The document provides an overview of oxidation-reduction (redox) reactions and titrations, detailing the principles, types of redox titrations, and relevant indicators. It explains oxidation as the loss of electrons and reduction as the gain of electrons in reactions, along with examples of common redox titrations such as permanganometry, iodimetry, and iodometry. The document also includes methods for preparing solutions and detecting endpoints during these titrations.

2. CONTENTS
• INTRODUCTION – CONCEPT OF OXIDATION AND REDUCTION
• THEORY AND PRINCIPLE
• TYPES OF REDOX TITRATION
• REDOX INDICATORS
2

3. INTRODUCTION
•Chemical reactions in which electrons are transferred from one atom to another
atom are known as Oxidation – Reduction reactions.
•Titration which involve transfer of electrons between the titrant and analyte are
referred to as Oxidation - Reduction reactions.
•Titrations involving both oxidation and reduction reactions are referred to as
Redox titrations.
3

4. •Originally, Oxidation mean addition of oxygen to a substance and also used for
removal of hydrogen.
•Reduction mean removal of oxygen to a substance and addition of hydrogen.
SO2 O SO3
+
H2S O S H2O
+ +
- Addition of Oxygen
- Removal of hydrogen
CuO 2H Cu H2O
+ +
C2H2 2H C2H4
+
- Removal of oxygen
- Addition of Hydrogen
4

5. •Oxidation – reduction reaction as reaction in which electrons are transferred between
species or in which atoms change oxidation number.
•In any reaction in which oxidation occurs, reduction must also occur simultaneously.
•When a substance gives up electrons, there must be another substance to receive
them.
•The first substance is oxidised and the second substance is reduced in an Oxidation –
Reduction reaction.
5

6. OXIDATION
o Oxidation is defined as the part of an oxidation – reduction in which there is a loss of
electrons by a species or an increase in the oxidation number of an atom.
oAn oxidising agent is a species that oxidises another species; thus, oxidising agent is
itself getting reduced in the process and gains electrons from the reducing agents.
oExample:
Potassium permanganate
Potassium bromate
Potassium iodate
Ceric ammonium sulphate
Iodine
6

7. REDUCTION
o Reduction is defined as the part of an Oxidation – Reduction reaction in which there is
a gain of electrons by a species or a decrease in oxidation number of an atom.
oA reducing agent is a species that reduces another species namely oxidising agent, and
the reducing agent is itself oxidised.
oIn this reaction the reducing agent loses electrons which are gained by the oxidising
agent.
o Examples:
Sodium thiosulphate
Ferrous ammonium sulphate
Titanous chloride, sulphate etc.
7

8. Examples of Oxidation – Reduction reactions
•Ferric salt is reduced by a Titanous salt
This equation may be written in the ionic form as
Separated into 2 half equations, one represents the oxidation,
And other is reduction.
TiCl3 FeCl3 TiCl4 FeCl2
+ +
Ti3+
Fe3+
Ti4+
Fe2+
+ +
Ti3+
Ti4+
e
+
Fe3+
e Fe2+
+
8

9. 9

10. 10

11. THEORY AND PRINCIPLE
•In Oxidation – Reduction reactions or Redox titrations, when one substance is
oxidised, some other substance must be correspondingly reduced and when one
substance is reduced, some other substance must be correspondingly oxidised.
•This relationship may be expressed according to Faraday’s law : A change in
charge of one is equivalent to the gain or loss of 96,500 C of electricity for each
formula weight of element or group of elements involved.
11

12. •Since in every oxidation – reduction reaction the charge lost or gained by one
substance must necessarily be gained or lost by another, if follows that there is
always a transfer of electrons in oxidation – reduction reactions.
•The reactant which loses electrons in an oxidation – reduction reaction is the
Reducing agents and can be identified in the equation for the reaction as the reactant
containing a constituent atom or atoms converted to a higher state of oxidation.
Fe2+
Fe3+
e
+
2I-
I2 2e
+
12

13. •The electrons lost by the reducing agent are gained by the Oxidising agent, the
reactant containing a constituent atom or atoms which are converted to a lower state
of oxidation.
OXIDATION NUMBER OR OXIDATION STATE
o Oxidation number is defined to the charge an atom in a substance would have if the
pairs of electrons in each bond belonged to the more electronegative atom.
Ce4+
e Ce3+
+
Fe3+
e Fe2+
+
13

14. TYPES OF REDOX TITRATION
1. Titration involving KMnO4 (Permanganometry/ permanganometric titration)
Eg: Estimation of H2O2
2. Titration involving Iodine
(a) Direct titration/ Iodimetry
Eg: Estimation of ascorbic acid
(b) Indirect titration/ Iodometry
(i) Back titration of excess iodine with sodium thiosulphate
eg: Estimation of anhydrous sodium sodium sulphite
14

15. (ii) Release of iodine and its titration with sodium thiosulphate
eg: Estimation of copper sulphate, chlorinated lime etc..
3. Titration involving Dichromate/ Dichrometry
Eg: Estimation of iron, chromium etc..
4. Titration involving bromine/ Bromimetry
Eg: Estimation of Phenol, isoniazid etc..
5. Titration involving Potassium iodate
Eg: Estimation of KI, weak iodine solution
6. Titration involving Ceric ammonium sulphate (Cerimetry/ Ceriometric titration)
Eg: Estimation of Ferrous sulphate etc..
7. Titanometry
15

16. PERMAGNOMETRY
• Premanganometric titrations are those that use a solution of KMnO4 as an oxidizing agent.
• The solution of KMnO4 is intensely purple coloured.
• As KMnO4 is a strong oxidant, it reduces the intensely purple coloured MnO4
- ions to colourless
Mn2+ ions.
• During the titration (i.e while KMnO4 added to sample) the solution becomes colourless due to
the presence of excess Mn2+ ions.
• Once all the Mn2+ ions have completely reacted, addition of KMnO4 changes the colour of
solution to pink.
• This change in colour indicates the end point, at which both the ions are at equilibrium.
16

17. • KMnO4 act as a self indicator, hence the reaction does not require the addition of any other indicator to
detect the end point.
• KMnO4 act as an oxidising agent not only in acidic solutions but also in alkaline and neutral media. It
exhibits different oxidation states in different media.
1. Acidic media
KMnO4 itself act as an indicator.
MnO4
- is reduced to Mn2+
MnO4
-
8H+
5e-
Mn2+ 4H2O
+ + +
Reduction
2KMnO4 3H2SO4
K2SO4
2MnSO4
3H2O 5[O]
+ + + +
17

18. 2. Basic and Neutral Media
MnO4
- is reduced to MnO2.
Diphenylamine or phenylanthranilic acid – used as indicators in basic and neutral solutions,
because of formation of MnO2.
KMnO4 is more stable in neutral solution than in acidic and basic solutions.
pH determines the stability of KMnO4.
(a) Moderate basic medium
(b) Neutral medium
MnO4
-
2H2O 3e-
MnO2 4OH-
+ + +
Reduced
MnO4
-
4H+
3e-
MnO2
2H2O
+ + +
18

19. Preparation of 0.1N KMnO4 Solution
 3.161 gm of potassium permanganate is needed to prepare 1 litre of 0.1N solution.
Standardisation of 0.1N KMnO4 Solution
Weight accurately about 0.63gm oxalic acid + dissolve in 30ml distilled water and make
up the volume to 100ml in volumetric flask.
Transfer 20 ml in to conical flask + 20ml of dil. Sulphuric acid, heated to 70oC
Titrated hot soln. with 0.1N KMnO4
End point: Appearance of permanent pink colour
19

20. Detection of End-point
• Potassium permanganate is a good example for self indicator in redox titration.
Examples of Permanganometry
• Hydrogen peroxide
• Ferrous sulphate
• Benzoyl peroxide
• Lithium oxalate
20

21. IODIMETRY
•Determination involving direct titration with iodine due to the oxidising power
of iodine in aqueous solution, such type of titrations are called as Iodimetry.
•Iodimetric titration is a type of redox titration where we determine the
concentration of a reducing agents by titrating against a standard iodine
solution.
21

22. PRINCIPLE
• I2 is a weak oxidant and it can be reduced by the reductants.
• Potassium iodide is used to dissolve iodine.
• The excess of iodide ion is present react to form tri-iodide.
• The tri-iodide react with the strong and weak reducing agents like sodium thiosulphate etc..
I2 2I-
KI K+
I-
+
I2 I-
I3
-
+
2S2O3
2-
I2 S4O6
2-
2I-
As3+
I2
As5+
2I-
+ +
+ +
22

23. Preparation of Standard 0.05M Iodine solution
• 0.335gm of iodine dissolves in 20ml KI and make up volume to 1000ml by water at 25oC.
Standardization of 0.05M Iodine solution
• Standardization using Arsenic trioxide
Primary standard substance for standardizing iodine solution.
Substance present in aqu. Soln. as Arsenous acid is oxidized by Iodine.
Arsenic trioxide is slowly soluble in cold water more rapidly soluble in boiling water and
readily soluble in NaOH solution form sodium arsenite.
H3AsO3 I2 H2O H3AsO4 2HI
H3AsO3 I2 H2O H3AsO4 2H+
2I-
+ + +
+ + + +
As2O3 6NaOH 2Na3AsO3 3H2O
+ +
23

24. • If iodine is added to this alkaline solution – forms sodium hypoiodite.
• The excess of sodium hydroxide is neutralized with HCl using methyl orange as the indicator
until yellow colour is changed to pink.
• Sodium bicarbonate is added to neutralize the hydriodic acid, formed in the reversible reaction.
• NaOH or Na2CO3 cannot remove HI because they react with iodine.
• Finally the reaction is,
2NaOH I2 NaIO NaI H2O
+ + +
Sodium
hypoiodide
H3AsO3 H2O I2
H3AsO4 2HI
+ + +
6NaOH 3I2 5NaI NaIO3 3H2O
3Na2CO3 3I2 5NaI NaIO3 3CO2
+
+
+ +
+ +
Na3AsO3 I2 2NaHCO3
Na3AsO4 2NaI 2CO2 H2
+ + + + +
24

25. • Procedure
Weight 0.15gm of arsenic trioxide previously dired @105oC for 1hr + 20ml of 1M NaOH, warming
Dilute with 40ml water + 0.1ml methyl orange + add dil HCl drop wise
Until yellow colour changed to pink.
Add 2gm of Na2CO3 + 50ml water + 3ml starch solution
Titrated with 0.05M Iodine
Permanent blue colour produced
25

26. Detection of End point
1. Self indicator: A solution of iodine in aqu. Iodide has an intense yellow to brown colour.
One drop of 0.05M Iodine solution imparts a perceptible pale yellow colour to 100ml if
water.
2. Starch solution: Starch imparts a deep blue colour to Iodine solution even at very low
concentration. But starch cannot be used in strongly acidic solution due to hydrolysis
of starch molecule. Starch indicator should be freshly prepared or stabilized by addition
of a preservative such as Mercuric iodide (HgI) or Formamide.
26

27. Examples of Iodimetry
(i) Direct Titration with Iodine solution
• Ascorbic acid
• Acetarsol
• Carbarsone
• Analgin
(ii) Back Titration
• Sodium metabisulphite
• Dimercaprol
Assay of Ascorbic acid I.P
• Determination depends upon the quantitative oxidation of Ascorbic acid to
dehydroascorbic acid with iodine in acid solution in presence of H2SO4.
27

28. Procedure
Weigh accurately abt 0.1g dissolve in mixture of 100ml of freshly boiled and cooled water + 25ml
of 1M H2SO4.
Titrated immediately with 0.05M Iodine using starch solution as indicator.
O
C
CH2
H OH
O
H OH
O
O
H
2HI
O
C
CH2
H OH
O
O O
O
H
I2
+ +
H2SO4
Ascorbic acid Dehydro
ascorbic acid
28

29. IODOMETRY
•The titrations in which the equivalent amount of iodine is liberated from the
potassium iodide by the sample and the liberated iodine is titrated against
standard sodium thiosulphate solution, such a type of indirect determination of
strong oxidising agent is called as Iodometry.
29

30. PRINCIPLE
• In iodometric analysis of oxidising agents, an equivalent amt of Iodine is produced when the
sample oxidises with potassium iodide in the presence of mineral acid.
• The equivalent amt of iodine liberated is measured by titration with standard sodium thiosulphate
solution using starch mucilage solution as indicator added towards the end of the titration.
Preparation of 0.1N Sodium Thiosulphate
• 25gm of A.R crystallised sodium thiosulphate is dissolved in one litre of water.
• Sodium thiosulphate is readily obtainable in a state of high purity and efflorescent in nature.
• It’s a reducing agent.
2S2O3
2-
S4O6
2-
2e
+
30

31. Standardization of 0.1N Sodium Thiosulphate
• Standardization using Potassium bromate
Potassium bromate is a primary standard and can be obtained in high state of purity.
This can be determined by the addition of potassium iodide and dil. HCl with potassium
bromate.
After reaction between potassium bromate and potassium iodide, excess KI in acid solution
result in liberation of Iodine .
Its titrated with sodium thiosulphate solution using starch solution as indicator which is added
towards the end of the titration.
KI H2O HI KOH
+ +
2KBrO3 HI HIO3 KBr KIO3 HBr
+ + + +
31

32. • Procedure
Dissolve 0.278gm in 100ml standard flask with water and pipette 20ml from the solution + 2gm of
KI + 3ml of 2N HCl
Titrate with 0.1N Sodium thiosulphate
Using starch solution as indicator added towards the end of titration.
End point: disappearance of blue colour.
HIO3 5KI 6HCl 3I2 6KCl 3H2O
+ + + +
KBrO3 = KIO3 = 3I2
32

33. Detection of End point
1. Starch solution
• Most commonly used indicator in titration involving iodine.
• 2 types of starch preparations available as indicator :
(i) Starch mucilage
(ii) Starch solution
i. Preparation of Starch mucilage
Triturate 0.5gm of starch or soluble starch with 5ml water + sufficient water to produce 100ml
Boil for few mins, cool and filter
• It must be freshly prepared to produce blue colour with free iodine
33

34. ii. Preparation of Starch solution
Triturate 1gm of soluble starch + 5ml water , stirring continuously
Add to 100ml of boiling water containing 10mg of mercuric iodide
• Aqu starch suspension decompose within few days because of bacterial action – adding
Mercuric iodide or chloroform as a bacteriostat.
• Starch indicator solution – added towards end of the titration.
• Disadvantage
 Insolubility in cold water.
Instability of suspensions in water
Gives water insoluble complex with iodine.
 Sometimes drift endpoint which is marked when solutions are dilute.
34

35. 2. Sodium starch glycollate indicator
• Preparation
Dissolve 5gm of finely powdered solid sodium starch glycollate by mixing 1 to 2ml of ethanol + 100ml
of cold water
Boiling for few mins with vigorous shaking
Finely opalescent solution.
• Advantage
Does not form a water insoluble complex with iodine.
With excess of iodine the colour of solution containing 1ml of indicator is green – when iodine
conc. Diminishes the colour changes to blue.
35

36. Examples for Iodometry
1. Copper sulphate
2. Chlorinated lime
3. Phenol
4. Chloramine
36

37. TITRATION INVOLVING POTASSIUM IODATE
• KIO3 is a powerful oxidizing agent.
• React with reducing agents like iodide, iodine, arsenic trioxide.
• Iodate titration performed in presence of alcohol and standard organic acids.
37

38. PRINCIPLE
1. Reaction between KIO3 and reducing agents in moderately acidic solution stop at state iodate
reduced to iodine.
2. Reaction with more powerful reductant like Titanium chloride Iodate reduced to
Iodide.
3. In more strongly acid solution reduction to Iodine monochioride. [more widely used]
KIO3 – act as more powerful oxidizing agents.
KIO3 5KI 6HCl 3I2 6KCl 3H2O
2KIO3
5H3AsO3 2HCl I2 5H3AsO4 H2O 2KCl
+ + + +
+ + + + +
IO3
-
6Ti3+
6H+
I-
6Ti4+
3H2O
+ + + +
IO3
-
6H+
Cl-
4e ICl 3H2O
+ + + +
38

39. Preparation of 0.05M KIO3
Dry A.R KIO3 @ 120oC for 1hr cool in desiccator
Weigh 10.7gm and dissolve in water to 1000ml
Standardization of 0.05M KIO3 Solution
• KIO3 obtained in high state of purity.
• 2 methods:
1.Using Sod thiosulphate
2.Using KI
39

40. 1. Using Sodium thiosulphate
KIO3 + excess KI + dil. HCl
Liberate Iodine
Titrated with Sodium thiosulphate solution using starch solution as indicator
End point: Disappearence of blue colour
Detection of End Point
• Starch solution cannot be used characteristic blue colour of starch-iodine complex is not
formed @ high conc. of acid.
40

41. • Immiscible solvent chloroform/ carbon tetrachloride are added disappearance of last
trace of violet colour due to iodine from solvent is taken as end point.
• Dyes used instead of immiscible solvent
• p-ethoxy chrysoidine – satisfactory reversible indicator.
Dyes Colour change
Amaranth Red to colourless
p-ethoxy chrysoidine Red to orange
Naphthalene black Green to faint pink
41

42. Examples of Titration Involving Potassium Iodate
• Potassium iodide IP’69
• Weak Iodine solution
• Aqueous Iodine solution
• Sodium diatrizoate
42

43. DICROMETRY
• Titration involving Potassium dichromate in acidic solution by using dilute sulphuric acid.
• It is a direct titration method.
• Mainly used for titration of Ferrous sulphate.
• Potassium dichromate is a very strong oxidizing agent.
43

44. PRINCIPLE
• Potassium dichromate is high purity and high stable compound.
• Aqueous solution not attacked by oxidisable impurities composition does not change on
keeping because of this property its an alternative for potassium permanganate.
• Aqueous solution stable towards light.
• Excellent primary standard.
• Titration only take place in acidic solution.
• In neutral solution Potassium dichromate turns out to be a mixture of dichromate and chromate
ion – Hydrolysis of dichromate ions to orange yellow chromate ions [weak oxidising agents].
Cr2O7
2-
14H+
6e-
2Cr3+
7H2O
Cr2O7
2-
6Fe2+
14H+
2Cr3+
6Fe3+
7H2O
+ + +
+ + + +
44

45. • Thus oxidising strength of dichromate is reduced in neutral solutions.
Preparation of 0.0167M Potassium Dichromate
• 4.9g Potassium dichromate dried and kept in dessicator for 4hrs and dissolved in 1000ml distilled
water.
Standardization of 0.0167M Potassium Dichromate
20ml solution + 1g KI + 7ml 2M HCl + 250ml water
Liberated iodine is titrated with 0.1M Sodium thiosulphate
Using 3ml starch solution as indicator added towards end point
Cr2O7
2-
H2O 2CrO4
2-
2H+
orange
yellow colour
+ +
45

46. End point: colour change from blue to light green.
Detection of End Point
o Dichromate not used as a self indicator – because its reduction product (Cr3+) is green which
hinders in the visual detection of end point by masking dichromate colour.
o Indicators used in dichrometry:
1.Diphenyl amine in presence of ortho phosphoric acid
2.N-phenyl anthranilic acid
3.Potassium ferricyanide
o End point: appearance of reddish pink colour.
Cr2O7
2-
14H+
6e-
2Cr3+
7H2O
+ + +
46

47. Advantage of Dichromate over Permanganate
• It is obtainable in a state of high purity and can be used as primary standard solutions of
dichromate in water.
• Very stable towards light and atmospheric conditions.
Examples of Dichrometry
• Determination of iron in sample of iron wire.
• Determine ferrous and ferric ions in a solution.
• Determine ferric ion in solution of ferric alum.
• Determine various salts of iron.
47

48. CERIMETRY
• Redox titrations involving ceric ammonium sulphate as an oxidizing agent is called as Cerimetry.
• Also known as Cerate oximetry.
• Cerium sulphate is a strong oxidizing agent and can be used only in acidic solution – H2SO4.
• Since ceric ammonium sulphate is a strong oxidizing agent we can determine various reducing
substances by simple titration.
48

49. PRINCIPLE
• Sulphuric acid is used to prevent hydrolysis and precipitation of basic salts as in neutral solution
ceric hydroxide or in basic medium as basic salts precipitate.
• Can replace KMnO4 as its very stable.
• Use only in acid solution
• Solution is intense yellow colour in hot solution – act as self indicator.
• Oxidation reaction
• Simple one electron reaction is the main advantage of Cerimetry – complication due to unstable
intermediates eliminated.
Ce3+
Ce4+
reduction
Ce4+
e Ce3+
+
49

50. • In the presence of reducing agent it undergoes reduction to the Cerrous state .
Preparation of 0.1M Ceric Ammonium Sulphate
65gm Ceric ammonium sulphate + 30ml H2SO4, gentle heat + 500ml water
Cool solution and dilute to 1000ml with water
Ce4+
Fe2+
Ce3+
Fe3+
Ce(SO4)2 2FeSO4
Ce(SO4)3 2Fe(SO4)3
+ +
+ +
50

51. Standardization of 0.1N Ceric Ammonium Sulphate
1. Using Arsenic Trioxide (As2O3)
• Reaction between Ceric(IV) Sulphate and As2O3 is very slow @ ordinary temperature Add
Osmium tetroxide soln which catalyst to increase the reaction rate.
• Oxidation of Sodium arsenite [As2O3 + NaOH] arsenate by ceric compound.
• Ferroin used as indicator give end point as colour change from pink to pale blue.
As2O3 6NaOH 2Na3AsO3 3H2O
+ +
Ce(SO4)2 Na3AsO3 H2O 2Ce2(SO4)3 Na3AsO4 H2SO4
4[Ce4+
e Ce3+
]
4Ce4+
Na3AsO3
+ + + +
+
=
51

52. • Procedure
0.2g Arsenic trioxide, dried @ 105oC, 1hr + 25ml 8% NaOH + 100ml water + 30ml dil H2SO4 +
0.15ml Osmic acid
Titrate with 0.1M Ceric ammonium sulphate
Using 0.1ml Ferroin sulphate solution as indicator
End point: Pink to pale blue
Advantages of Ceric(IV) Sulphate
Stable over prolonged periods.
Need not be protected from light
52

53. Employed in determination of reducing agents in presence of a high concentration of HCl.
Cerium(III) ion is colourless and does not obscuse the indicator end point.
Very versatile oxidising agent.
Best standardised with primary standards Arsenic trioxide.
0.1M of Cerium(IV) sulphate solution not highly coloured and does not obscuse vision when
reading meniscus in burettes.
No intermediate products are formed.
Ferroin is very satisfactory indicator in filtration with ceric salts.
53

54. Disadvantage of Ceric(IV) Sulphate
Until recently cerate methods were limited to electrometric titrations till the development of a
satisfactory indicator.
The change of colour of Ceric to Cerrous is from light yellow to colourless thus the end point is
not easily discriminable.
Cerate methods are less economical.
Detection of End Point
Self indicator – bright yellow in colour
Internal indicators used:
a. N-phenyl anthranilic
54

55. b. Ferroin
c. 5,6 –dimethyl Ferroin
Examples for Cerimetry
o Ferrous fumarate
o Ferrous gluconate
o Paracetamol
o Ferrous sulphate
55

56. BROMIMETRY
• Redox titration which involve Potassium bromate is known as Bromimetry.
• Titration process in which bromination of a chemical indicator is observed.
• Potassium bromate – strong oxidising agent.
• Reaction take place in presence of acidic medium 1M HCl.
56

57. PRINCIPLE
• Potassium bromate is reduced smoothly to bromide in presence of HCl and which is the oxidised
to give free bromine by more bromate.
• At the end of the titration free bromine appears.
• Bromine is very volatile, such operation should be conducted @ low temperature in conical flask
fitted with ground glass stopper.
BrO3
-
6H+
6e Br-
3H2O
BrO3
-
5Br-
6H+
3Br2 3H2O
+ + +
+ + +
KBrO3 5KBr 6HCl 3Br2 6KCl 3H2O
+ + + +
57

58. • Excess bromine is determined iodometrically by addition of excess of KI and titration of liberated
iodine with standard thiosulphate solution.
Preparation of 0.0167M Potassium bromate solution
• Dissolve 3.34g of KBrO3 which is dried @ 120oC for 1 – 2 hrs + 15g KBr in 1000ml of distilled
water.
Standardization of 0.0167M Potassium bromate solution
Pipette 25ml of solution in iodine flask + 120ml water + 5ml HCl
2I-
Br2 I2 2Br-
I2 Na2S2O3 Na2S2O6 2NaI
+ +
+ +
58

59. Insert stopper in flask and shake it gently + 5ml KI, shake gently
Allow to stand for 5 mins in dark
Liberated iodine titrated with 0.1M Sodium thiosulphate solution
Using starch solution as indicator added towards the end
End point: disappearance of blue colour.
Detection of End Point
• Indicators used:
a. Methyl orange
59

60. b. Methyl red
c. Naphthalene black
d. Xylidine ponceau
Example of Bromimetry
• Assay of Isoniazide
• Hydroxyl amine
• Determination of arsenic or antimony
1. Assay of Isoniazid I.P
Determined by addition of Potassium bromide and direct titration with potassium bromate in
presence of HCl.
60

61. Bromine is released and reacts with isoniazid
Azo dye Methyl red solution used as indicator becomes decolourised by oxidation at the endpoint.
KBrO3 5KBr 6HCl 3Br2 6KCl 3H2O
+ + + +
N
C
O NH NH2
Br Br O
H2
N
COOH
N2 4HBr
Isoniazid
+ + + +
61

62. TITANOMETRY
• Titration which involve a strong reducing agent Titanous chloride (TiCl3) is called as
Titanometry.
• TiCl3 – strong reducing agents.
• Introduced by Knecht and Hibbert.
62

63. PRINCIPLE
1. It is used for direct titration of ferric salts.
2. Analysis of Dye stuffs and in analysis of organic compounds.
• Dye stuffs – Azo dyes
Eg: Indigo carmine
Methylene blue
Brilliant green
Crystal violet
• Nitro compounds: Nitroglycerine, Nitroguanidine
63
FeCl3
TiCl3 FeCl2 TiCl4
+ +
RNO2
6TiCl3
6HCl RNH2
6TiCl4
2H2O
+ + + +

64. • Nitraso and Azo compounds and Quinones
• Oxidation potential
H+ ions involved in oxidation and potential of system is strongly dependent on acidity of solution.
Smaller the hydrogen ion conc. greater reducing power
64
RNO 4TiCl3 4HCl RNH2
4TiCl4 H2O
RN=NR' 4TiCl3 4HCl RNH2 R'NH2 4TiCl4
+ + + +
+ + + +
Ti3+
3H2O TiO2+
2H+
e
+ + +
2H+
2Ti3+
2Ti4+
H2O
+ +

65. • Great disadvantage error cannot be exactly determined in blank determination.
• TiCl3 – rapidly oxidised on exposure to air due to atmospheric oxygen.
• Precaution: Maintain CO2 or H above liquid.
Preparation of Titanous chloride solution
 In market TiCl3 available as 15 – 20% .
Solution mixed with equal volume of conc. HCl
To remove traces of sulphides boiled for a min and diluted with water to about 10 times the original
volume.
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66. Give a solution approximately N/10 and which should be immediately transferred to the apparatus
in such quantity as to fill the bottle up to the neck.
Burette can be filled as required by releasing the upper of the two clips.
As solution is drawn from the bottle, its place is taken automatically by CO2 from a kipp apparatus
charged with marble and HCl.
First burette – full of solution be rejected and a free current of CO2 should then be allowed to flow
through the apparatus with the lower clip open for 2 mins, in order to make sure that all oxygen has
been displaced
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67. Standardization of 0.1N TiCl3 solution
• Standardized against Ferric ammonium sulphate
• Procedure
Prepare 0.1N solution of ferric ammonium sulphate in 10 % H2SO4.
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68. Pipette 25ml this solution into a conical flask
Pass a current of CO2 through the liquid to replace the air in the flask and maintain a stream of CO2
through the liquid throughout the titration
Add 10% NH4SCN (indicator), solution turns intense red colour
Titrate with 0.1N TiCl3 till red colour disappears.
Stream of CO2 must be maintained throughout the titration
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69. Examples for Titanometry
• Estimation of Ferric ion
• Estimation of Methylene blue
1. Estimation of Ferric Ion
Thiocyanate is a suitable indicator
 Ferric iron has been reduced, the red-brown colour of ferric thiocyanate complex disappears.
 Procedure
50ml ferric solution + 5-10ml of 4N HCl or H2SO4, heat to 50 – 60oC
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70. Add 0.5 – 1ml of 10% KSCN solution + 3 portions of about 300mg of NaHCO3 to drive air from
solution
Titrate with Titanous chloride solution
End point: colour intensity disappears
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FeCl3
TiCl3 FeCl2 TiCl4
+ +

71. REDOX INDICATORS
• The ideal oxidation – reduction indicator will be one with an oxidation potential intermediate
between that of the solution titrated and that of the titrant and which exhibits a sharp, readily
detectable colour change.
• An oxidation – reduction indicator is a compound which exhibits different colours in the
oxidized and reduced forms.
• Ideally, the reactions should be reversible and should give a precise and easily observable colour
change the end-point and donot depend on the specific properties of the oxidizing or reducing
agents used – such indicators are Redox indicators.
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72. A. Internal Indicators
• Ferroin sulphate
• Diphenylamine
1. Ferroin Sulphate
o 1,10 – phenanthroline – Iron (II) complex sulphate.
o Ferroin sulphate is a bright red complex formed by combination of the base.
o Ferroin sulphate complex indicator is readily oxidized reversibly with strong oxidizing agents.
o End point – pale blue in colour.
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N N
Ferroin sulphate

73. oSubstituted Ferroin indicators: Nitroferroin, Dimethyl ferroin.
2. Diphenyl amine
o 1%w/v solution of diphenylamine in conc. Sulphuric acid.
o It will give intense blue – violet colouration at the endpoint.
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NH
Diphenyl amine

74. B. External Indicators
1. Potassium hexacyano Ferrate(III)
o Used in titration of Iron(II) by potassium dichromate.
o Drops of the solution removed to a spotting tile during the titration will give a deep Prussian blue
colour with potassium hexacyano ferrate (III) because ferrous ions are still present.
C. Self Indicator
1. Potassium permanganate
oVarious reducing agents are oxidised by potassium permanganate in a acid solution.
2. Iodine Solution
o Reducing agents can be titrated with Iodine solution without using indicator, because the dark
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75. brown colour of iodine disappears as the result of reduction of I2 to I- ions.
3. Cerium (IV) Sulpate
oYellow in colour reduced to colourless ions.
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