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Solids

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Solids We can think of solids as falling into two groups: Crystalline particles are in highly ordered arrangement. Solids Amorphous no particular order in the ... – PowerPoint PPT presentation

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Title: Solids


1
Solids
  • We can think of solids as falling into two
    groups
  • Crystallineparticles are in highly ordered
    arrangement.

2
Solids
  • Amorphousno particular order in the arrangement
    of particles.

3
Attractions in Ionic Crystals
  • In ionic crystals, ions pack themselves so as to
    maximize the attractions and minimize repulsions
    between the ions.

4
Crystalline Solids
  • Because of the order in a crystal, we can focus
    on the repeating pattern of arrangement called
    the unit cell.

5
Crystalline Solids
  • There are several types of basic arrangements in
    crystals, such as the ones shown above.

6
Crystalline Solids
  • We can determine the empirical formula of an
    ionic solid by determining how many ions of each
    element fall within the unit cell.

7
Ionic Solids
  • What are the empirical formulas for these
    compounds?
  • (a) Green chlorine Gray cesium
  • (b) Yellow sulfur Gray zinc
  • (c) Green calcium Gray fluorine

(a)
(b)
(c)
CsCl
ZnS
CaF2
8
Types of Bonding in Crystalline Solids
9
Covalent-Network andMolecular Solids
  • Diamonds are an example of a covalent-network
    solid in which atoms are covalently bonded to
    each other.
  • They tend to be hard and have high melting points.

10
Covalent-Network andMolecular Solids
  • Graphite is an example of a molecular solid in
    which atoms are held together with van der Waals
    forces.
  • They tend to be softer and have lower melting
    points.

11
Metallic Solids
  • Metals are not covalently bonded, but the
    attractions between atoms are too strong to be
    van der Waals forces.
  • In metals, valence electrons are delocalized
    throughout the solid.

12
Metallic Bonding
"SEA OF MOBILE VALENCE ELECTRONS" The animation
you are looking at attempts to help you
understand the nature of a metal bond. The gray
spheres represent metal cations (positively
charged ions), and the red moving spheres
represent electrons.  Metals have low ionization
energies, thus they do not have a tight hold on
their valence electrons.  These outer electrons
easily move around, as they do not "belong" to
any one atom, but are part of the whole metal
crystal.  The negatively charged electrons act as
a "cement" that hold the positively charged metal
ions in their relatively fixed positions.
13
The fact that the electrons flow easily helps to
explain some of the characteristics of metals -
Metals are good conductors of heat and
electricity. This is directly due to the mobility
of the electrons. - The "cement" effect of the
electrons determines the hardness of the metal. 
Some metals are harder than others the strength
of the "cement" varies from metal to metal. - 
Metals are lustrous.  This is due to the uniform
way that the valence electrons of the metal
absorb and re-emit light energy. - Metals are
malleable (can be flattened) and ductile (can be
drawn into wires) because of the way the metal
cations and electrons can "flow" around each
other, without breaking the crystal
structure. Metallic bonds are best characterized
by the phrase "a sea of electrons"
14
  • Metallic Bonding
  • atoms in metals are packed very closely in an
    orderly arrangement
  • each atom loses its valence electrons to become a
    positive ion

15
Phase Diagrams Equilibrium can exist not only
between the liquid and vapor phase of a substance
but also between the solid and liquid phases, and
the solid and gas phases of a substance. A phase
diagram is a graphical way to depict the effects
of pressure and temperature on the phase of a
substance                                     
                                                  
          
16
Phase Diagrams
  • Phase diagrams display the state of a substance
    at various pressures and temperatures and the
    places where equilibria exist between phases.

17
Phase Diagrams
  • The AB line is the liquid-vapor interface.
  • It starts at the triple point (A), the point at
    which all three states are in equilibrium.

18
Phase Diagrams
  • It ends at the critical point (B) above this
    critical temperature and critical pressure the
    liquid and vapor are indistinguishable from each
    other.

19
Phase Diagrams
  • Each point along this line is the boiling point
    of the substance at that pressure.

20
Phase Diagrams
  • The AD line is the interface between liquid and
    solid.
  • The melting point at each pressure can be found
    along this line.

21
Phase Diagrams
  • Below A the substance cannot exist in the liquid
    state.
  • Along the AC line the solid and gas phases are in
    equilibrium the sublimation point at each
    pressure is along this line.

22
Phase Diagram of Water
  • Note the high critical temperature and critical
    pressure
  • These are due to the strong van der Waals forces
    between water molecules.

23
Phase Diagram of Water
  • The slope of the solidliquid line is negative.
  • This means that as the pressure is increased at a
    temperature just below the melting point, water
    goes from a solid to a liquid.

24
Phase Diagram of Carbon Dioxide
  • Carbon dioxide cannot exist in the liquid state
    at pressures below 5.11 atm CO2 sublimes at
    normal pressures.

25
Phase Diagram of Carbon Dioxide
  • The low critical temperature and critical
    pressure for CO2 make supercritical CO2 a good
    solvent for extracting nonpolar substances (such
    as caffeine).
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