1. GAS PROPERTIES Video 5.1

2. Kinetic Molecular Theory Review  Ideal Gases are perfect gases. They have:  No mass  No volume  No attractive forces  When will real gases behave this way?  When they are spread out  Temperature is high  Pressure is low

3. Boyle’s Law  Pressure and volume have an indirect relationship.  P 1 V 1 =P 2 V 2  What happens to a balloon if you put pressure on it? If you release it into the atmosphere?

6. Charles’ Law  Volume and temperature have a direct relationship.  V 1 /T 1 = V 2 /T 2  What happens to a car tire in the winter?

8. Gay-Lussac’s Law  Pressure and temperature have a direct relationship.  P 1 /T 1 = P 2 /T 2

9. Combined Gas Law P 1 V 1 = P 2 V 2 T 1 T 2 Only used when discussing initial and final conditions of a gas. If something is held constant, cross it out of the equation. Temperature must be in Kelvin!

10. Question 1 P 1 V 1 = P 2 V 2 T 1 T 2 P 1 V 1 = P 2 V 2 (5.6x10 3 )(1.53) = (1.5x10 4 )(x) X =. 571 L

11. Question 2  A 32.9L sample of a gas at constant pressure increases in temperature from 25 to 45C. Should the volume increase or decrease? Calculate the new volume. P 1 V 1 = P 2 V 2 T 1 T 2 V 1 = V 2 T 1 T 2 32.9 = x. 298 318 X= 35.1 L

12. Question 3  A gas in a rigid container has a pressure of 3.5 atmospheres at 200K. Calculate the pressure at 273K. P 1 V 1 = P 2 V 2 T 1 T 2 P 1 = P 2 T 1 T 2 3.5 = x. 200 273 X = 4.78 atm

13. GAS LAWS Video 5.2

14. Avogadro’s Law  Two different gases as the same temperature, volume, and pressure have the same number of molecules.

15. Graham’s Law of diffusion  Gases move from high to low concentrations. Lighter gases diffuse faster.  Which diffuses the fastest: He or N 2 ?  Why will F 2 and Ar diffuse at almost the same rate?

16. Graham’s Law of Diffusion

17. Dalton’s Law of Partial Pressures  The total pressure in a system of gases equals the pressure of each individual gas combined.

18. Example If the total pressure of a container of gases is 3 atmospheres and it contains 1 atmosphere of oxygen, 0.5 atm of nitrogen, 0.75 atm of methane, what is the partial pressure of the remaining gas? 3 – 1 -.5 -.75 =.75 atm

19. How do we breathe?

20. Why is Boyle’s law important to scuba divers?

21. Why is Charles’ Law important in hot air ballooning?

22. SOLUTIONS Video 5.3

23. Vapor Pressure  When a partially filled container of a liquid is sealed, some particles vaporize while some condense. The gas particles exert a pressure on it’s own liquid, called the vapor pressure.  If temperature increases, more gas evolves and Vapor Pressure increases.

24. Vapor Pressure  The boiling point is the temperature at which the vapor pressure equals the external pressure and the liquid becomes a gas.  Normal Boiling Point is the temperature at which a substance boils at 1 atm pressure.

25. Vapor Pressure  At higher altitudes, such as Denver Colorado, air pressure is much lower due to decreases in amount of air molecules. Therefore, water boils at a lower temperature and food takes longer to cook.  At lower altitudes the opposite is true.

26. Vapor Pressure Using Table H answer the following questions: 1. At which temperature will water have the highest vapor pressure: 25, 75, 90 or 120? 2. If the VP of propanone is 90kPa, what is the boiling point? 3. When the temperature of ethanoic acid is changed from 55C to 110C the change in the VP is _______? 4. What are the normal boiling points of all gases listed in table H? 120 52C 80-10 = 70 kPa 55, 79, 100, 118 respectively

27. Properties of Liquids Surface Tension: elastic force on the liquid surface. The stronger the intermolecular force, the higher the surface tension. surface tension

28. Properties of Liquids  Viscocity: measure of a liquid’s resistance to flow.  The greater its intermolecular forces of attraction, the greater the viscosity, the slower it flows.  If you increase the temperature, viscosity decreases.

29. Solutions  Solutions are homogeneous physical mixtures of two or more pure substances. They should be transparent and pass through a filter.  In a solution, the solute is dispersed uniformly throughout the solvent.  Solute: The minor part, get dissolved.  Solvent: The major part, does the dissolving. Water is a universal solvent.

30. Solutions  Aqueous solutions have a water solvent.  The solute can change phase in order to dissolve.  Solids and gases can be solutions too like air and alloys (steel, brass, jewelry).

31. Solubility  If a solute dissolves in the solvent, it is soluble.  If the solute does not dissolve in the solvent it is not soluble and forms a solid (precipitate).  Solubility refers to the amount of solute that can dissolve in 100 grams of water.

32. Factors that affect solubility:  Chemical Nature: Remember the rule: like dissolves like. Polar solvents dissolve polar solutes only.  Temperature: Increasing the temperature will increase the rate of solubility. (Think of iced tea dissolving sugar versus hot tea; hot tea dissolves faster). However, for gases, the relationship is reversed. If you increase the temperature of a gas, the gases will escape (their solubility will decrease).

33. Factors that affect solubility:  Pressure: For gases only: if the pressure is increased, the solubility increases. This is why soda companies package soda under high pressures.

34. Soluble or insoluble? Use Table F  Potassium fluoride  Ammonium hydroxide  Calcium carbonate  Lithium phosphate  Lead acetate  Strontium sulfate  Magnesium chromate  Barium hydroxide  Silver chloride S S I S S I S S I

35. Types of Solutions Saturated: solvent holds as much solute as is possible at that temperature.  In Table G, saturated solutions are on the line.  Dissolved solute is in equilibrium with solid solute particles (some may dissolve and some may precipitate out at the same time).

36. Types of Solutions  Unsaturated: solvent holds less than the maximum amount of solute for that temperature is dissolved in the solvent.  In table G, unsaturated solutions fall under the line.

37. Types of Solutions  Supersaturated: solvent holds more solute than is normally possible at that temperature.  In table G, supersaturated solutions are found above the line.  These solutions are unstable; crystallization can usually be stimulated by adding a “seed crystal” or scratching the side of the flask.

38. Temperature Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.

39. Temperature  The opposite is true of gases:  Carbonated soft drinks are more “bubbly” if stored in the refrigerator.  Warm lakes have less O 2 dissolved in them than cool lakes.

40. Gases in Solution  The solubility of liquids and solids does not change a lot with pressure.  The solubility of a gas in a liquid is directly proportional to its pressure (increase the pressure on a gas, increase the solubility.)

41. CONCENTRATION Video 5.4

42. Concentration  Concentrated solutions contain large amounts of solutes dissolved in the solvent.  Dilute solutions contain small amounts of solutes dissolved in solvent. Which is concentrated?

43. Concentration  On Table G, concentrated solutions are near the top and dilute solutions are near the bottom.  You can add more solvent to a solution in order to dilute the solution.  There are many ways to calculate the concentration of a solution.

44. mol of solute L of solution M = Molarity (M)  Because volume is temperature dependent, molarity can change with temperature.  The equation can be found on table T.

45. Examples 1. If 3 moles of NaCl are dissolved in 6L of water, what is the molarity? 2. If 29 grams of NaCl are dissolved in one liter of water, what is the molarity? 3moles/6L = 29g/58g = 0.5moles 0.5moles/1L = 0.5M

46. Examples 3. If 100 grams of KF are dissolved in 300ml of water, what is the molarity? 4. Calculate the volume needed to create a 2M solution with 3.5 moles of Li2O. 100g/58.1g =1.72mol X = 1.75L 1.72mol/.3L = 2M=3.5mol/x 5.7M

47. Mass Percentage Mass % of A = mass of solute total mass of solution  100

48. Calculate the % by mass: 1. 25 grams of solute in 250 grams of solution. 2. 75.5 grams of NaCl in 255 grams of water. 3. 2 moles of NaOH in 250 grams of solution. 25/250 x 100 = 75.5/(255+75.5) x 100 = 2x40g = 80g/250 x 100 = 10% 22% 32%

49. Parts per Million CO concentration over 70 ppm can cause illness. Over 150ppm can cause death. Ocean water contains over 35000 ppm of salt. 350ppm of CO2 is considered safe. We are at 380ppm in our atmosphere. Global warming?

50. ppm = mass of A in solute total mass of solution  10 6 Parts per Million

51. Find the concentration in ppm:  125 grams of solution with 12 grams of solute.  32.5 grams of NH3 in 200 grams of water.  A solution with 25% solute.  A solution with 60% solvent. 12/125 x 10 6 = 40/100 x 10 6 = 32.5/232.5 x 10 6 = 25/100 x 10 6 = 1.0 x 10 5 ppm 1.4 x 10 5 ppm 2.5 x 10 5 ppm 4.0 x 10 5 ppm

52. COLLIGATIVE PROPERTIES Video 5.5

53. Colligative Properties  Changes in colligative properties depend only on the number of solute particles present, not on the identity of the solute particles.

54. Colligative Properties of Electrolytes Since these properties depend on the number of particles dissolved, solutions of electrolytes (which break down into ions in solution) should show greater changes than those of nonelectrolytes. But it also depends on the Molarity.

55. Boiling Point Elevation and Freezing Point Depression Nonvolatile solute-solvent interactions also cause solutions to have higher boiling points and lower freezing points than the pure solvent.

56. Examples 1. Which compound when dissolved in water, will have the highest boiling point? a. CaCl 2 b. NaI c. C 6 H 12 2. Which compound when dissolved in water, will have the lowest freezing point? a. 1M NaOH b. 2M NH 3 c. 0.5M Mg 3 N 2