1

sp

Many important ionic compounds are only slightly soluble in water (we used to call them “insoluble” – Chapter 4)
 An equation can represent the equilibrium between the compound and the ions present in a saturated aqueous solution:

Solubility product constant, K sp
:
 K sp
for the BaSO
4
example above:
 Write the dissociation reaction for each solid: o Al(OH)
3 o Mg
3
(PO
4
)
2
2

3

Try It Out: Write a chemical equation AND the solubility constant expression representing solubility equilibrium for Ag
3
AsO
4
, K sp
=
1.0 x 10 -22 .
4

sp

K sp
is an ____________________________________
 Molar solubility is the

Molar solubility is ______________________ to the value of
K sp
, but molar solubility and K sp
are ___________ the same thing.

In fact, ____________________________________ doesn’t always mean ________________________________________
 Solubility depends on both K sp
and the form of the equilibrium constant expression
Example 16.2 At 20 °C, a saturated aqueous solution of silver carbonate contains 32 mg of Ag
2
CO
3
per liter of solution. Calculate
K sp
for Ag
2
CO
3
at 20 °C. The balanced equation is
Ag
2
CO
3
(s) ↔ 2 Ag + (aq) + CO
3
2– (aq) K sp
= ?
5

Example 16.3 From the K sp
value for silver sulfate, calculate its molar solubility at 25 °C.
Ag
2
SO
4
(s) ↔ 2 Ag + (aq) + SO
4
2– (aq) K sp
= 1.4 x 10 –5 at 25 °C
Try It Out: Calculate the K sp
value for cadmium iodate, solubility
0.097 g Cd(IO
3
)
2
/100 mL H
2
O.
6


The common ion effect affects solubility equilibria as it does other aqueous equilibria.

The solubility of a slightly soluble ionic compound is
______________________ when a second solute that furnishes a common ion is added to the solution.
Try It Out: Calculate the molar solubility of Mg(OH)
2
in 0.25 M
NaOH (aq).
7


Ions that are ____________ common to the precipitate can also affect solubility o CaF
2
is ____________ soluble in 0.010 M Na
2
SO
4
than it is in water.

Increased solubility occurs because of interionic attractions.
 Each…

The effective concentrations, or _____________________, of
Ca +2 and F -1 are lower than their actual concentrations.

Q sp
is the

Q ip
can then be compared to K sp
. o Precipitation should occur if ______________. o Precipitation cannot occur if ______________. o A solution is just saturated if ______________.
 In applying the precipitation criteria, the effect of dilution when solutions are mixed must be considered.
8


Example 16.6 If 1.00 mg of Na
2
CrO
4
is added to 225 mL of
0.00015 M AgNO
3
, will a precipitate form?
Ag
2
CrO
4
(s) ↔ 2 Ag + (aq) + CrO
4
2– (aq) K sp
= 1.1 x 10 –12

A slightly soluble solid does not precipitate _______________ from solution, but we generally consider precipitation to be
“complete” if 99.9% of the target ion is precipitated (0.1% or less is left in solution)

Three conditions generally _____________________________
__________________________ of precipitation:
1.
A _____________________ value of K sp
.
2.
A _________________________ of the target ion.
3.
A concentration of common ion that _________________
________________________that of the target ion.
9

Example 16.9 To a solution with [Ca 2+ ] = 0.0050 M, we add sufficient solid ammonium oxalate, (NH
4
)
2
C
2
O
4
(s), to make the initial [C
2
O
4
2– ] = 0.0051 M. Will precipitation of Ca 2+ as CaC
2
O
4
(s) be complete?
CaC
2
O
4
(s) ↔ Ca 2+ (aq) + C
2
O
4
2– (aq) K sp
= 2.7 x 10 –9
10

• If the anion of a precipitate is that of a weak acid, the precipitate will dissolve somewhat when the pH is lowered:
CaF
2
(s) ↔ Ca 2+ (aq) + 2 F – (aq)

If, however, the anion of the precipitate is that of a strong acid, lowering the pH will have no effect on the precipitate.
AgCl(s) ↔ Ag + (aq) + Cl – (aq)
Will pH effect the solubility of: o CaCO
3 o Mg(OH)
2 o AgCl o KNO
3
Example 16.11 What is the molar solubility of Mg(OH)
2
(s) in a buffer solution having [OH – ] = 1.0 x 10 –5 M, that is, pH = 9.00?
Mg(OH)
2
(s) ↔ Mg 2+ (aq) + 2 OH – (aq) K sp
= 1.8 x 10 –11
11

A _____________________ consists of a central metal atom or ion, with other groups called _________________bonded to it.
The __________ ion acts as a Lewis acid (accepts electron pairs).
__________________ act as Lewis bases (donate electron pairs).
The equilibrium involving a complex ion, the metal ion, and the ligands may be described through a ______________________
Examples
• A complex ion contains cobalt (III) and 6 ammonia atoms.
Write the formula for this complex ion.
• A complex ion contains zinc and 4 cyanide ions. Write the formula for this complex ion.
12

13

Example 16.14 If 1.00 g KBr is added to 1.00 L of the solution described in Example 16.13, should any AgBr(s) precipitate from the solution?
AgBr(s) ↔ Ag + (aq) + Br – (aq) K sp
= 5.0 x 10 –13
14

Example 16.15 What is the molar solubility of AgBr(s) in 3.0 M
NH
3
?
AgBr(s) + 2 NH
3
(aq) ↔ [Ag(NH
3
)
2
] + (aq) + Br – (aq)
K c
= 8.0 x 10 –6
 Certain metal hydroxides, insoluble in water, are amphoteric; they will react with both strong acids and strong bases

________________, ___________________, and
________________ are amphoteric.
15


________________________________ examines the relationship between heat and work.

Why study thermodynamics? o With a knowledge of thermodynamics and by making a few calculations before embarking on a new venture, scientists and engineers can save themselves a great deal of time, money, and frustration. o Thermodynamics tells us what processes are
_______________.
 Kinetics tells us whether the process is
_____________
nd
 Says that any spontaneous process…

More order requires a great deal of work and is often impossible to achieve o To put one system in order requires that you put others in __________________.

_______________________ is a mathematical concept describing the distribution of energy within a system.
16

o It is a measure of molecular randomness, or disorder.
 ________________________ means not requiring an outside source of energy to proceed. o It doesn’t mean that the process happens
________________. It just means that it can happen without outside energy _________________________.
o A __________________________ process is one that cannot take place in a system left to itself. o If a process is spontaneous… o Example: o Kinetics determines the ____________________ by which equilibrium is reached. o A high ____________________
______________________ can effectively block a reaction that is thermodynamically favored. o Example: o Early chemists proposed that spontaneous chemical reactions should occur in the direction of
_________________________
17

o It is true that many ____________________________ processes are spontaneous and that many
__________________________ reactions are nonspontaneous o However, enthalpy change is not a
______________________ criterion for predicting spontaneous change
 The idea that exothermic reactions are spontaneous and that endothermic reactions are not works in many cases, not all cases.
 Example:

We need to look at factors other than enthalpy (H), which is where entropy (S) comes in.

The gases in the figure on p. 19 mix because of entropy.
18


The ___________________ of the system remains unchanged in the mixing of the gases …
 … but the number of ____________________for the
______________________of that energy increases.
 The spreading of the energy among states, and increase of entropy, often correspond to a greater physical disorder at the microscopic level (however, entropy is not “disorder”).
 There are two driving forces behind spontaneous processes: the tendency to achieve a lower energy state (enthalpy change) and the tendency for energy to be distributed among states (entropy).
 In many cases, however, the two factors work in opposition.
One may increase and the other decrease or vice versa. In these cases, we must determine which factor predominates.

The difference in entropy (S) between two states is the
____________________________.
 The greater the number of configurations of the microscopic particles (atoms, ions, molecules) among the energy levels in a particular state of a system, the ___________________ is the entropy of the system.
19


Entropy generally increases when:
1. Solids melt to form liquids.
2. Solids or liquids vaporize to form gases.
3. Solids or liquids dissolve in a solvent to form nonelectrolyte solutions.
4. A chemical reaction produces an increase in the number of molecules of gases
5. A substance is heated.

Examples: o Melting Ice o Crystallization o Evaporation o Dissolving a Solid o Freezing
Example 17.2 Predict whether each of the following leads to an increase or decrease in the entropy of a system. If in doubt, explain why.
(a) The synthesis of ammonia:
N
2
(g) + 3 H
2
(g)  2 NH
3
(g)
20

(b) Preparation of a sucrose solution:
C
12
H
22
O
11
(s) → C
12
H
22
O
11
(aq)
(c) Evaporation to dryness of a solution of urea, CO(NH
2
)
2
, in water:
CO(NH
2
)
2
(aq)  CO(NH
2
)
2
(s)
Try It Out: Which reactions below would have an INCREASE in entropy?
1.
H
2
O (s)  H
2
O (g)
2.
CO
2
(g)  CO
2
(l)
3.
C
2
H
4
(g) + H
2
(g)  C
2
H
6
(g)
4.
2 KClO
3
(s)  2 KCl (s) + 3 O
2
(g)

So far, Spontaneity is more likely if o Enthalpy ( ∆ H) is ____________________________ o Entropy ( ∆ S) is __________________________
21

• According to the _______________________________, the entropy of a pure, perfect crystal can be taken to be zero at 0
K.
• The ________________________________________ is the entropy of one mole of a substance in its standard state.
• Since entropy increases with temperature, standard molar entropies are positive—even for elements.
• Entropy is a state function, so it is calculated in the same way that we calculated enthalpy first semester!
S =
v p
v r
22

Try It Out: The following reaction may occur in the high-temperature, catalyzed oxidation of NH
3
(g). Use data from Appendix C to calculate the standard molar entropy change for this reaction at 25°C.
4 NH
3
(g) + 3 O
2
(g)  2 N
2
(g) + 6 H
2
O (g)

In general, the more atoms in its molecules, o Molecules with more atoms have more ways of moving with respect to each other – more vibrational modes.
The more ways that molecules can absorb energy, the more energy levels available for dispersal of this energy, and the greater the entropy.
• Entropy can be used as a sole criterion for spontaneous change, but the entropy change of both the ____________ and its _______________________must be considered.
23

• ____________________________________ establishes that all spontaneous processes increase the entropy of the
______________________(system and surroundings).
• If entropy increases in both the _______________ and the
__________________, the process is spontaneous. o Is it possible for a spontaneous process to exhibit a
decrease in entropy?
• We do not have an effective way to measure the entropy of the universe so we have another thermodynamic function –
• Gibbs Free Energy is the amount of energy in a system that is available to do _____________________________.
• For a process at constant temperature and pressure:

_______________________ = enthalpy-driven o The flow of thermal energy provides most of the free energy for the reaction.

_______________________ = entropy-driven o Reaction disorder provides most of the free energy for the reaction.
24

o If __________(negative), a process is _________________. o If ________ (positive), a process is ___________________. o If ___________, neither the forward nor the reverse process is favored; there is no net change, and the process is at _________________________.
Example 17.4 Predict which of the four cases in Table 17.1 you expect to apply to the following reactions.
(a) C
6
H
12
O
6
(s) + 6 O
2
(g)  6 CO
2
(g) + 6 H
2
O(g) ΔH = –2540 kJ
(b) Cl
2
(g)  2 Cl(g)
25

Example 17.4 A Predict which of the four cases in Table 17.1 you expect to apply to the following reactions.
(a) N
2
(g) + 2 F
2
(g)  N
2
F
4
(g) + 6 H
2
O(g) ΔH = –7.1 kJ
(b) COCl
2
(g)  CO(g) + Cl
2
(g) ΔH = +110.4 kJ
• The _______________________________, of a reaction is the free energy change when reactants and products are in their standard states.
• The _______________________________ is the free energy change for the formation of 1 mol of a substance in its standard state from the elements in their standard states.
26

27

Use the standard-state form of the Gibbs equation to determine the standard free energy change at 25 ᵒ C for the reaction:
2 NO (g) + O
2
(g)  2 NO
2
(g) ∆ H ᵒ = -114.1 kJ ∆ S ᵒ = -146.2 J/K
Use standard free energies of formation from Appendix C to determine the standard free energy change at 25 ᵒ C for the reaction:
CS
2
(l) + 2 S
2
Cl
2
(g)  CCl
4
(l) + 6 S (s)
28

eq
• ______________ is a criterion for equilibrium at any temperature.
• ______________ is a criterion for equilibrium at a single
temperature, that temperature at which the equilibrium state has all reactants and products in their standard states.
• ΔG and ∆ G o are related through the reaction quotient, Q:
• When ∆ G = 0, then Q = K eq
, and the equation becomes:
• The concentrations and partial pressures we have used in K eq are approximations.
• ___________________ (a) are the correct variables for K eq
.
But activities are very difficult to determine. That is why we use approximations. o For pure solid and liquid phases: a = 1. o For gases: Assume ideal gas behavior, and replace the activity by the numerical value of the gas partial pressure (in atm). o For solutes in aqueous solution: Replace solute activity by the numerical value of the solute molarity (M).
29

30

Example 17.9
Determine the value of K eq
at 25 ᵒ C for the reaction:
2 NO
2
(g) ↔ N
2
O
4
(g)
31

Example 17.9 A
Use data from Appendix C to determine K eq
at 25 ᵒ C for the reaction:
2 HgO (s) ↔ 2 Hg (l) + O
2
(g)
32