Acidimetric Titration
Theory of acidimetry:
Acidimetry, essentially involves the direct or residual titrimetric analysis of alkaline
substances (bases) employing an aliquot of acid and is provided usually in the analytical control
of a large number of substances. Examples :
(a) Organic substances : urea, sodium salicylate, diphenhydramine, emetine hydrochloride,
meprobamate, paramethadione, pyrazinamide etc.
(b) Inorganic substances : sodium bicarbonate, milk of magnesia, ammonium chloride,
calcium hydroxide, lithium carbonate, zinc oxide etc.
Direct titration method:
It is an usual practice that when a solid substance is to be assayed, an aliquot quantity of the
same may be weighed accurately and dissolved in sufficient water so that the resulting solution
should have more or less the same equivalent concentration as that of the acid used in the
titration. Methyl orange (pH range = 3.0 to 4.4) is the indicator of choice for obvious reasons,
as phenolphthalein and most other indicators are instantly affected by the carbonic acid
(H2CO3) generated in the reaction which ultimately cause a change in colour even before the
reaction attains completion.
Residual titration method:
Residual titration or back titration is normally employed in the following two situations,
namely :
Case I : when a chemical reaction proceeds rather slowly or sluggishly.
Case II : when the substance under determination fails to give a sharp and distinctly visible
end-point with an indicator by direct titration.
In usual practice, the residual titration is accomplished by allowing to dissolve the substance
under estimation in an accurately measured quantity of a standard solution of known strength
present in excess and subsequently titrating the excess of the latter with another previously
standardized solution.
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Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
practical lab 1
.
Acidimetric Titration
A: Direct Titration Method
Assay of Sodium Carbonate
Introduction
Sodium carbonate, Na₂CO₃, is a sodium salt of carbonic acid. It most commonly occurs
as a crystalline heptahydrate, which readily effloresces to form a white powder, the
monohydrate. Formula: Na2CO3, Molar mass: 105.9885 g/mol, Melting point: 851 °C,
Density: 2.54 g/cm³, Boiling point: 1,633 °C ,Soluble in water and very slightly soluble in
alcohol, odorless powder that absorbs moisture from the air, has an alkaline taste, and forms
a strongly alkaline water solution. It is one of the most basic industrial chemicals.
Sodium carbonate decahydrate, Na2CO3·10H2O, is a colorless, transparent crystalline
compound commonly called Sal soda or washing soda. Because seaweed ashes were an
early source of sodium carbonate, it is often called soda ash or, simply, soda.
Used topically for dermatitides, mouthwash, vaginal douche; veterinary use as emergency
emetic . In solution to cleanse skin, in eczema, to soften scabs of ringworm. For the temporary
relief of symptoms due to constipation: abdominal bloating, abdominal discomfort,
irregularity of bowel movements. ...
principle:
Direct acid base titration of sodium carbonate against 0.5N sulphuric acid and by using
methyl orange solution as indicator.
The equation of reaction is:
2Na2CO3 + H2SO4
Na2SO4 + 2H2O + 2CO2
Materials Required :
1 g of sodium carbonate ; 0.5 N sulphuric acid.
Procedure :
1. Weigh accurately about 1 g, of sodium carbonate.
2. dissolve it in 20 ml of water (DW)
3. titrate with 0.5 N sulphuric acid, using methyl orange solution as indicator.
.
Each ml of 0.5 N sulphuric acid is equivalent to 0.53 gm of Na2CO3.
Calculation:
% Na2CO3 = V x N x meq. wt. x 100 / wt of sample
Cognate Assays :Sodium bicarbonate ; sodium salicylate tablets.
2
Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
practical lab 2
Assay of borax
sodium borate or sodium tetraborate
Na2B4O7.10H2O
introduction:
Borax, also known as sodium borate, sodium tetraborate, or disodium tetraborate, is an
important boron compound, a mineral, and a salt of boric acid. Powdered borax is white,
consisting of soft colorless crystals that dissolve easily in water. Its molecular weight 381.4
,melting point 743ºC ,and boiling point 1,575 ºC.
Borax has a wide variety of uses. It is a component of many detergents, cosmetics,
and enamel glazes Borax .It should be kept in well closed container. Borax is used as antiseptic
and astringent in mouth paints.. It is also used to make buffer solutions in biochemistry, as
a fire retardant, as an anti-fungal compound for fiberglass, as a flux in metallurgy, neutroncapture shields for radioactive sources, a texturing agent in cooking, and as a precursor for
other boron compounds.
. Borax is generally described as Na2B4O7·10H2O. However, it is better formulated as
Na2[B4O5(OH)4]·8H2O, since borax contains the [B4O5(OH)4]2− ion. In this structure, there are
two four-coordinate boron atoms (two BO4 tetrahedra) and two three-coordinate boron atoms
(two BO3triangles).
Borax is also easily converted to boric acid and other borates, which have many applications.
Its reaction with hydrochloric acid to form boric acid is:
Na2B4O7·10H2O + 2 HCl → 4 H3BO3 + 2 NaCl + 5 H2O
The "decahydrate" is sufficiently stable to find use as a primary standard for acid
base titrimetry
When borax is added to a flame, it produces a yellow green color Borax is not used for this
purpose in fireworks due to the overwhelming yellow color of sodium. Boric acid is used to
color methanol flames a transparent green.
principle:
Borax is a salt derived from a weak acid and a strong base, so its aqueous solution can be
assayed using a standard N/2 Hydrochloric acid solution in an acid base titration.
Na2B4O7.H2O+HCl
Borax
2NaCl +4H3BO3 +5H20
Boric acid
Borax samples may sometimes be contaminated with boric acid or sodium carbonate, Thus
two titrations are carried out;
1. borax and (sodium carbonate if present) with Hydrochloric acid.
2. boric acid with sodium Hydroxide.
The correct volumes of the two titrations can then tell the presence and type of impurities,
besides the quantities of each.
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Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
Reaction between borax and Hydrochloric acid liberates boric acid a very weak acid that
needs to increase its acidity ( by converting it into a strong acid) so that it can react with
sodium hydroxide by direct titration. This is achieved by the addition of excess of mannitol, a
poly hydroxylated alcohol.
Mannityl boric acid complex is formed which is a strong acid.
mannitol
4H3BO3 + 4NaOH
4NaBO2 + 8H2O
procedure:
1. Put 10ml of borax solution in conical flask and add to it 10ml of D.W .
2. Add 2 drops of methyl red solution.
3. Titrate with 0.5N HCl drop wise with shaking until the colour changes from yellow
(pH6) to pink (pH 4.4).
4. Boil the titration liquid to expel the generated carbon dioxide (if any) and cool.
5. If the colour changes back into yellow, resume titration with HCl until you get the pink
colour. Record the volume of HCl used.
6. Add 4gm of mannitol and 2 drop of phenolphthalein solution (pH-8 colourless ;pH-10
pink).
7. Start the second titration with 1N NaOH solution drop wise with shaking until the
solution changes from pink to yellow and pink again. Record the volume of NaOH
solution used.
The solution of the titration should be boiled after the end point to expel carbon dioxide
generated (if sodium carbonate is present):
Na2CO3 + 2HCl
2NaCl + H2O +CO2
This is to prevent the formation of carbonic acid which may affect the result of the second
titration with sodium hydroxide solution.
Mannitol is added in excess quantitey to prevent backward hydrolysis of mannityl boric acid
complex.
Calculations:
V1 = Volume of 0.5N HCl
V2 = Volume of 1N NaOH
1. If V1 = V2
Then the sample of borax is pure. Sometimes a differences of not more than 0.3ml
between the two volumes is allowed. in this case use the average.
Calculate the weight and percentage w /v of borax in your sample.
Each 0.09535 gm of borax is equivalent to 1ml of 0.5N HCl solution.
2. If the volume of 0.5N HCl solution consumed is more than the volume of 1N NaOH
solution consumed:
V1 > V2
Then the sample of borax is impure and contains, in addition to borax, sodium carbonate.
V1 - V2 = the volume of 0.5N HCl solution consumed by sodium carbonate.
Calculate the weight and percentage w/v of bath borax and sodium carbonate in your
sample.
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Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
Each 0.09535gm of borax is equivalent to 1ml of 1N NaOH solution.
Each 0.0265gm of sodium carbonate equivalent to 1ml 0f 0.5 N HCl solution.
3. The volume of 1N NaOH solution consumed is more than the volume of 0.5N HCl
solution consumed:
V2 > V1
mannitol
H3BO3 +NaOH
liberated and
contaminant boric acid
NaBO2 + H2O
Then the sample of borax is impure and contains in addition to borax, boric acid.
V2 -V1 = the volume of 1N NaOH solution consumed by the contaminant boric acid.
Calculate the weight and percentage w/v of bath borax and contaminant boric acid in your
sample.
Each 0.06184 gm of boric acid is equivalent to 1 ml of 1N NaOH solution.
Each 0.09535 gm of borax is equivalent to 1ml of 0.5N HCl solution.
5
Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
practical lab 3
B: Residual titration method
Assay of Zinc Oxide
introduction:
Zinc oxide is an inorganic compound with the formula ZnO. ZnO is a white powder that is
insoluble in water, melting point 1975C ,molar mass 81.408 . it is widely used as an additive in
numerous materials and products including rubbers, plastics, ceramics, glass, cement,
lubricants, paints, ointments, adhesives, sealants, pigments, foods (source of Zn nutrient),
batteries, ferrites, fire retardants, and first-aid tapes. It occurs naturally as the mineral zincite,
but most zinc oxide is produced synthetically.
Zinc oxide as a mixture with about 0.5% iron(III) oxide (Fe2O3) is called calamine and is
used in calamine lotion. Zinc oxide can be used in ointments, creams, and lotions to protect
against sunburn and other damage to the skin caused by ultraviolet light (see sunscreen).
principle:
Equation :
ZnO + H2SO4
(81.38)
ZnSO4 + H2O
The requisite quantity of ZnO gets dissolved in the sulphuric acid thereby neutralizing an
equivalent amount as shown by the above equation. Thus, the amount of sulphuric acid
neutralized by the ZnO is estimated by subtracting, from the total amount of sulphuric acid
utilized, the quantity neutralized by the standard NaOH in the back titration
Materials Required :
1.5 g of zinc oxide ; 1 N sulphuric acid ; 1 N sodium hydroxide ; 2.5 g ammonium chloride.
Procedure :
1.Weight accurately 1.5 g of freshly ignited and cooled zinc oxide.
2. dissolve it with 2.5 g of ammonium chloride in 50 ml of 1 N sulphuric acid with the help of
gentle heating.
3. After complete dissolution, add methyl orange and titrate the excess of sulphuric acid with
1 N sodium hydoxide.
Each millilitre of 1 N sulphuric acid is equivalent to 40.6 mg of ZnO.
. The equivalent weight of ZnO, as shown in the above equation comes out to be 40.69 g
(i.e., 81.38/2 = 40.69). Hence, each millilitre of 1 N sulphuric acid, 1 meq neautralized by the
ZnO, is equivalent to 40.68 mg or 1 meq of ZnO.
Thus, the percentage of zinc oxide present in the sample may be calculated as follows :
% ZnO =(V1- V2) x meq.wt x100 / wt of sample
V1 =Tatal volume of H2SO4
V2=Volume of H2SO4 react with NaOH
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Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
Oxidation- Reduction Methods
Permanganate, Dichromate And Sulphate Titration Methods
The oxidation and reduction processes essentially take place simultaneously in a reaction, thus
one entity gets reduced in the process of oxidizing the second. ‘Redox’—is the abbreviated form
of reduction— oxidation systems. In the oxidation—reduction methods of analysis a change in
valence of the reacting products is a must which is contrary to precipitation and neutralization
methods of analysis where no change in valence occur. The major oxidizing agents normally
employed in volumetric titrations include, potassium permanganate, potassium dichromate, and
ceric sulphate
Assay Methods
The quantitative estimations of a number of pharmaceutical substances may be carried out by
using a variety of potential oxidizing agents as stated below :
(i) Permanganate Methods :
(a) Direct Titration Methods,
(b) Indirect Titration Methods.
(c) Residual Titration Methods.
(ii) Dichromate Methods :
Direct titrations with Potassium Dichromate.
(iii) Ceric Sulphate Titration Methods :
Direct Titrations with Ceric Sulphate
permanganate methods
The vital application of potassium permanganate as a potential oxidizing agent in an acidic
medium mainly rests on the reactions designated by the following equations :
Chemically we have :
2KMnO 4 + 3H2SO4 K2SO4 + 2MnSO4 + 3H2O + 5(O)
Ionically we have :
MnO4– + 8H+ + 5e Mn2 + + 4 H2O
Therefore, KMnO4 5e
or 158.0 g KMnO4 5000 ml N
or 31.60 g KMnO4 1000 ml N
or 3.16 g KMnO4 1000 ml 0.1 N KMnO4
7
Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
practical lab 4
Preparation And standardization of 0.1 N Potassium Permanganate
Solution:
introduction:
Potassium permanganate is an inorganic chemical compound with the formula KMnO4. It is
a salt consisting
of K+ and MnO−4 ions.
Formerly
known
as permanganate
of
potash or Condy's crystals, it is a strong oxidizing agent. It dissolves in water to give
intensely pink or purple solutions, the evaporation of which leaves prismatic purplish-black
glistening crystals. In 2000, worldwide production was estimated at 30,000 tonnes. In this
compound, manganese is in the +7 oxidation state.
1. preparation:
Materials Required :
Potassium permanganate : 3.5 g.
Procedure :
1. Weigh accurately about 3.2 g of potassium permanganate on a watch-glass. Transfer the
contents to a 250 ml beaker containing cold water and stir vigorously with a glass rod
to effect rapid dissolution.
2. Decant the solution through a small plug of glass wool supported by a funnel, into a 1
liter volumetric flask thereby leaving the un dissolved residues in the beaker.
3. Add more DW to the beaker and repeat the above process till all the potassium
permanganate gets dissolved.
4. Finally make up the volume to the graduated mark and shake well so as to effect
uniform mixing.
Note :
(i) KMnO4 must be weighed on a watch-glass and not on any kind of paper since
cellulose fibers are corrosively attacked by it,
(ii) Likewise, filtration of KMnO4 solution must be done though cleaned glass wool
and not cotton wool, and
(iii) Avoid heat in the preparation of KMnO4 solution because traces of grease or
other possible contaminants on the glass vessels used can catalyse its decomposition.
2. standardization:
Materials Required :
Oxalic acid : 6.3 g ; sulphuric acid concentrated : 5 ml.
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Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
Procedure :
1. Weigh accurately about 6.3 g of pure oxalic acid (AnalaR-Grade) into a 1 liter
volumetric flask, dissolve it in sufficient DW and make up the volume up to the
mark.
2. Pipette out 25 ml of this solution, add to it 5 ml of concentrated sulphuric acid
along the side of the flask, swirl the contents carefully and warm up to 70°C.
3. Titrate this against the potassium permanganate solution from the burette till the
pink color persists for about 20 seconds.
Precautions :
(i) Sufficient acid must be present, otherwise formation of a brown color during titration
may be observed,
(ii) Similar brown coloration can also be observed by using too high a temperature or by
using a dirty flask
(iii) To avoid such anomalies always rinse the flask with solution of H2O2 and dilute H2SO4
before performing the titrations.
Safety and Waste:
• Wear safety goggles at all times while in the laboratory.
• Potassium permanganate is a strong oxidizing agent and can be damaging to skin, eyes, and
clothing. Wash thoroughly after handling.
• Potassium permanganate solutions will stain skin and clothing if spilled.
• Sulfuric acid solutions are damaging to the skin, eyes, and clothing–especially if allowed to
concentrate through evaporation of water. If the sulfuric acid solution is spilled on the skin or
clothes, wash immediately and inform the instructor.
• Iron salts may be irritating to the skin. Wash after handling.
• Dispose of any excess KMnO4 in the appropriate waste container after insuring that no other
groups need additional solution.
Direct Titration Methods
Hydrogen peroxide solution and potassium bromide are two pharmaceutical substances that
may be estimated by employing 0.1 N potassium permanganate solution and adopting the direct
titration method.
9
Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
practical lab 5
Assay of Hydrogen Peroxide Solution
Redox titration
introduction:
Hydrogen peroxide is a chemical compound with the formula (H2O2).its molecular mass
(34.017g/mo), Melting point (-0.43ºC), boiling point( 150.2ºC), pka 11.75. It is the
simplest peroxide (a compound with an oxygen-oxygen single bond) and in its pure form is a
colorless liquid, slightly more viscous than water. For safety reasons it is normally encountered
as an aqueous solution, also colorless. Hydrogen peroxide is a strong oxidizer and is used as
a bleaching agent and disinfectant. Concentrated hydrogen peroxide, or 'high-test peroxide' is
a reactive oxygen species and has been used as a propellant in rocketry.
Organisms naturally produce trace quantities of hydrogen peroxide, most notably by
a respiratory burst as part of the immune response. Hydrogen peroxide exhibits oxidizing and
reducing properties, depending on pH.
In acidic solutions, H2O2 is one of the most powerful oxidizers known—strong than
chlorine, chlorine dioxide, and potassium permanganate. Also, through catalysis, H2O2 can be
converted into hydroxyl radicals (•OH), which are highly reactive.
Hydrogen peroxide
M.wt
34.017g/mol
principle:
This method utilizes the reduction of potassium permanganate (KMnO4) by hydrogen peroxide
in sulfuric acid. Hydrogen peroxide is usually treated as a strong oxidizer, but in the presence
of even stronger oxidizer it can become a reducing agent:
H2O2 → O2 + 2H+ + 2ePermanganate in low pH is strong enough to quantitatively oxidize hydrogen peroxide to
oxygen. This reaction is used for the determination of hydrogen peroxide concentration.
Reaction taking place during titration is
2MnO4- + 5H2O2 + 6H+ → 2Mn2+ + 5O2 + 8H2O
10
Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
Materials Required :
Hydrogen peroxide solution : 10 ml ; 5 N sulphuric acid : 5 ml ; 0.1 N potassium permanganate.
Procedure :
1. Dilute 10 ml of hydrogen peroxide solution to 250 ml with DW in a volumetric flask.
2. To 25.0 ml of this solution add 5 ml of 5 N sulphuric acid
3. titrate with 0.1 N KMnO4 to a permanent pink endpoint.
Each ml of 0.1 N potassium permanganate is equivalent to 0.001701 g of H2O2
.
Equations :
Chemically, we have:
2KMnO4 + 3H2SO4
5H2O2 + 5(O)
K2SO4 + 2MnSO4 + 3H2O + 5(O)
5O2 + 5H2O
Summing up : 5H2O2 + 2KMnO4 + 3H2SO4
Ionically we have :
K2SO4 + 2MnSO4 + 8H2O + 5O2
2MnO–4 + 6H+ + 5H2O2
2Mn2+ + 8H2O + 5O2
Therefore, 5H2O2 2MnO4–l0e
or 34.02 g H2O2 2000 ml N
or 17.01 g H2O2 1000 ml N
or 0.001701 g H2O2 1 ml 0.1 N KMnO4
Calculations :
(For % w/v of H2O2)
The ‘volume strength’ of the hydrogen peroxide solution is the number of ml of oxygen at
NTP* which may be produced by the complete thermal decomposition of 1 ml of H2O2
solution. Hence, decomposition takes place as designated by the following equation :
2H2O2
2H2O + O2
or 68.04 g H2O2 22400 **ml O2
or 1 g H2O2 329.2 ml O2
The IP limit of H2O2 solution is 5-7% w/v.
Now, let us consider a sample which contains 6.25 per cent w/v H2O2 :
Therefore, 100 ml sample ≡ 6.25 g H2O2
or 1 ml sample  0.0625 g H2O2  0.0625 × 329.2 ml O2
 20.58 ml O2
* NTP = Normal temperature and pressure
** At standard temperature and pressure (STP) 1 mole of O2  22.4 L
11
Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
practical lab 6
Redox Titration – Potassium Permanganate with an Iron(II) Salt
Assay of ferrous sulfate
(FeSO4.7H2O)
Introduction:
ferrous sulfate is the chemical compound with the formula FeSO4.Its molecular weight
278.05,melting point 329.9ºC. It has a pale green colour as a powder, and as crystals it has a
bluish green colour, freely soluble in water. It is used medically to treat iron deficiency, and
also for industrial applications. Known since ancient times as copperas and as green vitriol, the
blue-green heptahydrate is the most common form of this material.
All iron sulfates dissolve in water to give the same aquo complex[Fe(H2O)6]2+, which
has octahedral molecular geometry and is paramagnetic.
Together with other iron compounds, ferrous sulfate is used to fortify foods and to treat irondeficiency anemia. Constipation is a frequent and uncomfortable side effect associated with the
administration of oral iron supplements. Stool softeners often are prescribed to prevent
constipation.
principle:
. In acidic solution, potassium permanganate rapidly and quantitatively oxidizes iron(II) to
iron(III), while itself being reduced to manganese(II). The half reactions for the process are:
When these half-reactions are combined to give the overall balanced chemical reaction
equation, a factor of five must be used with the iron half-reaction so that the number of
electrons lost in the overall oxidation will equal the number of electrons gained in the
reduction:
Potassium permanganate is one of the most commonly used oxidizing agents because it is
extremely powerful, inexpensive, and readily available. Potassium permanganate is particularly
useful among titrants since it requires no indicator to signal the endpoint of a titration.
Potassium permanganate solutions–even at fairly dilute concentrations–are intensely colored
purple. The product of the permanganate reduction half-reaction, manganese(II), in dilute
solution shows almost no color. Therefore, during a titration using KMnO4 , when one drop
excess of potassium permanganate has been added to the sample, the sample will take on a pale
red/pink color (since there are no more analyte ions remaining to convert the purple MnO 4ions to the colorless Mn2+ ions.
12
Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
procedure:
1. carefully weigh ~0.5 g of ferrous sulfate in a beaker , Add 15 ml of distilled water and
stir to dissolve the solid.
2. Add 8 ml of 3 N sulfuric acid, H2SO4 , to the sample (Caution!). Sulfuric acid is added
to the samples to provide the hydrogen ions, H+, required for the reduction of the
permanganate ion.
3. Titrate with Potassium Permanganate Solution to the endpoint. The end point is the first
appearance of a permanent, pale pink colour.
calculation:
1ml of 0.1N KMnO4  0.02780 gm FeSO4 .7H2O



13
Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali

Iodimetric and iodometric titration
ASSAY METHODS
Assay methods involving iodine can be categorized under the following heads namely :
A. Iodimetric Assays :
(a) Direct titration with iodine,
(b) Residual titration method : i.e., excess of iodine is titrated with sodium thiosuphate.
B . Iodometric Assays : i.e., release of iodine and subsequent titration with sodium
thiosulphate.
Iodimetric Assay:
In such estimations, the pharmaceutical substances can be measured either directly or back
titration of excess iodine with sodium thiosulphate solution
14
Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
practical lab 7
Preparation and standardization of 0.1 Iodine Solution
Introduction:
Iodine is a chemical element with symbol I and atomic number 53. The name is
from Greek ἰοειδής ioeidēs, meaning violet or purple, due to the color of elemental iodine
vapor. Melting point: 113.7 °C, boiling point: 184.3 °C, atomic mass: 126.90447 g/mol.
Iodine and its compounds are primarily used in nutrition, and industrially in the production
of acetic acid and certain polymers. Iodine's relatively high atomic number, low toxicity, and
ease of attachment to organic compounds have made it a part of many X-ray contrast materials
in modern medicine. Iodine has only one stable isotope. A number of iodine radioisotopes are
also used in medical applications.
Iodine is required by higher animals for synthesizing thyroid hormones, which contain the
element. Because of this function ,radioisotopes of iodine are concentrated in the thyroid
gland along with nonradioactive iodine. If inhaled, the radioisotope iodine-131, which has a
high fission product yield, concentrates in the thyroid, but is easily remedied with nonradioactive potassium iodide treatment.
Elemental iodine is used as a disinfectant in various forms. The iodine exists as the element, or
as the water-soluble triiodide anion I3− generated in situ by adding iodide to poorly watersoluble elemental iodine (the reverse chemical reaction makes some free elemental iodine
available for antisepsis). In alternative fashion, iodine may come from iodophors, which
contain iodine complexed with a solubilizing agent (iodide ion may be thought of loosely as the
iodophor in triiodide water solutions).
Examples of such preparations include:



Tincture of iodine: iodine in ethanol, or iodine and sodium iodide in a mixture of ethanol
and water.
Lugol's iodine: iodine and iodide in water alone, forming mostly triiodide. Unlike tincture
of iodine, Lugol's has a minimized amount of the free iodine (I2) component.
Povidone iodine (an iodophor).
1. Preparation of 0.1 Iodine Solution:
Theory : Iodine in aqueous solution acts as an oxidizing agent which forms the basis of assay
method involving direct titration with iodine. Thus, we have:
I2 + 2e →2I–
or I2 ≡2e
or 126.9 g I2 ≡1000 ml N
15
Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
or 12.69 g I2 ≡1000 ml 0.1 N
or 3.17 g I2 ≡250 ml 0.1 N
Materials Required :
Iodine : 3.2 g ; potassium iodide : 7.5 g.
Procedure :
Weigh accurately 3.2 g of crushed iodine crystals on a watch glass and transfer to a beaker
containing potassium iodide (7.5 g) and water (10 ml).
Dissolve the contents of the beaker with the help of a glass rod and frequent swirling.
Transfer the contents of the beaker quantitatively to a 250 ml volumetric flask and make up the
volume with DW.
Explanation :
Iodine is sparingly soluble in water but undergoes rapid dissolution in the presence of
potassium iodide due to the formation of the corresponding triiodide ion :
I2 + I-–→I-3
Thus, potassium iodide plays dual role, viz., in iodimetry—to solubilize iodine in aqueous KI
solution, and in iodometry—as reducing agent, the excess KI helps in retaining liberated I2 in
solution through interaction with KI.
2. Standardization of 0.1 Iodine Solution by the aid of Sodium Thiosulphate
Theory :
Iodine solution may be standardized by using sodium thiosulphate (AR-Grade) whereby
the latter gets oxidized to sodium tetrathionate as expressed below :
or
o
or
or
or
2S2O3 -2S4O6-2 + 2e
I2 + 2e–2I–
2S2O -2 I2 2e
2 × 248.2 g Na2S2O3.5H2O ≡ 2000 ml N
248.2 g Na2S2O3.5H2O ≡ 1000 ml N
24.82 g Na2S2O3.5H2O ≡ 1000 ml 0.1 N Iodine
Materials Required :
Sodium thiosulphate (AR) : 6. 25 g ; 0.1 N I2 solution.
Procedure :
1. Weigh accurately 6. 25 g of sodium thiosulphate (AR) to a 250 ml volumetric flask.
2. Dissolve it in DW, shake well and make up the volume to the mark with DW.
3. Pipette 25 ml of 0.1 iodine solution into an iodine flask.
4. Titrate with the standard sodium thiosulphate solution, before the E.p add 5 drops of
starch solution and titrate until the solution becomes almost colorless.
Note : Stock solutions of sodium thiosulphate may be preserved by the addition of a few
drops of sodium hydroxide solution (20% w/v) which serves as stabilizer as well as
prevents decomposition.
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Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
Preparation of Starch Solution
Material Required :
Starch 1.0 g.
Procedure :
1. Weigh 1.0 g starch in a glass in a glass pestle-mortar and triturate thoroughly with 10
ml of cold DW.
2. Boil separately 200 ml of DW in a beaker and add the starch paste to it with vigorous
stirring.
3. The resulting mixture is boiled gently for a further period of 30 minutes till a
transluscent and thin liquid having an uniform consistency is obtained.
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Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
practical lab 8
Assay of copper sulfate
CuSO4.5H2O
Introduction:
Copper(II) sulfate, also known as cupric sulfate or copper sulphate, is the chemical
compound with the chemical formula CuSO4. This salt exists as a series of compounds that
differ in their degree of hydration. The anhydrous form is a pale green or gray-white powder,
whereas the pentahydrate (CuSO4·5H2O), the most commonly encountered salt, is bright blue.
Copper(II) sulfate exothermically dissolves in water to give the aquo complex [Cu(H2O)6]2+,
which has octahedral molecular geometry and is paramagnetic. its molar mass:159.6 g/mol
anhydrous, and 249.6g/mol pentahydral .It used as a herbicide, fungicide and pesticide, It is
used in swimming pools as an algicide.
Several chemical tests utilize copper sulfate. It is used in Fehling's solution and Benedict's
solution to test for reducing sugars, which reduce the soluble blue copper(II) sulfate to
insoluble red copper(I) oxide. Copper(II) sulfate is also used in the Biuret reagent to test for
proteins.
Copper sulfate is also used to test blood for anemia. The blood is tested by dropping it into a
solution of copper sulfate of known specific gravity – blood which contains
sufficient hemoglobin sinks rapidly due to its density, whereas blood which does not sink or
sinks slowly has insufficient amount of hemoglobin.
In a flame test, its copper ions emit a deep green light, a much deeper green than the flame
test for barium.In the presence of chlorine, copper ions emit a deep blue light.
principle:
Iodometric determination of copper is based on the oxidation of iodides to iodine by copper
(II) ions, which get reduced to Cu+.
Comparison of standard potentials for both half reactions (Cu2+/Cu+ E0=0.17 V, I2/IE0=0.54 V) suggests that it is iodine that should be acting as oxidizer. However, that's not the
case, as copper (I) iodide CuI is very weakly soluble (Ksp = 10-12). That means concentration of
Cu+ in the solution is very low and the standard potential of the half reaction Cu2+/Cu+ in the
presence of iodides is much higher (around 0.88 V).
In effect reaction taking place in the solution is
2Cu2+ + 4I- → 2CuI(s) + I2
and produced equivalent amount of iodine can be titrated with thiosulfate solution.
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Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
For the best results reaction should take place in the slightly acidic solution (pH around 4-5),
correct pH is obtained by addition of ammonia and acetic acid, effectively creating acetic
buffer. Lab practice shows that the end point is sharper when we add some thiocyanate to the
solution. Copper (I) thiocyanate is slightly less soluble than iodide, which makes concentration
of Cu+ even lower, increasing the oxidation potential of the Cu2+/Cu+ system.
Solution should be free of other substances that can oxidize iodides to iodine (for example
Fe3+ or nitrates).
As it was already explained, first reaction taking place is:
2Cu2+ + 4I- → 2CuI(s) + I2
This is followed during titration by the reaction of the iodine with the thiosulfate:
2S2O32- + I2 → S4O62- + 2I-
solutions used
To perform the determination we will need concentrated ammonia and concentrated acetic acid
solutions, solid potassium iodide, titrant - 0.1 M thiosulfate solution, and indicator - starch.
procedure
Procedure below assumes that original solution is acidic or neutral.
1. Pipette aliquot containing copper (II) into 250 mL conical flask with a glass stopper.
2. Add concentrated ammonia till solution turns dark blue.
3. Add concentrated acetic acid till solution loses dark blue color, and then about 3mL.
4. Add 2 g of solid potassium iodide, swirl well.
5. Put stopper on the flask and put solution in a dark place for 5 minutes.
6. Titrate swirling the flask, until a pale yellow.
7. Add 5 ml of the starch solution.
8. Add 1 g of potassium thiocyanate.
9. Titrate swirling the flask, until blue color disappears.
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Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
calculations:
1ml 0f 0.1N Na2S2O3  0.02497 gm of CUSO4.5H2O
V x ch.f. = gm of CUSO4 IN 10 ml of unknown
gms of CUSO4 /10 x100 =% w /v for liquid unknown.
Discussion:
Explain why?
1. we used acidic media.
2. we add KSCN.
3. The mixture was turbid
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Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
Argentometric methods
Introduction:
In general, titrations governed by precipitation reactions do not really constitute an
appreciable number in volumetric determinations in comparison to either redox or acid-base
reactions. The interaction between silver-nitrate and sodium chloride in solutions result into the
precipitation of silver chloride as shown below:
NaCl + AgNO3 → AgCl ↓ + NaNO3
In actual practice, however, such titrations are more or less restricted to those involving
precipitation of Ag+ with anions, for instance : halogens (Cl–, Br–, I–) and thiocyanate (SCN–).
Generally, it is quite difficult and tedious to locate the exact point at which further addition of
reagent affords no more precipitation. Therefore, the choice and wisdom of a chemical
reaction is preferably sought so as to result in either a colored solution or a colored precipitate
at the end point.
A typical instance may be cited by application of potassium chromate
(K2CrO4) solution in the above case whereby any extra drop of silver nitrate, after all the
chloride has been precipitated, immediately causes precipitation of red chromate showing that
the end point has been duly achieved.
It is, however, interesting to observe here that such reactions do offer limited usage
because of the following two facts:
(i) Co-precipitation effects do not give a real composition of the precipitate.
(ii) Choice of appropriate indicator is very much limited.
Besides, the foregoing facts another vital aspect to be taken into consideration is the solubility
product that plays a major role in such titration. Hence, the equilibrium constant of the reaction
giving the precipitate
The four cardinal parameters that may be considered for a feasible argentometric analysis
are:
(i) Precipitate formed must be insoluble,
(ii) Precipitation process should be fast and rapid,
(iii) Co-precipitation effects must be minimal,
(iv) Detection of equivalence point must be apparently visible.
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Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
Assay methods
Argentometric titrations may be divided into two broad categories:
(i) Direct titration with silver-nitrate.
(ii) Ammonium thiocyanate-silver nitrate titrations (Volhard’s Method).
Direct titration with silver nitrate
Pharmaceutical substances essentially containing halides may be estimated by direct titration
with silver nitrate solution as a titrant.
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Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
practical lab 9
1. Preparation of 0.1 N Silver Nitrate Solution
Materials Required:
Silver nitrate (AR): 16.989 g.
Procedure: Weigh accurately 16.989 g of silver nitrate on a watch-glass and transfer
quantitatively into a 1 litre volumetric flask. Add freshly prepared DW and make up the volume
to 1000 ml. Thus, we have:
AgNO3 + NaCl → AgCl ↓ + NaNO3
or AgNO3 ≡ NaCl ≡ H
or 169.89 g AgNO3 ≡ 58.45 g NaCl ≡ 1000 ml N
or 0.01699 g AgNO3 ≡ 0.005845 g NaCl ≡ 1 ml 0.1 N AgNO3
2. Standardization of 0.1 N Silver Nitrate Solution
Materials Required :
Sodium chloride : 0.1 g ; acetic acid (33% w/v) : 5 ml ; methyl alcohol (95%) : 50 ml ; eosin
solution (0.5% w/v in water) : 5 ml ; 0.1 N silver nitrate solution.
Procedure:
1. Weigh accurately about 0.1 g of sodium chloride, previously dried at 110°C for 2
hours,
2. dissolve in 5 ml of water. Add 5 ml of acetic acid, 50 ml of methyl alcohol and three
drops of eosin solution.
3. Stir thoroughly on a magnetic stirrer and titrate with the silver nitrate solution till the
white particles of AgCl change from white to pink.
4. Each 0.005844 g of sodium chloride is equivalent to 1 ml of 0.1N silver nitrate.
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Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali
Practical lab 10
Assay of Potassium Chloride
introduction:
The chemical compound potassium chloride is a metal halide salt composed of potassium
and chlorine. In its pure state, it is odorless and has a white or colorless vitreous crystal
appearance, , with a crystal structure that cleaves easily in three directions.
Its formula: KCl molar mass: 74.5513 g/mol, melting point: 770 °C, density: 1.98 g/cm³,
boiling point: 1,420 °. Freely soluble in water; insoluble in ethanol
KCl is used in medicine, scientific applications, and food processing. It occurs naturally
as the mineral sylvite and in combination with sodium chloride as sylvinit.
principle:
AgNO3 + KCl →AgCl ↓+ KNO3
or AgNO3 ≡KCl ≡H
or 169.89 g AgNO3 ≡74.55 g KCl ≡1000 ml N
or 0.01699 g AgNO3 ≡0.007455 g KCl ≡1 ml of 0.1 N AgNO3
Materials Required :
Potassium chloride : 0.25 g ; potassium chromate solution (5% w/v in water) :
10 ml ; 0.1 N silver nitrate solution.
Procedure :
1. Weigh accurately about 0.25 g of potassium chloride in a conical flask.
2. Dissolve it in 50 ml of DW .
3. titrate with 0.1 N silver nitrate solution, using 2-3 drops of potassium chromate solution
as indicator till precipitation of red chromate is indicated.
Each ml of 0.1 N silver nitrate solution is equivalent to 0.007455 g of KCl.
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Practical pharmaceutical chemistry
Assist. Prof. Karima F. Ali