Phase Changes
• Matter can change from one
form to another. As this occurs,
energy also changes.
*As one proceeds from ice to water
to water vapor, there is an increase
in kinetic energy.
* The changes of phase are not
chemical, they are physical
changes.
Heating and Cooling Curves
Heating Curve
If there is a change in the
temperature there is a change in
kinetic energy because there is
a change in the average motion
of the particles.
During a phase change there is no
change in temperature, therefore
no change in kinetic energy.
Instead, energy goes to breaking
bonds so it is a potential energy
change.
Energy Changes
Along
AB KE, no PE
BC no KE, PE
CD KE, no PE
DE no KE, PE
EF KE, no PE
A cooling curve would be the opposite.
Gases
• Gases (g): Transparent,
compressible, expand without limit,
have no shape/volume. **Take the
shape and volume of their container.
Gases exert pressure: STP:
defined as standard temperature
and pressure
*Found on Table A 101kPa or
1atm
*Pressure can also be 760 torr
or 760 mm Hg
Liquids
• Liquids: no definite shape/ but
definite volume with very low
compressibility.
• Compressibility is the ability to
occupy less space.
Boiling Point
• The temperature at which the
vapor pressure of a liquid
reaches atmospheric
pressure; therefore allowing
particles to escape as a gas.
*When vapor pressure of a
liquid = atmospheric pressure
As atmospheric pressure increases
one must raise the vapor
pressure of the liquid by
increasing its temperature.
Normal Boiling Point is measured at
standard pressure
For water it is: 100C or 373K
Vapor Pressure
•
•
•
See Table H
The pressure exerted by the vapor
evaporating off the surface of a liquid.
Each liquid has its own vapor pressure.
• As the temperature of the liquid
increases, the vapor pressure
of that liquid increases.
Evaporation
• The change of phase from liquid
to gas.
Heat of Vaporization: The amount
of heat energy required to
vaporize a given mass of a liquid
to gas at a constant
temperature. This is an
endothermic process, energy is
being absorbed.
Each substance has its own Heat of
Vaporization.
For water at its normal boiling
temperature of 100C and
standard pressure the heat of
vaporization is 2260 joules per
gram. (Table B)
Condensation
• The change of phase from
gas to liquid. This is an
exothermic process.
• The Heat of Condensation is
the direct opposite of the heat
of vaporization. The quantity of
heat energy is the same as for
the heat of vaporization, but
instead of being absorbed the
heat energy is being released.
Solids
• Definite shape/definite
volume.
***Regular Geometric Pattern***
 Melting (fusion):
An
endothermic process in which
a solid becomes a liquid. For
water at Standard Pressure the
temperature at which melting
occurs is 0C / 273K.
Freezing (Solidification)
• Freezing is the direct opposite of
melting, but instead of being an
endothermic process where
energy is absorbed, it is an
exothermic process where
energy is released. Water
freezes at 0C / 273K
Heat of Fusion
• The amount of heat energy required
to change a given mass of solid to
liquid at a constant temperature.
Each substance has its own
Heat of Fusion.
For water at standard
pressure, this quantity of heat
is 334 J/gK (Table B)
Heat of Solidification (Crystallization)
• The direct opposite of the
Heat of Fusion. Since
solidification is an
exothermic process, the heat
energy is released instead of
absorbed.
Sublimation
• The change of phase from
solid to gas, completely
skipping the liquid phase.
• This generally occurs only in
solids with high vapor
pressures and weak
intermolecular forces of
attractions.
• Examples: Dry ice (CO2),
• paradichlorobenzene (moth
balls), I2