Inorganic Pharmacy -II Lab MANIK

Laboratory Manual
Course Code: PHRM 2114
Course Title: Inorganic Pharmacy-II Lab
Prepared By
Shadid Uz Zaman and Md. Imran Nur Manik
Lecturer
Department of Pharmacy
Northern University Bangladesh
Edited By
Somaia Haque Chadni
Lecturer
Department of Pharmacy
Northern University Bangladesh

2114: Inorganic Pharmacy-II Lab
Prepared By: Shadid Uz Zaman and Md. Imran Nur Manik
Edited By: Somaia Haque Chadni
Lecturer; Department of Pharmacy; Northern University Bangladesh (NUB).
Sl.
No.
Date Name of the experiment
Page
No.
01
Conversion of different water insoluble or sparingly soluble
substances into water soluble form. 01 – 07
02
Name of the Experiment: Manufacture of Aluminium Hydroxide Gel
from the supplied material. 08 – 10
03-A
Qualitative analysis (identification) of anions from inorganic salt
solutions. 11 – 15
03-B Qualitative analysis (identification) of group I & group II cations. 16 – 20
03-C Qualitative analysis (identification) of group III cation. 21 – 27
03-D Qualitative analysis (identification) of group IV & group V cations. 28 – 31
APPENDIX
How to test the Ammonium with Litmus paper?
How to Clean a Platinum Wire?
What is Qualitative Analysis?
32
SAFETY RULES FOR CHEMICAL LABORATORY 33
Version: Halim-01 01 1426
LAST UPDATE:05 05 2019

Phrm 2114: Inorganic Pharmacy -II Lab
Prepared By: Shadid Uz Zaman and Md. Imran Nur Manik Edited By: Somaia Haque Chadni Page 1
Experiment No. 01
Name of the Experiment: Conversion of different water insoluble or sparingly soluble
substances into water soluble form.
Introduction:
A solution may be defined as a homogeneous mixture of two or more components that
form a single phase.
The component that determines the phase of the solution is termed the solvent and
usually constitutes the largest proportion of the system. The other components of solution
which is dispersed as molecules or ions throughout the solvent are termed solutes.
The transfer of molecules or ions from a solid state into solution is known as dissolution.
The extent to which the dissolution proceeds under a given set of experimental conditions
is referred to as the solubility of the solute in the solvent.
Solution Process
In order to dissolve a solid, the particles of the solvent must be able to separate the
particles of the solute and occupy the intervening space. This happens when the positive
end of the solvent molecule approaches the negative end of the solute molecule. The
solute molecule is pooled in the solution when the force between the solvent solute
molecules overcomes the attractive forces between the solute molecules themselves.
Therefore, the solute will be soluble in the solvent if the solute-solvent attraction is
stronger than the solvent-solvent attraction and solute-solute attraction.
H1 H2
H3
An ionic solid dissolves in water because the number of water molecules around the
surface is greater than the number of other ions of the solid. The attraction between polar
water molecules and a charged ion enables the water molecules to pull ions away from the
crystal, a process called dissolving.
A General Solubility Rule:
Polar molecules are soluble in polar solvents e.g. C2H5OH in H2O. Thus polar compounds,
like table sugar (C12H22O11), are soluble in polar solvents and insoluble in non-polar
solvents. On the other hand non-polar molecules are soluble in non-polar solvents e.g.
CCl4 in C6H6.
Accordingly non-polar compounds, like naphthalene (C10H8), are soluble in non-polar
solvents and insoluble in polar solvents.
This it the like dissolves like rule.
Figure: A molecular view of the solution process. The
solute molecules and solvent molecules first being
spread apart, and then being mixed together. The
relative strength of forces holding solvent molecules
togetherH1, solute particles togetherH2, and the
forces between solvent and solute molecules H3 in
the solution are important in determining the
solubility. Md.
Imran
Nur
Manik

Factors Affecting Solubility of a Solute
Temperature
Generally in many cases solubility increases with the rise in temperature and decreases
with the fall of temperature but it is not necessary in all cases. However we must follow
two behaviours:
In endothermic process solubility increases with the increase in
temperature and vice versa. For example: solubility of potassium nitrate increases with
the increase in temperature.
In exothermic process solubility decrease with the increase in temperature.
For example: solubility of calcium oxide decreases with the increase in temperature.
Effect of pressure
The effect of pressure is observed only in the case of gases. An increase in pressure
increases of solubility of a gas in a liquid. For example CO2 is filled in cold drink bottles
(such as coca cola, Pepsi 7up etc.) under pressure.
Chemical natures of the solute and solvent
A polar solute will dissolve in a polar solvent but not in a non-polar solvent. The adage
"like dissolves like" is very useful. Example: Alcohol (polar substance) dissolves in water
(polar substance). Water (polar substance) does not dissolve in oil (non-polar substance)
Stirring
With liquid and solid solute, stirring brings fresh portions of the solvent in contact with
the solute, thereby increasing the rate of dissolution.
Amount of solute already dissolved
When there is little solute already in solution, dissolving takes place relatively rapidly. As
the solution approaches the point where no solute can be dissolved, dissolving takes place
more slowly.
Molecular size
The larger the molecules of the solute are, the larger is their molecular weight and their
size. If the pressure and temperature are the same than out of two solutes of the same
polarity, the one with smaller particles is usually more soluble.
Importance of Solubility in Pharmacy
Solubility is one of the important parameters to achieve desired concentration of drug in
systemic circulation for achieving required pharmacological response. Poorly water soluble
drugs often require high doses in order to reach therapeutic plasma concentrations after
oral administration.
Most of the drugs are either weakly acidic or weakly basic having poor aqueous solubility.
These poorly water soluble drugs having slow drug absorption leads to inadequate and
variable bioavailability and gastrointestinal mucosal toxicity.
For orally administered drugs solubility is the most important one rate limiting parameter
to achieve their desired concentration in systemic circulation for pharmacological
response.Solubility also plays a major role for other dosage forms like parenteral
formulations as well.
Md.
Imran
Nur
Manik

Methods for Expression of Concentration
Expression Symbol Definition
Mole Fraction X
Ratio of moles of one constituent of a solution to
the total moles of all constituents (solute & solvent)
Mole percent X×100
Moles of one constituent in 100 moles of the
solution. It is obtained by multiplying by 100X
Percent by weigh %W/W 100
SolutionofMass
SoluteofMass

Percent by volume %V/V 100
SolutionofVolume
SoluteofVolume

Percent by weight-in-volume %W/V 100
SolutionofVolume
SoluteofMass

Percent by volume-in-weight %V/W 100
SolutionofWeight
SoluteofVolume

Parts By Mass (Solid in solid) FactortionMultiplica
SolventofMass
SoluteofMass

Parts per million ppm Multiplication Factor=106
Parts per billion ppb Multiplication Factor=109
Principle:
Solubility, the phenomenon of dissolution of solute in solvent to give a homogenous
system, is one of the important parameters to achieve desired concentration of drug in
systemic circulation for desired (anticipated) pharmacological response. Low aqueous
solubility is the major problem encountered with formulation development of new
chemical entities as well as for the generic development. More than 40% new chemical
entities developed in pharmaceutical industry are practically insoluble in water. Any drug
to be absorbed must be present in the form of solution at the site of absorption. Various
techniques are used for the enhancement of the solubility of poorly soluble drugs which
include physical and chemical modifications of drug and other methods like particle size
reduction, crystal engineering, salt formation, solid dispersion, use of surfactant,
complexation, and so forth. Selection of solubility improving method depends on drug
property, site of absorption, and required dosage form characteristics.
Solubility Definitions
Description Parts of Solvent required to dissolve 1 part of Solute
Very Soluble Less than 1 part
Freely Soluble from 1 to 10 parts
Soluble from 10 to 30 parts
Sparingly Soluble from 30 to 100 parts
Slightly Soluble from 100 to 1000 parts
Very Slightly Soluble from 1000 to 10,000 parts
Practically Insoluble more than 10,000 parts
Md.
Imran
Nur
Manik

Apparatus
1. Test tubes
2. Pipette
3. Dropper
4. Bunsen Burner
5. Watch glass
Reagents
1. Solid Sample
a. Salicylic acid
b. Benzoic acid
2. Chemicals
a. Sodium Carbonate (Na2CO3)
b. Sodium Hydroxide (NaOH)
3. Distilled water
Salicylic acid
Solubility: 1 gm in 460m ml water
Conversion Reaction:
2 + Na2CO3  2 + CO2 + H2O
Procedure:
1. Take a small amount of salicylic acid (0.5 g) in a dry test tube.
2. Add a small volume of water (1 ml) and shake well. Observe that whether it is
dissolving or not.
3. Now add small amount of Sodium carbonate (Na2CO3) and shake well.
4. Observe that whether it is dissolving or not.
5. Continue the shaking for complete dissolution.
Table-1: Data for solubility of salicylic acid only:
No. of
Observations
Amount of
Salicylic
Acid Taken
(gm)
Volume of Solvent
Required to
Dissolve Salicylic
Acid (mL)
Types of Solubility
Observed
(as per solubility definition)
1
2
Md.
Imran
Nur
Manik

Table-2: Data for solubility of salicylic acid & sodium carbonate (salt):
No. of
Observations
Amount of
Salicylic &
Na2CO3 Taken
(gm)
Volume of solvent
Required to
Dissolve Salt
(mL)
Types of Solubility
Observed
1
2
Pharmaceutical uses of Salicylic Acid (Salicylic Acid Topical)
Uses
This medication is used on the skin to treat common skin and foot
(plantar) warts. Salicylic acid helps cause the wart to gradually peel off.
This medication is also used to help remove corns and calluses.
Topical salicylic acid is used to help clear and prevent pimples and skin blemishes
in people who have acne.
Topical salicylic acid is also used to treat skin conditions that involve scaling or
overgrowth of skin cells such as psoriasis (a skin disease in which red, scaly
patches form on some areas of the body), ichthyoses (inborn conditions that cause
skin dryness and scaling), dandruff, corns, calluses, and warts on the hands or
feet.
Topical salicylic acid should not be used to treat genital warts, warts on the face,
warts with hair growing from them, warts in the nose or mouth, moles, or
birthmarks.
Salicylic acid is in a class of medications called keratolytic agents.
Topical salicylic acid treats acne by reducing swelling and redness and unplugging
blocked skin pores to allow pimples to shrink.
It treats other skin conditions by softening and loosening dry, scaly, or thickened
skin so that it falls off or can be removed easily.
Market preparations of Salicylic Acid
SL. Brand Name Company
01
02
03
04
05
Md.
Imran
Nur
Manik

Benzoic acid
Solubility: 1 gm in 300 ml water
Conversion Reaction:
Procedure:
1. Take a small amount of benzoic acid (0.5 g) in a dry test tube.
2. Add a small volume of water (1 ml) and shake well. Observe that whether it is
dissolving or not.
3. Now add small amount of Sodium hydroxide (NaOH) and shake well.
4. Observe that whether it is dissolving or not.
5. Continue the shaking for complete dissolution.
Table-3: Data for solubility of benzoic acid only:
No. of
Observations
Amount of
Benzoic
Acid taken
(gm)
Volume of Solvent
Required to
Dissolve Benzoic
Acid (mL)
Types of Solubility
Observed
1
2
Table-4: Data for solubility of benzoic acid & sodium hydroxide (salt):
No. of
Observations
Amount of
Benzoic & NaOH
Taken (gm)
Volume of Solvent
Required to
Dissolve Salt (mL)
Types of Solubility
Observed
1
2
Md.
Imran
Nur
Manik

Pharmaceutical uses of Benzoic Acid
Uses
1. Benzoic acid helps prevent infection caused by bacteria.
2. Benzoic acid and salicylic acid topical (for the skin) is a combination medicine used
to treat skin irritation and inflammation caused by burns, insect bites, fungal
infections, or eczema.
3. As an inactive ingredient in the pharmaceutical industry, it is used as
a. Antimicrobial preservative,
b. Antifungal, and
c. Tablet and capsule lubricant.
4. Benzoic acid has been used in combination with salicylic acid, as in Whitfield's
ointment, for use as an antifungal for athlete's foot and ringworm.
Brand names of Benzoic Acid
SL. Brand Name Company
01
02
03
04
05
Precautions
Comments
Md.
Imran
Nur
Manik

Experiment No. 02
Name of the Experiment: Manufacture of Aluminium Hydroxide Gel from the
supplied material.
Introduction:
Aluminum is a naturally occurring mineral. Aluminum hydroxide is an antacid.
Aluminum hydroxide is used to treat heartburn, upset stomach, sour stomach, or acid
indigestion. Aluminum hydroxide is also used to reduce phosphate levels in people with
certain kidney conditions.
Principle:
Aluminum hydroxide is prepared by reacting potassium alum with sodium carbonate. Sodium
carbonate and potash alam react and produces insoluble aluminium hydroxide and other
soluble salts of sodium, potassium and aluminium etc. With the help of filter paper, this
insoluble gel like aluminium hydroxide product is collected from the upper layer. Finally
the product is flavouried, preserved and labelled suitably.
For preparing aluminium hydroxide gel, a hot solution of potash alum is added slowly to a
hot solution of sodium carbonate and not vice versa. The precipitate of aluminium
hydroxide is washed thoroughly with hot water till it is free from sulphate. The gel is then
adjusted to the required volume with distilled water.
3Na2CO3 + 2KAl(SO4)2+3H2O = 3Na2SO4+K2SO4+2Al(OH)3 +3CO2
Potash alum
[3Na2CO3 + K2SO4.Al2(SO4)3.24H2O+3H2O = 3Na2SO4+K2SO4+2AI(OH)3 +3CO2+21H2O]
If sodium carbonate solution is added to potash alum solution, then it is difficult to wash
out the sulphate completely. Due to adsorption by aluminium hydroxide, some carbonate
may be present.
Apparatus:
1. Beaker (50 mL)
2. Wash Bottle
3. Funnel
4. Filter paper
5. Measuring cylinder
6. Spatula
7. Stirrer
8. Mortar & Pestle (If powdering of the reagent is needed)
9. Balance
Reagents:
1. Na2CO3
2. Potash alum [ Alum, Potassium alum, Potassium aluminum sulfate dodecahydrate AlK(SO4)2·12H2O
or KAl(SO4)2 or AlK(SO4)2 or K2SO4.Al2(SO4)3.24H2Oor AlKO8S2]
3. Hot Water
4. Cold Water
5. Purified Water[ Distilled Water]
Note: Al(OH)3 is a very light, voluminous precipitate, and settles very slowly if at all, and is therefore difficult to
wash. It will settle less readily in a more dense liquid than in one less dense. Hence the addition of hot water greatly
facilitates the washing by decantation. However, prolonged contact with boiling water changes the constitution of
the hydroxide from Al(OH)3 to Al(OH)O[Aluminum hydroxide oxide], and the latter compounds is insoluble in acids.
Hence, the liquid in which the hydroxide is suspended must not be boiling, but only hot.
Md.
Imran
Nur
Manik

Experimental Procedure:
1. Dissolve 2.5 gm Na2CO3 in 15 mL hot distilled water in a beaker.
Note: Ensure complete solubilisation (Filter if required).
2. Dissolve 2.5 gm potash alum in 25 mL hot distilled water in a beaker.
Note: Ensure complete solubilisation (Filter if required).
3. Mix the solution of step-1 and step-2 with continuous stirring.
Note: Slowly add potash alum solution to the sodium carbonate solution.
4. Afterwards add 2 mL hot distilled water and remove gas with stirrer.
5. Dilute the solution with cold water (Q.S).
6. Filter, Collect and Wash the PPT.
7. Dry the filtrate at a temperature between 55-65°C. This will yield white, light,
amorphous, odourless and tasteless powder.
Calculation of Yield Value
Amount of Na2CO3 Taken=2.5 gm
Moles of Na2CO3 Taken= mmole23.6mole0.0236
gm106
gm2.5

Now, [From the reaction we can see that]
3 mole Na2CO3 =2 mole Al(OH)3
Thus, 3 mmole Na2CO3 =2 mmole Al(OH)3
So, 23.6 mmole Na2CO3= mmole733.51
3
23.62


Theoretical yield value=15.73 mmole
Practical amount of Al(OH)3 obtained = x gram
mmole
78
1000
mole
78
x
x


Therefore the yield value= %100
valuelTheoritica
valuePractical
 = %100
15.7378
x1000
%100
73.15
78
x1000
























Md.
Imran
Nur
Manik

Brand names of Aluminium hydroxide gel
Sl. Brand Name Company
01
02
03
04
05
Precautions:
1. Alum solution must be poured into the alkali solution, and not vice versa.
2. The injunction of Na2CO3 to the alum solution must be done slowly and with constant
stirring.
3. The liquid in which the hydroxide is suspended must not be boiling, but only hot.
4. The washed hydroxide is to be dried without much heat in order to avoid it’s becoming
hard and gritty.
Md.
Imran
Nur
Manik

Experiment No. 03-A
1. Name of the Experiment
Qualitative analysis (identification) of anions from inorganic salt solutions.
2. Principle
Salts are compounds which are formed from the reaction between acids and bases.
An inorganic salt is just a salt which has no organic part in its molecular structure.
Any salt is composed of two parts – cation and anion. The cation part usually comes from
the base and the anion part usually comes from the acid. As a result, most of the common
anions are also called acid radicals.
Some of the most common acid radicals are sulfate (SO4
2–), chloride (Cl–), bromide (Br–),
iodide (I–) and nitrate (NO3
–) which come from H2SO4, HCl, HBr, HI, HNO3 respectively.
Identification of these radicals in an unknown salt sample is done by taking advantage of
the fact that, not all salts of the radicals are highly soluble in water. For example, BaSO4
is poorly water-soluble but BaCl2 has good water solubility. So, if sulfate ion is present in
a sample, addition of BaCl2 will cause formation of precipitate.
Test Principle Reactions
Sulphate
test
principle
A. Barium nitrate (/Barium
chloride) test:
Sulphate salt of barium is
insoluble in water and forms
white precipitate, but other
salts of barium are soluble in
water.
So, if addition of BaCl2 causes
formation of precipitate, then
sulphate is present.
B. Lead acetate test:
Sulphate salt of Lead is
insoluble in water as well as in
dil HNO3 and forms white
precipitate, but other salts of
barium are soluble in water.
So, if addition of Pb(CH3COO)2
causes formation of precipitate,
then sulphate is present.
SO4
2 + BaCl2 BaSO4 + 2 Cl
(soluble) (white ppt.)
HCl is required to remove any carbonate ion
that may be present (because carbonate ion
also produce white PPT with BaCl2).
Instead of HCl, HNO3 can be used but not
H2SO4.
SO42−+ Pb(CH3COO)2→ PbSO4↓+ 2CH3COO−
White PPT
PbSO4↓+ HNO3(Dil.)→ No reaction
Md.
Imran
Nur
Manik

Halide
test
principle
Silver salts of chloride,
bromide and iodide are
insoluble in water, with
solubility decreasing from
chloride salt to iodide salt. So,
addition of silver nitrate will
cause formation of precipitate
(white/cream/yellow).
When ammonium hydroxide is
added, some of the silver ions
will form soluble
diamminesilver(I) ions.
Thus the insoluble salt will
become soluble. But silver
iodide salt is so insoluble that
even concentrated ammonium
hydroxide can’t dissolve it.
(soluble) (ppt.)
X + AgNO3 AgX + NO3

Chloride, bromide, and iodide ions will give
white, pale cream, and pale yellow PPT
will be formed.
When ammonium hydroxide is added, the
following reaction moves to the right:
AgX+2NH4OH ⇌ [Ag(NH3)2]X+2H2O
Soluble Complex
As a result, the amount of free ions in the
solution will decrease. So, the following
reaction will move to the right.
AgX Ag+ + X
(insoluble) (soluble)
As a result, the salt will be solubilized.
But, in case of silver bromide, concentrated
ammonium hydroxide is required to
significantly dissolve the salt.
And for silver iodide, it is not possible to
dissolve it even with conc. ammonium
hydroxide.
Nitrate
test
This is known as the Ring test of
Nitrate.
In this test, a brown colored
ring is formed between two
layers of acid and ferrous
sulphate solution in a test tube
if nitrate ion is present.
With H2SO4
NO3
−
+H2SO4→HSO4
−
+HNO3
6FeSO4+3H2SO4+2HNO3→3Fe2(SO4)3+4H2O+2NO
FeSO4+NO+5H2O→[Fe(NO)(H2O)5]SO4
NO3
- + 3 Fe2+ + 4 H+ NO + 3 Fe3+ + 2H2O
[Fe(H2O)6]2+SO4
2- + NO [Fe(H2O)5NO]2+SO4
2- + H2O
Nitrosoferrous sulfate
(brown)
Addition of acid (such as sulphuric acid or
hydrochloric acid) will reduce the nitrate to
NO.
This NO will then attack the ferrous ions to
form nitrosoferrous complex which resides in
the interface of the acid layer and ferrous
sulphate layer.
Acetate
ion
(Ferric
chloride
test)
When ferric chloride solution is
added to the salt solution
reddish brown colour (brick
red) forms.
[Divide the filtrate into two portions.]
To one part addition of dil. HCl
the color disappears (PPT
forms).
To second part adding water
and boiling, reddish brown PPT
forms.
This confirms the presence of
acetate ion.
The reddish color is due to the formation of
(CH3COO)3Fe.
3CH3COO− +FeCl3→(CH3COO)3Fe +3Cl−
Reddish colour
Reddish color disappears.
The reddish brown precipitate is due to the
formation of (CH3COO)(OH)2Fe.
(CH3COO)3Fe+2H2O→(CH3COO)(OH)2Fe↓(s)+
2CH3COOH Reddish brown PPT
Md.
Imran
Nur
Manik

3. Required apparatus
1. Test tubes
2. Pipette (5mL)
3. Volumetric flask
5. Beaker
6. Stirrer
7. Filter paper
8. Water bath and ice bath
4. Required reagents
1. For sulphate (SO42−) identification
a. Dilute (6M) Hydrochloric acid (HCl)
b. 0.1M Barium chloride (BaCl2) or Barium nitrate {Ba(NO3)2} solution
2. For halide (X−) ions identification
a. Dilute nitric acid (2M)
b. 5% Silver nitrate (AgNO3) solution
c. Dilute ammonium hydroxide solution (2M)
d. Concentrated ammonium hydroxide solution (~9M)
3. For nitrate ion (NO3−) identification
a. Concentrated Hydrochloric acid (37%) {/Conc. Sulphuric acid }
b. 5% Ferrous sulphate (FeSO4) solution
4. For Acetate ion (CH3COO−) identification
a. Dilute Hydrochloric acid
b. 0.2M Ferric chloride (FeCl3) solution
5. Procedure
5.1 Test for sulphate ion
A. Barium nitrate (/Barium chloride) test:
1. Take 3-4mL of the supplied salt solution in a test tube.
2. Add 2mL of dilute (6M) hydrochloric acid solution to the test tube.
CAUTION: Take special care when handling hydrochloric acid, it is extremely
corrosive. Add slowly to the side wall of the test tube.
3. Add 0.5mL of 0.1M Barium chloride or Barium nitrate solution slowly using the
pipette. Check to see if white precipitate is formed.
4. Add more Barium salt solution if precipitate doesn’t form initially or only small
precipitate is obtained initially.
B. Lead acetate test:
2. Slowly add few drops of Lead acetate solution using the pipette. Check to see if
white precipitate is formed. This is insoluble in dilute HNO3.
Md.
Imran
Nur
Manik

5.2 Test for halide ions
1. Take 2mL of the supplied salt solution in a test tube.
2. Add 2mL of dilute nitric acid to the test tube.
CAUTION: Be careful when handling nitric acid. Add slowly to the side wall of the
test tube.
3. Add 0.5mL of 5% silver nitrate solution. Check to see if White (Cl−)/ Pale
Cream(Br−)/Pale Yellow(I−) precipitate is formed.
4. Add more silver nitrate solution if precipitate doesn’t form initially or only small
precipitate is obtained initially.
5. If precipitate is formed, add dilute ammonium hydroxide solution. Check to see if
precipitate dissolves.
6. If precipitate doesn’t dissolve, add concentrated ammonium hydroxide solution.
Check to see if precipitate dissolves.
5.3 Test for nitrate ion
2. Add concentrated hydrochloric acid as such that salt solution: Conc. HCl is 2:1.
3. Cool it. (Use ice bath if required)
NOTE: You can use a water bath or ice bath for this.
4. Add 5mL of 5% ferrous sulphate solution very carefully with a pipette to the side
wall of the test tube.
5. The ferrous sulphate solution should form a layer above the acidic solution.
6. Check to see if a brown ring is formed at the interface of the two layers.
5.3 Test for acetate ion
2. Add 0.2M ferricchloride solution drop-wise.
3. Reddish color (brick red) forms.
4. (If PPT appears) Filter and divide the filtrate into two portions.
5. To one part, add dilute hydrochloride acid solution
6. Reddish color disappears.
7. To second part, add water and boil.
8. Reddish brown precipitate appears. This confirms the presence of acetate ion.
Md.
Imran
Nur
Manik

6. Inference
Name of
the test
Observation Inference
Sample 1 Sample 2 Sample 3
Sulphate
Halide
Nitrate
Acetate
7. Result
The supplied sample solution contained: ion(s)
8. Precautions
1. The test tubes must be thoroughly cleaned with DW before use.
2. The reagents must be prepared with DW.
3. The supplied sample solution must not be touched with bare hands.
4. Never taste any salt, it may be poisonous. Salts of arsenic and mercury are highly
poisonous.
Specimen: Writing inference
Confirmation test for Acid Radicals
Experiment Observations Inference
2 ml of the supplied sample solution was taken in
a test tube. Then 0.5 ml 5% AgNO3 was added to
the sample solution, followed by 2ml dil. HNO3.
White PPTwas formed.
Reaction: Cl–
(aq)+AgNO3(aq)AgCl(s)+NO3
–
White PPT
Cl–
Present
Then excess dilute ammonium [aq. NH3 (2M) ]
solution was added to it.
PPT dissolved in excess dilute (2M) NH4OH solution.
Reaction:
AgCl(s)+2NH4OH(aq) ⇌ [Ag(NH3)2]Cl(aq)+ 2H2O
Soluble complex
Md.
Imran
Nur
Manik

Experiment No. 03-B
Qualitative analysis (identification) of group I & group II cations.
2. Principle
the base and the anion part usually comes from the acid.
Based on solubility characteristics, the common inorganic cations (usually metal in
nature) are classified into five groups. Cation of each group is precipitated by a common
reagent.
So, by step-wise addition of those reagents, we can first identify which group the cation
belongs to. Then by using specific reagents we can confirm the cation.
Group Cations Specific precipitating agents
I Hg+, Pb2+, Ag+ Dilute HCl
II
Bi3+, Cu2+, Hg2+, Cd2+, Pb2+, Sn2+, Sn4+,
Sb3+, Sb5+, As3+, As5+
H2S
III Fe2+, Fe3+, Al3+, Cr3+, Zn2+, Ni2+, Co2+, Mn2+ NH4Cl + excess NH4OH(NaOH + NH4OH)
IV Ca2+, Ba2+, Sr2+ (NH4)2CO3(& Na2CO3)
V Na+, K+, Mg2+, NH4
+ -
The whole process of group identification can be shown using the following flow-chart.
Md.
Imran
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The cations of group I are separated when dilute HCl is added because of the formation of
insoluble chloride salts. Thus, to identify whether the cation belongs to group I or not, one
should add dilute HCl. If precipitate is formed, then the specific cation must be identified.
Ag+ + HCl→ AgCl↓(s)+H+
White PPT
Pb2+ + 2HCl → PbCl2↓(s)+2H+ [PPT dissolves when boiled with water, since PbCl2 is soluble in HOT water ]
White PPT
For the confirmation of the specific cation of group I following process should be utilized.
Confirmation
test
Description Reactions
Ag+ test
When NH3 solution is added to
silver chloride, the white
precipitate dissolves.
Ag++ NH4OH⇌[Ag(NH3)2]+
{AgCl(s)+2NH4OH(aq) ⇌ [Ag(NH3)2]Cl(aq)
Soluble
+2H2O}
Pb2+ test
Potassium chromate is added to
the stock solution.
A bright yellow precipitate
confirms Pb2+ ion.
PbCl2(aq)+ K2CrO4(aq)→
PbCrO4(s)↓+2KCl
Yellow PPT
If no precipitate is formed when dilute HCl is added to the stock solution, then group I
cations are not present. Then we have to add H2S which will form insoluble sulfide salts of
group II cations. H2S is not directly produced, rather thioacetamide is added which will
produce H2S in solution.
Hg2+ + H2S → HgS↓ (black) + 2H+ + 2Cl–
Cu2+ + H2S → CuS↓ (black) + 2H+ + 2Cl–
If precipitates are formed upon addition of thioacetamide, the next step is to confirm the
cation.
Confirmation
test
Description Reaction
Hg2+ test
When HNO3 is added, HgS doesn’t
dissolve, but sulfides of other cations
dissolve.
Then HCl is added (aqua regia),
which will cause formation of soluble
HgCl2 (mercuric chloride) salt.
Addition of SnCl2 will cause
formation of Hg2Cl2 and elemental Hg
both of which form precipitates.
HgS +HNO3 →No Reaction
HgS + HNO3 + 6Cl–+ 8H+  3HgCl2
+ 2NO + 3S + 4H2O
[HgS + HNO3 + HCl HgCl2 + NO + S + H2O]
Hg2Cl2 + Sn2+ 
Hg + Hg2Cl2 + Sn4+
Cu2+ test
When Sodium hydroxide is added to
a solution of cupric (Cu2+) ions, blue
gelatinous precipitate will form.
[This PPT is insoluble in excess alkali.]
Cu2+ + NaOH Cu(OH)2(s) ↓+Na+
(Blue PPT)
[Cu(OH)2+NaOH→ No Reaction ]
When aqueous ammonia is added to
a solution of cupric (Cu2+) ions, blue
precipitate will form. Excess addition
of ammonia will dissolve the blue
precipitate forming a deep blue color.
Cu2+ + NH4OH Cu(OH)2(s)↓+NH4+
(Blue PPT)
Cu(OH)2 + 4NH4OH  [Cu(NH3)4]2+
(aq) +4H2O (Deep Blue solution)
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Confirmation test Description Reaction
Pb2+
test
When 50% HNO3 solution is added PbS
dissolves.
3PbS+8HNO3→3Pb(NO3)2+4H2O
+2NO+3S
Cu2+
test
When 50% HNO3 is added, CuS dissolves. 3CuS+8HNO3→3Cu(NO3)2+4H2O
+2NO+3S
When potassium ferrocyanide is added to a
solution of cupric (Cu2+) ions, chocolate
brown (reddish brown) precipitate will form.
Cu2+ (aq) + K4[Fe(CN)6](aq) 
Cu2[Fe(CN)6]↓(s)+4K+ (aq)
(chocolate brown PPT)
When KI is added to a solution of cupric
(Cu2+) ions, white precipitate will form.
But there will be a brown coloration due to
the presence of I2.
Cu2+ (aq)+ 4KI→Cu2I2↓(s)+ I2 +4K+
(White PPT) (brown coloration)
1. Test tubes
2. Pipette (5mL)
3. Volumetric flask
5. Wash bottle
6. Beaker
7. Stirrer
8. Filter paper
10. H2S generating apparatus
A. For group I cations
a. Dilute hydrochloric acid (6M)
b. Ammonia solution (9M)
c. Potassium chromate (1M)
B. For group II cations
a. Ferrous sulfide (FeS)
b. Concentrated sulfuricacid
c. Nitric acid (6M)
d. Stannous chloride (0.1M)
e. Concentrated Ammonia solution (9M)
f. Dilute Ammonia solution (2M)
g. Dilute Sodium hydroxide solution (2M)
h. Potassium ferrocyanide, K4[Fe(CN)6]
i. Potassium Iodide, KI
5. Procedure
5.1 Test for group I cations
3. Check to see if precipitate is formed.
4. If precipitate is formed, go for confirmation tests.
5. If precipitate is not formed, go for test of group II cations.
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5.1.1 Confirmation test of Ag+
1. Take the solution containing the precipitate.(Obtained earlier from step 5.1)
2. Add 2-3mL of ammonia solution. If the precipitate is dissolved, then Ag+ is
confirmed.
5.1.2 Confirmation test of Pb2+
1. Take 2mL of the stock solution in a test tube.
2. Add 0.5mL of 1M potassium chromate solution. If precipitate starts to form, add
more 1M potassium chromate (up to 2mL).
3. Bright yellow precipitate confirms Pb2+ ions.
5.2 Test for group II cations
4. From 5.1, (Solution containing added HCl) take the solution and add H2S from the
Kipp’s apparatus.
5. Check to see if any precipitate is formed. If black precipitate is formed, first go for
Hg2+ or Cu2+ confirmation test.
5.2.1 Confirmation test of Hg2+
1. First separate the black precipitate from the solution by filtration.
2. Take the precipitate in another test tube and add 1mL distilled water.
3. Then add 1mL 6M Nitric acid & Shake vigorously.
4. If precipitate doesn’t dissolve, it is indicative of the presence of Hg2+.
5. Add 1mL 6M HCl and 1mL 6M HNO3 to the test tube. Heat in a boiling water bath.
6. Cool the solution in an ice bath.
7. Add 0.5mL of 0.1M SnCl2 drop wise.
8. Grey to black precipitates confirms the presence of Hg2+.
5.2.2 Confirmation test of Cu2+
With NaOH
2. Add 0.5 mL dilute NaOH solution drop wise.
3. Check to see if light blue gelatinous precipitate is formed.
4. If precipitate is observed, then add more NaOH solution. Precipitate insoluble in
excess alkali.
With NH4OH
1. Again take 2mL of the stock solution in a test tube
2. Add 0.5 mL dilute ammonia solution drop wise.
3. Check to see if light blue precipitate is formed.
4. If precipitate is observed, then add more ammonia solution. A deep blue colored
solution will be formed. This confirms the presence of Cu2+
.
With potassium ferrocyanide
2. Add 0.5 mL (1-2 drops) potassium ferrocyanide.
3. Check to see if chocolate brown( reddish brown) precipitate is formed.
4. If precipitate is observed, then this confirms the presence of Cu2+
.
Md.
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With KI
2. Add 0.5 mL (1-2 drops) of KI solution.
3. Check to see if white precipitate is formed. The PPT would have brown
coloration due to the presence of I2.
4. This also confirms the presence of Cu2+
.
6. Inference
Group
test
Confirmation
test
Observation
Inference
Group I
-
Hg+ test
Ag+ test
Pb2+ test
Group II
-
Hg2+ test
Cu2+ test
7. Result
8. Precautions
Confirmatory test for Group II cations
2ml stock solution was taken in a test tube and
2M NaOH solution was added to it.
Blue gelatinous PPTwas formed.
Reaction: Cu2+
(aq)+2OH–
(aq)Cu(OH)2(s)
Blue PPT Cu2+
Present
Then excess 2M NaOH solution was added to it. The Precipitate was insoluble in excess NaOH.
Again, 2ml stock solution was again taken in a
test tube and ammonia solution was added to it.
Blue gelatinous PPT formed.
Reaction:Cu2+
(aq)+2NH3(aq)+2H2O
Cu(OH)2(s)+ 2NH4
+
(aq)
Blue gelatinous PPT
Cu2+
ConfirmedAfterwards, excess ammonia solution (9M) was
added to it.
Precipitate was soluble in excess 9M NH4OH.
Excess aq. NH3
Reaction: Cu(OH)2 (s)+ 4NH3(aq) 
[Cu(NH3)4]2+
(aq)+ 2OH–
(aq)
Tetraamminecopper(II) ion
Md.
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Experiment No. 03-C
Qualitative analysis (identification) of group III cations.
2. Principle
nature) are classified into five groups. Cations of each group is precipitated a common
reagent. So, by step-wise addition of those reagents, we can first identify which group the
cation belongs to. Then by using specific reagents we can confirm the cation.
The common inorganic cations can be classified into five groups based on solubility
characteristics. Cations of each class are precipitated by a specific reagent.
II
Bi3+, Cu2+, Hg2+, Cd2+, Pb2+, Sn2+, Sn4+,
Sb3+, Sb5+, As3+, As5+
H2S
III Fe2+, Fe3+, Al3+, Cr3+, Zn2+, Ni2+, Co2+, Mn2+ NH4Cl + excess NH4OH(NaOH + NH4OH)
IV Ca2+, Ba2+, Sr2+ (NH4)2CO3(& Na2CO3)
+ -
Cations of group III are not precipitated by dilute HCl or H2S, but they can be precipitated
from a salt solution by the addition of NH4Cl in the presence of excess NH4OH.
With this reagent, hydroxide or sulphide salts of cation will be precipitated. If precipitate
is formed, then confirmation tests should be performed to identify specific cation.
Confirmation test Description Reactions
Fe2+
test
If the precipitate from the above reaction is greenish, then it may be ferrous.
If dilute sodium hydroxide is added to a solution
containing ferrous ions (Fe2+), a green gelatinous
precipitate of iron (II) hydroxide will form.
Excess addition of sodium hydroxide will not
dissolve the precipitate.
Fe2+(aq)+2OH–(aq)
Fe(OH)2(s)↓
Green precipitate
[FeSO4 + 2NaOH  Fe(OH)2 +
Na2SO4]
Fe(OH)2+ NaOH→NR
Excess
When aqueous ammonia (dil. /2M) is added to a
solution of ferrous (Fe2+) ions, a green gelatinous
precipitate will form.
Excess addition of ammonia will not dissolve the
precipitate.
Fe2+(aq)+2NH3(aq)
Fe(OH)2(s)↓ +2NH4
+(aq)
Green precipitate
[FeSO4 + 2NH4OH → Fe(OH)2 +
(NH4)2SO4 ]
Fe(OH)2+ NH4OH→NR
Excess
To confirm, potassium ferricyanide is added to the
stock solution, a deep blue precipitate is obtained.
Fe2+ + K3[Fe(CN)6] 
KFe[Fe(CN)6]↓ (blue PPT) +2K+
Md.
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Fe3+
test
If the precipitate from the above reaction is brownish, then it may be ferric.
If dilute sodium hydroxide is added to a solution
containing ferric ions (Fe3+), a brownish red
gelatinous precipitate of iron (III) hydroxide will
form.
Excess addition of sodium hydroxide will not
dissolve the precipitate.
[But it dissolves in HCl]
Fe3+(aq)+3OH–(aq)
Fe(OH)3(s)↓
Reddish brown precipitate
[FeCl3 + 3NaOH  Fe(OH)3 + 3NaCl]
Fe(OH)3+ NaOH→NR
Excess
[Fe(OH)3+3HCl→FeCl3+3H2O]
solution of ferric (Fe3+) ions, a brownish red
gelatinous precipitate will form.
precipitate.
Fe3+(aq)+3NH4+ (aq)
Fe(OH)3↓(s) +3NH4
+(aq)
Reddish brown precipitate
[FeCl3 + 3NH4OH → Fe(OH)3 + 3NH4Cl ]
Fe(OH)3+ NH4OH→NR
Excess
To confirm, potassium ferricyanide is added to the
stock solution, a red (greenish) precipitate is
obtained.
Fe3+ + K3[Fe(CN)6] 
F[Fe(CN)6]↓ +3K+
(Deep brown PPT)
When potassium ferrocyanide is added to a
solution of cupric (Fe3+) ions, deep blue (Prussian
blue) precipitate of ferric ferrrocyanide will form.
4Fe3+ (aq) + 3K4[Fe(CN)6](aq)
Fe4[Fe(CN)6]3↓+12K+
(Prussian blue)
In another test tube, addition of potassium
thiocyanate to stock solution will produce blood
red solution, due to the formation of ferric
sulphocyanide.
[Fe(H2O)6]3++ 3KCNS(aq)
[Fe(H2O)3(CNS)3]2+ +3K++H2O
Or, Fe3+ + 3KSCN →
Fe(SCN)3 + 3K+
(Blood-red complex)
Zn2+
test
When dilute sodium hydroxide is added to a
solution containing zinc ions (Zn2+), a gelatinous
white precipitate is formed.
Zn2++ 2NaOH
Zn(OH)2↓(s)+2Na+(aq)
White PPT
Addition of concentrated sodium hydroxide (excess
dil. NaOH) produces a soluble complex resulting in a
clear solution.
Or, Zn(OH)2 + 2NaOH  Na2ZnO2 +2H2O
Soluble Complex
Zn(OH)2 + 2NaOH 
Na2[Zn(OH)4](Sodium Zincate.)
Addition of H2S in this solution results in the
formation of a white PPT due to the formation of
ZnS.
Na2ZnO2 +2H2S→
ZnS↓(s)+Na2S+2H2O
White PPT
When ammonia (dil./2M) is added to a solution of
zinc (Zn2+) ions, white gelatinous precipitate will
form.
Zn2+(aq)+2NH3(aq)+2H2O
Zn(OH)2(s)↓+ 2NH4
+(aq)
White gelatinous PPT
Excess addition of ammonia will dissolve the
precipitate forming a water soluble colourless
tetraaminezinc ion.
Zn (OH)2 (s)+ 4NH3(aq) 
[Zn(NH3)4]2+ (aq)+ 2OH–(aq)
Tetraamminezinc(II) ion
+4H2O
Md.
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Manik

1. Test tubes
2. Pipette (5mL)
3. Volumetric flask
5. Wash bottle
6. Beaker
7. Stirrer
8. Filter paper
1. For group III cations
a. Ammonium chloride
b. Ammonia solution (9M)
c. Ammonia solution (2M)
d. Potassium ferricyanide (K3[Fe(CN)6]) [250 mg in 10 ml]
e. Potassium ferrocyanide (K4[Fe(CN)6]) [250 mg in 10 ml]
f. Potassium thiocyanate [250 mg in 10 ml]
g. Sodium hydroxide (2M)
h. Sodium hydroxide (10M)
A. 6M HCl { To test for the Gr. I cations}
Al3+
test
When dilute sodium hydroxide is added to a
solution containing aluminium ions (Al3+), a white
gelatinous precipitate is formed.
Al3++ 3NaOH–→
Al(OH)3 ↓(s)+3Na+
(White PPT)
Addition of concentrated sodium hydroxide (excess
dil. NaOH) produces a soluble complex resulting in a
clear solution.
[Al(OH)3 (s)+ 3NaOH(Conc.)→ Na[Al(OH)4](aq)]
NaAlO2 (Sodium-meta-Aluminate)
Al(OH)3 (s)+ OH– (aq)→
[Al(OH)4]–(aq)
Tetrahydroxoaluminate ion
Or, Al(OH)3 (s)+ NaOH→
NaAlO2 + 2H2O [1000°C]
Soluble complex
Addition of NH4Cl in this clear solution followed by
heating, results in the reformation of white PPT due
to the formation of Al(OH)3.
NaAlO2 + NH4Cl+H2O  ΔHeating
Al(OH)3 ↓(s)+NaCl(aq)+NH3↑(g)
(White PPT.)
solution of aluminium (Al3+) ions, white gelatinous
precipitate will form.
precipitate.
Al3+(aq)+3NH3(aq)+3H2O
Al(OH)3↓(s)+ 3NH4
+(aq)
White gelatinous PPT
Al(OH)3 (aq)+ NH4OH→ NR
Excess
Md.
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5. Procedure
5.1 Test for group III cations
4. If no precipitate is formed, add H2S in the test tube. Check again to see if
precipitate is formed.
5. If no precipitate is formed above, add 2mL concentrated NH4OH and 1gm of NH4Cl.
Check to see if precipitate is formed.
6. If precipitate is formed then go for confirmation tests of Group III cations. If no
precipitate is formed, go for ‘Test for group IV cations’.
5.1.1 Confirmation test of Fe2+cation
With NaOH
3. Check to see if light green gelatinous precipitate is formed.
excess alkali.
With NH4OH
3. Check to see if green gelatinous precipitate is formed.
4. If precipitate is observed, then add more ammonia solution. Precipitate insoluble in
excess aqueous ammonia.
To further confirm the presence of Fe2+
With Potassium ferricyanide (K3[Fe(CN)6])
1. Take 2ml of the stock the solution in a test tube.
2. Add 2mL potassium ferricyanide(just one drop) solution. Check to see if precipitate
is formed.
3. Light blue precipitate confirms the presence of Fe2+.
5.1.2 Confirmation test of Fe3+ cation
With NaOH
3. Check to see if light brownish red gelatinous precipitate is formed.
excess alkali.
Md.
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With NH4OH
3. Check to see if brownish red gelatinous precipitate is formed.
4. If precipitate is observed, then add more ammonia solution. Precipitate insoluble in
excess aqueous ammonia.
To further confirm the presence of Fe3+
With Potassium ferricyanide (K3[Fe(CN)6])
2. Add 2mL potassium ferricyanide(just one drop) solution. Check to see if precipitate
is formed.
3. Deep Brown (Green in our lab) colour indicates the presence of Fe3+.
With Potassium ferrocyanide (K4[Fe(CN)6])
2. Add 2mL potassium ferrocyanide(just one drop) solution. Check to see if precipitate
is formed.
3. Deep blue (Prussian blue) colour indicates the presence of Fe3+.
With Potassium thiocyanate (KSCN)
1. To confirm Fe3+, you can also use potassium thiocyanate.
2. Take 2ml of the stock the solution in a test tube and add 2mL of potassium
thiocyanate(just one drop) solution to stock solution, if the colour change to blood
red then Fe3+ is confirmed.
5.1.3. Confirmation test of Zn2+
With NaOH
2. First add 2M sodium hydroxide drop wise very slowly.
3. Check to see if white precipitate has been formed. If precipitate is formed, add
concentrated NaOH (10M) {excess 2M NaOH} to check if the precipitate dissolves.
4. If the precipitate dissolves, then Zn2+ is present.
5. Take this clear solution and add H2S to it.
6. White PPT is formed due to the formation of ZnS further confirms the presence of
Zn2+.
Md.
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With NH4OH
1. Again take 2mL of the stock solution in a test tube.
2. Now add ammonia (dil. /2M) drop wise very slowly.
concentrated ammonia (9M) {excess 2M ammonia} to check if the precipitate
dissolves.
4. If the precipitate dissolves, then Zn2+ is further confirmed.
(10dilution may be needed for the appearance and disappearance of PPT)
5.1.4 Confirmation test of Al3+
With NaOH
2. First add 2M sodium hydroxide drop wise very slowly.
3. Check to see if white gelatinous precipitate has been formed. If precipitate is
formed, add concentrated NaOH (10M) {excess 2M NaOH} to check if the precipitate
dissolves.
4. If the precipitate dissolves, then Al3+ is confirmed.
5. Now add NH4Cl in this clear solution and heat.
6. White PPT is formed due to the formation of Al(OH)3 which confirms the presence of
Al3+.
With NH4OH
1. Again take 2mL of the stock solution in a test tube.
2. Now add ammonia (dil. /2M) drop wise very slowly.
concentrated ammonia (9M) {excess 2M ammonia} to check if the precipitate
dissolves.
4. If the precipitate does not dissolves, then Al3+ is further confirmed.
6. Inference
Group
test
Confirmation
test
Group
III
-
Fe2+ test
Fe3+ test
Zn2+ test
Al3+ test
Md.
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7. Result
8. Precautions
This experiment requires the use of hazardous components and/or has the potential for
hazardous reactions.
1. Potassium ferricyanide, potassium ferrocyanide, and potassium thiocyanate are
dangerous if heated or in contact with concentrated acids since toxic hydrogen
cyanide gas may be liberated.
2. Potassium thiocyanate is moderately toxic by ingestion.
3. Potassium ferricyanide, potassium ferrocyanide, and ferrous sulfate are slightly toxic by
ingestion.
4. Iron(II) sulfate is corrosive to skin, eyes, and mucous membranes.
5. Iron(III) chloride may be a skin and tissue irritant.
6. Avoid body contact with all chemicals.
7. Wear i. Chemical splash goggles,
ii. Chemical-resistant gloves, and
iii. Chemical-resistant apron.
8. Wash hands thoroughly with soap and water before leaving the laboratory.
Analysis for Group III cations
2ml stock solution was taken in a test tube and
2M NaOH solution was added to it.
Green gelatinous PPTwas formed.
Reaction:
Fe2+
(aq)+2OH–
(aq)Fe(OH)2(s)
Green jelly like PPT
Fe2+
Present
Then excess 2M NaOH solution was added to it. Precipitate was insoluble in excess NaOH.
Again 2ml stock solution was taken in a test tube
and ammonia solution was added to it.
Green gelatinous PPT formed.
Reaction: Fe2+
(aq)+2NH3(aq)+2H2O
Fe(OH)2(s)+ 2NH4
+
(aq)
Green gelatinous PPT
Fe2+
Present
Afterwards, excess ammonia solution was added
to it.
Precipitate was insoluble in excess aqueous ammonia
solution.
Again 2mL of the stock solution was taken in a
test tube and one drop potassium ferricyanide
solution was added to it.
Light blue precipitate was formed.
Fe2+
+ K3[Fe(CN)6] KFe[Fe(CN)6] +3K+
(blue PPT)
Fe2+
Confirmed
Md.
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Experiment No. 03-D
Qualitative analysis (identification) of group IV & group V cations.
2. Principle
nature) are classified into five groups. Cations of each group is precipitated a common
reagent. So, by step-wise addition of those reagents, we can first identify which group the
cation belongs to. Then by using specific reagents we can confirm the cation.
The common inorganic cations can be classified into five groups based on solubility
characteristics. Cations of each class are precipitated by a specific reagent.
II Bi3+, Cu2+, Hg2+, Cd2+, Pb2+, Sn2+, Sn4+,
Sb3+, Sb5+, As3+, As5+
H2S
III Fe2+, Fe3+, Al3+, Cr3+, Zn2+, Ni2+, Co2+, Mn2+ NH4Cl + excess NH4OH (NaOH + NH4OH)
IV Ca2+, Ba2+, Sr2+ (NH4)2CO3 (& Na2CO3)
+ -
Cations of group IV are not precipitated by dilute HCl or H2S, as well as by the addition of
NH4Cl in the presence of excess NH4OH.
If no precipitate is formed with the addition of NH4OH and NH4Cl to the stock solution,
then (NH4)2CO3 is added to it. If precipitate is formed, then group IV cations are present.
When (NH4)2CO3 is added to a salt solution containing NH4Cl and NH4OH, the carbonates
of Ba2+, Sr2+ and Ca2+ are precipitated. Ca2+ +(NH4)2CO3→ CaCO3↓(s)+2NH4
+
White PPT
In that case, confirmation tests are carried out as follows:
Confirmation
test
Description Reaction
Ca2+ test
When (NH4)2CO3 is added to a salt
solution containing NH4Cl and
NH4OH, the carbonate of Ca2+ is
precipitated.
The white PPT of CaCO3 dissolves in
hot acetic acid.
Ca2+ +(NH4)2CO3→CaCO3↓(s)+2NH4
+
White PPT
CaCO3+2CH3COOH→(CH3COO)2Ca
+ H2O + CO2↑ Soluble
Md.
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Solution of Ca ions gives Yellow ppt
with K2CrO4 which is also soluble in
acetic acid.
Ca++ K2CrO4 → CaCrO4↓(s)+ 2K+
Yellow PPT
When the sample salt (solid form or solution) is heated in Bunsen burner,
a brick-red flame confirms presence of Ca2+.
NH4+ test
If Nessler’s reagent is added to a
solution containing ammonium ions
(NH4
+), a brown precipitate is
formed.
2K2[HgI4]+NH4+3KOH →
HgO·Hg(NH2)I(s)↓+7KI− +2H2O
(Yellow or brown)
When sodium hydroxide is added to
a solution containing ammonium
ions (NH4
+), and the mixture is
heated, a colourless gas (Characteristic
ammoniacal smell) will be produced
which turns moist red litmus paper
blue.
NH4
+ (aq)+ NaOH 
NH3(g)↑+H2O+Na+
Ammonia
Bringing a glass rod moistened with
concentrated HCl in contact with the
evolved gas produces dense white
fumes due to the formation of
NH4Cl.
NH3 + HCl (Conc.)→NH4Cl(g)↑
White fumes
1. Test tubes
2. Pipette (5mL)
3. Volumetric flask
5. Wash bottle
6. Beaker
7. Stirrer
8. Filter paper
A. 6M HCl
B. 9M NH4OH
C. NH4Cl
1. For group IV cations
a) Ammonium carbonate (5% solution)
b) Potassium chromate
c) Acetic acid
2. For group V cations
a) Litmus paper (red)
b) Conc. hydrochloric acid
c) Nessler's Reagent.
[a 0.09 mol/L solution of potassium tetraiodomercurate(II) (K2[HgI4]) in 2.5 mol/L potassium hydroxide]
Md.
Imran
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Manik

5. Procedure
5.1 Test for group IVcations
4. If no precipitate is formed, add H2S in the test tube. Check again to see if
precipitate is formed.
5. If no precipitate is formed above, add 2mL concentrated NH4OH and 1gm of NH4Cl.
Check to see if precipitate is formed. If precipitate is formed then go for
confirmation tests.
6. If no precipitate is formed, then add (NH4)2CO3.If precipitate is formed, then go for
‘Test for group IV cations’.
7. If no precipitate is formed, then go for ‘Test for group V cations’.
5.2 Test for group IVcations
1. From 5.1, take the solution and add 5% ammonium carbonate solution.
2. Check to see if any precipitate is formed. If precipitate is formed,
go for confirmation test.
5.2.1 Confirmation test of Ca2+
With (NH4)2CO3
1. Take 2mL Sample solution in a test tube.
2. Add 5% ammonium carbonate solution into the test tube.
3. Check to see if any white precipitate is formed.
4. If precipitate is formed then add acetic acid into the PPT.
5. The PPT gets solubilised and confirms the presence of Ca2+ in the sample.
With K2CrO4
1. Take 2mL Sample solution in a test tube.
2. Add Potassium chromate solution into the test tube.
3. Check to see if any yellow precipitate is formed.
4. If precipitate is formed then add acetic acid into the PPT.
5. The PPT gets solubilised and confirms the presence of Ca2+ in the sample.
With Flame Test
1. Take a platinum or nichrome wire and dip it in concentrated hydrochloric acid.
2. Then heat the wire in Bunsen burner until flame is no longer visible.
3. Then take the wire and dip it in the sample salt solution. Conversely, you can dip it
in a solid sample of the salt.
4. Place the wire in the Bunsen burner. Observe the flame colour.
5. Brick-red flame confirms the presence of Ca2+ in the sample.
Md.
Imran
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Manik

5.2.2 Confirmation test of NH4
+
With Nessler's Reagent
2. Add Nessler's reagent drop wise.
3. Check to see if brown precipitate has been formed. If precipitate is not forming or
disappears as quickly as it forms then add more Nessler’s reagent until the
precipitate forms.
4. Add more Nessler’s reagent to check if the precipitate dissolves.
5. If the precipitate does not dissolve, then NH4
+is confirmed.
With NaOH and Litmus paper
2. Add 3ml NaOH solution to it.
3. Then heat the solution carefully in Bunsen burner until a colourless gas evolves.
4. The gas would have Characteristic ammoniacal smell.
5. Place a moist red litmus paper in the mouth of the test tube.
6. Check to see if the red litmus paper turns into blue.
7. If the colourless gas turns the red litmus paper into blue then NH4
+is further
confirmed.
With NaOH and HCl
2. Add 3ml NaOH solution to it.
3. Then heat the solution carefully in Bunsen burner until a colourless gas evolves.
4. The gas would have Characteristic ammoniacal smell.
5. Place a glass rod moistened with Conc. HCl.
6. The gas gives white fumes and NH4
+is further confirmed.
6. Inference
Group
test
Confirmation
test
Group
IV
-
Ca2+ test
Group
V
-
NH4
+ test
7. Result
8. Precautions
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Imran
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Manik

How to test the Ammonium with Litmus paper?
Hold the tube containing the ammonium nitrate solution with a test-tube
holder. Gently warm the tube along its sides using a back-and-fort
motion through a burner flame. Do not allow the solution to boil.
CAUTION: At all times, make sure that the opening of the tube is
pointed away from other people. Hold a moistened piece of red litmus
paper near the mouth of the test tube, as shown in Figure. The test will
be spoiled if the solution contacts the litmus paper. Record the changes
you observe. Fan the vapours coming out of the tube toward your nose
with your hand. Cautiously sniff the vapours. Record your observations.
How to Clean a Platinum Wire?
Bend a small loop at the end of the platinum wire and dip the looped end in concentrated hydrochloric acid taken
in a test tube. Now bring it to the oxidising (non-luminous) flame of the burner. A change in colour of the flame
indicates that the wire is not clean and needs the above operation. The operation should be repeated till the wire is
clean. Platinum wire should always be cleaned after the flame test is performed.
Note : It should be noted that persistent golden yellow colour of sodium flame conceals all other flames, therefore,
flame has to be observed carefully in presence of sodium.
What is Qualitative Analysis?
Qualitative analysis is a method of Analytical chemistry that deals with the determination of elemental
composition of inorganic salts. It is mainly concerned with the detection of ions in an aqueous solution of the
salt.
The common procedure for testing any unknown sample is to make its solution and test this solution with
various reagents for the ions present in it. Testing with various reagents gives characteristic reaction of certain
ions, which may be a colour change, a solid formation or any other visible changes. There are separate
procedures for detecting cations and anions, called the Cation Analysis and Anion Analysis
Md.
Imran
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Manik

1. Never work alone in the laboratory.
2. Never attempt unauthorised experiment.
3. Never taste any chemical. Should a chemical enter the mouth, rinse out with water at
once and report the matter to the teacher.
4. Never inhale poisonous gases or volatile liquids such as chloroform, carbon
tetra-chloride etc. Smell them carefully.
5. When heating any substance in a test tube, be careful to point the mouth of the test tube
away from yourself and your neighbours.
6. Never add more than the quantity of the reagent you have been told to add.
7. When mixing acid with water, add acid slowly to water and not water to acid.
8. Do not pour hot conc. acid directly into the sink. Allow it to cool and then pour carefully
while the water tap is kept open.
9. Always allow a hot test tube to cool before replacing it in the rack, if the rack is wet.
10. Be cautious with heated glass.
11. Never heat a funnel or a watch glass over a flame.
12. Never use cracked or broken glass ware.
13. Reagents once taken out of the bottle should not be poured back.
14. Replace the stoppers of the reagent bottles immediately after use. Do not interchange
them.
15. The test tube must be kept at eye-level and away from you when making an
observation.
16. Open the water taps gently and close when not in use. Close the gap taps immediately
after use.
17. At the end of the experiment wash the used glass ware and arrange them in order.
18. Wash your hands before leaving the lab.
Visit Website: Essential Pharma Documents
ISC Practical Chemistry Volume 1 for Class XI - S P Sharma & Dr. Ajaya Baboo
Chemistry-Qualitative-Analysis
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Imran
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Manik

Inorganic Pharmacy -II Lab MANIK

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