Understanding States of Matter: Gases and Kinetic-Molecular Theory

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: • 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 • Assumptions of Kinetic-molecular theory: • (3) Particle Energy: • Kinetic Energy (KE) of particle represented by : • KE = mν2 • m = mass • ν = velocity (speed & direction) • All particles have same mass but different velocities; do not have same kinetic energy • Temperature: measure of average kinetic energy of particles

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 Effusion: similar to diffusion; gas escapes through a tiny hole. Graham’s Law of Effusion: relationship between effusion rates and molar mass of gas Rate of Effusion =

8. Graham’s Law of Effusion Use it to relate rates of diffusion between different gases since heavier particles diffuse more slowly. =

9. Practice Problem Find the ratio of diffusion rates between ammonia (NH3) and hydrogen chloride (HCl). molar mass NH3: 17.0 g/mol molar mass HCl: 36.5 g/mol

10. Practice Problem Find the ratio of diffusion rates between ammonia (NH3) and hydrogen chloride (HCl). (molar mass NH3: 17.0 g/mol molar mass HCl: 36.5 g/mol) =

11. Practice Problem Find the ratio of diffusion rates between ammonia (NH3) and hydrogen chloride (HCl). (molar mass NH3: 17.0 g/mol molar mass HCl: 36.5 g/mol) = * Plug & Chug

12. Practice Problem Find the ratio of diffusion rates between ammonia (NH3) and hydrogen chloride (HCl). = = = 1.47 * NH3 diffuses ~1.5 times as fast as HCl

13. Practice Problem Calculate the ratio of effusion rates for nitrogen (N2) and neon (Ne).

14. Practice Problem Calculate the ratio of effusion rates for nitrogen (N2) and neon (Ne). =

15. Practice Problem Calculate the ratio of effusion rates for nitrogen (N2) and neon (Ne). =

16. Practice Problem Calculate the ratio of effusion rates for nitrogen (N2) and neon (Ne). = = 0.849

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, 1022 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 • 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 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. • Ptotal = P1 + P2 + P3 + ….Pn

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 O2, if the partial pressure of CO2 is 0.70 atm and the partial pressure of N2 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 O2, if the partial pressure of CO2 is 0.70 atm and the partial pressure of N2 is 0.12 atm? Ptotal = P1 + P2 + P3 + ….Pn

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 O2, if the partial pressure of CO2 is 0.70 atm and the partial pressure of N2 is 0.12 atm? Ptotal = PO2 + PN2 + PCO2

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 O2, if the partial pressure of CO2 is 0.70 atm and the partial pressure of N2 is 0.12 atm? Ptotal = PO2 + PN2 + PCO2 0.97 atm = PO2 + 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 O2, if the partial pressure of CO2 is 0.70 atm and the partial pressure of N2 is 0.12 atm? Ptotal = PO2 + PN2 + PCO2 0.97 atm = PO2 + 0.12 atm + 0.70 atm PO2 = 0.97 atm – 0.12 atm – 0.70 atm PO2 = 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: • Intramolecular forces: • Attractive forces that hold particles together in ionic, covalent, and metallic bonds • - “intra” means “within” • 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 + - + - - - + + + + + + + + + +

31. Intermolecular Forces Three types of intermolecular forces: Dispersion Forces Dipole-dipole Forces 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 (O2) 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+ temporary dipole 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 -

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 (H2O), or nitrogen (NH3). • Hydrogen bonds explain why water is liquid at room temperature, while other molecules of comparable masses are gases. (H2O vs. CH4)

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+

40. Intermolecular Forces Comparison of Intermolecular Forces

41. Practice Problem Determine the type of intermolecular forces in each molecule. F2 H2O HI HF

42. 13.3 Liquids: Fluidity • 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

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

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

45. Liquids: Viscosity • 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

46. Liquids: Viscosity • 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

47. Liquids: Surface Tension • 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

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

49. Liquids: Capillary Action • 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”

50. Solids: Density • Density of solids