States of Matter

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States of Matter
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Three Phases of Matter
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Phase Differences
Solid definite volume and shape particles
packed in fixed positions particles are not free
to move
Liquid definite volume but indefinite shape
particles close together but not in fixed
positions particles are free to move
Gas neither definite volume nor definite shape
particles are at great distances from one
another particles are free to move
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A Molecular Comparison of Liquids and Solids
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Phase Changes
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Phase Changes

• Evaporation
• molecules at the surface gain enough energy to
  overcome IMF
• Volatility
• measure of evaporation rate
• depends on temp IMF

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Phase Changes
Boltzmann Distribution
p. 477
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Phase Changes

• Equilibrium
• trapped molecules reach a balance between
  evaporation condensation

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Phase Changes

• Vapor Pressure
• pressure of vapor above a liquid at equilibrium

v.p.

• depends on temp IMF
• directly related to volatility

temp
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Phase Changes

• Boiling Point
• temp at which vapor pressure of liquid equals
  external pressure

• depends on Patm IMF
• Normal B.P. - b.p. at 1 atm

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Phase Changes

• Melting Point
• equal to freezing point

• Which has a higher m.p.?
• polar or nonpolar?
• covalent or ionic?

polar
ionic
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Phase Changes

• Sublimation
• solid ? gas
• v.p. of solid equals external pressure

• EX dry ice, mothballs, solid air fresheners

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Phase Changes
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Phase Changes

• Energy Changes Accompanying Phase Changes
• Sublimation ?Hsub gt 0 (endothermic).
• Vaporization ?Hvap gt 0 (endothermic).
• Melting or Fusion ?Hfus gt 0 (endothermic).
• Deposition ?Hdep lt 0 (exothermic).
• Condensation ?Hcon lt 0 (exothermic).
• Freezing ?Hfre lt 0 (exothermic).

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Phase Changes

• Energy Changes Accompanying Phase Changes
• All phase changes are possible under right
  conditions.
• heat solid ? melt ? heat liquid ? boil ? heat gas
• endothermic
• cool gas ? condense ? cool liquid ? freeze ? cool
  solid
• exothermic

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Phase Diagram

• Represents phases as a function of temperature
  and pressure.
• Triple point
• Critical point
• Critical temperature the minimum temperature for
  liquefying a gas using pressure
• Critical pressure pressure required for
  liquefaction

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Phase Diagrams

• Show the phases of a substance at different temps
  and pressures.

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Phase Changes
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Carbon dioxide
Carbon dioxide
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Water
Water
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Carbon
Carbon
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Liquids vs. Solids

• LIQUIDS
• Stronger than in gases
• Y
• high
• N
• slower than in gases

• SOLIDS
• Very strong
• N
• high
• N
• extremely slow

IMF Strength Fluid Density Compressible Diffusio
n
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Properties of a Liquid

• Compressibility/Density
• Ability to Diffuse
• Surface Tension
• Capillary Action
• Viscosity
• Compare these aspects to both solids and gases
• Which ones are similar, which are different?
• Why?

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Properties of a Liquid
Diffusion
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Liquid Properties

• Surface Tension
• attractive force between particles in a liquid
  that minimizes surface area

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Surface Tension
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Liquid Properties

• Capillary Action
• attractive force between the surface of a liquid
  and the surface of a solid

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Some Properties of a Liquid

• Viscosity Resistance to flow

• High viscosity is an
• indication of strong
• intermolecular forces

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Solute
A solute is the dissolved substance in a solution.
Salt in salt water
Sugar in soda drinks
Carbon dioxide in soda drinks
Solvent
A solvent is the dissolving medium in a solution.
Water in salt water
Water in soda
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Dissolution of sodium Chloride
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Concentrated vs. Dilute
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Types of Solids

• Crystalline - repeating geometric pattern
• covalent network
• metallic
• ionic
• covalent molecular
• Amorphous - no geometric pattern

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Types of Solids
Covalent Molecular (H2O)
Covalent Network (SiO2 - quartz)
Amorphous (SiO2 - glass)
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Bonding in Solids

• Covalent-Network Solids
• ALL COVALENT BONDS.
• Atoms held together in large networks.
• Examples diamond, graphite, quartz (SiO2),
  silicon carbide (SiC), and boron nitride (BN).
• In diamond
• each C atom is tetrahedral there is a
  three-dimensional array of atoms.
• Diamond is hard, and has a high melting point
  (3550 ?C).

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Network Atomic Solids
Some covalently bonded substances DO NOT form
separate molecules.
Diamond, a network of covalently bonded carbon
atoms
Graphite, a network of covalently bonded carbon
atoms
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Amorphous solids

• considerable disorder in their structures
  (glass and plastic).

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Types of Solids
Ionic (NaCl)
Metallic
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Bonding in Solids

• Ionic Solids
• Ions (spherical) held together by electrostatic
  forces of attraction.
• There are some simple classifications for ionic
  lattice types.

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Bonding in Solids

• Metallic Solids
• Problem the bonding is too strong for London
  dispersion and there are not enough electrons for
  covalent bonds.
• Resolution the metal nuclei float in a sea of
  electrons.
• Metals conduct because the electrons are
  delocalized and are mobile.

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Metal Alloysare solid solutions

• Substitutional Alloy some metal atoms replaced
  by others of similar size.
• brass Cu/Zn

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Metal Alloys(continued)

• Interstitial Alloy Interstices (holes) in
  closest packed metal structure are occupied by
  small atoms.
• steel iron carbon

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Bonding in Solids
Pg. 189-193 (Ch. 6-5)

• Molecular Solids
• Intermolecular forces dipole-dipole, London
  dispersion and H-bonds.
• Weak intermolecular forces give rise to low
  melting points.
• Room temperature gases and liquids usually form
  molecular solids and low temperature.
• Efficient packing of molecules is important
  (since they are not regular spheres).

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Bonding in Solids
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Definition of IMF

• Attractive forces between molecules.

• Much weaker than chemical bonds within
  molecules.

• a.k.a. Van der Waals forces

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Types of IMF
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London-Dispersion Forces

• LDF act between all atoms and molecules
• They are the ONLY IMF that acts among noble-gas
  atoms and non-polar molecules
• Strength of the force is directly related to
  number of interacting electrons
• Increase atomic/molar mass ? increase in numbers
  of electrons ? increase in LDF

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Types of IMF

• London Dispersion Forces

View animation online.
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Definitions

• Dipole-Dipole
• The forces of attraction between polar molecules
• Hydrogen Bonds
• The IMF in which a hydrogen atom that is bonded
  to a highly electronegative atom is attracted to
  an unshared pair of electrons of an
  electronegative atom in a nearby molecule

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Types of IMF

• Dipole-Dipole Forces

View animation online.
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Types of IMF

• Hydrogen Bonding

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Determining IMF

• NCl3
• polar dispersion, dipole-dipole
• CH4
• nonpolar dispersion
• HF
• H-F bond dispersion, dipole-dipole, hydrogen
  bonding

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Intermolecular Forces
Forces of attraction between different molecules
rather than bonding forces within the same
molecule.

• Dipole-dipole attraction
• Hydrogen bonds
• Dispersion forces

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Forces and States of Matter

• At STP, substances with
• very weak intermolecular attraction
• gases
• strong intermolecular attraction
• liquids
• very strong intermolecular attraction
• or ionic attraction
• solids