Complexometric titration

Complexometric Titration
Metal ion (M+) + EDTA (HY) Metal-EDTA complex (MY) + H+
M+n + H2Y-2 MYn-4 + 2H+
Na4-xHxY can be used to represent any species of EDTA, with x designating the
number of acidic protons bonded to the EDTA molecule.
Due to low solubility of EDTA in water, its disodium dihydrate
EDTA salt i.e. Na2H2Y.2H2O is used.
Complex is formed by the reaction of metal ion (Mn+) with either an anion e.g.[Ag(CN)2]- or
neutral molecule e.g.[Ag(NH3)2]+ The metal ion is known as Central metal atom. The anion
or neutral molecule is known as Ligand (L)

• It is a volumetric analysis in which formation of a colored complex is used to indicate the
end-point of the titration.
• It is used for the determination of a mixture of different metal ions present in the solution.
• Complexometric titrationChelatometry is based on essentially Lewis acid-base reactions
in which electron pair is donated from one chemical to another.
• Disodium EDTA (Na2H2Y ⋅ 2H2O: better soluble in water) is most common chelating
agent used for Complexometric titration because it creates water soluble stable metal
complex in alkaline medium. (except the alkali metals which form too weak complexes)
• All the complexes have exact 1:1 stoichiometry (regardless of the charge of the cation).
They react stoichiometrically and can be used to quantitatively determine the metal ions
in the sample by titration.
• The titration needs to be carried out at alkaline medium to consume H+ which produced
during reaction, so that the reaction goes to the right and stability of the complex formed
(MY) is increased.
• The versatility, sensitivity, and general convenience of complexometric titrations are
dependent on the correct choice of indicators for endpoint detection.
• In deprotonated form of EDTA , six binding sites—four negatively charged carboxylate
groups and two tertiary amino groups—that can donate six pairs of electrons to a metal
ion to form coordinate covalent bond. [EDTA =Hexadentate ligand]
Some terms
• Ligand: Compound having at least one pair of unshared electrons available for bond formation
• Chelate: It is a complex formed between the ligand containing two or more donor groups and metal
to form ring structure. (heterocyclic rings or chelate rings). Chelates are usually insoluble in water
but soluble in organic solvent.
• Chelating agents: Organic molecules/ligands containing two or more donor groups which combine
with metal to form complex having ring structure. Most chelating agents consist of N or O.
• Sequestering agent: Ligands which form water soluble chelates e.g. EDTA.

• The resulting metal–ligand complex, in which EDTA forms a cage-like structure around
the metal ion, is very stable.
• Central metal atom= acts as Lewis acid (electron acceptor)
• Ligand= acts as Lewis base (electron donor)
• Coordinate bond (dative)= The bond formed between central metal atom (ion)
(acceptor) and the Ligand (donor)
• Dative bond is similar to covalent bond(formed of two electrons) Buti n dative
bond the electrons pair are donated from one atom to the other.
• The atom gives electron pair is known as donor, while the atom accept electron
pair is known as acceptor.
• Coordination number =The no. of coordinate bonds formed to a metal ion by its
ligands

• In general, the fully ionized form of EDTA will form a 1:1 complex
with all metal ions in solution. Thus, we have :
• Ligands are any combination of anions that can donate an electron
pair to form coordinate covalent bonds such as EDTA.
• Some examples of polyvalent metal ions are given below :
Bivalent Metal ions : Ca2+ , Mg2+ , Zn2+ ,
Trivalent Metal ions : Fe3+ , Al3+ , Cr3+ ,
Tetravalent Metal ions : Sn 4+ ,Ce4+ , Cr4+ , Pt4+ .
There are two main classes of ligands:
1. Unidentate ligands: Ligands that are bound to a metal ion at one
place are called as unidentate ligands. Example: NH3
2. Bidentate or multidentate ligands: These ligands contain more than
one group which is capable of binding with the metal ions.
Example: ethylene diamine, EDTA

CHELATING AGENTS
• A chelating agent is a chemical compound that can bind with a metal ion and
prevent that metal from undergoing other chemical reactions.
• These chelating agents are composed of atoms with lone pairs. These lone pairs can
be donated to a metal ion.
• The donation of one lone electron pair to the metal atom forms a coordinate
covalent bond. The number of coordinate covalent bonds present in a
coordination complex is called the coordination number.
• EDTA is a common example of a chelating agent. It is a multidentate ligand. This
means it can bind with the metal ion via several atoms by forming coordinate
covalent bonds. There are some other chelating agents that are Bidentate. They
form only two coordinate covalent bonds.
• Chelating agents are very useful in isolating heavy metals to remove them from
drinking water, to deactivate metal ions that can cause precipitation, to limit
available metal ion content, etc.

Effect of pH on complex formation:
• During the formation of stable, 1:1 water soluble, colorless complex there
will be the release of two H+ ions, as the number (n) of M-EDTA complex
increases in the solution, there will be the removal of 2n[H+] ions, acidity
increases in the solution. So we need to maintain the pH of the solution
with suitable basic buffer in order to have a M-EDTA complex. Hence
complexometric titrations are generally carried out in basic pH.
• Therefore, the stability of the metal complex is pH dependent. The lower
the pH of the solution the more hydrogen ions are available to compete
with the metal ion for the ligand, and the equilibrium shifts to the left,
causing a decrease in the stability of the complex. This equation
illustrates the competition between the metal ion and the hydrogen ion
for the ligand: M2+ + H2Y2- = MY2- + 2H+

STABILITY OF COMPLEXES
• Generally, the formation of a 1 : 1 chelate complex (MY n − 4 )may be
designated by the following equation :
where, M = Metal ion, and X = Chelating ion.
Hence, the stability constant, Kf , may be expressed as :
Kf =[MYn-4]/[Mn+][Y4-]
The equilibrium constant for the reaction of a metal with a ligand is called the
formation constant, Kf , or the stability constant.
The formation constant Kf = [MY n − 4 ] /[Mn+][Y4− ] describes the reaction between
Y4− and a metal ion.
Generally, complex ions with polydentate ligands have much higher formation
constants than those with monodentate ligands. This additional stability is
known as the chelation effect

Stability of Metal-Ligand Complexes
The stability of complexes is influenced by a number
of factors related to the ligand and metal ions.
1. Nature of the metal ion: Small ions with high
charges lead to stronger complexes.
2. Nature of the ligand: The ligands forming
chelates impart extra stability (chelon effect). For
example the complex of nickel with a
multidentate ligand is more stable than the one
formed with ammonia.
3. Basicity of the ligand: Greater basicity of the
ligand results in greater stability of the complex.

4. Size of chelate ring: The formation of five- or
six - membered rings provides the maximum
stability.
5. Number of metal chelate rings: The stability
of the complex is directly related to the number
of chelate rings formed between the ligand and
metal ion. Greater the number of such rings,
greater is the stability.
6. Steric effects: These also play an important
role in the stability of the complexes.

Requirement for Metal ion indicator
1. The colour must be sufficiently intense, so that a
minimum amount of indicator can be used.
2. The colour contrast between the indicator and
Metal-indicator complex should be readily
observable.
3. The Metal-indicator complex should possess
sufficient stability to ensure a sharp colour
change, however it should be less stable than
Metal-EDTA complex.
4. The change in equilibrium from metal-indicator
complex to the Metal-EDTA complex should be
sharp and rapid.

5. The colour reaction of the indicator should be
selective.
6. The indicator must be very sensitive to metal ions
so that the colour change occurs at near the
equivalence point.
7. The indicator must be stable in the titration
medium.
8. The indicator must be stable on storage also.
9. All the above requirements must be fulfilled in
the pH range in which the proposed titration is to
carried out.
10. It should be commercially available in adequate
purity.

Types of complexometric titration:
1. Direct titration
2. Back titration
3. Replacement titration
4. Alkalimetric titration
5. Indirect titration
1. Direct titration:
• It is the simplest and most convenient method in which the standard
solution of EDTA is slowly added to the metal ion solution till the end
point is reached.
• The solution containing metal ion is buffered to a desired pH and titrated
directly with the standard EDTA solution.
• A blank titration may be performed by omitting the sample to check the
presence of impurities in reagent.
• Ca 2+, Mg2+ , Zn2+ ions are determined by this method.
Metal ion + EDTA metal-EDTA complex

2. Back titration
• A measured amount of EDTA is added in excess to react with the metal
sample to be examined.
• The resulting solution will contain unreacted EDTA which is then back
titrated with another metal ion standard solution in the presence of
indicator.
• ZnCl2, ZnSO4, MgCl2, MgSO4 is used as standard metal ion solution.
• Al+3, Co+2, Pb+2, Mn+2, Hg+2, and Ni+2 can be determine by using Back
titration method.
• Metal ion 1 + EDTA in excess metal-EDTA complex 1 +
unreacted standard EDTA
standard metal ion 2
• metal- EDTA complex 2

3. Replacement titration
• When direct or back titrations do not give sharp endpoints or when there is no
suitable indicator for the analyte the metal may be determined by this
method.
• In this method , determination of metal ion is performed by displacing
magnesium or zinc ions from EDTA complex with equivalent amount of
calcium ion and the liberated magnesium or zinc ions are then titrated with
standard EDTA solution.
• To a calcium salt solution, ammonia-ammonium buffer is added. To this,
standard known volume of Mg-EDTA solution is added. In this reaction, stable
Ca-EDTA complex is form and Mg ions are liberated which may be titrated with
a standard EDTA solution.
• Cadmium , lead and mercury can also be determined by this method.
• Ca2+ + Mg-EDTA Ca- EDTA + Mg2+
• Stable complex
• Mg2+ + EDTA Mg-EDTA

4. Alkalimetric titration
• In an unbuffered solution, when a solution of disodium edetate (Na2H2Y)
is added to the metal ion solution, complex is formed with the liberation
of two equivalents of H+ ions.
• The free H+ ion then is titrated with a standard solution of alkali , NaOH
using acid-base indicator or potentiometric method of detecting end
point.
• Only metals forming EDTA complex of high stability constant can be
determined by this method.

5. Indirect titration
• This method is used to determine the ions such as Halides, phospates, and
sulphates that do not form complex with EDTA.
• In the determination of sulphate ion, SO4
-2 ion solution is treated with
excess of standard solution of Barium ion.
• The formed precipitate of BaSO4 is filtered off and unreacted Barium ions
present in filtrate is titrated with EDTA.
• In this way, indirect determination of the amount of sulphate ion present
in the sample solution is possible.
• Sulphate ion + excess barium ion barium sulphate ppt +
(Known amount) unreacted barium ion
barium-EDTA COMPLEX
Filter out ppt and
Titrate barium ion
With EDTA

End point detection in complexometric titration
• The equivalence point of a complexation titration occurs when we react
equivalent amounts of the titrand and titrant.
• A variety of methods are available for locating the end point, including
indicators and sensors that respond to a change in the solution conditions.
a. Metallochromic indicator
• Most indicators for complexation titrations are organic dyes—known as
metallochromic indicators—that form stable complexes with metal ions.
• The indicator, Inm– , is added to the titrand’s solution where it forms a stable
complex with the metal ion, MInn– . As we add EDTA it reacts first with free
metal ions, and then displaces the indicator from MInn– .
• If MInn– and Inm– have different colors, then the change in color signals the end
point.

• At the onset of the titration the buffered reaction medium contains the
metal-indicator complex (MI) that has formed, and excess metal ions.
• When EDTA, the titrant, is added to the system a competitive reaction
takes place between the free metal ions and the EDTA.
• Since the metal-indicator complex is weaker than the metal-EDTA chelate
(K, > KMI), the EDTA which is being added to the system during the course
of the titration is progressively chelating the free metal ions in solution at
the expense of the metal-indicator complex.
• Finally, at the equivalence point, EDTA removes the last traces of metal
from the indicator, and the indicator changes from its complexed
color to its metal-free color.

Eriochrome Black T (EBT)
• Indicator Eriochrome Black T becomes wine red in colour when bound
with metal ions while remaining blue in colour when free from metal ions
• EDTA is colourless whether it’s bound to metal ions or not
• Eriochrome Black T binds with metal ions loosely while EDTA binds with
metal ions strongly.
• When all metal ions are bound to EDTA, indicator EBT remains free in the
sample and the solution turns blue.

b. Spectrophotometric detection
• Limitation of metallochromic indicator is that if the solution is already
colored, it will be difficult to observe the indicator’s color change at end
point.
• For example, when titrating Cu2+ with EDTA, ammonia is used to adjust the
titrand’s pH. The intensely colored Cu(NH 3)4
2+ complex unclear the
indicator’s color, making an accurate determination of the end point
difficult.
• If at least one species in a complexation titration absorbs electromagnetic
radiation, then we can identify the end point by monitoring the titrand’s
absorbance at a carefully selected wavelength.

• For example, we can identify the end point for a titration of Cu2+ with EDTA
in the presence of NH3 by monitoring the titrand’s absorbance at a
wavelength of 745 nm, where the Cu(NH3 )4
2+ complex absorbs strongly.
• At the beginning of the titration the absorbance is at a maximum. As we
add EDTA, however, the reaction decreases the concentration of
Cu(NH3 )4
2+ and decreases the absorbance until we reach the equivalence
point ,owing to the color change from an intense deep blue of the
copper(ll)-tetraamine complex to the less intense light-blue color of the
copper(ll)-EDTA chelate. The end point is detected when the absorbance
does not change upon the addition of more titrant.
absorbance
Volume of EDTA

c. Amperometric titration
• An amperometric method is used for end-point detection of
complexometric titrations of metal ions in unbuffered solutions with
disodium ethylenediaminetetraacetate (Na2H2Y).
• A platinum disc microelectrode is used as a sensor for monitoring the
hydrogen ions delivered during the titration. Synthetic solutions
containing Ca2+ and Mg2+ ions, each alone or in mixture, are examined
• Before the equivalence point, the overall current is mainly related to the
concentration of hydrogen ions, which increases up to the point where all
Ca2+ ions have been titrated. After the equivalence point, the overall
current is instead related to the concentrations of both H+ and H2Y2-
• The intersection point of the two straight lines should be assumed as the
end-point of the titration.

d. Potentiometric method
• End point can be detected by measuring the change of potential at an
indicator electrode, immersed in the solution being analyzed, as the
titrant is being added.
• The end point of the titration is indicated by a large and rapid change of
potential.
• The equivalence point is indicated by a marked change in the titration
curve, when potential is plotted versus volume of titrant added to the
solution.
• The mercury electrode system can be used to detect the end point in the
titration of almost all metal ions. When the electrode is immersed in a
solution containing a mercury(II) salt the following equilibrium is
established:

• The potential of the electrode is given by Nernst equation: When the
potential of the system is measured, the standard potential E0 is known,
and the concentration of free mercury(II) ion is calculated by:
• E= E0-(0.0591/2) log[Hg2+ ]
• In the titration procedure using mercury electrode system, the electrode is
immersed in the solution containing the metal ion to be titrated and a
small quantity of a mercury(II)-EDTA chetate (usually a few drops of a 0.01
M solution),

• Mercury(II) is displaced from the mercury(II)-EDTA chelate by the metal ion
(that is, a divalent ion M2+ ), and the following exchange equilibrium results:
• From this equilibrium it can be seen that the concentration of M2+ , which is
in excess of the Hg-EDTA2- at the start of the titration, determines the
amount of free mercury(II) in solution ;thus the potential of the mercury
electrode depends upon the ratio [M2+ ]/[M-EDTA 2-].
• During the early part of the titration with EDTA, the ratio [M 2+ ]/[M-EDTA2-]
changes only slightly; however, in the vicinity of the end point this ratio
undergoes its largest changes and the potential of the mercury electrode
changes accordingly. As in all potentiometric titrations, the equivalence
point is indicated by a sharp inflection in the titration curve.

TITRATION SELECTIVITY AND MASKING REAGENTS
Chelons are very unselective titrants because stable chelates are formed with
a great number of metal ions. When several metal ions are present in a
solution, the titration data usually give the total metal content. Thus, it is
necessary to find practical methods which increase the selectivity of the
titrant.
Several methods have been designed to increase titrant selectively :
1. pH adjustment of the reaction medium,
2. the use of selective metal indicators,
3. the use of selective precipitants,
4. the addition of masking agents to the reaction medium.

1. pH adjustment of the reaction medium
• The formation of a metal chelate is dependent on the pH of the reaction
medium. In weakly acidic solution the chelates of many metals, such as
the alkaline earths, are completely dissociated, whereas chelates of
bismuth, iron(II) or chromium readily form.
• Thus, in acidic solution, bismuth can be effectively titrated with a chelon
in the presence of the alkaline earth metals. This method is based upon
the differences in stability of the chelates formed between the metal
ions and the chelon.
2. Use of selective metal indicators
• These indicators are metal complexing agents which preferentially react
with different metal ions under various conditions.
• Pyrocatechol Violet s a well-known example of selective metal indicator.
This substance preferentially reacts with bismuth in acidic solutions (pH
2 to 3). When bismuth and zinc are present in a mixture, the bismuth is
directly titrated with chelon in an acidic solution using Pyrocatechol
Violet as indicator.

• At the completion of this titration, an ammonia buffer is added to adjust
the pH to 10 and the system is again titrated to the Pyrocatechol Violet
end point with the chelon solution.
• The volume of titrant employed in the second part of the titration is used
to calculate the zinc concentration.
3. Use of selective precipitant
• Some of the more common precipitating agents which separate out
interfering metal ions, or groups of metal ions are oxalate for removal of
calcium ions, sulfide for the heavy metals, fluoride for aluminum, and
hydroxide for some transition metals, iron, and magnesium.
4. The addition of a masking agent to the reaction medium
• These substances are usually auxiliary chelating or complexing
compounds, which selectively react with the interfering metal ions
forming more stable complexes than does the chelon.

• Therefore, these interfering ions are said to be masked and are not titrated by
the chelon.
• Since masking a cation in solution is an equilibrium process, it is possible to
reverse this reaction by the addition of suitable demasking agents; thereby, a,
series of metal ions can be titrated in one solution.
• One example of the use of masking and demasking reagents in chelometry is
the analysis of a mixture of three metals: copper, cadmium, and calcium.
• Direct titration of the mixture with EDTA gives the sum of the three metals.
• Copper and cadmium may be masked with the addition of cyanide to the
solution, leaving only the calcium ion titratable.
• When formaldehyde or chloral hydrate is added to the cyanide-containing
mixture, only the cadmium is demasked, and the EDTA titrates the sum of the
calcium and cadmium.
• In this manner the concentrations of all three ions are determined by three
individual titrations.

• The reaction sequences involved in the various steps of this titration
procedure are represented as:
Step 1: All three metals are titrated:
Step 2: Only calcium titrated (Cu and Cd masked by cyanide):

Step 3 : Calcium and cadmium titrated (Cu masked):

Common Masking Agents
Masking agent Metals masked
Cyanide Nickel, cobalt, zinc, cadmium, copper,
iron(II)
Iodide Mercury
Ascorbic acid Iron(III), copper
Triethanotamine BAL (2,3-dimercapto-1 -
propan-l-ol)
Iron(IlI), aluminum, manganese, Tin,
antimony, cadmium, mercury, zinc ,
Tiron (disodium catcchol-3,5-disulfonate) Aluminum, iron(III)

PREPARATION OF BUFFER SOLUTIONS:
1. Ammonia Buffer
Dissolve 54 g of ammonium chloride in 100 ml of distilled water, add 350 ml of ammonium hydroxide
solution, and dilute to 1000 ml with distilled water.
2. Acetate Buffer
Dissolve 136 g of sodium acetate trihydrate in 500 ml of water; add 29 ml of glacial acetic acid and
sufficient water to make 1 liter.
Preparation of 0.1 M EDTA solution
To prepare a 0.1 M EDTA solution, dissolve 37.22 g of the commercially available disodium EDTA
dihydrate in water. When solution is complete, dilute to 1 liter.
Standardization of 0.1M EDTA solution using calcium carbonate
• Dissolve an accurately weighed quantity (200 to 400 mg) of calcium carbonate in 100 ml of water
with the aid of a minimum amount of hydrochloric acid
• Adjust the pH of the solution to about 12 with a freshly prepared 0.1 N sodium hydroxide solution.
• Add 40 mg of murexide powder and 3 ml of Naphthol Green B indicator.
• Titrate to a deep-blue color with EDTA solution.
• Calculate the molarity by the formula:
MEDTA= WCaCO3/ ( MWCaCO3 * VolEDTA )
where WCaCO3 is the weight of calcium carbonate sample, MWCaCO3the molecular weight of calcium
carbonate (100.1), and VolEDTA is the volume, of EDTA solution required for the titration.
• The EDTA solution should be standardized against the metal ion for which it is to be used as titrant
and under the conditions of the actual titration. For example, if zinc is to be determined, the EDTA
solution should be standardized by titrating a zinc solution prepared from pure zinc.

1. DETERMINATION OF ALUMINUM (Back Titration)
• Dissolve an accurately weighed sample of the aluminum-containing
substance, in water with the aid of a few drops of hydrochloric acid, if
necessary.
• Dilute the solution to 100 ml with water. To 25 ml of this solution, add 25
ml of the standardized 0.1 M EDTA solution and 25 ml of water.
• Heat the solution on a water bath for 30 min, add 10 ml of acetate buffer,
and titrate with lead nitrate or zinc sulfate solution using a few drops of
Xylenol Orange as indicator.
• At the end point, the indicator changes from a lemon-yellow color to red.
• The volume of 0.1 M EDTA solution consumed by the aluminum is found
by subtracting the volume of lead nitrate or zinc sulfate from the total
volume of EDTA added to the system.
Each milliliter of 0.1 M EDTA is equivalent to 2.698 mg of aluminum.

2. DETERMINATION OF CALCIUM
Calcium ions may be determined by four various titrimetric procedures
a. Direct Titration
• Dissolve an accurately weighted sample of the calcium salt, equivalent to
about 100 mg of calcium, in 50 ml of water with the aid of a few drops of
- hydrochloric acid, if necessary. When solution is complete, dilute the
sample with water to volume in a 100 ml volumetric flask and mix
thoroughly.
• Accurately pipette 50 ml of the sample-solution into a suitable titration
vessel containing 50 ml of water, adjust the pH of the solution to 12 with
a freshly prepared 1.0 N sodium hydroxide solution, add 300mg Hydroxy
Naphthol Blue indicator, and titrate rapidly with 0.1 M EDTA until the
color changes to a deep blue.
Each milliliter of the titrant is equivalent to 4.008 mg of calcium.

b. Direct Titration In the Presence of a Magnesium Salt
• Prepare the sample solution as described under the direct titration procedure.
• Pipette 50 ml of the solution into titration vessel containing 50 ml of water;
add by means of a pipette 2 ml of 0.1 M magnesium sulfate solution, 5 ml of
ammonia buffer, and a few drops of Eriochrome Black T indicator. Titrate with
0.1 M EDTA solution to a deep-blue color.
• The volume of 0.1 M EDTA consumed by the calcium is calculated by
substracting the volume of 0.1 M magnesium sulfate, in milliliters, from the
total volume of 0.1 M EDTA required to obtain the deep-blue color at the end
point.
c. BackTitration
• Prepare a sample solution as previously described.
• 50 ml of the solution into a suitable vessel, add by means of a pipette 50 ml of
0.1M EDTA solution, add 5 ml of ammonia buffer, and a few drops of
Eriochrorne Black T indicator. Titrate with 0.1 M magnesium sulfate solution
until the color changes from blue to wine-red.
• The difference in volume between the 0.1 M EDTA solution and the
magnesium sulfate solution is the volume of 0.1 M EDTA solution consumed by
the calcium.
Each milliliter of this solution is equivalent to 4.008 mg of calcium.

d. Replacement Titration
• A solution of Mg-EDTA, equivalent to 50 mG /mI of the chelate is prepared
by dissolving 19.90 g of MgSO4.7H2O and 30.1 g of Na2 –EDTA.2H2O in
water and diluting to 1 liter.
• Prepare the sample solution as previously described.
• Pipette 5O ml of the sample solution into-a titration vessel containing 50
ml of water; add 25 ml of Mg-EDTA solution, 5 ml of ammonia buffer, and
a few drops of Eriochrome Black T indicator. Titrate to a deep blue with 0.1
M EDTA solution.
Each milliliter of titrant is equivalent to 4.008 mg of calcium.

3. DETERMINATION OF MAGNESIUM
• Dissolve an accurately weighed sample, equivalent to about 75 mg of mangesium. in
sufficient water to make 100 ml. Pipette 50 ml of this solution into a suitable vessel
containing 50 ml of water; add 5 ml of ammonia buffer, and a few drops of Eriochrome
Black T indicator, Titrate to a deep blue color with 0.1 M EDTA solution.
Each milliliter of titrant is equivalenL to 2.432 mg of magnesium.
4. DETERMINATION OF LEAD
• Dissolve an accurately weighed sample, equivalent to about 250 mg of lead, in
sufficient water to make 100 ml. Pipette 50 ml of the sample solution into a suitable
vessel, add 5 ml of acetate buffer, a few drops of Xylenol Orange indicator, and titrate
with 0 I M EDTA solution to a purple-yellow end point.
Each milliliter of titrant is equivalent to 20.72 mg of lead.
This procedure is similar to the BP assay for lead acetate.
5. DETERMINATION OF ZINC
Dissolve an accurately weighed sample of zinc salt, equivalent to about 200 mg of zinc, in
sufficient water to make 100 ml. Pipette 50 ml of the sample solution into a vessel
containing 50 ml of water, and add 5 ml of acetate buffer and a few drops of Xylenol
Orange indicator. Titrate to a purple-yellow end point with 0.1 M ETA solution.
Each milliliter of the titrant is equivalent to 6.538 mg of zinc.

Assay of calcium gluconate
Preparation of 0.05 M EDTA
Dissolve 18.6 g of Disodium EDTA dihydrate in sufficient water to produce 1000 ml.
Standardization of 0.05 M EDTA
Weigh accurately about 0.1 g of calcium carbonate and transfer it to a conical flask .
Dissolve in 3 ml of HCl and 10 ml of purified water. Boil for 10 min, cool and dilute to 50
ml with water. Add few drops of Erichrome Black T indicator and titrate with prepared
EDTA till end point.
Note down volume of EDTA solution consumed during titration. Repeat procedure two
more times. Calculate the molarity by:
Molarity= (W CaCO3* 0.05)/ (VEDTA *5.0045)
Each ml of 0.05 M EDTA is equivalent to 5.0045 mg calcium carbonate
Assay of Calcium gluconate
• Weigh 0.5 g of calcium gluconate and dissolve it in 20 ml of hot water containing 2 ml of
hydrochloric acid. Dilute it to 100 ml of distilled water. From it, pipette out 50 ml of
solution and add 10 ml of ammonium buffer solution.
• Titrate the solution against 0.05M EDTA using mordant black II indicator till color
changes from pink to blue.
Each mL of 0.05 M disodium edetate is equivalent to 22.42 mg of Calcium gluconate.

Complexometric titration

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