1. States of Matter and Intermolecular Forces Chapter 11 11-1 States and State Changes

2. Solids Particles have an orderly, fixed arrangement Fixed volumes and shapes

3. Liquids Particles move easily past one another (have more energy) Fixed volume, no fixed shape

4. Viscosity Ability to Flow Honey is very viscous

5. Surface Wetting Adhesion Stick to something else Cohesion Stick to each other

6. Capillary Action The movement of water up through a tube – because of adhesion and cohesion

7. Surface Tension Cohesive forces Causes liquids to minimize surface area That’s why water drops are round

8. Gas Particles are independent Far apart No fixed volume or shape Gases and liquids are fluids

9. Changing State Freezing – liquid becomes a solid Melting – solid becomes a liquid Evaporation – liquid becomes gas Condensation – gas becomes liquid Sublimation – solid becomes gas Deposition – gas becomes solid

10. Temperature, Energy, and State

11. Evaporation High energy particles change to gas Causes the substance to cool

12. Boiling Point The temperature at which bubbles of vapor rise to the surface Also depends on atmospheric pressure

13. Intermolecular Forces 11-2

14. Attraction between Particles Takes energy to separate particles (change state) The stronger the force, the more energy it takes The boiling and melting point is a good measure of the strength of the force Strong force of attraction = high boiling point

15. Force of attraction in Ions Higher force of attraction then between molecules High melting points Smaller ions  larger force (NaCl > KCl) Larger charge  larger force (CaF2 > NaCl)

16. Intermolecular Forces The Force of Attraction between molecules

17. Types of Intermolecular Forces Dipole-Dipole Forces Hydrogen Bonds London Dispersion Forces All are short range Little effect on gases Many gases have low boiling point (that is why they are gases)

18. Polar Molecul e A molecule that has an unequal distribution of charge One end slightly positive, One end slightly negative Caused by difference in electronegativity of the atoms

19. Dipole-Dipole Forces Interaction between polar molecules Positive end of one molecule attracts the negative end of another

20. Dipole-Dipole Forces and Boiling Point The more polar the molecules, the stronger the force between them, the higher the boiling point

21. Hydrogen Bonds When a hydrogen atom of one molecule is attracted to an atom of a different molecule Water

22. Hydrogen Bonds Can create a larger difference in electronegativity Also hydrogen is small and has only 1 electron Which increases the bond strength Which increases the boiling point

23. Hydrogen Bonds and Water Water has unique properties, because of hydrogen bonds Can form multiple hydrogen bonds  Strong intermolecular forces

24. Solid water is less dense than liquid water Ice Floats Ponds freeze from top down Expanding ice cracks rocks and concrete

25. London Dispersion Forces The force that hold non- polar molecules together The weakest of the intermolecular forces Explains why some non- polar molecules are not gases

26. London Dispersion Forces Nonpolar molecules can become temporary dipoles (electrons move from side to side) Causes molecules to attract each other

27. London Dispersion Forces Nearby molecules always attract The more electrons, the stronger the force

28. Energy of State Changes 11-3

29. Enthalpy The total energy of a system

30. Entropy A measure of system’s disorder

31. Enthalpy of Fusion The energy added during melting or removed during freezing AKA the heat of fusion

32. Entropy of Fusion The increase of entropy when a solid melts

33. Enthalpy of Vaporization The energy added during evaporation

34. Entropy of Vaporization The increase of entropy when a liquid evaporates Much larger than entropy of fusion

35. The molar enthalpy of fusion The heat energy needed to melt 1 mol of a substance For water it is 6.01 kJ/mol

36. The molar enthalpy of vaporization The heat energy needed to evaporate 1 mol of a substance For water it is 40.67 kJ/mol

38. Phase Equilibrium 11-4

39. System A set of components that are being studied

40. Phase A region that has the same composition and properties throughout Lava lamp – Two phases of liquid - Different chemical compositions

41. Phase Water – Two phases, same chemical composition - Different States

42. Dynamic Equilibrium The net amount of substance in a given phase stays the same Eg. The rate of evaporation equals the rate of condensation Which of these?

43. Vapor Pressure The pressure exerted by a gas in equilibrium with a liquid Boiling point – The temp at which vapor pressure equals the external pressure

44. As temperature increases, vapor pressure increases Normal Boiling Point – when vapor pressure equals the atmospheric pressure

45. Phase Diagrams A graph of the relationship between the state of a substance and its temperature and pressure

46. Phase Diagrams 3 lines Vapor pressure for liquid-gas equilibrium A-B Liquid-solid equilibrium A-D Solid gas equilibrium A-C

47. Triple Point The temperature and pressure at which all three states are in equilibrium

48. Critical Point The temperature and pressure at which the gas and liquid states become identical Called a supercritical fluid

49. Supercritical Fluid