States of Matter
Chapter 13
13-1 The Kinetic Molecular Theory
• KMT – describes the behavior of gases in
terms of particles in motion
Size:
• Gas particles are small
• Spaced very far apart
• No significant attractive or repulsive forces
between the gas particles
The Kinetic Molecular Theory
Motion:
• Gas particles are in constant random
motion
• Kinetic energy is transferred when
particles collide with each other or the
sides of their container
• The collisions are elastic (no energy is
lost)
The Kinetic Molecular Theory
Energy:
• Depends on the particles mass and
velocity
• KE = ½ mv2
• Temperature is a measure of the average
kinetic energy of the particles in a gas
• 2 gases at the same temperature have the
same average kinetic energy
Explaining the behavior of gases
Compression and expansion:
• A gas will expand to fill its container
• Random motion of gas particles fill the
empty spaces and it expands until the
container stops it.
• Large amounts of empty space between
gas particles allow it to be compressed
(squeezed into a smaller space)
Explaining the behavior of gases
• Diffusion refers to the movement of one
material through another
• Gas particles flow past each other easily
because there are no significant forces of
attraction between them.
• Particles diffuse from areas of high
concentration to areas of low concentration
Explaining the behavior of gases
Effusion – a gas escaping through a tiny
opening
• Gases with a higher mass effuse slower
than gases with less mass
Example: A tire
deflating from a
puncture
Gas pressure
Pressure is force per unit area.
• When gas particles collide with the walls of their
container, they exert pressure on the walls.
• SI unit for pressure is the Pascal (Pa)
• 1 atm = 101.3 kPa = 760 mm Hg = 760 Torr
• Convert 4.5 atm to kPa?
4.5 atm x
101.3kPa
1atm
= 455.85 kPa = 460 kPa
Gas pressure
Dalton’s Law of partial pressure:
Ptotal = P1 + P2 + P3 ….
Example:
• Air is made up of four main gases: N2, O2, Ar,
and CO2
• Air pressure at sea level is approximately 760
mm Hg.
• Calculate the partial pressure of oxygen, given
the following partial pressures: N2, 594 mm Hg;
Ar, 7.10 mm Hg; and CO2, 0.27 mm Hg.
Gas pressure
• 760 mmHg = Poxygen + 7.10 mm Hg + 0.27
mm Hg + 594 mm Hg
• The partial pressure of oxygen is about
159 mm Hg.
Homework:
• Section 13.1 # 4,5,6 & 7-12
pg. 392
Forces of Attraction
• The attractive forces that hold particles
together in ionic and covalent bonds are
called intramolecular forces. (within the
molecule)
• Intermolecular forces hold particles of
different molecules together. (between
molecules)
• There are three types of intermolecular
forces: dispersion forces, dipole–dipole
forces, and hydrogen bonds.
13.3 Liquids and Solids
Liquids:
• Fixed volume, no fixed shape
• Densities are much greater than that of
gases
• Particles are packed closer together
• Is a fluid – has the ability to flow and
change shape
13.3 Liquids and Solids
Viscosity – measure of the resistance of a
liquid to flow
• Example: water vs. molasses
• Viscosity is affected by the type of
intermolecular forces and by the
temperature
• As temperature increases,
viscosity decreases
13.3 Liquids and Solids
Surface Tension: The particles on the
surface of a liquid are being pulled down
by intermolecular forces (between two
molecules)
• The stronger the attraction between
particles the stronger the surface tension
13.3 Liquids and Solids
Capillary action occurs when adhesive
forces are greater than cohesive forces
• Adhesion is the force of attraction between
molecules that are different, such as water
molecules and the molecules of silicon
dioxide in glass.
• Cohesion is the force of attraction between
identical molecules, such as water
molecules
Example: water in a graduated cylinder
13.3 Liquids and Solids
Solids:
• Particles in a solid are in constant motion
(vibrating)
• Definite shape and volume
• Very strong attractive forces acting
between molecules
• More dense than most liquids
• Most solids are in the shape of crystals
(repeating geometric patterns)
13.3 Liquids and Solids
• The type of ions
and the ratio of
ions determine
the structure and
the shape of the
crystal.
13.3 Liquids and Solids
• The particles in an amorphous solid are
not arranged in a regular, repeating
pattern and do not form crystals (they can
take on different shapes).
• Examples of amorphous solids include
glass, rubber, and many plastics.
13.4 Phase Changes
• Most substances can exist in three
states— solid, liquid, and gas—depending
on the temperature and pressure.
• States of substances are called phases
when they coexist as physically distinct
parts of a mixture, such as ice water.
• When energy is added to or taken away
from a system, one phase can change into
another.
13.4 Phase Changes
Phase Changes that require energy:
• MELTING – solid to liquid
melting point – the point at which the
bonds holding the solid
together are broken
• VAPORIZATION – liquid to a gas or vapor
evaporation – vaporization only on the
surface of the liquid
13.4 Phase Changes
Phase Changes that require energy:
Boiling Point – the point at which vapor
pressure is greater than
air pressure
• SUBLIMATION – a solid changes directly
to a vapor
example: Dry Ice
13.4 Phase Changes
Phase Changes that release energy:
• CONDENSATION – gas to liquid
• FREEZING – liquid to solid
freezing point – temperature at which a
liquid turns to solid
• DEPOSITION – a substance changes
directly from a gas to a
solid
example: frost on a window
13.4 Phase Diagrams
• Temperature and pressure control the
phase of a substance.
• A phase diagram is a graph of pressure
versus temperature that shows in which
phase a substance exists under different
conditions of temperature and pressure.
• A phase diagram typically has three
regions, each representing a different
phase and three curves that separate
each phase.
13.4 Phase Diagrams
• The triple point is the point on a phase
diagram that represents the temperature
and pressure at which three phases of a
substance can coexist
• The critical point indicates the critical
pressure and the critical temperature
above which a substance cannot exist as
a liquid.
Phase Diagram for H2O