# 13 gas laws by jianglifang

VIEWS: 7 PAGES: 69

• pg 1
```									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refrigeration systems, but has now been
replaced by other compounds that do not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.
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
32. Deduce the most useful equation for
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.

```
To top