Solubility Rules
& Reference Tables
Components of a Solution
Solute: substance
being dissolved
Ex: Salt, Sugar
Solvent: substance
doing the dissolving
Ex: Water, Hexane
Solubility:
How much solute can
dissolve under certain
conditions of
temp. and pressure.
Factors Affecting
Solubility
Surface Area
More solute/solvent contact
means faster dissolving
Crush substance into fine
powder
Use mortar and pestle
Stirring or Agitation:
More solute/solvent
contact (solids/liquids)
However, stirring disturbs
dissolved gases and they
come out of solution.
Temperature of Solvent
Higher temperatures will allow more
solid solutes to dissolve
Gases dissolve better when
solvent temperature is colder.
Ex:
CO2 gas in hot soda (flat)
vs. cold soda (fizzy)
Pressure
Effects gas solubility only
Why?
Increasing pressure on
a gas above a liquid
causes more gas
molecules to be
“pushed” into solution.
Ex: CO2(aq) in soda
Nature of Solute and Solvent
 Polar solutes dissolve in polar solvents

Nonpolar solutes dissolve in nonpolar
solvents

Most ionics (but not all) dissolve in polar
solvents (molecule-ion attractions)
Amount of Solute already Dissolved
As particles dissolve in solution fewer
solvent molecules are available to dissolve
new solute.

Miscible:
2 liquids that dissolve
(ex: alcohol and water)

Immiscible:
2 liquids that do not
dissolve (ex: oil and water)
Electrolytes:
Conduct electricity when dissolved in water
Why do they Conduct?
Create mobile ions in solution.
 The more concentrated the solution the
more it conducts

Includes:
Soluble Ionic Compounds (ex: NaCl)
 Acids (ex: HCl)
 Bases (ex: NaOH)

Who Will Conduct?








Which of the following compounds will conduct
in solution? (ionic salt, acid, base?)
See Ref Tables for common acids/bases
C6H12O6
LiBr
KOH
CH4
H2SO4
NO2
C6H12O6
 LiBr
 KOH
 CH4
 H2SO4
 NO2

Will Not (Covalent)
Will (Ionic)
Will (Base)
Will Not (Covalent)
Will (Acid)
Will Not (Covalent)
Using Reference Table G

Shows solubility in
grams of solute per
100 grams of water at
different temps
Saturated Solutions:
hold max solute
possible at that temp.
Table G:
Solubility curve lines
show saturation levels
at different temps

Saturated Solutions are at EQUILIBRIUM.
Rate of dissolving = Rate of crystallization

Ex: How many grams
of NaNO3 are needed
to create a sat.
solution in 100g of
water at 50 °C?

Go to 50 °C and up to
NaNO3 and over.

Answer: 116 grams
Look at The Water!!

Table G is for 100 grams of water.

Amount of water in your problem may be
different and you need to adjust your
answer.
How many grams of NaNO3 are needed to
create a sat. solution in 300g of water at 50 °C?

Answer: 116 grams x 3 (three times as much
water!)

Or you can use a proportion:
116 grams
100 g H20
=
x grams
300g H20
Unsaturated Solutions
could still hold more
solute at that temp.
Would fall “below the line”
on Table G
Ex: 40 g of NaNO3 in
100g water at 50°
Supersaturated Solutions
hold more solute than they
should at that temp.
Would fall “above the line”
on Table G
Ex: 140 g of NaNO3 in
100g water at 50°
How do Supersaturated Solutions
Form?

Create a saturated solution
at a high temp. and slowly
let solution cool.

Certain solutes can stay in
solution.


Ex: sodium acetate
Supersaturated solutions are
unstable.

Add just one more “seed
crystal”, all excess solute will
precipitate leaving a saturated
solution behind
Supersaturated Sodium Acetate
solution after seed crystal added
Describe These Solutions
Saturated, Unsaturated or Supersaturated?
100 g NH4Cl at 70° in
100 g water
Falls above the line
(Supersaturated)
10g SO2 at 10° in 100g
water
Falls below the line
(Unsaturated)
40g NaCl at 90° in 100g
water
Falls on the line
(Saturated)
Concentrated Solutions:
have a lot of solute
dissolved in the solvent
Ex: Saturated solution of KI
at 10°
135 grams in 100 g water
= pretty concentrated
Dilute Solutions: only
have a little solute
dissolved.
Ex: Sat. solution of SO2
at 50°
4 grams in 100 g water
= relatively dilute
If Temp. Changes
 How much will
precipitate out of
solution if a
saturated NaNO3
solution at 60° is
cooled to 20° ?
Reference Table F

Describes which ionic compounds are
soluble or insoluble in water.

Certain combinations of ions hold together
so strongly that water cannot dissolve
them into solution (insoluble)

Is this soluble or not?
CaCO3

Carbonate (CO3-2) is insoluble and Ca+2 as
a partner is not an exception

Is this soluble or not?
NaNO3

Nitrate (NO3-1) is always soluble, there are
not exceptions

Is this soluble or not?
Li3PO4

Phosphate (PO4-3) is insoluble, however,
Li+1 is a Group 1 ion so it is an exception
and the compound is soluble.
Soluble or Not?
Look out for exceptions!
 CaSO4
 MgSO4
 PbCrO4
 Li2S
 NH4OH
Insoluble
 Soluble
 Insoluble
 Soluble
 Soluble

CaSO4
MgSO4
PbCrO4
Li2S
NH4OH
Precipitates

Precipitates are insoluble
ionic compounds formed in
double replacement reactions.

Determine which product is the
insoluble precipitate by using
Table F.

When a precipitate forms,
you create a heterogeneous
mixture.

You can separate a
precipitate by filtration.

The solid will stay on the
paper.