Liquids, Solids, and
Intermolecular Forces
Chapter 11
Intermolecular Forces (IMF’s)
• The state of a sample of matter (solid, liquid,
gas) depends on the magnitude of
intermolecular forces between the particles
relative to the amount of thermal energy in
the sample
– Remember, thermal energy is kinetic (motion,
vibrations, wiggling)
Review
• Solids
• Liquids
• Gases
Liquids
• Properties of liquids
– Liquids have high densities in comparison to gases
– Liquids have indefinite shape and assume the
shape of their container
– Liquids have a definite volume and are not easily
compressed
Solids
• Properties of solids
– Solids have high densities in comparison to gases
– Solids have a definite shape and do not assume
the shape of their containers
– Solids have a definite volume and are not easily
compressed
– Solids may be crystalline (ordered) or amorphous
(disordered)
A comparison of IMF’s
Properties of the Phases of Matter
Phase
Density
Shape
Volume
Strength of
IMF’s
Gas
Low
Indefinite
Indefinite
Weak
Liquid
High
Indefinite
Definite
Moderate
Solid
High
Definite
Definite
STRONG
Phase Changes
• You can change the phase of a sample by
changing the temp, pressure, or both
Add heat to a solid  liquid … add more heat (or increase
the pressure)  gas
Take away heat from a gas (or decrease the pressure) 
liquid … take away more  solid
Q?? Why does pressure only affect the (l) and (g)??
IMF’s
• The strength of the
intermolecular forces
between molecules or atoms
determine the phase (s, l, g)
• Strong IMFs result in solids
and liquids (high melting and
boiling points)
• Weak IMFs typically result in
gases (low Tm & Tb)
IMFs vs Bonds
• IMFs originate from the interactions between
charges, partial charges, and temporary
charges on molecules / atoms / ions
• Similar concept as bonds but are MUCH
WEAKER… they are not the bonds holding a
compound / molecule together, they are the
forces holding molecule A close to molecule B
Strengths
• Coulomb’s law:
E=
1
4πɛo
*
q1q2
r
• Bonds are strong because of large charged
attractions at a close range
• IMFs are much weaker due to weaker charges
at longer distances
• Water evaporates at 100⁰C
• But, to break the O-H bond would take
thousands of degrees
Different Types of IMFs
1.
2.
3.
4.
Dispersion Forces (London Force)
Dipole-Dipole Forces
Hydrogen Bonding
Ion-Dipole Forces
• Three of these can potentially occur in all
substances; H-Bonds can be found only in
mixtures
#1: Dispersion Forces (London Forces)
• Found in all molecules and atoms
• Caused by the fluctuations in electrons around
the atoms / within the molecules
– Causes an instantaneous dipole (+ side / - side)
• Think about your turn signal in your car…
Dispersion Forces
• CAN INDUCE and stick to / attract to other
neighboring molecules (IMF’s)
Magnitude of Dispersion Forces
• The magnitude of the force depends on two
things:
1. The size of the molecule
(simply put: the Molar Mass)
–the number of electrons
2. The shape / structure of the
molecule
1. Size
• Typically, the larger the molecule, the more
electrons and the more free “roaming”
around they are able to do.
• This opens the possibility of more electrons
‘clumping’ up and causing the +/• As the electron cloud increases, the dispersion
forces increase… causing a _________ B.P.
2. Structure
• The smaller the area of interaction, the lower
the forces will be
• The larger the area, the higher the forces
#2: Dipole-Dipole Forces
• The dipole-dipole force exists in all molecules
that are polar
• Polar molecules have permanent dipoles that
interact with perm dipoles of neighboring
molecules
• The (+) end of one molecule interacts with the
(-) of a neighboring molecule
Polar Molecules
• Remember, all molecules have dispersion forces.
• Polar molecules have Dipole-Dipole forces IN
ADDITION to these dispersion forces
– This raises their melting and boiling points
• The strength (magnitude) of a dipole moment is
measured in “Dipole moment” (D)
– The greater the D, the larger the dipole-dipole forces
Miscibility
• The polarity of molecules also determines the
miscibility of liquids
– The ability to mix without separating into two
phases
• All polar liquids are miscible with other polar
liquids
• All nonpolar liquids are miscible with other
nonpolar liquids
“LIKES DISSOLVE / MIX WITH LIKES”
Immiscibility
• Opposites generally do not mix together
– i.e. nonpolar + polar
Practice
• Which of the following molecules have dipoledipole forces?
(a) CO2
(b) CH2Cl2
(c) CH4
Draw the structure and determine if there are
polar bonds present… try to see if the molecule
is overall polar to have dipole-dipole forces
Practice
• Which of the following molecules have dipoledipole forces?
(a) C I4
(b) CH3Cl
(c) HCl
#3: Hydrogen Bonding
• Polar molecules with hydrogen can bond to
F,N,O
• CAN make Hydrogen bonds:
HF
NH3
H2O
• H-bonds are STRONG for IMFs!!
– LEADS TO HIGH MP’S AND BP’S!
• Sort of the “super dipole-dipole” force
Why??
• The large Δ in EN between H and N,O,F makes
a very (+) end on the H and a very (-) end on
the N, O, or F
• Also, because these elements are still quite
small, they are able to get pretty close to one
another, increasing the strength of attraction
(coulomb’s law)
What's going on here??
WARNING!!
H-BONDS ARE NOT BONDS!!!
STILL ONLY ABOUT 2-5% THE
STRENGTH OF A TYPICAL
COVALENT BOND!!
Practice #1
• One of the following is a liquid at room
temperature. Which one?
Formaldehyde (CH2O)
Fluoromethane (CH3F)
Hydrogen peroxide (H2O2)
**TRY drawing them out!!
Explanation to Practice #1
• All three have similar molar masses
– Meaning similar dispersion forces
• All three are also polar
– BUT… the hydrogen peroxide has H-O bonds
• So, there are hydrogen bonding in the H2O2
• Note: the other two DO have H and N,O,F… but
they need to be bonded directly to the H
– Why??
Practice #2
• What has the higher boiling point, HF or HCl??
• Why??
#4 Ion-Dipole Force
• The ion-dipole force occurs when an ionic
compound is mixed with a polar compound
– Occurs when something ionic is put into water and
dissociates into its ions
THESE ARE THE STRONGEST
INTERMOLECULAR FORCES
Ion-Dipole Force
• Think about NaCl in water
– The Na+ and Cl- ions break apart and float around
– The water SURROUNDS and attaches to the respective
opposite charges
IMF’s in Action
• Intermolecular forces is the reason and
determination of phase (s, l, or g)
• They also are the cause of other very
important (for life) properties
– Surface Tension
– Viscosity
– Capillary Action
Surface Tension
• Why does a fly fishermen’s fly float on the surface
of the water… the hook is metal and more dense
• The Surface Tension of a liquid is the energy
required to increase the surface are by unit
amount
• Water has a surface tension of 72.8 mJ/m2
– Meaning it takes 72.8 mJ of energy to increase the
surface area of water by one square meter
Why??
• Liquids always like to minimize their surface
area
• This creates kind of a “skin” on the top of the
liquid that resists penetration
• For the fishermen’s fly to sink it would have to
slightly increase the surface area of the
water… which is resisted by the water
Surface Tension and IMFs
• Surface Tension decreases with decreasing IMFs
• You can float a paper clip on water but cannot
float a paper clip on benzene (a nonpolar organic
solvent) because there are no H-bonds AND the
dispersion forces are not strong enough to
combat the increase in surface area
Water Droplets
• Surface Tension is also the reason why water
“beads” up and forms almost perfect spheres of
water – in space (with no gravity), large amounts
of water will ball up perfectly.
• Why? The sphere is the shape with the smallest
ratio of surface are to volume… minimizing the
number of molecules at the surface… minimizing
the potential energy of the system
Viscosity
• Viscosity is the resistance for liquid to flow
• Motor oil is more viscous than gasoline
• Viscosity is greater in substances with stronger
intermolecular forces because the molecules are
strongly attracted to each other, preventing them
from flowing past one another
Viscosity and Temperature
• How does temperature affect viscosity?
• Viscosity decreases with an increase in temperature
– The added kinetic energy will temporary overcome the
IMF’s and allow them to move more freely past one
another.
Capillary Action
• Capillary Action – the ability of a liquid to flow
against gravity up a narrow tube
– Caused by two forces:
1. Adhesion –
2. Cohesion –
Capillary Action
• Capillary Action – the ability of a liquid to flow
against gravity up a narrow tube
– Caused by two forces:
1. Adhesion – the attraction between liquid molecules
and the sides of the glass, root, etc. (to other things)
2. Cohesion – the attraction between molecules within
a liquid (to itself)
QUESTION:
What is a meniscus?? Why
does it happen??
Properties Affected by IMFs
• Phase Changes…
• Vaporization – the process by which thermal
energy can overcome IMFs and produce a
phase change from liquid to gas
• Removing energy will decrease their average
kinetic energy and bring them back down as a
liquid (condensation)
Open System
• Both vaporization and condensation occur at
EQUAL rates (as some go into the gaseous
phase, others fall back in the liquid phase)
– Average kinetic energy (some are lower than the
average, others are higher)
• In an open beaker / container
– The evaporation occurs more… the gases escape
into the atm and do not come back
Closed System
• In the same scenario,
the beaker or bottle
is not capped /
sealed
• The two processes
still occur at the
same rate
– Reach a happy
“dynamic
equilibrium”
Other Liquids
• What happens if you increase the
temperature?
• What happens to the vaporization process /
temperature if the liquid is not water… but
something with much weaker IMFs?
Volatility
• If a liquid evaporates easily it is volatile
• Those that do not evaporate easily are called
nonvolatile
• Example: Acetone (nail polish remover) is
more volatile than water.
– Motor oil is essentially nonvolatile at room temp
• Does not evaporate much or at all at room temp.
**Summary**
• The rate of vaporization increases with
increasing temperature
• The rate of vaporization increases with
increasing surface area!!
• The rate of vaporization increases with
decreasing strength of IMFs
Endo- or Exo- ??
• Vaporization vs. condensation
– Q: endo- or exo- ?
• If no additional energy is put into a system,
what happens to the E as the liquid is let to
evaporate?
Heat of Vaporization
• The amount of heat required to vaporize one
mole of a liquid to gas is called the heat of
vaporization (ΔHvap)
– The ΔHvap is always positive because the process is
always endothermic
– It is temperature / pressure dependent
Heat of Vaporization
• The ΔHvap for water at 100 ⁰ C is +40.7 kJ/mol
• When a liquid condensates, that same amount of
energy is released
The ΔH for water condensing at 100⁰C is:
ΔH = - ΔHvap = -40.7 kJ/mol (@100⁰C)
(releases 40.7 kJ per mole of H2O)
 (-) means leaves the system
Practice #1
• Calculate the mass of water (in grams) that
can be vaporized at its boiling point with 155
kJ of heat.
68.6 grams H2O
Practice #2
• (pg 480)
• Calculate the amount of heat (in kJ) required to
vaporize 2.58 kg of water at its boiling point.
More Practice (good TQ)
• Suppose that 0.48 g of water at 25⁰C condenses on the
surface of a 55-g block of aluminum that is initially at 25⁰C.
If the heat released during condensation goes only toward
heating the metal, what is the final temperature (in ⁰C) of
the metal block?
• The specific heat capacity of aluminum is 0.903 J/g ⁰C)
Vapor Pressure –
Back to a SEALED System!!
• Once the liquid/gas is allowed to reach this
dynamic equilibrium:
– The pressure of the gas with its liquid is called
vapor pressure
• The vapor pressure depends on the IMFs!!
– The weaker the IMFs, the more volatile the liquid,
the more molecules in the gaseous state, the
higher the vapor pressure!!
Volume and Vapor Pressure
• What happens to the vapor pressure in these
scenarios?
Systems WANT Equilibrium!!
• When a system in dynamic equilibrium is
disturbed, the system will respond so to
minimize the disturbance and reestablish its
state of equilibrium
Practice
• What happens to the vapor pressure of a substance
when its surface area is increased at constant
temperature?
a) The vapor pressure increases
b) The vapor pressure remains the same
c) The vapor pressure decrease
B!! The VP is independent of surface area… with the
increase in surface area DOES cause an increase in
vaporization… but also an increase in the condensation!!
Vapor Pressure and Temp
• When the temperature of a substance is
increased, the vapor pressure rises
– The thermal energy enables more molecules to
break free and enter the gaseous state
– A small change in T will cause a big increase… but
this curve relationship does level out and flatten
What do we call this??
• The boiling point!!!
Boiling Point
• The boiling point is the temperature at which its
vapor pressure equals the external pressure.
• This is the temperature where molecules, EVEN
THOSE IN THE MIDDLE, have the energy to break
free (overcome the IMFs between them and their
neighbors) ….
– Not just those on the surface bouncing and springing
out
The Normal Boiling Point
• The normal boiling point is the temperature
that the vapor pressure equals 1 atm (standard)
• This changes!! How?
• Increasing or decreasing the pressure (deviating
from the P = 1 atm) will change this boiling
point (no longer “normal” boiling point)
Supercritical Points
• Sealed system… but NOW increasing the Temperature
• The critical temperature (Tc) and critical pressure
(Pc) are the conditions needed to cause this change
–
SUPERCRITICAL FLUID
Supercritical Fluids
• Supercritical Fluids have properties of both liquids
and gases.
– Kind of an in between phase
• They are particularly good
solvents… dissolving a large
range of molecules.
• Supercritical CO2 dissolves
caffeine, making it easy to be
removed for de-caf
Phase Changes Involving Solids
• Think about a block of ice
– The thermal energy is significantly less than that
of liquids or gases
• However, at the surface, molecules may still
have enough energy to break free from their
neighboring IMFs holding them together and
float off into the gaseous phase
– What is this called?
Sublimation
• Sublimation - The molecules on the surface of
a solid that have enough energy to go directly
to the gaseous phase
• Some of the water vapor molecules may also
collide with the surface of the ice and be
‘pulled’ in / captured by the IMFs of the
solids… this process is known as deposition
Open System
• Sublimation usually occurs at a faster rate
than deposition because the gaseous particles
float away and never come back
• Think about ice cubes left in the freezer for
too long… they shrink… why?
• Notice the gradual growth of ice crystals on
the walls around the freezer??
Interesting
• Freezer burn is when foods are left in the
freezer too long and dry out
• The frozen water molecules are left to sublime
and never return
Heating Curves
• Thinking back to our block of ice…
– If we heat the cube up (from, say -10⁰C) the
molecules move faster and faster and the
temperature will increase
– However, you soon realize the temperature stops
increase even though you are adding in more heat
energy. What is happening?
Heating Curve for Water, H2O
Terms for Heating Curves
• Fusion – the phase change between solid and
liquid (aka: melting and freezing)
• Vaporization – (we already know) is the
change between liquid and gas (aka:
evaporation and condensation)
Heat of Fusion (ΔHfus)
• The fastest way to cool a drink is by adding ice
cubes to it
– They melt and pull absorb energy (endothermic)
H2O(s)  H2O(l) ΔHfus = 6.02 kJ/mol
• Freezing is the opposite (exothermic) –
releases energy to cool and freeze itself
H2O(l)  H2O(s) ΔH = -ΔHfus = - 6.02 kJ/mol
Compare ΔHfus and ΔHvap
• Typically, the ΔHfus is much less than the ΔHvap
• This is because the solid and liquid phases are
much closer together and require less energy
to exhibit this phase change.
– Gases are much more energized and will usually
require a much higher ΔH (in this case, ΔHvap)
Mathematical Relevance
Five Parts
• When there is NOT a phase change taking
place, you can calculate the heat (q) using
q = mCΔT
• Like Parts 1, 3, and 5
Five Parts
• When there IS A CHANGE taking place, you
need to use the ΔH fus/vap
• Remembering, these values are PER MOLE
q = nΔHfus
Like Part (2)
q = nΔHvap
Like Part (4)
Phase Diagrams
• Phase
Diagrams are
maps of the
phases of a
substance as a
function of
pressure (yaxis) and
temperature
(on the x-axis)
Important Parts
• Regions – solid, liquid, or gas
• Lines – each line represents a T and P that the
two phases are at dynamic equilibrium. The two
phases are equally stable and coexist.
• The Triple Point – a unique set of conditions
where all three phases are actually equally stable
and in equilibrium
• The Critical Point – the set of P and T above
which a supercritical fluid exists
Phase Diagram for CO2
Phase Diagram for Iodine
WATER IS SPECIAL!!