1. Gases
Chapter 11

2. Gases Gases are made of _______________________.
These particles move in _______________ at varying speeds until they collide with another particle or a barrier.

3. Gases A gas has ______________________
The Gas Laws pertain to an _____________, an imaginary gas that serves as a model for gas behaviors.

4. Is it real or ideal?
Ideal Conditions/Ideal Gas -
imaginary gas that fits the all assumptions of the kinetic-molecular theory
Real Gas –
particles have size and intermolecular attraction to other particles, so do not conform to gas laws under very _____ pressure or very _____ temperatures.

5. Kinetic Theory based on idea that particles are always in ________
At _______ temperature, all gases have the ________ kinetic energy (KE)
But gases do not all have same ______

6. Rearrange to fastest to slowest
KE = __________
m = mass
v = velocity
Rearrange to fastest to slowest
O CO H N2
_____________________
Why??
Well who’s faster, a 350 lb football player or a 120 lb runner?

7. ALL gases at the SAME temperature have the same average KE
If the KE = 20, note the speed for the gases . H2
KE = ½ m v2
20 = ½ 2g v2
20 = 1 v2
20/1 = v2
20 = v2
v2 = /20
v = Ar
KE = ½ m v2
20 = ½ 40g v2
20 = 20 v2
20/20 = v2
1 = v2
v2 = /1
v = 1

8. Kinetic molecular theory assumptions:
1.   Gases consist of large # of ___ molecules that are ___ apart relative to their size. Thus they have ___________ and lots of empty space between them.

9. Kinetic molecular theory assumptions:
2.    Lots of _________________
- don’t stick together after collisions.
No net loss of _____________
3.    NO _________ or __________ for each other (ideal gases do not condense to a liquid or a solid as no attraction)

10. Kinetic molecular theory assumptions: (cont)
4.    ____________ = energy of motion
Gas particles are in constant, rapid motion that is random. Therefore, passes “kinetic energy” = energy of motion.
5.    Average kinetic energy depends on the _______________ of the gas.

11. The Nature of Gases for Ideal Gases
1. Expand to fit container - no definite ___________ or __________
2. ______ - slide past each other – fluid as attraction forces not significant
3. ____________ - particles far apart
4. _______________ – particles can be pressed close together.

12. Nature cont. Rate of Diffusion depends on: a. ____________
5. ___________ - random mixing caused by random motion
Rate of Diffusion depends on:
a.   ____________
b.   ______________
c. ______________ between particles
O vs H2O

13. Nature cont. Smaller molecule → __________________________
6. _____________ - gas particles under pressure pass thru tiny openings
Smaller molecule → __________________________
(Does a balloon stay full forever?
Why not?)

14. To describe what a gas is doing,
4 measurable qualitites are used
          ________________
         ________________
       ________________
        ________________

15. Pressure Pressure = _______ (Force is in _____________
Force is caused by gas molecules hitting wall - their collisions
Air around us exerts a pressure on us.
What would happen to your eyeballs if you were to go out into outer space??
Where is the most pressure - sea level, ocean bottom, mountain top?

16. Barometer Note air pushing down
Vacuum has ________ inside, so is not pushing down

17. A simple manometer.

18. Standard Temperature & Pressure (STP) -- must be standard to compare things STP = ______________________________________________________ pressure = __________________________________________________________

19. Conversions Standard Pressure with different units 1 atmosphere (atm)
=760 mm Hg (millimeter of mercury)
=760 torr =14.7 lb/in2=76 cm Hg
=29.9 in Hg = 1013 millibar (mbar)
= 101,325 Pascal = KPa (SI unit)
Pascal - the SI unit = 1N/ m2

20. These all equal each other, so can use as conversion factors
Given 700 mm Hg, how many atm?

21. Covert the following
860 mm Hg = ______ atm
720 torr = ______ in Hg
0.89 atm = ______ mm Hg
1.2 atm = ______ in Hg

22. The Gas Laws
______________ -the volume of a fixed mass of gas varies ___________ with the pressure at constant temperature.
So as P __________  V ________ !

23. As pressure increases, the volume decreases
Temperature and number of moles are constant

24. The effects of decreasing the volume of a sample of gas at constant temperature. Molecules hit more often, so Pressure ____________

25. _____________________
Boyle’s Law
P1V1 = P2V2 initial = final
_____________________

26. Boyle’s law
Ex. A sample of O2 has a volume of 100 ml and a pressure of 200 torr. What will it’s volume be if the pressure is increased to 300 torr (T is kept constant)
First think – what is the relationship?
P ___________, then V ___________

27. Boyle’s law
P1 x V1 = P2 x V2
V2 =

28. Plotting Boyle's data from Table 5. 1
Plotting Boyle's data from Table 5.1. (a) A plot of P versus V shows that the volume doubles as the pressure is halved. (b) A plot of V versus 1/P gives a straight line. The slope of this line equals the value of the constant k.

29. A plot of PV versus P for several gases at pressures below 1 atm.

30. deals with ________ and _____________ (Pressure constant).
Charles Law
deals with ________ and _____________
(Pressure constant).
V and T (in Kevin) are ________________
As you _______________ the temperature - molecules MOVE __________
KE (Kinetic energy) _____________ as molecules jump around more

31. Charles Law Found gases had “zero” volume at -273 oC
Temperature MUST be in ____________
Found gases had “zero” volume at -273 oC
So named this _________________, which equals 0 Kelvin = 0 K
K = ____ oC
oC = ___ K

32. Charles Law V1 = T1 _________________ V2 T2 Can rewrite: V1T2 = V2T1
(Watch 1's & 2's)

33. The effects of increasing the temperature of a sample of gas at constant pressure. Molecules are moving ____________ so they hit the container harder, thus size must _________ to keep P same

34. Charles
Ex. A sample of neon gas occupies a volume of 752 ml at 25 oC. What volume will it occupy at 50 oC.
Pressure is constant.
Think first - as Temperature increases, Volume will ____________

35. Charles MUST change oC to Kelvin Given : V1 = 752 ml
T1 = 25  oC = 298 K
V2 = ?
T2 = 50 oC = 323 K

36. Charles
V1 x T2 = V2 x T1
V2 =

37. Gay- Lussac’s Law
the pressure of a fixed mass of gas at constant volume varies ___________ with the Kelvin temperature
P = T1
P T _________________
Or P1 x T2 = P2 x T1

38. The effects of increasing the temperature of a sample of gas at constant volume. Molecules _____ _______, so pressure __________

39. Ex. A gas in an aerosol can is at a pressure of 3. 00 atm at 25 ° C
Ex. A gas in an aerosol can is at a pressure of 3.00 atm at 25 ° C. Directions warn the user not to keep the can in a place where temperature exceeds 52 ° C. What would the pressure in the can be at 52 oC?
Given
P1 = atm
P2 = ?
T1 = 25 oC or 298 K
T2 = 52 oC or 325 K

40. Combined Gas Law P1V1 = P2V2 T1 T2 or P1V1 T2 = P2V2 T1
Expresses the relationship between ________, _______, and __________ of a fixed amount of gas.
P1V = P2V2
T T2 or
P1V1 T2 = P2V2 T1

41. Ex. A helium balloon has a volume of 50. 0 L at 25 ° C and 1. 08 atm
Ex. A helium balloon has a volume of 50.0 L at 25 ° C and 1.08 atm. What volume will it have at 10 oC and atm?
Given:
V1 = 50.0L
V2 = ?
P1 = atm
P2 = atm V2 =
T1 = 25 ° C = K
T2 = 10 ° C = 283 K

42. Ideal Gas Equation
_____________ - equal volumes of gases at the same temperature and pressure contain equal numbers of molecules
Thus, the volume occupied by one mole of gas at ______ = standard molar volume of a gas = ________

43. Ideal Gas Equation
Ex. A chemical reaction produces mol of O2. What volume will it occupy at STP?
Given: mol O2
1 mol = L at STP

44. Ex. A chemical reaction produces 98. 0 mL of SO2 at STP
Ex. A chemical reaction produces 98.0 mL of SO2 at STP. What mass in grams was produced?
Given: Volume SO2 at STP 98.0 mL = L
0.098 L SO mol = mol SO2
22.4L
mol SO g = g SO2
1 mole
OR L SO mole g = g SO2
22.4 L mole

45. the constant, R, is the ideal gas constant has a value of .0821 L atm
Ideal Gas Law
the mathematical relationship of pressure, volume, temperature, and the number of moles of gas.
___________________
pressure x volume = # of moles x constant x temperature
the constant, R, is the ideal gas constant has a value of L atm
mole K
WATCH units !!

46. The effects of increasing the number of moles of gas particles at constant temperature and pressure.

47. Plots of PV/nRT versus P for several gases (200 K).

48. Plots of PV/nRT versus P for nitrogen gas at three temperatures.

49. Ex. What is the pressure in atm exerted by 0
Ex. What is the pressure in atm exerted by 0.50 mol sample of N2 in a 10.0L container at 298 K?
Given: V of N2 = 10.0L
n of N2 = mole
T = 298 K
P = ?
PV = nRT therefore, P = nRT V

50. Combining the ideal gas law with density - to find molar mass or gas density.
Ex. What is the density of ammonia gas at
63 oC and 705 mm Hg?
Given: P = 705 mmHg
V = ?
n = assume 1 mole
R = L atm /mole K
T = 63 oC

51. Combining the ideal gas law with density - to find molar mass or gas density.
V = 1mole( L atm / mole K) ( 336 K)
.928atm =
D = mass mass of 1 mol NH3 = 17g
volume (from Periodic table)
D = 17g
29.73L =

52. Gas Stoichiometry
Stoichiometry – the mass relationship between reactants and products in a chemical reaction. In general, follow these rules:
a. Use stoichiometry to do mole or mass conversions (or if STP, 1 mole = 22.4 L).
b. Use the Ideal Gas Law (PV=nRT) to convert between moles and volume.

53. Gas Stoichiometry Practice
CaCO3 → CaO + CO2
How many grams of calcium carbonate must be decomposed to produce 5.00 L of carbon dioxide at STP?
Do this two ways:
1st way: Use PV = nRT first, solve for n, then stoichiometry.
2nd way: Use stoich & 1 mole = 22.4 liters

54. Gas Stoichiometry Practice
Calculate the volume of H2 gas that can be obtained under laboratory conditions of temperature and pressure of 25 C and atm when 5.00 g of sodium is reacted with water:
2 Na + H2O  H2 + 2 NaOH

55. Gas Mixtures and Partial Pressures
Many gases are a mixture of gases; air being a prime example. We define the ______________________ of a gas which is the pressure a component of a gas mixture would have if it were all by itself in the same container.

56. Dalton’s Law
John Dalton (atomic theory) extended the gas laws simply as:
Ptotal = P1 + P2 + P3 + …
This simply says that the sum of the ____________________ of the gases will add up to the total pressure. This way we can treat a mixture of gases just like a pure gas.

57. Partial Pressure Practice
What would be the final pressure of the mixture of gases for the processes depicted in each of the following illustration?

58. Effusion and Diffusion
__________________________ - the rates of effusion of gases at the same temperature and pressure are inversely proportional to the square roots of their molar masses.
Rate of effusion A = Mb
Rate of effusion B Ma

59. The effusion of a gas into an evacuated chamber.

60. Ex. Compare the rates of effusion of H2 and O2 at same temperature and pressure.
Given:
H2 mol weight 2
O2 mol weight 32
Rate of effusion H2 = √ Mass O2
Rate of effusion O √ Mass H2

61. Important Gas relationships
As volume increases, pressure ____________ at constant temperature
As temperature increases, pressure _________ at constant volume
As temperature increases, volume _________ at constant pressure
Standard Temperature and Pressure - to compare gases use (standard temperature & Pressure _____)
Standard temperature = _____ C
Standard Pressure = _____ mm Hg = ___ atm average barometric pressure at sea level