LeChateliers Lab

Stresses Applied to Chemical Equilibrium
Many chemical reactions do not go to completion. Rather, they come to a point of
chemical equilibrium before the reactants are fully converted to products. At the point of
equilibrium, the concentrations of all reactants remain constant with time. In this experiment, you
will investigate how outside forces acting on a system at equilibrium provoke changes within the
system (Le Chatelier’s principle).
Early in the study of chemical reactions, it was noted that many chemical reactions do not
produce as much product as might be expected, based on the amounts of reactants taken
originally. These reactions appeared to have stopped before the reaction was complete. Closer
examination of these systems (after the reaction had seemed to stop) indicated that there were
still significant amounts of all the original reactants present. Quite naturally, chemists wondered
why the reaction had seemed to stop, when all the necessary ingredients for further reaction were
still present.
Some reactions appear to stop because the products produced by the original reaction
themselves begin to react, in the reverse direction to the original process. As the concentration of
the products begins to build up, product molecules will react more and more frequently.
Eventually, as the speed of the forward reaction decreases while the speed of the reverse
reaction increases, the forward and reverse processes will be going on at exactly the same rate.
Once the forward and reverse rates of reaction are identical, there can be no further net change
in the concentrations of any of the species in the system, at this point, a dynamic state of
equilibrium has been reached. The original reaction is still taking place but is opposed by the
reverse of the original reaction also taking place. In this experiment, you will study changes
made in a system already in equilibrium, by the reference to Le Chatelier’s principle.
Le Chatelier’s principle states that, if we disturb a system that is already in equilibrium,
then the system will react so as to minimize the effect of the disturbance. This is most easily
demonstrated in cases where additional reagent is added to a system in equilibrium, or when one
of the reagents is removed from the system in equilibrium.
Solubility Equilibria
Suppose we have a solution that has been saturated with a solute: This means that the
solution has already dissolved as much solute as possible. If we try to dissolve additional solute,
no more will dissolve because the saturated solution is in equilibrium with the solute:
Solute + Solvent  Solution
Le Chatelier’s principle is most easily seen when an ionic solute is used: Suppose we have a
saturated solution of sodium chloride, NaCl. Then
NaCl  Na+ (aq) + Cl- (aq)
will describe the equilibrium that exists. Suppose we then try adding an additional amount of one
of the ions involved in the equilibrium: For example, suppose we added several drops of
concentrated HCl solution (which contains the chloride ion at high concentration). According to
Le Chatelier’s principle, the equilibrium will shift so as to consume some of the added chloride
ion. This would result in a net decrease in the amount of NaCl that could dissolve. If we watched
the saturated NaCl solution as the HCl was added, we should see some of the NaCl precipitate
as a solid.
LeChatelier’s Principle
Complex Ion Equilibria
Oftentimes, dissolved metal ions will react with certain substances to produce brightly colored
species called complex ions. For example, ion (III)reacts with the thiocyanate ion (SCN -) to
produce a bright red complex ion:
Fe3+ + SCN-  [FeNCS2+]
This is an equilibrium process that is easy to study because we can monitor the bright red color of
[FeNCS2+] as an indication of the position of the equilibrium: If the solution is very red, there is a
lot of [FeNCS2+] present; if the solution is not very red, then there must be very little [FeNCS 2+]
Using this equilibrium, we can try adding additional Fe 3+ or additional SCN- to see what
effect this has on the red color according to the Le Chatelier’s principle. We will also ass a
reagent (silver ion) that removes SCN- from the system to see what effect this has on the red
color, and NaOH, which precipitates Fe3+ as iron (III) hydroxide and removes iron from the
Acid/Base Equilibria
Many acids and bases exist in solution in equilibrium sorts of conditions. This is particularly true
for the weak acids and bases. For example, the weak base ammonia is involved in an
equilibrium in aqueous solution
NH3 (aq) + H+ (aq)  NH4+ (aq)
Once again we can use Le Chatelier’s principle to play around with this equilibrium: We will try
adding more ammonium ion or hydrogen ion to see what happens. Since none of the
components of this system is itself colored, we will be adding an acid/base indicator that changes
color with pH, to have an index of the position of the ammonia equilibrium.
The indicator we will use is phenolphthalein, which is pink in basic solution and colorless
in acidic solution.
HIn (aq)  H+ (aq) + In- (aq)
Phenolphthalein itself is a weak acid in aqueous solution, in which the colorless (HIn) and
colored (In-) forms are in equilibrium. Changes in the ammonia system are reflected by a shift in
the phenolphthalein equilibrium.
Another easily studied equilibrium which is dependent on the acidity of the system is that
between chromate ion, CrO4 2-, and dichromate ion, Cr2O7 2-. If a strong acid is added to a
solution of potassium chromate, the intensely yellow chromate ion is converted to the bright
orange dichromate ion
2CrO4 2- (aq) + 2H+ (aq)  Cr2O7 2- (aq) + H2O
Similarly, if a strong base is added to a solution containing the range dichromate ion, the
equilibrium is shifted in favor of the yellow chromate ion
Cr2O7 2- (aq) + 2OH- (aq)  2CrO4 2- (aq) + H2O
Beginning students often think these two reactions represent oxidation – reduction processes
(chromate and dichromate often are encountered in redox reactions); a careful examination of
redox numbers, however, will show that there is no transfer of electrons involved in either
LeChatelier’s Principle
Oleic acid (C18H34O2) is a weak, long chain organic acid found in fats and oils. It is not
very soluble in water. The effect of added strong acid and strong base on the ionization
equilibrium of oleic acid will be investigated.
Safety Precautions
 Safety eyewear approved by your institution must be worn at all times while you are in
the laboratory, whether or not you are working on an experiment.
 Concentrated ammonia is a strong respiratory and cardiac stimulant. Use concentrated
ammonia only in the fume exhaust hood.
 Concentrated hydrochloric acid is severely damaging to skin, eyes, and clothing, and its
vapor is highly toxic and irritating. Use concentrated HCl only in the exhaust hood, and
handle the bottle with a towel or gloves to protect your hands. If HCl is spilled on the
skin, wash immediately and inform the instructor.
 Iron (III) chloride, copper (II) sulfate, and potassium thiocyanate may be toxic in ingested.
Wash after use.
 Chromium compounds (dichromate and chromate ions) are highly toxic, may burn skin
and eyes, and are suspected mutagens/carcinogens. Wear gloves during their use.
Inform the instructor of any spills.
 Dispose of all reagents as directed by the instructor
Apparatus/Reagents Required
Saturated sodium chloride solution
12 M HCl
concentrated ammonia solution
ammonium chloride
phenolphthalein indicator
0.1 M solutions of potassium chromate and potassium dichromate
0.1 M barium chloride solution, oleic acid.
0.1 M KSCN,
0.1 M AgNO3,
6 M NaOH
Record all data and observations directly on the report pages in ink.
A. Solubility Equilibria
Obtain 2 mL of saturated sodium chloride solution in a small test tube. This solution was
prepared by adding solid NaCl to water until no more would dissolve. Then the clear
solution was filtered from any undissolved solid NaCl.
Add 10 drops of 12 M HCl (caution!) to the NaCl solution. A small amount of solid NaCl
should form and precipitate out of the solution. The crystals may form slowly, and may
be very small. Examine the test tube carefully.
On the lab report sheet, describe what happens in terms of LeChatelier’s principle. What
is the “stress” applied? In which direction does the equilibrium shift?
LeChatelier’s Principle
B. Complex Ion Equilibria
Prepare a stock sample of the bright red complex ion [FeNCS2+] by mixing 2 mL of 0.1 M
iron (III) chloride and 2 mL of 0.1 M KSCN solutions. The color of this mixture is too
intense to use as it is, so dilute this mixture with 100 mL of water.
Pour about 5 mL of the diluted red stock solution into each of the five test tubes. Label
the test tubes as 1,2,3,4, and 5.
Test tube 1 will have no change made in it so that you can use it to compare its color with
what will be happening in the other test tubes.
To test tube 2, add about 1 mL of 0.1 M FeCl3 solution.
To test tube 3, add about 1 mL of 0.1 M KSCN solution.
To test tube 4, add AgNO3 solution dropwise until a change becomes evident. Ag+ ion
removes SCN- ion from the solution as a solid (silver thiocyanate).
To test tube 5, add 6 M NaOH until a change is evident. What is the precipitate that
forms? Why did the red color of the iron/thoicyanate complex fade when NaOH was
Describe the intensification or fading of the red color in each test tube in terms of
LeChatelier’s principle. What is the “stress” applied in each case? In which direction
does the equilibrium shift?
C. Acid Base Equlibria
1. Ammonia/ Phenolphthalein
 In the exhaust hood, prepare a dilute ammonia solution by adding 2
drops of concentrated ammonia (caution!) to 25 mL of water.
Add 3 drops of phenolphthalein to the dilute ammonia solution, which
will turn pink (ammonia is a weak base, and phenolphthalein is pink in
basic solution).
Place about 5mL of the pink dilute ammonia solution into each of two
test tubes.
To one of the test tubes, add several small crystals of ammonium
chloride (which contains the ammonium ion, HN4+). What happens to
the color of the solution?
To the other test tube, add a few drops of 12 M HCl. What happens to
the color of the solution?
Describe what happens to the pink color in terms of how LeChatelier’s
principle is affecting the ammonia and phenolphthalein equilibria.
What is the “stress” applied in each situation? In which direction do
the equilibria involved shift?
LeChatelier’s Principle
2. Chromate/Dichromate
Wear gloves when performing this section.
Place 5 drops of 0.1 M potassium chromate in each of two semimicro test tubes.
To one sample of chromate ion, ass 1 drop of 12 M HCl and stir. To the second test
tube, ass one drop of 6 M NaOH and mix. Account for any color changes that occur in
terms of LeChatelier’s principle. What is the “stress” applied? In which direction does the
equilibrium shift?
Dispose of the chromium compounds as directed by the instructor. Do not pour down the
Place 5 drops of 0.1 M potassium chromate in a semimicro test tube. Add 2 drops of 0.1
M barium chloride solution. What happens? Barium chromate is highly insoluble in
water. Add 2 drops of 12 M HCl to the test tube and mix. What happens? Is barium
dichromate more soluble in water than is barium chromate? Explain.
3. Oleic acid
Oleic acid is a weak organic acid which is not very soluble in water. When oleic acid is placed in
water, most of the oleic acid will not mix with the water, but the portion that does dissolve ionizes
in a manner similar to other weak acids
HA  H+ (aq) + A- (aq)
Place about 10 drops of water in a semimicro test tube, and add a drop of oleic acid. Mix
and allow to settle. Make a note of the appropriate thickness of the oleic acid layer in the
test tube.
Add 2 drops of 12 M HCl to the test tube and mix. What happens to the thickness of the
oleic acid layer? Explain your observation in terms of the effect of LeChatelier’s principle
on the ionization equilibrium for oleic acid.
Add 6-8 drops of 6 M NaOH to the test tube and mix. What happens to the thickness of
the oleic acid layer? Explain your observations in terms of the effect of LeChatelier’s
principle on the ionization equilibrium for oleic acid.
LeChatelier’s Principle
Name ______________________________
Stress Applied to Chemical Equilibrium
1. Perhaps the most studied equilibrium system is that by which gaseous ammonia, NH3, is
produced from elemental hydrogen and nitrogen gases. Write the balanced chemical
equation for this reaction.
2. Suppose the ammonia synthetic reaction in Question 1 has reached a state of equilibrium at
a particular temperature. Tell what effect on the net amount of ammonia obtained each of the
following disturbances to the equilibrium system will have, compared to a system in which no
change is made.
a. ...additional nitrogen gas is pumped into the system
b. ...ammonia gas is liquefied and removed from the system as it forms
...the reaction is compressed to a smaller total volume
d. ...a very efficient catalyst is found for the reaction
LeChatelier’s Principle
Stress Applied to Chemical Equilibrium
A. Solubility Equilibria
Effect of adding HCl to saturated NaCl observation
B. Complex Ion Equilibria
Effect of adding additional Fe3+ to [FeNCS2+] observation
Effect of adding Ag+ to [FeNCS2+] observation
C. Acid/Base Equilibria
1. Ammonia/Phenolphthalein
Effect of adding NH4+ observation
2. Chromate/Dichromate
Potassium chromate
Effect of adding HCl observation
LeChatelier’s Principle
Effect of adding NaOH observation
Potassium Dichromate
Effect of adding HCl observation
Effect of adding NaOH observation
Effect of adding barium chloride observation
3. Oleic Acid
Effect of adding HCl to oleic acid observation
Effect of adding NaOH to oleic acid observation
LeChatelier’s Principle
1. Explain how the solubility of a salt represents a situation of dynamic equilibrium. What is it
that is in equilibrium?
2. Students often think that the reactions of dichromate and chromate ion discussed in this
experiment are oxidation/reduction reactions (chromate and dichromate are often involved in
redox). Show, by determining the oxidation states of the reactants and products, that there is
no transfer of electrons in these reactions.
3. Oleic acid is a long chain (18 carbon atoms) organic acid called a “fatty acid”. You found that
it did not dissolve very well in water, but did dissolve readily in sodium hydroxide.
+ NaOH (aq)  Na+A- (aq) + H2O (l)
The sodium salts of fatty acids (such as sodium oleate) are given a special name and have very
useful properties. Use your textbook or an encyclopedia of chemistry to describe what is special
about such compounds.
LeChatelier’s Principle
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