Concept of solubility euilibria deals with the extent of solubility of different salts or compounds in water or other solvents and also tells about different factors which control the solubility of different salts.

1. Speaker
Shaheen Praveen
Department of Soil Science & Agricultural Chemistry
COA, IGKV
Raipur, C.G.

2. Precipitation
Precipitation - The formation of a solid from solution.
The reverse of dissolution.
Dissolution - The process by which a substance dissolves.
The reverse of precipitation.
Importance - (a) Selective precipitation is an important
industrial purification process, especially when crystals
are formed. (b) Scales that form on boilers and teeth are
to be prevented, as are kidney stones. (c) Precipitation
forms minerals - dissolution removes them.

3. Saturation
Saturated Solution - One in which a dissolution -
precipitation equilibrium exist between a solid and its
dissolved form. Here the equilibrium is dynamic and the
rate of dissolution is equal to the rate of precipitation.
Unsaturated Solution - One in which the concentration of
dissolved solid is not sufficient to cause precipitation.
Obviously a quantitative description of this type of
heterogeneous equilibrium is subject to the law of mass
action, and equilibrium expressions can be written and
deductions made concerning the concentration of various
species at equilibrium.

4. Solubility
Solubility - The greatest amount of a substance that will dissolve in
equilibrium in a specified volume of solvent at a particular
temperature.
 Molar solubility (mol/L) is the number of moles of solute dissolved
in 1 L of a saturated solution.
 Solubility (g/L) is the number of grams of solute dissolved in 1 L of
a saturated solution.
Example - The solubility of silver chloride in water at 25 o
C is .0018
g/L or 1.3 x 10-5
M.
Most solubility's increase with temperature.

5. Classification of Ionic Materials by Solubility
Soluble Ionic Materials - have solubilities in excess of 10
gL-1
.
Insoluble Ionic Materials - have solubilities less than 0.1
gL-1
.
Slightly Soluble Materials - have solubilities between 0.1
and 10 gL-1
.

6. Solubility Rules
Salts are generally more soluble in HOT water (Gases are
more soluble in COLD water).
Alkali Metal salts are very soluble in water.
NaCl, KOH, Li3PO4, Na2SO4 etc...
Ammonium salts are very soluble in water.
NH4Br, (NH4)2CO3 etc…
Salts containing the nitrate ion, NO3
-
, are very soluble in water.
Most salts of Cl-
, Br-
and I-
are very soluble in water -
exceptions are salts containing Ag+
and Pb2+
.
soluble salts: FeCl2, AlBr3, MgI2 etc...
“insoluble” salts: AgCl, PbBr2 etc...

7. Dissolving a Salt…
A salt is an ionic compound -
usually a metal cation bonded to a
non-metal anion.
The dissolving of a salt is an
example of equilibrium.
The cations and anions are attracted
to each other in the salt.
They are also attracted to the water
molecules.
The water molecules will start to
pull out some of the ions from the
salt crystal.

8. At first, the only process occurring
is the dissolving of the salt - the
dissociation of the salt into its ions.
However, soon the ions floating in
the water begin to collide with the
salt crystal and are “pulled back in”
to the salt. (precipitation)
Eventually the rate of dissociation is
equal to the rate of precipitation.
The solution is now “saturated”. It
has reached equilibrium

9. Solubility Equilibrium:
Dissociation = Precipitation
Na+
and Cl -
ions
surrounded by
water
molecules
NaCl Crystal
In a saturated solution, there
is no change in amount of
solid precipitate at the
bottom of the beaker.
Concentration of the solution
is constant.
The rate at which the salt is
dissolving into solution
equals the rate of
precipitation.

10. Nature of Ionic Equilibria
Most salts dissociate into ions when they dissolve. Equilibrium
then exists between the solid salt and its aquated ions, and not
between the solid salt and dissolved molecules of the salt. For
example:
 PbSO4(s) = Pb2+
(aq) + SO4
2-
(aq)
This equilibrium system may be described by the mass-action
expression:
Ksp = [Pb2+
][SO4
2-
]
Note that the pure solid does not enter into the equilibrium.

11. Solubility and the Ksp
• One may provide solubility information as the solubility,
S or as the solubility product, Ksp.
• These two quantities are obviously related to each other.
• Saturated solutions of salts are another type of chemical
equilibrium.
• Slightly soluble salts establish a dynamic equilibrium with
the hydrated cations and anions in solution.

12. When the solid is first added to water, no ions are initially
present.
As dissociation continues, the concentration of aqueous
ions increases until equilibrium is reached.
This process can be represented by the solubility product
constant or Ksp expression.
Ksp = [M² ][X⁺ ˉ]²
(aq)2X(aq)2M(s)
2
MX −++↔

13. Even “insoluble” salts dissociate a little – their Ksp values
range from 10-10
to 10-50
.
A Ksp value is unique to a given salt at a given
temperature.
Why would a change in temperature
alter the value of Ksp?
Solubility unit- mol/L or, g/L or, mg/L

14. Solubility indicates the amount of salt that dissociates to
form a saturated solution – think solubility curve.
In essence, it indicates the equilibrium position for a
given set of conditions.
We can have different solubilities with the same Ksp.

15. Ksp Utilization
Calculate the effect of a common ion or pH on solubility
Determine if a precipitate will form given concentrations
of ions and Ksp
Determine the order of precipitation in a mixture of ions

16. How can We Calculate Ksp
Write the dissociation equation first.
Write the Ksp expression – leaving out the solid
(Note: some of these have quite large exponents).
For a saturated solution of AgCl, the equation
would be:
AgCl (s) ↔ Ag+
(aq) + Cl-
(aq)
The solubility product expression would be:
Ksp = [Ag+
] [Cl-
]

17. For example (solubility given)-
Lead (II) chloride dissolves to a slight extent in water
according to the equation:
PbCl2 ↔ Pb+2
+ 2Cl-
Calculate the Ksp if the lead ion concentration has been
found to be 1.62 x 10-2
M
Solution:- Considering the equation, if lead’s concentration
is “x” , then chloride’s concentration is “2x”.
So. . . . Ksp = [Pb+2
] [Cl-
]2
Ksp = (1.62 x 10-2
)(3.24 x 10-2
)2
= 1.70 x 10-5

18. Calculation of Solubility
Write the dissociation equation.
Use the equation to consider the amount of ions given that
‘s’ of the solid dissociates.
Write the Ksp expression and substitute ‘s’ values and
solve (Pay attention to powers and roots).
For example:- In a saturated solution of silver carbonate,
what is the molar solubility of the salt? Ksp = 8.1 X 10-12

19. Continued…
Ag2CO3(s) ↔ 2 Ag+
+ CO3
-2
-s +2s +s
Ksp = [Ag+
]2
[CO3
-2
] = (2s)2
(s) = 4s3
Ksp = 8.1 X 10-12
= 4s3
s = 1.3 X 10-4
M

20. Comparing Solubilities
The relative solubilities can be deduced by
comparing values of Ksp.
These comparisons can only be made for salts
having the same ION:ION ratio.

21. Which Salt is More Soluble?
Ag2S Ksp = 1.0 X 10-49
Ni(CN)2 Ksp = 3.0 X 10-23
Ag2S ↔ 2 Ag+
+ S-2
Ni(CN)2 ↔ Ni+2
+ 2 CN-
* Since both make 3 ions (4s3
) – the larger Ksp is
the more soluble salt - Ni(CN)2

22. Common Ion Effect
The presence of a common ion in a solution will lower the
solubility of a salt.
LeChatelier’s Principle:
The addition of the common ion will shift the
solubility equilibrium backwards. This means that
there is more solid salt in the solution and therefore
the solubility is lower.

23. Effect of a Common Ion on pH or
Solubility
The pH of a solution can also affect solubility if H+
or OH-
can
interact with the salt’s ions.
Q.) How will the solubility of calcium carbonate be affected
if it is dissolved in a solution of calcium chloride?
presence of Ca2+
ions will shift equilibrium position to left –
CaCO3 will be less soluble
Q.) How will the solubility of silver phosphate be affected
by an decrease in pH?
H+
react with PO4
3-
thus causing a shift right and increasing
solubility

24. Q.) Would magnesium hydroxide (milk of magnesia) be more
soluble in an acid or a base? Why?
Mg(OH)2(s) ↔ Mg2+
(aq) + 2 OH-
(aq)
How will the solubility of magnesium hydroxide be affected by
an increase in pH?
more OH-
will cause shift left – decrease sol.
How will the solubility of magnesium hydroxide be affected by
an decrease in pH?
H+
will react with OH-
cause shift right – increase sol.

25. The Common Ion Effect on Solubility
The solubility of MgF2 in pure water is 2.6 x 10-4
mol/L.
What happens to the solubility if we dissolve the MgF2 in
a solution of NaF, instead of pure water?

26. Q.) Calculate the solubility of MgF2 in a solution of 0.080 M
NaF.
MgF2 (s)  Mg2+
(aq) + 2 F-
(aq)
---- 0.080 M
+ s + 2s
s 0.080 + 2s
Ksp = 7.4 x 10-11
= [Mg2+
][F-
]2
= (s)(0.080 + 2s)2
Since Ksp is so small…assume that 2s << 0.080
7.4 x 10-11
= (s)(0.080)2
s = 1.2 x 10-8
mol/L

27. Solubility and pH
Q.) Calculate the pH of a saturated solution of silver
hydroxide, AgOH. Refer to the table in your booklet for
the Ksp of AgOH.
AgOH (s)  Ag+
(aq) + OH-
(aq)
---- ----
+ s + s
s s
Ksp = 2.0 x 10-8
= [Ag+
][OH-
] = s2
s = 1.4 x 10-4
M = [OH-
]
pOH = - log (1.4 x 10-4
) = 3.85
pH = 14.00 - pOH = 10.15

28. Determine if a Precipitate will Form
With some knowledge of the reaction quotient
(Q), we can decide
1) whether a ppt will form, and
2) what concentrations of ions
are required to begin the ppt. of
an insoluble salt.

29. 1. Q = Ksp, the system is at
equil. (saturated)
2. Q < Ksp, the system is not at
equil. (unsaturated – shift
right)
3. Q > Ksp, the system is not at
equil. (supersaturated – shift
left)
*Precipitates form when
the solution is
supersaturated.

30. Quantitative Analysis of Cations