Solids, Liquids, Gases (and Solutions)

1
Solids, Liquids, Gases (and Solutions)
2
Three Phases of Matter
3
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
4
A Molecular Comparison of Liquids and Solids
5
Phase Changes
6
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).

7
Phase Changes

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

8
(No Transcript)
9
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

10
Phase Changes
11
Carbon dioxide
Carbon dioxide
12
Water
Water
13
Carbon
Carbon
14
Types of Solids

• Crystalline Solids highly regular arrangement
  of their components table salt (NaCl), pyrite
  (FeS2).

15
Representation of Components in a Crystalline
Solid

• Lattice A 3-dimensional system of points
  designating the centers of components (atoms,
  ions, or molecules) that make up the substance.

16
Ionic Solids
17
Bonding in Solids

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

18
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).

19
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
20
Amorphous solids

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

21
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.

22
(No Transcript)
23
MetalsClosest Packing of Atoms
24
Metal Alloysare solid solutions

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

25
Metal Alloys(continued)

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

26
Molecular Solids
Strong covalent forces within molecules
Weak covalent forces between molecules
Sulfur, S8
Phosphorus, P4
27
Bonding in Solids

• 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).

28
Intermolecular Forces
Forces of attraction between different molecules
rather than bonding forces within the same
molecule.

• Dipole-dipole attraction
• Hydrogen bonds
• Dispersion forces

29
Intermolecular Forces
Hydrogen Bonding
30
Intermolecular Forces
Dipole-Dipole Forces
31
Intermolecular Forces

• London Dispersion Forces
• One instantaneous dipole can induce another
  instantaneous dipole in an adjacent molecule (or
  atom).
• The forces between instantaneous dipoles are
  called London dispersion forces.

32
Intermolecular Forces
London Dispersion Forces
33
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

34
Bonding in Solids
35
Classification of Matter
Solutions are homogeneous mixtures
36
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
37
Dissolution of sodium Chloride
38
Concentrated vs. Dilute
39
Some Properties of a Liquid

• Surface Tension The resistance to an increase
  in its surface area (polar molecules, liquid
  metals).

• Capillary Action Spontaneous rising of a liquid
  in a narrow tube.

40
Surface Tension
41
Some Properties of a Liquid

• Viscosity Resistance to flow

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