states of matter gases, liquids and solids. kinetic molecular theory of gases describes the motion...
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Brownian MotionTRANSCRIPT
States of MatterGases, Liquids and Solids
Kinetic Molecular Theory of Gases Describes the motion of gas particles Points of Kinetic Molecular Theory:
Gas particles are point masses Explains low density and compressibility
Gas particles are in constant motion 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 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
Kinetic energy of a particle E = ½mv2
For the same energy, a heavier particle moves more slowly [v = (2E/m)]
Graham’s Law Comparing two particles with the same
energy:½mAvA
2 = ½mBvB2
Rearrange and cancel:vA
2/vB2 = mB/mA
orvA/vB = (mB/mA)
Graham’s Law 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 Points of Kinetic Molecular Theory
continued Gas particles are point masses Gas particles are in constant motion All collisions between gas particles are perfectly
elastic No attractive forces between particles
Gases at the same temperature have the same average kinetic energy
Gas Pressure 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
Mercurybarometer
Open armmanometer
Open arm manometer
Closed arm manometer
Dalton’s Law of Partial Pressures 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 London dispersion forces (van der Waal’s
forces) 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 Magnitude of London dispersion forces
depends on 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 Dipole-dipole interactions – in polar
molecules opposite poles attract.
Stronger than LDF
Interparticle forces Hydrogen bonding – strong interaction
between hydrogens (+) and electronegative atoms (-) 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 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 Viscosity increases with molecular surface
area (chain length) Viscosity decreases with increasing
temperature Surface tension depends on strength of
interparticle interactions Defined as the energy required to increase the
surface area of a liquid by a certain amount
Liquids Capillary action – ability of a liquid to rise in
a narrow tube 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
Most dense state (except water) Particles vibrate in place Molecular solids
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
Molecular solids Low MP/BP Insulators Usually crystalline Examples: water, sugar, caffeine
Network solids Covalent network solids
Covalent bond throughout
Solids Highest MP/BP No smaller units Examples: C (diamond), Si, quartz
Ionic solids 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
Metals 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 No regular arrangement No sharp melting point Examples: rubber, glass
Crystals 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 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 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 Solid/gas: sublimation and deposition Energy and phase changes
Melting is endothermic, freezing is exothermic Boiling is endothermic, condensation is
exothermic Sublimation is endothermic, deposition is
exothermic
Phase diagrams 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 Shows change in temperature as heat is
added Slope of curve gives specific heat No change in temperature during phase
changes
Heating curves