1. Liquids & Solids
Chapter 10

2. Heat of Fusion/Vaporization
H2O(s) ----> H2O(l) Hfo = 6.02 kj/mol
H2O(l) ----> H2O(g) Hvo = 40.7 kj/mol
From the Ho values above, which two states are most similar?
How do the attractive forces between the molecules compare in these two states to the third state?

3. Three States of Matter

4. Types of Bonding Intermolecular Intramolecular between molecules
dipole-dipole forces
hydrogen bonding
London Dispersion Forces
Intramolecular
within the molecule
covalent bonding
ionic bonding
When ice changes to liquid and then to vapor, the
intramolecular forces (covalent bonds) stay intact, only
the weaker hydrogen bonds between molecules weaken
and break.

5. Intermolecular Forces
Forces between (rather than within) molecules.
dipole-dipole attraction: molecules with dipoles orient themselves so that “+” and “” ends of the dipoles are close to each other. (1 % as strong as covalent or ionic.)
hydrogen bonds: dipole-dipole attraction in which hydrogen is bound to a highly electronegative atom. (F, O, N)

6. Electrostatic interaction of two polar molecules.

7. The polar water molecule and hydrogen bonds
among water molecules.

8. The boiling points of the covalent hydrides of the
elements in Groups 4A, 5A, 6A, & 7A.

9. Instantaneous and induced dipole moments
between nonpolar molecules -- London
Dispersion Forces.

10. London Dispersion Forces
relatively weak forces that exist among noble gas atoms and nonpolar molecules. (Ar, C8H18)
caused by instantaneous dipole, in which electron distribution becomes asymmetrical.
the ease with which electron “cloud” of an atom can be distorted is called polarizability.

11. Some Properties of a Liquid
Surface Tension: The resistance to an increase in its surface area (polar molecules). A sphere has the maximum volume for the minimum surface area.

12. Some Properties of a Liquid
Capillary Action: Spontaneous rising of a liquid in a narrow tube.
Viscosity: Resistance to flow (molecules with
large intermolecular forces).

13. Some Properties of a Liquid
Cohesive forces exist between molecules of a liquid. Adhesive forces exist between the liquid and its container.

14. Types of Solids
Crystalline Solids: highly regular arrangement of their components [table salt (NaCl), pyrite (FeS2)].
Amorphous solids: considerable disorder in their structures (glass).

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. Representation of Components in a Crystalline Solid
Unit Cell: The smallest repeating unit of the lattice.
simple cubic -- 1 atom/cell
body-centered cubic -- 2 atoms/cell
face-centered cubic -- 4 atoms/cell

17. Three cubic unit cells and the corresponding
lattices.

18. Simple Cubic Cell
1 atom per cell
side length (do) = 2 r
do = 2 r

19. Body-Centered Cell 2 atoms per cell Body diagonal = do 3 = 4r
do2 -- diagonal through the base of cube.
4r do
do2

20. Face-Centered Cell 4 atoms per cell Face diagonal = do 2 = 4r
do2 -- diagonal through the face of cube.
4r do do

21. Face-centered cubic unit cell.

22. Bragg Equation
Used for analysis of crystal structures and to calculate the distance between planes in crystals.
n = 2d sin 
d = distance between atoms
n = an integer
 = wavelength of the x-rays

23. Reinforcement or cancellation of X-rays.

24. Reflection of X-rays of wavelength  from a pair
of atoms in two different layers of a crystal.

25. Types of Crystalline Solids
Atomic Solid: contains atoms at the lattice points (diamond).
Ionic Solid: contains ions at the points of the lattice that describe the structure of the solid (NaCl).
Molecular Solid: discrete covalently bonded molecules at each of its lattice points (sucrose, ice).

26. Three crystalline solids -- a) atomic solid, b) ionic
solid, and c) molecular solid.

28. Packing in Metals
Model: Packing uniform, hard spheres to best use available space. This is called closest packing. Each atom has 12 nearest neighbors.
hexagonal closest packed (“aba”)
cubic closest packed (“abc”)

29. Closest packing arrangement of uniform spheres --
aba. This forms hexagonal closest packed -- hcp.

30. Atoms arranged in aba pattern forming hexagonal
10_218
(a)
(b)
(a))
Top view
Atom in third layer
lies over atom in
first layer.
Atoms arranged in aba pattern forming hexagonal
closest packed (hcp) structure -- 2 atoms/cell.

31. Hexagonal closest packed structure -- central
atom has 12 nearest neighbors.

32. Face-centered cubic is cubic closest packed
(ccp). The spheres are packed in an abc
arrangement.

33. Bonding Models for Metals
Electron Sea Model: A regular array of metals in a “sea” of electrons.
Band (Molecular Orbital) Model: Electrons assumed to travel around metal crystal in MOs formed from valence atomic orbitals of metal atoms.
Conduction Bands: closely spaced empty molecular orbitals allow conductivity of heat and electricity.

34. Representation of the energy levels (bands) in a
magnesium crystal. 1s, 2s, & 2p orbitals are
localized, but 3s & 3p orbitals are delocalized to
make molecular orbitals.

35. Metal Alloys
Substances that have a mixture of elements and metallic properties.
1. Substitutional Alloy: some metal atoms replaced by others of similar size.
brass = Cu/Zn

36. Metal Alloys (continued)
2. Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms.
steel = iron + carbon
3. Both types: Alloy steels contain a mix of substitutional (Cr, Mo) and interstitial (Carbon) alloys.

37. Substitutional
Alloy
Interstitial Alloy

38. graphite, diamond, ceramics, glass
Network Solids
Composed of strong directional covalent bonds that are best viewed as a “giant molecule”.
brittle
do not conduct heat or electricity
carbon, silicon-based
graphite, diamond, ceramics, glass

39. Network solid structure of diamond.

40. Semiconductors
A substance in which some electrons can cross the band gap.
Conductivity is enhanced by doping with group 3a or group 5a elements.
n-type semiconductor -- doped with atoms having more valence electrons -- Phosphorus.
p-type semiconductor -- doped with atoms having fewer valence electrons -- Boron.
See Figure on page 477 in Zumdahl.

41. Molecular Solids molecular units at each lattice position.
strong covalent bonding within molecules.
relatively weak forces between molecules.
London Dispersion Forces -- CO2, I2, P4, & S8.
Hydrogen Bonding -- H2O, NH3, & HF.

43. Trigonal, Tetrahedral, & Octahedral Holes
Trigonal holes -- formed by three spheres in the same layer.
Tetrahedral holes -- formed when a sphere sits in the dimple of three spheres in an adjacent layer.
Octahedral holes -- formed between two sets of spheres in adjoining layers of closest packed structures.

44. Trigonal, Tetrahedral, and Octahedral holes.

45. Hexagonal & Cubic Closest Packed
1 octahedral hole for each atom or ion.
2 tetrahedral holes for each atom or ion.
Simple cubic and body-centered cubic are not closest packed structures!

46. The location (x) of a tetrahedral hole in the face-
centered cubic unit cell. The S2- ions are
closest packed with the Zn2+ ions in alternating
tetrahedral holes.

47. The location (x) of an octahedral hole in the face-
centered cubic unit cell. The Cl- ions have a ccp
arrangement with the Na+ ions in all the
octahedral holes.

48. Volatile liquids have high vapor pressures.
. . . is the pressure of the vapor present at equilibrium.
. . . is determined principally by the size of the intermolecular forces in the liquid.
. . . increases significantly with temperature.
Volatile liquids have high vapor pressures.

49. Vapor Pressure Low boiling point high vapor pressure.
weak intermolecular forces.
Low vapor pressure
high molar masses.
strong intermolecular forces.

50. Boltzman Distribution -- number of molecules in
a liquid with a given energy versus kinetic energy
at two different temperatures.

51. Natural Log of Vapor Pressure Versus Reciprocal Kelvin Temperature
y = m x + b
Slope = If the slope is known, then H can be calculated.

52. Clausius-Clayperon Equation
Temperatures must be expressed in Kelvin.
See Example 10.6 on page 488 in Zumdahl.

53. Sublimation
Change of a solid directly to a vapor without passing through the liquid state.
Iodine
Dry Ice
Moth Balls

54. vapor pressure of solid = vapor pressure of liquid
Melting Point
Molecules break loose from lattice points and solid changes to liquid. (Temperature is constant as melting occurs.)
vapor pressure of solid = vapor pressure of liquid

55. Boiling Point
Constant temperature when added energy is used to vaporize the liquid.
vapor pressure of liquid = pressure of
surrounding atmosphere

56. Phase Diagram
Represents phases as a function of temperature and pressure.
critical temperature: temperature above which the vapor can not be liquefied.
critical pressure: pressure required to liquefy AT the critical temperature.
critical point: critical temperature and pressure (for water, Tc = 374°C and 218 atm).

57. Heating curve for water.
H = (ms t)ice + m Hf + (ms t) water + m Hv + (mst)steam
E = KE & PE + PE KE & PE PE PE & KE

58. Solid and liquid water interact only through the
vapor state.

59. Phase diagram for water -- Tm is the regular melting
point. The solid/liquid line has a negative slope.

60. Phase diagram for carbon dioxide -- the solid/
liquid line has a positive slope.

61. Phase diagram for sulfur -- note the two different
solid forms of rhombic and monoclinic sulfur.

62. Phase diagram for carbon -- note the two solid
forms of diamond and graphite.