7.3
Qualitative Changes in
Equilibrium Systems
Consider again the student who is attempting to maintain her position while walking up
the “down” escalator. What happens if the escalator suddenly starts moving faster? She
notices that her steps are no longer keeping up with the escalator’s new speed. She is
drifting downward. To reestablish equilibrium she must increase her stepping rate to
match the movement of the escalator. After doing so, she is again in dynamic equilibrium,
but at a lower level on the escalator. The student, wanting to remain at rest, counteracts
changes in the system in order to maintain her state of dynamic equilibrium. In doing
so, she establishes a new equilibrium at a different level.
Similar adjustments occur when chemical systems at equilibrium are disturbed.
Le Châtelier’s Principle
In 1884, the brilliant French chemist Henry Louis Le Châtelier (Figure 1) made observations of chemical systems at equilibrium. As a result of his studies, he developed a
generalization that has become one of the most powerful laws of science:
Figure 1
Henry-Louis Le Châtelier,
1850–1936, a French chemist and
engineer, worked in chemical industries. To maximize the yield of products, Le Châtelier used systematic
trial and error. After measuring
properties of many equilibrium
states in chemical systems, he discovered a pattern and stated it as a
generalization. This generalization
has been supported extensively by
evidence and is now considered an
extremely useful tool in chemistry. It
has become known as Le Châtelier’s
principle.
equilibrium shift movement of a
system at equilibrium, resulting in a
change in the concentrations of
reactants and products
Le Châtelier’s Principle
When a chemical system at equilibrium is disturbed by a
change in a property, the system adjusts in a way that
opposes the change.
The application of Le Châtelier’s principle involves an initial equilibrium state, a
shifting “non-equilibrium” state, and a new equilibrium state.
Le Châtelier’s principle provides a method of predicting the response of a chemical
system to a change of conditions. Using this simple approach, chemical engineers can produce more of the desired product, making technological processes more efficient and more
economical. For example, Fritz Haber used Le Châtelier’s principle to devise a process
for the economical production of ammonia from atmospheric nitrogen. (See the Haber
process, Figure 6, Section 7.2.)
Le Châtelier’s Principle and Concentration Changes
Le Châtelier’s principle predicts that, if more of a reactant is added to a system at equilibrium, then that system will undergo an equilibrium shift.
The effect of adding more of a reactant is that we first observe the reactant concentration decreasing, as some of the added reactant changes to products. This period of
change ends with the establishment of a new equilibrium state, in which concentrations
are usually different from their original values. The system changes in a way that opposes
the change. For example, the production of Freon-12, a chlorofluorocarbon (CFC),
involves the following equilibrium reaction:
CCl4(l) 2 HF(g) e CCl2F2(g) 2 HCl(g)
Freon-12
To improve the yield of Freon-12, more hydrogen fluoride is added to the equilibrium
system. The additional reactant disturbs the equilibrium state and the system shifts to the
right, consuming some of the added hydrogen fluoride by reaction with carbon tetrachloride. As a result, more Freon-12 is produced and a new equilibrium state is obtained.
450 Chapter 7
NEL
Section 7.3
Concentration (mol/L)
Concentration (mol/L)
2 CO(g) + O2(g)
2 CO2(g)
In chemical equilibrium shifts, the
imposed concentration change is normally
only partially counteracted, and the conCO2(g) added
centrations of the reactants and products in
the final equilibrium are usually different
from the values in the original equilibrium
[CO2(g)]
state. This can be seen in the concentra[CO(g)]
tion–time graph in Figure 2. Notice that
[O2(g)]
the graph has three lines representing
changes in the concentrations of CO2(g),
CO(g) and O2(g). The amount of CO2(g) in
Time
the vessel is reduced in the shift.
The removal of a product (if the removal
2 CO(g) + O2(g)
2 CO2(g)
decreases concentration) will also shift an
equilibrium forward, to the right. The
CO(g) removed
carbon dioxide reaction can be shifted forward by removing either gaseous product
(Figure 3), since decreasing the amount of
a gas lowers its concentration in the reac[CO2(g)]
tion container.
[CO(g)]
The effects of forward and reverse shifts
in equilibrium are shown in the iron (III)
[O2(g)]
thiocyanate system in Figure 4 (p. 452).
Adjusting an equilibrium state by adding
Time
and/or removing a substance is a common
application of Le Châtelier’s Principle. For industrial chemical reactions, engineers strive
to design processes where reactants are added continuously and products are continuously removed, so that an equilibrium is never allowed to establish. If the reaction is
always shifting forward, the process is always making product (and presumably, the
industry is always making money).
Consider an industrial example, the final step in the production of nitric acid, which
is represented by the reaction
Figure 2
The reaction establishes an equilibrium that is then disturbed (at the
time indicated by the vertical dotted
line) by the addition of CO2(g). Some
of the added CO2(g) reacts,
decreasing its concentration, while
the concentration of both products
increases until a new equilibrium is
established. The concentrations eventually become constant again, at a
new level. However, the initial K value
and the final K value are the same.
Figure 3
The reaction establishes an equilibrium that is disturbed (at the time
indicated by the vertical dotted line)
by the removal of CO(g). The equilibrium shifts forward; the concentration of O2(g) increases while the
concentration of CO2(g) decreases,
until a new equilibrium is established. The initial K value and the
final K value are the same.
air
CO2
O2
air-filled
sacs in
lungs
3 NO2(g) H2O(l) e 2 HNO3(aq) NO(g)
In this process, nitrogen monoxide gas is removed from the chemical system by a further
reaction with oxygen gas. The removal of the nitrogen monoxide causes the system to shift
to the right—more nitrogen dioxide and water react, replacing some of the removed nitrogen
monoxide. As the system shifts, more of the desired product, nitric acid, is produced.
A vitally important biological equilibrium is that of hemoglobin, Hb(aq), (a protein complex in red blood cells that transports oxygen around the body), oxygen, and oxygenated
hemoglobin, HbO2(aq).
Hb(aq) O2(aq) e HbO2(aq)
As blood circulates to the lungs, the high concentration of dissolved oxygen shifts the
equilibrium to the right and the hemoglobin becomes oxygenated (Figure 5). As the
blood circulates throughout the body, cellular respiration consumes oxygen. This removal
of oxygen shifts the equilibrium to the left, and more oxygen is released.
NEL
heart
capillary
blood
vessels
body cells
Figure 5
Oxygenated blood from the lungs is
pumped by the heart to body tissues. The deoxygenated blood
returns to the heart and is pumped
to the lungs. Shifts in equilibrium in
this system occur continuously.
Chemical Systems in Equilibrium 451
Figure 4
Disturbing the iron(III) thiocyanate
equilibrium
(a)
3
Fe(aq)
added
2
FeSCN(aq)
Fe3
(aq) SCN(aq)
SCN(aq) added
(c)
2
FeSCN(aq)
Fe3
(aq) SCN(aq)
(d)
2
FeSCN(aq)
added
2
FeSCN(aq)
Fe3
(aq) SCN(aq)
2
FeSCN(aq)
Concentration (mol/L)
Fe3
(aq) SCN(aq)
(b)
[Fe3
(aq) ]
[SCN
(aq) ]
2
[FeSCN(aq)
]
Time
(a)
(b)
(c)
(d)
When solutions containing Fe3
(aq) (colourless)
and SCN
(aq) (brown) are
mixed, an equilibrium is
reached with the product,
FeSCN
(aq) (deep red), as
shown by the constant,
uniform light brown
colour of the equilibrium
solution. On the graph,
notice that the concentrations of Fe3
(aq) and
SCN
(aq) drop after mixing
as they react to form
FeSCN2
(aq). All three concentrations become constant when equilibrium is
reached (flat lines).
3 is added. In
Fe(aq)
response the system shifts
to the right, producing
2 .
more red FeSCN(aq)
Notice the spike in the
3 when more
graph of Fe(aq)
is added, and that the
3
concentration of Fe(aq)
subsequently drops. The
2
concentration of FeSCN(aq)
rises as more is produced.
ions are used
As SCN(aq)
up, the concentration
drops. Equilibrium is
reestablished at a new
level (flat lines).
The addition of more
solution containing
SCN
(aq) shifts the equilibrium to the right, producing more of the dark
red FeSCN2
(aq) ions. Note
the corresponding
changes in the graph.
ions to
Adding FeSCN(aq)
the mixture forces the
equilibrium to shift
toward the reactants,
giving the solution a
paler colour. Note the
corresponding changes
in the graph
2
Rate Theory and Concentration Changes
Kinetic theory provides a simple explanation of the equilibrium shift that occurs when
a reactant concentration is increased. We assume that when reactant is added, with more
reactant particles present per unit volume, collisions are suddenly much more frequent
for the forward reaction. This increases the forward rate significantly. Since the reverse
reaction rate is not immediately changed, the rates are no longer equal, and for a time
the difference in rates results in an observed increase of products.
Of course, as the concentration of products increases, so does the reverse reaction rate.
The forward rate decreases as reactant is consumed, until eventually the two rates (forward and reverse) become equal to each other again. The rates at the new equilibrium state
are faster than those at the original state, because the system now contains a larger number
of particles (and therefore, there are more collisions) in dynamic equilibrium.
452 Chapter 7
NEL
Section 7.3
If a substance is removed, causing an equilibrium shift, the explanation is similar
except that the initial effect is to suddenly decrease one of the equilibrium reaction rates.
Remember that the addition or removal of a substance in solid or liquid state does not
change the concentration of that substance. The reaction of condensed phases (solids
and liquids) takes place only at an exposed surface—and if the surface area exposed is
changed it is always exactly the same change in available area for both forward and
reverse reaction collisions. The forward and reverse rates will change by exactly the same
amount if they change at all, so equilibrium is not disturbed and no shift occurs.
Le Châtelier’s Principle and Temperature Changes
The energy in a chemical equilibrium equation can be treated as though it were a
reactant or a product.
Endothermic reaction:
reactants energy e products
Exothermic reaction:
reactants e products energy
Energy can be added to or removed from a system by heating or cooling the container.
In either situation, the equilibrium shifts to minimize the change. If the system is cooled,
the system tries to “warm” itself and the equilibrium shifts in the direction that produces
heat. If heat is added, the equilibrium shifts in the direction that absorbs heat.
The equilibrium between two oxides of nitrogen illustrates the effect of temperature
on a system at equilibrium. The equation for this reaction is
hot water (85°)
N204(g)
2N02(g)
ice water (0°)
Figure 6
Each of these flasks contains an
equilibrium mixture of dinitrogen
tetroxide and nitrogen dioxide. Shifts
in equilibrium can be seen when one
of the flasks is heated or cooled.
N2O4(g) energy e 2 NO2(g)
colourless
reddish-brown
2 SO2(g) O2(g) e 2 SO3(g) energy
Removing energy (cooling) causes the system to shift to the right. This shift yields
more sulfur trioxide while at the same time partially replacing the energy that was removed.
Rate Theory and Energy Changes
Kinetic theory explains the equilibrium shift that occurs when the energy of a system at
equilibrium is increased or decreased. Consider the contact process reaction equation mentioned above—a typical exothermic reaction.
2 SO2(g) O2(g) e 2 SO3(g) energy
Rate theory explains the result of cooling this exothermic system by assuming that
both forward and reverse reaction rates are slower at the lower temperature, but that
the reverse rate decreases more than the forward rate. While the rates remain unequal,
the observed result is the production of more product and more energy. The shift causes
concentration changes that will increase the reverse rate and decrease the forward rate
until they become equal again, at the new, lower temperature (Figure 7).
NEL
Concentration (mol/L)
When the system at equilibrium is heated, the reaction shifts to the right, increasing
the concentration of nitrogen dioxide. This is made visible by the intensification of the
reddish-brown colour of the reaction mixture (Figure 6). The energy that is added
causes the system to shift to the right, absorbing some of the added energy.
2 SO2(g) + O2(g)
In the exothermic production of sulfur trioxide, as part of the contact process for
making sulfuric acid, the product is favoured if the temperature of the system is kept low.
2 SO3(g) + energy
temperature
decreased
[SO2(g)]
[O2(g)]
[SO3(g)]
Time
Figure 7
The reaction establishes an equilibrium that is disturbed (at the time
indicated by the vertical dotted line)
by a decrease in temperature. The
equilibrium shifts forward,
increasing the concentration of
SO3(g) product while decreasing the
concentration of both reactants, until
a new equilibrium is established.
Chemical Systems in Equilibrium 453
CO2(g) + energy
2 CO(g) + O2(g)
Concentration (mol/L)
temperature
decreased
For endothermic reactions, the situation is reversed. Consider an endothermic reaction such as the decomposition of carbon dioxide into carbon monoxide and oxygen.
CO2(g) 566 kJ e 2 CO(g) O2(g)
[CO2(g)]
[CO(g)]
[O2(g)]
Cooling the system, and maintaining it at a lower temperature, causes the rate of the
forward and reverse reactions to decrease, but the rate of the forward reaction decreases
more than the reverse reaction. While the rates remain unequal, the observed result is the
production of more reactant and more energy. The shift causes concentration changes
that will increase the reverse rate and decrease the forward rate until they become equal
again, at the new, lower temperature (Figure 8).
Time
Figure 8
The reaction establishes an equilibrium that is then disturbed (at the
time indicated by the vertical dotted
line) by a decrease in temperature.
The equilibrium shifts in the reverse
direction, increasing the concentration of CO2(g) while decreasing the
concentrations of CO(g) and O2(g),
until a new equilibrium is established.
Le Châtelier’s Principle and Gas Volume Changes
According to Boyle’s law, the concentration of a gas in a container is directly related to
the pressure of the gas.
Decreasing the volume to half its original value doubles the concentration of every gas
in the container. Changing the volume of any equilibrium system involving gases may
cause a shift in the equilibrium. To predict whether a change in pressure will affect the
equilibrium state, you must consider the total number of moles of gas reactants and the
total number of moles of gas products. For example, in the equilibrium reaction where
nitrogen and hydrogen form ammonia, four moles of gaseous reactants produce two
moles of gaseous product.
N2(g) 3 H2(g) e 2 NH3(g)
1 mol 3 mol
4 mol
2 mol
If the volume of the vessel containing this reaction mixture is decreased, the overall
pressure is increased. Le Châtelier’s principle suggests that the system will react in a way
that resists the change—i.e., in a way that reduces the pressure. In this case, the equilibrium will shift to the right, which decreases the number of gas molecules in the container and reduces the pressure (Figure 9).
Key:
N2(g)
H2(g)
NH3(g)
(a)
Figure 9
(a) An equilibrium mixture containing N2(g), H2(g), and NH3(g)
(b) The volume is decreased,
increasing the pressure.
(c) The reaction shifts to the right,
toward the side with fewer molecules, to relieve pressure.
454 Chapter 7
(b)
(c)
If the volume of the vessel is increased, the pressure is decreased, and the shift is in the
opposite direction, to the left, which counteracts the change by producing more gas
molecules.
NEL
Section 7.3
A system with equal numbers of gas molecules on each side of the equation will not shift
after a change in volume, since no shift can change the pressure in the vessel. Consider the
equilibrium reaction between hydrogen and iodine to produce hydrogen iodide.
H2(g) I2(g) e 2 HI(g)
2 mol
Systems involving only liquids or solids are not affected appreciably by changes in
pressure. Substances in these condensed states are virtually incompressible, and so reactions involving them cannot counteract pressure changes.
Rate Theory and Gas Volume Changes
We can again explain the equilibrium shift observed when the volume of a system
involving gaseous reactants and products is changed, as an imbalance of reaction rates.
2 SO2(g) O2(g) e 2 SO3(g) 198 kJ
2 mol + 1 mol
3 mol
2 mol
Kinetic theory explains the effect of a decrease in volume by assuming that both the forward and reverse reaction rates increase because the concentrations (partial pressures) of
reactants and products increase. However, the forward rate increases more than the reverse
rate because there are more particles involved in the forward reaction. This means that the
increase in the total number of effective collisions is greater for the forward reaction.
Again, while the rates remain unequal, the observed result is the production of more
product. The shift causes concentration changes that eventually increase the reverse rate
and decrease the forward rate until they become equal again (Figure 10).
2 SO2(g) + O2(g)
Concentration (mol/L)
1 mol + 1 mol
2 mol
2 SO3(g) + energy
Volume of container
decreased
[SO2(g)]
[O2(g)]
[SO3(g)]
Time
Figure 10
The equilibrium is disturbed by a
decrease in the volume of the container (at the time indicated by the
vertical dotted line). The equilibrium
shifts forward (i.e., toward products). The concentration of SO3(g)
increases while the concentration of
reactants decreases, until a new
equilibrium state is established.
Changes That Do Not Affect the Position of
Equilibrium Systems
We have looked at three changes that have an effect on the equilibrium of a chemical
system—concentration changes, energy changes, and pressure (volume) changes. There
are other changes that have no effect whatever on the equilibrium position of a chemical system. Here is a brief look at these changes, and why they have no effect.
Adding Catalysts
Catalysts are used in most industrial chemical reactions and biological systems. A catalyst decreases the time required to reach the equilibrium position, but does not affect the
final position of equilibrium. The presence of a catalyst in a chemical reaction system
lowers the activation energy for both forward and reverse reactions by an equal amount,
so the equilibrium establishes much more rapidly but at the same position as it would
without the catalyst (Figure 11). Forward and reverse rates are increased equally. The value
of catalysts in industrial processes is to decrease the time required for equilibrium shifts
to occur, allowing a more rapid overall production of the desired product.
NEL
Chemical Systems in Equilibrium 455
(a)
(b)
uncatalyzed reaction
catalyzed reaction
without
catalyst
activated
complexes
slow
activation energy
of uncatalyzed
products
reaction
Ep
reactants
activation energy
of catalyzed
reaction
before equilibrium
∆H
with
catalyst
Reaction Progress
Figure 11
(a) A catalyst reduces the activation energy by the same amount
whether the reaction proceeds to the right or to the left. (b) It does
not affect the relative concentrations of entities.
INVESTIGATION 7.3.1
Testing Le Châtelier’s Principle
(p. 514)
Le Châtelier’s Principle provides
guidance when predicting how a
chemical system at equilibrium will
respond to disturbance.
at equilibrium
fast
before equilibrium
at equilibrium
Adding Inert Gases
The pressure of a gaseous system at equilibrium can be changed by adding a gas while
keeping the volume constant. If the gas is inert in the system, for example, if it is a noble
gas or if it cannot react with the entities in the system, the equilibrium position of the
system will not change. We can explain this by turning to rate theory: The presence of
the inert gas changes the probability of successful collisions for both the reactants and
the products equally, resulting in no shift in the equilibrium system (Figure 12).
H2 molecule
N2 molecule
SUMMARY
Variables Affecting Chemical Equilibria
Table 1
NH3 molecule
Variable
Type of Change
Response of System
concentration
increase
shifts to consume some of the added reactant
or product
decrease
shifts to replace some of the removed
reactant or product
increase
shifts to consume some of the added thermal
energy
decrease
shifts to replace some of the removed thermal
energy
increase
(decrease in pressure)
shifts toward the side with the larger total
amount of gaseous entities
decrease
(increase in pressure)
shifts toward the side with the smaller total
amount of gaseous entities
temperature
He atom
volume
Figure 12
Adding an inert gas such as helium
to a system at equilibrium increases
the pressure of the system (at constant volume) but does not cause a
shift in the equilibrium position.
456 Chapter 7
Variables That Do Not Affect Chemical Equilibria
catalysts
—
no effect
inert gases
—
no effect
NEL
Section 7.3
DID YOU
Practice
Understanding Concepts
1. How will the following system at equilibrium shift in each of the following cases?
2 SO3(g)
(a)
(b)
(c)
(d)
(e)
e 2 SO2(g) O2(g)
H ° 197 kJ
SO2(g) is added
the pressure is decreased by increasing the volume of the container
the pressure is increased by adding Ne(g)
the temperature is decreased
O2(g) is removed
2. (a) Draw a concentration–time graph (similar to Figure 10) that illustrates the
changes in concentration that occur when F2(g) is added to a sealed vessel
containing the following equilibrium:
Br2(g) 5 F2(g)
e 2 BrF5(g)
(b) Draw a concentration–time graph for the removal of some HOCl(g) from the following equilibrium:
H2O(g) Cl2O(g)
e HOCl(g)
3. The following reaction is used in the commercial production of hydrogen gas.
CH4(g) 2 H2O(g)
e CO2(g) 4 H2(g)
(a) In a closed system, how would a catalyst affect the establishment of equilibrium
in the system?
(b) How would the concentration of H2(g) at equilibrium be affected by the use of a
Ni(s) catalyst?
4. Much of the brown haze hanging over large cities is nitrogen dioxide, NO2(g). Nitrogen
dioxide reacts to form dinitrogen tetroxide, N 2O4(g), according to the equilibrium:
2 NO2(g)
(brown)
e N2O4(g) 57.2 kJ
(colourless)
Use this equilibrium to explain why the brownish haze over a large city disappears in
the winter, only to reappear again in the spring. (Assume that the atmosphere over
the city constitutes a closed system.)
5. A land-based vehicle will be an important part of any
future exploration of the planet Mars. One proposed
design for a Mars rover uses a methane gas fuel cell as
its power supply. The methane fuel can be made on
Mars using a chemical reaction that has been known
for over 100 a — the Sabatier methanation reaction:
CO2(g) 4 H2(g)
KNOW
?
No Sweat
Chickens cannot perspire. When a
chicken gets hot, it pants like a
dog. Farmers have known for a long
time that chickens lay eggs with
thin shells in hot weather. These
fragile eggs are easily damaged.
Eggshell is primarily composed of
calcium carbonate, CaCO3(s). The
source of the carbonate portion of
this chalky material is carbon
dioxide, CO2, produced as a waste
product of cellular respiration. The
carbon dioxide dissolves in body
fluids forming the following equilibrium system:
CO2(g) e CO2(aq) e H2CO3(aq)
(chicken breath)
(in the blood)
H2CO3(aq) e H
(aq) HCO3(aq)
(in the blood)
2
HCO3
(aq) e CO3(aq) H(aq)
2
CO32
(aq) Ca3(aq) e CaCO3(s)
(in the blood)
(eggshell)
When chickens pant, blood carbon
dioxide concentrations are
reduced, causing a shift through
all four equilibria to the left and a
reduction in the amount of calcium
carbonate available for making
eggshells. Ted Odom, a graduate
student at the University of Illinois,
found a simple solution to the
problem: Give the chickens carbonated water to drink in the
summer. This shifts the equilibria
to the right, compensating for the
leftward shift caused by panting.
Rumour has it that chickens quite
enjoy sipping on their Perrier on
steamy summer evenings.
e CH4(g) 2 H2O(g)
H 165 kJ at 250°C
Predict, using Le Châtelier’s principle, the conditions
required in a closed Sabatier reactor to produce the
maximum amount of methane.
6. In caves we sometimes see large structures known as
stalactites and stalagmites (Figure 13).
These formations are made of crystals of the insoluble Figure 13
compound calcium carbonate, CaCO3(s), also known as
limestone. Calcium carbonate forms when dissolved
NEL
Chemical Systems in Equilibrium 457
carbon dioxide and calcium ions combine, according to the following equilibrium:
CaCO3(s) 2 H+(aq)
e Ca2+
(aq) H2O(l) CO2(g)
Use Le Châtelier’s principle to answer the following:
(a) How would the stalagmites and stalactites be affected if the water in the cave
became more acidic?
(b) How would the hardness of the water [Ca2+
(aq)] affect the growth of stalagmites
and stalactites?
Applying Inquiry Skills
7. A student designed an experiment to measure the effect of the addition of chloride
ions on the equilibrium point of the following system.
Cu2+
(aq) 4 Cl(aq)
2
e CuCl4(aq)
(blue)
(green)
Question
What effect does the addition of chloride ions have on the system at equilibrium?
Prediction
(a) Predict whether an equilibrium shift will occur and the effect of the shift on the
system.
Experimental Design
Test tubes containing 10 mL of a copper(II) chloride solution are combined with 5-mL,
10-mL, and 15-mL samples of 1 mol/L hydrochloric acid, HCl(aq). All test tubes were
stoppered, shaken, and allowed to reach equilibrium.
Evidence
Table 2 Observations on Cu2+
(aq) Equilibrium Investigation
Test tube
(10 mL of CuCl2(aq)
in each)
Vol. HCl(aq)
added
(mL)
Total vol.
after mixing
(mL)
Colour at
equilibrium
1
5 mL
15 mL
blue
2
10 mL
20 mL
blue-green
3
15 mL
25 mL
green
Analysis
(b) Use the Evidence to answer the Question.
Evaluation
(c) Critique the Experimental Design.
Making Connections
8. The digestion of some high-protein foods, such as red meat, beans, lentils, and shell-
fish, releases uric acid, HC5H3N4O3(aq), which ionizes into hydrogen ions, H+(aq) and
urate ions, C5H3N4O3
(aq), in the bloodstream. People whose kidneys do not function
properly cannot excrete urate ion sufficiently quickly, leading to an increased concentration of urate in the blood. This sometimes leads to a painful form of arthritis known
as gout, characterized by the formation of tiny needle-like crystals of sodium urate,
NaC5H3N4O3(s), in joints and tissues, according to the equation
NaC5H3N4O3(s)
e Na+(aq) C5H3N4O3
(aq)
(a) Suppose you were a nutritionist. What advice could you give to your patients who
suffer from gout? Explain why following the advice would be effective.
458 Chapter 7
NEL
Section 7.3
(b) Many women take calcium supplements on a daily basis to prevent the loss of
bone mass (a condition known as osteoporosis). If a woman suffering from
osteoporosis has gout too, she may also develop kidney stones (which can consist of calcium urate). Write a chemical equilibrium equation for this reaction and
explain why this happens.
(c) Research and report on other non-dietary treatments of gout.
GO
www.science.nelson.com
Section 7.3 Questions
Understanding Concepts
position of a chemical equilibrium. Briefly explain how
rate theory explains the effect in each case.
(b) List two factors that do not affect the position of
equilibrium.
Predict the shift in the equilibrium and draw a graph of
concentration versus time for relevant reactants to communicate the shift after the following stresses are applied to
the system:
(a) Hydrochloric acid is added.
(b) Silver nitrate is added.
2. For each of the following chemical systems at equilibrium,
5. The two oxyanions of chromium(IV) are the orange dichro-
1. (a) List three environmental factors that may affect the
3. The following equation is important in the
industrial production of nitric acid. Predict the
direction of the equilibrium shift for each of the
following changes in a closed vessel. Explain
any shift in terms of changes in rates of the forward and reverse reactions.
Concentration (mol/L)
use Le Châtelier’s principle to predict the effect of the
change imposed on the chemical system. Indicate the
direction in which the equilibrium is expected to shift, if at
all. For each example, sketch a graph of concentration
versus time, plotted from just before the possible change to
the established equilibrium.
(a) H2O(l) energy e H2O(g)
The container is heated.
(b) H2O(l) e H+(aq) OH–(aq)
A few crystals of NaOH(s) are added to the container.
(c) CaCO3(s) energy e
CaO(s) CO2(g)
CO2(g) is removed from the container.
C2H6(g)
(d) CH3COOH(aq) e H+(aq) CH3COO–(aq)
A few drops of pure CH3COOH(l) are
C2H4(g)
added to the system.
(a)
(b)
(c)
(d)
O2(g) is added to the system.
The temperature of the system is increased.
NO(g) is removed from the system.
The pressure of the system is increased by decreasing
the volume of the reaction vessel.
(e) Argon gas is added to the system without changing
the volume.
4. In a solution of copper(II) chloride, the following equilib2– 4 H O
CuCl4(aq)
2 (l)
dark green
NEL
–
e Cu(H2O)42+
(aq) 4 Cl(aq)
2– H O
Cr2O7(aq)
2 (l)
2– 2 H+
e 2 CrO4(aq)
(aq)
orange
yellow
6. Identify the nature of the change imposed on the equilib-
rium system (Figure 14) at each of the times indicated A,
B, C, D, and E.
C2H4(g) + H2(g) e C2H6(g) + energy
H2(g)
NH3(g) 5 O2(g) e 4 NO(g) 6 H2O(g) energy
rium exists:
2– , and the yellow chromate ion, CrO 2– .
mate ion, Cr2O7(aq)
4(aq)
Explain why a solution containing the following equilibrium
system turns yellow when sodium hydroxide is added.
A
B
C
Time
D
E
Figure 14
Applying Inquiry Skills
7. A student plans to test Le Châtelier’s principle by
increasing the pressure in a closed system containing a
nitrogen dioxide–dinitrogen tetroxide equilibrium,
NO2(g)–N2O4(g), and measuring changes in colour intensity.
Question
(a) State a Question for this experiment.
Prediction
(b) Predict an answer to the Question.
Experimental Design
(c) Propose an experimental design.
blue
Chemical Systems in Equilibrium 459
Evidence
ment is known as the Claus process, which involves the following reaction:
e 3 S(s) 2 H2O(g) heat
2 H2S(g) SO2(g)
The Claus process is capable of removing up to 95% of the
sulfur emissions from petroleum-processing plants.
(a) Research and report on the Claus process.
(b) Describe why it is advantageous to remove the sulfur
from the process as quickly as it forms.
GO
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10. An air purification system involving lithium hydroxide, LiOH,
was used in NASA’s Apollo missions to the moon. LiOH
absorbs carbon dioxide.
2 LiOH(s) CO2(g)
e Li2CO3(s) H2O(l)
Use Le Châtelier’s principle to explain why the amount of
time astronauts can spend in a spacecraft is limited.
Evaluation
(d) Evaluate the experimental design.
8. A student plans an investigation into the effects of various
changes on a chromate–dichromate equilibrium. The chro2–
mate ion, CrO4(aq)
, reacts with oxygen to form the following
2– , and water.
equilibrium with the dichromate ion, Cr2O7(aq)
2–
2 H+(aq) 2 CrO4(aq)
2– H O
e Cr2O7(aq)
2 (l)
(yellow)
(orange)
Experimental Design
The effects of the following changes on a chromate–dichromate equilibrium are noted:
(i) the addition of 0.1 mol/L hydrochloric acid
(ii) the addition of 0.1 mol/L sodium hydroxide solution
(iii) the addition of 0.1 mol/L barium nitrate
Prediction
(a) Predict the effect that each change would have on the
colour of a chromate-dichromate equilibrium mixture.
Synthesis
(b) After disposing of the contents of the test tube, a student discovers that the inside of the test tube is coated
with a light yellow precipitate that cannot be easily
washed off. What chemical could be added to the test
tube to remove the precipitate?
Making Connections
9. Hydrogen sulfide is a foul-smelling and toxic byproduct of
the processing of crude oil and natural gas. One method to
recover H2S so that it does not contaminate the environ-
460 Chapter 7
11. When the Olympic Games were held in Mexico in 1968,
many athletes arrived early to train in the higher altitude
(2.3 km above sea level) and lower atmospheric pressure of
Mexico City. Exertion at high altitudes, for people who are
not acclimatized, may make them “lightheaded” from lack
of oxygen. A similar effect occurred at the 2002 Winter
Olympics in Salt Lake City, Utah (1.3 km above sea level).
Use the theory of dynamic equilibrium and Le Châtelier’s
principle to explain this observation. How are people who
normally live at high altitudes physiologically adapted to
their reduced-pressure environment?
GO
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12. Hemoglobin, Hb, a protein molecule found in red blood
cells, attracts and binds inhaled oxygen, which can then be
transported throughout the body.
Hb(aq) O2(aq)
e HbO2(aq)
Carbon monoxide, CO(g), binds more readily to hemoglobin
than oxygen and can displace oxygen according to this
equilibrium:
HbO2(aq) CO(g)
e HbCO(aq) O2 (g)
K 200 at 37°C
Consider this scenario:
A patient, unconscious due to suspected carbon monoxide
poisoning, has just been brought to the hospital emergency
ward where you are the doctor in charge. Based on your
knowledge of Le Châtelier’s principle, what treatment
would you recommend
GO
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