0.50 x
0.020
2x
(0.50 - x) = 0.28 x
x = 0.39
0.14
Concentration
Initial
Change
Equilibrium
[HBr(g)]
(mol/L)
0.00
0.78
0.78
[H2(g)]
(mol/L)
0.50
-0.39
0.11
[Br2(g)]
(mol/L)
0.50
-0.39
0.11
At equilibrium, [H2(g)] = [Br2(g)] = 0.11 mol/L, and [HBr(g)] = 0.78 mol/L.
0.11 mol
(b) nH2 = nBr2 = 0.500 L u
0.055 mol
1 L
0.78 mol
nHBr = 0.500 L u
= 0.39 mol
1 L
11. Ɣ Atomic theory is necessary because equation reactions involve the collision of reactants
and products, which are made up of atoms, or ions, or combinations of atoms in
molecules.
Ɣ The kinetic molecular theory states that all matter is composed of tiny particles that are in
constant motion. Further, the extent of this motion is dependent upon the amount of
energy possessed by the particles. Molecules with greater energy will move faster. This
movement is essential for chemical reactions to occur.
Ɣ According to collisionreaction theory, collisions that take place between reactant
molecules must occur with enough energy and in the proper orientation for products to
form. Only some percentage of reactant entity collisionscalled “effective”
collisionswill result in a change in the entities.
Ɣ The position of equilibrium is dependent on reaction rates for the forward and reverse
reactions. Equilibrium is a state in which these rates have become equal.
Extension
12. It is universally accepted that the concentration of carbon dioxide in the atmosphere is
increasing, and it is widely believed that this supports the process of global warming (by
increasing the “greenhouse effect”). Recent research suggests that ocean absorption of CO2 is
having significant acidifying effects. This could be a grave problem for marine life. Some
scientists are considering ways to remove carbon dioxide from the atmosphere, to slow the
greenhouse effect. One suggestion is to “inject"” carbon dioxide deep into ocean waters—
which would exacerbate the acidification of the oceans. Other significant effects on CO2
concentration include production by volcanoes and sequestration by marine organisms that
build calcium carbonate shells.
15.2 QUALITATIVE CHANGE IN EQUILIBRIUM SYSTEMS
Investigation 15.2: Equilibrium Shifts (Demonstration)
(Pages 691, 700–701)
Purpose
The purpose of this demonstration is to test Le Châtelier’s principle by studying two chemical
equilibrium systems: the equilibrium between two oxides of nitrogen, and the equilibrium of
carbon dioxide gas and carbonic acid.
606
Unit 8 Solutions Manual
Copyright © 2007 Thomson Nelson
Problem
How does a change in temperature affect the nitrogen dioxidedinitrogen tetraoxide equilibrium
system? How does a change in pressure affect the carbon dioxide-carbonic acid equilibrium system?
Prediction
According to Le Châtelier’s principle, an increase in temperature of the nitrogen oxide
equilibrium system will shift the equilibrium to produce more nitrogen dioxide, while a decrease
in temperature will shift the equilibrium to produce more dinitrogen tetraoxide. Reducing the
pressure on the carbon dioxide equilibrium system will shift the equilibrium toward the carbon
dioxide side of the equilibrium equation.
Evidence
Ɣ When the temperature of the nitrogen oxide equilibrium was increased, the colour of the
mixture darkened. When the temperature was decreased, the colour of the chemical system
lightened.
Ɣ When the pressure on the carbon dioxide equilibrium system was decreased, the
bromothymol blue indicator in the system turned blue. When the pressure was increased, the
indicator turned yellow.
Analysis
On the basis of the evidence gathered in this experiment, an increase in temperature of the
nitrogen oxide equilibrium system (as represented by the equilibrium equation) causes the
equilibrium to shift to the right, to increase the concentration of nitrogen dioxide. This answer is
based on the interpretation that a darkening of the brown colour means an increase in the
concentration of nitrogen dioxide present. A decrease in pressure of the carbon dioxide
equilibrium system (as represented by the equilibrium equation) shifts the equilibrium to the left
to increase the concentration (partial pressure) of the carbon dioxide. This answer is based on the
interpretation that a change in the colour of bromothymol blue from yellow to blue means an
increase in the pH and thus a decrease in the hydrogen ion concentration.
Evaluation
The Prediction is judged to be verified, because the evidence agrees with the predicted results of the test.
Le Châtelier’s principle is judged to be acceptable as an authority, because the prediction was verified.
The Purpose was accomplished. This investigation was a good test of Le Châtelier’s
principle, but it was quite limited in scope.
Career Connection: Chemical Process Engineer
(Page 693)
Chemical process engineers specialize in the products and processes of a particular industry such
as pulp and paper manufacturing, pharmaceuticals, petroleum refining, energy processing,
plastics, metal extraction and refining, or adhesives and coatings production. They may also
specialize in areas within various industries such as process control, pollution control, or
fermentation processes.
Education requirements:
Ɣ A bachelor's degree in chemical engineering or in a related engineering discipline is required.
Ɣ A master's degree or doctorate in a related engineering discipline may be required.
Ɣ Licensing by a provincial or territorial association of professional engineers is required to
approve engineering drawings and reports and to practice as a Professional Engineer (P.Eng.).
Ɣ Engineers are eligible for registration following graduation from an accredited educational
program, and after three or four years of supervised work experience in engineering and
passing a professional practice examination.
Ɣ Supervisory and senior positions in this unit group require experience.
Copyright © 2007 Thomson Nelson
Unit 8 Solutions Manual
607
Job prospects and work assignments:
Ɣ Work prospects in the short term are fair, largely because of current environmental concerns.
Ɣ This is a mobile profession, with skills being transferable across provincial and even national
boundaries. There is considerable mobility between chemical engineering specializations at
the less senior levels.
Ɣ Employment growth may, in the longer term, be below average.
Ɣ Engineers often work in a multidisciplinary environment and acquire knowledge and skills
through work experience that may allow them to practise in associated areas of science,
engineering, sales, marketing, or management.
Ɣ Chemical engineers work closely with chemists and other scientists and engineers, and
mobility is possible between some fields of specialization.
Practice
(Page 695)
1. Changes in concentration (in liquid solution), temperature, and pressure (in gaseous systems)
can cause a shift of equilibrium position.
2. (a) The equilibrium shifts to the right, producing more H2O(g).
(b) Adding solid NaOH causes an increase in the hydroxide ion concentration as it dissolves.
The equilibrium shifts to the left, reducing [H+(aq)] and [OH–(aq)], although the final
[OH–(aq)] is higher than the original equilibrium concentration (before NaOH(s) was
added).
(c)
608
Unit 8 Solutions Manual
Copyright © 2007 Thomson Nelson
(d) The equilibrium shifts to the right, increasing [H+(aq)] and [CH3COO–(aq)] and partially
reducing the introduced increase in [CH3COOH(aq)].
3. (a) There are three moles of gaseous reactant molecules for each one mole of gaseous
product molecules in this reaction. Increasing the total pressure would cause the
equilibrium to shift right, toward the formation of the desired product.
(b) Since the reaction is exothermic, a high temperature would shift the reaction equilibrium
left.
(c) Although increasing the temperature will shift the equilibrium position to the left, it will
also speed up all reactions, causing the system to establish equilibrium more quickly. If
the goal is to produce the most product in the least time, then the fact that a higher
temperature increases the reaction rate is the more important point.
(d) Adding more reactants continuously and removing product continuously will constantly
shift the reaction equilibrium to the product side, creating more efficient production of
the product.
Investigation 15.3: Testing Le Châtelier’s Principle
(Pages 696, 701–703)
Purpose
The purpose of this investigation is to test Le Châtelier’s principle by applying stress to four
different chemical equilibria.
Problem
How does applying stresses to particular chemical equilibria affect the systems?
Part I CoCl42–(alc) + 6 H2O(alc) p Co(H2O)62+(alc) + 4 Cl–(alc) + energy
Part II H2Tb(aq) + p H+(aq) + HTb–(aq)
HTb–(aq) + p H+(aq) + Tb2–(aq)
Part III Fe3+(aq) + SCN–(aq) p FeSCN2+(aq)
Part IV Cu(H2O)42+(aq) + 4 NH3(aq) p Cu(NH3)42+(aq) + 4 H2O(l)
Prediction
Part I Cobalt(II) Complexes
According to Le Châtelier’s principle,
(a) adding water will shift the equilibrium to the right
(b) adding saturated silver nitrate will shift the equilibrium to the right (by removing Cl– as
AgCl(s)).
(c) adding heat will shift the equilibrium to the left
Copyright © 2007 Thomson Nelson
Unit 8 Solutions Manual
609
Part II Thymol Blue Indicator
According to Le Châtelier’s principle,
(a) adding HCl(aq) (increasing [H+(aq)]) will shift the equilibrium to the left.
(b) adding NaOH(aq) (decreasing [H+(aq)] by reaction with OH–(aq)) will shift the
equilibrium to the right.
Part III Iron(III)–Thiocyanate Equilibrium
According to Le Châtelier’s principle,
(a) adding Fe(NO3)3(aq) will shift the equilibrium to the right.
(b) adding KSCN(aq) will shift the equilibrium to the right.
(c) adding NaOH(aq) will shift the equilibrium to the left (by removing Fe3+(aq) as
Fe(OH)3(s)).
Part IV Copper(II) Complexes
According to Le Châtelier’s principle,
(a) adding NH3(aq) will shift the equilibrium to the right.
(b) adding HCl(aq) will shift the equilibrium to the left (by removing some NH3(aq) in an
acidbase reaction with HCl(aq)).
Evidence
Part I Cobalt(II) Complexes
(a) Adding water turned the solution more pink.
(b) Adding saturated silver nitrate turned the solution pink.
(c) Adding heat turned the solution more blue.
Part II Thymol Blue Indicator
(a) Adding HCl(aq) turned the yellow solution red.
(b) Adding NaOH(aq) turned the yellow solution blue.
Part III Iron(III)–Thiocyanate Equilibrium
(a) Adding Fe(NO3)3(aq) turned the solution more red.
(b) Adding KSCN(aq) turned the solution more red.
(c) Adding NaOH(aq) turned the solution more yellow.
Part IV Copper(II) Complexes
(a) Adding NH3(aq) turned the solution more blue.
(b) Adding HCl(aq) turned the solution less blue.
Analysis
According to the colour-change evidence gathered in this experiment, the systems are affected in
the same way as predicted above using Le Châtelier’s principle.
(Note that Part 1(a) is a special solution equilibrium situation, using alcohol as the solution
solvent. In this case, adding liquid water increases the concentration of water and causes an
equilibrium shift —unlike the common situation in dilute aqueous solutions, where the
concentration of water remains essentially constant.)
Evaluation
All parts of the Prediction are judged to be verified because the evidence (colour changes) was as
expected. Part I (b) is a little uncertain because the water in the saturated silver nitrate solution
could be responsible for at least part of the colour change. Le Châtelier’s principle is judged to be
acceptable because the prediction was verified in all cases.
The Purpose was accomplished. This investigation was a good test, lending further
confidence in the use of Le Châtelier’s principle.
610
Unit 8 Solutions Manual
Copyright © 2007 Thomson Nelson
Case Study: Urea Production in Alberta
(Pages 696–697)
1. CO2(g) p CO2(l)
NH3(g) p NH3(l)
In both cases, increased pressure (higher gas concentration) on the left forces the equilibrium
to the right. Pressure does not affect pure liquids so the equilibrium shifts to the right to try to
reduce the gas pressure of the reactant in each equilibrium. Neither equilibrium position is
affected if the liquid phases react to form ammonia.
2. (a) Heat energy would be written on the product side as this should be an exothermic
process. This is according to the generalization that most solids dissolve better at higher
temperatures—that is, the dissolving of most solids is endothermic.
(b) Urea is predicted to be highly soluble because it is a polar molecule with four N—H
bonds and two lone pairs on N atoms. All six of these sites can ‘hydrogen bond’ to water
molecules.
(c) A series of sieves, with successively finer holes, could be used to separate out first the
largest granules, and then successively smaller ones. The smallest granules would be
collected last.
Extension
3. Friedrich Wöhler’s classic experiment was the production of an organic compound (urea) by
heating the inorganic reactant, ammonium cyanate. The reaction is:
NH4OCN(s) + heat o NH2CONH2(s)
4. (Students’ presentations will vary, but should include the following information.)
The current cost of urea is about $200/t.
Lab Exercise 15.C: The Nitrogen Dioxide–Dinitrogen Tetraoxide Equilibrium
(Page 698)
Purpose
The purpose of this exercise is to use Le Châtelier’s principle to predict the response of an
equilibrium to an introduced change in conditions.
Problem
How does increasing the pressure affect the nitrogen dioxide-dinitrogen tetraoxide equilibrium?
Prediction
The nitrogen dioxide-dinitrogen tetraoxide equilibrium is represented by the equation:
p
2 NO2(g)
N2O4(g) + energy
colourless
reddish brown
Copyright © 2007 Thomson Nelson
Unit 8 Solutions Manual
611
According to Le Châtelier’s principle, increasing the pressure on the nitrogen dioxidedinitrogen tetraoxide equilibrium causes the equilibrium to shift to the left, which decreases the
total number of gas molecules in the system (and thus, the total pressure). (Two moles of N2O4(g)
produce one mole of NO2(g).)
Analysis
When the plunger on the syringe is depressed, the intensity of the orange-brown colour of the gas
in the syringe at first increases and then decreases. The initial increase in colour intensity is likely
due to forcing the gases into a smaller volume, which increases the concentration (pressure) of
NO2(g) and intensifies its colour. The orange-brown colour fades as the system changes to
establish a new equilibrium state by shifting to the left, consuming some of the reddish-brown
NO2(g) to form more of the colourless N2O4(g). The decrease in NO2(g) concentration during the
equilibrium shift does not completely counteract the original increase, however, so the colour at
the final equilibrium is slightly darker than it was at the initial equilibrium.
Evaluation
The experimental Design, compressing a sample of nitrogen dioxide gas in a syringe, is judged to
be adequate because this experiment produced the type of evidence needed to answer the
problem. I am quite certain of the evidence, which is a very simple observation.
The Prediction is verified because the qualitative observations clearly indicate that the
equilibrium shifts to the left when the pressure is increased on the nitrogen dioxidedinitrogen
tetraoxide equilibrium. Le Châtelier’s principle is judged to be acceptable because the prediction
was verified.
This investigation is an acceptable (if limited) test of Le Châtelier’s principle, and
therefore accomplishes the Purpose.
Investigation 15.4: Studying a Chemical Equilibrium System
(Pages 698, 703)
Purpose
The purpose of this investigation is to use Le Châtelier’s principle to solve a problem concerning
the effect of an energy change on the following equilibrium system.
+
SCN–(aq) p
FeSCN2-(aq)
Fe3+(aq)
(almost colourless)
(colourless)
(red)
Problem
Is the iron(II) thiocyanate equilibrium endothermic or exothermic?
Prediction
On the basis of the theory that forming bonds releases energy and that breaking bonds requires
energy, the equilibrium system as written is exothermic.
Fe3+(aq) + SCN–(aq) p FeSCN2+(aq) + energy
Design
Aqueous solutions containing iron(III) ions and thiocyanate ions are mixed to obtain an
equilibrium system with an intermediate colour. Separate samples of this mixture are then heated
and cooled while any colour change is observed compared with an unaltered sample (the control).
The temperature of the system is the manipulated variable and the colour of the system is the
responding variable.
Materials
Ɣ lab apron
Ɣ eye protection
612
Unit 8 Solutions Manual
Ɣ distilled water
Ɣ ice water
Copyright © 2007 Thomson Nelson
Ɣ
Ɣ
Ɣ
iron(III) nitrate solution
potassium thiocyanate solution
100 mL beaker
Ɣ hot water
Ɣ three medium test tubes with rack
Procedure
1. Mix the iron(III) nitrate solution with the potassium thiocyanate solution to obtain an
equilibrium mixture that is light red. Dilute the mixture if necessary.
2. Half fill each of the three test tubes with the equilibrium mixture.
3. Place one test tube of the mixture into hot water and the other into ice water.
4. Observe and record any colour changes compared with the third test tube.
5. Dispose of the solutions into the sink.
Evidence
Ɣ The equilibrium mixture placed in hot water decreased in colour intensity compared with the
control.
Ɣ The equilibrium mixture placed in ice water increased in colour intensity, that is, the solution
became a darker red than the control.
Analysis
According to the evidence gathered in this investigation and Le Châtelier’s principle, the reaction
as written above is exothermic. The decrease in colour intensity when the temperature is
increased indicates that the equilibrium system shifts to the left to counteract the added energy.
This would only happen if energy is a product of the forward reaction. An interpretation of the
evidence from the cold-water bath provides the same answer.
Evaluation
The Designusing hot and cold water to shift the equilibrium systemwas adequate because it
provided an answer to the problem. This Design was simple and efficient. The Procedure,
Materials, and skills are judged to be adequate because they were simple and easy to use. Using a
light or a white background could have helped in observing colour changes. I am quite confident
in the evidence obtained.
The Prediction, based on bond energy theory, is judged to be verified because the
experimental results clearly agree with the predicted results. Because the prediction was verified,
the bond energy concept used to make the prediction appears to be acceptable.
The Purpose of this investigation was accomplished, but many more problems concerning
the effect of an energy change could be investigated for other systems.
Web Activity: Web QuestPoison Afloat
(Page 698)
[Students’ reports should cover the following information.]
What was the most likely cause of death?
Ɣ The symptoms that the victim exhibitedcherry-red blood, dizziness, and nauseaare
consistent with both cyanide and carbon monoxide poisoning. The fact that the victim spent a
long time in an enclosed space in which a propane heater was operating suggests that carbon
monoxide poisoning was a more likely cause of death than cyanide poisoning. In addition, the
absence of the smell of bitter almonds points away from cyanide poisoning.
What is the likely source of the substance that killed the victim?
Ɣ Carbon monoxide is an odourless, colourless gas produced during the incomplete combustion
of hydrocarbon fuels such as propane. While the victim slept, the propane heater most likely
started to produce carbon monoxide. Because the cabin was enclosed, with little additional
oxygen entering the space, carbon monoxide accumulated until it reached lethal
concentrations.
Copyright © 2007 Thomson Nelson
Unit 8 Solutions Manual
613
Use Le Châtelier’s principle to explain why this substance is dangerous.
Ɣ Hemoglobin, an iron-containing protein in red blood cells, is responsible for oxygen transport
within the body. Hemoglobin transport of oxygen is necessary because the solubility of
oxygen in blood is too low to permit the transport of sufficient oxygen required by body cells.
Hemoglobin (Hb) binds with oxygen to form oxyhemoglobin, as given in the following
equilibrium:
Hb(aq) + O2(aq) p HbO2(aq)
(1)
Carbon monoxide binds even more readily with hemoglobin to produce carboxyhemoglobin:
(2)
Hb(aq) + CO(aq) p HbCO(aq)
When oxygen and carbon monoxide are both present, the preferential binding of carbon
monoxide for hemoglobin effectively displaces oxygen, as given by:
(3)
HbO2(aq) + CO(g) p HbCO(aq) + O2(g)
The value of the equilibrium constant for equilibrium (3) is 210, indicating that the position
of the equilibrium strongly favours carboxyhemoglobin. Consequently, carbon monoxide
poisoning prevents hemoglobin from transporting oxygen, which eventually results in death.
What is the recommended treatment for poisoning involving this substance?
Ɣ Patients suffering from carbon monoxide poisoning are often treated by having them breathe
100% oxygen, either through a mask or in a hyperbaric chamber. The increase in the oxygen
concentration shifts equilibrium (3) to the left, favouring oxyhemoglobin, displacing carbon
monoxide from the body.
Describe three other scenarios that could result in the same type of poisoning.
Ɣ Other scenarios that could result in carbon monoxide poisoning include:
A clogged chimneyRising warm air in a chimney gradually cools, causing the water vapour
that it is carrying to condense. In the absence of a chimney liner, the moisture can damage the
mortar and brick of the chimney, causing bits of material to fall. As this debris accumulates, it
can clog the chimney. Consequently, carbon monoxide present in the chimney gases will be
sent back into the house.
A faulty furnaceAs a furnace ages and deteriorates, it can develop cracks in its heat
exchanger (the device responsible for transmitting heat from the flame to the air that is
circulated throughout the house). Cracks in the heat exchanger can cause the production of
carbon monoxide.
Automobile exhaustAutomobile emissions contain carbon monoxide. Running a vehicle in
an enclosed space, such as a garage attached to a house, can allow carbon monoxide to
diffuse into the house.
Section 15.2 Questions
(Page 699)
1. (a) The equilibrium shifts to the right in an attempt to remove the added oxygen. Adding
O2(g) increases its concentration and thus increases the number of collisions that form
products, so the forward reaction rate is increased while the reverse rate is unchanged.
(b) The equilibrium shifts to the left in an attempt to remove the added energy. Both reaction
rates increase as faster moving molecules collide more often, but the reverse rate is
increased more than the forward rate.
(c) The equilibrium shifts to the right in an attempt to replace the removed nitrogen
monoxide. Removing NO(g) decreases the number of reverse reaction collisions, thus
slowing the reverse reaction rate without changing the forward rate.
(d) The equilibrium shifts to the left in an attempt to reduce the pressure by shifting to the
side with the fewer molecules. Reducing the volume concentrates all the reaction
614
Unit 8 Solutions Manual
Copyright © 2007 Thomson Nelson
constituents, making both forward and reverse reaction collisions more frequent; but the
reverse reaction rate is increased more, because more molecules are involved.
2. Decrease the temperature, decrease the container volume to increase the overall pressure,
increase [N2(g)], increase [H2(g)], and remove NH3(g) from the system.
3. (a)
The equilibrium shifts to the left, reducing [Cu(H2O)42+(aq)] and [Cl–(aq)], although the
final [Cl–(aq)] is higher than the original equilibrium concentration.
(b)
Adding silver nitrate precipitates AgCl(s), thereby removing Cl–(aq) from the system.
The equilibrium shifts to the right, increasing [Cu(H2O)42+(aq)] and [Cl–(aq)], although
the final [Cl–(aq)] is lower than the original equilibrium concentration.
4. At A, the concentration (or pressure) of every chemical in the system is decreased by
increasing the container volume.
At B, the temperature is increased.
At C, C2H6(g) is added to the system.
At D, no shift in equilibrium position is apparent; the change imposed must be addition of a
catalyst, or of a substance that is not involved in the equilibrium reaction.
5. Percent yield would increase in (a) and (d). (Note that the forward process in both cases is
endothermic.)
6. (a) The concentration of methane is increased. This results in more frequent favourable
collisions with chlorine molecules. As a result, the equilibrium would shift to the right, in
the direction of products.
Copyright © 2007 Thomson Nelson
Unit 8 Solutions Manual
615
(b) Since the number of moles of gas entities on the reactant and product side is equal, this
system is insensitive to overall changes in pressure due to volume changes. Both forward
and reverse rates are affected equally. As a result, increasing the volume of the container
would have no effect on the position of the equilibrium.
(c) Removing energy (decreasing temperature) from this exothermic process would cause the
equilibrium to shift to the right, to the product side. Both the forward and reverse rates
will decrease, but the reverse rate will decrease more.
(d) Introducing a catalyst serves only to increase the rate of both the forward and reverse
reaction rates equally. It has no effect on the position of equilibrium.
7. (a) Kc is unchanged.
(b) Kc is unchanged.
(c) Kc will decrease because the forward reaction is endothermic and the equilibrium shifts to
the left. (Only temperature affects the value of an equilibrium constant.)
(d) Kc is unchanged.
Extension
8. Nitrogen “narcosis” is a loss of rationality (an effect on the brain) caused by too high a
nitrogen concentration in the blood at high pressures. The “bends” are an agonizing and
potentially lethal condition caused by dissolved nitrogen leaving the blood as bubbles of gas
when pressure is removed too quickly (when a diver rises). The bubbles block capillaries in
the circulatory system.
Chapter 15 SUMMARY
Make a Summary
(Page 704)
1.
2. (1) The rate at which effective collisions of reactants are occurring (resulting in the
formation of products) exceeds the rate at which effective collisions of products are
occurring (resulting in the formation of reactants).
616
Unit 8 Solutions Manual
Copyright © 2007 Thomson Nelson