Le Châtelier’s Principle: Examples of Chemical Equilibria
PURPOSE
To observe a number of interesting and colorful chemical reactions that are examples of chemicals systems at
equilibrium. Observe how these systems respond to changes in the concentrations of reactants or products or to
changes in temperature, and to see that the direction of the shift in the equilibrium tends to at least partially offset the
change in conditions, a principle first clearly stated by Le Châtelier.
PRELAB PREPARATION
A state of chemical equilibrium is the point of balance where the rate of the forward reaction equals the rate of the
reverse reaction. Many chemical reactions go essentially to completion, but some do not, stopping at a point that lies
between no reaction and an essentially complete reaction. The same point of equilibrium can be reached by mixing
together products or reactants, clearly indicating that chemical reactions can go either forward or backward! For a
reaction where equilibrium is reached, the point of equilibrium may be approached from either side of the reaction,
which emphasizes the dynamic nature of chemical reactions. Every chemical reaction is in principle a reversible
reaction, but if the point of equilibrium greatly favors the reactants, we state that there is no reaction. If the point of
equilibrium favors the products, we state that the reaction goes to completion.
Nearly every chemical reaction consumes or releases energy (endothermic and exothermic reactions, respectively).
For these reactions, we can regard energy as if it were a reactant or product, and by adding or removing energy (by
heating or cooling), we can produce a shift in the chemical equilibrium. The concentration of the reactants and
products change to reflect the new equilibrium.
There is an added complication; some chemical reactions are so slow that you might be fooled into thinking that no
reaction occurs, whereas the equilibrium point when it is finally reached might greatly favor the products! Fortunately,
all reactions you observe in this experiment are rapid, so you will be able to observe almost immediately the effects of
changing the concentration of either reactants or products. Notably, you will observe that when you make such
changes, the point of equilibrium shifts in a direction that tends to offset the change. The principle emphasized by this
behavior was first fully stated by the French chemist Henri Louis Le Châtelier in 1884: A chemical reaction that is
displaced from equilibrium by a change in conditions (concentration, temperature, pressure, volume) proceeds toward
a new equilibrium state in the direction that at least partially offsets the change in conditions.
PROCEDURE
A. Shifting of Equilibria in Acid-Base Reactions: The Common Ion Effect
2 CrO4-2 + 2 H+ ! Cr2O7-2 + H2O
Yellow
Chromate
Orange
Dichromate
1. To 1mL of 1M Potassium Chromate (K2CrO4), add several drops of 3M Sulfuric Acid (H2SO4). Record your
observations.
2. Add several drops of 6M Sodium Hydroxide (NaOH) to the mixture until a change occurs. Record observations.
B. Weak Acid-Base Indicator Equilibria
Methyl Orange:
Phenolphthalein:
HIn + H2O ! H3O+ + InRed
Clear
Yellow-Orange
Pink
-
HIn represents the protonated (acid) form of the methyl orange indicator, and In represents the deprotonated (basic)
form.
1. To 1mL of water, add a drop of methyl orange.
2. Add 2 drops of 6M Hydrochloric Acid (HCl), followed by 4 drops of 6M NaOH. Record observations.
3. Repeat the test, using phenolphthalein in the place of methyl orange. Record observations.
C. Weak Acid-Base Equilibria
CH3COOH + H2O ! H3O+ + CH3COOAcetic Acid
Acetate Ion
1. To each of two 1mL samples of 0.1M Acetic Acid (HC2H3O2), add a drop of methyl orange.
2. To this, add 1M Sodium Acetate (NaC2H3O2) dropwise until a change in color is observed.
The added salt, sodium acetate, has an ion in common with acetic acid, a weak acid that dissociates in water to give
acetate ion. Note that adding acetate ion produces a change in the color of the indicator. This indicates a change in
the indicator from its acid form to its base form, which in turn must mean that the hydrogen ion concentration became
smaller when the sodium acetate was added. The effect on the dissociation of acetic acid that is produced by adding
sodium acetate is called the common ion effect.
3. To each of two 1mL samples of 0.1M Ammonium Hydroxide (NH4OH), add a drop of phenolphthalein. Record
observations.
4. To one of the samples, add 1M Ammonium Chloride (NH4Cl) dropwise until a change is observed.
5. To the other sample, add 6M HCl dropwise until a change is observed. Record observations of both samples.
D. Temperature-Dependent Equilibrium of Cobalt (II) Complex Ions
It is common for cations to attract ions or molecules with free electron pairs to form aggregates called complexes.
-2
The CHLORO complex of cobalt (II), CoCl4 , is TETRAHEDRAL and has a BLUE color. The AQUO complex,
+2
Co(H2O)6 , is OCTAHEDRAL and has a PINK color. Because conversion of one form to another involves a
considerable energy change, the equilibrium is temperature dependent.
CoCl4-2 + 6 H2O ! 4 Cl- + Co(H2O)6+2 + E
Blue
Pink
Le Châtelier’s principle applied to this reaction predicts that if energy is removed (by cooling the system), the
equilibrium tends to shift toward the aquo complex, because a shift in the equilibrium to produce more aquo
complexes produces some energy, thus partially offsetting the change.
1. Put 3mL of 0.15M Cobalt (II) Chloride (CoCl2) solution in Methanol in a test tube. Add deionized water dropwise to
the blue solution until when it just turns pink. (If it’s already pink, proceed to the next step)
2. Divide the pink solution into two equal portions.
3. Add 12M HCl dropwise to one test tube until a change is observed. Record observations.
o
4. Heat the other test tube in a beaker of hot water (65-70 C) on a hot plate in a fume hood until a change is
observed. (CAUTION: Methanol vapors are toxic and highly flammable.)
5. Remove the test tube from the hot water bath and allow it to cool until a change is observed.
E. Equilibria of Saturated Solutions
In a saturated solution of a salt such as sodium chloride, an equilibrium exists between the solution and any
undissolved salt that may be present. Although not visually apparent, sodium chloride ions are going into and coming
out of solution at the same rate. Since the solution contains as much dissolved Na+ and Cl- ions as it can hold,
adding a solution containing a common ion to the saturated solution can cause salt to precipitate, indicating that a
stress has been applied to the system. The equilibrium has shifts to compensate for this stress.
NaCl + H2O ! Na+ + Cl1. To 2mL of saturated (5.4M) sodium chloride (NaCl), add several drops of 12M HCl. Record observations.
Name: ______________________________
Class/ Section: ________________________
Report: Le Châtelier’s Principle
A. Shifting of Equilibria in Acid-Base Reactions: The Common Ion Effect
Describe changes upon addition of:
1. 3M H2 SO4 : ______________________________________________________________________
2. NaOH: _________________________________________________________________________
3. How did the equilibrium shift in response to the added reagents (toward the products, reactants, etc.)?
4. How does OH- exert an effect?
B. Weak Acid-Base Indicator Equilibria
Describe the effects of adding HCl and NaOH to:
5. Methyl Orange: __________________________________________________________________
6. Phenolphthalein: _________________________________________________________________
7. Explain your observations in terms of the equilibria (using the chemical equation for this section) and Le Châtelier’s
principle.
C. Weak Acid-Base Equilibria
8. Complete the equation for the dissociation equilibrium of acetic acid in water.
CH3COOH + H2O !
9. Explain the observed changes when 1M Sodium Acetate is added to 0.1M Acetic Acid in the presence of Methyl
Orange indicator.
Describe any observed odor or changes when 0.1M Ammonium Hydroxide (with 2-3 drops of phenolphthalein) is
treated with the following:
10. NH4Cl: __________________________________________________________________________
11. HCl: ____________________________________________________________________________
In which direction (left or right) does each reagent shift the equilibrium for the dissociation of NH3 in the equation
below?
NH3 + H2O ! NH4+ + OH-
12. NH4Cl: ____________________
13. HCl:
____________________
14. Write the net ionic equation for the reaction of NH 4OH with HCl.
D. Temperature-Dependent Equilibrium of Cobalt (II) Complex Ions
15. Rewrite the equation for this equilibrium:
Describe what happens when you treat the pink (AQUO) complex in the following ways:
16. Add 12M HCl to the AQUO complex:
17. Heat the AQUO complex:
18. Cool the AQUO complex:
19. Interpret these results in terms of the equation describing the equilibrium and Le Châtelier’s principle:
E. Equilibria of Saturated Solutions
20. Describe what happens when 12M HCl is added to saturated (5.4M) NaCl:
-
21. What is the Cl concentration in the original solution? ________________________
-
22. What is the Cl concentration after adding 12M HCl? ________________________
-
23. Explain this behavior in terms of the relative Cl concentrations in the two solutions, and in terms of the
equilibrium equation:
NaCl(s) ! Na+ + Cl-