Transcript Chapter 12: Intermolecular Attractions and the Properties

• A change in state is called a phase change
• Evaporation is the change in state from
liquid to gas
• Sublimation is the change from solid to gas
• Both deal with the motion of molecules
• You have also probably noticed that the
evaporation of liquids produce a cooling
effect
Molecules that are able
to escape from the
liquid have kinetic
energies larger than the
average. When they
leave, the average
kinetic energy of the
remaining molecules is
less, so the temperature
is lower.
• The rate of evaporation depends on the
temperature, surface area, and strength of
the intermolecular attractions
At higher
temperature, the
total fraction of
molecules with
kinetic energy
large enough to
escape is larger so
the rate of
evaporation is
larger.
• For a given liquid, the rate of evaporation
per unit surface area is greater at a higher
temperature
Kinetic energy distribution in two different liquids, A and B, at
the same temperature. The minimum kinetic energy required by
molecules A to escape is less than for B because the
intermolecular attractions in A are weaker than in B. This
causes A to evaporate faster than B.
• As soon as a liquid is placed in an empty
container, it begins to evaporate
• Once in the gas phase, molecules can
condense by striking the surface of the
liquid and giving up some kinetic energy
• The rate of evaporation equals the rate of
condensation at equilibrium
• This can occur in a system where the
molecules are constrained to remain close to
the liquid surface
(a) The liquid begins to evaporate in the closed container.
(b) Dynamic equilibrium is reached when the rate of
evaporation and condensation are equal.
• Similar equilibria are reached in melting
and sublimation
At the melting point a
solid begins to change
into a liquid as heat is
added. As long no heat
is added or removed
melting (red arrows)
and freezing (black
arrows) occur at the
same rate an the
number of particles in
the solid remains
constant.
At equilibrium, molecules evaporate from the solid at
the same rate as molecules condense from the vapor.
• When molecules evaporate, the molecules
that enter the vapor phase exert a pressure
called the vapor pressure
• The equilibrium vapor pressure is the
vapor pressure once dynamic equilibrium
has been reached
• The equilibrium vapor pressure is usually
referred to as simply the vapor pressure
• Vapor pressures can be measured using a
manometer
• Measuring the (equilibrium) vapor pressure
of a liquid at a specific temperature
Variation of vapor pressure with temperature. Ether is
said to be volatile because it has a high vapor pressure
near room temperature.
• Volume changes can effect vapor pressure
(a) Equilibrium exists between liquid and vapor. (b) The volume
is increased, the pressure drops, and the rate of condensation
drops. (c) Once more liquid evaporates, equilibrium is reestablished and the vapor pressure returns to its initial value.
• Solids also have vapor pressures
• At a given temperature, some of the
particles at the solid have enough kinetic
energy and escape into the vapor phase
• When vapor particle collide with the
surface, they can be captured
• The pressure of the vapor that is in
equilibrium with the solid is called the
equilibrium vapor pressure of the solid
• The boiling point of a liquid can be defined
as the temperature at which the vapor
pressure of the liquid is equal to the
prevailing atmospheric pressure
• The normal boiling point is the
temperature at which the vapor pressure is 1
atm
• Molecules with higher intermolecular forces
have higher boiling points
Boiling points of
the hydrogen
compounds of
elements of Groups
IVA, VA, VIA, and
VIIA of the periodic
table. The boiling
points of molecules
with hydrogen
bonding are higher
that expected.
• Heating and cooling curves can be used to
determine melting and boiling points
(a) A heating curve observed when heat is added at a constant
rate. (b) A cooling curve observed when heat is removed at a
constant rate. The “flat” regions of the curves identify the
melting and boiling points. Supercooling is seen hear as the
temperature of the liquid dips below its freezing point.
• The energy associated with the phase
changes can be expressed per mole
• The molar heat of fusion is the heat
absorbed by one mole of solid when it melts
to give a liquid at the same temperature and
pressure
• The molar heat of vaporization is the heat
absorbed when one mole of the liquid is
changed to one mole of vapor at constant
temperature and pressure
• The molar heat of sublimation is the heat
absorbed by one mole of a solid when it
sublimes to give one mole of vapor at
constant temperature and pressure
• All of these quantities tend to increase with
increasing intermolecular forces
• The concept of equilibrium is important and
will be encountered again
• Equilibria are often disturbed or upset
• According to Le Chatelier’s Principle
– When a dynamic equilibrium in a system is
upset by a disturbance, the system responds in a
direction that tends to counteract the
disturbance and, if possible, restore equilibrium
• The term position of equilibrium is used to
refer to the relative amounts of the
substance on each side of the double
(equilibrium) arrows
• Consider the liquid vapor equilibrium
liquid  heat


vapor
• Increasing the temperature increases the
amount of vapor and decreases the amount
of liquid
• We say that the equilibrium has shifted
• This can also be referred to as a right shift
because more vapor is produced at the
expense of the liquid
• Temperature-pressure relationships can be
represented using a phase diagram
The phase diagram of water. The line AB is the vapor pressure curve
for ice; BD the vapor pressure curve for liquid water; BC the melting
point line; point B the triple point (the temperature where all three
phases are in equilibrium); and point D labels the critical point (and
the critical temperature and pressure). Above the critical temperature
a distinct liquid phase does not exist, regardless of pressure.
• A substance that has a temperature above its
critical temperature and a density near its
liquid density is called a supercritical fluid
• Supercritical fluids have some unique
properties that make them excellent solvents
• The values of the critical temperature tends
to increase with increased intermolecular
attractions between the particles