1. CHAPTER 12
Liquids and Solids

2. Intermolecular Forces
Forces of attraction between different molecules rather than bonding forces within the same molecule.
Dipole-dipole attraction
Hydrogen bonds
Dispersion forces

3. Dipole-Dipole Attraction
dipole-dipole attraction: molecules with dipoles orient themselves so that “+” and “-” ends of the dipoles are close to each other.

4. Dipole Forces

5. Hydrogen Bonding
hydrogen bonds: dipole-dipole attraction in which hydrogen is bound to a highly electronegative atom. (F, O, N)

6. Hydrogen Bonding in Water

7. Hydrogen Bonding in DNA

8. London Dispersion Forces
Dispersion forces: relatively weak forces caused by instantaneous dipole, in which electron distribution becomes asymmetrical.
The ONLY forces of attraction that exist among noble gas atoms and nonpolar molecules. (Ar, C8H18)

9. London Dispersion Forces

10. Boiling point as a measure of intermolecular attractive forces

11. Some Properties of a Liquid
Surface Tension: The resistance to an increase in its surface area (polar molecules, liquid metals).

12. Some Properties of a Liquid
Capillary Action: Spontaneous rising of a liquid in a narrow tube.

13. Some Properties of a Liquid
Viscosity: Resistance to flow (molecules with large intermolecular forces).

14. Types of Solids
Crystalline Solids: highly regular arrangement of their components [table salt (NaCl), pyrite (FeS2)].

15. Types of Solids
Amorphous solids: considerable disorder in their structures (glass).

16. 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.

17. Types of Crystalline Solids
Ionic Solid: contains ions at the points of the lattice that describe the structure of the solid (NaCl).

18. Types of Crystalline Solids
Molecular Solid: discrete covalently bonded molecules at each of its lattice points (sucrose, ice).

19. Packing in Metals
Model: Packing uniform, hard spheres to best use available space. This is called closest packing. Each atom has 12 nearest neighbors.

20. Closest Packing Holes

21. Metal Alloys
Substitutional Alloy: some metal atoms replaced by others of similar size.
brass = Cu/Zn

22. Metal Alloys (continued)
Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms.
steel = iron + carbon

23. 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

24. Sulfur – S8

25. Phosphorus – P4

26. Diamond

27. Graphite

28. Zirconia

30. Equilibrium Vapor Pressure
The pressure of the vapor present at equilibrium.
Determined principally by the size of the intermolecular forces in the liquid.
Increases significantly with temperature.
Volatile liquids have high vapor pressures.

31. LeChatelier’s Principle
When a system at equilibrium is placed under stress, the system will undergo a change in such a way as to relieve that stress.

32. Translation:
When you take something away from a system at equilibrium, the system shifts in such a way as to replace what you’ve taken away.
When you add something to a system at equilibrium, the system shifts in such a way as to use up what you’ve added.

33. LeChatelier’s Example #1
A closed container of ice and water at equilibrium. The temperature is raised.
Ice + Energy  Water
The equilibrium of the system shifts to the _______ to use up the added energy.
right

34. LeChatelier’s Example #2
A closed container of N2O4 and NO2 at equilibrium. NO2 is added to the container.
N2O4 + Energy  2 NO2
The equilibrium of the system shifts to the _______ to use up the added NO2.
left

35. LeChatelier’s Example #3
A closed container of water and its vapor at equilibrium. Vapor is removed from the system.
water + Energy  vapor
The equilibrium of the system shifts to the _______ to make more vapor.
right

36. Water phase changes constant Temperature remains __________
during a phase change.
constant
Water phase changes

37. Kinetic Energy

38. 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).

39. Water
Water

40. Phase changes by Name

41. Carbon
dioxide
Carbon dioxide

42. Carbon
Carbon

43. Sulfur