1. AP Chem Today: Gas Behavior and Gas Laws Review
Test Corrections due by next Fri 10/19!

2. Kinetic Molecular Theory
Explains properties of gases, liquids, and solids in terms of energy using an ideal gas, an imaginary gas which fits all the assumptions of kinetic molecular theory

3. Kinetic Molecular Theory
Assumptions:
Gas particles are in continuous, rapid, random motion
The size of the gas particles is negligible compared to the distance between gas particles. Therefore, gases are mostly empty space
There are no attractive or repulsive forces between gas particles.

4. Kinetic Molecular Theory
Elastic collisions occur between molecules during collisions, meaning there is no loss of kinetic energy during collisions
The average kinetic energy of the molecules is proportional to the absolute temperature of the gas.
Note: An IDEAL GAS is THEORETICAL and is used to PREDICT the behavior of REAL GASES. The ASSUMPTIONS above are not true of REAL GASES.

5. Problems with KMT: Why gases are not ideal…
1. Gas atoms/molecules take up space (they are matter)! 2. Some intermolecular (attractive) forces exist between gas molecules 3. Small volume containers result in more collisions (b/c there’s less space to move) between gas molecules

6. Conditions in which a REAL GAS behaves MOST like an IDEAL GAS:
1. Low pressure (moves around more freely) 2. High temperature (fast moving) 3. Low intermolecular forces (not attracted to each other) 4. Small size 5. Large volume container (more space to move, less likely to collide)

7. Physical Characteristics of gases:
Gases do not have a definite shape or volume, they expand to fill a container
Gases have low density
Gases can be compressed

8. Variables that Define a Gas
Volume (V) - may be expressed in liters, milliliters, cm3, dm3.
Temperature (T) – Always expressed in Kelvin!!
= 298 K

9. Variables that Define a Gas
Number of Moles (n) - how many particles are present in the sample of gas
Pressure-the force per unit area on a surface
Exerted by all gases on any surface they collide with
Collisions cause pressure!

10. Pressure units: STANDARD PRESSURE:
1atm = 760mmHg = 760 torr = kPa = 1.01x105 Pa
𝟏𝟐𝟑 𝒌𝑷𝒂 𝟏 𝒂𝒕𝒎 𝟏𝟎𝟏.𝟑 𝒌𝑷𝒂
=𝟏.𝟐𝟏 𝒂𝒕𝒎

11. Measuring Pressure: Barometer
used to measure atmospheric (air) pressure
The higher the altitude the lower the atmospheric pressure and the lower the height of the mercury in the barometer
Whatever the height is, that is your pressure

12. Measuring Pressure of a Gas
Manometer-
measures the pressure of an enclosed sample
Can be open or closed
Closed Manometer

13. Pressure Open Manometer -h +h
Gas pressure is less than atmospheric pressure when the height of the liquid in the manometer is higher on the _______________. Therefore you will ________________ the height and the atmospheric pressure.
Gas pressure is greater than atmospheric pressure when the height of the liquid in the manometer is higher on the _______________. Therefore you will ________________ the height and the atmospheric pressure.
left side of the U
right side of the U
subtract
add

14. Examples P of gas sample P of gas sample = 757.8 - 655 = 757.8 + 105
1. If the atmospheric pressure is 757.8mmHg, what is the pressure of the gas in each of the following manometers?
P of gas sample =
= mm Hg
P of gas sample =
= mm Hg

15. Gas Laws Summary Table Name of Law Equation/Definition
Type of Relationship
Constant
Dalton’s Law
----
Graham’s Law
Boyle’s Law T
*Charles’ Law P
*Gay-Lussac’s Law V
*Combined Gas Law
Avogadro’s Law
*Ideal Gas Law

16. Dalton’s Law Ptotal = P1 + P2 + P3 + …
each individual gas behaves as if it were independent of the others.
the total pressure exerted by a mixture of gases is the sum of the individual pressures of each gas
Ptotal = P1 + P2 + P3 + …

17. Partial Pressures Ptotal = Pgas + PH2O
When one collects a gas over water, there is water vapor mixed in with the gas.
To find only the pressure of the desired gas, one must subtract the vapor pressure of water from the total pressure
Ptotal = Pgas + PH2O
Value depends on temperature

18. Partial Pressures and Mole Fraction
The mole fraction of an individual gas component in an ideal gas mixture can be expressed in terms of the component's partial pressure or the moles of the component:
The partial pressure of an individual gas component in an ideal gas can be obtained using this expression:

19. Dalton’s Law Ptotal=PN + PO Ptotal=250 + 300 Ptotal=550 mm Hg
1. Two gases such as oxygen and nitrogen are present in a flask at the following pressures. When combined, what is the pressure of the flask? PNitrogen =250. mm Hg, POxygen=300. mm Hg
Ptotal=PN + PO
Ptotal=
Ptotal=550 mm Hg

20. Dalton’s Law Examples Ptotal=Poxygen + PH2O 731=Poxygen + 17.5
3. Oxygen gas from the decomposition of potassium chlorate was collected by water displacement. The barometric pressure and the temperature during the experiment were 731.0torr and 20.0˚C respectively. What was the partial pressure of the oxygen collected? The vapor pressure of water at 20°C is 17.5 torr   
Ptotal=Poxygen + PH2O
731=Poxygen
Poxygen =713.5 torr

21. 𝟒𝟖𝟎 𝒈 𝑶 𝟐 𝟏 𝒎𝒐𝒍 𝑶𝟐 𝟑𝟐 𝒈 𝑶𝟐 =𝟏𝟓 𝒎𝒐𝒍
𝟖𝟎 𝒈 𝑯𝒆 𝟏 𝒎𝒐𝒍 𝑯𝒆 𝟒 𝒈 𝑯𝒆 =𝟐𝟎 𝒎𝒐𝒍
𝟑𝟓𝒎𝒐𝒍
𝟏𝟓 𝟑𝟓
𝟐𝟎 𝟑𝟓
𝟕 𝒙 𝟏𝟓 𝟑𝟓 =𝟑 𝒂𝒕𝒎

22. Graham’s Law of Diffusion
Under ideal conditions, the rates at which different gases diffuse (spread out) are inversely proportional to their molar masses.

23. Graham’s Law Examples
Compare the rates of diffusion of H2 and O2 gases at the same temperature and pressure.
H2 will diffuse more quickly than O2 because it has a smaller molar mass (H2 = 2.02 g/mol, O2 = g/mol)

24. Graham’s Law Examples 𝑀𝑜𝑙𝑎𝑟 𝑀𝑎𝑠𝑠 𝑈𝑛𝑘𝑛𝑜𝑤𝑛 𝐺𝑎𝑠=𝟏𝟔𝟑 𝒈/𝒎𝒐𝒍
4. A sample of hydrogen gas effuses through a porous container about 9 times faster than an unknown gas. Estimate the molar mass of the unknown gas.
𝑅𝑎𝑡𝑒 𝐻𝑦𝑑𝑜𝑔𝑒𝑛 𝑅𝑎𝑡𝑒 𝑈𝑛𝑘𝑛𝑜𝑤𝑛 𝐺𝑎𝑠 = 𝑀𝑜𝑙𝑎𝑟 𝑀𝑎𝑠𝑠 𝑈𝑛𝑘𝑛𝑜𝑤𝑛 𝐺𝑎𝑠 𝑀𝑜𝑙𝑎𝑟 𝑀𝑎𝑠𝑠 𝐻𝑦𝑑𝑟𝑜𝑔𝑒𝑛
9= 𝑀𝑜𝑙𝑎𝑟 𝑀𝑎𝑠𝑠 𝑈𝑛𝑘𝑛𝑜𝑤𝑛 𝐺𝑎𝑠 2.02
𝑀𝑜𝑙𝑎𝑟 𝑀𝑎𝑠𝑠 𝑈𝑛𝑘𝑛𝑜𝑤𝑛 𝐺𝑎𝑠=𝟏𝟔𝟑 𝒈/𝒎𝒐𝒍

25. Boyle’s Law P1V1= P2V2 Inversely proportional
Molecular explanation: decreasing volume reduces the amount of space gas molecules have to move  increases the # of collisions  increases the pressure

26. 𝑽𝟏 𝑻𝟏 = 𝑽𝟐 𝑻𝟐 Charles’s Law Directly proportional
Temperature must be in Kelvin!
Molecular explanation: increasing temperature causes the gas molecules to move faster ( kinetic energy)  molecules hit the container with more force  causes container to expand

27. 𝑷𝟏 𝑻𝟏 = 𝑷𝟐 𝑻𝟐 Gay-Lussac’s Law Directly Proportional
Temperature must be in Kelvin!!
Molecular explanation: increasing temperature causes the gas molecules to move faster ( kinetic energy)  molecules collide more frequently  pressure increases

28. Combined Gas Law
Relates pressure, temperature, and volume in one equation

29. STP
STP= Standard Temperature & Pressure 273 K & 1 atm

30. 𝟏𝟐.𝟑 𝑳 𝟐𝟎+𝟐𝟕𝟑 𝑲 = 𝑽𝟐 −𝟏𝟐+𝟐𝟕𝟑 𝑲
𝟏𝟐.𝟑 𝑳 𝟐𝟗𝟑 𝑲 = 𝑽𝟐 𝟐𝟔𝟏 𝑲
V2 = L

31. Avogadro’s Law 𝑛 1 𝑉 1 = 𝑛 2 𝑉 2 (n = # of moles)
The volume of a gas at constant temperature and pressure is directly proportional to the number of moles of the gas.
Equal volumes of gases at the same pressure and temperature contain the same # of moles and molecules
1 mole of any gas at STP is 22.4L

32. PV=nRT Ideal Gas Law n= # moles R=universal gas constant
Describes the state of a hypothetical ideal gas.
Good approximation of the behavior of many gases under many conditions.
PV=nRT
n= # moles
R=universal gas constant
Will be GIVEN on tests/quizzes!
R is the universal gas constant.
To determine which R to use in your equation you have to look at the _____________.
Pressure units being used!!

33. (1.046 atm)V= (4.5 mol)(0.0821)( K)
V = 105 L