Saratha International Sr.Sec.
School
Kodiveri , Gobi-Sathy Main Road Gobichettipalayam
CHEMISTRY-043
[2023-2024]
PROJECT REPORT ON
MEASURING THE SOLUBILITY OF
SATURATED SOLUTIONS
PREPARED BY :
STANDARD : XII
ROLL NO :
Index
1. Certificate
2. Acknowledgement
3. Objective
4. Introduction
5. Basic concepts
6. Materials and Equipment
7. Experimental Procedure
8. Observation
9. Conclusion
10. Result
11. Precautions
12. Bibliography
CERTIFICATE
This is to certify that the Project titled 'Saturated
solutions: Measuring Solubility' was completed under
my guidance and supervision by Roll
No. a student of XII SCI, Faith
Academy within the stipulated time as prescribed by
the CBSE.
Mrs. Sasheela Jose Head, Department of Chemistry
Faith Academy New Delhi
ACKNOWLEDGEMENTS
I gratefully acknowledge my sincere thanks to our
respected chemistry teacher Mrs.Sasheela Jose for
her remarkable, valuable guidance and supervision
throughout the project work. I ' m also most indebted
to Mrs.Rao for her encouragement, help, suggestion
and readily helpful service in performing the
experiment.
Parichay Saxena
Roll NO :
Objective:
The goal of this project is to measure the
solubilities of some common chemicals:
• Table salt (NaCl)
• Epsom salts (MgSO4)
• sugar (sucrose, C12H22O11).
Introduction
A good part of the substances we deal with in daily life, such as
milk, gasoline, shampoo, wood, steel and air are mixtures. When
the mixture is homogenous, that is to say, when its components
are intermingled evenly, it is called a solution. There are various
types of solutions, and these can be categorized by state (gas,
liquid, or solid).
The chart below gives some examples of solutions in different
states. Many essential chemical reactions and natural processes
occur in liquid solutions, particularly those containing water
(aqueous solutions) because so many things dissolve in water. In
fact, water is sometimes referred to as the universal solvent. The
electrical charges in water molecules help dissolve different kinds
of substances. Solutions form when the force of attraction
between solute and solvent is greater than the force of attraction
between the particles in the solute.
Two examples of such important processes are the uptake of
nutrients by plants, and the chemical weathering of minerals.
Chemical weathering begins to take place when carbon dioxide in
the air dissolves in rainwater. A solution called carbonic acid is
formed. The process is then completed as the acidic water seeps
into rocks and dissolves underground limestone deposits.
Sometimes, the dissolving of soluble minerals in rocks can even
lead to the formation of caves.
Types of Solutions
State of State of State of
Solute Solvent Solution
Air, natural gas gas gas gas
Alcohol in water,
liquid liquid liquid
antifreeze
Brass, steel solid solid solid
Carbonated water,
gas liquid liquid
soda
Sea water, sugar solid liquid liquid
solution 1
Hydrogen in platinum gas solid solid
If one takes a moment to consider aqueous solutions, one quickly
observes that they exhibit many interesting properties. For
example, the tap water in your kitchen sink does not freeze at
exactly 0°C. This is because tap water is not pure water; it
contains dissolved solutes. Some tap water, commonly known as
hard water, contains mineral solutes such as calcium carbonate,
magnesium sulfate, calcium chloride, and iron sulfate. Another
interesting solution property is exhibited with salt and ice.
Another example comes from the fact that salt is spread on ice
collected on roads in winters. When the ice begins to melt, the
salt dissolves in the water and forms salt water. The reason is
that with the addition of salt the melting point of water increases
and as a result the snow melts away faster.
Even some organisms have evolved to survive freezing water
temperatures with natural "antifreeze." Certain arctic fish have
blood containing a high concentration of a specific protein. This
protein behaves like a solute in a solution and lowers the freezing
point of the blood. Going to the other end of the spectrum, one
can also observe that the boiling point of a solution is affected by
the addition of a solute. These two properties, namely freezing-
point depression and boiling-point elevation, are called colligative
properties (properties that depend on the number of molecules,
but not on their chemical nature).
Removing snow from blocked
roads. Before manually removing
it, salt is spread on the snow cover
to ease the job.
Basic Concepts
A saturated solution is a mixture in which no more solute can be
practically dissolved in a solvent at a given temperature. It is said
practical because theoretically infinite amount of solute can be
added to a solvent, but after a certain limit the earlier dissolved
solute particles start rearranging and come out at a constant
rate. Hence overall it appears that no solute is dissolved after a
given amount of solute is dissolved. This is known as a saturated
solution.
In an unsaturated solution, if solute is dissolved in a solvent the
solute particles dissociate and mix with the solvent without the
re-arrangement of earlier dissolved solute particles.
Solubility depends on various factors like the Ksp of the salt,
bond strength between the cation and anion, covalency of the
bond, extent of inter and intramolecular hydrogen bonding,
polarity, dipole moment etc. Out of these the concepts of H-
bonding, covalency, ionic bond strength and polarity play a major
role if water is taken as a solvent.
Also physical conditions like temperature and pressure also play
very important roles as they affect the kinetic energy of the
molecules.
Materials and Equipment
To do this experiment following materials and
equipment are required:
• Distilled water
• Metric liquid measuring cup (or graduated
cylinder)
• Three clean glass jars or beakers
• Non-iodized table salt (NaCl)
• Epsom salts (MgSO4)
• Sugar (sucrose, C12H22O11)
• Disposable plastic spoons
• Thermometer
• Three shallow plates or saucers
• Oven
• Electronic kitchen balance (accurate to 0.1 g)
Experimental Procedure
Determining Solubility
1. Measure 100 mL of distilled water and pour into a clean, empty
beaker or jar.
2. Use the kitchen balance to weigh out the suggested amount (see
below) of the solute to be tested.
a. 50 g Non-iodized table salt (NaCl)
b. 50 g Epsom salts (MgSO4)
c. 250 g Sugar (sucrose, C12H22O11)
3. Add a small amount of the solute to the water and stir with a
clean disposable spoon until dissolved.
4. Repeat this process, always adding a small amount until the
solute will no longer dissolve.
5. Weigh the amount of solute remaining to determine how much
was added to the solution.
6. Try and add more solute at the same temperature and observe
changes if any.
7. Now heat the solutions and add more solute to the
solutions.
Observations:
Salt Amount of salt Moles dissolved
dissolved in 100mL
water to make
saturated solution.
NaCl (Non-iodized 36.8 grams 0.7
common salt)
MgSO4 32.7 grams 0.255
C12H22O11 (sucrose) 51.3 grams 0.15
Adding more solute at the same temperature to the saturated
solutions yielded no significant changes in NaCl and Epsom salt.
However at all temperatures the saturation point of sucrose
could not be obtained exactly as due to the large size of the
molecule the solution became thick and refraction was more
prominent. Neglecting this observation in the room for error, the
experiments agreed with the theory.
Adding more solute to heated solutions increased the solubility in
all the 3 cases. The largest increase was shown by NaCl,
followed by Epsom salt and sucrose. These facts too agreed with
the theory as at high temperatures the kinetic energy of
molecules increases and the collisions are more effective.
Conclusion:
The solubility of NaCl is the highest as it an ionic salt and easily
dissociates in water. Also since the size of both the cation and
anion are small, the collisions are more and hence probability of
dissociation is high. The solubility of MgSO 4 is also high as it is
also an ionic salt, but due to a larger anion, collisions are not
very effective. The solubility of C12H22O11 is the least as it a very
large molecule due to which hydrogen bonding with the water
molecules is not very effective. Also due to the large number of
carbon and oxygen atoms, inter molecular H-bonding is more
dominant than intramolecular H-bonding.
Solution of NaCl
Solution of sucrose MgSO4 solution
Precautions:
1. While adding the solute to the solvent, the solution should be stirred
slowly so as to avoid the formation of any globules.
2. Stirring should not be vigorous as the kinetic energy of the molecules
might change due to which solubility can increase.
3. While stirring, contact with the walls of the container should be
avoided as with every collision, an impulse is generated which makes
the dissolved solute particles rearrange themselves. As a result
solubility can decrease.
4. The temperature while conducting all the three experiments should
be approximately same.
5.Epsom salt should be first dried in order to remove the water of
crystallization (MgSO4.7H2O).
Result:
The saturated solutions of NaCl, MgSO4 and C12H22O11 were made and
observed. The observations agreed with the related theory within the
range of experimental error.
Bibliography:
www.google .com
https://en.wikipedia.org/wiki/Link
N.C.E.R.T Class 12 Chemistry
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