1
MIDLANDS STATE UNIVERSITY
DEPARTMENT OF CHEMICAL TECHNOLOGY
ANALYTICAL TECHNIQUES PRACTICAL CT104
Compiled by Mambanda Isaac(Chief Technician) & Gonzo Muriel(Snr Lab Technician)
2
DETERMINATION OF CHLORIDE BY THE MOHR TITRATION
Titrimetric methods in which silver nitrate is used as the titrant are called ‘argentimetric ‘ methods .
Argentimetric methods are the most important of those based on the precipitation processes. They are
used for the determination of halides ,cyanide .thiocyanate and less frequently for arsenate , phosphate,
sulphide ,and certain other anions .There are tree general methods of argentimetry: Mohr’s method with
potassium chromate as indicator f the adsorption indictor method; and the Volhard’s method.
Ag+ ion reacts with Cl- ion according to the equation:
Ag+(aq) +
Cl(aq) 
AgCl (g)
The principle of Mohr’s method depends on the fractional precipitation of silver chloride and silver
chromate .Since silver chloride is considerably less soluble (AgCl ;K sp=1.88x10 -10 and
Ag2CrO4;Ksp=1.5x10-10) silver chromate does not precipitate until (for practical purposes)the solution is
free from Cl- ions.Red silver chromate then precipitates and imparts a faint red-orange color to the
precipitate of AgCl.
The amount of indicator is important. Silver chromate does not precipitate until its solubility product has
been exceeded .Accordingly , the point at which this occurs is influenced by the amount of chromate
present .If insufficient indicator is present , the end-point comes too late because an undue excess of Ag+
ions is necessary to exceed the solubility product .On the other end an excessive amount not only tends to
mask the end-point , but diminishes the solubility of silver chromate to such an extend that the end –point
appears too soon. However it can be shown by suitable calculation that for 0.1M solutions a fair amount of
latitude is permissible .Thus when 25ml of 0.1M NaCl are titrated with 0.1M AgNO3 over a range of
chromate ion between 1,4x10-1 and 1.4x10-3 M ,the error only varies from -0.015 to +0.015ml.Generally, 12ml of a5% solution for each 100ml of solution ([CrO 42-] =5x10-3) is used.
The p H of the solution should be carefully controlled .The second dissociation of chromic acid is
small and chromate ions react with hydrogen ions as follows:
CrO42- + H + <=> HCrO 4-
Hence the sensitivity falls with increasing acidity , eventually a point is reached where silver chromate
does not precipitate at all.
In alkaline solution the hydroxyl ion concentration may be sufficient to cause precipitation of silver
hydroxide before silver chromate. A pH range of 6.5 to 10.5 is recommended. Acid solutions can be
neutralized with CaCO3 ,NaHCO3 , MgO or borax. Alkaline solutions may be made slightly acidic with
acetic acid ; a slight excess of CaCO3 is then added.The solubity of silver chromate increases with
increasing temperatures , hence the sensitivity of the method decreases.Working at room temperature is
recommended.
Procedure
Compiled by Mambanda Isaac(Chief Technician) & Gonzo Muriel(Snr Lab Technician)
3
Standardization of AgNO3 solution
Weigh accurately about 1.5g of reagent grade sodium chloride which has been dried at 120 0C FOR 1-2
hours, dissolve in distilled water and dilute to the mark in a 250 ml volumetric flask, mix well. Calculate
the molarity of the solution.
Pipette 25ml portions of the solution into a 250ml conical flask , add 1ml of indicator and titrate slowly ,
whilst swirling the to the first color change from the yellow color of the solution .When the end-point
appears , shake well , if the color reverts continue the titration until a permanent change s obtained .
Calculate the molarity of the AgNO 3 solution.
Determination of chloride
Weigh out 1.5-2.0 g of the sample , dissolve in distilled water and dilute to 250ml in a graduated flask. Add
1ml of indicator and proceed as above. Calculate the percentage of chloride in the sample.
Compiled by Mambanda Isaac(Chief Technician) & Gonzo Muriel(Snr Lab Technician)
4
DETERMINATION OF CHLORIDE BY THE VOLHARD METHOD
Ag+ ion reacts with Cl- ion according to the equation ;
Ag+(aq)
+

Cl-(aq)
AgCl(g)
In the volhard method an excess of standard silver nitrate is added to the halide solution and the unreacted
portion is back titrated with standard thiocyanate solution; ferric alum is used as indicator.The latter forms
the ferrithiocyanate complex at the end point.This method is not only sensitive, but is also of wide
applicability because it can be operated in a fairly strong acid solutioon. The method is for the direct
titration of silver.
Although Volhard mantained that the back titration could be carried out in the presence of the precipitated
silver chloride , other investigators indicated that thiocyanate reacted with silver chloride.
AgCl
+
CNS-  AgCNS + Cl-
Indeed this is to be expected from the wide difference in solubilities of the silver salt:
CN=
[CNS-]
KSAgCl
KSAgCNS
=
1.1x10-10
7.0x10-10
The minute amount of thiocyanate in excess, which would otherwise mark the end-point, reacts with the
precipitate as above , the reaction continuing with further additions of thiocyanate until the concentration of
chloride ions liberated during the reaction is 157 times the concentration of thiocyanate ions. Hence , the
color fades rapidly after each addition and the end point occurs some millimeters too late .
To overcome this effect some workers have made an extensive study of the titration , and reccomended
filtering off the silver chloride before back titrating . Immediate filtration can cause errors owing to
adsorption of excess of silver ions , hence it is necessary to shake the mixture for about five minutes before
filtering or to boil for a short time to coagulate the silver chloride.
Several workers have recommended the addition of organic solvents to minimize the reactionn between
thiocyanate and silver chloride and thus avoid filtration . The best solvent is nitrobenzene. The effect of the
solvent is obscure. It seems likely, however that it forms a water proof coating around the silver chloride
precipitate and thus prevents it’s dissolution. Adsorption of silver ions on the precipitate is also
inhibited.
Some workers ascribe the fading end-point obtained in the presence of silver chloride to the slow
reaction of adsorbed silver ions with thiocyanate. They recommend that after addition of silver nitrate
potassium nitrate also be added and the mixture boiled for a few minutes , cooled and then titrated
immeditely . Desorption of silver ions occurs , and. On cooling re-adsorption cannot occur in the presence
of potassium nitrate.
However it has been verified that no fading occurs in the presence of silver chloride when this procedure is
used . Nevertheless, it seems likely that the widely differing solubilities of silver chloride and thiocyanate
play at least a major part in the fading effect. It may be that boiling in the presence of potassium nitrate
causes the precipitate to age rapidly and dissolutoin of this aged precipitate is so slow that no fading effect
can be observed.
Compiled by Mambanda Isaac(Chief Technician) & Gonzo Muriel(Snr Lab Technician)
5
Because silver bromide and iodide are less soluble than silver thiocyanate , no reaction can take place
between the precipitate and excess of thiocyanate , hence it is necessary to add nitrobenzene or to filter
before the back-titration.
In the case of iodide, the indicator should not be added until access silver is present, otherwise oxidation of
iodide occurs according to the following equation:
2Fe3+
+
2I- 
2Fe2+ +
I2
OPTIMUM CONDITIONS FOR THE TITRATION
1. Indicator error with the method is negligible when 0,14M solutions are used, a pale orange color
being obtained with 0,01ml of 0,1M thiocyanate in a volume of 100ml.This is equivalent to a
thiocyanate concentration of 10-5 M. The end-point occurs too soon however, owing to adsorption
of silver ions by freshly precipitated silver thiocyanate. Accordingly, when the first perceptible
color is obtained, the solution should be shaken vigorously and the titration continued until a
permanent color is obtained.
2. Hydrogen Ion Concentration.
The titration is usually carried out at an acid concentration of about 0,5M although the range
generally specified is 0,2-1,6M.When arsenate or phosphate is present it better to use a
concentration of about 2M to prevent their precipitating.
3.
Temperature
Low results are obtained if the temperature is much above 25 oC, because nitric acid then bleaches
the ferric thiocyanate. If, therefore, the solution has been heated, care should be taken to ensure it
has been cooled below 25o C before the titration is carried out.
Standardisation of Potassium Thiocyanate
You are provided with a 0,1 M standard silver nitrate solution. Weigh about 3,0g of potassium
thiocyanate, transfer to a 500ml volumetric flask, dissolve in distilled water and dilute to the mark.
Pipette 10ml of 0,1M silver nitrate solution to a 250ml conical flask.Add 5ml of 6M HNO 3 and
1ml of Fe(III) indicator and titrate with thiocyanate solution. When the first change to orange-red
occurs, shake well and continue the titration carefully until a permanent color remains after
shaking . Calculate the molarity of the thiocyanate solution.
Procedure 1
Weigh out accurately 1.5-2.0g of the sample provided , dissolve in distilled water and dilute to
250ml in a graduated flask . Mix well . Pipette 25ml of the solution into a 250ml beaker , add
5ml of 6M HNO3 and then 2-6 ml excess of 0.1 AgNO3 solution. Stir well and filter into a 250ml
conical flask . Wash with 1:100 nitric acid .Add 1-2ml of indicator to the combined filtrate and
washings and titrate as above. Calculate the percentage of chloride in the sample.
Procedure 2
Proceed as above , adding 2-5 of 0.1 M silver nitrate solution in excess. Add ml of nitrobenzene
And 1ml of indicator and shake well to coagulate the precipitate .Titrate with 0.1M thiocyanate
solution until a permanent color appears which does not fade after 5 minutes.
Compiled by Mambanda Isaac(Chief Technician) & Gonzo Muriel(Snr Lab Technician)
6
IODIMETRY-DETERMINATION OF AVAILABLE CHLORINE IN
BLEACHING POWDER
Iodimetry can be divided into two parts: direct and indirect iodimetry. Direct iodimetry involves titration
with standard iodine solutions; sometimes it is necessary to add an excess of iodine and back- titrate with
thiosulphate or arsenic(III) solution .Indirect iodimetry which is of greater importance in analytical work ,
involves liberation of iodine from iodide by an oxidant and this iodine is then titrated with a standard
thiosulphate solution . This practical is an example of indirect iodimetry.
INDICATORS
One drop of 0.1N iodine solution imparts an intense yellow color to 100ml of solution , hence iodine , like
potassium permanganate, can serve as its own indicator, but is more usual to employ starch .In the presence
of iodide starch should be freshly prepared or contain a preservative ;mercury (II)iodide, thymol and
salicylic acid have been used , but it is necessary to ensure that the preservative does not interfere in the
particular the particular titration undertaken .Old starch solutions give reddish colored complexes rather
than blue complexes and sharp end-points cannot be obtained.
A solid 5% solution of starch in urea dissolves instantaneously in aqueous solutions and gives excellent
end-points; the solid preparation is almost indefinitely stable .The commercial product , Thyodene , is now
commonly used.
AVAILABLE CHLORINE IN BLEACHING POWDER
Bleaching powder has an indefinite composition approximately to the formula CaCl(OCl), and contains a
number of impurities, only part of the chlorine is available for oxidizing purposes, this being known as
available chlorine. It is determined by treating the bleaching powder with acetic acid and potassium iodide,
the free iodine so liberated being titrated with sodium thiosulphate solution:
CaCl(OCl)
+
2I- + 2H+  Ca2+ + 2Cl- +H2O +
I2
The liberated iodine is then titrated with standard thiosulphate solution :
2S2O32-
+ I2  S4O62- + 2I-
starch being used as indicator. Note that when titrating iodine with thiosulphate , starch should not be added
until immediately before the equivalence (as detected visually , by fading of the iodine color).This is
because some iodine tends to remain bound to starch particles after the equivalence point is reached.
Procedure
Standardization of Sodium Thiosulphate Solution
You will be provided with approximately 0.1M sodium thiosulphate solution. Weigh accurately about
0.89g of potassium iodate (RMM=214.00) previously dried at 1800C, dissolve in distilled water and dilute
to 250ml in a volumetric flask Calculate the molarity of the solution. Pipette 25ml portions into each of
250ml conical flasks and add 1g of potassium iodide ( a rough balance will do) and add 5ml of 4M H2SO4.
Titrate with thiosulphate solution to a faint straw color, add starch preparation and continue the titration to
end-point ,which is just the disappearance of the blue color to colorless.
Compiled by Mambanda Isaac(Chief Technician) & Gonzo Muriel(Snr Lab Technician)
7
Determination of available chlorine in bleaching powder
Weigh 2.5g of bleaching powder into a mortar , add a little distilled water and rub to a paste .Add more
distilled water , allow to settle and decant the liquid into a 250ml; volumetric flask Repeat until the whole
powder has been transferred .Dilute the flask to the mark with distilled water , shake well and pipette 50ml
into a conical .Add 2g of potassium iodide and 10ml glacial acetic acid .Titrate the liberated iodine with the
standardized thiosulphate solution . Repeat with further portions of the solution .Calculate the percentage of
available chlorine.
Compiled by Mambanda Isaac(Chief Technician) & Gonzo Muriel(Snr Lab Technician)
8
EDTA TITRATIONS
Ethylenedinitrilotetraacetic acid (EDTA:1,2 diaminoethanetetra-acetic acid) is one of a class of
aminopolycarboxylic acids which form soluable strong complexes with many metal ions. When a solution
containing asuitable divalent metal ion is titrated with EDTA , the reaction that occurs can be formulated:
M2+
H2Y2 -
+
MY2-
2H+
+
which is a special case of the general equation:
Mm+
+
HnYn-4
MYm-4
+
nH+
Where Y4- is the fully deprotonated anion of EDTA.
The free acid has the formula:
-
OOC-CH2
CH2-COOH
N-CH2-CH2-N
HOOC-CH2
CH2-COO-
It is a hexadentate ligand and one of its most interesting features is that it forms 1:1 complexes with metal
ions regardless of their charge. The first practical application of the EDTA titration was the determination
of water hardness (calcium and magnesium salts) but it can be used for many other ions as well. In this
practical EDTA titrations will be illustrated by using it to titrate zinc, calcium, magnesium and nickel.
The end-points of titrations with EDTA are usually detected by metallochromic indicators. These are
organic colourship matters that undergo a colour change when they form metal complexes. Generally
speaking, not only the dyestuff but also its metal complex are coloured, so that the end-point of the
titration is not characterized by e appearance or disappearance of a color , but rather a change in color. Two
of the most common metallochromic indicators are murexide and eriochrome black T (Erio T for short).
Compiled by Mambanda Isaac(Chief Technician) & Gonzo Muriel(Snr Lab Technician)
9
Murexide
Murexide , which is blue in alkaline solution , forms a pink complex with calcium ions.When a solution
containing calcium and murexide is titrated with EDTA, the original blue color of the indicator is restored
when all the calcium ions have been complexed .The anion of murexide has the following constitution.
HN
O
C
C O
C
HN C O
O C NH
N C
C O
O C NH
Several other metal ions such as cadmium , cobalt , copper , mercury , nickel and zinc give a color with
murexide. Magnesium is without effect ; hence , at a suitable pH , calcium can be titrated in the presence of
magnesium when murexide is used as indicator.
Eriochrome Black T
This indicator is blue between pH 6.3 and 11.3.The magnesium complex is wine-red ; when magnesium is
titrated with EDTA in the presence of this indicator between these pH limits , the latter reverts to its
original blue color when all the magnesium ions have been complexed.The optimum pH is 10.
The complex formed between calcium and this dye is too unstable for indicator purposes.When a solutuion
containing both calcium and magnesium os titrated with EDTA , calcium is preferentially complexed by
EDTA , but an end point is not obtained until all the magnesium has also been complexed. Hence the sum
of magnesium and calcium is obtained .Several other metals give similar color reactions and can be titrated
in the same way as magnesium, e.g. zinc ,cadmium ,lead and manganese.
Zinc
In this practical EDTA will be standardized with a zinc solution.This titration presents no difficulties and
involves titration in an ammoniacal solution with Erio T as indicator.
Magnesium
The indicator Erio T is used in this titration.The stability of the complexes of Mg with EDTA and with the
indicator are just high enough to permit an accurate titration and in point of the fact that color change at the
end-point (wine-red) is less sharp than in many other complexometric titrations. The titration must be
continued until the last suggestion of a reddish blue has gone.Interfering metals such as Cu,Ni,Co,Fe etc are
masked by KCN- sodium sulphide is equally effective as a precipitant.
Calcium
The titration of calcium with EDTA can be carried out in very dilute solutions with murexide as indicator.
However,the the color change from red to blue- violet which takes place place in strongly alkaline medium
(Ph 12) is not so sharp as in metallochromic indicators.
Nickel
Compiled by Mambanda Isaac(Chief Technician) & Gonzo Muriel(Snr Lab Technician)
10
Nickel can be titrated directly in strongly ammoniacal solution using murexide as indicator .With all nickel
titrations ,, it is necessary to titrate slowly near the end-point, as the rate of formation of the Ni-complex is
not very high.
Ni can also be determined by way of back-titrations, and this has the advantage that the lower rate of
reaction is now of no consequence.The back-titration can be carried out with either copper or zinc.
Procedure
Standardization of EDTA solution
Weigh accurately about 0.25g of zinc oxide and transfer into a conical flask .Add 0.5M HCl dropwise until
all the zinc oxide has dissolved .Cool and transfer to a 250ml volumetric flask with distilled water .Pipette
10ml portions into each of three 250ml conical flasks ;add concentrated ammonia solution in excess to
dissolve the precipitate formed .Titrate with the EDTA solution using Erio T indicator .Calculate the
molarity of the EDTA solution.
HARDNESS OF WATER
Hardness of water is due calcium and magnesium salts which are present as hydrogen carbonates
(temporary hardness), and sulphates, chlorides and nitrates (permanent hardness). These salts are leached
out by natural water, usually containing dissolved CO2, as it flows among rock strata underground. Hard
water has poor washing properties and when heated leaves deposits of solid CaCO 3. The determination of
water hardness is, therefore, important in many industrial processes where water is an important raw
material, so that where necessary, appropriate steps can be taken.
DETERMINATION OF Ca HARDNESS IN WATER
Pipette 100ml of bore-hole water into a 250ml conical flask, add 2ml of the ammonia buffer pH 10 and add
2 drops of freshly prepared 1% sodium sulphide solution. After 1 minute add a tip of a spatula Erio T
indicator and titrate with the standardized EDTA solution until the last traces of red have disappeared from
the solution.
Calculate the total hardness as ppm CaCO3.
DETERMINATION OF NICKEL
Pipette 100ml of the neutral sample solution into a 250ml conical flask. Add murexide indicator. Then add
10 ml of 1M NH4Cl. If the pH of the solution is below 7 the indicator will have an orange-yellow colour
(due to NiH4D+) and dilute ammonia must be added dropwise until the colour changes to yellow (due to
NiH2D-). Start the titration and continue until the end-point is near. If the colour reverts to orange because
of a fall in pH, add a few drops of ammonia and continue the titration. Just before the end-point make the
solution strongly ammonical by adding 10ml of concentrated ammonia and titrate to a brilliant colour
change from yellow to bluish-violet.
Calculate the concentration of Ni in the sample solution.
Compiled by Mambanda Isaac(Chief Technician) & Gonzo Muriel(Snr Lab Technician)
11
CT104 PRCTICAL SCHEDULE 2: GRAVIMETRY
Title: Gravimetric determination of sulphate as barium sulphate.
Theory
The precipitation of barium sulphate is the most widely used process for the gravimetric determination of
the sulphate ion. The reaction is:
SO42-(aq) + Ba2+(aq)→ BaSO4(s)
This determination is fraught with many possible sources of error of which coprecipitation and occlusion of
solution are among the greatest. The digestion process described is necessary to provide a readly filterable
precipitate. The ageing of the precipitate depends on its solubility in water; the fine particles dissolve more
readily and the larger particles grow larger by crystal growth from the saturated or even super-saturated
solution of the barium sulphate.
It is usual to filter the precipitate on either a filter paper (fine grade) or pulp pad. Drying the precipitate at
110-120
degrees celcius is not satisfactory and ignition at 800-1000 C is necessary to remove the water which is
firmly bound in the crystal lattice. During the ignition of the precipitate especially in the presence of carbon
from the charring paper, free access of air is essential to avoid reduction of the sulphate to sulphide.
Procedure.
Weigh accurately an amount of the sample provided between 0.3-0.5g into each of the two 400 ml beakers,
dessolve in about 200 ml of distilled water and add 2.5 ml of 4M HCL. Heat to boiling and add, with
stirring, and in aslow stream, 10 ml of a 10% barium chloride solution. Cover the beaker with a watch glass
and maintain the solution just belowits boiling point for about 30 minutes. Allow the precipitate to settle
and add a few drops of precipitant to confirm that precipitation is complete. Leave to cool in your
cupboard. Decant the solution through a No. 42 ashless filter paper. Wash the precipitate and filter with hot
distilled water until the filtrate is essentially chloride-free.
Transfer the filter paper to a pre-weighed porcelain crucible and dry over a very low flame from a bunsen
burner (or in a drying oven). Char the paper gently allowing the air to circulate in the crucible but taking
care to prevent the paper catching fire. Use the crucible lid if necessary to stifle the flames. Then heat the
crucible very strongly with a bunsen flame until incineration and ingnition is complete. Transfer to
desiccator. Weigh when cool.
Calculate the percentage of sulphate in the sample.
Compiled by Mambanda Isaac(Chief Technician) & Gonzo Muriel(Snr Lab Technician)
12
CT104 PRCTICAL SCHEDULE 3: COMPLEXOMETRIC TITRATION
Title: Determination of the percentage of zinc in hydrated zinc sulphate.
NOTES:
Metal-indicator + EDTA → Metal-EDTA + Indicator
(Red)
(Blue)
MATERIALS
0.1M EDTA solution
Zinc Sulphate crystals
Solid Eriochrome Black T
Ammonia-ammonium chloride buffer solution (pH 10)
Procedure
Weigh out about 7g of zinc sulphate crystals, dissolve in distilled water and dilute to exactly 250 ml.
Pipette 25 ml of this and add 2 ml of ammonia-ammonium chloride buffer solution (pH 10) and enough
solid Eriochrome Black T indicator to give a clearly visible colour. Titrate with the standard EDTA
solution (0.1M) until the colour changes from red to blue.
The reaction is essentially: Zn2+ + Y4-→ ZnY2Calculate the percentage of zinc in the sample.
Compiled by Mambanda Isaac(Chief Technician) & Gonzo Muriel(Snr Lab Technician)