2. Introduction
Oxidation-reduction reactions are reactions in which electrons are transferred
from one reactant to another reactant.
Oxidation is defined as the loss of electrons
Reduction as the gain of electrons.
They must occur simultaneously, when a substance gives up electrons,
There must be another substance to receive them
The first substance is oxidized and the other is reduced
Substances which lose electrons are reducing agents or reductants
Those which gain electrons are oxidising agents or oxidants.
LEO the
lion says
GER
3. Electrons carry negative charge
So more numbers of electrons more negative charge
Example:
Na + cl = Nacl
(no charge) (no charge) (Na⁺ cl⁻)
Na donates electron Cl Receives electrons
Sodium is reducing and cl is oxidising
As Na is losing electrons the negative charge decreases and
converts to positive +
Cl is receiving electrons so more negative charge as more
electrons are adding.
How?
4. No reaction can happen in half
If there is reduction there should be oxidation.
Half reactions:
Na + cl = Nacl
How to write half reactions…let’s see..
For chloride
Cl + e⁻ = cl⁻
For sodium can we write like this. Let’s see..
Na - e⁻ = Na⁺
Something is not
right here…can
you guess?
5. Exactly…….
Is there any “-” symbol in a reaction phase…
No…..
So how to write the half reaction then?
Ok here we go…..
Na Na⁺ + e⁻
Here the reaction conveys the same information as it is donating electrons.
Na Na⁺ + e⁻
cl + e⁻ cl⁻
Na + cl Na⁺cl⁻ (Nacl)
7. Oxidation Number
The oxidation number is basically the count of electrons that atoms in a molecule
can share
Rules:
Elements by itself = 0 eg: Ag
Group 1A = always +1
Group 2A = always +2
Halogens = usually -1, positive with oxygen
Monoatomic ion = ion charge is the oxidation number (cl⁻, cu²⁺)
H = +1 with nonmetals (oxygen, carbon)
-1 with metals (cu, iron)
o = usually -2
-1 in peroxide
F = always -1
sum of oxidation number for a neutral compound = 0
sum of oxidation number for a polyatomic ion = ion charge
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17. Oxidising agent
A substance that tends to bring about oxidation by being reduced and
gaining electrons.
Na + cl = Nacl
Chlorine is reducing by gaining electrons but it is an oxidising agent
Cl + e⁻ = cl⁻
Reducing agent
A substance that tends to bring about reduction by being oxidized and
losing electrons.
Sodium is oxidising by loosing electrons but is a reducing agent
Na Na⁺ + e⁻
18. Equivalent weight
Equivalent weight of a substance (oxidant or reductant) is equal to molecular
weight divided by number of electrons lost or gained by one molecule of the
substance.
It is not a constant quantity but depends up on the reaction it is taking place.
Equivalent weight of oxidising agents = Molecular weight
No.of electrons gained by one molecule
Equivalent weight of reducing agents = Molecular weight
No. of electrons lost by one molecule.
19. Example :
KMno₄ K⁺ + Mno₄⁻
Basic medium
Mno₄⁻ + e⁻ Mno₄⁻ ²
Oxygen oxidation number = -2
o₄⁻ = -2x4 = -8
Mn = ?
As it is a polyatomic ion oxidation number = ion charge
Mno₄⁻ + e⁻ Mno₄⁻ ²
(+7) (-8) = -1 (+6) (-8) = -2
1 electron
Equivalent weight = Molecular weight
No of electrons gained
158
1
20. Theory of redox Titrations
Redox titration consists of two different types of electrodes.
1. Indicator Electrode
2. Reference Electrode
Indicator Electrode:
Used to sense the presence or change in concentration of the oxidized and
reduced forms of a redox couple
Usually, the indicator electrode is an inert noble metal, such as Pt
Pt half reactions at the electrode:
Fe³⁺ + e⁻ Fe²⁺ Eº = 0.767 V
Reference Electrode:
Standard hydrogen electrode and standard calomel electrode used as
reference electrode.
It has accurately maintained potential
Redox potential (also
known as oxidation /
reduction potential
'ORP', pe, E0', or. )
Is a measure of the
tendency of a chemical
species to acquire
electrons from or lose
electrons to an
electrode and thereby
be reduced or oxidized
respectively
21. Redox Indicators
A redox indicator is an indicator compound that changes color at specific
potential differences
A redox indicator compound must have a reduced and oxidized form with
different colors and the redox process must be reversible.
In(oxidation) + ne⁻ = In(red)
Types of Indicators:
Self Indicator :
Potassium permanganate is a good example for the self indicator.
Cerric sulphate and Iodine are other examples
After equivalence point, the titrant will impart a definite pink color at the end of the
titration.
External Indicator :
Based on some visible reactions of the titrated substances with suitable reagent.
End point is marked by failure to elicit reaction
22. Eg: potassium ferricyanide
Titration of ferrous ions with potassium dichromate.
Drops removed during titration on to a tile gives Prussian blue
colour because ferrous ions still present.
At the end point ferric ions are present and does not give colour
Internal or redox Indicators:
These Indicators have different colours in oxidised or reduced form
Most of these are dyes.
Eg: Diphenylamine, Diphenyl Benzidine
Potentiometric method:
This method is useful when suitable indicators are not available
and also when visual indicators fail or have limited accuracy.
23. Cell Representation :
Cu(s) cuso₄ (0.100M) Zncl₂ (0.200M) Zn
Copper electrode immersed in 0.100 M cuso₄ ( First half cell electrode)
Zinc electrode immersed in 0.200 M zncl₂ (second half cell cathode)
If Eº is positive it is spontaneous reaction
If Eº is negative it is non spontaneous reaction which has to be reversed for spontaneous
reaction
24. Measurement of electrode potential
Nernst equation
More positive half cell reaction is by oxidizing agent (anode)
Less positive half cell reaction is by reducing agent ( cathode)
Nernst Equation is the relationship represented between the
concentration and electrode potential for the half cell reaction.
E= Reduction potential
Eº= Standard potential
25. Cerimetric Titration
It is a redox titration in which an iron color change indicates the end point.
26. The potential difference is caused by the ability of electrons to flow from one
half cell to the other.
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30. Iodimetry and Iodometry
Iodimetry:
Principle:
Standard Iodine solution is used as standard.
Iodine is a weak oxidant and it can be reduced by reductants to Iodide ions
I₂ ↔ 2I⁻
Strong reducing agents examples
Stannous chloride, sodium thiosulphate
Sn²⁺ +I₂ Sn⁴⁺ + 2I⁻ (stannous)
2S₂O₃ +I₂ S₄O₆²⁻ + 2I⁻ (sodium thiosulphate)
Weak reducing agents
Arsenic
As³⁺ + I₂ As⁵⁺ +2I⁻
This method is used to quantify oxidising agnets
31. Steps involved:
1. Take a standard solution of Iodine in the Iodine
flask.
2. Add 1ml of Indicator solution
Eg: Starch or sodium starch glycolate
I₂ + Indicator Blue color
3. Titrate the above solution using analyte
solution in burette
Eg : sodium thio sulphate
4. At the equivalence point all the I₂ will react
with the sodium thio sulphate
The solution present in Iodine flask is colorless
In + 2I⁻ No reaction
(colorless)
As the indicator does not react with the Iodide
ions there is no reaction and it will be colourless.
32. Iodometry : principle:
A redox titration where the appearance or disappearance of elementary
iodine indicates the end point.
Liberated Iodine from Iodide is used for titration and the method is
considered as Indirect titration.
KI ↔ K⁺ + I⁻
2I⁻↔ I₂ ↑ + 2e⁻
When we have solutions of strong oxidant CuSO4, KMnO₄ add excess KI solution
in acidic medium so that Iodide ions are oxidized to Iodine ions.
Reaction with Iodide ions with analyte as follows:
(cupric ions) (Iodide) (copper Iodide) (Iodine)
Step 1: 2 Cu²⁺ + I⁻ ↔ 2CuI + I₂ ↑ (Iodine is liberated from Iodide)
Step 2: I₂ + 2S₂O₃ ↔ S₄O₆²⁻ + 2I⁻ (Titrated with sodium thio
sulphate)
2 Cu²⁺ ≡ I₂ ≡ 2S₂O₃
33. Steps involved in Iodometry:
1. Take the analyte solution in the Iodine flask - Cu²⁺ solution
2. Add excess of KI solution so reaction between Cu²⁺ and KI takes
place and it will liberate I₂
3. Add indicator solution in to the Iodine flask which gives blue colour
Starch or Sodium starch glycolate
4. Titrate the above solution by using standard sodium thio sulphate till
the appearance of colourless solution.
I₂ + 2S₂O₃ ↔ S₄O₆²⁻ + 2I⁻
I⁻ + Indicator No reaction
Applications:
Iodometry in its many variations is extremely useful in volumetric
analysis. Examples include the determination of copper(II), chlorate,
hydrogen peroxide, and dissolved oxygen
34. Bromatometry
Principle: Potassium bromate is a strong oxidizing agent in acidic medium.
Bromatometry is a titration process in which the bromination of a chemical
indicator is observed.
Reaction takes place generally in presence of acidic medium 1M Hcl.
The liberated Bro₃⁻(bromate ions) reacts with analyte directly which is a direct
titration.
KBro₃ ↔ K⁺ + Bro₃⁻
Arsenite
3AsO₃³⁻ ( Analyte example) + Bro₃⁻ 3AsO₄³⁻ + Br⁻
This liberated Br⁻ (bromine ion) further reacts with Bro₃⁻ in presence of acidic medium
H⁺ ions.
5 Br⁻ + Bro₃⁻ + 6H⁺ 3 Br₂ + 3 H2O
Br₂ reacts with indicator which is methyl orange, methyl red
This Br₂ oxidizes indicator and solution is colorless.
35. Steps:
1. Take analyte in stoppered conical flask at low temp
2. Add 1M Hcl to analyte to make it acidic medium.
3. Two to three drops of Indicator solution- color becomes red because of
acidic medium
4. Titrate the analyte with standard in the burette
5. End point is red color to colorless
Applications:
Bromination of Indicators can be analysed
It is used to determine arsenic, antimony ,iodide compunds
36. Dichromatometry
Potassium Dichromate is used as standard K2Cr2O7
It is an Oxidizing agent in presence of acidic medium and used as primary
standard
K2Cr2O₇ is used only in acidic medium Cr₂O₇²⁻(dichromate) is rapidly reduced to
Cr³⁺(chromium) which is green in colour.
K ₂Cr₂ O₇ 2K⁺ + Cr₂O₇²⁻
Cr₂O₇²⁻ +14 H⁺ 6e⁻ 2 Cr³⁺ +7H₂O
Iron II salt is used as analyte
6Fe²⁺ + Cr₂O₇²⁻ + 14H⁺ 2 Cr³⁺ + 6Fe³⁺ + 7H₂O
Cr³⁺ is green in color after reduction of Cr₂O₇²⁻ ions
By using simple indicator method end point can’t be determined so we need to use
external indicator method.
Eg: potassium ferricyanide
37. Steps involved in dichromatometry:
1.Take sample solution in conical flask Fe²⁺
2. Add sulphuric acid for acidic medium
Titrate with the potassium dichromate in burette.
Take one drop of solution from the conical flask near the end point and put it
in the external indicator
Before the ed point the ferrous ions reacts with potassium ferricyanide and
converts to ferric so Prussian blue colour will appear
At the end point this reaction will not occur so the colour will not change.
Applications:
Used to determine Iron II salt.
