Project on
Saturated
Solutions:
Measuring
Solubility
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. ******** Head,
Department of Chemistry
ABCDEFaith
ACKNOWLEDGEMENTS
I gratefully
acknowledge my sincere thanks to our
respected chemistry
teacher Mrs/Mr.####### for
her remarkable,
valuable guidance and supervision
throughout the
project work. I 'm also most indebted
to Mrs/Mr.********
for her encouragement, help, suggestion
and readily helpful
service in performing the
experiment.
XYZ
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
Example State of State of State of
Solute Solvent Solution
Air, natural gas gas gas gas
Alcohol in water,
antifreeze
liquid liquid liquid
Brass, steel solid solid solid
Carbonated water,
soda
gas liquid liquid
Sea water, sugar
solution
solid
1
liquid liquid
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 adition of salt the melting
point of water increases
and as aresult 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 freezingpoint
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 Hbonding,
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, C1
2 H 2 2 O 1 1)
• 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
dissolved in 100mL
water to make
saturated solution.
Moles dissolved
NaCl (Non-iodized
common salt)
36.8 grams 0.7
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.
Howerver 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 incrrease was
shown by NaCl,
followed by Epsom salt and sucrose. These
facts too agreed with
the theory as at high temperatures the
kinetic enery of
molecules increases and the collisions are
more effective.
Conclusions:
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
MgSO4 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 (actual photo)
Solution of scucrose MgSO4 solution
(unsaturated and
saturated)
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:
Google.com
Wikipidia.com
