Contents
content
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page
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Abstract
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2
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Introduction
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3
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Literature review
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5
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Objectives
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7
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Methodology
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8
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Results
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11
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Discussion
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14
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Conclusion and
recommendation
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15
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Reference
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16
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Appendix
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17
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Abstract
This is an
experiment to study chemical equilibrium and Le Chatelier’s Principles.
Chemical equilibrium is the stable state a chemical reaction reaches when there
is no further undergoing reaction or change. Le Chatelier’s Principles state
that when a chemical equilibrium is disturbed by certain changes, it will shift
the position of equilibrium towards the side to minimize the effect of the
changes applied.
In this
experiment, six sets of chemical reaction were carried out to study the effect
of adding particles or ions of reactants or products to the position of
equilibrium.
(i)
The saturated sodium chloride solution
equilibrium
(ii)
The iron (III) thiocyanate ion
equilibrium
(iii)
The acetic acid equilibrium
(iv)
The chromate-bichromate equilibrium
(v)
The bismuth chloride-water equilibrium
(vi)
The cobalt (II) chloride equilibrium
The shifting of equilibrium can be determined
from the colour of the solution or mixture since the reactions chosen in this
experiment have reactants of different colours from corresponding products, or
undergo other distinct changes.For example, in the first part of the
experiment, concentrated hydrochloric acid was added to a saturated solution of
sodium chloride. The chloride ions yield by concentrated hydrochloric acid had
increased the chloride ion concentration which is a product of hydration of
sodium chloride. Hence, the changes were observed and shifting of this
equilibrium position was deduced.
This
experiment must be conducted very carefully and the substances must be measured
and added very accurately because any slight error might influent the outcome
of the experiment. Besides, the mixture of substances must be stirred well
until the substances are will mixed and form a mixture of same physical state
(homogeneous solution), in this experiment, liquid.
The
experiment outcome supports Le Chatelier’s Principles in the aspect of
concentration. When the concentration of a particle in an equilibrium is
altered, the equilibrium position will shift to the direction to counteract the
change. In this experiment, the concentrations of product or reactant particles
were increased by adding like particles (the same substance or substance
containing like ions). The equilibriums conducted shifted to the opposite
direction of the increased concentration, and produced more particles of the
substances it shifted to.
Introduction
This
is a chemical equation. A and B are reactant chemical species, S and T are
product species, and α, β, σ, and τ are the stoichiometric coefficients of the
respective reactants and products.
The change
from left to right in the equation is known as the forward reaction. The change
from right to left is the backward reaction. In a chemical system such as the
equation shown, when the forward reaction is happening at a same rate with the
backward reaction, the system is said to be in chemical equilibrium.
Chemical
equilibrium applies only to reversible reactions. A reversible reaction is the
one which can be made to go in either direction under certain conditions, in a
closed system, meaning that no substance is added to or extracted from the
reaction. In a closed system, energy is, however, can be transferred in and out
at will. The harpoon arrows pointing both ways indicate equilibrium. The equilibrium
position of a reaction is said to lay far to the right if, at equilibrium,
nearly all the reactants are used up and far to the left if hardly any product
is formed from the reactants.
Le
Chatelier’s Principles state that when a change or a stress is applied to an
equilibrium system, the position of equilibrium will shift in a way to
counteract the stress applied, if possible. Le Chatelier’s Principles are used
to make predictions about the shifting of equilibrium position in a system, and
the effect of stress in the system. The factors that affect an equilibrium
system are concentration, temperature, pressure and volume, pressure, and the
presence of catalyst. When a change in one of these factors is applied to a
system, the system will attempt to oppose the change by shifting the position
of equilibrium to the side which reduce the effect of the change.
Changing
the concentration of either reactants or products will shift the equilibrium
position to the side that would reduce that change in concentration. If the
concentration of the reactant is increased, the equilibrium position will shift
to the right to counteract the increase in amount of reactant, thus yielding
more products. On the other hand, if the concentration of the reactant is reduced,
the equilibrium position will shift to the left, and decreases the products.
Changes in pressure are only applied to
reactions involving gases:
The same
theory as changes in concentration is applied since pressure affects the number
of mole of gaseous particle in a unit volume.
Changes in
volume attribute to changes in pressure. The equilibrium concentrations of the
products and reactants do not directly depend on the pressure subjected to the
system. However, a change in the pressure due to a change in volume of the
system will cause a shift of the equilibrium position. Similar to the effects
of other factors, the equilibrium position will shift to the side with lowered
pressure, of shift away from the side with increased pressure.
The effect
of changing the temperature in the equilibrium system can be made clear by
incorporating heat as either a reactant or a product. When the reaction is
exothermic (∆H is negative, the reaction releases energy), we include heat as a
product. When the reaction is exothermic (∆H is positive, the reaction absorbs
energy), we include heat as a reactant. Hence, we can determine whether
increasing or decreasing the temperature will favour the forward or reverse
reaction by applying the same principles as concentration changes.
A catalyst
has no effect on the position of equilibrium; it only increases the rate to
which the equilibrium can be reached. A catalyst speeds up the rate of reaction
by providing additional mechanism(s) to the reaction. Adding a catalyst allows alternative
pathways to be made, where the particles can be absorbed onto the catalyst
temporarily before being rebounded into a new arrangement. The absorption of
particles to the catalyst requires lower activation energy than the rebonding
of particles directly, hence the activation energy of the whole reaction can be
lowered, and this frequently increases the rate of reaction since lower
activation energy can be reached more readily.
Literature review
Chemical
equilibrium is the situation at which the forward and backward reaction of a
reversible reaction proceeds at a same rate. Thought the concentrations of the
reactants and products remain as a constant, the reactions are actually
continuing in a way for every unit amount of products formed by the forward
reaction, the same amount of the products are converted to the reactants again
by the backward reaction. At such point, dynamic equilibrium is achieved, where
the word ‘dynamic’ indicates that the reactions continue.
Such reactions
must be carried out in a closed system. A closed system means that while the
substances are left to react under certain conditions, no substance can be
added to or extracted from the system. However, energy transfer is allowed at
will. For example, heat can be provided to the reaction.
The
above graph shows the reaction rate of a reversible reaction, let say, A +
B C + D. The reaction rate of A
+ B is initially the highest, because they are present in maximum amount. Then,
the reaction rate decreases with time, because the amount of the reactants
decreases. At the same time, C and D are formed. C and D react with each other
to form A and B in the reverse or backward reaction. More C and D are produced
as time proceeds, therefore the rate of backward reaction increases. This
carries on until the rates of forward and backward reactions become equal, then
the reactions will not affect the concentration or amount of the substances in
the system any further. Chemical equilibrium is reached.
Le Chatelier’s
Principles generally state that if a dynamic equilibrium is disturbed by
changing the conditions, the position of equilibrium moves to counteract the
change. Le Chatelier's principle gives a qualitative idea of an equilibrium
system's response to changes in reaction conditions. It does not give any
explanation on why the changes occur.
The first
factor to be predicted by Le Chatelier’s Principle is concentration. When the
concentration of a substance in equilibrium is increased, the system tends to
react the substance into products in order to reduce the concentration of the
substance again. This in turn increases the product of the reaction in that
direction of the chemical reaction. According to the statement or definition of
Le Chatelier’s principle, the other changes of concentration in an equilibrium
system can be predicted in the same way.
For the
second factor, i.e. pressure, it is applied only to the equilibrium systems
involving gas. The pressure of gas is determined by the number of moles of
gaseous particles in a unit volume. Therefore, if the pressure of a system is
increased, the position of equilibrium will shift to the direction with the
less number of moles of gaseous particles to decrease the pressure. The
opposite situation will happen if the pressure of the system is decreased. If
the number of molecules of both side of the equilibrium reaction is the same,
pressure does not affect the equilibrium position. This is because no matter
which direction the equilibrium position shift to, the number of moles of the
gas will remain the same and thus the pressure will not be altered.
The other
factor which affects the equilibrium position is temperature. However it works
differently for endothermic and exothermic reactions. Increasing the temperature
of a system in dynamic equilibrium favours the endothermic reaction. The system
counteracts the change by absorbing the extra heat. Hence endothermic reactions
will shift their equilibrium position to the right if the temperature is
increased, thus yielding more products. Decreasing the temperature of a system
in dynamic equilibrium favours the exothermic reaction. The system counteracts
the change by producing more heat. The equilibrium position will therefore
shift to the right when temperature decreases.
It is a
common mistake to assume that catalyst affects the equilibrium position of a
dynamic equilibrium. In fact, adding catalyst has no effect to a system in
equilibrium. This is because a catalyst speeds up both the forward and backward
reaction in the same extent. As the relative rate of the forward and backward
reactions does not change, the dynamic equilibrium will remain at the same
equilibrium position. Catalyst, however, can increase the rate at which the
dynamic equilibrium is reached.
An inert
gas is a gas which does not react with other elements or other like-element.
Inert gases are positioned in Group 18 of the periodic table, for example,
helium. Adding an inert gas into a gas phase equilibrium system at constant volume
does not result in a shift of the equilibrium position. This is because the
addition of a non-reactive gas does not change the partial pressures of the
other gases in the container. Therefore, the equilibrium constant does not
change because it depends on the partial pressure of each gas instead of the
total pressure exerted. Similarly, equilibrium position is also affected by the
partial pressures only. As a result, the equilibrium position is not affected
by the addition of noble gases into a system of equilibrium, provided that the
volume remains constant. If, however, the volume is allowed to increase in the
process, the partial pressures of all gases would be decreased resulting in a
shift towards the side with the greater number of moles of gas.
Objectives
1.
To understand dynamic equilibrium and
Le Chatelier’s Principle.
2.
To observe the change of an
equilibrium when the concentration of a reactant or product is altered.
3.
To be able to predict the effect of
concentration change on chemical equilibrium.
Methodology
Materials
A set of 4” and 6” test tubes, test tube rack,
pipettes, 100 mL graduated cylinder, 250 mL beaker, saturated sodium chloride
solution, concentrated hydrochloric acid, 0.1 M iron (III) chloride, 0.1 M
potassium thiocyanate, 6 M NaOH, 50% NaOH solution, 0.1 M acetic acid, methyl
orange, sodium acetate, sodium chloride, 0.1 M potassium chromate, 6 M nitric
acid, bismuth chloride, 1 M cobalt (II) chloride.
Methods
A. The Saturated Sodium Chloride Solution
Equilibrium
- To a 4" test tube, 5mL of saturated sodium chloride (NaCl) solution was added.
Its appearance recorded.
- To this solution, several drops of concentrated hydrochloric
acid (HCl) were added. Our observations recorded.
B. The Iron (III)
Thiocyanate Ion Equilibrium
- To 100 mL of water in a 250 mL beaker, 2 mL of 0.1 M
iron (III) chloride (FeCl) solution and 2 mL of 0.1 M potassium
thiocyanate (KSCN) solution was added.
This stock solution was stirred until it is homogeneous.
Observations were recorded.
- To a 4" test tube (Tube 1); 5 mL of
the stock solution was added. To this solution, 20 drops of 0.1 M iron
(III) chloride (FeCl3) solution was added. The observations
were recorded.
- To a 4" test tube (Tube 2), 5 mL of
the stock solution was added. To this solution, 20 drops of 0.1 M
potassium thiocyanate (KSCN) solution was added. The observations were
recorded.
- To a 4" test tube (Tube 3), 5 mL of
the stock solution was added and then 5 drops of 6 M sodium hydroxide
(NaOH) solution added. Observations recorded.
- To Tube 2, 2 drops of a 50% sodium
hydroxide (NaOH) solution was added. What is happening was recorded. Then
5 drops of 12 M hydrochloric acid (HC1) was added and the observations
recorded.
C. The Acetic Acid Equilibrium
- To each of three 4" test tubes, 3 mL of a 0.1 M acetic
acid (HC2H3O2) solution was added. To
each of the tubes, a few drops of methyl orange solution was added, and
the tubes agitated until the solutions are homogeneous and the
observations recorded.
- To the first test tube (Tube 1); a few crystals of sodium
acetate (NaC2H3O2) was added. The tube
was agitated in order to dissolve the solid. The observations recorded.
- To the second test tube (Tube 2), a few crystals of sodium
chloride (NaCl) was added. The tube was agitated in order to dissolve the
solid. The observations recorded.
- To the third test tube (Tube 3), a few drops of 6 M sodium
hydroxide (NaOH) ws added. The tube was agitated and the observations
recorded.
D. The Chromate-bichromate Equilibrium
- To a 4" test tube, 5 mL of a 0.1 M potassium chromate (K2CrO4)
solution was added. Its color was observed and recorded.
- To this solution, 6 M nitric acid (HNO3) solution
was added one drop at a time, until a distinct change is noted. The
observations recorded.
- Then, to the same test tube, 6 M sodium hydroxide (NaOH)
solution was added one drop at a time, until once again a distinct change
has been observed. Again, the observations recorded.
E. The Bismuth Chloride-Water Equilibrium
- To a 6" test tube, 2 mL of distilled water was added. A
small crystal of bismuth chloride (BiCl3) was added to the
water and the tube was agitated. The observations recorded.
- To this mixture, 12 M of hydrochloric acid (HCl) was added
one drop at a time while agitating the test tube, until a distinct change
observed and the observations of this change recorded.
- Then, water was added one drop at a time, with agitation, to
the test tube, until once again see a distinct change .These observation
recorded.
F. The Cobalt
(II) chloride Equilibrium:
- To a 4" test
tube, 5 drops of a 1 M cobalt (II) chloride (CoCl2) solution
was added. Its color was noted and this observation recorded.
- 12 mL of hydrochloric
acid (HC1) was added one drop at a time with agitation until a significant
change noticed. The observations recorded.
- Then,
water was added one drop at a time, with agitation, and the change is
observed occurring in the test tube recorded.
Results
- The Saturated Sodium Chlorine
Solution Equilibrium
SOLUTION
|
OBSERVATION
|
5
ml of saturated sodium chloride (NaCl ) solution
|
The
solution is colourless
|
5
ml of saturated sodium chloride (NaCl) solution + several drops of
concentrated hydrochloride acid (HCl)
|
White
precipitate
|
B. The Iron (III)
Thiocyanate Ion Equilibrium
SOLUTION
|
OBSERVATION
|
100
ml water + 2 ml of 0.1M iron (iii) chloride (FeCl3) solution + 2 ml of 0.1M
potassium thiocynate (KSCN ) solution
|
The
solution turns from pale yellow to red
|
NOTE: All the tubes in this part of the
experiment contain 5 ml of the stock solution above.
(Tube
1 ) + 20 drops of 0.1M iron (iii) chloride (FeCl3) solution.
|
The
intensity of red colour increases. A dark red solution is formed.
|
(Tube
2) + 20 drops of 0.1M potassium thiocyanate (KSCN ) solution.
|
The
intensity of red colour increases. A dark red solution is formed.
|
(Tube
3) + 5 drops of GM sodium hydroxide (NaOH) solution.
|
The
intensity of red colour decreases. A pale yellow solution is formed.
|
(Tube
2) + 20 drops of 0.1M potassium thiocyanate (KSCN ) solution + 2 drops of
50%
sodium hydroxide (NaOH) solution,
------------------------------------------------------------
+
drops of 12M hydrochloride acid (HCl ).
|
The
intensity of red colour decreases. A pale yellow solution is formed.
-----------------------------------------------------------
The
solution turns form pale yellow to red again.
|
C. The Acetic
Acid Equilibrium
SOLUTION
|
OBSERVATION
|
3 ml of a 0.1M
acetic acid (HC2H302) solution + a few drops of methyl orange solution
|
The solution
turns from colourless to orange upon addition of methyl orange solution.
|
NOTE: Each test
tube in this part ofthe experiment contains the above solution.
(Tube
1) + a few crystals of sodium acetate (NaC2H3O2)
|
The
intensity of orange colour decreases. A pale orange solution is formed.
|
(Tube
2) + a few crystals of sodium chloride (NaCl)
|
The
solution turns from orange to red.
|
(Tube
3) + a few drops of 6M sodium hydroxide (NaOH)
|
The
intensity of orange colour decreases. A pale orange solution is formed.
|
D. The Chromate-bichromate
Equilibrium
NOTE: A same
solution is used throughout this part of experiment.
SOLUTION
|
OBSERVATION
|
5 ml of a 0.1M
potassium chromate (K2CrO4) solution
|
The solution is
yellow
|
+ 6M nitric
acid (HN03) solution
|
The solution
turns from yellow to orange.
|
+ 6M sodium
hydroxide (NaOH) solution.
|
The solution
turns from orange to light yellow.
|
E. The Bismuth Chloride-Water Equilibrium
NOTE: A same
mixture is used through this part of the experiment
MIXTURE
|
OBSERVATION
|
2 ml of
distilled water + a small crystal of bismuth chloride (BiCl3)
|
The mixture is
milky
|
+ 12M of
hydrochloric acid (HCl)
|
The mixture
turns from milky to clear colour solution.
|
+ water
|
The mixture
turns milky again.
|
F. The Cobalt (II) Chloride Equlibrium
Note: A same
solution is used throughout this part of the experiment.
SOLUTION
|
OBSERVATION
|
5 drops of a 1M
cobalt (ii) chloride (COCl2) solution
|
The solution is
pale red
|
+ 12 ml of
hydrochloric acid (HCl)
|
The solution
turn from pale red to pale blue
|
+ water
|
The solution
turns from pale blue to pale purple. A
purple ring is formed on the surface.
|
Discussion
1. In your own words, explain Le
Chatelier’s Principle.
Le
Chatelier’s Principle states that when a change or stress is applied to an
equilibrium system, the position of the equilibrium will shift in the way to
counteract the stress applied, if possible.
The factors
affecting a system in equilibrium which Le Chatelier’s Principle can be applied
to are temperature, pressure and volume, concentration, and the presence of
catalyst in the reaction.
This
principle is used to predict the shifting of the equilibrium position in a
system, and thus the effect of stress in the system.
2. What would occur if a few drops of
saturated Na2SO4 solution were added
to a saturated NaCl solution?
Na2SO4 ↔ 2Na+ + SO42- (equation
1)
NaCl
→ Na+ + Cl- (equation
2)
When a few
drops of Na2SO4 solution were added
to a saturated NaCl solution containing Na+ ions and Cl- ions, the
concentration of Na+ ions will increase.
Hence, the equilibrium position of equation 1 will shift to the left according
to Le Chatelier’s Principle to counteract the increased concentration of the
products. This results in formation of more Na2SO4 since the reaction is reversible.
3. In procedure D3, explain the reason
for the change you observed.
2
CrO42- (aq) + 2H+ (aq) 
Cr2O72- (aq)
+ H2O (l)
Cr2O72- (aq)
+ H2O (l)
(yellow) (orange)
NaOH
undergoes neutralization with HNO3, resulting in the decrease of H+inos.
The equilibrium position will shift to the left and more CrO42-
ions will be produced. Hence, the solution changes colour from orange to
yellow. However, the yellow solution is lighter than the potassium chromate
solution in Step 1 because during the neutralization, H2O is
produces, diluting the solution.
Conclusion and recommendations
Le Chatelier's principle states, if a
chemical system at equilibrium experiences a change in concentration,
temperature, volume, or partial pressure, then the equilibrium shifts to
counteract the imposed change and a new equilibrium is established.
This experiment has proven that
when a system at equilibrium is subjected to a change in concentration of a
reactant or product, the product will, if possible, shift its equilibrium
position so as to counteract the effect of the change.
In obtaining the results of this experiment, the observer must
take the observations carefully and record the colour of the solution
accordingly. The stock solution must be agitate
until it is homogeneous before we use it. Be extra careful while handling the
chemicals because some of it is corrosive, flammable or carcinogenic.
References
·
Chang, Raymond (2003). General
Chemistry.3rded. New York: McGraw-Hill
·
(2000). A-Level Chemistry. UK: Nelson Thornes
·
P.W. Atkins, Elements of Physical Chemistry, 3rd Edition,
Oxford University Press, 1993.
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