CHEMISTRY P
MRS. BENIPAL, MRS. EDWARDS, MR. NELSON, DR. SELGRATH
NAME:
SPRING 2012
EXPERIMENT #11: CHEMICAL EQUILIBRIUM AND LE CHATELIER'S PRINCIPLE
PURPOSE:
To study chemical reactions at equilibrium and their responses to stress as described by
Le Chatlier's Principle.
INTRODUCTION:
You will study three different equilibrium systems involving ions in solution. The first system is the conversion
of the indicator thymol blue from its blue form to its yellow form. The extent to which the forward reaction is
favored depends upon the concentration of hydrogen ions, [H+], in solution:
thymol blue (blue form) + H+(aq) ⇌ thymol blue (yellow form)
The second reaction involves the light yellow iron (III) ion, Fe3+, and thiocyanate ion, SCN–, which forms the
colored complex, FeSCN2+:
Fe3+(aq) + SCN-(aq) ⇌ FeSCN2+(aq)
The third system involves the hydrated cobalt (II) ion, Co(H2O)62+, which can be converted to the chlorinated
complex ion, Co(H2O)4Cl2:
Co(H2O)62+(aq) + 2Cl–(aq) ⇌ Co(H2O)4Cl2(aq) + 2H2O(l)
You will study the effects of changing the temperature on this system. Your data should allow you to predict
whether the forward or reverse reaction is endothermic.
PRE-LAB QUESTIONS: (Complete in your science notebook in advance of entering the laboratory)
1. Write an equation for the equilibrium reaction in Part I.
2. Write the equation for the equilibrium reaction involving the formation of the iron (III) thiocyanate ion from
Part II.
3. Write the equation for the equilibrium reaction between the hydrated cobalt (II) ion and the hydrated cobalt (II)
chloride ion from Part III.
4. In Part I, what is the purpose of adding thymol blue to two Erlenmeyer flasks if additional reagents are only to
be added to the first flask?
5. What precautions should you observe when heating the cobalt chloride and acid solution in Part III?
6. What is the purpose of adding distilled water to the beaker in Part II?
7. What is the function of test tube “A” in Part II?
8. Write and balance the equation for the reaction between hydrogen ions and hydroxide ions.
9. Referring to your answer to question #8, how might the hydrogen ion concentration be decreased?
10. What is a spectator ion?
©BHS Chemistry, 202
CHEMISTRY P
MRS. BENIPAL, MRS. EDWARDS, MR. NELSON, DR. SELGRATH
NAME:
SPRING 2012
PROCEDURE: (Include with your laboratory report, or refer to this page)
PART I: Thymol Blue Equilibrium
1. Clean two 125 mL Erlenmeyer flasks. To each flask, add approximately 20 mL of distilled water and 5 drops
of thymol blue solution. Record the color of this solution in the data section.
CAUTION: Hydrochloric acid (HCl) and sodium hydroxide (NaOH) are extremely corrosive. Handle these
solutions with care.
2. To the first flask, add a single drop of 0.1 M NaOH. Swirl the contents of the flask and continue the dropby-drop addition until a definite color change is observed. The second flask serves as a control. Compare
the solution colors. Record the new color of the solution.
3. To the first flask, add 0.1 M HCl drop-by-drop until a definite color change is observed. Record the new
color.
PART II: Ferro-thiocyanate Ion Equilibrium
4. Place 5 drops of 0.2 M FeCl3 into a 100 mL beaker. Record the color of the iron (III) chloride solution.
Record the color of a 0.2 M KSCN solution, and then add 5 drops of the potassium cyanate to the beaker
containing the FeCl3. Swirl the mixture and note the color.
5. Add enough distilled water to the solution to dilute the intense color to a light amber color (50 to 70 mL).
6. Pour approximately 5 mL of this solution into each of five test tubes labeled A to E.
7. To test tube A, add 10 drops of distilled water. Test tube A will serve as a control.
8. To test tube B, add 10 drops of 0.2 M KCl.
9. To test tube C, add 10 drops of 0.2 M Fe(NO3)3.
10. To test tube D, add 10 drops of 0.2 KSCN.
11. To test tube E, add 10 drops of 6.0 M NaOH.
12. Record the color of the solution after the addition of drops to test tubes B through E; compare to the
control test tube A.
PART III: Hydrated Cobalt (II) Chloride Equilibrium
13. Place a small crystal (approximately 0.2 g) of cobalt (II) chloride hexahydrate, CoCl2·6H2O, into each of
two 100 mL beakers.
CAUTION: CoCl2·6H2O is very poisonous, and can be fatal if taken internally. Do not get any in your
mouth, and wash your hands at the conclusion of this lab.
14. Using a 25 mL graduated cylinder, add 10 mL of 6.0 M HCl to the first beaker. Record the color of the
resulting solution.
15. To the second beaker, add 10 mL of distilled water. Record the color of the resulting solution.
16. Gradually add distilled water to the solution in the first beaker until it appears to be the same color as that in
the second beaker. Record the color.
17. Set-up ring stand assembly and light Buddy Bunsen. Gently heat the first beaker over a low flame. After a
definite color change is observed, shut off the burner. Remove the beaker from the ring stand, and record
the color of the solution.
18. Set-up a cool-water bath in a 250 mL beaker, fill with approximately 50 to 100 mL of tap water, and lower the
first beaker into the bath. As it cools, record any changes in color.
©BHS Chemistry, 202
CHEMISTRY P
MRS. BENIPAL, MRS. EDWARDS, MR. NELSON, DR. SELGRATH
NAME:
SPRING 2012
EXPERIMENT #18: Equilibrium and Le Chatelier’s Principle (continued)
DATA AND OBSERVATIONS: (Record on a separate page)
PART I: Thymol Blue Equilibrium
a) Initial color of thymol blue solution:
Reagent Added
Stress
(ion added)
________________________
Color Observed
Equilibrium Shift
(forward or reverse)
b) NaOH(aq)
c) HCl(aq)
PART II: Ferro-thiocyanate Ion Equilibrium
d) Color of 0.2 M FeCl3 solution:
________________________
e) Color of 0.2 M KSCN solution:
________________________
f) Color of mixture; FeSCN2+ complex:
(test tube A)
________________________
Reagent Added
(test tube)
g) distilled water (tt A)
h) KCl (tt B)
i) Fe(NO3)3 (tt C)
j) KSCN (tt D)
k) NaOH (tt E)
Stress
Ion
Spectator
Ion
Observed
Color
Equilibrium Shift
(forward or reverse)
PART III: Hydrated Cobalt (II) Chloride Equilibrium
l) Color of Co(H2O)4Cl2 solution:
.
m) Color of Co(H2O)62+ solution:
Step
n) Step 16
o) Step 17
p) Step 18
©BHS Chemistry, 202
pink
Stress
Observed Color
.
Equilibrium Shift
(forward or reverse )
CHEMISTRY P
MRS. BENIPAL, MRS. EDWARDS, MR. NELSON, DR. SELGRATH
NAME:
SPRING 2012
POST-LAB DISCUSSION:
The equilibrium system in Part I can be affected by the addition of any reagent supplying H+ or OH-. Hydrogen
ions are characteristic of acids and hydroxide ions are characteristic of bases. The addition of an acid should
favor the conversion into the yellow form of the indicator.
Another factor to consider in explaining your observations is the fact that H+ will react with OH- to form water
molecules. As a result of the formation of water, the concentrations of H+ and OH- in the equilibrium reaction are
reduced.
In reviewing your data from Part II, it is important to note that the formation of FeSCN2+ is heavily favored over the
reverse decomposition reaction. Diluting the solution with water allows you to observe small changes in the
equilibrium more readily. In explaining your observations of this equilibrium, remember that spectator ions do not
participate in the net reaction.
In step 11 of Part II, you noted the formation of a precipitate. In this case, the precipitate is Fe(OH)3. In
explaining your results, remember that the formation of this precipitate reduces the Fe3+ concentration of the
solution.
For the reaction in Part III, recall that increasing the temperature of an equilibrium reaction favors the endothermic
reaction. From your observations, you should be able to predict which of the two reactions, forward or reverse, is
endothermic.
QUESTIONS: (Answer on a separate page)
1. Apply Le Chatelier's Principle to explain the results you would expect to obtain when 6.0 M NaOH(aq) was
introduced into the iron (III) thiocyanate ion (FeSCN2+) equilibrium system. Explain your answer.
2. If the hydrated cobalt (II) ion complex (Co(H2O)62+) was placed in a refrigerator, what would you predict as the
color of the refrigerated solution? Explain your answer.
3. Analyze the equilibrium equation for the cobalt (II) ion. Using your lab results determine which reaction is
exothermic: the forward reaction or the reverse? Explain your decision.
4. Predict how the addition of NaCl would affect the cobalt (II) ion equilibrium. Use Le Chatelier's Principle to
explain your answer.
5. Predict how the addition of AgNO3 would affect the cobalt (II) ion equilibrium. Use Le Chatelier's Principle to
explain your answer.
6. If heating the solution would cause the color to shift from blue to yellow, which direction is the exothermic
reaction for the equilibrium examined in Part I, the forward reaction or the reverse reaction?
7. Write the equilibrium law (mass action) expression for the systems explored in Part II and Part III.
8. In Part II you added Fe(NO3)3(aq) to test tube C (initially at equilibrium).
a) What effect did this have on the rate of the forward reaction (increase, decrease, no change)?
b) What effect did this have on the rate of the reverse reaction (increase, decrease, no change)?
c) How will this effect the concentration of FeSCN2+(aq) ion (will more be produced by the forward reaction, or
will more be consumed by the reverse reaction, or will it be unaffected)?
©BHS Chemistry, 202