Ask a chemist, they always have
Solutions
Solutes and Solvents
 Solution:
a homogenous mixture
 Solute: thing that dissolves
 Solvent: thing that does the dissolving
(found in the largest amounts)

If the solvent is water, then it is called an
aqueous solution
Solubility: Ionic Compounds
 Ions
form, separate (dissociate) and move
throughout the solution

The forces that hold the ions together are
overcome by the ions’ attractions to polar
water.
• Ion- dipole interaction
 Because
ions are present, ionic solutions
can conduct a current

Current is just movement of electrons
Figure 15.1: Dissolving of solid
sodium chloride.
Solvation animation
Animation with Audio
Figure 15.2: Polar water molecules interacting with positive
and negative ions of a salt.
Solubility: Polar Compounds
 “Like
dissolves like”
 Typically, hydrogen bonding occurs
between the substance being dissolved
and the polar water molecules
 Example:


Sugar in water
Ethanol in water
Figure 15.3: The polar water molecule interacts strongly with
.
the polar O—H bond in ethanol
Figure 15.4: Structure of
common table sugar.
Get interactions between water molecules the polar
regions on the sugar (the Os) , and some hydrogen
bonding at the -OH groups
“They go together like oil and water.”
(things that don’t dissolve or mix)

Anything nonpolar will
not mix well with
anything polar

Examples:
• Oil spill
• Salad dressing

Can mix when shaken
(LDF) and then may
separate out (other
forces)
How much is too much?
 There
is a limit to the amount of a
substance dissolved
 Saturated: the solution holds as much
solute as possible at that temperature.

Unsaturated: solution has not reached the
limit
 Can



you have too much? YES!
Supersaturated: have as much solute
dissolved as possible, then cooled and all the
solute stays dissolved.
In other words…the solution contains more
dissolved solid than a saturated solution
created at the same temperature.
These can begin crystallization of the solute
at the slightest change
Energy changes and Solvation
 Any
chemical change (including solvation)
requires a change in energy


Energy removed from or added to the
reactants from the surroundings
NaOH(s) Na+ (aq)+ OH- (aq) ΔH= -44.5 kJ/mol
(that’s 44.5kJ released, so exothermic, per mole of NaOH)
• Because you are breaking the ionic bond, energy must be
either released when breaking the bond, or consumed when
making the new ions


ALL changes in formula indicate a change in energy.
However, sometimes the energy change is so small, you can’t
tell that a change has occurred
Molarity (M)

Most common way to express concentration
 Molarity is the number of moles of solute
dissolved in each liter of solution
 Formula

M = moles of solute
liters of solution

Dependent on temperature
 The higher the molarity the stronger the
concentration
Molality (M )
 Another
way to calculate concentration
 Formula

M=
moles solute
.
kilograms of solvent
 Not
dependent on temperature
 The higher the molality the stronger the
concentration
Normality (N)


3rd way to calculate concentration
Typically used with acids and bases


Formula


Indicates amount of H+ and OH- available for
acid/base reactions
N=
equivalents
.
1 liter of solution
Equivalent weight: the mass in grams of
acid/base that gives 1 mole of H+/OH The higher the normality the stronger the
concentration
Colligative properties
 In
the winter, why do we throw salt
when it snows?
 Why
does Emeril add salt to boiling
water when cooking pasta?
 It
has to do with Intermolecular
Forces (IMFs)
Intramolecular forces (bonds)
 Hold
atoms together in molecules
 Have high energy associated with them

it’s difficult to break molecules into their
individual atoms
 Different
types based upon what is going
on with the electrons (electron clouds)


Ionic: electrons are transferred between atoms
Covalent
• electrons are shared between atoms
• Can be polar or nonpolar
Intermolecular forces (IMFs)

Hold molecules together
 Much weaker than intramolecular forces

Intramolecular bonds are usually 100x or even 1000x
stronger
*(kJ are units of energy like Calories; 1Cal= 4.184kJ)


1000cal= 1Cal
1cal =4.184J
Figure 14.2: Intermolecular forces exist between
molecules. Bonds exist within molecules.
Why do we care?
 The
strength of the IMFs determine
the state of matter


Solid, liquid, or gas*
*Not plasma- intramolecular bonds are broken to get
plasmas
Solids, Liquids, and Gases
shape
Gas
Liquid
Solid
indefinite
indefinite
definite
volume
variable
with
P and T
constant
**
constant
**
density
variable
with
volume
change
constant
**
constant
**
motion with
Energy*
level of
organization*
strength
of IMFs*
high
high; molecules
freely moving
with great
distance
compared to
molecular size
between them
very low
low
moderate
high; molecules
freely moving
past each other
but in close
proximity to each
other
low
moderate
low
low; vibration only
as molecules are
basically fixed in
place
high
high
energy*
*all at room temperature, ~25C
**small variations occur due to temperature changes, very little variable with pressure changes

Things with strong IMFs tend to be solids at
room temperature
 Things with weak IMFs tend to be gases at room
temperature
 Medium IMFs tend to be in between

liquids, yes, but with varying characteristics
Amorphous solids: long transition between solid
and liquid states- gets soft, then melts (like wax)
 Crystalline solids: definite, clear melting point
(no soft transition- ie: ice)
Energy Changes Accompanying
Changes of State

Think back: Each change of state is accompanied by a
change in the energy of the system

Whenever the change involves the disruption of intermolecular
forces, energy must be supplied
Energy Changes Accompanying
Changes of State

The disruption of intermolecular forces accompanies the
state going towards a less ordered state (higher
entropy)

As the strengths of the intermolecular forces increase, greater
amounts of energy are required to overcome them during a
change in state
• Takes more energy to go from



a liquid to a gas
• than
from a solid to a liquid
Removing energy allows the molecules to “selforganize”, and results in an more ordered state

Lower entropy
Heat of Fusion
 The
melting process for a solid is also
referred to as fusion

The enthalpy change associated with
melting a solid is often called the heat of
fusion (Δ Hfus)
• Ice ΔHfus = 6.01 kJ/mol

Δ H is a change (Δ) in enthalpy (H), a measure of energy that
is much like heat, but takes into account a few other factors
Heat of Vaporization
 The
heat needed for the vaporization of
a liquid is called the heat of
vaporization (Δ Hvap)
• Water Δ Hvap = 40.67 kJ/mol
 Vaporization
energy
requires the input of heat
 Less
energy is needed to allow
molecules to move past each other
than to separate them totally,

so ΔHfus < Δ Hvap
The heating/cooling curve for water heated or cooled at a constant rate.
 Think
of IMFs like magnets: stronger
magnets hold things more firmly
together


The more firm the connections, the less
molecular motion can occur with the
same amount of Energy added
Adding (or removing) energy from the
system can overcome (or increase) the
IMFs, and cause a change in state
• Add Energy, move from S -> L -> G
• Remove Energy, move from G -> L -> S
Vaporization and Vapor Pressure
 The
molecules in a sample of a liquid
move at various speeds

(average speed is the temperature; some
have more energy, some have less, but the
overall KE is temperature)
 Sometimes
molecules at the surface have
sufficient speed to overcome the attractive
forces and leave the liquid surface (thus
vaporizing)
Figure 14.9: Microscopic view
of a liquid near its surface.
Dynamic equilibrium
 Dynamic
equilibrium is the state
where there is simultaneous
and equal vaporization and
condensation of the substance
 In a closed container, at some
pressure, the amount that
vaporizes will equal the amount
condensing on the surface of
the liquid

This is the equilibrium vapor
pressure
VP and IMFs

Stronger IMFs equal lower vapor
pressures


Less likely to evaporate
Weaker IMFs equal higher vapor
pressures

Substance with very low IMFs and
therefore high vapor pressures evaporate
very quickly and easily
• Called volatile substance

Mass and shape important, just like with
boiling point

Heavier = lower VP
• ex: oil

Lighter= higher VP
• ex: alcohol
• More volatile

Think propane (C3H8) v. gasoline (C8H18)
VP and Boiling
 Vaporization
occurs at any temperature, but
occurs more rapidly as temperature
increases


Molecules at the surface would have to have
more speed to overcome the IMFs
Boiling is the point at which the vapor pressure
equals the external pressure on the surface of
the liquid
• Molecules are able to “escape” liquid phase b/c they
have enough Energy to break the IMFs

Convert PE of IMFs to KE of motion in a gas
Boiling and VP, con’t



Liquids have some air dissolved in them in tiny invisible bubbles
As water vaporizes in the liquid, it is added to the bubbles
Also, the gas bubbles are expanding because they are being
heated; this causes an increase in volume, but not mass

At this point, 2 things are going on:
• This decreases density, causing the bubbles to float to the surface
• Also, as gas expands, the pressure increases


When the pressure of the bubble increases to greater than the
vapor pressure at the surface, the liquid is boiling
All molecules must be vaporized before a further increase in
temperature can occur

Need to break all IMFs (convert all PE of IMFs before increasing KE of
molecules)
Figure 15.10: Pure water.
Boiling Point and Elevation

As elevation on the Earth’s surface increases,
the atmospheric pressure decreases

(smaller column of air pushing down on the area;
therefore less pressure)

Boiling point changes as the atmospheric
pressure changes
 If you could decrease the pressure without
changing temperature, the substance would boil
at a lower temperature

A decrease in pressure results in a decrease in BP
Figure 14.14: The formation of the bubble is
opposed by atmospheric pressure.
Boiling point elevation

By adding salt (or other compounds) to water,
the temperature of boiling goes up it boils at
a higher temperature
• Interrupts H bonding
• Need more vapor molecules and greater pressure
to get bubbles to form
• Takes more time to get vapors to add to bubbles
• The molecules that do get into the bubbles need
more energy

Dependent on how much solute is added
Figure 15.9: A bubble in the interior of liquid water surrounded by
solute particles and water molecules.
Figure 15.10: Solution (contains solute).
Vapor Pressure Reduction
 Vapor
pressure changes as IMFs change
 For the same reasons boiling point is
disturbed
 What would evaporate faster:


Salt water
Distilled water

WHY?
Boiling Point Elevation Calculations

Water with salt added boils at a higher
temperature than pure water. By how much will
the boiling point change if 100.g of salt is added
to 500. g of water? Kbwater = 0.52 oC/m
 Formula:
Tb = Kbm i

Kb : Molal Boiling Point elevation constant (oC/m)

i= = Pieces that the material dissociates into (for ionic compounds only)
(Keep I at 1 (one) for covalent compounds)

Freezing Point Depression and Boiling
Point Elevation
Formula
Melting
Point
(°C)
Boiling
Point
(°C)
Kf(°C/
m) (
Kb(°C
/m)
Water
H2O
0.000
100.000
1.858
0.521
Acetic acid
HC2H3O2
16.60
118.5
3.59
3.08
Benzene
C6H6
5.455
80.2
5.065
2.61
Camphor
C10H16O
179.5
...
40
...
Carbon disulfide
CS2
...
46.3
...
2.40
Cyclohexane
C6H12
6.55
80.74
20.0
2.79
Ethanol
C2H5OH
...
78.3
...
1.07
Solvent
Freezing point depression

By adding salt (or other solutes) to water, the
temperature of freezing drops it freezes at a
lower temperature
• Because H bonding is disturbed
• Dependent on how much solute is added
Freezing Point Depression Calcs

Antifreeze protects cars from freezing and
overheating. Calculate the freezing point
depression of a solution of 100. g of ethylene
glycol (C2H6O2) antifreeze in 0.500 kg of water.
Kf water = 1.86 oC/m
 Formula:
Tf = Kfm i

Kf : Molal Freezing Point depression constant (oC/m)

i= Pieces that the material dissociates into (for ionic compounds only)
(Keep I at 1 (one) for covalent compounds)

Freezing Point Depression and Boiling Point Elevation
Formula
Melting
Point
(°C)
Boiling
Point
(°C)
Kf
(°C/m)
Kb
(°C/m
)
Water
H2O
0.000
100.000
1.858
0.521
Acetic acid
HC2H3O2
16.60
118.5
3.59
3.08
Benzene
C6H6
5.455
80.2
5.065
2.61
Camphor
C10H16O
179.5
...
40
...
Carbon disulfide
CS2
...
46.3
...
2.40
Cyclohexane
C6H12
6.55
80.74
20.0
2.79
Ethanol
C2H5OH
...
78.3
...
1.07
Solvent
 Colligative
properties interactive