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97
Volumetric Analysis
THEORY Sulphadimidine is a white or almost white crystals or powder, very
slightly soluble in water. It dissolves in solutions of alkali hydroxides and in
dilute mineral acids. In this method sodium nitrite reacts with sulphadimidine in
nitrous acid form to give diazonium salts. At the end point excess of nitrous acid
depolarizes the electrode which causes the flow of current in the galvanometer
and the permanent deflection of galvanometer needle is observed.
Equivalent factor
1 NaNO
2
{ 1 C
12
H
14
N
4
O
2
S
1000 mL of 1 M NaNO
2
{ {(12 u 12) (14 u 1) (4 u 14) (2 u 16)
(1 u 32)} g of C
12
H
14
N
4
O
2
S
1000 mL of 1 M NaNO
2
{ (144 14 56 32 32) g of C
12
H
14
N
4
O
2
S
1000 mL of 1 M NaNO
2
{ 278 g of C
12
H
14
N
4
O
2
S
1 mL of 1 M NaNO
2
{ (278/1000) g of C
12
H
14
N
4
O
2
S 0.278 of
C
12
H
14
N
4
O
2
S
Calculation
1 mL of 1 M NaNO
2
0.278 g of C
12
H
14
N
4
O
2
S
a mL of 1 M NaNO
2
{ (0.278 u a) g of C
12
H
14
N
4
O
2
S
a mL of y M NaNO
2
{ (0.278 u a u y) g of C
12
H
14
N
4
O
2
S
Since 10 mL of C
12
H
14
N
4
O
2
S solution contains (0.278 u a u y) g of
sulphadimidine
1 mL will contain (0.278 u a u y)/(10) g of C
12
H
14
N
4
O
2
S
100 mL will contain (0.278 u a u y u 100)/(10) g of C
12
H
14
N
4
O
2
S
So
g /100 mL of Sulphadimidine
Equivalent factor Burette reading
Mo
=
×
×
llarity NaNO 100
10
2
()
×
Category: Antibacterial
Result The given sulphadimidine solution was found to contain _______
g/100 mL of sulphadimidine.
2.5 Complexometric Titration
The technique which involves titration of metal ions with a complexing agent
or chelating agent (ligand) is commonly referred as complexometric titration.
98
Practical Pharmaceutical Analytical Chemistry
This method represents the analytical application of a complexation reaction. In
this method, a simple ion is transformed into a complex ion and the equivalence
point is determined by using metal indicators or electrometrically.
Various other names such as chilometric titrations, chilometry, chilatometric
titrations and EDTA titrations have been used to describe this method. The
complex formed is known as chelate.
Metal ions in the solution are always solvated, i.e. a definite number of solvent
molecules (usually 2, 4 or 6) are firmly bound to the metal ion. However, these
bound solvent molecules are replaced by other solvent molecules or ions during
the formation of a metal complex or metal coordination compound.
The molecules or ions which displace the solvent molecules are called ligands.
Ligands or complexing agents or chelating agents can be any electron donating
entity, which has the ability to bind to the metal ion and produce a complex
ion. An example of a complexation reaction between Cu (II) ion and four
ammonium molecules in an aqueous solution may be expressed by the following
chemical equation:
OH
2
OH
2
Cu
OH
2
OH
2
NH
3
NH
3
Cu
NH
3
NH
3
2
2
4 NH
3
The bonds in the complex formed are either ordinary covalent bonds in which
the metal and the ligand contribute one electron each, or coordinate bonds in
which both electrons are contributed by the ligand. Thus, the hexacyanoferrate
ion may be considered to consist of three ordinary covalent bonds and three
coordinate bonds, although in the complex the bonds are identical hybrid bonds
which have been shown to be directed towards the apices of a regular octahedron.
Fe
CN
CN
CN
CN
CN
CN
Fe
CN
CN
CN
NC
NC
CN
Fe
CN
CN
CN
NC
NC
CN
3
(a)
(b)
(c)
Hybrid bonds
The hexacyanoferrate iron (III) ion
The negative charge on the complex ion is equal to the total number of the
negative groups minus the valency of the metal ion. When the neutral groups are
only involved, the charge on the complex is positive and is equal to the metal
ion, e.g., [Cu (NH
3
)
4
]
2
.
99
Volumetric Analysis
2.5.1 Classification of Ligands
1. Unidentate ligands:
Ligands that are bound to metal ion only at one place are called unidentate
ligands (one toothed). Example: NH
3
, halide ions, cyanide ions, etc.
Cu NH Cu NH
2
33
2
+
+
+
⎯→⎯⎯⎯
←⎯⎯⎯⎯
()
2. Bidentate and multidentate ligands:
Ligands that are known contain more than one group, capable of binding
with metal ions are known as multidentate ligands or chelating agents. They
include bidentate ligands (two donor atoms), tridentate ligands (three donor
atoms), quadridentate ligands, etc.
Example: Ethylene diamine (bidentate ligand)
Chelate Compound or Chelate
Complexes involving simple ligands, i.e. those forming only one bond are
described as coordination compound. A complex of a metal ion with two
or more groups on a m ultidentate ligand is called as a chelate or a chelate
compound. There is no fundamental difference between coordination
compound and a chelate compound except that in a chelate compound,
ring influence the stability of compound. Thus, a chelate can be described
as a heterocyclic ring structure in which a metal atom is a member of ring.
The stability of a chelate is usually much greater than that of corresponding
unidentate metal complex.
Reagent EDTA
Disodium salt of EDTA is a water soluble chelating agent and is always preferred.
It is nonhygroscopic and a very stable sequestering agent (Ligands which form
water soluble chelates are called as sequestering agents).
There are chelating agents that form water insoluble chelates with metal ions,
e.g., oxime or 8-hydroxy quinoline.
OH
N
EDTA and 8-hydroxy quinoline are important reagents used in analytical
chemistry. Sequestering agents are used to liberate or solubilize metal ions. The
agents which form water insoluble chelates are used to remove the metal ions
from solution by precipitation.
EDTA has the widest general application in analyses because of the following
important properties:
100
Practical Pharmaceutical Analytical Chemistry
s It has low price.
s The special structure of its anion which has six ligand atoms.
s It forms strainless five–membered rings.
Factors influencing EDTA reactions are:
1. The nature and activity of metal ion.
2. The pH at which the titration is carried out.
3. The presence of interfering ions such as CN
, citrate, tartrate, F
and other
complex forming agents.
4. Organic solvents also increase the stability of complex.
2.5.2 Effect of pH on Complex Formation
Edetic acid ionizes in four stages (pK
1
2.0, pK
2
2.67, pK
3
6.16 and
pK
4
10.26) and, since the actual complexing species is Y
4
, complexes will
form more efficiently and be more stable in alkaline solution. If, however, the
solubility product of the metal hydroxide is low, it may be precipitated if the
hydroxyl ion concentration is increased too much. On the other hand, at lower
pH values when the concentration of Y
4
is lower, the stability constant of the
complexes will not be so high. Complexes of most divalent metals are stable in
ammoniacal solution. Those of the alkaline earth metals, such as copper, lead
and nickel, are stable down to pH 3 and hence can be titrated selectively in
the presence of alkaline earth metals. Trivalent metal complexes are usually still
more firmly bound and stable in strongly acid solutions. For example, the cobalt
(III) edetate complex is stable in concentrated hydrochloric acid. Although most
complexes are stable over a fair range of pH, solutions are usually buffered at
a pH at which the complex is stable and at which the colour change of the
indicator is most distinct.
2.5.3 Principle of Complexometric Titration
Many principles of acidbase titrations are used in complexometric titration.
In complexometric titration, the free metal ions disappear as they are changed
into complex ions. In acidbase titrations, the end point is marked by sudden
change in pH. Similarly, in EDTA titration, if we plot pM (negative log of metal
ion concentration) vs volume of titrant, the pM rapidly increases (Fig. 1). This
sudden pM raise results from removal of traces of metal ions from solution by
EDTA.
Any method, which can determine this disappearance of free metal ions,
can be used to detect end point in complexometric titrations. End point can
be detected usually with an indicator or instrumentally by potentiometric or
conductometric (electrometric) method.
There are three factors that are important in determining the magnitude of
break in titration curve at end point.
101
Volumetric Analysis
1. The stability of complex formed: The greater the stability constant for
complex formed, larger the charge in free metal concentration (pM) at
equivalent point and more clear would be the end point.
2. The number of steps involved in the complex formation: Fewer the number
of steps required in the formation of complex, greater would be the break
in titration curve at equivalent point and clear would be the end point.
3. Effect of pH: During a complexometric titration, the pH must be a constant
by use of a buffer solution. Control of pH is important since the H
ion plays
an important role in chelation. Most ligands are basic and bind to H
ions
throughout a wide range of pH. Some of these H
ions are frequently displaced
from the ligands (chelating agents) by the metal during chelate formation.
2.5.4 Methods of End Point Detection
End point in complexometric titration can be detected by the following two
methods:
Indicators
The end point in complexometric titrations is shown by means of pM
indicators. The concept of pM arises as follows:
If K is the stability constant,
K [MX]/[M][X]
then [M] [MX]/[X]K
or log [M] log [MX]/[X] log K
and pM log [X]/[MX] pK
Therefore, if a solution is made such that [X] [MX],
Then, pM pK
or pM pK
where K is the dissociation constant
This means that, in a solution containing equal activities of metal complex
and free chelating agent, the concentration of metal ions will remain roughly
constant and will be buffered in the same way as hydrogen ions in a pH buffer.
Since, however, chelating agents are also bases; equilibrium in a metalbuffer
solution is often greatly affected by a change in pH. In general, for chelating
agents of the amino acid type (e.g., edetic acid and ammonia triacetic acid),
it may be said that when [X] [MX], pM increases with pH until about pH
10, when it attains a constant value. This pH is, therefore, usually chosen for
carrying out titrations of metals with chelating agents in buffered solutions.
The pM indicator is a dye which is capable of acting as a chelating agent to
give a dyemetal complex. The latter is different in colour from the dye itself
102
Practical Pharmaceutical Analytical Chemistry
and also has a low stability constant than the chelatemetal complex. The colour
of the solution, therefore, remains that of the dye complex until the end point,
when an equivalent amount of sodium edetate has been added. As soon as there
is a slightest excess of edetate, the metaldye complex decomposes to produce
free dye; this is accomplished by a change in colour.
Metal indicators must comply with the following requirements:
x Compound must be chemically stable throughout the titration.
x It should form 1:1 complex which must be weaker than the metalchelate
complex.
x Colour of the indicator and the metal complexed indicator must be
sufficiently different.
x Colour reaction should be selective for the metal being titrated.
x The indicator should not compete with the EDTA.
Mechanism of action of indicator
Let the metal be denoted by M, indicator by I and chelate by EDTA. At the onset
of the titration, the reaction medium contains the metalindicator complex
(MI) and excess of metal ion. When EDTA titrant is added to the system, a
competitive reaction takes place between the free metal ions and EDTA. Since
the metalindicator complex (MI) is weaker than the metalEDTA chelate,
the EDTA which is being added during the course of the titration is chelating
the free metal ions in the solution at the expense of the MI complex. Finally,
at the end point, EDTA removes the last traces of the metal from the indicator
and the indicator changes from its complexed colour to its metal free colour
(Table 2.5).
Fig. 2.1
Volume of EDTA
Log [M]
End point
103
Volumetric Analysis
Table 2.5
Indicators used in complexometric titrations
S.No.
Name of the
indicator
Colour change pH range Metals detected
1.
Mordant black II
Red to blue
67
Ca, Ba, Mg, Zn, Cd, Mn,
Pb, Hg
Eriochrome blackT
Solochrome blackT
2.
Murexide
or
Ammonium purpurate
Violet to blue 12 Ca, Cu, Co
3. Catechol-violet Violet to red
810
Mn, Mg, Fe, Co, Pb
4.
Methyl blue Blue to yellow
45
Pb, Zn, Cd, Hg
Thymol blue Blue to grey
1012
5. Alizarin Red to yellow 4.3 Pb, Zn, Co, Mg, Cu
6.
Sodium alizarin
sulphonate
Blue to red 4 Al, thorium
7. Xylenol orange Lemon to yellow
13
Bi, thorium
45
Pb, Zn
56
Cd, Hg
Example:
N
OH
O
3
S
O
3
S
O
2
N
N
HO
M
2
N
O
M
O
N
O
2
N
2H
Mordant black II
(Eriochrome black T, Solochrome black T)
Blue (pH 10)
Pink
Structures of some pM indicators
OH
OH
CH
2
N(CH
2
COOH)
2
O
O
Alizarine fluorine blue
OH
NaO
2
.O
2
S
N
N
HO
Calcone
104
Practical Pharmaceutical Analytical Chemistry
N
N
N
O
OH
OH
N
N
OH
HO
OH
Murexide
O
O
OH
OH
SO
3
Na
Sodium alizarin sulphonate
Instrumental Methods of End Point Detection
1. Spectrophotometric detection: The change in the absorption spectrum when
a metal ion of a complexing agent is converted to the metal complex,
or when one complex is converted to another can usually be detected
more accurately and in more dilute solution by spectrophotometry
than by visual methods. In practice, an indicator giving a colour change
in the visible region is generally, employed, but coloured ions may be
titrated without an indicator using spectrophotometric methods.
2. Amperometric titration: The effect of complex formation on the half-wave
potential of an ion is to render it more negative. If the electrode potential
is adjusted to a value between that of the half-wave potential of the free
cation and that of the complex, and disodium edetate solution is added
slowly, the diffusion current will fall steadily until it equals the residual
current, i.e. until the last trace of free cation has been complexed. This
is the end point and the amount of standard disodium edetate solution
added is equivalent to the amount of metal present.
3. Potentiometric titration
4. High frequency titrator
2.5.5 Types of Complexometric Titrations
Complexometric titrations are of four types:
Direct Titration
It is the simplest and the most convenient method used in chelometry. In
this method, the standard chelon solution is added to the metal ion solution
until the end point is detected. This method is analogous to simple acidbase
titrations.
Example: Calcium gluconate injection and calcium lactate tablets, etc.
Limitations:
Slow complexation reaction
Interference due to presence of other ions
105
Volumetric Analysis
Back Titration
In this method, excess of a standard EDTA solution is added to the metal solution,
which is to be analyzed and the excess is back titrated with a standard solution
of a second metal ion.
Example: Determination of Mn. This metal cannot be directly titrated with
EDTA because of precipitation of Mn(OH)
2
.
Replacement titration
In this method the metal, which is to be analyzed, displaces quantitatively the
metal from the complex. When direct or back titrations do not give sharp end
points, the metal may be determined by the displacement of an equivalent
amount of Mg or Zn from a less stable edetate complex.
Mn Mg EDTA Mg Mn EDTA
+− +−
+⎯→⎯⎯⎯ +
22 22
Mn displaces Mg from Mn EDTA solution. The freed Mg metal is then
directly titrated with a standard EDTA solution. In this method, excess
quantity of Mg EDTA chelate is added to Mn solution. Mn quantitatively
displaces Mg from Mg EDTA chelate. This displacement takes place because
Mn forms a more stable complex with EDTA. By this method Ca, Pb, Hg
may be determined using Eriochrome blackT indicator.
Indirect Titration
This is also known as alkalimetric titration. It is used for the determination
of ions such as anions, which do not react with EDTA chelate. Protons from
disodium edetate are displaced by a heavy metal and titrated with sodium
alkali.
Mn H X MX H
n
n
+
()
+
+⎯→⎯⎯⎯ +
2
2
4
2
-
-
Example: Barbiturates do not react with EDTA but are quantitatively pre-
cipitated from alkaline solution by mercuric ions as 1:1 complex.
Experiment 1 To prepare and standardize 0.05 M disodium edetate.
Primary standard Granulated zinc (Zn) or calcium carbonate (CaCO
3
)
Indicator: Mordant black II or hydroxyl naphthol blue
Method
x Preparation: Weigh accurately 18.61 g of disodium edetate and dissolve it in
water. Finally adjust the volume to 1 L to give 0.05 M solution.
x Standardization: Weigh accurately 0.8 g of granulated zinc and dissolve it in
a mixture of 12 mL dilute hydrochloric acid and 0.1 mL bromine water by
gentle warming. Boil to remove excess of bromine, cool and add sufficient
106
Practical Pharmaceutical Analytical Chemistry
water to produce 200 mL. Pipette out 20 mL of the solution and neutralize
it with 2 M sodium hydroxide solutions. Dilute to about 150 mL with water
and add sufficient ammonia buffer (pH 10) to dissolve the precipitate and
add 5 mL of the buffer in excess. Add 50 mg of mordant black II mixture
and titrate with disodium edetate solution until the solution turns green.
or
Weigh accurately about 200 mg of CaCO
3
in a titration flask. Add 50 mL
of water and minimum quantity of dilute HCl to dissolve CaCO
3
. Adjust
the pH of the solution to 12 by adding NaOH. Add 300 mg of hydroxyl
naphthol blue indicator and titrate it with the prepared M/20 disodium
edetate solution, until the solution is deep blue in colour.
End point: Appearance of green (with Zn) or deep blue (with CaCO
3
) colour.
Reactions involved
With zinc:
Zn HCl ZnCl H
Zinc
Hydrochloric acid
Zinc chloride Hydrog
+⎯→⎯⎯⎯ +2
22
een
Zinc chloride
Disodiume
ate
ZnCl C H N Na C H
2 1014 2 2 1014
+⎯→⎯⎯⎯
det
NNZn NaCl
Zinc salt of disodiume
ate
Sodium chloride
2
2
det
+
With calcium carbonate:
2
32
HCl CaCO CaCl
Hydrochloric
acid
Calcium
carbanate
Calciu
+⎯→⎯⎯⎯
mm
chloride
Water Carbon
dioxide
Calcium
chloride
HO CO
CaCl I
++
+
22
2
nnCaIn
Ca In
Indicator
Calcium indicator
complex
Calcium in
⎯→⎯⎯⎯ −
−
ddicator
complex
Disodiume ate
CHNNa CHNCa+⎯→⎯⎯⎯ +
10 14 2 2 10 14 2
det
IIn
Indicator
free
()
THEORY The zinc reacts with hydrochloric acid to give zinc chloride. Excess
of hydrochloric acid is removed by using sodium hydroxide. The zinc chloride
further reacts with disodium edetate to give zinc salt of edentate. The reaction
of zinc chloride with disodium edetate requires a pH 10 which is maintained by
ammonia buffer.
or
The hydochloric acid solubilizes the calcium carbonate by converting it
to calcium chloride. The sodium hydroxide makes the solution alkaline and
maintains the pH at about 12 so that the CaEDTA complex would be stable