Standardization is the process of determining the concentration of a solution by reacting it quantitatively with a standard solution. Primary standards are highly pure, stable reagents used to standardize and prepare secondary standards. They have properties like high purity, stability, and molecular weight. Primary standards like sodium chloride and sodium carbonate are used to standardize silver nitrate and acid solutions respectively. Accurate standardization is important for titration calculations which rely on knowing the true concentration of the standard solution.

1. Standardization
It is the determination of strength of a given
sample solution and can be determined by
reacting the solution quantitatively with
standard solution.

2. Primary standard:
In chemistry, a primary standard is a reagent
that is very pure, stable, not hygroscopic, and
has a high molecular weight.
Ideally, it’s also non-toxic, inexpensive, and
readily available.
A primary standard provides a reference to
find unknown concentrations in titrations and is
used to prepare secondary standards

3. Primary Standard Properties
High purity
High stability/low reactivity
High equivalent weight (to reduce mass
measurement error)
Not hygroscopic (to reduce mass changes
from water absorption)
Non-toxic or low toxicity
Inexpensive
Readily available

4. Why Primary Standards Are Used
Chemicals react according to mole ratios. A
titration determines the concentration of an
unknown solution based on volume of a solution
with known concentration needed to react with
the solution of unknown concentration. But, the
accuracy of the calculation relies on truly
knowing the concentration of one solution.

5. So, for example, sodium hydroxide (NaOH) reacts in a 1:1 ratio
with hydrochloric acid (HCl). But, sodium hydroxide is not a
primary standard because it’s typically impure.
Sodium hydroxide is highly hygroscopic and absorbs carbon
dioxide from air, so if you weigh out a sample, some of that mass
is actually water and carbon dioxide. That throws off any
calculation involving the mole ratio because you really have less
sodium hydroxide in the solution than you think.
Meanwhile, sodium carbonate (Na2CO3) is a good primary
standard for a reaction with hydrochloric acid because it’s
available at high purity, has a higher molecular weight than
sodium hydroxide, and isn’t as hygroscopic

6. Examples of Primary Standards
Sodium chloride (NaCl): for silver nitrate (AgNO3)
reactions
Sodium carbonate: for titrating acids
Potassium hydrogen phthalate or : for titrating
bases
Potassium hydrogen iodate [KH(IO3)2 : for
titrating bases
Potassium dichromate (K2Cr2O7 ): for redox
reactions
Zinc powder (after dissolving it in hydrochloric
acid or sulfuric acid): to standardize EDTA (ethylene
diamine tetraacetic acid) solutions

7. Secondary Standards
A secondary standard is a reagent that has
been standardized against a primary standard.
In other words, a secondary standard’s
concentration is known by titrating it against a
measured volume of a primary standard instead
of by weighing it out and dissolving it in a solvent.
A secondary standard may be less pure and
more reactive than a primary standard, but it still
upholds some of the properties of a standard.
It’s stable enough that its concentration
remains known for a long time. Sodium hydroxide
(NaOH) is a common secondary standard.

9. Errors
The difference between the measured value &
True value is called Errors

10. Sources of Errors
1. Human source
2. Instrumantal,glassware & apparatus
3. Experimental condition
4. Constituent used in analysis
5. Procedure

11. Types of Errors
1. Systemic Errors
- Personal errors
- -instrumental errors
- Reagent errors
- Method errors
- Proportional errors (impure NaoH in titration of
HCl)
- constant Error
2. Random error

12. Methods of minimizing Errors
1. Proper calibration of instrument, apparatus & applying
correction
2. By running control determination: running a control
determination parallerly to the sample by taking
standard under same experimental condition.
- However standard should contain same weight of the
constituent present in unknown sample

13. 3. By running blank determination : (to know the
effect of impurities)
4. By comparing result of independent methods
(HCl titration with standard NaoH &
precipitation titration with AgCl2)
5. By using standard addition

14. Accuracy
Refers to concordance between the
measured & true value
(how much close the measured value with
true value)
Accuracy = 1 / error
Example: Two analyst showing their
experimental result ( calculation of % drug
content in PCM tablet ) (standard value is 98
%)

15. • Analyst A: 98,97,93,92
• Analyst B: 91,95,97,93
• For analyst A: 98+97+93+92 = 380/4= 95%
4
• For analyst B: 91+95+97+93 = 376/4= 94%
4
Conclusion: result of Analyst A is more accuracy

16. Precision
Refers to the degree of closeness between the
several measurement of same quantity of
sample by same method
Precision and accuracy are correlated to each other.
Accuracy represents correctness of measurement while
precision represent reproducibility
Precision always represent accuracy but higher degree of
accuracy doesn’t necessarily represent that the method is
precise.

17. Example: Two analyst performed analysis on a
sample that contain 70.10±0.04 % of a constituent.
Analyst A: 70.07,70.18,70.14,70.08 (average value
is 70.12)
Analyst B: 70.27,70.28,70.26,70.28 (average value
is 70.27)
Conclusion:
Analyst A result is more accuracy but not precise
Analyst B result is more precise but not accurate

18. Precision can be classified into 3 types:
1. Repeatability : same analysis is performed
several times in same condition & same time
(same day)
2. Intermediate precision: same analysis is
performed by changing analyst, chemical,
instrument etc.(separate day)
3. Reproducibility: same analysis is performed
in different laboratory by same method

19. Significant figures
Significant figures are the digit which are necessary to
express the result.
Accuracy of a measurement depends upon the number
of significant figures.
More the significant figures more will be the accuracy
Example: the value showing rough balance is 98.2 & the
sample value showing in analytical balance is 98.2365
In 98.2365 the figures 98.236 known with certainty
while sixth digit with some uncertainty
Significant figures are independent of decimal point.
Example 12.3,1.23,0.123 all contain three significant
figures

20. Rules about significant figure
• Those digits which are non-zero are significant.
For example, in 6575 cm there are four significant figures and in 0.543
there are three significant figures.
• If any zero precedes the non-zero digit then it is not significant. The
preceding zero indicates the location of the decimal point, in 0.005
there is only one and the number 0.00232 has 3 figures.
• If there is a zero between two non-zero digits then it is also a
significant figure.
For example; 4.5006 have five significant figures.
• Zeroes at the end or on the right side of the number are also
significant.
For example; 0.500 has three significant figures
• Zeros placed on the right of the last non-zero digit after the decimal
point are significant. For example, 0.00153100 contains six significant
digits.

21. For correction of significant number, rounding
of the number made by adding 1 to the last figure
if the rejected figure is 5 or above.
Example: if u need only first three digit from
1.236 & 1.233,then after round off we get 1.24 &
1.23

23. Solved Examples on Significant Figures
Example 1: Write down the significant figures of
the list of numbers 367, 0.0075, 56.004, 98.70,
and 230.00
Solution: From the list of numbers, let us find out
the significant figures of each number.
367 - Three significant figures
0.0075 - Two significant figures
56.004 - Five significant figures
98.70 - Four significant figures
230.00 - Five significant figures

24. UNIT-II
Acid Base Titration
1. The Arrhenius theory of acids and bases states
that “an acid generates H+ ions in a solution
whereas a base produces an OH– ion in its solution”.
2. The Bronsted-Lowry theory defines “an acid as a
proton donor and a base as a proton acceptor”.
Acid Base+ H+
3. Finally, the Lewis definition of acids and bases
describes “acids as electron-pair acceptors and
bases as electron-pair donors”.

25. An acid-base titration is a quantitative analysis of acids and bases;
through this process, an acid or base of known concentration
neutralizes an acid or base of unknown concentration.
The titration progress can be monitored by visual indicators, pH
electrodes, or both.
The reaction’s equivalence point is the point at which the titrant
has exactly neutralized the acid or base in the unknown analyte; if
you know the volume and concentration of the titrant at the
equivalence point, you can calculate the concentration of a base
or acid in the unknown solution.

26. Theory of Acid Base Indicator
Indicators(acid-base indicator) are substances that are used
in quantitative analysis to obtain information about the
completion of a chemical reaction.
In other words, they are used to determine the end point
of the reactions.
Acid base indicators are also called hydrogen ion
indicators or neutralization indicators. The color of acid base
indicators varies in acidic and alkaline mediums, that is, their
color changes with appropriate change in pH value.

27. Example:
The pH range of phenolphthalein is 8.0 – 9.8.
In solutions with pH less than 8.0, it is colorless and
in solutions with pH of more than 9.8, its color is
pink. Phenolphthalein is colorless in weak alkaline
solution and pink in strong alkaline solution
“Acid base indicators are those substances which
have one color in acidic solution and another color
in alkaline solution i.e. their color changes with
proper change in pH value.”

28. It is clear that it is not necessary that the color of
an acid base indicator is the same in all alkaline
solutions. Likewise, it is also not necessary that
the color of an acid base indicator be the same in
all acidic solution
Theory of Indicators: (explain the behavior of
acid base indicators.)
1. Ostwarld Theory
2. Quinonoid Theory

29. Ostwarld Theory
Theory of indicator action was suggested by W.ostwald.
Most of acid base indicator are weak organic acids or bases
& colour changes due to ionization of indicator
Example: phenolphthalein is a weak acid. It is colorless.
Consider HPh is weak organic acid indicator & in aqueous
solution it dissociate as
HPh(colorless) ⇋ H(colorless) + Ph(Pink)
(undissociated) (dissociated)
In the above reaction, dissociated indicator have different
colour than undissociated indicator

30. In acid base titration , if the indicator placed in
acidic solution of acid HA, following equilibrium
exist
HPh ⇋ H+ + Ph-
HA ⇋ H+ + A-
Hph is a weak acid ,so it poorly dissociate & due to
common ion effect the dissociation further
decreased. The equilibrium is shifts towards left
side.
Therefore in presence of acid the dissociation of
Hph is almost negligible & indicator is mostly in
undissociated form & show undissociated indicator
colour ( colourless)

31. But if the same indicator added in alkali solution
of base BOH, the dissociation of Hph will be
increased
HPh ⇋ H+ + Ph- (dissociation increased)
BOH ⇋ OH- + B+
The hydrogen ion removed by OH- by formation
of water molecules & equilibrium is shifted
towards right side & concentration of Ph- will give
dissociated indicator colour. i.e phenolphthalein
give pink colour

32. Similarly if the indicator is weak organic base , it poorly
dissociated in alkali solution but dissociation
increased in acidic solution.
Example :
Thus the color change of methyl orange can also be
explained. Methyl orange is a weak base and thus dissociates
–
MeOH(Yello) ⇋ Me+(Red) + OH-(colorless)
In basic medium:
Due to the common ion effect in the presence of OH ions in
alkaline medium, ionization of MeOH is very low, i.e the
concentration of Me is very low and the color of the solution
remains yellow.

33. In acidic medium:
H ion obtained from acid in acidic medium
combines with OH ions derived from MeOH to
form water.
The amount of ionization of MeOH increases to
restore equilibrium. Thus the concentration of Me
ions in the solution increases and the color of the
solution becomes red.

34. MeOH ⇋ Me ++ OH-
HCl ⇋ H ++ Cl-
H ++ OH- ⇋ H2O
Me+ + Cl- ⇋ MeCl

35. Quinonoid theory
Quinonoid theory suggest that the colour change
of indicator is due to structural changes of Indicator.
In acidic solution indicator exist in different
structural form & in basic solution it exist in
different structural form
These structural form have different colours,
therefore in acidic solution indicator give different
colour and in basic solution different colour.

36. According to this theory, indicator are aromatic carbonic
compounds used in acid base titrations.
They are a mixture of at least two movable forms
Among these, one form is in acidic medium and the
other form in alkaline medium in greater proportion.
These two forms are called benzenoid and quinonoid
forms.
Both these forms have different colors. The color of the
quinonoid form is darker than the color of the benzenoid
form

37. Therefore, when the pH value of a solution changes, there is a
change in the benzenoid form to the quinonoid form or the
quinonoid form to the benzenoid form causing a change in color.

38. Example:
The acidic solution of phenolphthalein, being in its benzenoid form, is
colorless.
When this solution is added to such a base that the solution becomes
alkaline, the entire benzenoid form turns into a quinonoid form.
The color of phenolphthalein’s alkaline solution also becomes red as
the color of the anions obtained from the quinonoid form is red.

39. Similarly methyl orange is also found in two tautomeric
forms.
In acidic solution, this compound remains in quinonoid
form and in alkaline solution it remains in benzenoid form.
The color of quinonoid form is red and the color of
benzenoid form is yellow orange.
Therefore, the color of acidic solution of methyl orange is
red and the color of alkaline solution is yellow orange.

40. Classification of acid base Titration
1. Titration of strong acid with strong base
2. Titration of weak acid with strong base
3. Titration of weak base with strong acid
4. Titration of weak base with weak acid
5. Titration of polybasic acid with strong base

51. H
a
n

52. Neutralization curve
Neutralization curve represents the neutralization reaction
between acid & base in neutralization titration.
It is plotted between the pH and volume of acid or base
added.
The following information derived from neutralization curve
1. The curve is useful in studying the neutralization process
by studying the change in hydrogen ion concentration
during titration
2. It represent the progress of acid base titration
3. It indicate the end point of titration
4. It indicate the sensitivity of titration & chances of error
5. On the basis of pH conditions near inclination point the
selection of indicator for particular titration is made.

53. Non Aqueous Titration
Non aqueous titration are the titration which carried out using non
aqueous solvent (other than water)
Why choose non aqueous titration:
1. If the reactants or productants are insoluble in water.
2. React with water
3. Too weak base or acid that cannot be titrated in water due to their
poor infliction in pH at end point in water
Example of non-aqueous solvent:
- Acetone
- Glacial acetic acid
- Formic acid
- Ethylene diamine
- chloroform

54. Concept of Non Aqueous Titration

55. Advantages:
1. Weak acid & base give poor end points in aqueous titration but
in non aqueous titration they give good end point.
2. The substance which are poorly soluble in water but soluble in
non aqueous organic solvent can be determined by non-aqueous
titration
3. Mixture of two or more acids can be determined
Disadvantages:
1.Solvents used in titration are expensive
2.temperature, moisture & carbon dioxide should be controlled
3. Indicators need to be prepared in non aqueous medium
4. It is not an environment friendly method

56. Non aqueous solvent
1. Protophilic solvents: These are the solvent which have high affinity for proton &
are basic in nature.
HB + S SH+ + B-
Acid solvent solvated proton conjugate base
Example: Liquid ammonia, ketones and amines
2. Protogenic solvents: which readily donate protons & are acidic in nature
Example: sulphuric acid, hydrogen fluoride
3. Amphiprotic solvents: Solvents which are able to donate as well as accept proton(
Protogenic & protophilic solvent)
Example: Acetic acid
CH3COOH CH3COO- + H+
While in presence of strong acid like perchloric acid, it act as base & accept proton
CH3COOH + HClO4 CH3COOH2 + ClO4
4. Aprotic solvent: solvents which are chemically inert & does not donate or accept
proton , therefore they don't have any basic or acidic nature.
Example: Toluene, chloroform, benzene etc.

57. Effects of solvents

58. ALKALIMETRY & ACIDIMETRY
Alkalimetry refers to the determination of
strength of basic substances by titrating them
with standard acid solution.
Acidimetry refers to the determination of
strength of acidic substances by titrating them
with standard base solution.
In non-aqueous titration weak base & weak
acid are determined.

59. Titration of weak base:
Titration of weak base done by standard acid
mainly perchloric acid.
Acetic acid used as solvent for dissolving weak
base.
Reaction involved can be explained as follows.

62. Titration of weak acids:

63. 1.
Estimation of Sodium Benzoate :-
Sodium Benzoate :
Formula : C7H5NaO2 Mol. Wt. 144.1
Sodium Benzoate contains not less than 99.0 per cent and not more than 100.5 per cent of C7H5NaO2, calculated on the
dried basis.
Description : A white, crystalline or granular powder or flakes; odourless or with a faint odour; hygroscopic.
PRINCIPLE:
Sodium benzoate is a preservative that comes in the form of a white granular or crystalline powder. The base sodium
benzoate is dissolved in glacial acetic acid that increases the strength of weak base sodium benzoate. . As an indication,
the 1-naphtholbenzein solution is used to detect the endpoint of the reaction.
This practical is divided into two parts:
A: Preparation and standardization of perchloric acid (0.1M).
B: To perform the assay of sodium benzoate.

64. Preparation of 1-naphtholbenzein indicator
solution
Take 200 mg of 1-naphtholbenzein and dissolve in 50 ml
of anhydrous glacial acetic acid in a volumetric flask, and
properly mixing it.
Once it has completely dissolved, make up the volume to
100 ml.

65. Preparation of 0.1M solution of HClO4 and its standardization: Dissolve
8.5 ml of 72% HClO4 in about 900 ml glacial acetic acid with constant stirring,
add about 30 ml acetic anhydride and make up the volume (1000 ml) with
glacial acetic acid and keep the mixture for 24 hour.
Acetic anhydride absorbed all the water from HClO4 and glacial acetic acid
and renders the solution virtually anhydrous. HClO4 must be well diluted with
glacial acetic acid before adding acetic anhydride because reaction between
HClO4 and acetic anhydride is explosive.
Standardisation of HClO4 :To 500 mg of potassium acid phthalate add 25 ml
of glacial acetic acid and add few drops of 5% w/v crystal violet in glacial acetic
acid as indicator. This solution is titrated with 0.1 M HClO4. The colour
changes from blue to blue green.
• Equivalent factor:
• 0.01441 g of sodium benzoate would be equivalent
to one ml of 0.1 M perchloric acid.

66. TITRATION PROCEDURE
All glassware should be cleaned and dried according to standard
laboratory procedures.
Before filling the burette for the titration, rinse it with distilled water
and then pre-rinse it with a portion of the titrant solution(PERCHLORIC
ACID).
Pre-rinsing is required to make sure that all solution in the burette is
the desired solution, not a contaminated or diluted solution.
Note: Since it is a non-aqueous titration, it needs to be extra careful
during the cleaning and drying of glassware. Glassware is required to be
clean with acetone and dried for purpose.
Take the unknown stock solution of titrant in a clean and dry beaker
then fill the burette using the funnel.
Remove air bubbles from the burette and adjust the reading to zero.
Take 0.25 gm of sodium benzoate and pour it into a conical flask.
Add 20.00 ml anhydrous glacial acetic acid, sonicate to dissolve, and if
required warm it at 50°C.
Then add 2 drops of 1-naphtholbenzein indicator solution.

67. Titrate the sample solution with the perchloric
acid solution until the endpoint is reached.
The actual endpoint of the titration is indicated
by a change of blue color to emerald green.
To get accurate results, repeat the titration three
times.
Properly record the readings of the burette.
Take their mean and calculate the percentage
purity of sodium benzoate.
For the blank determination, repeat the titration
as directed above but without the sodium
benzoate (B).

68. Calculations:
% purity of sodium benzoate = V x E x AM x 100 / W x RM
Where,
V is the volume of perchloric acid solution used
V= A-B
A is a volume of perchloric acid solution used in titration with sodium benzoate
B is a volume of perchloric acid solution used in titration without sodium benzoate
E is an equivalent factor
AM is actual molarity
RM is a required molarity (MOLARITY GIVEN)
W is the weight of the sample

69. 2.
Estimation of Ephedrine Hydrochloride :-
Ephedrine Hydrochloride :
Formula : C10H15NO,HCl Mol. Wt. 201.7
Ephedrine Hydrochloride contains not less than 99.0 per cent and not more than 101.0 per cent of C10H15NO,HCl
calculated on the dried basis.
Description : Colourless crystals or a white, crystalline powder; odourless. It is affected by light.
For the Estimation of Ephedrine Hydrochloride :
Preparation of 0.1N solution of HClO4 and its standardization: Dissolve 8.5 ml of 72%
HClO4 in about 900 ml glacial acetic acid with constant stirring, add about 30 ml acetic
anhydride and make up the volume (1000 ml) with glacial acetic acid and keep the mixture
for 24 hour.
Acetic anhydride absorbed all the water from HClO4 and glacial acetic acid and renders the solution
virtually anhydrous.
HClO4 must be well diluted with glacial acetic acid before adding acetic anhydride because reaction
between HClO4 and acetic anhydride is explosive.

70. Standardization of HClO4 :To 500 mg of potassium acid
phthalate add 25 ml of glacial acetic acid and add few
drops of 5% w/v crystal violet in glacial acetic acid as
indicator. This solution is titrated with 0.1 HClO4.
The colour changes from blue to blue green.
Assay Procedure : Weigh accurately about 0.17 g of
Ephedrine Hydrochloride, dissolve in 10 ml of mercuric
acetate solution, warming gently, add 50 ml of acetone
and mix.
Titrate with 0.1 M perchloric acid, using 1 ml of a
saturated solution of methyl orange in acetone as
indicator, until a red colour is obtained. Carry out a blank
titration.
Equivalent factor : 1 ml of 0.1 M perchloric acid is
equivalent to 0.02017 g of C10H15NO HCl.

71. Calculations :
% Ephedrine Hydrochloride =
% purity of Ephedrine hydrochloride= V x E x AM x 100 / W x RM
Where,
V is the volume of perchloric acid solution used
V= A-B
A is a volume of perchloric acid solution used in titration with sodium benzoate
B is a volume of perchloric acid solution used in titration without sodium benzoate
E is an equivalent factor
AM is actual molarity
RM is a required molarity (MOLARITY GIVEN)
W is the weight of the sample