1. Liquids and Solids

2. Relative Magnitudes of Forces The types of bonding forces vary in their strength as measured by average bond energy. Covalent bonds (400 kcal/mol) Hydrogen bonding (12-16 kcal/mol ) Dipole-dipole interactions (2-0.5 kcal/mol) London forces (less than 1 kcal/mol) Strongest Weakest

3. Hydrogen Bonding Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen Hydrogen bonding between ammonia and water

4. Hydrogen Bonding in DNA TA Thymine hydrogen bonds to Adenine

5. Hydrogen Bonding in DNA CG Cytosine hydrogen bonds to Guanine

6. London Dispersion Forces The temporary separations of charge that lead to the London force attractions are what attract one nonpolar molecule to its neighbors. Fritz London 1900-1954 London forces increase with the size of the molecules.

7. London Dispersion Forces

8. London Forces in Hydrocarbons

9. Boiling point as a measure of intermolecular attractive forces

10. 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.

11. Some Properties of a Liquid  Viscosity: Resistance to flow  High viscosity is an indication of strong indication of strong intermolecular forces intermolecular forces

12. Types of Solids  Crystalline Solids: highly regular arrangement of their components

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

14. 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.

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

16. Closest Packing Holes

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

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

19. Bonds Between Atoms Covalent Ionic Molecular Substance Like dissolves like Network Solids Metallic nonpolar Polar - ions held together by electrostatic attraction -conducts electricity only in solution or in the molten state -forms as solids -If comparing ionic compounds argue charge. If charges are the same argue atomic radius. (Ex. MgCl 2, NaF NaCl IMF – LD, maybe HB -As MM goes up LD forces become stronger IMF – LD, DD, DD could be HB If two polar molecules do not contain HB, argue MM goes up so LD are stronger -highest melting point -chemical bonds broken when melted -does not dissolve in most solvents Ex. SiO 2 (sand), Quartz, C (diamond), C (graphite) -sea of mobile electrons -conducts electricity in solid and liquid state -melting points vary from high to low

20. Network Covalent Atomic Solids Some covalently bonded substances DO NOT form discrete molecules. Diamond, a network of covalently bonded carbon atoms Graphite, a network of covalently bonded carbon atoms

21. Covalent Network Structures A covalent network is when atoms are bonded together covalently to form giant macromolecular structures. The element Carbon can form Covalent Network structures Two examples are: Diamond, Graphite, & silicon dioxide (sand)

22. Diamond Each Carbon atom bonds to 4 others Strong covalent bonds throughout High melting points and Boiling points No free electrons so does not conduct electricity.

23. Graphite Each Carbon atom is bonded to 3 others The bonding within the layers is strong covalent. The spare fourth electron of each C atom is delocalised so graphite can conduct electricity There are weak intermolecular forces of attraction between the layers.

24. Covalent Networks Another example of a covalent network is silicon dioxide. This is a compound though it resembles diamond.

25. Molecular Solids Strong covalent forces within molecules Weak covalent forces between molecules Sulfur, S 8 Phosphorus, P 4

26. Properties of Ionic Compounds solid crystalsUsually solid crystals @ room temp (types of crystal structure & coordination number) Generally have high melting points conduct electricityCan conduct electricity when melted or dissolved in water ionsFormula represents number of ions in a representative unit

27. Metallic Bonds cationsMetals are made up of closely packed cations rather than neutral atoms (sea of electrons)Valence electrons are mobile (sea of electrons) free- floating valence electronsMetallic bonds consist of attractions of the free- floating valence electrons for the positively charged metal ions Good conductors of electricity as solids or liquids

28. Metallic Properties & Structures: Good conductors of electricityGood conductors of electricity (ductile)Can be drawn into wires (ductile) (malleable)Can be hammered or forced into shapes (malleable) Arranged in very compact & orderly patternsArranged in very compact & orderly patterns AlloysAlloys are a mixture of 2 or more elements, at least one of which is a metal

29. Forces in solids:  Metallic: metallic bonds (metal ions in electron sea)  Ionic: electrostatic (anions and cations)  Network: covalent bonds  Molecular: London, dipole-dipole, hydrogen bonds (all crystalline!)

30. Place the following from highest melting point to Lowest. Explain each choice and if any two are the Same type of substance explain why one was chosen Over the other. H 2 O Xe Ne NH 3 SO 2 PCl 3 SiO 2 MgCl 2 NaCl

31. Water phase changes Temperature remains __________ during a phase change. constant Energy

32. 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.

36. Calculating Vapor Pressure p. 487

37. LoctionFeet Above Sea Level P (torr) Boiling Point (˚C) Top Mt. Everest 29,02824070 Leadville CO10,15051089 Madison, WI 90073094 New York City, N.Y. 10760100 Death Valley, CA -282770100.3

39. 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, T c = 374°C and 218 atm).

40. Phase changes by Name

41. Water

42. Explain How Pressure affects Melting Points.

44. Carbon dioxide

45. Carbon

46. Sulfur