Transcript Kinetic Molecular Theory

Kinetic Molecular Theory
(Do not take notes)
• Composition & structure of molecules
affect the chemical & physical properties
of matter.
• Solids & liquids have a lot of variation
between both physical & chemical
properties.
• Gases, however, have very similar
physical and chemical properties.
Kinetic Molecular Theory
• Explains the behavior of gases in terms of
particles in motion
• Assumptions:
– Particles are very small with huge amounts of empty
space between them (so no attractive or repulsive
forces between them).
– Gas particles are in constant, random motion. Move
in straight lines until they collide with each other or
container walls. Collisions are elastic – no energy is
lost.
– Mass & velocity affect the kinetic energy of individual
gas particles (K.E. = ½ mv2)
Using KMT to explain gases
• Low density: gases have extremely low densities
because of the large spaces between gas
molecules
• Compression & expansion: gases can be
compressed if you decrease the space between
gas molecules; the random motion of gas
particles will cause gases to expand to any space
available
• Diffusion & effusion: both of these properties
occur because of the constant random motion of
gas particles and the lack of attractive or
repulsive forces among gas particles.
Gas Pressure
• Pressure is the force exerted compared to
the area so lying down on ice spreads out
your weight.
• Air pressure or atmospheric pressure is
exerted in all directions since air is all
around us moving randomly
• Measured with a barometer – one end is a
vacuum; the other is open to the air
• Many units are used for pressure: pascal,
torr, psi and atmospheres are all common
Intermolecular Forces of Attraction
• Occur between identical particles & help
explain how solids, liquids and gases exist
at the same temperature
• Three main types: dispersion, dipoledipole and hydrogen bonds.
• All intermolecular forces are weaker than
intramolecular forces (they have to be)
Dispersion Forces
• The weakest intermolecular force
• Significant in nonpolar substances
because no other intermolecular forces
exist
• Explains why non-polar liquids are hard to
pour and evaporate quickly (alcohols,
gasoline, etc)
Dipole-Dipole Forces
• Occur in polar molecules because
permanent dipoles exist
• Molecules line up head to tail or positive
region near a negative region
• Stronger than dispersion forces but
weaker than hydrogen bonds
Hydrogen bonds
• A special type of dipole-dipole bond
• Occur only between H and fluorine,
oxygen or nitrogen
• Explains why water (18.00 g/mol) has
such a huge surface tension and high
boiling point while methane (16.05 g/mol)
has very low surface tension and is a gas
at room temperature
Properties of liquids
• Liquid particles have random motion but
greater attractive forces than gases, so
liquids have a volume but no shape
• Density is also greater than gases while
compression is much less since liquid
particles are already close
• Liquids are less fluid than gases because
they diffuse much slower due to the
intermolecular forces interfering
Properties of liquids
• Viscosity is the measure of the resistance
of a liquid to flowing
• Viscosity is influenced by intermolecular
forces, shape of the particle and
temperature
• Year round oils actually change shape
from spheres in cold weather to long
strands when hot to increase viscosity
Properties of liquids – surface tension
• Particles at the surface have a greater
downward pull than particles in the middle
• Greater intermolecular forces usually means
greater surface tension
• Water forms a drop because of its high
surface tension
• Surfactants are compounds that lower
surface tension in a substance, like detergent
in water
Properties of liquids – capillary action
• Occurs when water is placed in a container
or substance that it’s highly attracted to
• If the attraction (adhesion) is greater than
water’s attraction to itself (cohesion),
capillary action occurs.
• Like the downward curve you see when a
liquid is in a glass
• When water is drawn up between the
cellulose fibers of paper towels or the
crystals in a diaper
Properties of Solids
• According to KMT solids have as much
kinetic energy at room temp as gases or
liquids, but attractive forces are so great
that particles in a solid move around a
fixed point
• Most solids are more dense than liquids or
gases (water is the exception)
Phase Changes
• Most substances can exist in 1 of 3 states
on Earth: solid, liquid or gas depending on
the pressure and temperature
• When energy is added or removed a
substance may change from one phase to
another
• KMT predicts this because as temperature
increases, motion increases and forces of
attraction can be overcome.
Phase Changes that Require NRG
• Includes melting, vaporizing (boiling),
evaporating and sublimation
• For melting, boiling and sublimation, the
temperature plateaus where the phase
change occurs because all incoming
energy is being used to break bonds
• Stronger bond = more energy = higher
boiling point or lower freezing point
Evaporation
• When a liquid changes to a gas but only
molecules at the surface escape to
become vapor
• Occurs at a lower temperature than boiling
point and is also a slower process than
vaporization
• Explains why sweating cools us or other
animals down
Sublimation
• Process in which a solid goes directly to a
gas without a liquid phase
• Iodine, dry ice, moth balls and air
fresheners all tend to sublime
• Water (ice) will sublime if pressure is
decreased – like when ice cubes shrivel
up in your freezer
• Freeze dried food is when food is frozen in
a vacuum so the water sublimes. Makes
food lighter and unable to grow bacteria.
Phase Changes that release NRG
• Includes condensation, freezing, &
deposition
• Occurs when a vapor or liquid comes into
contact with a cooler substance. The gas
or liquid loses heat and if enough heat is
lost, the forces of attraction become great
enough to form a liquid or solid.
• This is why clouds form and why it warms
up when it starts snowing
• Deposition is when a gas goes right to a
solid like when it snows or frost forms
Phase Diagrams
• Graphs pressure versus temperature
• Lines on diagram indicate where more than
one phase exists
• Triple point is the temperature and pressure
where all 3 states would exist or all 6 phase
changes can occur
• Critical point = pressure and temperature
where water can not exist as a liquid;
regardless of how much pressure you put on
the vapor beyond this point it will be a gas