Solutions
– Review of Solutions
– Intermolecular Forces and Solutions
(or Like Dissolves Like)
– Factors in Solubility
– Types of Solutions
– Factors in Rate of Dissolution
What are Solutions?
• Homogeneous mixtures
• Substances are soluble or miscible
• They can be in all phases
– Solid Solutions
– Liquid Solutions
– Gas Solutions
Why do Solutions Form?
Can we predict when a
solution will form?
What will happen if…?
Solutions
• You know that oil and water don’t mix
together to form a solution.
• The burning question is WHY???
• They are both liquids, why shouldn’t they mix
together easily to form a solution?
• And why does NaCl dissolve in water but
NOT in oil?
Rehash of Intermolecular Forces
• Dipole-Dipole (as in HCl)
• Hydrogen Bonding (as in HF or H2O)
• Induced Dipole-Induced Dipole or
London Dispersion Forces (as in H2
or He)
Intermolecular Forces
• Ion-Dipole
– Soluble salt dissolves
into ions
– Ions are attracted to
partial charges
Intermolecular Forces
Intermolecular Forces and
Solutions
• You’ve learned to predict which type of force
predominates based on the polarity of the
molecule.
• You learned that the intermolecular forces are
responsible for the boiling point (and melting
point) of a substance.
• But that’s not all that intermolecular forces do!
Intermolecular Forces and
Solutions
• Intermolecular forces drive the solution
process.
• If we understand how intermolecular
forces work, we can predict whether two
substances will mix or not, AND explain
WHY!
The Solution Process
• What happens when we mix two
substances together?
• We can visualize the mixing as
occurring in three steps:
Solvation
Or Hydration
The Solution Process
• The overall three
steps.
Hsln = H1 + H2 + H3
• But only H3 is exothermic!
The Solution Process
• So the solution process can be
exothermic or endothermic.
• Exothermic processes are favored by
nature, but endothermic processes
also occur naturally.
• How can we tell whether they will
occur naturally?
• The magnitude of Hsln is crucial!
The Solution Process
• If the energy released forming the
solute-solvent intermolecular forces
is small compared to the energy
required to break the solute-solute
and solvent-solvent forces, then the
solution process is highly
endothermic.
• The two substances won’t mix!
The Solution Process
• If the energy released forming the
solute-solvent intermolecular forces
is comparable to the energy required
to break the solute-solute and
solvent-solvent forces, then the
solution process is exothermic or
slightly endothermic.
• The two substances will mix!
The Solution Process
The Solution Process
• What this also means is the strength
and thus the “type” of intermolecular
forces must be similar between the
solute and solvent.
The Solution Process
• If the solute and solvent have similar
kinds of intermolecular forces, then
-H3 is similar to H1 + H2, and they
will mix.
• If the solute and solvent have
different types of intermolecular
forces, then -H3 is much less than
H1 + H2, and they will NOT mix.
Like Dissolves Like
• This is summed up as the rule “Like
dissolves like.”
• What types of forces are similar?
Like Dissolves Like
• Ionic compounds are held together
by ion-ion attractions. Ions have full
charges.
• Polar molecules are held together by
dipole-dipole or hydrogen bonding.
The molecules have partial charges.
• Nonpolar molecules are held
together by London forces (induced
dipoles), the weakest intermolecular
force.
Like Dissolves Like
• So polar molecules tend to mix with
other polar molecules.
• Ionic compounds tend to mix with
polar compounds.
• Nonpolar molecules (or atoms) tend
to mix with other nonpolar
substances.
Like Dissolves Like
• It’s actually even more complicated.
• Here’s an example:
• Chloroform, CHCl3, is a polar
molecule, with a dipole moment of
1.04 D.
• But it is not very water soluble. (1 mL
dissolves in 200 mL water)
Like Dissolves Like
• Methyl chloride, CH3Cl, is polar with
a dipole moment of 1.9 D.
• It is considered to be slightly water
soluble (100x more than chloroform).
• Ethyl methyl ether, CH2CH3OCH3, is
polar with a dipole moment of 1.12 D.
• It is water soluble.
Like Dissolves Like
• What’s going on?
• Why aren’t the polar molecules
methyl chloride and chloroform
water soluble?
• Because “like dissolves like” can be
stated more specifically:
Like Dissolves Like
• Substances which dissolve in each
other usually have similar types of
intermolecular forces.
• This can explain why chloroform and
methyl chloride aren’t very water
soluble.
Like Dissolves Like
• Chloroform, methyl chloride, and
ethyl methyl ether have dipole-dipole
intermolecular forces.
• Water has H-bonding.
• So they don’t seem very “like” water.
• Why is ethyl methyl ether water
soluble?
Like Dissolves Like
• Ethyl methyl ether can H-bond to
water as it contains an O.
• Chloroform and methyl chloride
can’t.
• So ethyl methyl ether is more “like”
water and is water soluble.
Like Dissolves Like
• It also gets more complicated for
larger molecules with polar and
nonpolar regions.
• If the molecule has lots of polar
regions (particularly H-bonding), it is
more likely to be water soluble.
Like Dissolves Like
Like Dissolves Like
• Some molecules like soap or ethanol
are soluble in both polar and
nonpolar solvents.
• This is because they have both polar
and nonpolar regions and so can
dissolve in both types of solvents.
Like Dissolves Like
Like Dissolves Like
• Now you can answer the questions:
• Why doesn’t water (a polar molecule)
mix with oil (a nonpolar molecule)?
• Why does NaCl (an ionic compound)
dissolve in water but NOT in oil?
Other Solubility Factors
• Now you can understand and predict
whether 2 substances will mix.
• But are there any other factors in
solubility?
– Temperature
– Pressure (for gas solubility)
Temperature & Solubility
• For most substances, the solubility
increases with increasing
temperature.
• This is not true for ALL compounds.
• For gases, the reverse is true: the
solubility decreases with increasing
T.
• This is why sodas go “flat” when
open at room temperature.
Temperature & Solubility
Temperature & Solubility
Pressure & Solubility of Gases
• As gases are compressible, we can
affect the solubility by changing the
pressure.
• The higher the vapor pressure of a
gas, the more it dissolves.
Pressure & Solubility of Gases
Pressure & Solubility of Gases
• This is Henry’s Law:
s = kHP
Where s is the solubility in M, P is the
partial pressure of the gas, and kH is
Henry’s Constant.
• Soda and sparkling wine
manufacturers rely on Henry’s Law!
• Deep sea divers HATE this law!
Why?
Solution Types
• Unsaturated
• Saturated
• Supersaturated
Solution Types
• Unsaturated
– Solutions where LESS solute is
dissolved than is possible at that
temperature.
– So the solute is not at its solubility
limit (usually in M or g/mL or g/L)
– Solution is CLEAR!
Solution Types
• Saturated
– Solutions where as much solute is
dissolved as is possible at that
temperature.
– Solute is at the solubility limit
– Solution may be clear or there may
be solute at the bottom.
Solution Types
• Supersaturated
– Solutions where MORE solute is
dissolved as is possible at that
temperature.
– Solute is over the solubility limit!
– Solution is clear.
– Unstable sln; may crystallize or
“crash” easily.
Solution Types
• How do you tell them apart?
– If there is undissolved solute, it IS
saturated.
– If there is no undissolved solute,
then it could be any of the 3.
Solution Types
• How do you tell them apart?
– Add a small amount of the solute
– If it “crashes”, then it was
supersaturated.
– If the solute doesn’t dissolve, then
it was saturated.
– If the solute dissolves, then it was
unsaturated.
Solution Types
Factors in Rate of
Dissolution
• Now you know the factors which
determine whether substances will
mix.
• But what factors help determine how
FAST they will mix?
• You actually know all of these
already!
• How do YOU get sugar to dissolve in
water?
Factors in Rate of
Dissolution
• Stirring
• Temperature
• Surface Area
Solution Concentration
Units
•
•
•
•
•
•
Mass %
Volume %
Mass/Volume %
Molarity (you know)
Mole Fraction (you know for gas)
Molality (abbrev. m)