VIEWS: 7 PAGES: 69 POSTED ON: 2/13/2012
Unit 14 Gases Book Chapter 13 Steve Fossett flies his balloon Solo Spirit, over the east coast of Australia during his attempt to make the first solo balloon flight around the world. In this episode • We will look at the properties of gases • We will study the Gas Laws and how they are related to each other • We will construct a model to predict why gases act the way they do • We put it all together along with math to make it practical Pressure 1. What is gas pressure again? • The push that molecules exert when they collide against objects 2. How are we able to expand a balloon? 3. What device is used to measure gas pressure? • barometer 4. What causes air pressure? • As molecules are drawn down toward the earth they will push on the surface and anything around it 5. What happens to air pressure as you go up in elevation? Why? • Decreases. Less molecules less pressure 6. What happens to air pressure as you go down into Death Valley? • It increases. Closer to center of earth means greater amount air between you and space, greater pressure 7. What is air pressure at sea level? • 760 mm of Hg, 1 atmosphere, 101.3 kpa, 14.7 psi(pounds per square inch) 8. Why so many units? • Mainly to confuse you 9. Will you have to use all these units? • Unfortunately, yes Ex. 1 The pressure of air in a tire is measured to be 28 psi. Represent this pressure in atmospheres, torr, and pascals Ex. 2 On a summer day in Breckinridge, Co, the air pressure is 525 mm Hg. What is the air pressure in atm? Boyles Law 10. Who conducted the first experiments on gases? • Robert Boyle 11. What did he work with? • J tube, mercury, stair case 12. What can we deduce from his data? • As pressure went up, Volume went down? 13. What happens when you double the pressure? • You halve the volume 14. What do we call this relationship? • Inverse proportional 15. What mathematical expression can we come up with for this relationship? 16. So if we know a gas’ pressure and volume and we then change one of the variables can we determine how the other variable was changed? • Yes 17. Deduce Boyle’s Law. Ex.3 Freon 12(the common name for the compound CCl2F2) was widely used in refrigeration systems, but has now been replaced by other compounds that do not lead to the breakdown of the protective ozone in the upper atmosphere. Consider a 1.5 L sample of gaseous CCl2F2) at a pressure of 56 torrs. If the pressure is changed to 150 torr at a constant temperature, (a) will the volume of the gas increase or decrease and (b) what will be the new volume of the gas? Ex. 4. In an auto engine the gaseous fuel air mixture enters the cylinder and is compressed by a moving piston before it is ignited. In a certain engine the initial cylinder volume is 0.725 L. After the piston moves up the volume is 0.075 L. The fuel air mix initially has a pressure of 1.00 atm. Calculate the pressure fo the compressed fuel air mix, assuming both T and amount of gas remain constant. Charles’ Law 18. After Robert Boyle, who picked up the gassy torch? • Jacque Charles 19. What is he famous for? • 1st solo balloon flight, first Hydrogen filled balloon 20. What is Charles’ Law? • If you increase the temperature of a gas, its volume will increase and vice versa 21. What kind of relationship is this? • Direct relationship • If one goes up by x amount the other does too • If one goes down by x amnt the other does too Researchers take samples from a steaming volcanic vent at Mount Baker in Washington. 22. What kind of graph do we get if we plot V vs T? • Straight diagonal line 23. No matter the gas, what temperature do we get when we extrapolate back to where all the line meet? • -273 C 24. When using the Gas Laws, what unit do we use? Why? • Kelvin, no negative numbers 25. Why would it be important to use a scale without negative numbers? • Insure that you do not end up with negative volumes or pressures 26. What is the mathematical definition of Charles’ Law? 27. Deduce the most useful equation for Charles’ Law Ex 5 A 2.0 L sample of air is collected at 298 K and then cooled to 278 K. The pressure is held constant at 1.0 atm. Does the volume increase or decrease? Calculate the volume of the air at 278 K. Ex. 6 A sample of gas at 15 ºC(at 1 atm) has a volume of 2.58 L. The temperature is then raised to 28º C(at 1 atm). Does the volume of the gas increase or decrease? Calculate the new volume. Ex. 7 In former times, gas volume was used as a way to measure temp using devices called gas thermometers. Consider a gas that has a volume of 0.675 L at 35 ºC and 1 atm pressure. What is the temp(in units of ºC) of a room where this gas has a volume of 0.535 L at 1 atm. Avogadro’s Law 28. What is the relationship between the number of molecules and volume? • The greater the number of molecules, the greater the volume and vice versa 29. What would be a good analogy for this law? • Inflating a balloon 30. What equation can we use to describe this? 31. What is another way of describing Avogadro’s Law? 32. Deduce the most useful equation for Avogadro’s Law? Ex. 8 Suppose we have a 12.2 L sample containing 0.50 mol of oxygen gas at a pressure of 1 atm and a temp of 25 ºC. If all of this O2 is converted to ozone, O3, at the same temp and pressure, what will be the volume of the ozone formed? Ideal Gas Law 33. What is so difficult about looking at real gases with the Gas Laws? • It is difficult to keep all the variables constant 34. What would be a good solution for this problem? • One equation that uses all the variables 35. Deduce the Ideal Gas Law • Read Chem in Focus for quiz on p 406 Ex. 9 A sample of Hydrogen gas, H2, has a volume of 8.56 L at a temperature of 0C and a pressure of 1.5 atm. Calculate the number of moles of H2 present in this gas sample. (assume the gas behaves ideally) Ex. 10 What volume is occupied by 0.250 mol of CO2 gas at 25 C and 371 torr. Ex. 11 Suppose we have a 0.240 mol sample of ammonia gas at 25 ºC with a volume of 3.5 L at a pressure of 1.68 atm. The gas is compressed to a volume of 1.35 L at 25 ºC. Use the Ideal Gas Law to calculate the final pressure. Ex. 12 A sample of diborane gas, B2H6, a substance that bursts into flames when exposed to air, has a pressure of 0.454 atm at a temperature fo -15 ºC and a volume of 3.48 L. If conditions are changed so that the temp is 36ºC and the pressure is 0.616 atm, what will be the new volume of the sample? Dalton’s Law of Partial Pressures 36. What kind of gas mixture do deep sea scuba divers use in their tanks? • Helium and O2 37. Why do they use this instead of a nitrogen mixture like real air? • Nitrogen dissolves in the blood under these pressure and then forms bubbles on the way up 38. How do gases act when they are in mixture? • The same as if they were alone, they ignore each other 39. Who first studied this concept? • Our old friend, Juan Dalton 40. What is Dalton’s Law? • For a mixture of gases, the total pressure exerted is the sum of the partial pressures of the gases present 41. What is partial pressure? • The pressure exerted by a gas as if it were by itself 42. What equation can we use to more succinctly explain Dalton’s Law? 43. Why does this equation work? Derive another useful incarnation of Dalton’s Law. Ex. 13. Mixtures of helium and oxygen are used in the “air” tanks of underwater divers for deep dives. For a particular dive, 12 L of O2 at 25ºC and 1.0 atm and 46 L of He at 25 ºC and 1.0 atm were both pumped into a 5.0 L tank. Calculate the partial pressures of each gas and the total pressure in the tank at 25ºC. 44. In real life, the most practical way of collecting gas is through the water displacement method. Describe. 45. What is the important problem with the water displacement method? • Its impossible to stop water from evaporating and joining the gas we are collecting 46. How can we factor out this problem? • Subtract out the partial pressure of water for that particular temperature. You can find the PP of water in your ChemOut Ex. 14. A sample of solid potassium chlorate, KClO3, was heated in a test tube and decomposed according to the equation: 2KClO3 --> 2KCl + 3O2(g). The O2 produced was collected over water at 22 ºC. The resulting mixture of O2 and H2O vapor had a total pressure of 754 torr and a volume of 0.650 L. Calculate the partial pressure of O2 in the gas collected and the number of moles of O2 present. The vapor pressure of water at 22 ºC is 21 torr. Laws and Models: A Review 47. What is the most practical gas law? • Ideal gas law 48. What is its basic assumption? • That gases cannot liquefy or interact in any way 49. What is the problem with this? • They do liquefy and interact 50. Under what conditions do Ideal Gases do not act like Real Gases? • High pressure and or low temperatures 51. What can we use to understand the relationships between the phases at different temp and pressures? • Phase diagrams 52. Draw and explain the phase diagram for water. 53. Draw and explain the phase diagram for CO2 54. What is triple point? • The temp and pressure where all phases can exist 55. What is critical temperature? • The temperature above which it does not matter how much pressure you add, it ain’t gonna liquefy 56. What is critical pressure? • The pressure below which it don’t matter how low the temp is, it ain’t gonna liquefy Graham’s Law 57. So basically, assuming a gas is ideal, how fast a molecule moves is dependent on what? • Its size 58. Why is it that the smaller the gas, the faster it will move? • It will collide less frequently Ex. 15 Which of the following molecules will diffuse faster? CH4, CO2, SO2, O2 Ex. 16 Which of the following molecules will diffuse faster? O2, N2, F2, NH3 Gas Stoichiometry • Thanks to the ideal gas law we can expand our love for stoichiometry from Mass to Mass Problems to problems involving gases Ex. 16 Calculate the volume of oxygen gas produced at 1.00 atm and 25 C by the complete decomposition of 10.5 g of potassium chlorate. The balanced equation for the reaction is; 2KClO3(s) --> 2KCl(s) + 3O2(g) Ex. 17 A sample of nitrogen gas has a volume of 1.75 L at STP. How many molecules of N2 are present? Ex. 18 Quicklime, CaO, is produced by heating calcium carbonate, CaCO3. Calculate the volume of CO2 produced at STP from the decomposition of 152 g of CaCO3, according to the reaction: CaCO3(s) --> CaO(s) + CO2(g) Ex. 19. Calculate the volume of hydrogen produced at 1.50 atm and 19 C by the reaction of 26.5 of Zinc with excess hydrochloric acid.