, NaNO2, can act as both oxidising and

reducting agents.

Example: Na ------------ Na+ + e-

S2- ------------ S + 2e-

During oxidation there is an increase in positive valency or decrease in negative valency.

The substances which can lose electrons during redox reactions are called reducing agents.

3.) Reduction is a process in which an atom or ion gains one or more electrons. It is also

called electronation.

Example: Cl + e- ---------------- Cl-

Sn4+ + 2e- ---------------- Sn2+

During oxidation there is an increase in negative valency or decrease in positive valency. The

substance which can gain the electrons during redox reactions are called oxidizing agents.

4.) The Oxidation State or oxidation number is the Charge an atom of an element would

have if it existed as an ion in a compound .

IN TERM OF

OXIDATION

REDUCTION

ELECTRONS

The Loss of electrons

The Gain of electrons

HYDROGEN

The LOSS of hydrogen

The GAIN of hydrogen

OXYGEN

The Gain of oxygen

The Loss of oxygen

OXIDATION STATE

The INCREASE in the oxidation state of the element

The DECREASE in the oxidation state of the element

5.) An Oxidizing Agent is a substance that cause oxidation on another substance.

6.) A Reducing Agent is a substance that cause reduction on another substance.

7.) A Redox Reaction is a chemical reaction in which oxidation and reduction both takes place.

8.) A substance is Oxidised if it there is a Gain of Oxygens in a chemical reaction.

9.) A substance is Oxidised if it there is Loss of Hydrogens in a reaction.

10.) A substance is Oxidised if there is loss of Electrons in a chemical reaction.

11.) A substance is Reduced if there is Loss of Oxygens in a chemical reaction.

12.) A substance is Reduced if it Gains Hydrogen in a reaction.

13.) A substance is Reduced if there is Gain of Electrons in a chemical reaction.

14.) OXIDATION STATE:

The Oxidation State (oxidation number) is the charge an atom of an element would have

if it existed as an ion in a compound.

Oxidation state is also called Oxidation Number.

The oxidation state can be a positive number (e.g.+3), a negative number (e.g.-3) or

zero.

The oxidation state of an element can be determined by using the following rules.

15.) Oxidation number is defined as the charge present on the atom of an element which is

present in the combined state. It may have a positive or negative value. The oxidation number

may be real or apparent charge.

16.) The algebraic sum of all oxidation numbers of elements in a compound is zero.

Example in K2MnO4

17.) The algebraic sum of all oxidation numbers of elements in a radical is equal to net

charge on that radical.

18.) Maximum oxidation number (except O and F) = Group number.

Minimum oxidation number (except metals) = group number – 8.

(Note group number is an Mendeleev’s modern periodic table).

KEY POINTS:

i)Oxidation state is not written in the same way as the charge on an ion .

ii)For example, in PbCl2, oxidation state of Pb is +2, but the charge on Pb2+ is 2+.

Similarly, the oxidation state of phosphorus in Na3P is -3, but the charge on P3- is 3- .

19.) Oxidation occurs when the oxidation state of an element increases .

Reduction occurs when the oxidation state of an element decreases .

20.) Oxidation and reduction always occur together. If one reactant is oxidised, the other

reactant must be reduced. We call the combined process the redox reaction.

21.) A substance that oxidizes other substances is called an Oxidizing Agent.

An oxidising agent is reduced when it oxidises another substance.

22.) A substance that reduces other substances is called an reducing Agent .

An reducing agent is oxidised when it reduces another substance.