Complexometric TITRATION FOR PG IST SEM

1. Dr.C.SREEDHAR
HEAD OF DEPARTMENT
PHARMACEUTICAL ANALYSIS
KARNATAKA COLLEGE OF PHARMACY,BANGALORE
M.PHARM 1st YEAR

2. 2
•INTRODUCTION:
Complexometric titration is a type of
titration based on complex formation
between the analyte(metal ion) and titrant
(complexing agents).
• Complexometric titrations are particularly
useful for determination of a mixture of
different metal ions in solution. An
indicator with a marked color change is
usually used to detect the end-point of the
titration.

3. 3
Complexometric titration is the volumetric
method of analysis in which the simple metal
ion is converted into complex by addition of
reagent known as ligand or complexing agent
and is used to estimate polyvalent ions i.e.,
divalent, trivalent.
E.g.:-

4. LEWIS DEFENITION FOR COMPLEXOMETRIC
TITRATION
 By Lewis defenition ,complexation is an acid-base
reaction between a ligand ,a Lewis base or electron
donor , and the metal ion, a Lewis acid or electron
acceptor.
 The nature of the chemichal bond between the metal
ion and the ligand may vary from covalent to ionic.
 There is a specific number of ligands that are bound to
each metal ion . This is termed as the co-ordination
number and usually a characteristic for each metal ion

5. Metal ion Co-ordination number
Silver I
Copper I
Mercury I
2
2
2
Silver II 4
Zinc II
Magnesium II
Cobalt II
Nickel
4 or 6
Aluminium III 6
Calcium II 6
Lead IV 6
Iron II and Iron III 6
Barium II 6

6. Cu2+
+ C C
H
NH2
O
OHH2
Glycine
Cu
O
N
H
O
N
H
CH2H2C
C C
O
O
+ 2H+
• Chelate: when a metal ion coordinates with
two or more donor groups of a single ligand to
form a five- or six-member heterocyclic ring.
• Unidentate: a ligand that has a single donor
group Ex: NH3
• Bidentate: a ligand that has two groups
available for covalent bonding Ex: glycine
• Tridentate, tetradentate, ……

7. Structure of H4Y and its
dissociation products.
Note that the fully
protonated species H4Y
exists as the double
zwitterion with the
amine nitrogens and
two of the carboxylic
acid groups protonated.
The first two protons
dissociate from the
carboxyl groups, while
the last two come from
the amine groups.

8. Structure of a
metal/EDTA
complex. Note
that EDTA
behaves here as a
hexadentate
ligand in that six
donor atoms are
involved in
bonding the
divalent metal
cation.

9. § Complexes of EDTA and Metal Ions
The reagent combines with metal ions in a 1:1
ratio regardless of the charge on the cation.
Ag+
+ Y4-
AgY3-
Al3+
+ Y4-
AlY-
Mn+
+ Y4-
MY(n-4)+
KMY=
[MY(n-4)+
]
[Mn+
] [Y4-
]

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Some of the basic considerations
 Complexometric titrations are particularly useful for the
determination of a mixture of different metal ions in solutions.
 If Complex formed is water soluble and stable those Water soluble
complexes are called sequestering agent.
 If a complexing agent can form more than one bond with polyvalent
ion, then it is considered as polydentate and called as chelating agent.
 The difference between sequestering agent and chelating agent is that
the sequestering agent remains in the solution in the chelated
condition where as chelating agent forms an insoluble complex with
the metal ion.
 In all EDTA Complexometric reactions the ratio of EDTA to the
metal ion is 1:1
 E.g.; dimethylglyoxime and salicylaldoxime are e.g. for chelating
agent where as EDTA is an e.g. for sequestering agent.

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TYPES OF COMPLEXOMETRIC
TITRATIONS
1. Direct titration
2. Back titration
3. Replacement of one complex
by another
4. Alkalimetric titration of metals

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Direct titration
A standard solution is added directly to the solution of the
substance being determined are referred to as a direct titration.
 The solution containing the metal ion to be determined is
buffered to the desired pH and few drops of indicator is added.
 The contents are titrated against standardized disodium edetate
solution till the end point shown by the colour change.
 A blank titration is carried out omitting the substance to be
determined, but contains all the other solutions like buffer &
indicator.
 The volume of edetate consumed in blank titration is substracted
from that obtained in the original estimation
Examples of such estimation:
Bismuth - Nitrate, carbonate, oxynitrate, sub nitrate
Calcium - Chloride, gluconate, lactate, carbonate,
Magnesium - Carbonate, oxide, stearate, sulphate,

13. 13
1. Assay of magnesium sulphate
Principle:
 Magnesium sulphate forms stable complex with disodium edetate in presence of
strong ammonia-ammonium chloride buffer.
 It is assayed by direct titration method.
 Mordant black mixture is used as indicator for detection of end point {red to blue}
and standard 0.05 M disodium edetate is employed as titrant.
Mg2+ + EDTA → Mg EDTA
Materials required:
 Standard disodium edetate 0.05 M, Magnesium sulphate sample (0.4 g),
 Ammonia-ammonium chloride buffer (10ml), Mordant black mixture as indicator.

14. Preparation of reagents
 0.05M Disodium EDTA – 18.6g EDTA Disodium is
added to AR grade water to produce 1000ml.
 Ammonia-ammonium chloride buffer solution- dissolve
67.5g of ammonium chloride in 650ml of AR grade water
and add strong ammonia solution .Make up the volume to
1000ml with AR grade water.
 If it is MgSO4 7H20-0.4g OR 0.3g of anhydrous
magnesium sulphate.

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 Assay:
 Weigh accurately about 0.3 g anhydrous magnesium sulphate
,transfer into 250ml conical flask and dissolve in 50 ml of
water. Add 10 ml of strong ammonia-ammonium chloride
solution, and titrate with 0.05M disodium
ethylenediaminetetraacetate using 0.1 g of mordant black II
mixture as indicator, until the pink colour is discharged from
the blue.
Equivalent factor:
0.006019g Mg SO4 ≡ 1 ml 0.05M disodium EDTA
Percentage purity w/w = 0.006019*volume of Edetate*M of edetate *100
Weight taken * 0.05

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 2. Assay of calcium carbonate:
 Principle:
 This is based on the principle of complexometric
titration. Disodium edetate is used as sequestering
agent. Dil HCl is used in order to dissolve calcium
carbonate in water because as such it is insoluble in
water.
 Acid will generate Co2 which is formed as follows.
 2HCl + CaCo3 → CaCl2 + H2CO3
 H2CO3 → H2O +CO2
 Ca CO3+H20+CO2→ Ca (HCO3)2
Insoluble Soluble

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PROCEDURE
 Weigh accurately about 0.1g and dissolve in 3ml of
dilute hydrochloric acid and 10 ml of water. Boil for
ten minutes, cool, dilute to 50 ml with water. Titrate
with 0.05 M disodium ethylenediamine tetra acetate to
within a few ml of the expected end-point. Add NaOH
and a mixture of calcon and anhydrous sodium
sulphate as indicator and continue the titration until
the colour changes to a full blue colour.
 Each ml of 0.05M disodium ethylenediamine tetra
acetate is equivalent to 0.005004 g of CaC03.

18. Back Titration
 Is necessary if the analyte precipitates in the absence of
EDTA, if it reacts too slowly with EDTA, or if it blocks the
indicator (i.e. cannot be displaced by EDTA from the
indicator). A known excess of EDTA is added to the analyte
solution; the excess is titrated with a standard solution of
another metal ion, which must not displace analyte from
EDTA (often Mg2+ is used)! Determination of the excess not
required by the sample hence the amount of volumetric
solution used by the substance is determined.
 If the analyte precipitates without EDTA, an acidified
solution of analyte (metals in general dissolve better at low
pH) is treated with excess EDTA, then adjusted to the
required higher pH.

19. In general this method is used for :
1.Volatile substances eg: ammonia
2.Insoluble substances eg: calcium carbona
3.Substances which need excess of reagent
eg: lactic acid
19

20. 20
PROCEDURE:
 Weigh accurately about 80mg of sample.
 To the sample add 45ml of water,0.2g of sodium chloride
and 20ml ethanol.
 Heat to boiling and add 0.05M lead nitrate, intially
dropwise and then more rapidly, with constant stirring.
 Heat to coagulate the precipitate, cool to room
temperature.
PROCEDURE FOR ASSAY OF ALUM

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Filter and wash the residue with small
volumes of alcohol.
To the filtrate and washings, add
hexamine(1g) and titrate the excess lead
nitrate with 0.05M disodium edetate using
xylenol orange solution as indicator.
The end point is yellow.

22. Assay of aluminium hydroxide gel
Procedure:
 Weigh accurately about 5g of the substance in a 100ml of
conical flask, add 3ml of 6M HCL .
 Dissolve the suspension by a slight warming on a water
bath.
 Cool to room temperature and transfer to a 100ml
volumetric flask and make up to the mark with water.
22

23.  Pipette out 25ml of this solution in to a 250ml conical flask
and add 40ml of the standard 0.05MEDTA solution .
 Dilute with 80ml of water and add 2-3 drops of methyl red
indicator
 Add 1M NaOH to neutralise the mixture as will be
recognised by change of colour of the indicator from red to
yellow.
 This solution is warmed on a water bath for 30mns then add
3g of hexamine and xylenol orange as indicator,the titration
mixture becomes yellow in colour.
23

24. Now the mixture is titrated with a
standard 0.05M lead nitrate solution.
End point is violet colour.
24

25. ASSAY OF Manganese II
 Mn II cannot be titrated with EDTA alkaline
solution,owing to precipitation of the manganese
hydroxide .An excess of EDTA is added to an acidic
solution of the manganese salt,an ammonia buffer is
used to adjust the solution PH to 10 and the excess
EDTA remaining after chelation is titrated with a
standard Zn II solution using Erichrome black T as
indicator .

26. Replacement Titration
 The method of displacement titration involves the
quantitative displacement of a second metal ( M II)
from a complex by the metal ( M1) being
determined.The freed second metal is then directly
titrated by a standard chelon solution .

27. ASSAY OF CALCIUM GLUCONATE
 Calcium gluconate is assayed by complexometric
replacement titration method.In the estimation of
calcium ions,ammonia-ammonium chloride buffer
and a known volume of magnesium sulphate is added
.This forms Mg-EDTA complex.In the reaction a stable
calcium EDTA complex is formed and magnesium ions
are liberated which is titrated with standard disodium
edetate solution.
 BLANK-Another titration is carried out with same
quantities of reagent under same condition without
sample(calcium gluconate)

28. b) The difference between (A-B) gives the amount of disodium
edetate consumed by the calcium gluconate.
 ca2+ + Mg-EDTA Ca-EDTA+ Mg2+
 Mg2+ + EDTA Mg-EDTA

29. 470
Figure 17-6
EDTA titration curves for
50.0 mL of 0.00500 M Ca2+
(K’CaY=1.75×1010) and Mg2+
(K’MgY=1.72×108) at pH 10.0.
Note that because of the larger
formation constant, the
reaction of calcium ion with
EDTA is more complete, and a
larger change occurs in the
equivalence-point region. The
shaded areas show the
transition range for the
indicator Eriochrome Black T.

30. PROCEDURE
 Transfer an accurately weighed quantity of about 0.5g
of calcium gluconate to a 250ml conical flask ,add 50ml
of water to dissolve it.
 Add 5ml of magnesium sulphate (0.5M) and 10ml strong
ammonia-ammonium chloride solution and titrate the
mixture with the 0.5M disodium edetate solution using
mordant black mixture as indicator to blue colour.Note the
reading as A.
 Repeat the experiment with the same reagents except
sample.Blank reading B.

31. The difference between (A-B) gives the amount of disodium
edetate required by the sample.
 Each ml of 0.5M disodium edetate = 0.02242g of
calcium gluconate.
 Percentage w/w=
---
Volume of EDTA required (A-B)* M of EDTA *0.002242*100
Weight of sample* 0.05

32. PM INDICATORS
 The indicators used in Complexometric titrations are called as
pM indicators, since these indicators are responsive to the
concentration of the metal ions in solution.
 pM indicators are chelating agents which give specific colour on
the formation of complex with metal ions.
 The free chelate and the complex with metal ions have two
different colours.
 Hence based upon the concentration of the metal ion, the colour
of the indicator changes and so the end point of such titration
can be known.
 pM is the negative logarithm of metal ion
concentration.
 pM = -log [M]

33. MUREXIDE

34. DESCRIPTION
 This is the ammonium salt of purpuric acid, and its anion has the structure.
 It was the first metal ion indicator to be employed in the EDT A titration.
 Murexide solutions are reddish-violet up to pH = 9 (H4D-),
 Violet from pH 9 to pH 11 (H3D2 -), and
 Blue-violet (or blue) above pH 11(H2D3 -).

 These colour changes are probably due to the progressive displacement of
protons from the imido groups.
 Murexide forms complexes with many metal ions: only those with Cu, Ni, Co,
and the lanthanoids and their colours in alkaline solution are orange (copper),
yellow (nickel and cobalt), and red (calcium); the colours vary somewhat with
the pH of the solution.
 Murexide may be employed for the direct EDT A titration of calcium at pH =11,
the colour change at the end point is from red to blue-violet, similarly the
colour change in the direct titration of nickel at pH 10-11 is from yellow to blue-
violet.

35. PREPARATION
 Aqueous solutions of Murexide are unstable and must
be prepared each day. The indicator solution may be
prepared by suspending 0.5 g of the powdered dyestuff
in water, shaking thoroughly, and allowing the
undissolved portion to settle. The saturated
supernatant liquid is used for titrations.

36. ERICHROME BLACK
T(SOLOCHROME BLACK T)

37. DESCRIPTION
 This substance is sodium 1-(1-hydroxy-2-naphthylazo)-
6-nitro-2-naphthol-4-sulphonate).
 In strongly acidic solutions the dye tends to polymerise
to a red-brown product, and therefore the indicator is
rarely applied in the EDT A titration of solutions more
acidic than pH = 6.5.
 The sulphonic acid group gives up its proton long
before the pH range of 7-12, which is of more interest
for metal-ion indicator use.

38. CONTINUED……
 This colour change can be observed with the ions of Mg, Mn, Zn, Cd,
Hg, Pb, Cu, and the Pt metals.
 To maintain the pH constant buffer mixture is added, and to keep the
above metals in solution weak complexing reagent such as ammonia or
tartarate is added.
 The cations of Cu, Co, Ni, Al, forms such stable indicator complexes
that the dye can not be liberated by adding EDT A: direct titration of
these ions using Solochrome Black indicator is not possible because
the metallic ions are said to 'block' the indicator.
 However, with Cu, Co, Ni, and Al a back-titration can be carried out, by
titrating the excess of EDTA with standard me or magnesium ion
solution.
 The indicator solution is prepared by dissolving 0.2 g of the dyestuff in
15 ml of triethanolamine with the addition of 5 ml of absolute ethanol
to reduce the viscosity.

39. Determine the transition ranges for Eriochrome
Black T in titrations of Mg2+ and Ca2+ at pH 10.0,
given that (a) the second acid dissociation
constant for the indicator is
(b) The formation constant for MgIn- is
(c) Ca2+ Kf = 2.5x105
H2O + HIn2
-
In3-
+ H3O+
K2 = 2.8 x 10-12
Mg2+
+ In3-
MgIn-
Kf = 1.0 x 107

40. SOLOCHROME DARK BLUE OR
CALCON

41. CONTINUED…..
 This is sometimes referred to as Eriochrome Blue Black RC.
 It is sodium 1-(2-hydroxy-l-naphthylazo)-2-naphthol-4-
sulphonate.
 The dyestuff has, two ionisable phenolic hydrogen atoms; the
protons ionise stepwise with pK’s of 7.4 and 13.5 respectively.
 An important application of the indicator is in the
Complexometric titration of calcium in the presence of
magnesium is carried out at a pH of about 12.3 in order to avoid
the interference magnesium.
 The magnesium is precipitated quantitatively as the hydroxide.
The colour change is from pink to pure blue.
 The indicator solution is prepared by dissolving 0.2 g of the
dyestuff in 50 ml of methanol.

42. XYLENOL ORANGE

43. CONTINUED…..
 This indicator is prepared by the condensation of O-
cresolsulphonephthalein (Cresol Red) with
formaldehyde and iminodiacetic acid.
 This is 3 ,3' -bis [N,N -di( carboxymethyl )-amino
methyl ]-o-cresolsulphonephthalein
 Special feature is dyestuff retains the acid-base
properties of Cresol Red and displays metal indicator
properties even in acid solution (pH = 3-5).
 Acidic solutions of the indicator are coloured lemon-
yellow and those of the metal complexes coloured red.

44. CONTINUED…..
 Direct EDTA titrations of Bi, Th, Zn, Cd, Pb, Co, etc.,
are readily carried out and the colour change is sharp.
 By appropriate pH adjustment certain pairs of metals
may be titrated successfully in a single sample
solution.
 Thus bismuth may be titrated at pH = 1-2 and zinc or
lead after adjustment to pH = 5 by addition of
hexamine.
 The indicator solution is prepared by dissolving 0.5 g
of Xylenol Orange in 100 ml of water.

45. CALMAGALITE

46. CONTINUED….
 This indicator, 1-(1-hydroxyl-4-methyl-2-phenylazo)-2-
napthol-4-sulphonic acid, has the same colour change
as Solochrome.
 It may be substituted for Solochrome Black without
change in the experimental procedures for the
titration of calcium Plus magnesium.
 Calmagite functions as an acid-base indicator.
 The hydrogen of the sulphonic acid group plays no
part in the functioning of the dye as a metal ion
indicator.

47. CONTINUED…
 The acid properties of the hydroxyl groups are expressed by
pKl = 8.14 and pK2 = 12.35.
 The blue colour of Calmagite at pH= 10 is changed to red by
the addition of magnesium ions.


 The pH = 10 is attained by the use of an aqueous ammonia-
ammonium chloride buffer mixture.
 The indicator solution is prepared by dissolving 0.05 g of
Calmagite in 100 ml of water. It is stable for at least 12
months when stored in a polythene bottle.

48. CLEAR BLUE RED

49. MORDANT BLACK II

50. CONTINUED….
 Indicator is blue in colour at pH 10.
 On complexation with metal ions pink colour is formed.
 Below pH 6.3 and above 11.5 it is reddish in colour hence it
is used at pH 10.
 It is used in the estimation of metal ions like Calcium,
Magnesium, Zinc, Cadmium, Lead and Mercury.
 It can’t be used with oxidizing ions like ferric, cerric or with
reducing ions like stannous and titannous.
 It can’t be used with ions like copper, aluminum, cobalt,
silver as these ions form more stable complexes with
indicator than with chelating agent.

52. 52
APPLICATIONS
Determination of water hardness
Determination of metal ions
Complexometric titrations are widely
used in the medical industry.
The traces of zinc in water can be
determined with complexometric titration

53. •'Soft' water has a hardness < 60 mg CaCO3/L, and can lead to corrosion of
pipes. 'Hard' water has a hardness > 270 mg CaCO3/L, and can lead to
deposition of Ca and Mg minerals (limescale in water heaters, pipes etc.). Ca2+
also precipitates soap: Ca2+ + 2 RCOO- → Ca(OOCR)2 Temporary hardness is
caused by Ca(HCO3)2 and can be removed by boiling:
• CaCO3 ↓ + CO2 ↑ + H2O Ca(HCO3)2 CaCO3
precipitates and the water left is now softer, but CaCO3 has to be removed or
(some of) it will dissolve again.
• Permanent hardness is mainly caused by Ca2+/Mg2+ sulfates and chlorides –
cannot be removed by boiling, but can be removed by ion exchange
(replacement of Ca2+/Mg2+ by Na+). Determination of the hardness of water
involves direct titration with EDTA of Ca2+ and Mg2+ (together) at pH 10
(NH3 buffer). Fe3+ may have to be reduced with ascorbic acid to Fe2+ first, and
then Fe2+ is masked with CN- (which also masks several other minor metal
ions).

54. REFERENCES
 Pharmaceutical chemistry theory and application
.Leslie G. Chatten.volume 1
 Practical pharmaceutical chemistry by
A.H.Beckett,J.B.Stenlake
 Practical pharmaceutical chemistry –I,
A.V.Kasture,S.G.Wododkhar,S.B.Gokhale
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