CHEMICAL
EQUILIBRIUM
Teacher's Guide
Steady Unsteadiness
Objectives
Students should be able to:
1 . State the Second Law of Thermodynamics.
2. Identify the implications of the regularity "energy runs downhill" to
chemical reactions.
3. Identify systems in a steady state.
4. Identify closed systems.
5. Describe at least two explanations for macroscopic properties
becoming constant during a chemical reaction.
6. Describe what is meant by the term dynamic chemical equilibrium.
Program Description
When the flask is shaken, oxygen dissolves in the water and oxidizes the
colorless reduced form of methylene blue, reforming the blue dye. Below
is a simplified mechanism for the system.
Details can be obtained from J.A. Campbell's article "Kinetics - Early
and Often" listed on page 5.
Explanation for Activity 2: The Thionine System
Program 1 begins with a toboggan sliding down a hill, illustrating the
tendency for most changes to involve movement from a state of high
energy to one of low energy. The application of this regularity to
chemical systems suggests that only exothermic reactions are possible.
It also suggests that chemical reactions must stop. The fallacy in their
conclusions is illustrated and the concepts of steady state and dynamic
equilibrium are introduced.
Thionine is protonated by the sulfuric acid, and then is electronically
excited by the light source. The excited thionine ion oxidizes the iron(II)
to iron(III). The reduced form of thionine is colorless. When the light
source is turned off the excited thionine decays to its ground state. This
reverses the iron reaction; iron(III) is reduced to iron(II), and the purple
color of the oxidized thionine returns.
A more detailed explanation is available in L.J. Hardt's article "The
Photochemical Reduction of Thionine" listed on page 5.
Before Viewing
After Viewing
Use Activity 1, The Blue Bottle, and/or Activity 2, The Thionine System,
to demonstrate that chemical reactions can be reversed. Most students
believe this cannot be done, so this is time well spent. Note that the
development of the accepted model for either of these activities is
beyond the ability of most students at this stage. However, if the activity
is done in groups of four, most groups should come up with a model that
i nvolves a set of reversible reactions.
Use Activity 3 to show that reactions proceed and reach a point
where macroscopic properties are constant. Ask students to suggest
plausible explanations for the constant macroscopic properties, then
show Program 1.
Use Activities 4 and 5 to illustrate the difference between steady state
and dynamic equilibrium. Demo 1 shows a typical steady state system,
with the rate of water flowing into the can equal to the rate of water
flowing out. The candle and burner flames also are steady state systems,
but the input and output are not as obvious. Have students explain the
"balance" in the systems.
Demo 2 involves two long-term experiments. The best way to treat
this demonstration is to set it up for next year and use the given observations for your current classes.
Activity 5 is a logical stepping stone to Program 2. It is an excellent
analogy of dynamic equilibrium. Working in pairs, students transfer
colored water from their container into their partner's. By starting with
different volumes of water and using different capacity vessels to
transfer the water, the situation can be used to illustrate many of the
features of a chemical system.
Explanation for Activity 1: The Blue Bottle
Glucose, a weak acid reacts with the hydroxide to form the glucoside ion
which reduces the blue form of methylene blue to a colorless form.
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2
Activity 5 is written as a student experiment. If you wish to demonstrate this activity, use beakers of different sizes to transfer the water
and aquaria or large battery jars filled with different volumes of colored
water.
Explanation for Activity 4, Demo 2:
Long-Term Experiments
Set-up 1
Since the volumes continued to decrease, water must have continuously
evaporated from the solutions. Thus the system must have been open to
gases, in this case water vapor escaping from the jar. Since the
phenolphthalein reddened in the cylinder labelled "B," then some of the
sodium hydroxide was neutralized, presumably by HCI(g) escaping from
cylinder "A."
Note: Phenolphthalein has two color changes as pH increases: colorless
to pink between 8-10 and then back to colorless around pH 14.
Set-up 2
because the solution in beaker "C" was absorbing water at a much faster
rate than "D." This was caused by the large difference in concentration.
The solution in beaker "C" was 615 times more concentrated than that
in "D" at the start. Thus the beakers tended to exchange water. The
exchange should have continued until the concentrations were equal.
The rate of evaporation from "D" should have started out fast and then
decreased until the concentrations were equal. But in this case the
situation was complicated by the water vapor diffusing into the jar. When
the rate of diffusion was greater than the rate of absorption by the
solution in beaker "C," the volume in "D" should have slowly begun to
rise. As the concentration decreased, the rate of absorption also should
have decreased; however, the concentration of the water was getting
larger at the same time so the rate of evaporation should have increased
until, at 21 months, the rate of evaporation and condensation became
equal, or perhaps all evaporation and condensation stopped. More
details can be obtained from L.W. Bixby's article "Long-Term Chemical
Reactions" listed on page 5.
Since the volumes of both beakers overflowed, the system must have
been open to water vapor. The level in beaker "D" dropped initially
Activities
Activity 1: The Blue Bottle
Apparatus
20 g of sodium hydroxide
20 g of glucose
1.5 mL of 1 % alcoholic methylene blue
3. Using "OH' - " to represent the sodium
hydroxide, "G" to represent the glucose, and
"MB" to represent the methylene blue, write
equations that will account for your
observations.
4. Make predictions and design experiments to
test the models you developed in step 3.
Check with your teacher before carrying these
out.
1 L of water
250 mL Erlenmeyer flask with rubber stopper
Note: Dissolve the ingredients in the litre of water.
Add the sodium hydroxide just before using.
Method
1 . Fill the Erlenmeyer flask a little less than half
full with the solution. Stopper the flask
i mmediately.
2. Shake the flask vigorously, then place it on a
white sheet of paper and observe what
happens. Repeat several times.
10 mL graduated cylinder
600 mL beaker
Light source (overhead projector or goose neck
lamp)
Method
Discussion
2.
Develop a model for the blue bottle experiment.
I ndicate the evidence on which your model is
based.
3.
Activity 2: The Thionine System
4.
Apparatus
I ron(II) sulfate heptahydrate
0.001 mol/L thionine
5.
Add 10 mL of the thionine solution, 10 mL of
the sulfuric acid solution, and sufficient water
to bring the solution to 500 mL.
Thoroughly mix in 2.0 g of iron(II) sulfate
heptahydrate.
Use the overhead projector to light the solution
from below or use the goose neck lamp with a
250 W photoflood bulb to light it from above.
Turn the light source on and observe what
happens. Then turn it off and observe. Repeat
several times.
Using "S" to represent the sulfuric acid, "T" to
represent the thionine, and "Fe 2+" to
represent the iron(II) sulfate in solution, write
equations that will account for your observations.
6. Make predictions and design experiments to
test the models you developed in step 5.
Check with your teacher before carrying these
out.
Discussion
Develop a model for the thionine system.
I ndicate the evidence on which your model is
based.
Activity 3: Acid/Base
Demonstrations
Demo 1
Slowly pour equal volumes of 1 mol/L NaOH and
HCI solutions containing the indicator bromthymol
blue into a beaker placed on the overhead
projector. Repeat for different but equal volumes
of NaOH and HCI. Try to explain why the color of
the mixture remains the same.
in the manometer. Again, the same two explanations are plausible.
Use a syringe to inject stock hydrochloric acid
containing the indicator bromthymol blue into
another example of the set-up. Explain the
changes in the manometer. Again, two explanations are plausible.
Now add a little of the hydrochloric acid to the
apparatus containing the ammonium hydroxide.
Similarly, add a little of the ammonium hydroxide to
the apparatus containing the hydrochloric acid. In
this case the chemical reaction is obvious when
the white smoke forms. Account for any changes
i n the indicator and manometer.
Note: It will be necessary to fiddle with the
concentrations and volumes of the acid and base
to get reasonable changes in the levels of the
manometer. Be sure to try this experiment before
you demonstrate it for the class.
Activity 4: Steady State/Dynamic
Equilibrium Demonstrations
Demo 1
Place a burning candle and a lit Bunsen burner on
the demonstration desk. Thirdly, set up the
apparatus as shown in Figure 2. Ask students to
account for the "balance" in the three systems.
Note: Use a fairly crude knife edge for the
fulcrum, otherwise it is very difficult to maintain the
balance. Ask students to compare the three
systems. How are they the same? How are they
different?
Figure
2
Note: There are two plausible explanations: that
the reaction has stopped, and that there are
opposing reactions occurring at equal rates.
Demo 2
Use a syringe to inject stock ammonium hydroxide
containing the indicator bromthymol blue into the
apparatus shown in Figure 1. Explain the changes
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Demo 2
Make a transparency of the two experimental setups described below, or better still, actually set
them up. Include the following:
Set-up 1
Forty millilitres of 0.3 mol/L HCI was added to a
100 mL graduated cylinder labelled "A." Ten millilitres of a 1 mol/L NaOH solution containing the
indicator phenolphthalein was added to a second
1 00 mL graduated cylinder labelled "B." The two
cylinders were sealed in a large jar and observed
weekly for two years.
The solutions in both graduated cylinders
decreased continually over the period of 24
months. After 13 months the solution in cylinder
B reddened.
Set-up 2
One hundred and fifty millilitres of saturated
( 6.15 mol/L) NaCl was placed in a 250 mL beaker
l abelled "C." One hundred and fifty millilitres of
0.01 mol/L NaCl was placed in a second 250 mL
beaker labelled "D." The two beakers were sealed
i n a large jar and observed weekly for two years.
After 13 months beaker "C" was filled and then
overflowed by 30 mL or more into the jar. The
l evel in beaker "D" initially dropped, but it too also
eventually overflowed. After 21 months there was
no apparent change.
Ask students to come up with plausible explanations for these observations. Help them focus
their attention on the principles learned in Program 1 with questions such as "How is the system
closed?" "How is the system open?" "Were the
systems at equilibrium?" "At a steady state?"
products (p). I n this way you can simulate the
reaction where one reactant decomposes into two
products.
If possible, move the fulcrum position to one
side. This enables you to simulate a situation
where there is more product than reactant at
equilibrium.
Add the appropriate number of tiles to each pan
to achieve a balance. Then challenge the students
to think of another way the balance could be
maintained. Obviously, if one (r) tile were removed
from one side and two (p) tiles were removed from
the other side, the system would remain balanced.
But that is not how a chemical system works.
Reactants are changed into products, and as soon
as that happens the system is no longer balanced.
However, if as one (r) tile were transferred to the
product side, two (p) tiles were removed and one
(r) tile was added to the reactant side, balance
would be maintained. This would simulate a steady
state with the continuous removal of products and
the continuous addition of reactants.
Another way to maintain balance is, of course,
to switch one reactant tile for two product tiles. In
this way both the reaction and its reverse would
have to occur and their rates would have to be
equal.
Activity 5: A Model for
Chemical Equilibrium
I n this activity you will simulate a system
approaching chemical equilibrium. Your task is to
determine if the simulation is a good model for a
chemical system at equilibrium.
Apparatus
Demo 3
Use an equal arm balance and some small bathroom tiles to simulate a system at equilibrium.
Label one set of tiles reactants (r) and, if possible,
another set - half the mass of the former =
Two 150 mL beakers
Two plastic rulers
Two pieces of glass tubing of different diameters,
10-15 cm long
1 0 mL graduated cylinder
Water with blue food coloring
Water with yellow food coloring
Method
1 . Put some blue colored water in one beaker
and some yellow colored water in the other.
The only restriction is that the total volume of
the two is 100 mL.
2. Work in pairs. Transfer water from B, the blue
beaker, using one of the pieces of glass
tubing, to Y, the yellow beaker. At the same
ti me your partner should transfer water from Y
to B using the other piece of glass tubing. Be
sure to keep the glass tube vertical at all times.
3. Use the rulers to record the height of the
water in each beaker after every five
exchanges.
4. Continue transferring water until you have
three successive readings that are the same.
5. When the heights remain constant measure
the volume transferred by each glass tube.
6. Repeat steps 1-5 for different volumes of blue
and yellow water, but keep the total volume
100 mL.
Observations
TIME
HEIGHT OF WATER HEIGHT OF WATER
BLUE BEAKER
YELLOW BEAKER
( 5 exchanges)
( mm)
( mm)
1
2
3
etc.
Further Reading
Discussion
1. What are the following analogous to in a
chemical reaction? The height of the water in
the two beakers; the transfer of water from
one beaker to another; two pieces of glass
tubing of different diameters; the coloring in
the water; the volume of water transferred; the
volume of water transferred when the height
remains constant.
2. What evidence suggests that you reached a
point of equilibrium in this activity? You should
be able to think of two key observations.
3. What property of the system determines the
final height of the water in the beakers? What
would this be analogous to in a chemical
reaction?
4. Plot height of water versus time in units of five
exchanges.
(a) Describe the graph in words.
(b) What is the significance of the horizontal
part of the graph?
(c) What would be the significance of the
slope of a tangent drawn at any point on
this graph?
(d) What happens to the slope of the tangent
as the system approaches equilibrium?
What does this imply?
(e) The implications of (d) contradict what is
actually happening; you can still see water
being transferred from one beaker to
another. Explain how it is possible for the
graph to suggest that the reaction has
stopped even though you can see it going
on.
5. Consider the following gaseous reaction:
Bixby, L.W. "Long-Term Chemical Reactions.
Parts I & If. " Chemistry. September and
October 1976.
Campbell, J.A. "Kinetics - Early and Often."
Journal of Chemical Education. Vol 40. No.
11. November 1963.
(a)
(b)
(c)
Plot a graph of the data.
Describe the graph in words.
Explain why the graph has the shape it
has.
(d) What is the significance of the horizontal
part of the graph?
(e) What would be the significance of the
slope of a tangent drawn at any point on
this graph?
(f) What happens to the slope of the tangent
as the system approaches equilibrium?
What does this imply?
( g) Use the analogy to explain the changes in
the slope of the tangents.
6. Is this analogy a good model for chemical
equilibrium? Explain your answer.
Hardt, L.J. "The Photochemical Reduction of
Thionine: A Reversible Reaction." Journal of
Chemical Education. Vol 26. No. 25. 1949.
Parry, Robert W. et al. Chemistry: Experimental
Foundations. Englewood Cliffs, New Jersey:
Prentice-Hall, 1970.
Rowley, Wayne R. E. Matter in Balance: Chemical
Equilibrium. Toronto: Wiley, 1979.
Sienko, Michell J., and Robert A. Plane.
Chemistry. New York: McGraw-Hill, 1976.
Toon, E.R., and G.P. Ellis. Foundations of
Chemistry. New York: Holt, Rinehart and
Winston, 1973.
One mole of C12 was added to 2 mol of CO in a
1 L container. The concentration of the three
gases was measured using a spectrophotometer every minute. The following data were
recorded:
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6
Dynamic Equilibrium
Objectives
Students should be able to:
1. Describe a system at equilibrium macroscopically and microscopically.
2. Use the Kinetic Molecular Theory to explain why a "dynamic" model
for chemical equilibrium is the most reasonable explanation for the
observations.
3. Describe the types of motion of molecules in the gaseous phase.
4. Describe the changes in concentration of the reactants and products
as a chemical system approaches and reaches equilibrium.
5. Interpret the "double arrow" conventions.
6. Identify data that supports a dynamic model for chemical equilibrium.
Program Description
Program 2 begins with a review of the tendency towards minimum
energy and the observation that in a closed system a chemical reaction
appears to have stopped. The reaction betwee n N02 and Nz04 i s used to
l ustrate the reversibility of chemical reactions. A quantitative look at the
eaction between hydrogen and iodine follows. The analysis of the
, ystem shows both reactants and products are present when the
eaction appears to have stopped. It also shows that this point - the
Activities
Activity 1: Testing the
Dynamic Model
I n this activity you will gather indirect evidence to
support the dynamic model for chemical
equilibrium. It is not possible to see what is
happening on a microscopic scale, but it is
possibie to check if the reaction and its reverse
occur and if both reactants and products are
present at equilibrium. The two systems you will
study are:
point of equilibrium - can be reached by starting with the product,
hydrogen iodide. But the puzzle of how to describe the system at this
point remains. Has the reaction stopped, as the macroscopic properties
suggest, or is there a dynamic equilibrium with reactions occurring in two
directions? To resolve this puzzle students are reminded that a kinetic
model is used to describe matter which supports the dynamic
description of the equilibrium point. In addition, they are presented with
empirical data that is difficult to explain without assuming a dynamic
equilibrium between the forward and reverse reaction for the
hydrogen/iodine/hydrogen iodide system.
Before Viewing
Use the discussion of Demonstration 3 in Activity 4, and Activity 5 in
Program 1 as an introduction to Program 2. The questions that remain
should be "Are the analogies a good way to view chemical equilibrium?"
and "What evidence is there to support a dynamic model?"
After Viewing
Have students complete Activity 1. This enables them to gather indirect
evidence for the dynamic model and to practise using it to describe a
physical change and a chemical change.
The formation of a black precipitate when
sodium sulfide is added is used to identify lead
i ons in solution. The addition of silver nitrate and
the formation of a pale yellow precipitate is used
to identify the bromide ions in solution. The silver
bromide that is formed darkens when exposed to
li ght. Iodine vapor can be identified by its
characteristic pink color.
Apparatus
1 mol/L NaBr solution (102.9 g NaBr/L) * in
dropping bottles
Short-stemmed funnel and filter paper
Small test tube
150 mL beaker
Wash bottle filled with distilled water
Piece of acetate and a sheet of white paper
A glass rod
* All solutions must be made with distilled water.
Part B
Small corked test tube containing a few crystals of
iodine
Large beaker containing hot (70-90°C) water
Hot plate
Method
Part A
1 . Add approximately 1 cm of the sodium
bromide solution to the test tube.
2. Add an equivalent volume of lead nitrate
solution to the test tube. Caution: Wear
disposable gloves. The white precipitate that
forms is lead bromide.
3. Fold the filter paper into a cone and place it in
the funnel. The funnel is held by the 150 mL
beaker.
4. Stir the contents of the test tube and transfer
them onto the filter paper. Use the wash
bottle to ensure that all of the solid is
transferred to the filter paper.
5. Wash the lead bromide on the filter paper
twice using approximately 2 cm of distilled
water in the test tube each time.
6. Put the acetate sheet over a piece of white
paper.
7. Use the glass rod to transfer as much of the
l ead bromide as possible from the filter paper
to the acetate sheet.
8. Cover the solid lead bromide with distilled
water. Mix the solid and water with the glass
rod. Call this mixture 1. Let the mixture sit
until you need it in step 13.
9. On another section of the acetate sheet put a
drop of the lead nitrate solution. Add to it a
drop of sodium sulfide solution. Record any
changes.
10. On another section of the acetate sheet put a
drop of the sodium bromide solution. Add to it
a drop of the silver nitrate solution. Record
any changes.
1 On another section of the acetate mix a drop
of the sodium bromide and lead nitrate
solutions. Call this mixture 2. Allow the solid
to settle and then divide the mixture into two
equal parts using the glass rod.
12. To one part of the mixture add a drop of the
sodium sulfide solution. To the other part add
a drop of the silver nitrate solution. Record
any changes.
13. Divide the mixture produced in step 8
( mixture 1) into two equal parts. To one part
add a drop of the sodium sulfide solution. To
the other part add a drop of the silver nitrate
solution. Record any changes.
14. Dispose of the filter paper as directed by
your teacher. Rinse the acetate sheet, test
tube, beaker, and glass rod with running
water. Do not forget to wash your hands
after cleaning up. Lead compounds are
toxic.
Part B
1 . Use the hot plate to heat water until it is
between 70 0 and 90 1C. Several pairs of
students can use the same hot water bath.
2. Hold the bottom of the test tube in the hot
water until the amount of iodine vapor remains
constant.
3. Remove the test tube from the hot water and
observe what happens.
4. Repeat steps 2 and 3 several times.
5. Return the corked test tube to your teacher.
Do not attempt to clean out the test tube.
Observations
Part A
Mixture 1 is formed by adding distilled water to
l ead bromide.
Mixture 2 is formed by mixing lead nitrate and
sodium bromide.
Part B
Describe the original contents of the test tube and
any changes that occurred as a result of heating
and cooling the tube.
Discussion
Part A
1 . Describe a test for lead ions in solution. Write
a balanced chemical and ionic equation for
the test.
2. Describe a test for bromide ions in solution.
Write a balanced chemical and ionic equation
for the test.
3. Write a balanced chemical and ionic equation
to describe the formation of a white precipi tate when solutions of lead nitrate and
sodium bromide are mixed.
4. What are two plausible explanations for the
fact that the amount of white precipitate
formed when lead nitrate and sodium bromide
are mixed does not seem to change on
standing?
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8
5. What is implied by the observations made
when sodium sulfide and silver nitrate are
added to mixture 2?
6. Explain why it is hard to accept the
"stopped" model for equilibrium based on the
observations made when sodium sulfide and
silver nitrate are added to mixture 2.
7. What is implied by the observations made
when sodium sulfide and silver nitrate are
added to mixture 1 ?
8. Explain why the indirect evidence supplied by
the observations made when sodium sulfide
and silver nitrate are added to mixture 1 build
our confidence in the dynamic model.
9. Write a balanced chemical and ionic equation
for a reaction that would account for the
observations made when sodium sulfide and
silver nitrate are added to mixture 1 .
10. Based on the observations write a balanced
chemical and ionic equation to describe the
two mixtures. Interpet all symbols used.
11. Describe what you think is happening, on a
microscopic scale, in these mixtures.
Part B
1 . Write an equation to describe the physical
change that produced the purple vapor when
the test tube was heated.
2. Write an equation to describe the physical
change that produced the needle-like crystals
in the upper region of the test tube when it
was removed from the hot water.
3. What evidence suggests the changes in 1 and
2 were changes of state and not chemical
changes?
4. What are two plausible explanations for the
fact that the amount of purple vapor remains
constant even though the test tube remains in
the hot water and there is still solid iodine at
the bottom of the test tube?
5. What evidence makes it hard to accept the
"stopped" model?
6. Write an equation to describe the system
when the amount of purple vapor remained
constant. Interpret all symbols used.
7. Describe what you think is happening on a
microscopic scale when the pink color in the
test tube remains constant.
Parts A and B
1. What is meant by the term "dynamic" when
applied to the concept of chemical equilibrium?
2. Why is a double arrow I used in a
chemical equation to describe a chemical
system at equilibrium?
3. What evidence suggests that the dynamic
model is the best description of a chemical
system at equilibrium?
Further Reading
Alyea, Hubert N., and F.B. Dutton. Tested
Demonstrations in Chemistry. Easton, Pennsylvania: Chemical Education Publishing Co.,
1965.
Parry, Robert W et al. Chemistry: Experimental
Foundations. Englewood Cliffs, New Jersey:
Prentice-Hall, 1970.
Rowley, Wayne R.E. Matter in Balance: Chemical
Equilibrium. Toronto: Wiley, 1979.
Toon, E.R., and G.P. Ellis. Foundations of
Chemistry. New York: Holt, Rinehart and
Winston, 1973.
Reaction Kinetics
Objectives
Students should be able to:
1 . Describe experimental evidence that supports a molecular distribution
of energies within a sample of gas.
2. Define the term activation energy.
3. Describe the effect of temperature on the molecular distribution of
energies.
4. Use an example to explain what is meant by the term chain
mechanism.
5. Draw energy pathways for endothermic and exothermic reactions.
6. Label the following on an energy pathway: activation energy; total
energy released; and the net energy required or released by a
chemical reaction.
7. Use the energy concepts to explain how a reaction and its reverse
end up in a dynamic equilibrium.
Program Description
Program 3 begins with a review of the meaning of the term dynamic
equilibrium. Then a problem is posed: "Why is it that under one set of
conditions the equilibrium point will involve very few product molecules,
but under other conditions it will involve almost complete conversion of
the reactants to products?" The example used is the chlorine/hydrogen/
hydrogen chloride system. A more detailed application of the kinetic
theory to chemical systems follows. The molecular distribution of kinetic
Activities
Activity 1: A Collision Model for
Chemical Reactions
Set up two stoppered test tubes, each with 2 cm
of solid potassium iodide on top of 2 cm of solid
lead nitrate. If the test tubes are set up several
days in advance, a yellow line will appear at the
point of contact between the two solids.
energy and its relationship to temperature, activation energy, and
reaction mechanisms is introduced and is used to explain how
hydrogen and chlorine can exist in the dark, forming very little hydrogen
chloride; but add a little daylight and almost all the hydrogen and chlorine
is transformed into hydrogen chloride. Energy graphs are used to
illustrate energy changes in exothermic and endothermic reactions and
how one influences the other in a system at equilibrium. The program
concludes by asking a question that opens the door to explore the
effects of other factors on the equilibrium point, such as concentration
and pressure. The question provides the logical link to the next program.
Before Viewing
Review the meaning of steady state and dynamic equilibrium and why
chemists believe most chemical reactions proceed to a state of dynamic
equilibrium. Use Activity 1 to show that a collision must occur before a
reaction can take place. Review the basic assumptions of the collision
theory and the factors that affect the rate of a chemical reaction. Use
Activity 2 to show that many chemical reactions involve a series of steps.
These demonstrations set the stage for the introduction of the collision
and mechanism models presented in Program 3.
After Viewing
Use Activity 3 to consolidate the ideas presented in the program.
Discuss the answers as they will help to form a bridge between
Programs 3 and 4.
Pass the two tubes around the class. Allow
students to shake one but not the other. They
should be able to see that yellow lead iodide only
forms when the two solids come in contact. This
point can be made dramatically by adding water to
the test tube, thus dissolving the solids. When
dissolved the contact area is increased
tremendously and the contents of the test tube
i nstantaneously turn bright yellow.
This is a good reaction to show students
because most believe reactions do not take place
between solids. It also illustrates the basic
assumption of the collision theory that reactants
must come in contact or collide before a reaction
takes place.
Activity 2: A Mechanism Model for
Chemical Reactions
Since most of the students' experience has been
with reactions that appear to go in a single step, it
is worth the time to show them some of the
evidence that suggests that a better way to
describe most chemical reactions is with a series
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of simple reactions or steps called the mechanism
of the reaction.
Demo 1
Add a drop of aqueous iron(III) chloride to a
solution of sodium thiosulfate (concentrations are
not critical) in a petrie dish on the overhead
forms and
projector. A black complex,
slowly disappears, leaving a pale yellow solid,
colloidal sulfur. Be prepared to show that the
observations are not due to simple diffusion. This
is accomplished by adding a drop of iron solution
to a second petrie dish containing just water. This
is an excellent opportunity to discuss plausible
explanations and then design experiments to test
them. Do not spend time presenting or trying to
develop the mechanism for this reaction. Just use
it to help establish that most chemical reactions
can be pictured as a series of simple steps which
add up to the overall reaction. The next
demonstration will help you drive this point home.
Demo 2
To avoid problems when conducting this
demonstration, only make as much solution as you
will need for your classes plus a couple of trials;
make up the solutions the night before you intend
to use them. The concentrations of the three
solutions are as follows:
To produce C make a slurry of the starch and
enough water to cover the starch in a mortar. Use
the pestle to mix the two well. Pour the slurry into
approximately 500 mL of boiling water. Turn off
the heat and stir well. Allow the colloidal starch
dispersion that forms to cool while you make up
solutions A and B. If necessary, add ice to cool
the dispersion to room temperature. Add sufficient
water to bring the total volume up to 1 L. Add the
15 g of sodium hydrogen sulfate(IV) and allow it
to dissolve.
Step
1.
2.
Caution: Do not add the sodium hydrogen
sulfate(IV) to the hot dispersion. If you plan to
keep solution C for more than a day or two add
1 g of salicylic acid and 10 mL of ethanol for
every litre of solution.
Label two Erlenmeyer flasks 1 and 2. Use the
chart below to obtain the volumes of A, B, C, and
water. Use four graduated cylinders, one for each
of the components, to avoid contamination.
When all of B is used up then step 3 can take
place:
Add the appropriate volumes of A, B, and water
to flasks 1 and 2. Have the students predict the
relative rates of the reaction in the two flasks.
They should predict flask 1 to be the faster of the
two because the concentration of B has been
reduced in flask 2.
Review why the total volume was kept
constant. Swirl the flasks to ensure that the
reactants are well mixed and add 20 mL of C
simultaneously to each of the flasks. Continue to
swirl. An orange precipitate will appear in flask 1
first. But just as the students begin to feel that
flush of success that comes from being right, flask
2 will turn jet black, followed shortly by flask 1.
Who won? How is it possible that the orange
precipitate could form first in flask 1 but the black
appear first in flask 2?
Divide the students into groups of four. Tell
them to represent the solutions by the letters A,
B, and C and ask them to develop a model that will
account for what they have seen. The following is
typical of what they will produce:
You could tell students the ingredients but that
would just stifle the creative process as they try to
recall appropriate chemical reactions. Remember
that the purpose of this exercise is not to teach
the mechanism of this reaction but to establish
that it is reasonable to think of a reaction as
i nvolving a series of steps.
Activity 3: Reaction Kinetics
The following questions will help you review the
terms and concepts of the rates of chemical
reactions. Some of them ask you to apply the rate
concepts to a chemical system at equilibrium.
Basic Terms and Concepts
1. Define the following terms:
endothermic; exothermic; activation energy;
activated complex; rate of reaction; and
mechanism of a chemical reaction.
2. Describe a chemical system at equilibrium in
terms of rates of reactions.
3. (a) List the four factors that determine the rate
of a reaction.
( b) Use the collision theory to explain on a
microscopic scale how each of the factors
affects a chemical reaction.
Energy Pathways
1. (a) Draw and label the energy pathway for the
following reactions:
Label the type of reaction (exo- or
endothermic); the position of the reactants
and products; the activated complex; the
activation energy; the total energy
released; and the net energy produced or
used by the reaction.
(b) Use the diagram to explain how a reaction
and its reverse end up in dynamic
equilibrium.
2. The following energy pathway, Figure 3,
describes the energy changes for the reaction:
(a) Complete the following chart for the steps
of the reaction mechanism. Actual values
are not required. Compare one step to the
other using terms such as larger and
slower.
(b) What is the rate-determining step for this
reaction?
(c) Explain why the reaction mechanism is a
better model for what happens on a
microscopic scale during this reaction than
the net equation.
Molecular Distribution of Kinetic Energy
1. (a) Draw and label the curve that shows the
distribution of kinetic energies in a sample
of gaseous molecules at room tempera
ture. (This distribution is called the
Maxwell-Boltzmann Distribution.) Show
what happens to the shape of the curve
when the temperature is increased. (Use a
different color.)
( b) Mark the activation energies of the
reactions in question 1(a) Energy
Pathways on the distribution curve.
(c) Use the diagram in 1(b) to explain why you
would expect more of the products of the
exothermic reaction at equilibrium than
those of the reverse endothermic reaction.
(d) Use the diagram in 1(b) to predict which of
two reactions will be most affected by an
i ncrease in temperature.
Implications for Equilibrium
1. Consider the effects on the rate of formation of
hydrogen iodide if the concentration of
hydrogen were increased. Use these changes
to predict what will happen if some hydrogen is
added to a container holding iodine and
hydrogen in equilibrium with hydrogen iodide.
In your prediction describe the changes in the
amounts of iodine and hydrogen iodide at
equilibrium.
2. Use an argument based on rates of reactions
to predict the effect of a catalyst on the
11
12
amounts of reactants and products at
equilibrium.
3. The formation and decomposition of hydrogen
i odide are second-order reactions.
(a) Write a rate expression for each reaction.
( b) How will the rates of the two reactions
compare when hydrogen iodide is in
equilibrium with hydrogen and iodine?
State your answer mathematically and in
words.
(c) Use the mathematical statement to show
that at equilibrium there exists a ratio of
concentrations of reactants and products
that equals a constant.
( d) What are the only variables that will alter
the value of the constant developed in (c)?
Further Reading
Alyea, Hubert N., and F.B. Dutton. Tested
Demonstrations in Chemistry. Easton, Pennsylvania: Chemical Education Publishing Co.,
1965.
Huff, George E. Molecules in Motion. Toronto:
Wiley, 1976.
Laidler, Keith J. Chemical Kinetics. New York:
McGraw-Hill, 1950.
Mahan, Bruce H. University Chemistry. Reading,
Massachusetts: Addison-Wesley, 1965.
Parry, Robert W. et al. Chemistry: Experimental
Foundations. Englewood Cliffs, New Jersey:
Prentice-Hall, 1970.
Toon, E.R., and G.P. Ellis. Foundations of
Chemistry. New York: Holt, Rinehart and
Winston, 1973.
Reaction Tendencies
Objectives
Before Viewing
Students should be able to:
Use the overhead to show the constant macroscopic properties in the
tubes. Have the students
describe the systems on both a macroscopic and microscopic scale.
Have them write an equation to describe the processes occurring in the
tubes. Review the effects of temperature, a change in volume, and
concentration on the rate of a chemical reaction, and use this data to
predict what will happen if these same variables are applied to a system
at equilibrium. Make the predictions and then demonstrate what
happens. As part of the predictions and demonstrations complete
Activity 1 . Ask the students to deduce a general principle based on what
they have observed, then show Program 4.
1. State Le ChMelier's Principle.
2. Apply Le Chatelier's Principle to predict and explain how the stresses
of heat and pressure affect a given system of equilibrium.
Program Description
Program 4 begins with a review of the concept of dynamic equilibrium.
The effect of a stress on a system at equilibrium is introduced using an
analogy. Credit for applying this common-sense approach to chemical
systems is given to Le Chatelier. The effects of the stresses of heat and
pressure are illustrated and accounted for. Le Chatelier's Principle is
used to predict the shift in the position of equilibrium. The program ends
by establishing a need for knowing the effect quantitatively.
Activities
Activity 1: Pressure Effects
The following can be demonstrated with a 50 mL
translucent plastic syringe on the overhead or it
can be a student activity using 1 mL disposable
syringes held over a white sheet of paper. The
syringes are sealed by heating the plastic at the
needle end in a flame and squeezing it shut with a
pair of tongs.
Write the following equation on the chalkboard:
After Viewing
Review Le Chatelier's Principle. Have the students review what happens
i n the demonstrations you showed them before the program. Redo them
if they cannot recall the observations. Complete Activities 2 and 3.
Activity 3 provides the link with Activity 1 in Program 5.
Ask the students for ways to change the
concentration of the participants. Obviously
keeping the volume constant and adding more of
one will change its concentration. However, it is
also possible to alter the concentration of both by
changing the volume of the container. Ask the
students which reaction, the forward or the
reverse, will be altered the most by a change in
volume. If their first guess is that both will be
changed equally, do not evaluate it but ask for the
rate expression for the formation of N02 and for
N204, assuming that the reactions occur microscopically as written. Write the following on the
chalkboard:
Now ask the question again. It should be more
obvious that a compression that halves the volume
would double the concentration of each species,
but that the rate of formation of N204 would be
four times what it was while the rate of formation
for N02 would only be doubled. Thus there would
be a build up of N204 until the rates of formation
were once again equal. At this new equilibrium
there would be more N204 and less N02 than
before the compression. Reasoning this way, the
students would predict that a compression,
13
14
increasing the pressure, would cause the intensi
of the reddish-brown color to fade.
Once the nrediction is estahlished fill the
To demonstrate the effect,
syringes with
quickly push the piston in halfway and hold it. At
first glance the only apparent change is an
increase in the intensity of the reddish-brown
color, the opposite to what was predicted. Repeat
the procedure. Push the piston in quickly, hold it
at the halfway point, pause until the color remains
constant, then quickly pull it out. Be careful not to
completely remove the piston. Careful observation
will reveal that the intensity increases when you
push the piston in and then fades. Similarly,
quickly pulling the piston out causes the intensity
to fade initially and then darken. Ask the students
for plausible explanations but don't evaluate them.
Program 4 uses an excellent analogy that will help
them account for their observations.
To generate the
use a gas generator or a
l arge test tube fitted with a one-hole rubber
stopper. Insert a glass tube that has been bent
approximately 100 ° and drawn out to a point into
the rubber stopper. This allows you to place the
glass tube into the barrel of the small syringe,
l etting it fill from the bottom up. Add a small
volume of concentrated nitric acid and small
pieces of copper wire to the test tube. Stopper
the test tube and fill the syringes with qas.
Caution: Nitric acid is corrosive; wear protective
gloves. NO2 is toxic; fill the syringes in a
fumehood. Add small pieces of copper so you
only produce as much gas as you want. The
reaction can be stopped by adding water to the
test tube. Be sure to wash your hands when you
are finished.
Because the syringes are very quickly stained
by the gas it is not possible to fill them in advance.
This also means that a disposable syringe can be
used only once or twice. Thus you will need a
large supply of them. If your doctor gives allergy
shots she or he may be willing to save used 1 mL
syringes for you. The large 50 mL syringes can
be purchased through a pharmacy.
Activity 2: Pressure Temperature
Effects
Demo 1
Review Le Chatelier's Principle. Have the students
give you the equation that describes the equilibrium in a water/ice mixture. Ask them to predict
the effect of increasing the pressure on this
system. Have a model of ice available so you can
ill ustrate why water expands when it freezes. This
i s a good opportunity to review some of the ideas
related to the Kinetic Molecular Theory.
After the prediction is established support a
block of ice between two iron rings on stands.
Hold a thin wire with 1 kg masses on each end
over the block. Make sure the masses will not rest
on the desk even after the wire passes through
the ice. To avoid a mess it is also advisable to
have some sort of catch basin for the water that is
produced as the ice melts.
Have the students apply their prediction to this
situation and then let the wire hang on the block.
The wire passes through the block and the
masses clatter to the desk surface, but the block
remains in one piece. Ask the students for
plausible explanations. The melting as a result of
i ncreased pressure is easy; students expect that.
But why does the water refreeze? If you have
i ncluded the energy requirements of this system in
the equation, students should soon realize that the
temperature of the ice drops when it melts. The
energy required for melting has to come from
somewhere. In this case it comes from the ice
itself. Be sure to point out that at this stage the
system is not at equilibrium but is attempting to
shift to a new equilibrium. Thus when the pressure
i s returned to normal, as the wire passes the
water will freeze at the lower temperature,
producing heat so the temperature returns to
0 ° C. It will take several minutes for the wire to
pass through, depending on the thickness of the
block. While you are waiting proceed with
Demonstration 2.
Demo 2
Place about 0.3 g (a few crystals) of crushed
in a test tube and add 5 mL of
ethanol. Shake vigorously until most of the solid
has dissolved. If the solution is not pink, add drops
of water until it just turns pink. Show the students
the color by holding the test tube over the stage
of the overhead projector or by pouring some of
the solution into a petrie dish on the overhead.
Write the following on the chalkboard:
Ask the students for evidence that suggests
the system is at equilibrium.
Ask the students to use Le Chatelier to predict
what will happen if the system is heated, then heat
the solution gently in a cool Bunsen flame until it
changes color.
Caution: Alcohol is flammable. Place a book over
the mouth of the test tube to smother the flames if
the ethanol catches fire.
Ask what will happen if the system is cooled,
then cool the test tube by immersing it in a beaker
of ice water.
Activity 3: Concentration Effects
This activity leads into Activity 1 in Program 5.
Thus it provides a natural bridge between the two
programs.
The demonstration is best done on an overhead
projector using petrie dishes. It can, of course,
be done on the demonstration desk with
test tubes or beakers but larger volumes of
reagents will be required. The solutions required
are the same as those required for Activity 1 in
Program 5.
Five petrie dishes
150 mL beaker
Stirring rod
Add 25 mL of KSCN solution to an equal
volume of water in the beaker. Add three or four
drops of the iron(III) solution into the beaker and
stir. The number of drops can vary. The important
point is that the color is intense enough to be
seen but not so intense that the equilibrium is
pushed essentially to completion. Ask the
students for the ions present in the mixture and
the possible combinations that might produce the
stress involved, and identify the reaction that
could reduce that stress. After the prediction is
established demonstrate what happens. Also point
out that the increase in color indicates that some
iron(III) ions were present in the solution.
Label the petrie dish used, "add SCN' - . " Use
the same line of questioning for the stress "add
Fe3+." After establishing that both reactants
are present at equilibrium ask students what
should be done to get the color to fade. They
should reply that removing one of the reactants
will shift the equilibrium to the left and the color will
fade. Again, review the reasoning behind the
students' answer. Point out that this would require
that the reaction he reversible. Tell them that
forms a complex with
that will
effectively tie it up, so it is removed from
participating in the thiocyanoiron(III) equilibrium.
Label the petrie dish, "remove Fe 3 +," and add a
few crystals of Na2HP04. Now ask if anyone has
noticed any other changes during the demonstration that you haven't discussed. Someone usually
notices that the reference has faded. Ask for plausible explanations. Some, of course, will think you
are imagining things. Show them the fifth petrie
dish to establish the effect of the stress of adding
heat from the projector on the thiocyanoiron(III)
equilibrium. If this is not dramatic enough warm
some of the solution in a Bunsen flame.
Further Reading
Alyea, Hubert N. and F.B. Dutton. Tested
Demonstrations in Chemistry. Easton, Pennsylvania: Chemical Education Publishing Co.,
1965.
Choppin, Gregory R. et al. Chemistry. Morristown,
New Jersey: Silver Burdett, 1978.
Othen, Clifford. Rates of Reaction and Equilibria.
London: Heinemann Educational Books,
1968.
Parry, Robert W. et al. Chemistry: Experimental
Foundations. Englewood Cliffs, New Jersey:
Prentice-Hall, 1970.
Parry, Robert W. et al. Chemistry: Experimental
Foundations. Teacher's Guide. Englewood
Cliffs, New Jersey: Prentice-Hall, 1975.
Rowley, Wayne R. E. Matter in Balance: Chemical
Equilibrium. Toronto: Wiley, 1979.
Toon, E.R., and G.P. Ellis. Foundations in
Chemistry. New York: Holt, Rinehart and
Winston, 1973.
15
16
The Equilibrium Constant
Objectives
Students should be able to:
1. Apply Le Chatelier's Principle to predict and explain how the stress of
concentration affects a given system of equilibrium.
2. State experimental evidence for Le Chatelier's Principle.
3. Use the "Dance Hall" analogy or the Kinetic Molecular Theory to
account for Le Chatelier's predictions.
4. Write the mass action expression for the equilibrium constant for a
given chemical system.
5. Identify evidence that supports the Law of Chemical Equilibrium.
6. State the implications of the size of the equilibrium constant.
Program Description
Program 5 reviews the application of Le Chatelier's Principle to chemical
systems by illustrating and accounting for the changes produced by
altering the concentration of the reactants or products. Le Chatelier's
predictions are then shown to be consistent with the dynamic model of
equilibrium. The program reestablishes the need for having a more
quantitative expression for the changes and then uses empirical data to
i ntroduce the equilibrium constant. The result is generalized as the Law
of Chemical Equilibrium. The law is then applied to several systems and
the meaning of the size of the equilibrium constant is introduced.
Before Viewing
Do Activities 1 and 2. Use Activity 3 in Program 4 as part of the prelab to
Activity 1. See the Teacher's Guide to Chemistry: Experimental
Foundations (see page 24) for suggestions to help students cope with
the serial dilution required in this lab. Use Program 5 as part of the postl ab for Activities 1 and 2. Before the post-lab have students plot their
depth ratio values versus the vial number on a set of axes drawn on a
ditto. (See the following graph of typical results.) Run the ditto off so
everyone has a set of the data. The points will vary over a fairly large
vertical range. Don't be concerned; there is bound to be considerable
variation due to the way the dilutions were carried out and the method
used to compare the colors. Use the general shape of the curve to draw
a smooth curve of "best fit." Read off the graph the "best" values for
each vial number and have all students use these ratios in the remainder
of the calculations.
Figure 4
0.9
0.8
0.7
0.6
0.5
0.4
0.3
Vial Number
Activity 1 enables students to discover that the mass action
expression for a system at equilibrium is independent of the concentrations of the reactants and products. It also introduces the idea of an
equilibrium constant using data gathered by the students. Activity 2 does
the same thing, using data gathered by others.
After Viewing
Do Activities 3 and 4. Activity 3 provides students with additional applications of Le Chatelier's Principle and another method for determining
the equilibrium constant. Activity 4 provides practice interpreting
equilibrium constants and using them to make predictions.
Activities
Activity 1: Chemical Equilibrium A Quantitative Description, Part 1
I n this experiment you will take a quantitative look
at the reaction
which was used qualitatively in an earlier
demonstration. This time you will determine the
concentration of each of the ions at equilibrium,
and then seek an expression that relates the
concentrations mathematically in a simple,
convenient manner. Such an expression would
enable the quantification of Le Chatelier's predictions, a necessity if you wish to be able to predict
profits and losses for an industrial process.
The concentrations are determined colonmetrically. Color intensity of a solution depends
upon the concentration and depth of the solution.
If you add water to a cup of tea the color intensity
remains constant. This is because the depth of the
tea increases as the concentration decreases.
The relationship is expressed as follows:
the depths of solution required to make the color
intensities equal.
To prepare the standard solution in this activity
you will use a small amount of thiocyanate ion,
SCN' - , and add a large excess of ferric ion,
Fe 3+ (aq). You can assume that all of the thiocyanate ion will be used in forming the complex
thiocyanoiron(III) ion, FeSCN 2 +(aq). Thus the
concentration of FeSCN 2 +(aq) in the standard will
be equal to the starting concentration of the
SCN' - ( aq). The validity of this assumption will be
discussed in the post-lab discussion.
Apparatus
Distilled water
Five flat bottom vials (must hold 10 mL)
Two 10 mL Mohr pipets
Pipet bulb
25 mL graduated cylinder
color intensity = kCd
150 mL beaker
C = concentration expressed in mol/L
Diffused light source or white paper
d = depth in mm
Ruler
k = proportionality constant
Medicine dropper
Thus if two solutions appear to have the same
color intensity
Method
Thus if you know the concentration of a solution in one situation (let's call it the standard
solution), you can calculate the concentration of
the solution in a second situation by comparing
1. Line up five clean, dry vials. Label them 1, 2,
etc.
2. Use a clean, dry Mohr pipet and pipet bulb tc
transfer 5.0 mL of 0.002 mol/L KSCN i nto
each of the vials.
3.
4. Add 10.0 mL of 0.2 mol/L Fe(N03)3 to the
25 mL graduated cylinder, then fill the
cylinder to the 25.0 mL mark with distilled
water. Pour the solution into a dry, clean 150
mL beaker to mix it. Calculate the concentration of this solution.
5. Rinse the second Mohr pipet with water and
finally with a little of the solution in the
beaker.
Caution: You only have 10 mL excess; only
use 2-3 mL to rinse the pipet. Do not return
the rinse solution to the beaker; discard it.
6. Use the second Mohr pipet and bulb to
transfer 5.0 mL of the solution in the beaker
to vial 2. Calculate the initial concentrations
of the Fe 3 +(aq) and SCN' - i n the 10 mL of
solution formed in the vial.
Note: One mole of
yields one mole
of
7. Rinse the graduated cylinder with water and
finally with a little (2-3 mL) solution from the
beaker.
8. Pour 10.0 mL of solution from the beaker
i nto the graduated cylinder then fill the
cylinder to the 25.0 mL mark with distilled
water.
9. Discard any solution remaining in the beaker
and rinse it with water. Dry the beaker with a
paper towel.
1 0. Pour the solution in the graduated cylinder
i nto the beaker to mix it.
1 1. Repeat steps 5 and 6, only this time transfer
the solution to vial 3, then repeat steps 7-10.
1 2. This process is called serial dilution. Continue
i t until you have 5.0 mL of successively more
diluted
solution in each vial.
Calculate the initial concentrations of
for each vial as part
of your preparation for the activity.
13. Wrap a strip of white paper around vials 1 and 2.
1 4. Look vertically down through the solutions at
a diffused light source. Use the medicine
dropper to remove a dropper full of the standard.
17
18
15. Return some of the standard in drops until the
colors match. Hold the vials close together
and blink your eyes between "looks" to help
avoid eye fatigue. Put the unused standard in
the dropper into a clean dry beaker, since
you may have to use some of this solution
later.
16. When the intensity of the color in the vials
matches, record the height of solution in
each tube to the nearest millimetre.
17. Repeat the matching procedure with vials 1
and 3, 1 and 4, and finally 1 and 5.
2. Calculate the following for vials 2 through 5:
(a) Calculate the initial concentrations of
Initial means
before any chemical reaction has taken
place. Subscript "I" is used to denote initial
concentrations.
(b) Calculate the radio of depths from the color
comparison. Subscript "x" is used to
denote the vial number.
( d) Calculate the equilibrium concentration of
Fe3+(aq) and SCN1 -(aq).
[SCN'"] E = [FeSCN2+]E
(e) Now try to find some simple mathematical
relationship between the equilibrium
concentrations that could be used to make
quantitative predictions. Try calculating the
following:
Activity 2: Chemical Equilibrium A Quantitative Description, Part 2
This activity provides you with data to discover a
general relationship for the concentrations of the
reactants and products at equilibrium. Perform the
Observations
Discussion
1. Use Le Chatelier's Principle to explain why the
assumption made to calculate
reasonable.
2. Which of the combinations, (i), (ii), or (iii), gives
the most constant value? To help you decide,
calculate the ratio of the largest value to the
smallest for each expression.
3. Restate the most constant expression in words
using the terms reactants and products. A
general statement of the regularity you have
noted is called the Law of Chemical
Equilibrium.
4. Assume that the reactions for this system
occur on a microscopic scale as written. Write
rate expressions for the forward and reverse
reactions. Show that the idea of a dynamic
equilibrium is consistent with the expression
found above.
necessary calculations to complete the data tables
for each of the following systems.
System I
System I involves the following equilibrium
between ethyl acetate, water, acetic acid, and
ethanol:
The data in the table was gathered by starting
with known amounts of ethyl acetate and water.
The two reactants were placed in an Erlenmeyer
flask and swirled. The concentration of acetic acid
at equilibrium was determined with the aid of a pH
meter. Since the production of one mole of acetic
acid also involves the production of one mole of
ethanol (ethyl alcohol), then the concentration of
acetic acid and ethanol at equilibrium would have
to be the same. All trials were done at room
temperature.
Note: The subscripts "I" and "E" used in the
table refer to initial and equilibrium concentrations
respectively.
For column "i" calculate:
The equilibrium concentration of N02 was
measured colorimetrically, similar to the method
used in Activity 1. The only difference is that a
spectrophotometer was used to measure the
color intensity.
Discussion
1. Show a sample calculation for
2. What stress caused the change in
from trial 1 to trial 8? Use Le Chatelier's
Principle to predict the effect of the stress.
Does the data in the table support your
prediction?
3. I n trials 4 and 5 1.00 mol of each of the
gases was put in the vessel (at least that is
the way it appears on the data table).
However, this was not the way the trial was
done. After equilibrium was reached in trial 4
the volume of the container was adjusted to 3 L.
For column "ii" calculate:
1. Show a sample calculation for
2. The two expressions of the concentrations of
the products and reactants that you calculated
are called mass action expressions. Which
mass action expression for equilibrium concentrations is not a function of concentration (i.e.,
it doesn't change when the concentrations
change)?
3. Which of the three expressions is the most
useful? Explain your choice.
4. Use Le Chatelier's Principle to predict the
effect when water is added to the ethyl
acetate system at equilibrium. Which trials
support your prediction?
Note: The subscripts "I" and "E" used in the
table refer to initial and equilibrium concentrations
respectively.
For column 'T' calculate:
For column "ii" calculate:
System If
System II involves the following equilibrium:
For column "iii" calculate:
( a) What would a change in the volume do to
the pressure exerted on the gases?
( b) Use Le Chatelier's Principle to predict the
effect of this stress on the
(c) Does the data in the table support your
prediction?
4. Which of the three mass action expressions
of concentrations at equilibrium is not a
function of concentration?
5. Which of the three expressions is the most
useful? Explain your choice.
6. What would be the mass action expression
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20
for the following system that would not be a
function of concentration?
Note: Follow the convention of putting the
concentration of the products over the
reactants:
7. The value of the mass action expression that
i s independent of concentration is called the
equilibrium constant. What is the equilibrium
constant for the ethyl acetate equilibrium?
What is it for the nitrogen dioxide equilibrium
at 142 42 0 C? At 11 5 0 C? At 70OC?
8. What two variables will alter the value of the
equilibrium constant? Use the data presented
Note: The subscripts "I" and "E" used in the
table refer to initial and equilibrium concentrations
respectively.
For column "i" calculate:
i n Systems I and II to answer this question.
reaction was written as an
9. The
endothermic reaction. In general, what effect
does an increase in temperature have on the
equilibrium constant for an endothermic
reaction? Write a generalized statement to
describe the effect of temperature on the
equilibrium constant. Is your statement
consistent with Le Chatelier's Principle?
Explain your reasoning.
1 0. Predict the mass action expression that will
equal the equilibrium constant for the
following equilibrium:
The equilibrium concentration of 12 was
measured colorimetrically, similar to the method
used in Activity 1. The only difference was
that a spectrophotometer was used to measure
the color intensity. All trials were carried out at
698 K.
Antimony(III) chloride reacts with water to form
white insoluble antimony(III) oxychloride. This
hydrolysis reaction can be represented by the
following equilibrium:
Discussion
1. Show a sample calculation for
2. (a) Which mass action expression was equal
to the equilibrium constant?
(b) Did you predict correctly in question 9 of
System II? If not, where did you go wrong?
3. The symbol for the equilibrium constant is
capital "K." Write the mass action expression
for the equilibrium constant for the following
equilibrium:
System III
System III involves the following equilibrium:
Activity 3: Chemical Equilibrium
Note: The convention is to put the concentration of the products over the reactants.
4. The mathematical expression you wrote in the
above question is a statement of the Law of
Chemical Equilibrium. State the law in words.
5. The equilibrium constant is independent of
concentration. What does this mean?
6. What two variables cause K, the equilibrium
constant, to vary?
The effects of adding water and hydrochloric
acid to the system will be investigated. Also K, the
equilibrium constant for this reaction, will be
determined.
One litre of the solution used contains 0.5 mol
and 6 mol HCI. That is, the solution concentrations are 0.5 mol/L with respect tc
and
6 mol/L with respect to HCI. In this solution you
can assume the concentration of SbOCI is almost
zero.
Do discussion questions 1-4 on this page as
part of your preparation for this activity.
Apparatus
10 mL graduated cylinder
50 mL buret
Two 125 mL Erlenmeyer flasks
20 mL 0.5 mol/L SbC13 in 6 mol/L HCI
25 mL 6 mol/L HCI
Caution: The soluble compounds of antimony are
almost as poisonous as the compounds of
arsenic. Wear protective gloves and wash your
hands thoroughly after this activity.
Method
1 . Transfer 5.0 mL of antimony(Ill) chloride
solution into an Erlenmeyer flask from a buret
at one of the dispensing stations your
teacher has set up.
2. Add 5 mL water from a graduated cylinder
and swirl the mixture. Record any changes.
3. Add an additional 10 mL of water in portions
of 2 mL at a time until a total of 15 mL have
been added. The 15 mL includes the 5 mL
added in step 2. Swirl the mixture and record
any changes with each addition.
4. Use the graduated cylinder to add 2 mL of
6 mol/L HCI. Swirl and record any changes.
5. Add an additional 6 mL in 2 mL portions,
swirling the contents and recording any
changes each time.
6. Dispose of the mixture as your teacher
directs.
7. Add 5.0 mL antimony(III) chloride solution
from the buret at the dispensing station to a
clean, dry 125 mL Erlenmeyer flask.
8. From another buret add water to the flask
until the white precipitate begins to form.
Swirl the flask continuously as you add the
water slowly. Try not to add too much. You
can use your results from steps 2 and 3 as a
guide. Record the volume of water required
to the nearest 0.1 mL.
9. Dispose of the mixture as your teacher
directs.
10. Calculate the concentration of HCI in the total
volume of solution when the precipitate was
just starting to form.
11. Calculate the concentration of antimony(III)
chloride present in the total volume,
assuming that a negligible amount of SbCI3
was used in the formation of the white
precipitate.
1 2. Determine the value of K', the modified
equilibrium constant (see discussion question
4), from the concentrations of HCI and
for the solution in which the SbOCI was just
starting to form.
7. Based on your observations, is the assumption
made in step 12 reasonable? If it isn't, how will
it affect your value of K'?
8. (a) What is the value for K'?
( b) What does the value for K' imply about the
hydrolysis of antimony(Ill) chloride?
(c) Pedict the relationship between your value
for K' and the accepted value. Explain
your reasoning.
Discussion
1. Explain why it is reasonable to assume that
[SbOCI] = 0 in the solution that contains
2. Predict the effect of adding water on the
antimony(III) chloride hydrolysis equilibrium.
Describe any changes you would expect to
see.
3. Predict the effect of adding hydrochloric acid
on the antimony(III) chloride hydrolysis
equilibrium. Describe any changes you would
expect to see.
4. (a) Write the mass action expression for the
equilibrium constant for the antimony(III)
chloride.
( b) Since the concentration of a pure solid or
li quid is directly proportional to its density it
will only change with temperature, Thus the
concentration of a pure solid or liquid in
equilibrium with other substances remains
constant. Since the concentration of the
water and the antimony(III) oxychloride do
not change they can be grouped with the
equilibrium constant to form a new constant
called K'. Write the mass action
expression for K'.
(c) What would K' be equal to in terms of K?
5. (a) At what point in the addition of water (steps
2 and 3) was the system at equilibrium?
(b) Why did the system not appear to change
after several additions of water?
6. Account for any change or lack of change that
occurred when the acid was added.
Activity 4: Equilibrium Applications
The following questions will help you learn to apply
the concept of equilibrium to chemical systems.
1. Use Le Chatelier's Principle to predict the
effects of heat, pressure, and concentration
on the following chemical systems:
(a) The production of methanol (methyl
alcohol). Stresses: add heat, increase
nressure by compressing the container.
removal of
methanol.
( b) The production of chlorine. Stresses:
remove heat, decrease pressure by
expanding the container, add oxygen, add
a catalyst, add a noble gas that doesn't
react with any participant in the
equilibrium.
(c) The formation of
Stresses: add heat. increase the pressure
of
add solid
21
what will be the concentration of
hydrogen iodide at equilibrium?
(b) At a certain temperature hydrogen iodide
i s 20 percent dissociated. Determine the
concentration of each component present
i n the mixture at equilibrium, if 0.50 mol of
HI are placed in a 1 .00 L vessel at this
temperature.
(c) Determine the equilibrium constant for the
formation of hydrogen iodide at this
certain temperature.
(d) Is the certain temperature lower or higher
than 699 K? Explain your reasoning.
( e) In a mixture of hydrogen iodide,
hydrogen, and iodine at 699 K, the partial
pressures are 70.9 kPa, 2.03 kPa, and
2.03 kPa respectively. Will there be any
change in the partial pressure if the
mixture is maintained at 699 K? If so, will
HI be consumed or formed?
(f) 1.00 mol of H2 and 2.00 mol of 12 gas are
put into a 10 L container at 699 K.
Determine equilibrium concentrations of
the three gases at equilibrium.
(g) What would happen to the equilibrium
concentrations in (f) if the volume were
compressed to 5.0 L?
is 243.1 kPa when the partial pressures are
expressed in kilopascals.
(a) 6.7 g of S02CI2 are placed in a 1.0 L bulb
at 375 K. Determine the pressure
exerted by the
assuming none of
it dissociated.
(b) Determine the partial pressures of
and
at equilibrium.
, c) Determine the partial pressures of
and
at equilibrium if
6.7 g of S02CI2 and 101.3 kPa of Cl2 are
placed in a 1.0 L bulb at 375 K.
( d) Compare parts (b) and (c). In order to get
the situation in (c), what stress must be
applied to the system in (b)? Use Le
Chatelier's Principle to predict the effect.
Compare your predictions with the
calculated concentrations. Are they
consistent?
9. Ammonium hydrosulfide decomposes in the
following manner:
( a) Determine the partial pressures of
ammonia and hydrogen sulfide at equilibrium when excess solid ammonium hydrosulfide decomposes in an excavated
chamber at 25 ° C.
( b) Use Le Chatelier's Principle to predict the
effect of injecting ammonia into the
system at equilibrium.
( c) Check your prediction by completing the
following problem:
Excess solid ammonium hydrosulfide is
placed in a flask with 50.7 kPa of
ammonia. What will be the partial
pressures of ammonia and hydrogen
sulfide at equilibrium?
1 0. The equilibrium constant for the reaction of
"A" and "B" to produce "C" is 4.0.
( a) Determine the equilibrium concentration
of each species if the starting conditions
are 0.50 mol of "A" and 1.0 mol of "B" in
a 10.0 L container.
( b) If it took 10 minutes to reach equilibrium,
accurately sketch a concentration time
graph for the attainment of equilibrium.
(c) On separate sketches of the graph, show
the effects of the following on the original
system at equilibrium:
(i) adding a catalyst
(ii) adding more "A"
(iii) raising the temperature of the system
(iv) reducing the volume of the container
to 5.0 L
(v) adding the noble gas neon.
Further Reading
Alyea, Hubert N., and F.B. Dutton. Tested
Demonstrations in Chemistry. Easton, Pennsylvania: Chemical Education Publishing Co.,
1965.
CHEM Study. Chemistry: An Experimental
Science. Teacher's Guide. San Francisco:
W.H. Freeman, 1963.
Choppin, Gregory R. et a/. Chemistry. Morristown.Silver Burdett, 1978.
NSCM Project Practical Activities Cl 0. Principles
of Chemical Equilibrium. Milton, Australia:
Jacaranda Press, 1973.
Othen, Clifford. Rates of Reaction and Equilibria.
London: Heinemann Educational Books,
1968.
Parry, Robert W. et a/. Chemistry: Experimental
Foundations. Englewood Cliffs, New Jersey:
Prentice-Hall, 1970.
Parry, Robert W. et al. Chemistry: Experimental
Foundations. Teacher's Guide. Englewood
Cliffs, New Jersey: Prentice-Hall, 1975.
Rowley, Wayne R. E. Matter in Balance: Chemical
Equilibrium. Toronto: Wiley, 1979.
Toon, E.R., and G.P. Ellis. Foundations of
Chemistry. New York: Holt, Rinehart and
Winston, 1973.
23
24
The Haber Process
Objectives
Before Viewing
Students should be able to:
Have students complete Activity 1. They are asked to apply Le
Chatelier's Principle in selecting conditions for an industrial process.
However, because an industrial process continually removes the wanted
products, the reaction never reaches equilibrium. Thus in most cases
their predictions will be quite different from the actual conditions used.
Because an industrial process is better described as a steady state than
an equilibrium system, an application of factors that affect the rate of a
reaction are also important. As the program points out, both concepts
must be applied.
Take up the students' answers to the questions but don't evaluate
their thoughts on question 6. Program 6 will help them with this one.
1. Describe the Haber Process for the production of ammonia.
2. Discuss the societal implications of the Haber Process.
3. Apply the concepts of rate and equilibrium to the development of the
conditions for an industrial process.
Program Description
Program 6 begins by establishing the historical events that created the
need for the Haber Process. It then goes through a development of the
process using the concepts from the previous programs. The development requires not only an application of the concepts of chemical
equilibrium but also of those of reaction kinetics. The program ends with
a brief description of the "rewards" reaped by Haber for his mastery of
chemical equilibrium.
After Viewing
Do Activity 2 and consider using Activity 3 as one of the projects that
students may choose to work on this term.
Activities
Activity 1: Le Chatelier and
the Haber Process
Experimental studies have shown that the
percentage of ammonia formed from hydrogen
and nitrogen at equilibrium varies as a function of
the pressure exerted on the system as a whole.
The results for a series of different temperatures
are shown in Figure 5.
Ammonia
Produced
( %)
Discussion
0
Figure
5
20
40
60
Pressure (mPa)
80
100
1 . Write an equation to describe the formation of
ammonia. The name given to the industrial
process using this chemical reaction is the
Haber Process.
2. Is the formation of ammonia an endo- or
exothermic reaction? Use data from the graph
to support your decision.
3. Use Le Chatelier's Principle to predict the
effects of the following stresses on the
ammonia equilibrium:
(a) removing ammonia
( b) increasing the applied pressure
(c) raising the temperature
(d) adding a catalyst
( e) adding an inert substance.
4. State evidence from the graph to support your
predictions in 3(b) and (c).
5. List in general terms the conditions that would
yield the most ammonia.
6. (a) At 20 mPa, what is the yield of ammonia
at:
(i) 473 K(200 ° C)
(ii) 873 K(600 ° C)
( b) At which temperature would you choose to
run the Haber Process? Why?
7. The conditions universally employed by
industry are 873 K(600 ° C), 20 mPa, and the
use of a catalyst. Suggest plausible reasons
for this.
Activity 2: Le Chatelier and
Other Industrial Processes
The Production of Polystyrene
As part of the reaction sequence in which
polystyrene is produced from benzene, ethyl
benzene is dehydrogenated to produce styrene,
the monomer of polystyrene. The following
equation describes the reaction:
1 . Use Le Chatelier's Principle to predict the
effects of the following on the styrene
equilibrium:
(a) increasing the applied pressure
(b) removing styrene
(c) reducing the temperature
( d) adding a catalyst
(e) adding an inert substance.
2. Based on your predictions in question 1,
select the conditions that would maximize the
yield of styrene.
3. (a) Why are the conditions selected in
question 2 unlikely to be those used by
i ndustry?
(b) What additional considerations must be
made when selecting conditions to
maximize the yield of an industrial process?
4. In practice, the pressure of the ethyl benzene
is kept low, as Le Chatelier would suggest, but
an inert substance - super-heated steam - is
added to keep the total pressure of the mixture
at atmospheric pressure. What advantage
does this method have over running the
reaction without the steam at a pressure less
than atmospheric?
5. The super-heated steam serves two other
functions: First it supplies energy to the
system. Why is this desirable? Second, it
reacts with any carbon formed as a byproduct
at high temperatures, preventing the carbon
from contaminating the catalyst. What products
will be formed as a result of the reaction
between carbon and steam?
6. What is the advantage of using super-heated
steam at 873 K(600°C) rather than ordinary
steam?
7. In practice, an iron oxide catalyst is used.
Explain why this is desirable.
8. The process is run at 873 K(600 ° C) with an
i ron oxide catalyst. The conversion is only
about 35 percent complete. The yield could
be increased by raising the temperature.
Suggest two plausible reasons why this is not
done.
9. The ethyl benzene (boiling point 136 ° C) that
has not reacted and the styrene (boiling point
146 ° C) are separated by fractional distillation.
The styrene, of course, is used to make polystyrene. If you were the industrialist, what
would you do with the ethyl benzene?
The Production of Sulfuric Acid The Contact Process
1 . Write a set of equations to describe the major
reactions taking place in the contact process.
2. The system sulfur dioxide, oxygen, and sulfur
trioxide is in equilibrium at 101.3 kPa and
400 ° C in the presence of a catalyst. State
whether the amount of sulfur trioxide would be
i ncreased, decreased, or unchanged by each
of the following. Give reasons for your
answers.
(a) decreasing the applied pressure at 400 °C
(b) adding 5 mol of oxygen at 400 ° C
(c) decreasing the concentration of SO2 at
400°C
(d) adding 2 mol of helium, V and T constant
(e) removing the catalyst.
3. Use Le Chatelier's Principle to suggest three
ways to increase the yield of sulfur trioxide.
4. State the conditions you would use as an
i ndustrialist. Explain the reason for choosing
these conditions if they differ from those you
suggested in question 3.
5. List the actual industrial conditions used (e.g.,
temperature; pressure; catalyst; for removing
sulfur trioxide). Explain the reason for choosing
these conditions if they differ from those you
suggested in question 4.
Activity 3: Industrial
Process Project
The task is to present, in an attractive fashion, the
chemistry involved in an industrial process.
25
26
Pictures, drawings, color, letter size, etc., can be
used to present the information in as attractive a
package as possible. In effect you are going to
produce an ad for the industrial process you
select.
Description of the Task
1. The industrial process chosen must involve
some chemical reaction2. The maximum area covered by your ad should
be 1500 cm2.
3. It must be possible to display your ad on the
wall.
4. The ad must contain the following information:
(a) name of industry
(b) chemical reactions involved
(c) conditions under which reactions are run
(d) location of plants
(e) source of raw materials.
5. The following may be included if desired:
(a) pictures of plants
(b) drawings or flow charts of the process
(c) labels used on final products
(d) anything else you feel makes your ad more
attractive.
Example Industrial Processes
Several industrial processes are listed. They are
i ncluded to give you an idea of the range of
processes from which you can select. The list is
not meant to be restrictive; the possibilities are
endless.
1. Production of iron
2. Production of aspirin
3. Film developing and printing
4. Electroplating
5. Brewing
Further Reading
Ashman, A., and G. Cremonesi. Sulfuric Acid.
London: LongmanslPenguin Books, 1968.
Bradford, Derek. Chemistry and the World Food
Problem. London: Heinemann Educational
Books, 1971.
Haber, L.F. The Nitrogen Problem. London:
LongmanslPenguin Books, 1966.
Parry, Robert W. et al. Chemistry: Experimental
Foundations. Englewood Cliffs, New Jersey:
Prentice-Hall, 1970.
Parry, Robert W. et al. Chemistry: Experimental
Foundations. Teacher's Guide. Englewood
Cliffs, New Jersey: Prentice-Hall, 1975.
Rowley, Wayne R. E. Matter in Balance: Chemical
Equilibrium. Toronto: Wiley, 1979.
Toon, E.R., and G.P. Ellis. Foundations of
Chemistry. New York: Holt, Rinehart and
Winston, 1973.
Ordering
information
To order the videotapes or this publication, or for
additional information, please contact one of the
following:
Ontario
TVOntario Sales and Licensing
Box 200, Station 0
Toronto, Ontario M4T 2T1
(416) 484-2613
Untied States
TVOntario
U.S. Sales Office
901 Kildaire Farm Road
Building A
Cary, North Carolina
27511
Phone: 800-331-9566
Fax: 919-380-0961
E -mail: [email protected] g
Videotapes
Program 1: Steady Unsteadiness
Program 2: Dynamic Equilibrium
Program 3: Reaction Kinetics
Program 4: Reaction Tendencies
Program 5: The Equilibrium Constant
Program 6: The Haber Process
BPN
240701
240702
240703
240704
240705
240706