1. Chapter 12 States of Matter

2. Gases Kinetic-molecular theory: Describes behavior of gases in terms of particles in motion Kinetic means “to move” Objects in motion have kinetic energy Proposed by Boltzmann & Maxwell Makes many assumptions

3. Gases Assumptions of Kinetic-molecular theory: (1)Particle Size: Gases are small particles separated from each other by empty space Volume of particle is small compared with volume of empty space Particles are far apart, have no attractive or repulsive forces between gas particles.

4. Gases Assumptions of Kinetic-molecular theory: (2) Particle Motion: Gas particles are in constant random motion Move in straight line until collision w/ walls of container or another gas particle Collisions are elastic (no lost kinetic energy)

5. Gases

6. Diffusion of gases Diffusion: describes movement of one material through another Smell food cooking in kitchen while watching TV in living room Rate of diffusion depends on mass of particles involved lighter particles diffuse more rapidly

7. Graham’s Law of Effusion

9. Practice Problem Find the ratio of diffusion rates between ammonia (NH 3 ) and hydrogen chloride (HCl). molar mass NH 3 : 17.0 g/mol molar mass HCl: 36.5 g/mol

10. Practice Problem

13. Calculate the ratio of effusion rates for nitrogen (N 2 ) and neon (Ne).

14. Practice Problem

17. Gas Pressure Pressure: force per unit area Gas particles exert force when they collide with walls of containers. Individual gas particles has little mass; exerts little pressure But, 10 22 gas particles in a liter container; pressure can be substantial How do we measure air pressure?

18. Gas Pressure How do we measure air pressure? A Barometer! Increase in air pressure: Hg rises Decrease in air pressure: Hg falls

19. Units of Pressure (1)Pascal (Pa): SI unit of pressure Derived unit from kilogram, meter, second One Pascal is equal to force of one newton per square meter: 1 Pa = 1 N/m2 (2) Pounds per square inch (psi) (3) Millimeter of mercury (mm Hg) (4) Torr (1 torr = 1 mm Hg) (4) Atmosphere (atm): represents air pressure 1 atm = 760 mm Hg = 760 torr

20. Units of Pressure UnitCompared w/ 1 atmCompared w/ 1 kPa kilopascal (kPa)1 atm = 101.3 kPa Millimeters of Hg (mm Hg) 1 atm = 760 mm Hg1 kPa = 7.501 mm Hg torr1 atm = 760 torr1 kPa = 7.501 psi Pounds per square inch (psi) 1 atm = 14.7 psi1 kPa = 0.145 psi Atmosphere (atm)1 kPa = 0.009 atm Comparison of Pressure Units

21. Dalton’s Law of Partial Pressure Dalton studied properties of gases & found that each gas in a mixture exerts pressure independently of other gases present. Dalton’s Law of Partial Pressure: total pressure of mixture of gases is equal to sum of the pressures of all the gases in mixture. P total = P 1 + P 2 + P 3 + ….P n

22. Dalton’s Law of Partial Pressure Example: A mixture of oxygen, carbon dioxide, and nitrogen has a total pressure of 0.97 atm. What is the partial pressure of O 2, if the partial pressure of CO 2 is 0.70 atm and the partial pressure of N 2 is 0.12 atm?

23. Dalton’s Law of Partial Pressure Example: A mixture of oxygen, carbon dioxide, and nitrogen has a total pressure of 0.97 atm. What is the partial pressure of O 2, if the partial pressure of CO 2 is 0.70 atm and the partial pressure of N 2 is 0.12 atm? P total = P 1 + P 2 + P 3 + ….P n

24. Dalton’s Law of Partial Pressure Example: A mixture of oxygen, carbon dioxide, and nitrogen has a total pressure of 0.97 atm. What is the partial pressure of O 2, if the partial pressure of CO 2 is 0.70 atm and the partial pressure of N 2 is 0.12 atm? P total = P O2 + P N2 + P CO2

25. Dalton’s Law of Partial Pressure Example: A mixture of oxygen, carbon dioxide, and nitrogen has a total pressure of 0.97 atm. What is the partial pressure of O 2, if the partial pressure of CO 2 is 0.70 atm and the partial pressure of N 2 is 0.12 atm? P total = P O2 + P N2 + P CO2 0.97 atm = P O2 + 0.12 atm + 0.70 atm

26. Dalton’s Law of Partial Pressure Example: A mixture of oxygen, carbon dioxide, and nitrogen has a total pressure of 0.97 atm. What is the partial pressure of O 2, if the partial pressure of CO 2 is 0.70 atm and the partial pressure of N 2 is 0.12 atm? P total = P O2 + P N2 + P CO2 0.97 atm = P O2 + 0.12 atm + 0.70 atm P O2 = 0.97 atm – 0.12 atm – 0.70 atm P O2 = 0.15 atm

27. Practice Problem What is the partial pressure of hydrogen gas in a mixture of hydrogen and helium if the total pressure is 600 mm Hg and the partial pressure of helium is 439 mm Hg?

28. 13.2 Forces of Attraction If all particles of matter have the same average kinetic energy (at room temp), why are some materials gases while others are liquids or solids?

29. Forces of Attraction The answer is the attractive forces between particles. Two types of attractive forces: 1)Intramolecular forces: -Attractive forces that hold particles together in ionic, covalent, and metallic bonds - “intra” means “within” 2)Intermolecular forces: -Attractive forces that hold together water molecules in drop of water, two different particles -“inter” means “between”

30. Intramolecular Forces Comparison of Intramolecular Forces ForceModelBasis of Attraction Example Ionic cations & anions NaCl Covalent positive nuclei & shared electrons H2H2 Metallic metal cations & mobile electrons Fe + - + - - + - + ++ +++ +++

31. Intermolecular Forces Three types of intermolecular forces: 1)Dispersion Forces 2)Dipole-dipole Forces 3)Hydrogen Bonds

32. Dispersion Forces Dispersion Forces: weak forces that result from temporary shifts in the density of electrons in electron cloud. Sometimes called London forces after German- American physicist who first described them, Fritz London.

33. Dispersion Forces Oxygen molecules (O 2 ) are nonpolar because electrons are equally distributed between equally electronegative oxygen atoms. But, electrons in the electron cloud are in constant motion. When two nonpolar molecules collide the electron cloud of one molecule repels the electron cloud of the other molecule. The electron density around one nucleus is greater in one region of each cloud, and forms temporary dipole.

34. Dispersion Forces Temporary dipole forms in each molecule. Temporary dipoles close together cause weak dispersion force between oppositely charged regions of dipole. δ+δ+δ+δ+δ - temporary dipole

35. Dispersion Forces Dispersion forces exist between all particles, but only play a significant role when no stronger forces of attraction acting on particles. Dominant force of attraction between identical nonpolar molecules!! Weakest forces of attraction. Forces have noticeable effect with increasing number of electrons involved (bigger dipole and hence larger forces)…explains why fluorine and chlorine are gases, bromine is a liquid, and iodine is solid at room temp.

36. Dipole-dipole Forces Dipole-dipole forces: attractions between oppositely charged regions of polar molecules. Polar molecules contain permanent dipoles; some regions of a polar molecule are always partially negative and some regions are always negative. Typically stronger than dispersion forces as long as the molecules being compared have similar masses.

37. Dipole-dipole Forces Neighboring polar molecules orient themselves so that oppositely charged regions line up. In HCl, the partially positive hydrogen atom is attracted to the partially negative chlorine atom. δ - δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ+δ+

38. Hydrogen Bonds Hydrogen bond: is a special type of dipole-dipole attraction that occurs between molecules containing a hydrogen atom bonded to a small, highly electronegative atom with at least one lone pair. Examples: hydrogen is bonded to fluorine (HF), oxygen (H 2 O), or nitrogen (NH 3 ). Hydrogen bonds explain why water is liquid at room temperature, while other molecules of comparable masses are gases. (H 2 O vs. CH 4 )

39. Hydrogen bonds Water molecules δ - δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ+δ+

40. Intermolecular Forces Comparison of Intermolecular Forces ForceType of molecule Strength of Force Example Dispersion nonpolar weak H2H2 Dipole- dipole polar, permanent dipole Stronger than dispersion HCl Hydrogen bonds polar, contains H-F, H-O, or H-N typically strongest NH 3

41. Practice Problem Determine the type of intermolecular forces in each molecule. 1)F 2 2)H 2 O 3)HI 4)HF

42. Fluidity: ability to flow Both liquids and gases flow Liquids can diffuse through another liquid (food coloring in water) Liquids diffuse more slowly than gases at the same temperature 13.3 Liquids: Fluidity

43. Which will diffuse more quickly 1)Food coloring in cold water? 2)Food coloring in hot water? Liquids: Fluidity

44. Viscosity: measure of the amount of resistance of a liquid to flow. What is more viscous… maple syrup OR milk? Liquids: Viscosity

45. Particles in liquid have attractive forces that slow movement as they flow past one another. Viscosity of liquid is determined by size and shape of particle Bigger particles have higher viscosity Longer particles have higher viscosity compared to shorter molecules Liquids: Viscosity

46. Viscosity is affected by temperature. What happens when you put oil into a hot pan? Increase in temp = increase in kinetic energy Higher temperature = Decrease in Viscosity Liquids: Viscosity

47. Intermolecular forces: attractions between particles Intermolecular forces not equal for all particles in a liquid. Particles in middle of liquid attracted to particles above and to side of them. Particles on surface have no attractions from above to balance attractions below Liquids: Surface Tension

48. Surface tension: measurement of inward pull by the particles in the interior. Allows spiders to “walk on water” Liquids: Surface Tension

49. Water in graduated cylinder has meniscus due to attractive forces between water molecules and glass molecules (silicon dioxide) Water molecules more attracted to silicon than other water molecules If narrow tubes, water will be drawn upward Called “Capillary Action” Liquids: Capillary Action

50. Density of solids Solids: Density

51. Phase changes of water: solid → liquid → gas These phase changes require energy Energy flows from higher temperature to lower temperature 13.4 Phase Changes

52. condensation vaporization melting freezing sublimation deposition gas liquid solid Phase Changes

53. Phase changes that require energy: Melting Vaporization Evaporation Sublimation Phase Changes

54. Melting: solid particles absorb enough heat to change to the liquid phase Phase Changes: Melting

55. Molecules have more kinetic energy w/ increasing temp Vaporization: process where a liquid changes to a gas or vapor Evaporation: vaporization occurs only at surface of liquid Your body controls its temp by evaporation Phase Changes: Vaporization

56. Evaporation in an open container (a) or evaporation in a closed container (b). Vapor pressure: pressure exerted by vapor over liquid Vapor Pressure

57. Boiling Point: temperature at which vapor pressure of a liquid equals the external or atmospheric pressure At boiling point, molecules have enough energy to vaporize Boiling Point

58. Sublimation: process which solid changes directly to a gas without first becoming a liquid “Dry Ice” is solid carbon dioxide Sublimation

59. Phase changes that release energy: Condensation Deposition Freezing Phase Changes

60. Water vapor molecules lose energy, velocity is reduced, interact with other water molecules Condensation: change from the vapor to the liquid phase Condensation is the reverse of vaporization Phase Changes: Condensation

61. Deposition: process by which substance changes from a gas or vapor to solid without becoming liquid. Reverse of sublimation Formation of snowflakes (energy released as crystals form) Phase Changes: Deposition

62. Freezing Point: temperature at which liquid is converted to crystalline solid Heat is removed from water, lose kinetic energy, molecules become fixed or “frozen” into set positions Phase Changes: Freezing

63. Phase Diagram: graph of pressure versus temperature; shows what phase a substance exists under different temperatures and pressures Triple point: point on graph that represents temp and pressure where all 3 phases exist Phase Diagram