Lab Activity
11 Le Chatelier’s Principle
Background
As a chemistry student, you are probably quite used to reading, writing, and balancing
equations that represent chemical reactions. Symbols that once seemed unfamiliar
may now be easy to read. For example, read the following equation:
H 2 g  + O 2 g   H 2O  l 
Hydrogen gas and oxygen gas are reacting, and the arrow symbol is read as “yields”
or “goes to”. The equation makes it clear that liquid water is the product of the
reaction. Usually when reading this type of equation you may assume that once the
reaction takes place, the reactants have been used up, and the product now exists in
place of the substances that were present at the start.
Through electrolysis, liquid water can also be decomposed back into hydrogen and
oxygen gas. The equation describing this process is the reverse of the equation
above:
H2O l   H2 g  + O2 g
Virtually all reactions can occur both forward and backward (if the conditions are
right). Generally we don't notice this because one of the two reactions is faster, or the
other reaction can only happen at extreme conditions, etc. Some reactions, however,
have forward and backward reactions that occur at similar speeds and under similar
conditions. In such a reaction, a situation will occur in which the rate of the forward
reaction is equal to the rate of the backward reaction. At that point the system is said
to be at equilibrium. In this state it appears that the reaction has stopped since the
amount of each reactant and product stops changing (you are creating the products
exactly as fast as you are using them).
You have already learned that the rate of a reaction depends on several factors
including the concentration of the reactants. As the concentration of the reactants
rises, the number of collisions between reactant particles increases and therefore the
rate increases as well. If the concentration is decreased, the number of collisions
decreases and the rate decreases as well. This idea applies to both forward and
backward reactions.
Henri Louis Le Chatelier put these two ideas together to explain what happens when a
system at equilibrium is disturbed. Le Chatelier's Principle states that when a system
is at equilibrium and something is done to stress the system, the reaction will shift in
the direction that relieves the stress. In English that means that if you change the
conditions of the system, the reaction will no longer be at equilibrium, and the reaction
will either be running forward faster than backward or vice versa.
A specific example will make it easier to understand. When the reaction
N2 (g) + 3 H2 (g) l 2 NH3 (g)
is at equilibrium, the reaction appears to stop because it is producing ammonia as fast
as the ammonia is being broken down. If extra nitrogen is added to the container,
there will be more N2/H2 collisions and the forward reaction will speed up. As a result
the reaction will be producing ammonia faster than the ammonia is broken down and
the amount of ammonia will rise. At the same time, as the reaction runs forward the
amount of nitrogen will decrease. As a result, the addition of nitrogen causes the
reaction to run forward and to use up the nitrogen added.
If nitrogen were removed from the same system, the forward reaction would slow
down, the amount of ammonia would decrease (it is being broken down faster than it
is being produced) and additional nitrogen would be made. In other words, if you
remove nitrogen the reaction will shift in the direction that produces more nitrogen (in
essence, replacing what you took away).
In this lab you will look at several equilibrium systems and will measure the result of
disturbing the equilibrium.
This exploration will focus on two different equilibrium systems: in part A, the
dissociation of acetic acid in water, and in part B the dissociation of monohydrogen
phosphate in water. In each case, after measuring the initial pH of the equilibrium
system, you will “stress” the equilibrium system by either increasing or decreasing the
concentration of one of the products.
Part A of this Lab will focus on the dissociation of acetic acid in water:
+
–
CH 3 COOH  aq  + H 2 O  l  l H 3 O  aq  + CH 3 COO  aq 
You will be adding the strong electrolyte sodium acetate (CH3COONa) to this solution.
Recall that strong electrolytes dissociate in solution. What ions will form? Which of the
ions that form will affect the equilibrium of acetic acid in water, and which ion will act
as a “spectator ion”, one that does not change or take part in the reaction?
Part B of this lab will focus on the dissociation of monohydrogen phosphate ion
(HPO4-2) in water. Write the reaction for the dissociation of this compound in water.
Adding zinc chloride, a soluble ionic compound, to this solution will supply zinc and
chloride ions:
+2
-
ZnCl 2  aq   Zn  aq  + 2Cl  aq 
Zinc phosphate is insoluble. Thus the addition of the zinc ions to the solution will
effectively remove the phosphate ions from the equilibrium. What will happen to the
equilibrium when one of the products is removed from the solution? Keep in mind
that there will be a spectator ion in this reaction as well.
You will also be adding sodium phosphate to the solution. Write the reaction for the
dissociation of this compound in water.
How do you think that this will effect the equilibrium? What will the spectator ion be?
Lastly, you will be adding sodium chloride to each solution. Write the reaction for the
dissociation of this strong electrolyte.
What effect will this addition have on each reaction and why?
Purpose
In this lab, you will Use a Chemistry Sensor and pH electrode to determine how a
reversible reaction in equilibrium responds to changes in chemical concentrations.
Materials
PASCO & Other Equipment
PASPORT Xplorer GLX
graduated cylinder, 25-mL
PASPORT Chemistry Sensor
wash bottle and waste container
pH electrode
protective gear
beaker, 150-mL
Consumables
0.5 M acetic acid (CH3CO2H), 50.0 mL
0.1 M zinc chloride (ZnCl2), 10.0ml
0.5 M sodium acetate (CH3CO2Na), 10.0 mL
0.1 M sodium phosphate (Na3PO4), 7.0ml
0.10 M sodium monohydrogenphosphate (sodium phosphate dibasic) (Na2HPO4) 50.0 mL
0.5M sodium chloride (NaCl)
Safety Precautions
Wear safety glasses and follow standard laboratory safety procedures.
•
Caution: Ammonia gives off strong fumes. Avoid breathing vapors; work in a
fume hood if possible.
•
Do not dispose of lead solutions or lead salts by pouring down the drain! Follow
your teacher’s instructions for correct disposal.
•
Be sure to wash your hands if any solutions are spilled onto your skin.
Pre-Lab Questions
Write and balance the equation for the dissociation of sodium acetate in water.
1)
Predict the effect of adding sodium acetate to the equilibrium system of acetic
acid dissociated in water solution.
2)
Write equation for the dissociation of monohydrogen phosphate in water.
3)
Write a balanced equation predicting the reaction that will occur with the available
hydroxide ions, and explain how the ammonium hydroxide equilibrium will be
reestablished.
4)
Identify the spectator ion for each equilibrium reaction condition.
5)
What effect will adding sodium chloride have on this equilibrium and why?
Procedure
Equipment Setup
Obtain 50.0 mL of 0.5 M acetic acid solution in a 150-mL beaker.
1)
Measure 10.0 mL of 0.5 M sodium acetate in a 25-mL graduated cylinder.
2) Measure 10.0 mL of 0.5M sodium chloride into a separate graduated cylinder
Xplorer GLX Setup
Connect the Chemistory Sensor to Port #1 on the GLX.
1)
Connect the pH electrode to the Chemistry Sensor.
2)
Calibrate the pH Sensor if necessary. Refer to the pH electrode Sensor
Information pages for instructions on calibrating the sensor.
3)
From the Home screen, press
to open the Graph display.
Record Data
Part A
Place the pH electrode in the 150-mL beaker containing the acetic acid.
1) Press
to begin recording the pH.
Watch the Graph display for the pH to stabilize. This is the initial equilibrium pH of the acetic acid solution.
2) Pour the sodium acetate solution from the graduated cylinder into the beaker and
stir gently. Watch the Graph display for the pH to stabilize. This is the second
equilibrium pH.
3) Pour the sodium chloride solution from the graduated cylinder into the beaker and
stir gently. Watch the Graph display for the pH to stabilize. This the final pH.
4) Press
again to end data collection
Part B
Follow the same steps above, except use 50.0mL of 0.1 M sodium monohydrogen phosphate, 10.0mL of 0.1 M zinc
chloride, 7.0 mL of 0.1M sodium phosphate and 10.0 mL of 0.5 M sodium chloride
1) Place the pH electrode in the 150-mL beaker containing the monohydrogen
phosphate.
2) Press
to begin recording the pH.
Watch the Graph display for the pH to stabilize. This is the initial equilibrium pH of the monohydrogen phosphate
solution.
3) Pour the zinc chloride solution from the graduated cylinder into the beaker and stir
gently. Watch the Graph display for the pH to stabilize. This is the second
equilibrium pH.
4) Pour the sodium phosphate solution from the graduated cylinder into the beaker
and stir gently. Watch the Graph display for the pH to stabilize. This is the third
equilibrium pH.
5) Pour the sodium chloride solution from the graduated cylinder into the beaker and
stir gently. Watch the Graph display for the pH to stabilize. This is the final
equilibrium pH.
6) Press
again to end data collection.
Analyze
Record calculations in your data table as you complete your analysis.
1)
In the Tools menu (
), select the Smart Tool. Use the Arrow Keys to move
the Smart Tool Cursor in order to determine the initial pH of the acetic acid, the
final equilibrium pH of the acetic acid, the initial pH of the monohydrogen
phosphate, and the final pH of the monohydrogen phosphate equilibrium. Record
these values in your Data Table.
Data Table
pH of initial equilibrium acetic acid
pH of second equilibrium acetic acid
pH of final equilibrium acetic acid
pH of initial equilibrium monohydrogen
phosphate
pH of second equilibrium monohydrogen
phosphate
pH of third equilibrium monohydrogen
phosphate
pH of final equilibrium monohydrogen
phosphate
Analysis and Synthesis Questions
By looking at the graph for Part 1, explain how the addition of the acetate ion shifted the equilibrium. What is the
evidence for this shift?
1)
By looking at the graph for Part 2, explain how the addition of the lead(II) ion
shifted the equilibrium. What is the evidence for this shift?
2)
What effect did the sodium chloride have and why?
3)
Below are some common misconceptions about chemical equilibrium. Rewrite
each statement to accurately reflect the nature of chemical equilibrium.
a
"Reactions always go to completion."
b
“Reactions at equilibrium have equal concentrations of reactants and
products.”
c
"Equilibrium is a fixed, static condition."
d
"The left hand side of a reaction in equilibrium happens first, then the right
hand side."