1. Metallic and Ionic Solids Sections 13.6-8
To play the movies and simulations included, view the presentation in Slide Show Mode.

2. Types of Solids Table 13.6 TYPE EXAMPLE FORCE
Ionic NaCl, CaF2, ZnS Ion-ion
Metallic Na, Fe Metallic
Molecular Ice, I2 Dipole Ind. dipole
Network Diamond Extended Graphite covalent

3. Network Solids
Diamond
Graphite

4. Network Solids
A comparison of diamond (pure carbon) with silicon.

5. Properties of Solids
1. Molecules, atoms or ions locked into a CRYSTAL LATTICE
2. Particles are CLOSE together
3. STRONG IM forces
4. Highly ordered, rigid, incompressible
ZnS, zinc sulfide

6. Crystal Lattices
Regular 3-D arrangements of equivalent LATTICE POINTS in space.
Lattice points define UNIT CELLS
smallest repeating internal unit that has the symmetry characteristic of the solid.

7. Cubic Unit Cells
There are 7 basic crystal systems, but we are only concerned with CUBIC.
All sides
equal length
All angles
are 90 degrees

8. Cubic Unit Cells of Metals Figure 13.24
Simple cubic (SC)
Body-centered cubic (BCC)
Face-centered cubic (FCC)

9. Units Cells for Metals
Figure 13.25

10. Simple Cubic Unit Cell Figure 13.28
Note that each atom is at a corner of a unit cell and is shared among 8 unit cells.

11. Body-Centered Cubic Unit Cell

12. Face Centered Cubic Unit Cell
Atom at each cube corner plus atom in each cube face.

13. Atom Packing in Unit Cells
Assume atoms are hard spheres and that crystals are built by PACKING of these spheres as efficiently as possible.

14. Atom Packing in Unit Cells

15. Crystal Lattices— Packing of Atoms or Ions
FCC is more efficient than either BC or SC.
Leads to layers of atoms.
See Closer Look, pp. 541

16. Packing of Atoms or Ions
Crystal Lattices—
Packing of Atoms or Ions
Packing of C60 molecules. They are arranged at the lattice points of a FCC lattice.

17. Number of Atoms per Unit Cell
Unit Cell Type Net Number Atoms SC
BCC
FCC 1 2 4

18. Atom Sharing at Cube Faces and Corners
Atom shared in corner
--> 1/8 inside each unit cell
Atom shared in face
--> 1/2 inside each unit cell

19. Simple Ionic Compounds
Lattices of many simple ionic solids are built by taking a SC or FCC lattice of ions of one type and placing ions of opposite charge in the holes in the lattice.
EXAMPLE: CsCl has a SC lattice of Cs+ ions with Cl- in the center.

20. Simple Ionic Compounds
CsCl has a SC lattice of Cs+ ions with Cl- in the center.
1 unit cell has 1 Cl- ion plus
(8 corners)(1/8 Cs+ per corner)
= 1 net Cs+ ion.

21. Simple Ionic Compounds
Salts with formula MX can have SC structure — but not salts with formula MX2 or M2X

22. Two Views of CsCl Unit Cell
Either arrangement leads to formula of 1 Cs+ and 1 Cl- per unit cell

23. NaCl Construction Na+ in octahedral holes
FCC lattice of Cl- with Na+ in holes

24. Octahedral Holes in FCC Lattice

25. The Sodium Chloride Lattice
Na+ ions are in OCTAHEDRAL holes in a face-centered cubic lattice of Cl- ions.

26. The Sodium Chloride Lattice
Many common salts have FCC arrangements of anions with cations in OCTAHEDRAL HOLES — e.g., salts such as CA = NaCl
• FCC lattice of anions ----> 4 A-/unit cell
• C+ in octahedral holes ---> 1 C+ at center
+ [12 edges • 1/4 C+ per edge]
= 4 C+ per unit cell

27. Comparing NaCl and CsCl
Even though their formulas have one cation and one anion, the lattices of CsCl and NaCl are different.
The different lattices arise from the fact that a Cs+ ion is much larger than a Na+ ion.

28. Common Ionic Solids Titanium dioxide, TiO2
There are 2 net Ti4+ ions and 4 net O2- ions per unit cell.

29. Common Ionic Solids Zinc sulfide, ZnS
The S2- ions are in TETRAHEDRAL holes in the Zn2+ FCC lattice.
This gives 4 net Zn2+ ions and 4 net S2- ions.

30. Common Ionic Solids Fluorite or CaF2 FCC lattice of Ca2+ ions
This gives 4 net Ca2+ ions.
F- ions in all 8 tetrahedral holes.
This gives 8 net F- ions.

31. Phase Diagrams

32. TRANSITIONS BETWEEN PHASES Section 13.10
Lines connect all conditions of T and P where EQUILIBRIUM exists between the phases on either side of the line.
(At equilibrium particles move from liquid to gas as fast as they move from gas to liquid, for example.)

33. Phase Diagram for Water
Liquid phase
Solid phase
Gas phase

34. Phase Equilibria — Water
Solid-liquid
Gas-Liquid
Gas-Solid

35. Triple Point — Water
At the TRIPLE POINT all three phases are in equilibrium.

36. Phases Diagrams— Important Points for Water
T(˚C) P(mmHg)
Normal boil point
Normal freeze point
Triple point

37. Critical T and P
As P and T increase, you finally reach the CRITICAL T and P
Above critical T no liquid exists no matter how high the pressure.

38. Critical T and P
COMPD Tc(oC) Pc(atm)
H2O CO CH
Freon (CCl2F2)
Notice that Tc and Pc depend on intermolecular forces.

39. Solid-Liquid Equilibria
In any system, if you increase P the DENSITY will go up.
Therefore — as P goes up, equilibrium favors phase with the larger density (or SMALLER volume/gram).
Liquid H2O Solid H2O
Density 1 g/cm g/cm3
cm3/gram

40. Solid-Liquid Equilibria
Raising the pressure at constant T causes water to melt.
The NEGATIVE SLOPE of the S/L line is unique to H2O. Almost everything else has positive slope.

41. Solid-Liquid Equilibria
The behavior of water under pressure is an example of
LE CHATELIER’S PRINCIPLE
At Solid/Liquid equilibrium, raising P squeezes the solid.
It responds by going to phase with greater density, i.e., the liquid phase.
Solid-Liquid Equilibria

42. Solid-Vapor Equilibria
At P < 4.58 mmHg and T < ˚C
solid H2O can go directly to vapor. This process is called SUBLIMATION
This is how a frost-free refrigerator works.