Transcript Chapter 13 States of Matter - Alamo Heights Independent

States of Matter
Kinetic Theory
Kinetic
Theory is based on the idea
that particles of matter are always in
motion.
Kinetic energy = Energy of motion
Kinetic Theory
 Gases—particles
are far apart with
no attraction or repulsion
 Liquids—particles can slide past
one another but do experience
attraction
 Solids—particles are not free to
move but tend to vibrate about fixed
points
Kinetic Energy and Temperature

Directly Related

Increase Temperature: Molecules move faster.
 Intermolecular

Decrease Temperature: Molecules move slower.
 Intermolecular

forces can’t hold them together.
forces can now hold them together.
Absolute zero (0 K) is the temperature at which the
motion of particles theoretically ceases.
The K.M.T. of Gases

5 basic assumptions of the kinetic theory
as it applies to ideal gases:
 Ideal
gas: hypothetical gas that fits all
assumptions of KMT.
1.
Gases consist of large numbers of
tiny particles that are far apart.
 Most
of the volume occupied by a gas
is empty space.
KMT of gases
2. Collisions
between gas particles
and between particles and
container walls are elastic
collisions.

An elastic collision: there is no net
loss of total kinetic energy.
3. Gas
particles are in continuous,
rapid, random motion.
KMT of gases
4.
5.
There are no forces of attraction
between gas particles.
The temperature of a gas
depends on the average kinetic
energy of the particles of the gas.
**The kinetic-molecular theory
applies only to ideal gases.**
KMT Key terms
Fluidity
 Particles glide easily past one another.
 Because liquids and gases flow, they
are both referred to as fluids.
Compressibility
 Particles, which are initially very far
apart are crowded closer together.
KMT key terms
Diffusion
 spontaneous mixing of the particles of two
substances caused by their random
motion
Effusion
 a process by which gas particles pass
through a tiny opening.
Real gas

gas that does not behave according to the
kinetic-molecular theory.
What is gas pressure?
 Gas
pressure is the force exerted by a
gas per unit surface area.
 This
pressure results from the
collisions of gas particles with objects.

The SI unit of pressure is the pascal (Pa)

1 atm = 760 mm Hg (or torr) = 101.3 kPa
An Early
Barometer
Barometers are used to
measure atmospheric
pressure.
The normal pressure
due to the atmosphere
at sea level can support
a column of mercury
that is 760 mm high.
The Nature of Liquids
 Unlike
gas particles, liquid
particles are attracted to each
other, but they are still able to
slide past one another.
 More
dense than gas
 Less compressible
 Able to diffuse
Diffusion of liquids
Liquid key terms
Surface tension:
 a force that tends to pull adjacent parts of
a liquid’s to decrease surface area to the
smallest possible size.
Capillary action
 attraction of the surface of a liquid to the
surface of a solid

Usually pulls upward
The Nature of Liquids

The conversion of a liquid or solid to a
gas is called vaporization.
 When
this occurs at the surface of a liquid
that is not boiling it is called evaporation.
 Volatile liquids evaporate at room temp.
A liquid will evaporate faster when
heated—more kinetic energy to overcome
intermolecular forces.
 But…evaporation is a cooling process
because particles with the highest energy
escape 1st!

The Nature of Liquids
 The
boiling point (bp) is the
temperature at which the vapor
within the liquid is visible.
 Bubbles form throughout the liquid,
rise to the surface, and escape into
the air.
The Nature of Liquids
 Since
the boiling point is where the
vapor pressure equals external
pressure, the bp changes if the
external pressure changes.
 Normal boiling point is the bp at
normal atmospheric pressure or
1atm.