2. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
• Typically when we think of what happens
Equilibrium during a chemical reaction we think of the
Introduction reactants getting totally used up so that
none are left and ending up with only
products. Also, we generally consider
chemical reactions as one-way events.
You may well have learned during earlier
science classes that this is one way to
distinguish chemical change from physical
changes - physical changes (such as the
melting and freezing of ice) are easily
reversed, but chemical changes cannot be
reversed (pretty tough to un-fry an egg).
• In this unit we will see that this isn't always
the case. We will see that many chemical
reactions are, in fact, reversible under the
right conditions. And because many
reactions can be reversed, our idea of a
reaction ending with no reactants left, only
products, will need to be modified.

3. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
• The amount of water on the earth
About (including the atmosphere) stays
Chemical relatively constant. Water evaporates
Equilibrium or sublimes into the gas phase. Water
in the gas phase returns to the earth
in the form of precipitation and dew.
There is a balance struck between
the various phases of water, solid,
liquid and gas. This is known as the
water cycle.
• A basketball team plays a game.
During the game, players enter and
leave the game. The number of
players on the floor never changes.
• What do these have in common?

4. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
Chemical
Equilibrium Equilibrium is the state
at which the rate of the
forward reaction equals
the rate of the reverse
reaction. At the point
of equilibrium, no more
measurable or
observable changes in
the system can be
noted.

5. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
• Equilibrium occurs when the rate
Review of the forward process is equal to
the rate of the reverse process in
a reversible reaction.
• Equilibrium can only occur in a
closed system.
• Equilibrium is dynamic.
• Liquid-vapour equilibrium occurs
when liquid molecules enter and
leave the liquid state at the same
rate.
• Solution equilibria occur in
saturated solutions in which solute
remains undissolved.
• Physical equilibria are those
occurring in physical changes.

6. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
Equilibrium

7. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
Reversible • In our study of physical
Reactions processes, equilibrium could
only be established in a closed
system. This also holds for
chemical reactions in
equilibrium.
• Let's take a look at the
conversion of nitrogen dioxide
gas, NO2, into dinitrogen
tetraoxide, N2O4 by the reaction:
2 NO2(g) ↔ N2O4(g)

8. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
• This reaction is then written as
Characteristics follows:
of Chemical
Equilibrium 2 NO2(g) ↔ N2O4(g)
• The double arrow indicates the
reaction is reversible.
• In reversible reactions:
–conversion to products is the
forward reaction
–conversion to reactants is the
reverse reaction
• For a system at equilibrium, the rate
of the forward reaction equals the
rate of the reverse reaction.

9. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
• For the reaction
Defining aA + bB ↔ cC + dD
Chemical
Equilibrium • at equilibrium, the graph of
concentration vs time may appear as
below:

10. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
• OR at equilibrium, the graph of
Defining concentration vs time may appear as
Chemical below:
Equilibrium
• the concentrations of reactants
and products remain constant over
time because the rate of the
forward reaction is equal to the
rate of the reverse reaction.

11. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
Equilibrium • In 1864, Cato Guldberg and
Law Peter Waage proposed the law
of mass action or the
equilibrium law. The studied
many systems at equilibrium
and found there was a
relationship between the
concentration of reactants and
products at equilibrium. They
suggested the equilibrium law
be a ratio of product
concentrations to reactant
concentrations. The value of this
ratio is called the equilibrium
constant.

12. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
• They proposed, for the reaction
Equilibrium aA + bB ↔ cC + dD
Law
• the forward and reverse reactions
were elementary reactions. This
means
rateforward = kf[C]c[D]d
ratereverse = kr[A]a[B]b
since
rateforward = ratereverse
kr[A]a[B]b = kf[C]c[D]d

13. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
• By rearranging the expression to
Equilibrium solve for rate constants,
Law
• The ratio of rate constants was
condensed to one constant, Kc, called
the equilibrium constant. The law of
mass action or equilibrium law then
becomes

14. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
Using the Example 1
For the reaction
Equilibrium
Law H2(g) + F2(g) ↔ 2HF(g)
1.00 moles of hydrogen and 1.00
moles of fluorine are sealed in a 1.00 L
flask at 150°C and allowed to react. At
equilibrium, 1.32 moles of HF are
present. Calculate the equilibrium
constant.

15. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
Using the Example 1
For the reaction
Equilibrium
Law H2(g) + F2(g) ↔ 2HF(g)
H2 + F2 2 HF
I 1.00 1.00 0 Insert initial concentrations
Since the equilibrium [HF] is 1.32 mol/L,
it increases by that amount. However,
C + 1.32 according to the stoichiometry [H2] and
[F2] will decrease by one half that
amount (see below)
E

16. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
Using the Example 2
For the reaction
Equilibrium
Law N2(g) + O2(g) ↔ 2 NO(g)
the equilibrium constant is 6.76. If 6.0
moles of nitrogen and oxygen gases
are placed in a 1.0 L container, what
are the concentrations of all reactants
and products at equilibrium?

17. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
Using the Example 2
For the reaction
Equilibrium
Law N2(g) + O2(g) ↔ 2 NO(g)
N2 + O2 2 NO
I 6.0 6.0 0 Insert initial concentrations
Since we do not know the equilibrium
concentrations of any of the species we
C insert x for the amount of N2 and O2
consumed and 2x for the amount of NO
produced.
E

18. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
• When reactants and products are
Reaction added into a container it is good to
Quotient, Q know whether equilibrium has been
reached. If equilibrium has not been
reached it is helpful to know which
reaction, forward or reverse, is
favoured in order for equilibrium to be
achieved.
• The reaction quotient , Q, or trial Kc
enables us to determine this
information. The reaction quotient is
determined by using the equilibrium
law and using either initial
concentrations or those determined
during experimental trials.

19. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
• To determine which reaction is
Reaction favoured and in which direction the
Quotient, Q system is moving, Q is compared to
Kc.
• If Q = K, the system is at equilibrium.
The forward and reverse rates are
equal and the reactant and product
concentrations remain constant.
• If Q > K, the system is NOT at
equilibrium. There is too much
product, so the reverse reaction is
favoured to bring the reactant-product
ratio to equal K by increasing reactant
concentration.

20. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
• If Q < K, the system is NOT at
Reaction equilibrium. The concentration of
Quotient, Q reactants is too large, so the forward
reaction is favoured. This results in
decreasing reactant concentrations
and increasing product
concentrations, bringing their ratio to
a value equal to K.

21. Outcome 4-01 Relate the concept of equilibrium to physical and chemical systems.
Include: conditions necessary to achieve equilibrium.
Using the Example 3
For the reaction
Equilibrium
Law N2(g) + O2(g) ↔ 2 NO(g)
It was found that 8.50 moles of
nitrogen, 11.0 moles of oxygen and
2.20 moles of nitrogen monoxide were
in a 5.00 L container. If the equilibrium
constant is 0.035, are the following
concentrations at equilibrium? If not,
which reaction is favored and which
concentrations are increasing and
which are decreasing?

22. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
Le
Chatelier’s
Principle

23. Le
Chatelier’s
Principle

24. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
• Henri Louis Le Chatelier (1850 -
Le 1936) was a French chemist and
Chatelier’s a mining engineer. He spent much
Principle of his time studying flames in
order to prevent mine explosions.
He also invented two new ways of
measuring very high
temperatures.
• In 1884 Le Chatelier proposed the
Law of Mobile Equilibrium, more
commonly called Le Chatelier's
Principle. The principle states:
When a system at equilibrium is
subjected to a stress, the system
will adjust so as to relieve the
stress.

25. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
Le
Chatelier’s
Principle

26. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
• Le Chatelier's Principle. The
Le principle states:
Chatelier’s
Principle When a system at equilibrium
is subjected to a stress, the
system will adjust so as to
relieve the stress.
• Stresses include:
–Changing concentration
–Changing pressure
–Changing temperature
–Adding a catalyst

27. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
• Le Chatelier's Principle. The
Le principle states:
Chatelier’s
Principle When a system at equilibrium
is subjected to a stress, the
system will adjust so as to
relieve the stress.
• Stresses include:
–Changing concentration
–Changing pressure
–Changing temperature
–Adding a catalyst

28. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
• In a system at equilibrium, a
change in the concentration of
Changing products or reactants present at
Concentration equilibrium, constitutes a stress.
• At equilibrium, the ratio of product
to reactant concentrations is
constant.
• Adding more reactant, or
removing product, upsets the
established equilibrium.
• The stress is relieved by forming
more product, or using up
reactant.

29. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
• If reactant is added or product is
Changing removed, we say that the
Concentration equilibrium "shifts to the right".
• The forward reaction rate
increases until equilibrium is
reestablished. That is, the forward
reaction is favoured until Kc is
attained again.
• Similarly, adding more product, or
removing reactant, causes the
system to shift the equilibrium left.
The reverse rate is favoured until
the product to reactant ratio is
equal to Kc once again.

30. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
Le
Chatelier’s
Principle

31. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
• The pressure of a system can be
Changing changed by increasing or
Pressure reducing the volume of the
reaction container. Increasing the
size of the container reduces the
pressure, while decreasing the
size of the container increases the
pressure of the system. Changing
the pressure of a system only
affects those equilibria with
gaseous reactants and/or
products.

32. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
• According to Le Chatelier's
Changing Principle, increasing the pressure
Pressure on a system at equilibrium causes
the system to shift to reduce its
pressure, by reducing the number
of particles in the system. That is,
shift to the side with fewer
molecules.
• Conversely, decreasing the
pressure on a system causes the
system to shift to increase the
pressure by increasing the
number of particles in the
container. That is, shift to the side
with more molecules.

33. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
Le
Chatelier’s
Principle

34. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
• Recall from Kinetics that
Changing increasing temperature always
Temperature increases the rate of a reaction.
However, increasing temperature
always increases the rate of an
endothermic reaction more than
the rate of an exothermic reaction.
• According to Le Chatelier's
Principle, a change in temperature
causes a stress on a system at
equilibrium. The system attempts
to relieve the stress by either
replacing lost heat or consuming
added heat.

35. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
• To solve equilibrium problems
Changing involving heat changes, consider
Temperature heat to be a product (exothermic
reactions) or a reactant
(endothermic reactions).

36. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
• Recall that adding a catalyst to a
Adding a system decreases the activation
energy of a reaction. This will cause
Catalyst the rate of a reaction to increase.
However, a catalyst lowers the
activation energy of BOTH forward
and reverse reactions equally.
• Therefore, adding a catalyst to a
system at equilibrium will NOT affect
the equilibrium position. However, if
a catalyst is added to a system
which is not at equilibrium, the
system will reach equilibrium much
quicker since forward and reverse
reaction rates are increased.

37. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
• The graphs we will be
Analyzing studying illustrate the rate
Experimental changes for the NO2 - N2O4
Data equilibrium system.
• We will assume that the
system was at equilibrium
initially before the stress was
applied, and after an
instantaneous change, the
system was allowed to
reestablish equilibrium. The
reaction is
2 NO2(g) ↔ N2O4(g)
ΔH = -58.0 kJ/mol N2O4

38. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
Rate vs Time Graphs For
Changing Reactant Concentration
If we add more NO2 to the If we remove some NO2
system the following from the system the
graph results. following graph results.

39. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
Rate vs Time Graphs For
Changing Product Concentration
This time see what the This time see what the
graphs look like when graphs look like when
product is added. product is removed.

40. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
Rate vs Time Graphs For
Changing Product Concentration
This time see what the This time see what the
graphs look like when graphs look like when
product is added. product is removed.

41. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
Rate vs Time Graphs For Temperature Changes
This time see what the This time see what the
graphs look like there is an graphs look like heat is
increase in temperature. removed.

42. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
Concentration vs Time Graphs For Changing Reactant Concentrations
The first graph shows the The next graph shows the
effect of adding more removal of the reactant.
reactant.

43. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
Concentration vs Time Graphs For Changing Product Concentrations
This time see what the This time see what the
graphs look like when graphs look like when
product is added. product is removed.

44. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.
Concentration vs Time Graphs for Temperature Changes
This time see what the This time see what the
graphs look like when graphs look like when
heat is added. heat is removed.

45. Outcome 4-06 Use Le Chatelier’s principle to predict and explain shifts in equilibrium.
Include: temperature changes, pressure/volume changes, changes in reactant/product concentration, the addition of
a catalyst, the addition of an inert gas, the effects of various stresses on the equilibrium constant.

46. Outcome 4-10 Write solubility product (Ksp) expressions from balanced chemical
equations for salts with low solubility.
• When a solution is saturated,
Solubility there exists an equilibrium
Equilibria between the dissolved solute
particles and the solid solute
particles.
• For an ionic compound, such as
sodium chloride, we express the
equilibrium in terms of a chemical
equation:
NaCl(s) ↔ Na+(aq) + Cl¯(aq)
• This equilibrium is dynamic, since
the rate of dissolving of each ion
is equal to the crystallization of
each ion.

47. Outcome 4-10 Write solubility product (Ksp) expressions from balanced chemical
equations for salts with low solubility.
• Substances chemists believed to be
The Solubility insoluble in water, by experiment,
were shown to be slightly soluble.
Product That is, with sensitive instruments,
chemists we able to show that
substances that would not dissolve
in water actually do dissolve to a
very small extent.
• When a sparingly soluble ionic solid
is dissolved in water to form a
saturated solution the general
equilibrium equation is
AaBb(s) ↔ aA+(aq) + bB¯(aq)
• Where A is a positively charged ion
and B is a negatively charged ion.

48. Outcome 4-10 Write solubility product (Ksp) expressions from balanced chemical
equations for salts with low solubility.
• At a given temperature, the equilibrium
law for this reaction is given as
The Solubility
Product
• However, the term KC[AaBb] can be
replaced by a new constant, Ksp, called
the solubility product.
Ksp = [A+]a[B¯]b
• The solubility product constant is the
product of ion concentrations in a
saturated solution. The solubility
product constant takes into account the
presence of the solid.

49. Outcome 4-10 Write solubility product (Ksp) expressions from balanced chemical
equations for salts with low solubility.
• Recall that the general form of the
solubility product expression is
The Solubility
Product Ksp = [A+]a[B¯]b
Expression
• for the dissociation equation
AaBb(s) ↔ aA+(aq) + bB¯(aq)
Example 1
• Write the dissociation equation
and the expression for the
solubility product constant for
calcium hydroxide.

50. Outcome 4-10 Write solubility product (Ksp) expressions from balanced chemical
equations for salts with low solubility.
Example 2
Calculating • Write the solubility product
Solubility expression for Pb3(PO4)2.
Product
Example 3
• If at equilibrium, the concentration
of silver ions is 1.3 x 10-5 mol/L
and the concentration of chloride
ions is 1.3 x 10-5 mol/L, what is the
Ksp of silver chloride?

51. Outcome 4-10 Write solubility product (Ksp) expressions from balanced chemical
equations for salts with low solubility.
• Solubility and solubility product are two
Solubility different terms.
• Solubility is the maximum amount of
solute that can dissolve in a certain
amount of solvent at a certain
temperature. Solubility has an infinite
number of possible values, depending
on temperature and other solutes
present.
• Solubility product is an equilibrium
constant and has only one value for a
given solid at a particular temperature.
Example 4
• The solubility of PbF2 is 0.466 g/L. What
is the value of the solubility product?

52. Outcome 4-10 Write solubility product (Ksp) expressions from balanced chemical
equations for salts with low solubility.
Example 5
Determining • The Ksp of magnesium hydroxide
Ion
Concentration is 8.9 x 10-12. What will be the
from Ksp equilibrium concentrations of the
dissolved ions in a saturated
solution of Mg(OH)2?

53. Outcome 4-11 Solve problems involving Ksp.
Include: common ion problems.
Common Ions
• When an ionic compound
dissolves in pure water, the
initial concentration of each
ion is zero. However, if an
ionic compound dissolves in a
solution that has an ion in
common with the compound,
this is not the case.
• Even though the starting
concentrations may not be
zero, the product of the ions
must still equal the solubility
product constant.

54. Outcome 4-11 Solve problems involving Ksp.
Include: common ion problems.
• For example, how would the solubility of
Common Ions silver chloride in pure water change if
we try dissolving it in tap water?
• Let's write out the dissociation equation:
AgCl(s) Ag+(aq) + Cl¯(aq)
• Tap water often has chlorine added to
kill bacteria. The chlorine exists as
chloride ions, so when we dissolve
silver chloride in tap water, chloride ions
are present.
• According to Le Chatelier's Principle,
Adding more chloride ions to a
saturated solution would cause the
equilibrium to shift to the left to use up
the excess product. This would result in
more solid formed, and a decreased
solubility.

55. Outcome 4-11 Solve problems involving Ksp.
Include: common ion problems.
Solubility in • Le Chatelier's Principle predicts
the Presence that the solubility of an ionic solid
of a Common in a solution containing a common
Ion ion decreases its solubility. Let's
see if this is supported by the
calculations.
Example 1
• Determine the solubility of silver
chloride in pure water and in a
solution of 0.10 mol/L sodium
chloride. The Ksp of AgCl is 1.7 x
10-10.

56. Outcome 4-11 Solve problems involving Ksp.
Include: common ion problems.
Solubility in Example 2
the Presence • The Ksp of lead (II) chloride, PbCl2,
of a Common is 1.6 x 10-5. What is the solubility
Ion of lead (II) chloride in a 0.10 mol/L
solution of magnesium chloride,
MgCl2?
Example 3
• The Ksp of lead (II) chloride is 1.6
x 10-5. What is the solubility of
lead (II) chloride in a 0.10 mol/L
solution of lead (II) nitrate,
Pb(NO3)2?