1. Liquids and Solids KMT of Liquids and Solids, Phase Diagram, Vapor Pressure Curve, Heating/Cooling Curve

2. Intermolecular Forces: Liquids, Solids, and Phase Changes 1.Types of Intermolecular Forces 2.Properties of liquids and solids 3.Phase change diagrams 4.Heating/cooling curve 5.Vapor pressure curve

3. Brainteaser!!!!  If substances at the same temperature have the same kinetic energy, why are they all not liquids, solids, or gases?

4. Intermolecular Forces  Inter molecular Forces are attractive forces between molecules. Think interstate!  Intra molecular Forces are attractive forces that hold molecules together  Inter vs. Intra  41 kJ to vaporize 1 mole of H 2 O  930 kJ to break all O-H bonds in one H 2 O molecule Which one is stronger????? Intramolecular forces are stronger than intermolecular forces!!!!

5. Dipole – Dipole Forces  Between polar molecules  What bond is the strongest?  Where is the intermolecular bond?  Dipole – molecule with a completely separate positively and negatively charged end

6. Ion – Dipole Forces  Between polar molecules and ions  Give me an example of an everyday solution between polar molecules and ions!!!!!!  Why are dipoles attracted to ions?

7. London-Dispersion Forces  Intermolecular forces are formed by temporarily induced dipole moments  How do dipoles become induced?  Electron clouds constantly move and when one molecule collides with another molecule the electrons are temporarily shifted to one side  This creates a momentary negative end and a positive end  Usually occurs between identical molecules (Example H 2 (g)

8. Hydrogen bonds  Force formed between molecules containing N–H, O–H, or F–H groups, and an electronegative O, N, or F atom.  10% of the energy in a covalent bond!!!!!!

9. Hydrogen Bonding H 2 O CH 3 OH NH 3

10. Phases of matter  Gases – molecules are widely separated and the “fluid” is compressible  Liquids – molecules are more tightly packed and liquids are relatively incompressible  Solids – molecules are tightly packed and solids are incompressible and rigid

11. Liquids  IMF’s limit the range of motion of particles in a liquid  Density – Liquids have a higher density at 25 °C than gases  Fluidity – Ability to flow  Viscosity – Measure of the resistance of a liquid to flow  Surface tension – The energy required to increase the surface area of a liquid by a given amount

12. Viscosity  Measure of a liquids resistance to flow  Inversely related to the size of the molecule and the type and strength of intermolecular forces  The higher the temperature the lower the viscosity  If temperature then viscosity  Here’s the tricky part:  If temperature then the liquid starts to flow

13. Surface Tension  The energy required to increase the surface area of a liquid by a given amount  Molecules in the center of a liquid are exposed to IMF from all sides  Molecules on the surface of a liquid are not exposed to IMF from all sides  In order to increase the surface area of a liquid the molecules in the interior of the liquid must move to the surface and the IMF’s must be broken

14. Capillary Action  Water molecules “cling” to the surface of the graduated cylinder by adhesion  Adhesion is the force of attraction between different types of molecules  Cohesion is the force of attraction between the same type of molecules  What force must be strongest for water to cling to the glass tube?  If adhesion forces are stronger than cohesion forces water will be drawn up the sides of the cylinder

15. Solids  Tightly packed molecules that are rigid and cannot be compressed  Density is highest in solids (except in water!!!)  Crystalline solid – solid whose atoms, ions, or molecules are arranged in an orderly, geometric, 3-D structure  Amorphous – atoms are randomly arranged because they typically cool too quickly. No order exists in the solid.

16. Types of Solids Crystalline – a well defined arrangement of atoms; this arrangement is often seen on a macroscopic level. (p.402) Atomic solidsAtomic solids Ionic solidsIonic solids Molecular solidMolecular solid Covalent networkCovalent network MetallicMetallic Units points that can be repeated in three dimensions to form a lattice

17. Phase Changes  Melting – the change from a solid to a liquid Melting Point – T at which forces holding lattice together are broken Melting Point – T at which forces holding lattice together are broken  Vaporization- the change from a liquid to a gas  Sublimation – the change from a solid to a gas  Condensation – the change from a gas to a liquid  Deposition – the change from a gas to a solid  Freezing – the change from a liquid to a solid GAS SOLIDLIQUID MELTING FREEZING CONDENSATION VAPORIZATION SUBLIMATION DEPOSITION

18. Phase Change Diagrams  Relationship between T and P  Triple point – P and T at which substance can coexists as a gas, liquid, and solid  Critical point – T at which a substance can no longer remain a liquid regardless of the pressure Look at the liquid solid line and its slope!!!!!

19. Phase Diagram for H 2 O  What is the difference between this diagram and the first?  The liquid solid line leans backwards! Normal Melting and boiling points Vapor pressure curve

20. Vapor Pressure  In a sealed container some water ( l ) changes phase to become water vapor and exerts a pressure over the surface of the liquid (if the container were open it would be considered partial pressure)

21. Heating Cooling Curve

22. 120 °C steam 100 °C water  steam 50°C liquid water 50°C liquid water 0 °C ice  liquid 0 °C ice  liquid -10 °C ice -10 °C ice Heat added  Heat added  Why does temperature “stand still”? Heating and cooling curve for H 2 O What bonds are broken?