Download States of Matter

States of Matter
Gases, Liquids and Solids
Kinetic Molecular Theory of Gases
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Describes the motion of gas particles
Points of Kinetic Molecular Theory:
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Gas particles are point masses
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Explains low density and compressibility
Gas particles are in constant motion
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They move in straight lines at 100-1000m/s
They change direction only when they run into something
Collisions with container walls cause pressure
Explains Brownian motion and diffusion/effusion
Brownian Motion
Diffusion and Effusion
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Diffusion is the mixing of gases
Effusion is the migration of a gas through a
tiny orifice into an evacuated space
Graham’s Law of effusion
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Kinetic energy of a particle E = ½mv2
For the same energy, a heavier particle moves
more slowly [v = (2E/m)]
Graham’s Law
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Comparing two particles with the same
energy:
½mAvA2 = ½mBvB2
Rearrange and cancel:
vA2/vB2 = mB/mA
or
vA/vB = (mB/mA)
Graham’s Law
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Molar mass can be used for m, and the effusion
rate is directly proportional to v
Example problem: Find the molar mass of a gas
that effuses at a rate twice as slow as helium.
Solution: vA/vB = (mB/mA)
If gas A is helium, vA/vB = 2, so (mB/4) = 2
(mB/4) = 4
and mB = 16g/mol
Kinetic Theory
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Points of Kinetic Molecular Theory
continued
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Gas particles are point masses
Gas particles are in constant motion
All collisions between gas particles are perfectly
elastic
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No attractive forces between particles
Gases at the same temperature have the same
average kinetic energy
Gas Pressure
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Pressure is force/area
Units: N/m2 (Pascal) (kPa = 1000 Pascals)
psi = pounds per square inch
Barometers – pressure measured as height
of a column of mercury in a mercury
barometer
Standard pressure = 760mmHg = 1 atm =
101.325kPa = 29.92inHg = 14.7psi
Mercury
barometer
Open arm
manometer
Open arm manometer
Closed arm manometer
Dalton’s Law of Partial Pressures
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When different gases are mixed, every
particle contributes equally to the total
pressure
In a gas mixture, the contribution of each
component depends on the mole fraction of
that component – more particles means
more pressure
PT = PA + PB + PC …
Interparticle forces
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London dispersion forces (van der Waal’s
forces)
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Present in all particles
Dominant attractive force in non-polar molecular
compounds
Weakest of all interparticle forces
Due to temporary dipoles formed by electron
dislocation
Interparticle forces
Interparticle forces
Interparticle forces
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Magnitude of London dispersion forces
depends on
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Size of molecule – more surface area means
stronger forces
Polarizability of electron cloud – larger atoms’
electron clouds are more polarizable due to
shielding
Interparticle forces
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Dipole-dipole interactions – in polar
molecules opposite poles attract.
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Stronger than LDF
Interparticle forces
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Hydrogen bonding – strong interaction
between hydrogens (d+) and
electronegative atoms (d-)
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Hydrogen must be attached to an
electronegative atom (usually O or N)
Strongest non-bonding interaction
Hydrogen bonding
Hydrogen bonding in ice
Hydrogen bonding in DNA
Liquids
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More dense than gases, less than solids
(except water)
Incompressible
Particles are in contact but able to move
past each other
Liquids (and gases) are fluids – able to flow
Viscosity – resistance to flow
Liquids
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Viscosity increases with molecular surface
area (chain length)
Viscosity decreases with increasing
temperature
Surface tension depends on strength of
interparticle interactions
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Defined as the energy required to increase the
surface area of a liquid by a certain amount
Liquids
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Capillary action – ability of a liquid to rise in
a narrow tube
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Happens when adhesive forces between tube
and liquid are greater than cohesive forces in
liquid
Height to which the liquid will rise is a measure of
the difference in adhesive and cohesive forces
Solids
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Most dense state (except water)
Particles vibrate in place
Molecular solids
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Smallest particle is a molecule
Molecules are composed of all nonmetals held
together by covalent bonds
Molecules are held next to each other by LDF,
dipole-dipole interactions or H-bonds
Solids
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Molecular solids
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Low MP/BP
Insulators
Usually crystalline
Examples: water, sugar, caffeine
Network solids
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Covalent network solids
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Covalent bond throughout
Solids
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Highest MP/BP
No smaller units
Examples: C (diamond), Si, quartz
Ionic solids
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All salts - composed of metal & nonmetal
High MP/BP
Simplest unit is “formula unit”
Ions are held in place by attraction to oppositely
charged ion
Insulators unless melted or in solution
Solids
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Metals
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Held together by nondirectional metallic bonds
“Electron sea” of shared electrons
Not usually brittle like crystalline solids
Conductors of heat/electricity
Ductile, malleable, luster
Amorphous solids
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No regular arrangement
No sharp melting point
Examples: rubber, glass
Crystals
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Have a regular, repeating pattern of atoms
Unit cell is simplest repeating pattern
Have sharp melting points
Ionic and metallic crystals have high
melting points
Molecular crystals have low melting points
Crystal types
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Cubic
HALITE
Cubic crystals
FACE CENTERED CUBIC
BODY CENTERED CUBIC
Face centered cubic - halite
Body centered cubic - CsCl
Tetragonal crystals
RUTILE (TiO2)
Other crystal types
HEXAGONAL (QUARTZ)
ORTHORHOMBIC (CALCITE)
Phase changes and energy
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Solid/liquid: melting and freezing
Melting point: temperature at which vapor
pressures of solid and liquid are equal
Liquid/gas: vaporization and condensation
Boiling point: temperature at which vapor
pressure of liquid = atmospheric pressure
Boiling: vaporization at BP
Evaporation: vaporization at a lower temperature
Boiling and evaporation
Freezing point
Phase changes
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Solid/gas: sublimation and deposition
Energy and phase changes
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Melting is endothermic, freezing is exothermic
Boiling is endothermic, condensation is
exothermic
Sublimation is endothermic, deposition is
exothermic
Phase diagrams
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Phase diagrams show conditions under which an
element, compound or mixture will exist in a given
state
Variables are pressure (y) and temperature (x)
Triple point: temperature and pressure where
solid, liquid and gas can exist in equilibrium
Critical temperature: temperature above which a
substance cannot be liquified
Critical pressure: pressure necessary to liquefy a
substance at the critical temperature (together
they make the critical point)
Phase diagrams
Phase diagrams – UF6
Phase diagrams – CO2
Phase diagram - water
Phase diagrams – water (detailed)
Heating curves
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Shows change in temperature as heat is
added
Slope of curve gives specific heat
No change in temperature during phase
changes
Heating curves