Unit 9 - States of Matter and Gas Behavior

1. Unit 9 - States of Matter and Gas Behavior Chemistry Chapters 12-13

2. The Behavior of Gases Gases expand, diffuse, exert pressure, and can be compressed because they are in a low density state consisting of tiny, constantly-moving particles. • Kinetic-molecular theory explains the different properties of solids, liquids, and gases. • Atomic composition affects physical and chemical properties. • The kinetic-molecular theory describes the behavior of matter in terms of particles in motion.

3. The Behavior of Gases • Gases consist of small particles separated by empty space. • Gas particles are too far apart to experience significant attractive or repulsive forces. • Gas particles are in constant random motion. • An elastic collision is one in which no kinetic energy is lost. • Kinetic energy of particles is directly proportional to temperature

4. The Behavior of Gases • Kinetic energy of a particle depends on mass and velocity. • Temperatureis a measure of the average kinetic energy of the particles in a sample of matter.

5. The Behavior of Gases • Great amounts of space exist between gas particles. • Compression reduces the empty spaces between particles.

6. Gas Pressure • Pressureis defined as force per unit area. • Gas particles exert pressure when they collide with the walls of their container. • The particles in the earth’s atmosphere exert pressure in all directions called air pressure. • There is less air pressure at high altitudes because there are fewer particles present, since the force of gravity is less.

7. Gas Pressure • Barometersare instruments used to measure atmospheric air pressure.

8. Gas Pressure • The average height of a mercury column in a barometer at sea level is 760 mm (760 mmHg) • There are other units of pressure as well with the following relationships: • 1.00 atm = 760 mmHg = 760 torr = 101.3 kPa

9. Gas Pressure • Dalton’s law of partial pressures states that the total pressure of a mixture of gases is equal to the sum of the pressures of all the gases of the mixture. • The partial pressure of a gas depends on the number of moles, size of the container, and temperature and is independent of the type of gas.

10. Gas Pressure • Ptotal = P1 + P2 + P3 +…Pn • Partial pressure can be used to calculate the amount of gas produced in a chemical reaction.

11. Gas Pressure • Gases are often collected by water displacement, but water vapor is introduced to the gas sample (mixture of gases) • You can use Dalton’s law to get the partial pressure of the gas without the water vapor

12. Forces of Attraction Intermolecular forces—including dispersion forces, dipole-dipole forces, and hydrogen bonds—determine a substance’s state at a given temperature. • Attractive forces between molecules cause some materials to be solids, some to be liquids, and some to be gases at the same temperature.

13. Forces of Attraction Prefix INTRA means “within”, where INTER means “between”

14. Forces of Attraction • Dispersion forces are weak forces that result from temporary shifts in the density of electrons in electron clouds. • Weakest of all intermolecular forces • Dispersion forces become increasingly stronger as molecule gets more electrons, which explains why iodine is a solid but fluorine is a gas at room temperature

15. Forces of Attraction • Dipole-dipole forces are attractions between oppositely charged regions of polar molecules.

16. Forces of Attraction • Hydrogen bondsare special dipole-dipole attractions that occur between molecules that contain a hydrogen atom bonded to a small, highly electronegative atom with at least one lone pair of electrons, typically fluorine, oxygen, or nitrogen.

17. Forces of Attraction

18. Liquids and Solids The particles in solids and liquids have a limited range of motion and are not easily compressed. • Liquids and solids are called “condensed phases” because their particles are very close together and do not have the energy necessary to “escape” the sample…their motion is limited

19. Liquids and Solids • Solid particles are arranged in a crystal arrangement where they are fixed in position. • This arrangement results in a substance that has a defined shape and volume.

20. Liquids and Solids • Liquid particles have enough energy to “slide” by other particles, but do not have the energy to escape the sample itself. • This results in a substance that has a fixed volume but takes the shape of its container.

21. Phase Changes Matter changes phase when energy is added or removed. • Heat is the transfer of energy from an object at a higher temperature to an object at a lower temperature.

22. Phase Changes • When ice is heated, the ice eventually absorbs enough energy to break the hydrogen bonds that hold the water molecules together. • When the hydrogen bonds break, the particles move apart and ice melts into water. • Themelting point of a crystalline solid is the temperature at which the forces holding the crystal lattice together are broken and it becomes a liquid.

23. Phase Changes • Vaporizationis the process by which a liquid changes to a gas or vapor. • Evaporationis vaporization only at the surface of a liquid.

24. Phase Changes • In a closed container, the pressure exerted by a vapor over a liquid is called vapor pressure.

25. Phase Changes • The boiling point is the temperature at which the vapor pressure of a liquid equals the atmospheric pressure.

26. Phase Changes • Evaporation vs Boiling Summary: • Evaporation is the result of surface level molecules obtaining enough energy to escape the sample as a gas • Boiling occurs throughout the sample when the vapor pressure of the liquid is the same as the atmospheric pressure

27. Phase Changes • Sublimationis the process by which a solid changes into a gas without becoming a liquid. • As heat flows from water to the surroundings, the particles lose energy. • Thefreezing point is the temperature at which a liquid is converted into a crystalline solid. • As energy flows from water vapor, the velocity decreases. • The process by which a gas or vapor becomes a liquid is called condensation. • Depositionis the process by which a gas or vapor changes directly to a solid, and is the reverse of sublimation.

28. Phase Changes • A phase diagram is a graph of pressure versus temperature that shows in which phase a substance will exist under different conditions of temperature and pressure.

29. Phase Changes

30. Phase Changes • The triple point is the point on a phase diagram that represents the temperature and pressure at which all three phases of a substance can coexist. • The critical point is the highest temperature the substance can exist as a liquid, regardless of pressure • Normal melting point and Normal boiling point are the temperatures at which these phase changes occur under standard pressure (760 torr)

31. Gas Laws For a fixed amount of gas, a change in one variable—pressure, temperature, or volume—affects the other two.

32. Gas Laws • Boyle’s law states that the volume of a fixed amount of gas held at a constant temperature varies inversely with the pressure. P1V1 = P2V2 where P = pressure and V = volume

33. Gas Laws • Example Problem. A sample of oxygen gas has a volume of 150. mL when its pressure is 0.947 atm. What will the volume of the gas be at a pressure of 0.987 atm if the temperature is constant?

34. Gas Laws • Givens and Unknown • P1=0.947 atm • V1=150. mL • P2=0.987 atm • V2=?

35. Gas Laws • Equation  P1V1=P2V2 • Solve for unknown, V2

36. Gas Laws • As temperature increases, so does the volume of gas when the amount of gas and pressure do not change. • Kinetic-molecular theory explains this property.

37. Gas Laws

38. Gas Laws • Celsius and Fahrenheit use ARBITRARY zero points, where any temperature lower than these points are considered to be negative • Absolute zerois zero on the Kelvin scale. • Remember - temperature is a measure of kinetic energy of the particles. Particles CAN NOT HAVE less kinetic energy than zero. Therefore, at absolute zero, all motion ceases.

39. Gas Laws • Conversions: • Kelvin T (K) = Celsius T (C) + 273 • Celsius T (C) = Kelvin T (K) - 273

40. Gas Laws • Charles’s law states that the volume of a given amount of gas is directly proportional to its kelvin temperature at constant pressure.

41. Gas Laws • IMPORTANT! Your temperatures MUST BE IN KELVIN when using Charles’ Law (or any gas law). • A rearrangement of the Charles’ Law equation that is often easier to use (no fraction) is V1T2=V2T1

42. Gas Laws • Example - A sample of oxygen gas has a volume of 150. mL when its temperature is 295 K. What will the volume of the gas be at a temperature of 309 K if the pressure is constant?

43. Gas Laws • Givens and Unknown • V1=150. mL • T1=295 K • V2=? • T2=309 K

44. Gas Laws • Equation  V1T2=V2T1 • Solve for V2

45. Gas Laws • Example # 2 - A sample of oxygen gas has a volume of 277 mL when its temperature is 52C. If the gas volume is increased to 344 mL, at what Celsius temperature will the gas be?

46. Gas Laws • Givens and Unknown • V1=277 mL • T1=52C - convert to K - 52 + 273 = 325 K • V2=344 mL • T2=?

47. Gas Laws • Equation - V1T2=V2T1 • Solve for T2 404K - 273 = 131C

48. Gas Laws • Gay-Lussac’s law states that the pressure of a fixed amount of gas varies directly with the kelvin temperature when the volume remains constant.

49. Gas Laws

50. Gas Laws • The combined gas lawstates the relationship among pressure, temperature, and volume of a fixed amount of gas. • We can rearrange the equation to remove the fraction and get P1V1T2=P2V2T1