1. Intermolecular Forces, Liquids, and Solids Chapter 11 – Brown & LeMay

2. Temperature Review Measure of kinetic energy What can you say about the KE of salt particles, water molecules, and oxygen particles at room temperature? State determined by strength of forces that keep particles together

3. Strength Compare energy needed for phase change vs. decomposition in HCl(l) Intermolecular (called weak) because they are weaker than ionic or covalent Boiling point reflects strength of bonds in liquid Melting point reflects strength of bonds in solids

4. Kinds of Intermolecular Forces Three major kinds: dipole-dipole, London dispersion, and hydrogen bonding In solutions, ion-dipole All are electrostatic in nature Approximately 15% of covalent or ionic strength

5. Ion - dipole When? Ionic solid + polar liquid Increases with increasing charge of ion or polarity of solvent Determines solubility

6. Dipole-Dipole forces Weaker than previous + end of one attracts – end of another If size is equal, more polar has stronger dipole attractions. (NH 3 vs H 2 O) If polarity is the same but masses differ, than smallest is stronger. (Able to orient better)

7. London Dispersion Forces All molecules have this Only attraction in nonpolar molecules How can Iodine be a solid? Temporary lopsided charge builds up from random motion of electrons - 1930 Increases with mass – we say it has greater polarizability Straight molecule is more polarizable than a curled up molecule – why? Halogen Family is a great essay

8. Hydrogen Bond Strongest of all “weak” forces Is caused when H is bonded to F, O, or N These are so electronegative that the H is a “naked nucleus” or bare proton Very attractive! Will bond to nearby electron pairs

9. Importance of Hydrogen Bonding Biological systems – DNA, proteins Water chemistry (MP, BP, specific heat, surface tension) Density of ice

10. Density Most solids are more dense than liquid Water is less dense because of hydrogen bonding At 4°C, water becomes less dense Important for life in winter Causes lake turnover Alum example

11. Practice Look at Flow Chart

12. Properties of Liquids Viscosity “Slower than….. Resistance of a liquid to flow Time it as it goes through a small tube with gravity acting upon it. Poise – 1g/cm-s Trends – same substance – decreases with increasing temperature series (same structure) – increases with increasing mass

13. Surface Tension How many drops on a penny? Uneven forces at surface Acts like pond scum Definition – energy needed to increase the surface area of a liquid by a certain amount  Water is high – why? Called “cohesive” force – together Water moving up a stem – adhesive force Capillary acion – rise up a thin tube Meniscus!

14. Phase Changes Solid to Liquid is called Heat of Fusion  H fus For water, 6 kJ/mol Liquid to Gas is called Heat of Vaporization  H vap For water, 40.7 kJ/mol  H sub is sum of each

15. Heating Curve Try a problem Remember - flat during phase change, temperature change when heating a single phase Cooling is opposite

16. Supercooling Happens with some liquids - remove heat and it doesn’t freeze when it should Very unstable May happen during hibernation

17. Critical Temperature Highest temperature at which a liquid can form from a gas when pressure is applied. Above this, the substance is called a supercritical fluid. Gas just becomes more compressed. Critical pressure - pressure at the critical temperature

18. Vapor pressure Vapor pressure forms above any liquid if container is closed – why? Equilibrium is reached This is vapor pressure Higher if forces holding liquid together are weak - called a volatile (fleeing) liquid

19. Boiling Point Temperature at which the VP equals atmospheric pressure Normal BP - boiling point at 1 atm Everest? Autoclave?

20. Phase Diagram Handout Look at lines Look at slope of AB Freeze-drying - library book example

21. Water vs. CO 2

22. Structure of Solids Amorphous (rubber, plastics) - large or mixtures - no true structure Crystalline - highly ordered structure Crystalline solids have true melting points

23. Unit Cell Repeating unit of a solid 7 types – (6-sided parallelograms) Ni, Na, NaCl Array of points in the crystal lattice

24. 3 cubic unit cells

25. Total Atoms for each unit cell

26. Packing Spheres naturally pack hexagonally Animation

27. Bonding Shown by x-ray diffraction Molecular - low MP If unit packs well, mp can be high Covalent Network Solid - very strong Many covalent bonds in 3-D Diamond, graphite, SiO 2, SiC, BN Ionic - greater charge, greater MP Metallic solids - hexagonal close packed, mp varies