Effects
of Concentration on Chemical Equilibrium
In
a chemical system at equilibrium, forward and reverse reactions are taking
place at equal rates, and concentrations of reactants and products remain constant.
The concentrations of reactants and products need not be equal, however.
Equilibrium can be reached at any relative concentrations of reactants and
products, depending upon the conditions and the specific reaction.
LeChatelier’s principle is a very
useful guide to predicting or explaining what happens when changes are made to
a system at equilibrium. In this activity, you will observe the reaction
between hydrated cobalt ion and chloride ion. With this reaction, the outcome
of changes in equilibrium conditions are visible.
The
reactant is pink and the product is blue. When the equilibrium conditions are
such that roughly equal amounts of reactant and product are present, a violet
color is seen. You will first observe how changing concentrations of reactants
and products affects equilibrium. Then you will investigate the effects of
different temperatures on the equilibrium.
OBJECTIVES
•
Set up an equilibrium system.
•
Observe the effects of changing concentrations and temperature on the
equilibrium.
HYPOTHESIS:
MATERIALS
apron
shallow pan
or
goggles
pneumatic
trough
96-well
microplate toothpicks
96-well
template distilled water
microtip pipets
(4) clear tape
hot
plate thermometer
thermal
mitt 400-mL beaker
PROCEDURE
Part
1: Effects of Concentration
1. Based
on LeChatelier’s principle, form a hypothesis about
the direction in which increasing the concentration of Cl– will
shift the reaction of hydrated cobalt ion and chloride ion. Record your
hypothesis.
2. Label
a 96-well template and a Microplate Data Form as
shown under Data and Observations. Place the microplate
on the template with the numbered columns away from you and the lettered rows
to the left.
3. CAUTION:
Many of the chemicals you will use are toxic. Follow proper
chemical hygiene procedures. Wash your hands thoroughly after completing this
laboratory activity. Place 4 drops of Co(NO3)2
solution in each well A1 through A8.
4. Add
1 drop of concentrated HCl to well A1. CAUTION: Hydrochloric
acid (HCl) is very corrosive. Avoid contact with skin
and eyes.
5. Add
2 drops of concentrated HCl to well A2.
6. Continue
to add HCl to each succeeding well, increasing the
amount for each well until you reach 8 drops in well A8. Stir each well with
the same toothpick, beginning with well A1.
7. Record
the color of the solution in each well in row A of the Microplate
Data Form.
8. Add
4 drops of concentrated HCl to each well.
9. Stir
each well with the same toothpick, as before. Record the color for the solution
in each well in the second row of the data form.
10. Add
5 drops of distilled water to each well.
11. Again,
stir each well with the same toothpick. Record the color for the solution in
each well in the third row of the data form.
12. Add
2 drops of 1M AgNO3 to each well.
13. Stir
each well with a single toothpick. Record the color for the solution in each
well in the fourth row of the data form.
14. Discard
the solutions in wells A1 through A8 according to the directions of your
teacher. Rinse the wells with distilled water and discard the rinse water in
the same manner as the solutions.
15. Rinse
the microtip pipets with
distilled water.
Part
2: Effects of Temperature
1. Set
up another set of equilibrium reactions by repeating steps 2–6 of the procedure
for Part 1.
2. Label
row G of the Microplate Data Form “Raise
Temperature.” Label Row H “Lower Temperature.”
3. Fill
a 400-mL beaker about one-half full with tap water. Heat the water to 10°C –
15°C above room temperature.
4. Prepare
a water bath by pouring the heated water into a shallow pan or pneumatic
trough.
5. Cover
the tops of wells A1 through A8 with a piece of clear tape. Seal each well by
running your finger over the top.
6. Float
the microplate on the surface of the hot water.
7. Allow
2 minutes for the plate to be heated by the water. Record the colors of the
solutions in the wells.
8. Replace
the water in the trough with water that is 10°C below room temperature.
9. Float
the microplate on the surface of the cold water.
10. Allow
2 minutes for the plate to be cooled by the water. Record the colors in the
wells.
11. Discard
the solutions in wells A1 through A8 according to the directions of your
teacher. Rinse the wells with distilled water and discard the rinse water in
the same manner as the solutions.
12. Rinse
the microtip pipets with
distilled water.
DATA AND OBSERVATIONS
Label
a Microplate Data Form as shown and use it to record
your observations.
Use
P to indicate a pink color, V for violet, and B for blue.
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PART 1 |
1 |
2 |
3 |
4 |
5 |
6 |
7 |
8 |
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Co 2+ + HCl |
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+ HCl |
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+ H2O |
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+Ag |
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PART 2 |
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RAISE TEMPERATURE |
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LOWER TEMPERATURE |
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ANALYSIS
Part
1: Effects of Concentration
1. How
does the concentration of Cl– vary in the initial solutions prepared
in row A? In which well is the indicator violet?
In
how many wells is the CoCl4 2– concentration higher than
the Co(H2O)6 2+ concentration?
2. What
condition changed when excess concentrated HCl was
added? How did this change affect the position of the violet indicator well and
the number of wells with higher CoCl4 2– concentration?
3. What
condition changed when water was added? How did this change affect the position
of the indicator well and the number of wells with higher CoCl4 2–
concentration?
4. What
condition changed when 1M AgNO3 was added? How did this
change affect the position of the indicator well and the number of wells with
higher CoCl4 2– concentration?
Part
2: Effects of Temperature
1. What
effect did a higher temperature have on the relative numbers of pink and blue
wells?
2. What
effect did a lower temperature have on the relative numbers of pink and blue
wells?
CONCLUSIONS
1. Using
LeChatelier’s principle and your analysis, explain
the results of adding concentrated HCl, H2O,
and AgNO3 to the reaction of Co(H2O)6 2+
with Cl– at equilibrium.
2. According
to LeChatelier’s principle, an increase in
temperature shifts an equilibrium in the direction of the endothermic reaction
and a decrease in temperature shifts it in the opposite direction of the
exothermic reaction. In the reaction of Co(H2O)6 2+
with Cl–, which direction of the equilibrium is exothermic and which
is endothermic?
3. Was
your hypothesis about the effect of Cl– concentration on the equilibrium
correct? Explain.
EXTENSION AND APPLICATION
1. Sodium
chloride would have been a safer source of Cl– for the reaction in this
experiment, but it could not be used. Explain why you think concentrated HCl was necessary.
2. The
active ingredient in swimming pool disinfectant is hypochlorous
acid, HClO, which is most effective in a narrow pH
range.
How
is the concentration of HClO affected when the
concentration of H+ is increased and decreased? Sometimes baking soda, NaHCO3,
is added to swimming pools. Why?