Chemical Equilibrium and Le Chatelier’s Principle
Pre-Lab Discussion
In most of the chemical reactions you have studied so far, at least one of the reactants
has been “used up.” The point at which a reactant is used up marks the end of the reaction, and
the reaction is said to have “gone to completion.” Under ordinary circumstances, the product(s)
of such reactions are not able to react to re-form the original reactants. Thus, these are “one
way” reactions. They proceed in one direction only.
Many other chemical reactions do not go to completion. Rather, the products of these
reactions remain in contact with each other and react to re-form the original reactants Such
reactants are said to he reversible. In a reversible reaction, the forward arid reverse reactions
proceed at the same time. When the rates of the two reactions are equal, a state of chemical
equilibrium is said to exist. Under such conditions, both the forward and reverse reactions
continue with no net change in the quantities of either products or reactants.
A state of equilibrium is affected by concentration and temperature and, if gases are
involved, by pressure. If a system at equilibrium is subjected to a change in one or more of
these factors, a stress is placed on the system. According to Le Chatelier’s principle, when a
stress is placed on a system at equilibrium, the equilibrium will shift in the direction that tends
to relieve the stress. Equilibrium will be reestablished at a different point, that is, with
different concentrations of reactants and products.
In this experiment, we will study two equilibrium systems. The equilibrium equation for
the reversible reaction of the first system is:
2CrO42- + 2H3O+  Cr2072- + 3H20
(yellow)
(orange)
The addition of an acid to this system increases the H3O+ concentration and causes the
equilibrium to shift to the right. The addition of any substance that causes a decrease in H3O+
concentration will have the opposite effect.
The equilibrium equation for the second system is:
CuCl42- + 4H2O  Cu(H2O)42+ + 4Cl(green)
(blue)
In this system, we will subject the system to a number of stresses (temperature and
concentration). It will be up to you to observe, through color changes, whether the system
shifts to the left (green) or to the right (blue)
By studying these two systems, you should achieve a better understanding of equilibrium
systems and their responses to stress.
Purpose
To study equilibrium systems and their responses to stress as described by Le Chatelier’s
principle.
Materials
beaker, 100 ml
dropper pipette
graduated cylinder, 10-mL
marking pencil
test tubes, 13x 100-mm (6)
test tube rack
hot plate
safety goggles
lab apron or coat
0.1 M K2CrO4
0.1 M K2Cr2O7
1.0 M HCl
1.0 M NaOH
16 M HCl
96% ethyl alcohol
copper (II) chloride solution
distilled water
Safety
Handle the HCl and NaOH solutions with care. They are corrosive substances and can injure the
skin or eyes. Flush any spills with cold water and report them to your teacher. Always wear
safety goggles and a lab apron or coat when working in the lab.
Procedure
1.
2.
3.
4.
5.
6.
1.
2.
3.
4.
5.
PART A
Using a marking pencil, number four test tubes 1 through 4. Stand the tubes in a rack.
Measure out 10 mL of 0.1 M K2CrO4. Pour 5 mL each into test tube 1 and test tube 2. Rinse
the graduated cylinder and measure out 10 mL of 0.1 M K2Cr2O7. Divide this equally
between test tube 3 and test tube 4.
Using a dropper pipette, add 1.0 M HCl drop by drop to test tube 1 until the color changes.
Record your observations.
Repeat step 3 with test tube 2. As soon as the color changes, rinse the pipette and use it
to add 1.0 M NaOH drop by drop to the solution until the color changes again. Record your
observations for this step.
Using the pipette, add 1.0 M NaOH to test tube 3 until the color changes. Record your
observations.
Repeat step 5 with test tube 4. As soon as the color changes, rinse the pipette and use it
to add 1.0 M HCl to the solution until the color changes again. Record your observations.
PART B
Fill test tube #1 to about 1/3 full with the copper (II) chloride solution. Put it in the rack;
this will be your control tube.
Fill test tube #2 to about 1/3 full with the copper (II) chloride solution. Place the tube in
hot water. Compare any color change to the control tube. Record your findings.
Fill test tube #3 to about 1/3 full with the copper (II) chloride solution. Place the tube in
ice water. Compare any color change to the control tube. Record your findings.
Fill test tube #4 to about ¼ full with the copper (II) chloride solution. Slowly add ethyl
alcohol to it. Compare any color change to the control tube. Record your findings.
Fill test tube #5 to about ¼ full with the copper (II) chloride solution. Slowly add distilled
6.
water. Compare any color change to the control tube. Record your findings.
Fill test tube #6 to about 1/3 full with the copper (II) chloride solution. Go to the
teacher’s desk. You will receive the 16 M HCl directly from the teacher. Compare any color
change to the control tube. Record your findings.
Observations and Data
Test Tube
1
2
3
4
PART A
Color Change
PART B
Test Tube
#2 – Heat Study
#3 – Cooling Study
#4 – Dehydration Study
#5 – Hydration Study
#6 – Common Ion Effect
Color Change
Conclusions and Questions
1. Write equilibrium equations for the reversible reactions that take place in Part A and Part B.
2. Using Le Chatelier’s principle, explain the color changes noted in Part A.
3. Using Le Chatelier’s principle, explain the color changes noted in Part B.
CITATIONS
Caiafa, M. (2006). “The Equilibrium of the Complex Salt.” Fresh Meadows, NY
Dorin, H., et al. (1989). Chemistry – The Study of Matter. Englewood Cliffs, NJ. Prentice Hall.
Pages 524-527
Wagner, M. (1989). Chemistry – The Study of Matter Laboratory Manual (Fourth Edition).
Englewood Cliffs, NJ. Prentice Hall. Pages 175-178.