Volumetric Analysis Titration Basics Reaction, End point & Indicators Types of Titrations Acid – Base Theory & Principles Acid Base titration Non- Aqueous Titration Precipitation Titration Complexometric Titration Oxidation- Reduction Titration Calculation General Information Errors

1. Volumetric Analysis
Dr. A. Amsavel M.Sc., B.Ed., Ph.D.,

2. Content
 Introduction
 Titration Basics
 Reaction, End point & Indicators
 Types of Titrations
 Acid – Base Theory & Principles
Acid Base titration Acid Base titration
 Non- Aqueous Titration
 Precipitation Titration
 Complexometric Titration
 Oxidation- Reduction Titration
 Calculation
 General Information
 Errors

3. Introduction
Tirtimetric / Volumetric analysis:
Volumetric analysis is performed for Quantitative determination of assay /
content. It is simple and commonly used technique in Chemical Industries.
Analysis conducted in Aqueous and non-aqueous medium.Analysis conducted in Aqueous and non-aqueous medium.
 Simple and easy
 Fast and can be done on site
 Less expensive
 Estimation of content or Assay of chemical
 Precise and accurate - depends on method and specificity

4. Requirements of a Titration Reaction
Reaction chose for titration must complete by 99.9 % ie < 0.1
can be un-reacted in the analysis
Reaction must be rapid: analysis shall be performed in a
reasonable time periodreasonable time period
The stoichiometry must be well defined
Reaction completion shall be predicted from equilibrium
constants
A method must be available to determine the equivalence
point

5. Types of Titration
1) Precipitation Titration:
 A(aq) + B (aq) = AB(salt)
2) Acid-Base Titration:
 H+ + OH¯ = H2O (strong acids or bases)
 HA + OH¯ = H2O + A¯ (weak acids)
 A¯ + H+ = H2O + HA (weak bases)
3) Complexometric Titration:
 Zn2+ + 4NH3 = Zn(NH3)42+
4) Redox Titration (Oxidation-Reduction)
 Fe2+ + Ce4+ = Fe3+ + Ce3+

6. Primary Standard
Importance of Primary /Reference to standard
 Primary standard is used to standardize Volumetric Solution (VS)
 The accuracy will be based Quality and Accuracy of the primary
standards used
 A standard is a reference material whose purity and A standard is a reference material whose purity and
composition are well known and well defined
Requirement of Primary Standard
 Usually solid to make it easier to weigh
 Easy to obtain, purify and store, and easy to dry
 Inert in the atmosphere
 High formula weight so that it can be weighed with high precision

7. Endpoint Detection
Endpoint Detection is critical; it is to know the completion of
reaction and accuracy of analysis ;
1 Visual indicators:
Observe a colour change or precipitation at the endpoint.
Reaction completion is identified by addition of external or self indicator
2 Photometry:
Use an instrument to find out the colour change or precipitation
3 Electrochemistry:
Potentiometry : Measure the potential change ( pH electrode)
Amperometry : Measure the change in current between electrodes in
Reaction solution
Conductance: Measure the conductivity changes of solution
Later two method can be used for coloured, turbid solution and
accurate end point

8. Acid-base titration
Understand the following shall be known for accurate
analysis
 Neutralization of reaction during titration
 Neutralization Indicators
 Indicators & mixed indicators
 Neutralization curve
 Non-aqueous titration

9. Principles of Acid & Base
Acids:
Arrhenius acid: Any substance that, when dissolved in water,
increases the concentration of hydronium ion (H3O+)
Bronsted-Lowry acid: A proton donor ie conjugate base
Lewis acid: An electron acceptorLewis acid: An electron acceptor
Bases:
Arrhenius base: Any substance that, when dissolved in water,
increases the concentration of hydroxide ion (OH-)
Bronsted-Lowery base: A proton acceptor ie conjugate acid
Lewis acid: An electron donor

10. Brønsted-Lowry Theory of Acids & Bases
The conjugate acid of a base is the base plus the attached proton .
The conjugate base of an acid is the acid minus the proton.

11. Lewis Theory of Acids & Bases
Lewis acid: An electron acceptor
Lewis acid: An electron donor

12. How to calculate pH ?
A solution contains [H+] of 0.1 mol/L or 10-1 (pH = -log(1 x 10-1)= pH -1
A solution contains [H+] of 0.001 mol/L or 10-3 (pH = -log(1 x 10-3)= pH -3
Q1: Calculate the pH of a solution if [H+] = 2.7 x 10-4 M
pH = -log[H+] pH = -log(2.7 x 10-4) = 3.57pH = -log[H+] pH = -log(2.7 x 10 ) = 3.57
Q2: Find the hydrogen ion concentration of a solution if its pH is 11.62.
[H+] = 10-pH [H+] = 10-11.62 = 2.4 x 10-12M
Q3: Find the pOH and the pH of a solution if its hydroxide ion concentration is
7.9 x 10-5M
pOH = -log[OH-] pOH = -log(7.9 x 10-5) = 4.10
pH + pOH = 14 pH = 14 - 4.10 pH = 9.9

13. An Equation for Buffer Solutions
 In certain applications, there is a need to repeat the
calculations of the pH of buffer solutions many times. This can
be done with a single, simple equation, but there are some
limitations.
 The Henderson–Hasselbalch equation:
• To use this equation, the ratio [conjugate base]/[weak acid]
must have a value between 0.10–10 and both concentrations
must exceed Ka by a factor of 100 or more.
[conjugate base]
pH = pKa + log ––––––––––––––
[weak acid]

14. What is strong acid and strong base?

15. General Knowledge: pH of various solutions
Stomach juice: pH = 1.0 – 3.0 Human blood: pH = 7.3 – 7.5
Lemon juice: pH = 2.2 – 2.4 Seawater: pH = 7.8 – 8.3
Vinegar: pH = 2.4 – 3.4 Ammonia: pH = 10.5 – 11.5
Carbonated drinks: pH = 2.0 – 4.0 0.1M Na2CO3: pH = 11.7
Orange juice: pH = 3.0 – 4.0 1.0M NaOH: pH = 14.0

16. What is an end Point?
Endpoint:
Point Of Neutralization = Equivalence Point
One Equivalent Of Acid reacts with One Equivalent Of
Base

17. Equilibrium Constant : Ka and Kb
The equilibrium constant for a Brønsted acid is represented
by Ka, and base is represented by Kb.
CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO– (aq)
[H O+][CH COO–][H3O+][CH3COO–]
Ka = –––––––––––––––––
[CH3COOH]
[NH4
+][OH–]
Kb = –––––––––––––
[NH3]
NH3(aq) + H2O(l) NH4
+(aq) + OH–(aq)
Notice that H2O is
not included in either
equilibrium
expression.
pH of 1M AcoH =2.4

19. Titration curve: HCl Vs NaOH solution
Volume of
NaOH
ml
1 M Sol
pH
0.1M sol
pH
0.0 0.0 1.0
50.0 0.5 1.5
75.0 0.8 1.8
90.0 1.3 2.3
98.0 2.0 3.0
99.0 2.3 3.3
10.0
12.0
14.0
pH curve of 100ml of HCL titrated against NaOH
of same normality
99.0 2.3 3.3
99.5 2.6 3.6
99.8 3.0 4.0
99.9 3.3 4.3
100.0 7.0 7.0
100.1 10.7 9.7
100.2 11.0 10.0
100.5 11.4 10.4
101.0 11.7 10.7
102.0 12.0 11.0
110.0 12.7 11.7
150.0 13.3 12.3
0.0
2.0
4.0
6.0
8.0
0.0 20.0 40.0 60.0 80.0 100.0 120.0 140.0 160.0
Series1 Series2

20. Titration curve: HCl Vs NaOH solution
NaOH 1 M Sol 0.1M sol
Vol ml pH pH
98.0 2.0 3.0
99.0 2.3 3.3
99.5 2.6 3.6
99.8 3.0 4.0
pH curve close to end point: 100ml of HCL
titrated against NaOH of same normality
99.8 3.0 4.0
99.9 3.3 4.3
100.0 7.0 7.0
100.1 10.7 9.7
100.2 11.0 10.0
100.5 11.4 10.4
101.0 11.7 10.7
102.0 12.0 11.0

21. Acid–Base Indicators
 An acid–base indicator is a weak acid or base.
 The acid form (HA) of the indicator has one color, the
conjugate base (A–) has a different color. One of the
“colors” may be colorless.
 In an acidic solution, [H3O+] is high. Because H3O+ is a
common ion, it suppresses the ionization of the indicator
acid, and we see the color of HA.
 In a basic solution, [OH–] is high, and it reacts with HA,
forming the color of A–.

22. Function of Indicators
How phenolphthalein is behaving in different pH
 Near pH 8, Indicator dissociates and gives red base
Human eye can detect it as a pink tinge at that pH
 Indicators must be carefully chosen so that their colour
changes take place at the pH values expected for an
aqueous solution of the salt produced in the titration.

23. Basis of Indicator selection
Indicator shall be chosen based on the neutralization pH and pKa/b

24. Colours of indicator at different pH

25. Indicators: Color changes against pH

26. Titration Curve: Strong Acid Vs Strong Base
Bromphenol blue,
At the equivalence point in an
acid–base titration, the acid
and base have been brought
together in precise
stoichiometric proportions.
(Endpoint)
Bromphenol blue,
bromthymol blue, and
phenolphthalein all
change color at very
nearly 20.0 mL
At about what volume
would we see a color
change if we used
methyl violet as the
indicator?

27. Titration Curve: Weak Acid Vs Strong Base
The equivalence-point
pH is NOT 7.00 here.
Why not??
Bromphenol blue was ok
for the strong acid/strong
base titration, but it
changes color far too
early to be useful here.

28. Titration Curve: Different combination

29. Non-Aqueous Titration

30. Limitation in Aqueous Titration
Titration in water solutions has limitation:
 To titrate week acids or weak bases
 To titrate separately for a mix of acids (bases) with near
dissociation constants.dissociation constants.
 To determine the substances which are insoluble in water.
 The substances, which are either to weakly acidic or too
weakly basics to give sharp end point in water
 The above can be overcome by non-aqueous to perform
easily and with accuracy

31. Non-Aqueous Titration (NAT)
Non aqueous titration: Titration performed in
solvent medium which does not contain water.
Substance is dissolved in a solvent and titrated
using acid or base as titrant.
 Theory is same as Acid-Base titration
 Reaction carry out in non-aqueous medium
 Extensively used for organic acids and bases
Principle is based on Brønsted-Lowry Theory

32. Where to use NAT
NAT is applied where;
◦ To titrate week acid or weak bases
 To titrate separately for a mix of acids (bases) with near
dissociation constants.dissociation constants.
 To determine the substances which are insoluble in water.
 The substances, which are not give sharp end point in
aqueous solutions, can be titrated non-aqueous solvent ( eg
too weakly acidic or basic)

33. Solvents used in NAT
Solvent which are used in non aqueous titration are called non
aqueous solvent.
Classified as four types:
 Aprotic solvents: Chemically neutral
 Eg. Toluene, carbon tetrachloride
 Protogenic solvents: Acidic nature readily donate protons, Protogenic solvents: Acidic nature readily donate protons,
 Eg. Anhyd. HF, H2SO4
 Amphiprotic solvent: Which are sly ionize and donate and accept
protons,
 Eg Alcohols, weak organic acids.
 Acetic acid makes weak acid into storing base
 Protophilc solvents: Posses high affinity for protons.
 Eg. Liq ammonia, Amine, Ketones
 Increases the acidic strength

34. Acetic Acid
Acetic acid slightly ionise and combine both protogenic and
protophilic propertiesamd able to donate and to accept protons
 Acetic acid is slightly ionize and dissociate to produce protons
CH3COOH ↔ CH3COO- + H+
But in the presence of perchloric acid, a far stronger acid, it willBut in the presence of perchloric acid, a far stronger acid, it will
accept a proton:
CH3COOH + HClO4 ↔ CH3COOH2+ + ClO4 –
The CH3COOH2+ ion can very readily give up its proton to
react with a base, so basic properties of a base is enhanced, so
titrations between weak base and perchloric acid can often be
accurately carried out using Acidic acid .

35. Levelling Solvents:
Levelling Solvents:
In general, strongly protophilic solvents are important to force
equilibrium equation to the right.
CH3COOH + HClO4 ↔ CH3COOH2+ + ClO4 –
This effect is so powerful that, in strongly protophillic solvents, all
acids act as of similar strength.
HB B- + H+
The converse occurs with strongly protogenic solvents, which cause
all bases to act as they were of similar strength.
Solvents, which act in this way, are known as Levelling Solvents.

36. Titration Of Bases
 The titrant should be a very strong acid. Ie Perchloric acid in
Dioxane
 The solvent should not be basic properties
 Aprotic solvents, such as benzene, chloroform, carbon
tetrachloride, chlorobenzene, either alone or mixed with glacial
acetic acid may sometimes be used for titration with acetous
perchloric acid
 To determine primary , secondary , tertrary amines,
heterocyclic amines

37. Titration Of Acids
 The titrant should be a solution of a strong base
 Solutions of quaternary ammonium hydroxides in organic solvents,
e.g. tetra-butylammonium hydroxide in benzene - methanol or IPA
or triethyl-n-butylammonium hydroxide in benzene – methanol.
 Solution of sodium or potassium methoxide in benzene - methanol
 Solvent (s):
 A mixture of benzene and methanol
 very weak acids (e.g., many phenols) usually require a more
strongly basic solvent, such as DMF, anhydrous ethylenediamine or
butylamine
 To determine week organic acids.
Precaution: Amine may absorb carbon dioxide from the atmosphere

38. Selection of Solvents and Titrant
 Acetic acid used for titration of weak bases, Nitrogen containing
compounds
 Acetonitrile / with ACOH: Metal ethanoates
 Alcohols (IPA, nBA) : Soaps and salts of organic acids,
 DMF: Benzoic acid, amides etc DMF: Benzoic acid, amides etc
 Perchloric acid in acetic acid
 Amines, amine salts, amino acids, salts of acids
 Potassium Methoxide in Toluene-Methanol
 Week organic acid
 Quaternary ammonium hydroxide in acetonitrile- pyridine
 Acids, enols, imides & sulphonamides

39. Endpoint Detection
End point detection is critical for titration, it is to know the
completion of reaction and accurate determination.
1) Visual indicators:
• Observe a colour change or precipitation at the endpoint.
– Reaction progress checked by addition of external or self indicator
Indicators: Crystal violet, Methyl red, Thymol blue, & 1-Naphthaol benzeinIndicators: Crystal violet, Methyl red, Thymol blue, & 1-Naphthaol benzein
2) Electrochemistry:
• Potentiometry - measure voltage change ( pH electrode)
• Amperometry - measure change in current between electrodes in
solution
• Conductance – measure conductivity changes of solution
Later two used for coloured, turbid & accurate end point

40. USP Titrimetry <541>

41. Precipitation Titration
 Chloride or Iodine can be titrated against Silver nitrate.
Precipitate of Silver halides formed and completion is detected
as end point
 Reagents used is based on Solubility products of precipitate
 Titration curve: pAg = -log [M n+ ] Vs Volume Titration curve: pAg = -log [M n+ ] Vs Volume
 Concentration of ions
 Eg. Ksol (AgCl) = [Ag+ ] X [Cl - ] = 1.2 X10-10
 Indicators:
 Formation of coloured compound (precipitate /complex)
 Adsorption indicators

42. Precipitation Titration Curve
Copy from Vogel

43. Complexometric titration
 Metal Ions can be titrated with EDTA.
 M + EDTA M(EDTA)
 Complex formation depend on Stability
constant & pH,constant & pH,
 Titration curve pM Vs Vol of EDTA
 Indicators (Metal / metal ion indicators):
 Eriochrome black T , P&R , Calmagite
 M-ln + EDTA M(EDTA) + In
 Eg. Ca & Mg estimation in water
pM = -log [M n+ ]

44. Types of Complexometric titration
 Direct Titration
 Back titration after formation of complex
 Replacement or Substitution Titration (using masking and
demasking, selective demasking agent)demasking, selective demasking agent)
 Separation by precipitation and solvent extraction
Application:
 Factor influence the titration, pH of solution, Concetration of
Metal ion, amount of indicator etc
 Determination of almost all the metals

45. Oxidation- Reduction titration
 Principle is based on Oxidation-Rduction reaction
 Reduction potential is calculated by
 Nernst equation
 E1= E’ + 0.591/n log (ox)/(red)
 E=(E1+E2)/2
 Equivalence point by redox potential Vs Volume
 Example of familiar titrations
 Potasium permanganate Vs Sod. Oxalate
 Sod dichromate Vs Ferric sulphate
 iodometric titration
 iodimetric titrations

46. Oxidation- Reduction Indicators

47. Oxidation- Reduction: Titration Curve
 Eg. Iron(II) can be titrated with
Ce (IV) in dil Sulphuric acid
medium

48. Calculation/ Formula
o Normality: Equivalent wt/1000ml or meq/mL
o Morality: Mole/1000ml
o V1 N1 = V2N2
o N1 = V2N2/V1o N1 = V2N2/V1
 Normality = Weight of sample x 1000 / Eq. wt x V
 Wt of sample (mg) = V x N x Eq. wt
 Assay = Qty estimated in sample x 100/ wt of sample
 Assay = V x N x Eq. wt x 100/ wt of sample x 1000
Where: V- Volume; N-Normality; wt- weight; Eq- Equivalent

49. Titration Error
Possibility of error in the test method:
 End point is critical in the volumetric analysis. Indicator or
other method to determine the end is important;
 The endpoint in the method is not identified exactly at the The endpoint in the method is not identified exactly at the
equivalence point due indicator or incomplete reactions;
Error = Vol. at endpoint – Vol. at equivalence point
 Negative error means endpoint is early or before equivalence
point ; Positive error is due to late or after equivalence point

50. Errors
Know the possible error in Volume and Weight measument
which affects the accuracy of analysis
 10 ml titre volume = 100 %
 If difference in volume is 0.1ml, then error is 1% If difference in volume is 0.1ml, then error is 1%
 5ml titre volume = 100 %
 If difference in volume is 0.1ml, then error is 2%
 Choice to reduce the error or optimum level is 25ml
 25 ml titre volume = 100 %
 0.1ml = 0.4% error

51. Volumetric apparatus
As per USP:
 Burette selection:
 NLT 30% nominal volume (15ml consumption in 50ml burette)
 Micro burette for < 10ml
 Limit of error:
 Volumetric flask: 25ml, 50ml, 100ml is 0.03, 0.05& 0.08ml
 Pipets:5, 10, 25 ml is 0.01, 0.02 &0.03ml
 Burets:10, 25, 50ml is 0.02, 0.1&0.1ml
Tips: out flow NMT 500µl per second for precise analysis

52. Operational & Personal Error
The following variables may affect the accuracy of measurement
Eg. When 10mL volumetric pipette used in the analysis
 Drain time and angle of drain
 Possible beads on the inner surface Possible beads on the inner surface
 Temperature
 Meniscus level
 Touching off last drop
 Rinsing of the pipet with the solution used
 Pipette calibration and etc.

53. Possible Error in Weighing
• Misreading of the balance,
• Balance not level,
• Not cleaning the surface of the balance first,
• Touching the weighed object with moist hands,• Touching the weighed object with moist hands,
• Leaving the balance doors open during weighing,
• Using a miscalibrated balance,
• Not cooling the sample down to near room temperature,
• Not removing a static charge from the sample,
• Excess vibration or air currents from people or nearby equipment, and
• Prolonged time sample left on pan adds/loses moisture.

54. Possible Contamination
Possible contamination in the Laboratory;
 Contaminate a sample during weighing by placing a
contaminated spatula
 Placing the sample on or into a contaminated holder during
weighing,weighing,
 Dropping some lint/hair/skin or sneeze into the sample while
weighing,
 Opening up a bottle of chemicals near the sample being
weighed.
 When performing trace analysis, it is possible for just a
microgram even massive fingerprint!

55. Units of measurement
Name Defining Units
Molarity
(e.g. 0.1200 M)
moles of solute/liter (solutions), or
millimoles/milliliter (solutions)
Percent
(e.g. 23.45 %)
(grams of substance/grams of sample) x
100%, or
centigrams/gram (seldom used)
(e.g. 23.45 %)
centigrams/gram (seldom used)
Parts per million
(e.g 2.34 ppm, 2.34
mg/L)
milligrams/liter (solutions), or
micrograms/milliliter (solutions)
milligrams/kilogram (solids), or
micrograms/gram (solids)
Parts per billion
(e.g. 0.45 ppb, 0.45
ug/L)
micrograms/liter (solutions), or
nanograms/gram (solids)

56. Reference
 Vogel's Text Book of Qualitative Inorganic Analysis 6th
Edition
 Qualitative Chemical Analysis Danial C. Harris
United States Pharmacopoeia United States Pharmacopoeia
 Metrohm – Manual

57. Thank YouThank You