Kinetic Theory and Gases: States of Matter and Particle Motion

1. Chapter 13 Notes States of Matter

2. Kinetic Theory and Gases • Kinetic Energy—Energy that an object has due to motion. • The Kinetic Theory is that tiny particles form all matter, and they are constantly in motion.

3. Kinetic energy vs. Potential Energy Kinetic refers to motion—so kinetic energy is the energy of motion; this is different from potential energy, which is the possible amount of energy stored in something. -the kinetic theory states that tiny particles form all matter, that are in constant motion.

4. Kinetic Theory and Gases • A gas is composed of particles that are small, hard spheres with insignificant volume and no particle interaction.

5. Kinetic Theory and Gases • Particles in a gas are in constant motion—they travel straight paths unless they collide with another particle or their container.

6. Kinetic Theory and Gases • All collisions are considered elastic—no energy is lost to friction.

7. Kinetic Theory and Gases 4.) No kinetic energy is lost when gas particles collide • elastic collisions occur w/ other gas particles or with the wall of the container • energy can be transferred in collision but the total kinetic energy of the 2 particles does not change

8. Kinetic Theory and Gases 5.) All gases have same average kinetic energy at the same temperature • kinetic energy of motion (molecules are always moving) • Temperature is a measure of the average kinetic energy of the particles in a sample of matter (at a given temp., all gases have the same avg. KE)

9. Temperature ↓ in temp = ↓ in K.E. (molecules slow down; theoretically, if you could lower the temp enough motion would cease)

10. Temperature Kelvin scale is a direct measure of average kinetic energy (eg. particles at 200 K have 2x as much nrg as at 100 K) K= oC + 273 (0 oC = 273 K) Which has more kinetic energy and does most damage to a brick wall - a big pickup truck or a Honda Prelude? wt. = 15,000 lbs wt. = 3000 lbs Kinetic Energy = ½ mv2 • big molecules move more slowly, lightweights move faster • gases move from hi concentration → lo concentration • rate they move depends on kinetic energy (in other words, the size and velocity of particles)

11. Kinetic Theory and Gases Effusion = gas escapes thru tiny opening ex: hole in tire, air effuses and tire goes flat ex: helium in balloon overnight vs air in balloon Diffusion = gas A mixing with (moving thru) gas B ex: perfume sprayed in one room, noticed in next rm ex: rotten egg

12. Kinetic Theory and Gases Graham’s law of effusion: Rate of effusion= 1/(sq. root of molar mass) Graham’s law of diffusion: Rate of diffusion= Rate A = (sq. root of molar mass B/ molar mass A) Rate B

13. Behavior of Gases • Kinetic-molecular theory  a great deal of space exists between gas particles • Large amount of empty space between the particles allows compressibility and expansion of gas particles

14. Gas Pressure • Kinetic theory explains the existence of gas pressure. • Gas pressure—the force exerted by a gas per unit surface area.

15. Gas Pressure • The force of one molecule hitting an object is relatively small, but the result of billions of particles of air hitting a surface at once is significant.

16. Gas Pressure pressure = force / unit area To increase pressure (force/area): 1. more particles per unit area a. decrease volume of container (↓ area) b. add more particles 2. increase temp: ↑ speed of particles causing ↑ collisions

17. What happens as you increase altitude (climb a mountain)? Gravity pulls air particles in toward earth. The air at higher altitudes has less air above pushing down and fewer air molecules in a given space. Atmospheric pressure decreases as you gain altitude. Pilots gauge their altitude by measuring pressure.

18. Atmospheric Pressure A barometer measures atmospheric pressure. The SI unit for pressure is the pascal (Pa). Atmospheric pressure at sea level is about 101.3 kilopascals (kPa). Other units of measurement are atmospheres (atm), mm Hg, and pounds per square inch (psi). 1 atm = 101.3 kPa = 760 mm Hg = 14.7 psi

19. Comparison of Pressure Units Units of Pressure (p390) 1 atm = the average atmospheric pressure at sea level kilopascal 1 atm = 101.3 kPa Torricelli 1 atm = 760 torr mm mercury 1 atm = 760 mm Hg inches mercury 1 atm = 29.9 in Hg pounds / in2 1 atm = 14.7 psi

20. Pressure conversion problems • Convert 190 mm Hg to atm 2. The pressure at the top of Mt Everest is 4.89 psi. How many mm of Hg is this? in. of Hg? How many atm?

21. What is an absence of particles called? • A vacuum! • No particles = no pressure • Atmospheric pressure is the amount of pressure from the particles in the atmosphere colliding with objects.

22. STP STP = Standard Temperature and Pressure Since temperature and air pressure may vary form place to place it is necessary to have standard reference conditions for testing purposes STP is commonly used to define standard conditions for temperature and pressure 0oC or 273K and 1 atm or 760 mm

23. Dalton’s Law of Partial Pressures There are mixtures of gases in a container • each type of gas contributes a fraction of the particles which will supply a similar fraction of the pressure

24. Dalton’s Law of Partial Pressures At constant vol. & temp., the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures Ptotal= P1 + P2 + P3 + ….Pn

25. Dalton’s Law of Partial Pressures example: Air contains oxygen, nitrogen, carbon dioxide and trace amounts of argon and other gases. What is the partial pressure of O2 at 1 atm of pressure if PN2 = 593.4 mm? PCO2 = 0.3 mm, and Pother = 7.1 mm ?

26. Dalton’s Law of Partial Pressures Partial Pressure * colliding particles → pressure * more particles → more pressure # of particles often measured in moles Does 1 mol O2 contain the same # of molecules as 1 mol H2? 1 mol = 6.02 x 1023 particles = 22.4 L Does 1 L of O2 contain the same # of molecules as 1 L H2 ?

27. End of Daily Notes

28. Liquids and Kinetic Theory Particles in a liquid still have kinetic energy—the particles vibrate and spin and slide past each other—but not as much as is present in a gas. One of the differences between the two is that particles in a liquid are attracted to one another. The attraction brings the particles closer together, and hold it together with other molecules. This also gives rise to surface tension.

29. Intermolecular forces • Intermolecular forces- hold together identical particles (drop of water), carbon atoms in graphite, and the cellulose particles in paper • All intramolecular, or bonding forces are stronger than intermolecular forces

30. Dispersion Forces • Dispersion forces weak forces that result from temporary shifts in the density of electrons in electron clouds (weakest intermolecular force) • Example: Oxygen molecules are nonpolar (b/c e- are evenly distributed); under the right conditions, oxygen molecules can be compressed into a liquid; the force of attraction between oxygen molecules is dispersion or London forces

31. Dispersion Forces • Dispersion forces cont: • e- in an e- cloud are in constant motion • When 2 nonpolar molecules are in close contact or when they collide, the e- cloud of one molecule repels the e- cloud of the other molecule. • The e- density around each nucleus is, for a moment, greater in one region of each cloud; each molecule forms a temporary dipole • When temporary dipoles are close together, a weak dispersion force exists between oppositely charged regions of the dipoles

32. Dispersion Forces • Recall your Halogen gases (F, Cl, Br, I) all exist as diatomic molecules. • The # of nonvalence e- from fluorine to chlorine to bromine, to iodine. B/c the larger halogens have more e-, there can be a greater difference between positive and negative dipoles and thus stronger dispersion forces

33. Dipole - Dipole Forces • Dipole – Dipole forces  attractions between oppositely charged regions of polar molecules since polar molecules contain permanent dipoles • Neighboring polar molecules orient themselves so that oppositely charged regions line up

34. Dipole - Dipole Forces For Example: When hydrogen chloride gas molecules approach, the partially positive hydrogen atom in one molecule is attracted to the partially negative chlorine atom in another molecule.

35. Hydrogen Bonds • Hydrogen Bonds  type of dipole-dipole attraction that occurs between molecules containing a hydrogen atom bonded to a small, highly electornegative atom with at least one lone e- pair

36. Hydrogen Bonds For example: • for a hydrogen bond to form, hydrogen must be bonded to either a fluorine, oxygen, or nitrogen atom • These atoms are electronegative enough to cause a large partial positive charge on the hydrogen atom, yet small enough that their lone pairs of e- can come close to hydrogen atoms

37. Rank the intermolecular forces in order of increasing strength Dispersion forces  dipole-dipole forces  hydrogen bond

38. Liquids • Kinetic-molecular theory predicts the constant motion of the liquid particles • Individual liquid molecules do not have fixed positions; forces of attraction between liquid molecules limit their range of motion so that the particles remain closely packed in a fixed volume

39. Liquids • Like gases, liquids can be compressed • The change in volume is much less than that of gases b/c liquid particles are already tightly packed together

40. Liquids • Fluidity  ability to flow • Liquids are less fluid than gases because of intermolecular attractions • Viscosity  measure of the resistance of a liquid to flow • As temp. increases, viscosity decreases

41. Solids (least KE) • The particles in the solid move, but don’t move around. They vibrate around a fixed point. • Most solids are crystalline—they have definite repeating structure. • Substances that have more than one crystalline structure are called allotropes.

42. Solids MOLECULAR solids: covalent molecules held together by intermolecular attractions only, weaker than ionic or metallic bonds so these have lower melting and boiling points ex: H2O, CO2, sugar, wax

43. Solids COVALENT NETWORK solid : a crystalline exception to the molecular norm ex: diamond

44. SOLIDS IONIC Solids: held together by strong attraction between + and – ions, hi melting pt., form crystals : ions arranged in orderly repeating pattern of unit cells ex: NaCl, KCl, MgSO4, NaOH

45. Solids METALLIC solids: cations in a sea of valence e-; most have strong bonds & crystalline structure & hi melting point

46. Solids There are some substances that have no crystalline structure at all. These are called amorphous solids. There atoms are randomly arranged with no pattern. Examples are rubber, plastic, glass, asphalt, etc.

47. Changes of State • We have discussed that the state of a substance does not just depend on the temperature of the substance, but also the pressure that it is under. • -A phase diagram shows the conditions at which a substance exists as a solid, liquid and gas.

48. Phase changes that require energy Vaporization = liquid turns to gas (vapor) Evaporation = vaporization occurring only at the surface (cooling process) Melting = solid becomes a liquid Vapor pressure = pressure exerted by a vapor over a liquid Boiling = vapor pressure equals atmospheric pressure (cooling process) Sublimation = solid changes directly into a gas

49. Evaporation evaporation—conversion of liquid to a gas when the surface of a liquid is not boiling evaporation is a cooling process – water (or sweat) absorbs heat kinetic energy rises surface water escapes the chaos and takes some of the kinetic energy (aka temp) with it leaving the cooler (slower moving) molecules behind to absorb more heat. They suck the heat out till they too escape.

50. Melting Point • The temperature at which a solid becomes a liquid is the melting point. • As kinetic energy is added to a solid, eventually the particles have so much energy that they overcome the interaction between particles and vibrate and spin themselves right out of their structure.