1. Chapter 10 and 11 Intermolecular forces and phases of matter Why does matter exist in different phases? What if there were no intermolecular forces? The ideal gas

2. Physical phases of matter Gas Liquid Solid Plasma

3. Physical properties of the states of matter Gases: 1.Highly compressible 2.Low density 3.Fill container completely 4.Assume shape of container 5.Rapid diffusion 6.High expansion on heating

4. Liquid (condensed phase) 1.Slightly compressible 2.High density 3.Definite volume, does not expand to fill container 4.Assumes shape of container 5.Slow diffusion 6.Low expansion on heating

5. Solid (condensed phase) 1.Slightly compressible 2.High density 3.Rigidly retains its volume 4.Retains its own shape 5.Extremely slow diffusion; occurs only at surfaces 6.Low expansion on heating

6. Why water exists in three phases? Kinetic energy(the state of substance at room temperature depends on the strength of attraction between its particles) Intermolecular forces stick molecules together (heating and cooling)

7. Intermolecular forces London Force or dispersion forces Dipole-dipole Hydrogen bond

8. London Force Weak intermolecular force exerted by molecules on each other, caused by constantly shifting electron imbalances. This forces exist between all molecules. Polar molecules experience both dipolar and London forces. Nonpolar molecules experience only London intermolecular forces

9. Dipole-dipole Intermolecular force exerted by polar molecules on each other. The name comes from the fact that a polar molecule is like an electrical dipole, with a + charge at one end and a - charge at the other end. The attraction between two polar molecules is thus a "dipole-dipole" attraction.

10. Hydrogen bond Intermolecular dipole-dipole attraction between partially positive H atom covalently bonded to either an O, N, or F atom in one molecule and an O, N, or F atom in another molecule.

11. To form hydrogen bonds, molecules must have at least one of these covalent bonds: H-N or H-N= H-O- H-F

14. Nonmolecular substances Solids that don’t consist of individual molecules. Ionic compounds(lattices of ions) They are held together by strong ionic bonds Melting points are high

15. Other compounds Silicon dioxide(quartz sand) and diamond (allotrope of carbon) These are not ionic and do not contain molecules They are network solids or network covalent substances

16. Real Gas Molecules travel fast Molecules are far apart Overcome weak attractive forces

17. Ideal Gas Gas that consists of particles that do not attract or repel each other. In ideal gases the molecules experience no intermolecular forces. Particles move in straight paths. Does not condense to a liquid or solid.

18. Ideal Gases Ideal gases are imaginary gases that perfectly fit all of the assumptions of the kinetic molecular theory. Gases consist of tiny particles that are far apart relative to their size. Collisions between gas particles and between particles and the walls of the container are elastic collisions No kinetic energy is lost in elastic collisions

19. Ideal Gases Ideal Gases (continued) Gas particles are in constant, rapid motion. They therefore possess kinetic energy, the energy of motion There are no forces of attraction between gas particles The average kinetic energy of gas particles depends on temperature, not on the identity of the particle.

20. Measurable properties used to describe a gas: Pressure (P) P=F/A Volume (V) Temperature (T) in Kelvins Amount (n) specified in moles

21. Pressure Is caused by the collisions of molecules with the walls of a container is equal to force/unit area SI units = Newton/meter 2 = 1 Pascal (Pa) 1 standard atmosphere = 101.3 kPa 1 standard atmosphere = 1 atm = 760 mm Hg = 760 torr

22. Measuring Pressure The first device for measuring atmospheric pressure was developed by Evangelista Torricelli during the 17 th century. The device was called a “barometer” Baro = weight Meter = measure

23. An Early Barometer The normal pressure due to the atmosphere at sea level can support a column of mercury that is 760 mm high.

24. Units of Pressure UnitSymbolDefinition/Relationship PascalPaSI pressure unit 1 Pa = 1 newton/meter 2 Millimeter of mercury mm HgPressure that supports a 1 mm column of mercury in a barometer AtmosphereatmAverage atmospheric pressure at sea level and 0  C Torrtorr1 torr = 1 mm Hg

25. Standard Temperature and Pressure “STP” P = 1 atmosphere, 760 torr, 101.3 kPa T =  C, 273 Kelvins The molar volume of an ideal gas is 22.4 liters at STP

26. Behavior of gases Rule 1: P is proportional to 1/V Rule 2: P is proportional to T Rule 3: P is proportional to n Combining all three: P is proportional to nT/V P=constant x nT/v R=constant= 0.0821 L atm/K mole

27. Boyle’s Law P inversely proportional to V PV= k Temperature and number of moles constant

28. Boyle’s Law Pressure is inversely proportional to volume when temperature is held constant.

29. A Graph of Boyle’s Law

30. Charles’s Law V directly proportional to T T= absolute temperature in kelvins V/T =k 2 Pressure and number of moles constant

31. Charles’s Law The volume of a gas is directly proportional to temperature, and extrapolates to zero at zero Kelvin. (P = constant) Temperature MUST be in KELVINS!

32. A Graph of Charles’ Law

33. Gay Lussac’s Law The pressure and temperature of a gas are directly related, provided that the volume remains constant. Temperature MUST be in KELVINS!

34. A Graph of Gay-Lussac’s Law

35. The Combined Gas Law The combined gas law expresses the relationship between pressure, volume and temperature of a fixed amount of gas. Boyle’s law, Gay-Lussac’s law, and Charles’ law are all derived from this by holding a variable constant.

36. Standard Molar Volume Equal volumes of all gases at the same temperature and pressure contain the same number of molecules. - Amedeo Avogadro

37. Avogadro’s Law V directly proportional to n V/n = k 3 Pressure and temperature are constant

38. Dalton’s Law of Partial Pressures For a mixture of gases in a container, P Total = P 1 + P 2 + P 3 +... This is particularly useful in calculating the pressure of gases collected over water.

39. Ideal Gas Law PV = nRT P = pressure in atm V = volume in liters n = moles R = proportionality constant = 0.0821 L atm/ mol·  T = temperature in Kelvins Holds closely at P < 1 atm

40. Gas Density … so at STP…

41. Density and the Ideal Gas Law Combining the formula for density with the Ideal Gas law, substituting and rearranging algebraically: M = Molar Mass P = Pressure R = Gas Constant T = Temperature in Kelvins

42. Diffusion: describes the mixing of gases. The rate of diffusion is the rate of gas mixing. Diffusion

43. Effusion Effusion: describes the passage of gas into an evacuated chamber.

44. Effusion: Diffusion: Graham’s Law Rates of Effusion and Diffusion