Добавил:
Sekretar
kiopkiopkiop18@yandex.ru
t.me/Prokururor I Вовсе не секретарь, но почту проверяю
Опубликованный материал нарушает ваши авторские права? Сообщите нам.
Вуз:
Предмет:
Файл:Ординатура / Хирургия / Библиотека им академика М.И. Перельмана / Книга_5708_Библиотеки_им_академика_М_И_Перельмана.pdf
7
Analytical Chemistry: An Introduction
Molarity is defined as the number of gram molecules of a solute present in 1 L
of a solution. A solution which contains 1 g molecule of the solute per litre of
the solution is known as a molar solution. It is expressed as 1 M. Gram molecule
is the weight of a substance in grams numerically equivalent to its molecular
weight. Thus, one gram molecular weight of H
2
SO
4
is 98.
Molality is defined as the number of gram molecules of a solute contained
in 1 kg of a solution. A solution which contains 1 g molecule of the solute per
kilogram of the solution is known as a molal solution.
Standardization deals with the determination of the strength in terms of nor-
mality, molarity or molality of a solution. It can be achieved by use of another
standard solution of a substance with which the solution reacts quantitatively
(called a secondary standard), or by use of a carefully weighed substance of a high
purity (called as primary standard). Thus a solution of sodium hydroxide may be
standardized using a standard solution of sulphuric acid (secondary standard)
or by using a sample of potassium hydrogen phthalate of high purity (primary
standard). A substance which is to be used as a primary standard should satisfy
the following conditions:
s
Should have been obtained in a pure state
s
Should have stability, i.e. no change on atmospheric exposure during
weighing, should not be hygroscopic and should not get oxidized by air
or affected by carbon dioxide
s
Should have relatively high molecular weight
s
Should be available in high solubility, i.e. should readily soluble in water
and react instantaneously with the standard solution
The point at which a substance to be analyzed react with the standard solu-
tion of the another substance in chemically equivalent amounts is called as stoi-
chiometric point commonly called as the end point. This change in some property of
the reaction mixture made detectable to the eyes with the help of some complex
organic substances called indicators. The end point can also be detected with the
help of the following:
s
A potentiometer, an instrument that measures the electrode potential of
the solution. These are used for titrations based on a redox reaction; the
potential of the working electrode will suddenly change as the endpoint
is reached.
s
A pH meter, a type of potentiometer that uses an electrode whose poten-
tial depends on the amount of H
+
ion present in the solution. This allows
the pH of the solution to be measured throughout the titration. At the
endpoint, there will be a sudden change in the measured pH. It can be
more accurate than the indicator method, and is very easily automated.
8
Practical Pharmaceutical Analytical Chemistry
s
Thermometric titrimetry, an extraordinarily versatile technique. This is
differentiated from calorimetric titrimetry by the fact that the heat of the
reaction (as indicated by temperature rise or fall) is not used to determine
the amount of analyte in the sample solution. Instead, the endpoint is
determined by the rate of temperature change.
s
Spectroscopy which measures the absorption of light by the solution
during the titration, if the spectrum of the reactant, titrant or product is
known. The relative amounts of the product and reactant can be used to
determine the end point.
The titration is actually stopped at this point.
Types of Titration
Titrations can be classified into five types depending on the type of reaction
involved:
1. Acidbase titrations: Acidbase titrations are based on the neutralization
reaction between the analyte and an acidic or basic titrant. This is also
known as neutralization reactions. This method most commonly uses a pH
indicator, a pH meter, or a conductance meter to determine the endpoint.
2. Redox titrations: Redox titrations are based on an oxidationreduction
reaction between the analyte and titrant. These titrations most commonly
use a potentiometer or a redox indicator to determine the endpoint.
3. Complexometric titrations: Complexometric titrations are based on the
formation of a complex between the analyte and the titrant. These titra-
tions generally require specialized indicators that form weaker complexes
with the analyte.
4. Precipitation titration: Precipitation titration is a type of analysis which involves
the formation of insoluble substance or precipitate. The end point in these
analyses is detected by complete formation of precipitate or by the use of
indicator or by some instrumental methods.
5. Diazotization titration: Diazotization is a reaction that converts an aro-
matic primary amine into diazonium compound. The end point in these
analyses is detected by complete formation of blue colour due to liberation
of iodine on treatment with starch iodide paper.
Chapter
2
Chapter
Volumetric Analysis
2.1 Acid–Base Titration
Acidbase titrations are the analytical procedures that are based on neutralization
reaction between an acid and a base. It involves the reaction of neutralization
between an acid and a base.
Acid
An acid may be defined as any chemical compound which when dissolved in
water undergoes dissociation with the formation of hydrogen ions, carried
as hydronium ion (Arrhenius concept) or a proton donor (BronstedLowry
concept) or an electron pair acceptor (Lewis concept).
It may be classified into
1. Two categories depending on their dissociation constant
a. Strong acid: having
high dissociation constant, e.g. Hydrochloric
acid
b. Weak acid: having low dissociation constant, e.g. Acetic acid
2. Two categories depending on the ability to donate protons
a. Monoprotic acid: which dissociate in water to give one proton only are
called as monoprotic or monobasic acids, e.g. HCl, HNO
3
, acetic acid, etc.
b. Polyprotic acid: which dissociate to give more than one proton like
diprotic or dibasic acids, e.g. H
2
SO
4
(two protons to donate), triprotic
or tribasic acids, e.g. phosphoric acid (three protons to donate).
Base
A base may be defined as any chemical compound which when dissolved in
water undergoes dissociation with the formation of hydroxyl ions (Arrhenius
concept) or a proton acceptor (BronstedLowry concept) or an electron pair
donor (Lewis concept).
10
Practical Pharmaceutical Analytical Chemistry
It may be classified into two categories depending on their dissociation
constant:
1. Strong base: having high dissociation constant, e.g. Sodium hydroxide
2. Weak base: having low dissociation constant, e.g. Ammonium
hydroxide
Neutralization Reaction
The neutralization reaction between acid and base gives salt and water as
product. The reaction between them is classified into four categories depending
on nature of an acid and a base:
1. Strong acid and strong base: which give neutral salt
HCl NaOH NaCl
Hydrochloric acid Sodium hydroxide Sodium chlo
+⎯→⎯⎯⎯
rride
Water
HO+
2
2. Strong acid and weak base: which give acidic salt
HCl NHOH NHCl
Hydrochloric
acid
Ammonium
hydroxide
Ammoni
+⎯→⎯⎯⎯
44
uum
chloride
Water
HO+
2
3. Weak acid and strong base: which give alkali salt
CH COOH NaOH CH COONa
Acetic acid
Sodium
hydroxide
Sodium ac
33
+⎯→⎯⎯⎯
eetate
Water
HO+
2
4. Weak acid and weak base: which give neutral salt
CH COOH NH OH CH COONH
Acetic acid Ammonium
hydroxide
Ammo
34 34
+⎯→⎯⎯⎯
nnium
acetate
Water
HO+
2
A volumetric analysis in which a volume of a sample solution is determined by
titrating it against acid or base using acidbase indicators is known as acidbase
titration.
2.1.1 Classification
It is classified into two main categories:
1. Acidimetry titration: The method to determine strength of a base with a
standard acid is known as acidimetry. In other words, the measurement of
the volume of a standard acid is acidimetry.
11
Volumetric Analysis
It is further classified into two parts:
a. Direct titration: This involves direct titration of an acid with a base, i.e.
titration of base present in the conical flask directly with an acid in the
burette. Example: An assay of sodium carbonate
22
23
HCl Na CO NaCl
Hydrochloric
acid
Sodium
carbonate
Sodium
+⎯→⎯⎯⎯
cchloride
Carbon
dioxide
Water
CO H O++
22
In this titration strength of sodium carbonate (base) is determined by direct
titration with hydrochloric acid (acid) present in the burette.
b. Residual titration: This involves the reaction of base with an excess of
standard acid and the unreacted excess acid is further titrated with
standard base. Example: An assay of zinc oxide.
ZnO H SO ZnSO H O
Zinc oxide
Sulphuric acid Zinc sulphate
W
+⎯→⎯⎯⎯ +
24 4 2
aater
Sodium hydroxide
Sulphuric acid Sod
NaOH H SO Na SO2
24 24
+⎯→⎯⎯⎯
iium sulphate
Water
HO+ 2
2
In this reaction zinc oxide first reacts with excess sulphuric acid (standard
acid) to give zinc sulphate. The excess of sulphuric acid (standard acid) further
reacts with sodium hydroxide (standard base).
2. Alkalimetry titration: The method to determine strength of an acid with a
standard base is called alkalimetry. In other words, the measurement of the
volume of a standard base is alkalimetry.
It is further classified into two parts:
a. Direct titration: This involves direct titration of a base with an acid, i.e.
titration of acid present in the conical flask directly with base in the
burette. Example: An assay of boric acid.
H BO NaOH NaBO
Boric acid
Sodium hydroxide
Sodium metabora
33 2
+⎯→⎯⎯⎯
tte Water
HO+ 2
2
b. Residual titration: This involves the reaction of acid with an excess of
standard base and the unreacted excess base is further titrated with
standard acid. Example: An assay of lactic acid.
CH CH OH COOH NaOH CH CH OH CO
Latic acid
Sodium
hydroxide
33
()
+⎯→⎯⎯⎯
()
OONa
HO
HCl NaOH
Sodium lactate
Water
Hydrochloric acid
Sodium hy
+
+
2
ddroxide
Sodium chloride
Water
NaCl H O⎯→⎯⎯⎯ +
2
12
Practical Pharmaceutical Analytical Chemistry
2.1.2 pH Value
It is defined as the logarithm of the reciprocal of the hydrogen ion concentration
or negative logarithm of hydrogen ion concentration. This term was given by
Sorensen.
The acidity, basicity or neutrality of a solution may be expressed in terms of
the molecular concentration of hydrogen ions.
pH H
log( / )1
where H
is the concentration of H
ions in gram-ion per litre.
Since, the logarithm of a quotient is the difference between the logarithm of
the dividend and the logarithm of the divisor:
pH log(1) log(H )
0 log(H )
log(H )
=−
+
=−
=−
+
+
Say, if a solution has (H
) 0.0001, i.e. (H
) 10
4
gram-ions per litre, the pH
may be calculated as follows:
pH log(H )
log(10 )
(4) 4
4
=−
=−
=−− =
+
−
pH scale ranges from 0 to 14. A neutral solution has pH 7, the acid solutions
have pH less than 7, and basic solutions have pH between 7 and 14.
2.1.3 Indicators
Indicators are the compounds which give visible colour change at the completion
of the reaction. In other words, a substance which indicates the equivalence
point is known as indicators or neutralization indicators or acid–base indicators. These
are organic dyes which change colour according to the pH (hydrogen ion
concentration) of the solution. They exist in different colours at different pH
values and these forms are mutually interconvertible. The point at which an
indicator shows a colour change is called the end point of titration. If the end point
of titration coincides with equivalence point, it is an ideal condition. But usually
there is a small difference between the equivalence point and the end point; this
difference is known as titration error or end point error.
13
Volumetric Analysis
Theories of Indicators
There are three theories of indicators that have been proposed.
1. Litmus theory
2. Ostwald theory
3. Quinonoid theory
Litmus theory According to this theory indicators behave as litmus. Litmus is a
weak acid which gives blue colour in ionic form and red colour in unionized form.
Addition of hydroxide ion in litmus solution causes reaction between hydroxide
and hydrogen ion of litmus releasing ionized litmus ion which gives blue colour.
On the other hand, addition of hydrogen ion gives excess hydrogen ion in
the system which favours the formation of litmus (unionized form) leads to
development of red colour.
Ostwald theory According to this theory acidbase indicator is a weak
organic acid or base which ionizes in aqueous solution to give different
colour.
This theory is based on the fact that unionized and ionized forms of indicator
have different colours which indicate the end point, whenever necessary.
Example: Methyl orange exist in red colour in ionized form (I
n
) and yellow
colour in unionized form (HI
n
).
N
CH
3
CH
3
N
S
O
O
N
O
Methyl orange (HI
n
)----- Yellow colour Methyl orange (I
n
)----- Red colour
O
N
S
O
O
N
N
CH
3
CH
3
H
_
_
Quinonoid theory According to this theory the two colour forms of indicators
are interconvertible, i.e. they are tautomers. One form is known as quinonoid form
which has deep colour and other form is known as Benzenoid form having a light
colour or no colour. The colour change of the indicator occurs when one form is
transformed into the other due to change in pH of the solution.
Benzenoid form
(colourless)
Quinonoid form
(pink Colour)
Example: Phenolphthalein exists as quinonoid form in the pH range of 8.3
10.0. Below and above that pH it exists in benzenoid form (Table 2.1).
14
Practical Pharmaceutical Analytical Chemistry
Table 2.1 Colour changes and pH range of some indicators
Indicator Chemical name pH range Colour (medium)
Acid Alkaline
Phenolphthalein Phenolphthalein 8.3–10.0 Colourless Red
Methyl orange Dimethylamino
phenylazobenzene-sulphonic acid
sodium salt
2.9–4.6 Red Orange
Methyl red 1-Carboxybenzene-
azodimethylaniline
4.2–6.3 Red Yellow
Phenol red Phenol-sulphonphthalein 6.8–8.4 Yellow Red
Cresol red (acid) 1-Cresol-sulphonphthalein 0.2–1.8 Red Yellow
Thymol blue (acid) Thymol-sulphonphthalein 1.2–2.8 Red Yellow
Bromophenol blue Tetrabromophenol-
sulphonphthalein
2.8–4.6 Yellow Blue
Bromocresol green Tetrabromo-m-
cresolsulphonphthalein
3.6–5.2 Yellow Blue
Thymol blue (base) Thymol-sulphonphthalein 8.0–9.6 Yellow Blue
Bromothymol blue Dibromothymol-
sulphonphthalein
6.0–7.6 Yellow Blue
Alizarin yellow R p-Nitrobenzeneazo-salicylic acid 10.1–12.1 Yellow Orange red
O
O
HO
OH
O
O
OH
O
H
H
Benzenoid form (acidic medium)
(colourless)
Quinonoid form (alkaline medium)
(pink colour)
Phenolic salt of carbinol
(strong alkaline medium)
(colourless)
O
O
O
OH
O
2.1.4 Equivalent Weights of Acid and Bases
The equivalent weight of an acid is that weight which contains 1.008 g of
replaceable hydrogen. The equivalent weight of a base is that weight which
contains one replaceable hydroxyl group.
Equivalent weight of an acid can be calculated by using the formula:
Equivalent weight of acid
Molecular weight
Number of replaceable h
yydrogen
15
Volumetric Analysis
For example, hydrochloric acid has one replaceable hydrogen so its equivalent
weight will be equal to its molecular weight. On the other hand, sulphuric acid
and oxalic acid have two replaceable hydrogen hence their equivalent weights
are equal to half the respective molecular weights.
Equivalent weight of hydrochloric acid
Molecular weight
1
36 5
1
.
336 5.
Equivalent weight of sulphuric acid
Molecular weight
2
98
2
49
Equivalent weight of a base can be calculated by using the formula:
Equivalent weight of base
Molecular weight
Number of replaceable h
yydroxyl group
For example, sodium hydroxide has one replaceable hydroxyl group so its
equivalent weight will be equal to its molecular weight (Table 2.2).
Equivalent weight of sodium hydroxide
Molecular weight
1
40
1
40
Table 2.2 Equivalent weights of some common acid and bases
S. no. Acid or base Molecular
weight
Acidity or
basicity
Equivalent
weight
1 Hydrochloric acid (HCl) 36.5 1 36.5
2 Sulphuric acid (H
2
SO
4
) 98.00 2 49.00
3 Nitric acid (HNO
3
) 63.00 1 63.00
4 Oxalic acid {(COOH)
2
·2H
2
O} 126.00 2 63.00
5 Potassium hydrogen phthalate
{C
6
H
4
(COOK)(COOH)}
204.00 1 204.00
6 Acetic acid (CH
3
COOH) 60.00 1 60.00
7 Sodium hydroxide (NaOH) 40.00 1 40.00
8 Potassium hydroxide (KOH) 56.00 1 56.00
9 Sodium carbonate (Na
2
CO
3
) 106.00 2 53.00
10 Sodium bicarbonate (NaHCO
3
) 84.00 1 84.00
16
Practical Pharmaceutical Analytical Chemistry
Experiment 1 To prepare and standardize XN sodium hydroxide.
Primary standard Oxalic acid and potassium hydrogen phthalate
Indicator: Phenolphthalein
Method
s Preparation: Weigh 40X g of sodium hydroxide and dissolve it in
minimum amount of water and finally make the volume 1 L to give
XN solution of sodium hydroxide.
Example: To prepare 0.1 N solution (X = 0.1) of sodium hydroxide,
weigh 40 u 0.1 g or 4 g of sodium hydroxide for 1000 mL.
s Standardization: Weigh accurately 200300 mg of potassium hydrogen
phthalate and dissolve it in minimum amount of water. Add 23
drops of phenolphthalein as an indicator and titrate it with sodium
hydroxide solution until permanent pink colour appears.
End point: Appearance of permanent pink colour.
Reaction involved
NaOH
COOK
COOH
Potassium hydrogen
phthalate
Sodium hydroxide
COOK
COONa
H
2
O
WaterPotassium sodium
phthalate
Calculation
N
W
EV
q
q
1000
where N is the normality of the given sodium hydroxide solution, W is the
weight of primary standard, E is the equivalent weight of primary standard and
V is the volume of sodium hydroxide consumed, i.e. burette reading.
Pharmaceutical uses
1. Sodium hydroxide is also used as a chemical that causes gelling of egg
whites in the production of century eggs.
2. Sodium hydroxide is used in washing or chemical peeling of fruits and veg-
etables, chocolate and cocoa processing, caramel colour production, poultry
scalding, soft drink processing, and thickening of ice cream.