states of matter gases, liquids and solids. kinetic molecular theory of gases describes the motion...

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