Chapter 13:
Vapor Pressure
Vapor Pressure is a measure of the force exerted by a gas above a liquid. Over time, the
number of particles of entering the vapor increases and some of the particles condense and
return to the liquid state.
In a system at constant vapor pressure, a dynamic equilibrium exists between the vapor and
the liquid. The system is in equilibrium because the rate of evaporation of liquid equals the rate
of condensation of vapor.
Vapor Pressure and Temperature Change
An increase in the temperature of a contained liquid increases the vapor pressure. This
happens because the particles in the warmed liquid have increased kinetic energy. As a result,
more of the particles will have the minimum kinetic energy necessary to escape the surface of
the liquid.
Boiling Point
Heating allows a greater number of particles at the liquid’s surface to overcome the attractive forces
that keep them in the liquid state.
Boiling point is the temperature at which the vapor pressure of the liquid is just equal to the external
pressure on the liquid.
When a liquid is heated to a temperature at which particles throughout the liquid have enough
kinetic energy to vaporize, the liquid begins to boil.
Boiling Point and Pressure Changes
Because a liquid boils when the vapor pressure is equal to the external pressure, liquids don’t always
boil at the same temperature. Because atmosphere pressure is lower at higher altitudes, boiling
points decrease at higher altitudes. At lower external pressure, the boiling point decrease, at higher
external pressure, the boiling point increases.
Normal boiling point is defined as the boiling point of a liquid at a pressure of 101.3 KPa.
Section 13.3
The general properties of solids reflect the orderly arrangement of their particles and the fixed
locations of their particles.
The melting point(mp) is the temperature at which a solid changes into a liquid. The melting
and freezing point of a substance are at the same temperature. At that temperature, the liquid
and solid phases are in equilibrium.
Equilibrium is a state in which opposing forces or influences are balanced,
Crystals are a solid in which the atoms, ions, or molecules are arranged in an orderly, repeating,
three-dimensional pattern called a crystal lattice.
The shape of a crystal reflects the arrangement of the particles within the solid.
Unit cells are the smallest group of particles within a crystal that retains the geometric shape of
the crystal.
Section 13.3 Cont.
Allotropes are two or more different molecular forms of the same element in the same physical
state.
Non-crystalline solids
Amorphous solid lacks an order internal structure. For example, rubber, plastic, and asphalts
are amorphous solids. Their atoms are randomly arranged.
Another example are glasses. A glass is a transparent fusion product of inorganic substances
that have cooled to a rigid state without crystallizing.
Section 13.4
Sublimation occurs in solids with vapor pressure that exceed atmospheric pressure at or near
room temperature. Iodine is an example of a substance that undergoes sublimation.
A phase diagram is the relationship among solid, liquid, and gas states of a substance in a
sealed container that can be represented on a single graph. It gives the conditions of
temperature and pressure at which a substances exists as a solid, liquid, and gas.
The conditions of pressure and temperature at which two phases exist in equilibrium are
indicated on a phase diagram by a line separating the phases.
The triple point describes the only set of conditions at which all three phases can exist in
equilibrium with one another.
Phase Diagram
Below the triple point, the vapor and liquid
cannot exist in equilibrium. Increasing the
pressure won’t change the vapor to a liquid.
The solid and the vapor are in equilibrium at
temperatures are low. With an increase in
pressure, the vapor begins to behave more like
a solid. For example, it is no longer easily
compressed.