Chapter 10
Gases
Conditions of ideal gases:
• Ideal gases have no attractive forces between the molecules.
• the atoms volume taken into account when looking at the volume a gas
occupies.
• Low pressure and high temperature conditions
Measurement of pressure:
• Barometer
• Manometer
• Units
Relationship of pressure and volume (Boyle’s Law)
Relationship of Temperature and Volume (Charles’ Law)
Relationship of Quantity and Volume (Avogadro’s Law)
Derivation of the Ideal Gas Law and the R constant
Values for the R constant:
Applications of the Ideal Gas Law
Gas Densities and Molar Masses
Volumes of Gases in Chemical Reactions (Stoichiometric relationships)
STP and same pressure/ temperature
Example1: The industrial synthesis of nitric acid involves reaction of nitrogen
dioxide gas with water. (Nitrogen monoxide gas is also produced) How many
liters of nitrogen dioxide can be produced if 5 liters of nitrogen dioxide react at
STP
Example 2:
The industrial synthesis of nitric acid involves reaction of nitrogen dioxide gas
with water. (Nitrogen monoxide gas is also produced) How many liters of
nitrogen dioxide can be produced if 5 liters of nitrogen dioxide react at 5.00 atm
and 298 K?
Varying temperature and pressure
Example 3:
Ammonia reacts with oxygen at 850C and 5.00 atm. The nitrogen monoxide
produced is sent across a collection tube to a container at a temperature of 25C
and 1atm. (NO remains a gas). How many liters of NO will be produced if 2 liters
of ammonia gas is used? (water is also a product in this reaction)
Dalton’s Law of Partial Pressures
Collecting gas over water
Example 4:
A sample of KClO3 is partially decomposed producing oxygen gas that is
collected over water. The volume of the gas collected is 0.250L at 26 and 765
torr.
a. How many moles of O2 are collected?
b. How many grams KClO3 were decomposed?
c. When dry, what volume would the collected oxygen gas occupy at the
same temperature and pressure?
Mixing gases
Example 5:
Consider the arrangement of bulbs as shown in the figure below. Each of the
bulbs contain a gas at the pressure shown. What is the pressure of the system
when all the stopcocks are opened, assuming the temperature remains constant.
We can neglect the volume of the capillary tubes connecting the bulbs.
Kinetic Molecular Theory
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Gases consist of molecules in continuous random motion.
The volume of the molecule negligible compared to the total volume the gas
occupies.
Attractive and repulsive forces negligible.
Energy can be transferred between molecules during collisions, but the
average KE doesn’t change.
The average KE is proportional to temperature.
Molecular effusion and diffusion (Graham’s Law)
Average kinetic energy of any molecule can be found by ½ mu2 where m = mass
of particle and u is the speed. If a particle is light, its rms should be high.
Therefore the KE should be relatively the same.
Rms can be calculated using the expression
Example 6:
Cory releases methane gas in class. Calculate the rms of these gas
particles.
Diffusion (spread of gas across a room)
Effusion (gas out of a small opening)
Graham’s Law of Effusion:
(Derive from rms of two gases)
Example 7:
An unknown gas of a homonuclear diatomic molecule effuses at arate that is
0.355 times that of oxygen at the same temperature. What is the identity of the
unknown gas?
Real Gases: The Van der Waals equation.
Real gases have finite volumes and they do have attractions.
We need to add corrections to the ideal gas law to account for the volume of
molecules AND molecular attractions. Since each gas has a different molecular
size and different intermolecular attractions, each gas needs its own correction
factor.
a and b are constants that vary for each gas.
(Table 10.3 on page 373 gives van der Waals constants for gas molecules.)
P = nrt
----- V – nb
n2a
----V2
Example 8: Consider a sample of 1.00 mol of CO2 confined to a volume of
3.00 L at 0.0 C. Calculate the pressure of the gas using
(a) ideal gas law
(b) van der Waals equation