Solids, Liquids and Gases Solid Fixed Volume Rigid, definite
shape Liquid Fixed Volume Indefinite Shape Particles easily glide
past one another Gas Particles farther apart than in liquids No
definite volume or shape Flowing and Compressible
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Particle Diagram
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The Kinetic Theory of Matter 1827: Robert Brown Scottish
Botanist Studied pollen grains on water Brownian Motion: Constant,
random motion of tiny chunks of matter Kinetic Energy: The energy
of moving objects
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Kinetic Model of Gases Particles move in a straight line until
they strike a container wall or another gas particle. Particles do
NOT lose speed when collisions are made (ELASTIC COLLISIONS) Gases
fill container because each particle moves until it hits a wall and
then changes direction
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Kinetic Model of Liquids Particles maintain volume, but not
shape Particles slide past one another, but they do not move as
straight or as quickly as particles of gases.
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Properties of Liquids Surface tension: the energy required to
increase the surface area of a liquid by a unit amount. Viscosity:
a measure of a liquids resistance to flow. Surface tension: The net
pull toward the interior of the liquid makes the surface tend to as
small a surface area as possible and a substance does not penetrate
it easily. Viscosity: Related to mobility of a molecule
(proportional to the size and types of interactions in the liquid).
Viscosity decreases as the temperature increases since increased
temperatures tend to cause increased mobility of the molecule.
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Intermolecular Forces Intermolecular forces: attractions and
repulsions between molecules that hold them together.
Intermolecular forces (van der Waals forces) hold molecules
together in liquid and solid phases. Ion-dipole force: interaction
between an ion and partial charges in a polar molecule.
Dipole-dipole force: attractive force between polar molecules with
positive end of one molecule is aligned with negative side of
other. London dispersion Forces: interactions between
instantaneously formed electric dipoles on neighboring polar or
nonpolar molecules.
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Kinetic Model of Solids Strong force between particles = rigid
structure Particles do not slide past each other, but instead
vibrate, or bounce back and forth between each other. Form a
CRYSTAL LATTICE, or repeating, fixed, 3-D arrangement
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Structure of Solids Types of solids: Crystalline a well defined
arrangement of atoms; this arrangement is often seen on a
macroscopic level. Ionic solids ionic bonds hold the solids in a
regular three dimensional arrangement. Molecular solid solids like
ice that are held together by intermolecular forces. Covalent
network a solid consists of atoms held together in large networks
or chains by covalent networks. Metallic similar to covalent
network except with metals. Provides high conductivity. Amorphous
atoms are randomly arranged. No order exists in the solid.
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Changing States Evaporation: particles of a liquid form a gas
by escaping from the surface VOLATILE liquids evaporate quickly
Vapor Pressure: Pressure at liquid/gas equilibrium Boiling: Vapor
Pressure is the pressure at the surface
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Changing States Condensation: Gaseous particles come closer
together and form a liquid Sublimation: particles of a solid escape
from the surface and form a gas Deposition: particles of a gas are
forced together under pressure and form a solid
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Changing States Melting: Particles of a solid begin to lose
their crystal lattice and slide past one another. Freezing:
particles form crystal lattice and form a rigid structure
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Changing States Heat of Vaporization Water vapor and Boiling
water are at the same temperature Gaseous particles form because
kinetic energy increases Joule: (J) SI unit of energy; energy
required to lift 1 kg mass by 1 meter against gravity Energy
absorbed when 1 kg of a liquid vaporizes at normal boiling
point.
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Changing States Heat of Fusion Melting Point: Temperature at
which crystal lattice begins to disintegrate. Freezing Point =
Melting Point Energy released as 1kg of a substance solidifies at
its freezing point.
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Heat of fusion for H 2 O: 80.0 cal/g A calorie is the amount of
heat needed to raise the temperature of one gram of water 1 0 C
1.Convert heat of fusion to Joules (J). 1.00 cal = 4.18J 4.18J/1
cal X 80.0 cal = 334J 2. What is the molar heat of fusion for water
( in cal & in J) a) 80.0 cal/g X 18.00g= 1440cal or 1.44 kcal
b) 334J X 18.00g= 6010J or 6.01kJ
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Sample Problems Heat of Vaporization of water: 540. cal/g
1.Convert heat of vaporization to J & kJ 540.cal/g X 4.18J/cal
= 2260J/g or 2.26kJ/g 2. What is the molar heat of vaporization for
water? (cal & J) a)540. cal/g x 18.00g/mol = 9720 cal or
9.72kcal b) 2260J/g x 18.00g/mol = 40.7 kJ
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Change of State Problems If you wanted to melt 16.00g ice at
25.0 0 C what would you have to do first? Change the temperature to
the melting temp: zero 0 Q = m t Cp Q= heat M= mass Cp= Specific
Heat Capacity
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Specific Heat Capacity = the amount of heat needed to raise the
temperature of 1.00g of a substance 1.00 0 C. Cp of water= 1.00
cal/g.K Cp of ice= 0.493 cal/g.K Cp of steam= 0.447 cal/g.K 1.To
melt 16.00g of ice at -25.0 its temperature must be raised to zero
C: Q = m t Cp Q = (16.00g)(25.0)(0.493) Q= 197cal
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2. How much energy is needed to melt 68.0 g of ice at 0 0 C
into water at 0 0 C ? 3. How much energy is needed to change 42.0g
of water at 100.0 0 C into steam at 100.0 0 C? 4. How much energy
is needed to raise the temperature of 57.0g of water at 25.0 0 C to
its boiling point? 5. How much energy is needed to convert 15.0g of
ice at -6.00 0 C to water at 25.0 0 C?
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The triple point of a substance indicates the temperature and
pressure at which a solid, liquid and vapor can coexist at
equilibrium. The critical point of a substance indicates critical
temperature and critical pressure. Critical temperature is the
temperature above which only the vapor can exist. Critical pressure
is the lowest pressure (at the critical temperature) at which the
substance can exist as a liquid. At any lower pressure only the
vapor exists.