Precipitation Titration
Presented by : Kaushik A. Bhakhar
(Pharmaceutical chemistry)
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The reaction involved in titrimetric analysis must occur quantitatively and must
proceed completely to form the product. Such a reaction may be of
1. Acid-base titration
2. Non-aqueous titration
3. Precipitation or argentometric titrations
4. Complexometric titration
5. Redox titration
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Must know about…
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Titrate : the substance to be analyzed
Titrant: the reagent of known concentration which is added to the solution
Titration curve: plot of solution pH vs volume of titrant
Titration: the process of finding out the volume of the titrant required to react
completely with a known volume of solution
Indicators: agents used to determine the end-point of titration
End-point: the system change the color of indicator by adding titrant to titrate
Equivalence point: a point in a titration when the moles per titrant equal the
substance being titrated
Standard solution: it is a solution of definite concentration or known strength
Standardization: it is a process where by the concentration of a solution is
determined by the known concentration of solution.
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• Precipitation Titration:
• The titration is based on the insoluble precipitate formation when the two reacting
substance are bought into contact are called as precipitation titration.
• A special type of titrimetric procedure involves the formation of precipitates
during the course of titration.
• The titrant react with the analyte forming an insoluble material and the titrations
continue till the very last amount of analyte is consumed.
• The first drop of titrant in excess will react with an indication resulting in a color
change and announcing the termination of the titration.
• Precipitations titration is commonly used to determine halide ions or silver
concentration using NaCl solution.
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The requirements for a reaction to be useful in titrametric analysis are :
1. The precipitate must be practically insoluble
2. The precipitation reaction should be rapid and quantitative
3. The titration results should not be hampered by adsorption (co-precipitation)
effects.
4. It must be possible to detect the equivalence point during the titration
5. Method based on precipitation of insoluble silver known as argentometry
6. Halogen can be determined by precipitation as sparingly soluble mercurous
salts HgCl2 & HgI2 is called as mercurometry.
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THEORY OF PRECIPITATION
The solution process and solubility product :
Solubility, which is dependent on the solvent and temperature, is the concentration
of dissolved solute in mole per litre when the solution is in equilibrium with a solid
state.
In the solid state the solute molecules occupy space in a fixed repeating pattern to
form what is called a crystal of solid. This repeating pattern depends upon the
molecular structure of the compound.
The solute molecules are held together in that pattern by intermolecular forces of
attraction.
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THEORY OF PRECIPITATION
Now in order to dissolve a solid, these forces of attraction must be overcome so that
solute-solute attraction is replaced by solute-solvent attraction.
The solvent should compete with crystal forces and overcome them, which often
means that the solvent environment must be similar to that provided by the crystal
structure. This is the basis for the simple rule "like dissolves like“.
During precipitation, however, the opposite condition is aspired for, where the
intermolecular forces between the molecules of product are high and solute-solute
forces replace the solute-solvent force.
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SOLUBILITY PRODUCT AND PRECIPITATION
Consider an aqueous solution of a slightly soluble salt BA in equilibrium with
excess of the solid at constant temperature. The equilibrium can be represented by:
where BA(s) represents the solid phase. In dilute aqueous solutions essentially no
undissociated BA will be present in the solution.
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SOLUBILITY PRODUCT AND PRECIPITATION
• The overall equilibrium is,
• Where B+ and A- represent ionic concentration
• The concentration of BA is constant in solid hence
Solubility constant product
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SOLUBILITY PRODUCT AND PRECIPITATION
• When the ionic product exceed the solubility product the solution is
saturated and precipitation occur
• When the ionic product is less than the solubility product the solution
is unsaturated.
• ex. In quantitative analysis excess precipitating agent is always
employed to ensure complete precipitating
• If the little excess if H2SO4 is employed, the ionic product far exceeds
the solubility product and there is complete precipitation.
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SOLUBILITY PRODUCT AND PRECIPITATION
• Oxalic acid cause complete precipitation of calcium-oxalate from
solution of calcium acetate but not from calcium chloride and calcium
nitrate
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SOLUBILITY PRODUCT AND PRECIPITATION
• Acetic acid is weak acid than oxalic acid thus it does not suppress the
dissociated oxalic acid. The concentration of oxalate ion is sufficient
to keep ionic product greater than solubility product of calcium
oxalate.
• In case of CaCl2 HCl is formed which is strong acid and highly
dissociated. It suppress the dissociation of oxalic acid by Common ion
effect.
• The oxalate ion concentration falls below the value required the
solubility product of calcium oxalate
• The precipitation is therefore incomplete that’s why Calcium oxalate
dissolve in HCl but not in Oxalic acid
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In a more complex case
Solubility product is of importance as it permits the calculation of one
of the ion concentrations if the other is known.
A substance precipitates out when the product of the ionic
concentrations exceeds the Ksp value, i. e. in equation (i) solid BA
will precipitate out when the product of [B+] and [A-] exceed Ksp
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Calculation of solubility and solubility product:
The solubility product of a slightly soluble electrolyte can be calculated if its
solubility is known and vice-versa.
(A) Calculation of solubility product :
1. Write the equation for the dissociation of the electrolyte, the solubility
product of which is to be calculated.
2. Express the solubility product (Ksp) of the electrolyte as the product of the
concentration of its ions.
3. Calculate the molar solubility of the substance i.e. the solubility in moles per
litre.
4. Calculate the concentration of each of the ions at equilibrium, using
dissociation equation and molar solubility found above.
5. Substitute the values for the concentration of the ions found in the expression
for solubility product and do the necessary calculations.
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Ex. 1. Calculate the solubility product of MgCO3 if 1 litre of its saturated solution
contains 0.533 g of MgCO3 at 20°C.
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The molar solubility of AgBr is 5.71×10-3 M. Calculate its solubility
product constant (Ksp).
1)When AgBr dissolves, it dissociates like this:
AgBr(s) ⇌ Ag+(aq) + Br¯(aq)
2) The Ksp expression is:
Ksp = [Ag+] [Br¯]
3) There is a 1:1 molar ratio between the AgBr that dissolves and Ag+ that is in solution.
In like manner, there is a 1:1 molar ratio between dissolved AgBr and Br¯ in solution.
This means that, when 5.71 x 10-7 mole per liter of AgBr dissolves, it produces 5.71 x 10-7 mole per liter
of Ag+ and 5.71 x 10-7 mole per liter of Br¯ in solution.
4) Putting the values into the Ksp expression, we obtain:
Ksp = (5.71 x 10-7) (5.71 x 10-7) = 3.26 x 10-13
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Determine the Ksp of calcium fluoride (CaF2), given that its molar
solubility is 2.14 x 10-4 moles per liter.
1) When CaF2 dissolves, it dissociates like this:
CaF2(s) ⇌ Ca2+(aq) + 2F¯(aq)
2) The Ksp expression is:
Ksp = [Ca2+] [F-]2
3) There is a 1:2 molar ratio between CaF2 and F¯. This means that, when 2.14 x 10-4 mole
per liter of CaF2 dissolves, it produces 2.14 x 10-4 mole per liter of Ca2+ and it produces
4.28 x 10¯4 mole per liter of F¯ in solution.
4) Putting the values into the Ksp expression, we obtain:
Ksp = (2.14 x 10-4) (4.28 x 10-8)
= 9.15 x 10-12
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Factors affecting solubility:
1. Common Ion Effect:
The solubility of any slightly soluble salt can be decreased by adding an excess of
either of its ions.
e.g. The dissociation of a slightly soluble salt BA is
This is the equilibrium condition. If, however, an excess of either B+ or A- are added
in the form of another salt (whose solubility) is greater than that of BA then the
product of ionic concentrations [B+] [A-] will exceed the solubility product and
hence BA will precipitate. The common ion effect provides a valuable method for
controlling the concentration of the ions furnished by a weak electrolyte.
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2. Effect of pH on solubility :
The solubility of a salt will be increased by decrease in pH, if the anion of the salt is
a conjugate base of a weak acid.
e.g. consider the slightly soluble salt BA, the anion of which (A-), is the conjugate
base of a weak acid HA.
Here, there will be two equilibriums in operation
• The A- from (1) will shift the equilibrium (2) towards the left while equilibrium
(1) will itself be shifted to the right. Hence, solubility of BA is increased with
increase in H+ or decrease in pH.
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The equilibrium expressions are
Molar solubility of BA is equal to [B+] which is equal to the total concentration of
A-, i.e. the A- dissolved from BA and that which is present in HA.
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3. Effect of temperature on solubility :
The solubility of the precipitate encountered in quantitative analysis increases with
the rise in temperature. With some substances the influence of temperature is small,
but with others it is quite appreciable. Thus, the solubility of AgCl at 10° and 100°
is 1.72 and 21.1 mg/litre, while that BaSO4 is 2.2 and 3.9 mg/litre respectively.
In many instances the common ion reduces the solubility to so small a value that
the temperature effect which is otherwise appreciable, becomes very small.
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4. Effect of the solvent upon the solubility:
The solubility of most organic compounds is reduced by the addition of organic
solvents such as methyl, ethyl and n-propyl alcohols etc.,
e.g. the addition of about 20% by volume of ethanol renders the solubility of lead
sulfate practically negligible, thus permitting quantitative separation.
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FRACTIONAL PRECIPITATION
We shall study the situation which arises when a precipitating reagent is added to a
solution containing two anions, both of which form slightly soluble salts with the
same cation
e.g. when silver nitrate solution is added to a solution containing both chloride and
iodide. The question which arises here is, which salt will precipitate first and how
completely will the first salt be precipitated before the second ions begins to react
with the reagent.
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The solubility products of silver chloride and silver iodide are 1.2 x 10-10 and
1.7 x 10-16 respectively.
It is evident that the solubility product of silver iodide being less will be exceeded
first and hence, Agl will be precipitated first
[Ag+]exceeds the value
before it exceeds
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AgCl will precipitate when the latter value is exceeded. After this both ions will be
precipitated simultaneously. The Ag+ ions will then be in equilibrium with both the
salts.
So, when the concentration of iodide ions is about one millionth part of the chloride
ion concentration, silver chloride will be precipitated.
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Thus, an almost complete separation is possible theoretically. This complete
separation is possible in practice if the point of complete precipitation of iodide is
detected. It can be possible either by use of adsorption indicator or by using
potentiometric titrations.
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Similarly, the fractional precipitation can be explained for a mixture of iodide and
bromide
When Ag+ will be equilibrium with both the salts.
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Precipitation of silver bromide will occur when concentration of bromide ion is
2 x 103 times that of iodide concentration. In this case, separation is not that
complete as in above case. It can be affected with accuracy by use of adsorption
indicator.
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Types
1. Mohr’s method
2. Volhard’s Method
3. Fajan’s method
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Mohr’s Method
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Mohr’s Method
Mohr’s method is a titration method is used to determined the concentration
of a solution an unknown concentration of a salt. It involving the solution with a
standard solution of AgNO3 to determine the concentration of the cation in the salt
Principle:
Mohr's method is based on the formation of a complex ion between the cation of the salt
and silver ions (Ag+). When the silver nitrate solution is added to the salt solution, the cation of
the salt reacts with the silver ions to form a complex ion. The endpoint of the titration is reached
when the excess silver ions are consumed and the concentration of the cation in the salt solution
can be determined
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Mohr’s Method
Karl Friedrich Mohr
1879, Bonn, Germany,
Field of pharmaceutical
and
chemical research;
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Mohr’s Method
Karl Friedrich Mohr
(1806-1879)
This method utilizes chromate as an indicator. Chromate forms a precipitate
with Ag+ but this precipitate has a greater solubility than that of AgCl, for
example. Therefore, AgCl is formed first and after all Cl- is consumed, the
first drop of Ag+ in excess will react with the chromate indicator giving a
reddish precipitate.
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The Mohr method uses chromate ions as an indicator in the titration of chloride ions
with a silver nitrate standard solution. After all the chloride has been precipitated as
white silver chloride, the first excess of titrant results in the formation of a silver
chromate precipitate, which signals the end point (1).
The reactions are:
By knowing the stoichiometry and moles consumed at the end point, the amount of
chloride is an unknown sample can be determined.
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Preparation of 5% K2CrO4 (indicator):
1.0 g of K2CrO4 was dissolved in 20 mL of distilled water.
Preparation of standard AgNO3 solution:
9.0 g of AgNO3 was weighed out, transferred to a 500 mL volumetric flask and
made up to volume with distilled water. The resulting solution was approximately
0.1 M. This solution was standardized against NaCl.
Reagent-grade NaCl was dried overnight and cooled to room temperature. 0.2500 g
portions of NaCl were weighed into Erlenmeyer flasks and dissolved in about
100 mL of distilled water. In order to adjust the pH of the solutions, small quantities
of NaHCO3 were added until effervescence ceased. About 2 mL of K2CrO4 was
added and the solution was titrated to the first permanent appearance of red
Ag2Cr2O4.
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Determination of Cl- in solid/solution sample:
1. Addition of silver nitrate (AgNO3) to the analyte containing chloride ions, lead
to formation of AgCl precipitate
2. Precipitate formation continue until all chloride ions have been completely
reacted with silver ions.
3. Now, further addition of AgNO3 to the analyte cause a reaction between AgNO3
& K2CrO4 leading to formation of silver chromate precipitate (Ag2CrO4)
4. Formation of Ag2CrO4 visualized by appearance of brick red colored
precipitates, that’s is the end-point of titration
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The titration was carried out at a pH between 7 and 10 because chromate ion is the
conjugate base of the weak chromic acid.
Therefore, when the pH is lower than 7, chromate ion is protonated and the chromic
acid form predominates in the solution. Consequently, in more acidic solutions the
chromate ion concentration is too low to produce the precipitate at the equivalence
point.
If the pH is above 10, brownish silver hydroxide forms and masks the end point. A
suitable pH was achieved by saturating the analyte solution with sodium hydrogen
carbonate.
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Applications:
1. It is used in determination of % of NaCl
2. In determination of NaCl and dextrose injection
3. It is used to estimate chloride ions in intraperitoneal dialysis fluid
4. In the determination of chromate ions & sodium nitroprusside
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Limitations:
1. Mohr’s method for precipitation titration is suitable in the PH range 6.5 to 10
 In highly acidic solution solubility of silver chromate increases leading to formation of
weak acid chromic acid
 So, end-point too late, to render this pure CaCO3, NaHCO3 & borax are added in
excess amount
 In strong alkaline solutions precipitates of silver hydroxide is formed
 Which disturb the continuous of the titration, to render this add dilute HNO3 or
ethanolic acid followed by addition of excess of CaCO3
2. Mohr’s method is not applicable for Iodine and thiocyanate titration due to adsorption of
chromate ion on AgI and AgSCN
3. During determination of BaCl2 , BaBr2 ,BaSO4 get precipitates first instead of Ag2CrO4
4. for Mohr's method room temp. is important as temp. increases solubility of Ag2CrO4 is also
increased
5. For titration of ammonium salt PH of solution should be maintained below 7
6. only applicable to salts that form complex ions with silver ions (Ag+).
7. The endpoint of the titration may be difficult to determine accurately, as the color
change of the indicator may be subtle.
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Volhard’s method
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Volhard Method (Formation of soluble
colored compound)
Jacob Volhard
(1834-1910)
• Volhard’s method was invented by JacobVolhard in 1874
• Estimation of halides and silver ions present in an acidic solution can be
achieved by titrating against standard solution
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• Principle: this method involved back titration
• Where excess of silver nitrate is back titrated using ammonium or potassium
thiocyanate as titrant
• Initially a precipitate of silver thiocyanate get formed
Once all silver ions have been reacted completely further addition of thiocyanate
solution results in formation of reddish brown complex due to reaction between
ferric ion and thiocyanate ion
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Preparation and standardization of 0.1M ammonium thiocyanate solution
Preparation :
• weight about 7.162 gm of NH4SCN followed by addition of small amount of
H2O and dissolved it
• Resultant solution made up to 1000 ml by water
• Standardization :
• Accurately measure 30 ml 0.1 M AgNO3 solution
• Dilute with 50 ml water followed by addition of 2 ml HNO3 and 2 ml ferric
ammonium sulphate solution
• Titrate the solution against 0.1 M NH4SCN solution titrate until color change
to reddish brown
• Each ml of 0.1 M AgNO3 = 0.007 g of NH4SCN
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• Procedure : to solution containing Cl- ions excess quantity of standard AgNO3 is
added which leads to formation of AgCl
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• Unreacted Ag+ ions (excess) back titrated with standard thiocyanate solution during
the white precipitate of silver thiocyanate is formed.
•
• When all silver ions reacted completely, further addition of thiocyanate solution
result formation of reddish brown colored precipitates
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2. Formation of a soluble coloured compound:
In this case, after the end point, the excess of precipitating reagent added reacts
with an indicator to form a soluble coloured complex. This procedure is
exemplified by the method of Volhard for the titration of silver in the presence of
free nitric acid with standard potassium or ammonium thiocyanate solution
using Fe+3 as indicator.
e.g. Determination of chloride ions. An excess of standard AgNO3 is added to the
chloride solution to be determined.
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1. The silver chloride precipitate is boiled for a few minutes to remove the adsobed
silver ions. Then the silver chloride precipitate is removed by filteration and the
cold filtrate is back titrated.
2. After addition of silver nitrate, potassium nitrate is added as a coagulant, the
suspension is boiled for 3 minutes, cooled and then titrated immediately.
Desorption of silver ions occurs and re-adsorption is prevented by the presence
of potassium nitrate.
3. The silver chloride particles are coated by an immiscible liquid e.g.
nitrobenzene and hence protected from reaction with thiocyanate.
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Applications:
1. It is used in determining the percentage of halogen like chloride, bromide and Iodide
2. It is used for determination of aminophylline tablets and injection
3. Used in determination of chlorbutanol, ethionamide, NaCl and lindane
4. It is used for determination or analysis of drugs like dimenhydrinate, chlorophetanol,
phenyl mercuric acetate and ox-phenonium bromide
5. Used for the analysis of NaCl hypertonic injection
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Limitation :
1. Performing Volhard’s titration above 20 oC results in fading of colour of the complex,
so should be perform below 20 oC
2. If nitric acid is contaminated with nitrous acid red colour complex formed due to
reaction with thiocyanic acid
3. During determination of iodine indicator should not be added until all iodine get ppt.
because iodine may undergo oxidation by Fe3+ ion to iodine
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Fajan's Method :
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3. Use of Adsorption Indicators or Fajan's Method :
K. Fajan, through his studies on nature of adsorption, introduced a useful type of
indicator for precipitation titrations. Such indicators are adsorbed on the surface of
the precipitate at the equivalence point and this adsorption is accompanied by a
colour change. These indicators are either acid dyes e.g. fluorescein, eosin etc. or
basic dyes e.g. rhodamine series.
• This method was introduced by K. fajan in 1923-24
• Based on adsorption property of a precipitates
• It involves direct titration of chloride with silver ions of AgNO3 using suitable
indicators
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Fajan’s Method
Kazimierz Fajans
(1887-1975)
Fluorescein and its derivatives are adsorbed to the surface of colloidal AgCl.
After all chloride is used, the first drop of Ag+ will react with fluorescein (FI-)
forming a reddish color.
Ag+ + FI-  AgF
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• Principle:
• The precipitate formed in titration has adsorption precipitates
• Precipitate initially adsorb its own ion which are in excess
• At the end point formed precipitate adsorbs oppositely charged of the indicator and
change color
• The indicator used in this method is called as adsorption indicator
• Most commonly used indicators in this method are fluorescene, dichlorofluorescene,
eosin, tartrazine, phenosafranine, diphenyl carbazone, rhodine
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• The various adsorption steps taking place in fajan’s method
The property of a colloidal precipitate to adsorb its own ions which are in excess, is
made use of in the case.
When a sodium chloride solution is titrated with silver nitrate the silver chloride
precipitate will adsorb chloride ions which are initially in excess. Thus the chloride
ions form the primary adsorbed layer, which in turn will hold the secondary
adsorbed layer of oppositely charged Na+ ions.
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Immediately after the equivalence point, Ag+ ions are in excess and hence silver
chloride ions now adsorb Ag+ ions as primary adsorbed layer and NO3- as secondary
adsorbed layer.
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Now, if the sodium salt of fluorescein is also present in the solution then the
negatively charged fluorescein ions would be adsorbed instead of NO3- as secondary
adsorbed layers and this adsorption occurs along with a change to pink colour due to
formation of a pink coloured complex of Ag+ and modified fluorescein ions.
An alternative view is that during the adsorption of fluorescein ion a rearrangement
of the structure of the ion occurs with the formation of a coloured substance.
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The conditions which govern the choice of an adsorption indicator are:
1. The indicator ion should have a charge opposite to that of the ion of the
precipitating agent.
2. Multivalent ions and other factors which have a coagulating effect should be
avoided as a colloidal state of the precipitate is desired.
3. The solution should be concentrated enough to give a sharp colour change.
4. The indicator should be secondarily adsorbed only after the equivalence point.
5. Precipitate particles should be of colloidal dimension to maximize quantity of
indicator adsorbed on these particles.
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Adsorption indicator
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Silver halides are sensitised to the action of light by the layer of adsorbed dye stuff
such as fluorescein and hence precipitation titrations of halides using argentometry
and adsorption indicators should be carried out with minimum exposure to sunlight.
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Comparison of Argentometric titration
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Precipitation titrations are useful in assay of various pharmaceuticals. The
following is the list of some drugs assayed by this technique
Aminophylline (for theophylline)
Aminophylline injection
Aminophylline tablets
Phenyl mercuric acetate
Sodium chloride
sodium chloride injection
Sodium chloride hypertonic injection
Sodium chloride and Dextrose injection
Volhard's Method
Volhard's Method
Volhard's Method
Volhard's Method
Volhard's Method
Volhard's Method
Volhard's Method
Mohrs method
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Self Study
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4. Turbidity Method: Gay Lussac's Method :
The fact that occurrence of turbidity accompanies precipitation reaction is made use
of in this method.
After the equivalence, the precipitation reaction ceases and addition of an extra drop
will not result in turbidity.
The procedure may be exemplified by the titration of silver nitrate with standard
sodium chloride in the presence of free nitric acid and a small quantity of pure
barium nitrate (to assist the process of coagulation).
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Weigh out accurately 0.4 g of silver nitrate in a well stoppered bottle. Add about 100
ml of water, a few drops of concentrated sulfuric acid and a small crystal of barium
nitrate.
Titrate with standard 0.1 M sodium chloride by adding 20 ml at once, stoppering the
bottle, and shaking vigorously until the precipitate of silver chloride has coagulated
and settled (a process aided by the barium sulfate) leaving a clear solution, Still the
silver ions are in excess.
Continue to add sodium chloride solution, 1 ml at a time, stoppering and shaking
after each addition till no turbidity is produced. Note the volume of sodium chloride
consumed, consider it as pilot reading.
Repeat the titration, adding 1 ml less than the pilot reading initially and continue
adding 0.02 ml after that. The end point can be determined within one drop.
Nephelo-turbidunetric methods can also be used.
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Repeat the titration, adding 1 ml less than the pilot reading initially and continue
adding 0.02 ml after that. The end point can be determined within one drop.
Nephelo-turbidunetric methods can also be used.
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TITRATION CURVES
Titration with precipitating agent are used for estimation of analytes, provided the
equilibria are rapid and method for detection of end point is available. Titration
curves based on p-values are useful for deducing the properties required of an
indicator and titration error that is likely to cause.
Consider the titration of chloride ions with standard silver nitrate solution. A
titration curve is plotted by pCl against the volume of silver nitrate.
Let us consider the changes in ionic concentrations which occur when 100 ml of 0.1
M NaCl is titrated with 0.1 M AgNO3. As we know the solubility product of AgCl
at laboratory temperature is 1.2 x 10-10.
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(A) Initially no Ag+ ions are present and the concentrtion of chloride ions will
be
(B) Upon addition of 50 ml of 0.1 M silver nitrate 50 ml of 0.1 M NaCl is
consumed and 50 ml of 0.1 M NaCl remains due to the reaction.
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AgCl immediately precipitates out till the equilibrium between
Ksp (AgCl) is reached
[Ag+] [Cl] and
where, Ksp is the solubility product, [Ag+] and[Cl-] are the concentrations
Ag+ and Cl- in solution.
Concentration of chloride now is
of
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A typical titration curve is prepared by plotting pCl against the volume of silver
nitrate. The titration curve can be illustrated as in figure
At the beginning of the titration on the pCl is 1. As the addition of titrant continues,
part of Cl- is removed from solution of AgCl precipitate and we can determine the
concentration of Cl- remaining. At the, equivalence point we have saturated
solution of AgCl and hence pCl = 4.96. Beyond the equivalence point, there is
excess of Ag+ and [Cl-] is determined from solubility product equation
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It will be seen by examining the chloride ion concentration that there is a
marked change in this value towards the equivalence point. Smaller the
solubility product of the precipitate, more pronounced the change will be.
Such precipitation titrations in which silver nitrate is used as a reagent are of
high importance and are termed argentometric processes.
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Thank you
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