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13 gas laws


									Unit 14 Gases

 Book Chapter 13
Steve Fossett
flies his
balloon Solo
Spirit, over the
east coast of
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
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
• 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
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
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
13. What happens when you double the
• 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 refrigeration systems, but has now been
 replaced by other compounds that do not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
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
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
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
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
• 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
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
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
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.

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