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Chapter 18  Reaction Rates and Equilibrium
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SMALL-SCALE EXPERIMENT
LE CHÂTELIER’S PRINCIPLE AND
CHEMICAL EQUILIBRIUM
Small-Scale Experiment for text Section 18.2
OBJECTIVES
 Observe and record how a chemical system at equilibrium responds to changes in concentration of reactants
or products.
 Describe shifts in equilibrium in terms of Le Châtelier’s principle.
INTRODUCTION
If you’ve ever watched a game of football, you know that it has become very specialized. Coaches design special
offenses and defenses for various down and distance situations, as well as for positions of the ball on the field, the
score, and the time remaining in the game. For example, offensive and defensive strategies are quite different for
a third down and one situation than for a first down and ten. As a result, several offensive and defensive players
often run on and off the field after each play. One thing remains constant: only 11 players per team are on the
field for any given play. This is an example of a dynamic equilibrium. A dynamic equilibrium is one in which
forward reactions (players running on the field) take place at the same rate as reverse reactions (players running
off the field). There is no net change in the number of players on the field.
You learned in Section 18.2 that chemical systems often reach a state of dynamic equilibrium. In a system at
chemical equilibrium, the rate of the forward reaction equals the rate of the reverse reaction. As a result, the
amounts of the reactants and products of equilibrium remain constant as long as no stress is placed on the system.
Le Châtelier’s principle says that if a system at equilibrium is stressed, the equilibrium balance will shift in a
direction that will relieve the stress. For example, if you add a reactant, the equilibrium will shift toward products
(to the right) so that there is a different balance of reactants and products. Similarly, if you add a product, the
equilibrium will shift toward reactants (to the left).
PURPOSE
In this experiment, you will investigate chemical systems at equilibrium. You will disturb them by adding or
subtracting reactants or products and observe how the equilibrium system responds. You will explain those
changes in terms of Le Châtelier’s principle.
SAFETY
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Wear safety goggles, an apron, and gloves when working with corrosive chemicals.
Use full small-scale pipets only for the controlled delivery of liquids.
Don’t chew gum, drink, or eat in the laboratory. Never taste a chemical in the laboratory.
Avoid inhaling substances that can irritate your respiratory system.
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MATERIALS
Small-scale pipets of the following solutions:
bromthymol blue (BTB)
potassium iodide (KI)
hydrochloric acid (HCl)
nitric acid (HNO3)
sodium hydroxide (NaOH)
silver nitrate (AgNO3)
ammonia (NH3)
sodium carbonate (Na2CO3)
copper(II) sulfate (CuSO4)
sodium thiosulfate (Na2S2O3)
lead(II) nitrate (Pb(NO3)2)
sodium phosphate (Na3PO4)
EQUIPMENT
small-scale reaction surface
empty pipet for stirring
plastic cup
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EXPERIMENTAL PAGE
Mix the following. Stir each mixture thoroughly by blowing air through a pipet. The chemical equations describe the
changes you observe.
1. Mix one drop BTB and one drop HCl (H+). Record your observations in Table 29.1.
Now add just enough NaOH to induce another change. Alternately add more HCl and
NaOH.
HBTB 
yellow
H+ + BTB–
blue
2. Mix one drop BTB and one drop NH3. Record your observations in Table 29.1. Now
add just enough HCl (H+) to effect a change.
NH4+ + OH–
NH3 + HOH
NH3 + H+
NH4+
3. Mix one drop NH3 and two drops CuSO4 (Cu2+). Record your observations in Table
29.1. Now add, with stirring, just enough NH3 to effect a change. Add HCl, with
stirring, until the light blue precipitate returns. Add more HCl until the precipitate
disappears. Repeat this procedure until you are sure of what you see.
Cu2+ + 2OH–
Cu2+ + 4NH3
H+ + OH–
Cu(OH)2(s) (light blue precipitate)
Cu(NH3)42+ (royal blue solution)
HOH
4. Mix one drop Pb(NO3)2 (Pb2+) and one drop KI (I–). Record your observations in Table
29.1. Now add, with stirring, enough NaOH to effect a change. Now add HNO3.
Pb2+ + 2I–
PbI2(s) (bright yellow precipitate)
Pb2+ + 2OH–
Pb(OH)2(s) (milky white precipitate)
5. Mix one drop Pb(NO3)2 and one drop NaOH with stirring. Record your observations in
Table 29.1. Now add more NaOH, drop by drop with stirring, untila change occurs.
Finally, add nitric acid, HNO3(H+), drop by drop, slowly, with stirring, until a change
occurs.
Pb2+ + 2OH–
H+ + OH–
Pb(OH)2(s)
HOH
Pb2+ + 3OH–
Pb(OH)3–
6. Add two drops AgNO3 to a plastic cup. Now add one drop of each of the chemicals in the left column, one at a
time. Stir and observe after each addition. When a change occurs, record your observations in Table 29.1. Go on
to the next chemical. Near the last addition, you may need to add a few more drops of AgNO3.
Add these chemicals in this order
Equations that account for changes
a. 1 drop Na2CO3
b. HNO3
c. NaOH
d. HCl
a. 2Ag+ + CO32–
Ag2CO3(s)
b. CO32– + 2H+
HOH + CO2(g)
c. 2Ag+ + 2OH–
Ag2O(s) + H2O
d. H+ + OH–
HOH
Ag+ + Cl–
AgCl
e. Ag+ + 2NH3
Ag(NH3)2+
f. Ag+ + I–
AgI(s)
g. Ag+ + S2O32–
AgS2O3–
h. 3Ag+ + PO43–
Ag3PO4(s)
e. NH3
f. KI
g. Na2S2O3
h. Na3PO4
Experiment 29 Le Châtelier’s Principle and Chemical Equilibrium
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Place this side of the Experimental Page facedown. Use the
other side under your small-scale reaction surface.
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EXPERIMENTAL DATA
Record your results in Table 29.1 or in a copy of the table in your notebook.
Table 29.1 Dynamic Equilibria
1. BTB + HCl
5. Pb(NO3)2 + NaOH
+ NaOH:
More NaOH
+ HNO3
2. BTB + NH3
+ HCl
3. NH3 + CuSO4
6. AgNO3 + Na2CO3
+ HNO3
+ NaOH
More NH3
+ HCl
+ HCl
+ NH3
More HCl
4. Pb(NO3)2 + KI
+ NaOH
+ KI
+ Na2S2O3
+ Na3PO4
+ HNO3
CLEANING UP
Avoid contamination by cleaning up in a way that protects you and your environment. Carefully clean the small-scale
reaction surface by absorbing the contents onto a paper towel. Rinse the small-scale reaction surface with a damp
paper towel and dry it. Dispose of the paper towels in the waste bin. Wash your hands thoroughly with soap and
water.
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QUESTIONS FOR ANALYSES
Use what you learned in this experiment to answer the following questions.
1. Bromthymol blue (BTB) is an acid, so it has a hydrogen ion (H +). For the purposes of this question, you will
write it HBTB. HBTB ionizes in water to produce hydrogen ion (H+) and bromthymol blue ion (BTB–).
HBTB
yellow
H+ + BTB–
blue
a. What color appeared when you added HCl to BTB? What color is HBTB? Explain the shift in equilibrium in
terms of Le Châtelier’s principle.
b. What color appeared when you added NaOH to the mixture? What color is BTB–? Explain the shift in
equilibrium in terms of Le Châtelier’s principle.
2. Ammonia reacts with water to produce hydroxide ion. Ammonia is neutralized by hydrogen ion:
NH3 + HOH
NH3 + H+
NH4+ + OH–
NH4+
a. What color did BTB change in the presence of ammonia (NH3)? Explain thecolor change in terms of Le
Châtelier’s principle. Include an equation in your explanation.
b. What happened when HCl was added? Explain in terms of Le Châtelier’s principle.
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3. Copper ions react differently in the presence of varying amounts of ammonia (NH3). The hydroxide produced
from a small amount of ammonia produces aprecipitate. Excess ammonia produces a highly colored solution:
Cu2+ + 2OH–
Cu2+ + 4NH3
Cu(OH)2(s) (precipitation equilibrium)
Cu(NH3)42+(complex ion in solution)
a. What color is the precipitate when only a little ammonia is added? What coloris the solution in the presence of
excess ammonia?
b. Explain in terms of Le Châtelier’s principle the disappearance of the precipitate when excess ammonia is
added.
c. What is the effect of the HCl? Explain in terms of Le Châtelier’s principle.
4. Lead(II) ion reacts with iodide ion to produce a bright yellow precipitate, and with hydroxide ion to form a milky
white precipitate.
Pb2+ + 2I–
Pb2+ + 2OH–
PbI2 (s)
Pb(OH2)(s)
bright yellow
milky white
a. Why does the yellow precipitate disappear when sodium hydroxide is added? Explain in terms of Le
Châtelier’s principle.
b. Why does the yellow precipitate reappear when nitric acid is added?
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5. Lead(II) ion reacts with varying amounts of hydroxide ions in different ways. A small amount of hydroxide
produces a precipitate; an excess of hydroxide ion produces a complex ion:
Pb2+ + 2OH–
Pb2+ + 3OH–
Pb(OH)2 (s)
Pb(OH)3– (complex ion in solution)
a. Why does the precipitate disappear when excess NaOH is added? Explain in terms of Le Châtelier’s principle.
b. Why does the precipitate reappear when a little nitric acid is added?
c. Why does the precipitate disappear when excess nitric acid is added?
6. Explain the observations in Step 6 by writing net ionic equations to describe each reaction of silver ion (Ag+) you
observed.
7. Explain why BTB is green in neutral water solution. In terms of Le Châtelier’s principle, why does it turn yellow
when acid is added, and blue when base is added?
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NOW IT’S YOUR TURN!
1. Repeat Step 1 with other indicators such as phenolphthalein (HPhen); bromphenyl blue (HBPB); methyl red
(HMR); alizarine yellow R (HAYR); and thymol blue (HTB). Explain in terms of Le Châletier’s principle what
happens with each when HCl and NaOH are added. Write net ionic equations to describe their reactions in water.
2. Repeat the precipitation reaction of lead nitrate with potassium iodide, as in Step 4. Can you alternately add
NaOH and HNO3 to change the precipitates? Explain.
3. Repeat the precipitation reaction of lead nitrate with sodium hydroxide, as in Step 5. Name two ways you can
make the precipitate disappear. Prove your answer by carrying out the experiments.
4. Perform the following experiment and explain your results in terms of Le Châtelier’s principle.
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a. Mix one drop FeCl3 with one drop KSCN. What happened? Write an equation to explain.
b. Add 10 drops water, separate into four parts with a soda straw, and add the indicated solutions. Tell what
happens in each case. Explain in terms of the respective equilibria.
c.
d.
e.
f.
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