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Honors Chemistry
Julien
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Gases
Chapter 11
Gases and Their Properties
4. The kinetic molecular theory describes the motion of atoms and molecules and explains the
properties of gases. As a basis for understanding this concept:
a. Students know the random motion of molecules and their collisions with a surface create the
observable pressure on that surface.
b. Students know the random motion of molecules explains the diffusion of gases.
c. Students know how to apply the gas laws to relations between the pressure, temperature, and volume
of any amount of an ideal gas or any mixture of ideal gases.
d. Students know the values and meanings of standard temperature and pressure (STP).
e. Students know how to convert between the Celsius and Kelvin temperature scales.
f. Students know there is no temperature lower than 0 Kelvin.
g.* Students know the kinetic theory of gases relates the absolute temperature of a gas to the average
kinetic energy of its molecules or atoms.
h.* Students know how to solve problems by using the ideal gas law in the form PV = nRT.
i.* Students know how to apply Dalton’s law of partial pressures to describe the composition of gases
and Graham’s law to predict diffusion of gases.
I. Properties of gases.
Read Section 11.1 in your eText! I will be able to use the kinetic molecular theory of gases to
describe the properties of gases.
Do the Properties of Gases Tutorial!
Do the Properties of Gases Self Study Activity!
A. Kinetic Molecular Theory of Gases.
1. A gas consists of small particles (atoms or molecules) that move randomly with high
velocities.
a. Gas molecules moving in random directions at high speeds cause a gas to fill the entire
volume.
2. The attractive forces between the particles of a gas are usually very small.
a. Gas particles are far apart and fill a container of any size and shape.
3. The actual volume occupied by gas molecules is extremely small compared to the volume
that the gas occupies.
a. The volume of the gas is considered equal to the volume of the container.
b. Most of the volume of a gas is empty space, which allows gases to be easily compressed.
4. Gas particles are in constant motion, moving rapidly in straight paths.
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Chapter 11 Notes from Basic Chemistry
a. When gas particles collide, they rebound and travel in new directions.
b. Every time they hit the walls of the container, they exert pressure.
c. An increase in the number or force of collisions against the walls of the container causes
an increase in the pressure of the gas.
5. The average kinetic energy of gas molecules is proportional to the Kelvin temperature.
a. Gas particles move faster as the temperature increases.
b. At higher temperatures, gas particles hit the walls of the container more often and with
more force, producing higher pressures.
6. Kinetic Molecular Theory of Gases—
a. Molecules from a perfume bottle move at about 1000 miles per hour when escaping from
the bottle at room temperature.
B. Pressure (P).
1. Gas particles collide with the walls of the chamber or surface that contains the gas, creating
pressure.
2. Atmospheric pressure—
C.
D.
E.
1.
2.
3. Gas pressure is measured atmospheres (atm), millimeters of mercury (mmHg), torr, Pascals (Pa),
kiloPascals (kPa), and inches of mercury.
Volume (V).
1. The volume of a gas is dependent on the size of the container or the collisions of the gas with the
walls of a flexible container.
2. The most common units of volume are liters (L) and milliliters (mL).
Temperature (T).
1. The temperature of a gas is related to the kinetic energy of its particles and based on temperature
that starts at absolute zero.
2. Temperature has only one unit and that is Kelvin (K).
Amount of gas (n).
Air can be measured by the number of particles and this is measured in the lab in terms of grams.
The grams must be converted to moles (mol) to relate it to the number of particles.
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Chapter 11 Notes from Basic Chemistry
Problem 1: When helium is added to a balloon, the mass of the gas, in grams, increases. What property of gas
is described?
II. Gas Pressure.
Read Section 11.2 in your eText! I will describe the units of measurement used for pressure,
and change from one unit to another.
Do the Converting Between Units of Pressure Tutorial!
A. Pressure.
1. Pressure—
Pressure =
force
area
2. Atmosphere (atm)—
3. If 760 mmHg equals 760 torr, then 1 mmHg equals 1 torr.
4. Using the SI system, pressure is measured in pascals (Pa), so 1 atm equals 1.01325 x 105 Pa
or 101.325 kPa or 14.7 pounds per square inch (psi).
B. Barometer.
1. Early barometers were based on a column of mercury that is 29.9 inches high under standard
conditions and taller under higher pressure and shorter under lower pressure.
2. Altimeters are based on barometric readings as are pressure gauges used by scuba divers
Do the Scuba Diving and Blood Gases Case Study!
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Chapter 11 Notes from Basic Chemistry
You are now able to do problems Converting Between Units of Pressure Tutorial in your first homework
assignment.
III. Pressure and volume (Boyle’s Law).
Read Section 11.3 in your eText! I will use the pressure-volume relationship (Boyle’s Law)
to determine the new pressure or volume of a certain amount of gas at a constant
temperature.
A. Boyle’s Law.
1. Boyle’s Law—
P1V1 = P2V2
Do the Pressure and Volume Tutorial!
Guide to Using the Gas Laws
STEP 1: Organize the data in a table of initial and final conditions.
STEP 2: Rearrange the gas law for the unknown.
STEP 3: Substitute values into the gas law and solve for the unknown.
Problem 2: A sample of helium gas has a volume of 312 mL at 648 torr. If the volume expands to 825 mL at
constant temperature, what is the new pressure in torr.
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Chapter 11 Notes from Basic Chemistry
Problem 3: A sample of methane gas, CH4, has a volume of 125 mL at 0.600 atm pressure and 25 °C. How
many milliliters will it occupy at a pressure of 1.50 atm and 25 °C?
You are now able to do problems 11.19 and 11.20 in your first homework assignment.
IV. Temperature and volume (Charles’ Law).
Read Section 11.4 in your eText! I will use the temperature-volume relationship (Charles’
Law) to determine the new temperature or volume of a certain amount of gas at a constant
temperature.
Do the Temperature and Volume Tutorial!
A. Charles’ Law.
1. Charles’ Law—
2. All temperatures used in gas law calculations must be converted to their corresponding Kelvin
(K) temperatures.
V1 V2
=
T1
T2
Problem 4: A mountain climber inhales 486 mL of air at a temperature of –8 C. What volume, in mL, will the
air occupy in the lungs if the climber’s body temperature is 37 °C?
You are now able to do problems 11.25 and 11.28 in your first homework assignment.
V. Temperature and pressure (Gay-Lussac’s Law).
Read Section 11.5 in your eText! I will use the temperature-pressure relationship (GayLussac’s Law) to determine the new temperature or pressure of a certain amount of gas at a
constant volume.
Do the Temperature and Pressure Tutorial!
A. Gay-Lussac’s Law.
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Chapter 11 Notes from Basic Chemistry
1. Gay-Lussac’s Law—
2. All temperatures used in gas law calculations must be converted to their corresponding Kelvin
(K) temperatures.
P1
P
= 2
T1
T2
Problem 5: In a storage area where the temperature has reached 55 °C, the pressure of oxygen gas in a 15.0-L
steel cylinder is 965 torr. To what Celsius temperature would the gas have to be cooled to reduce
the pressure to 850. torr?
B. Vapor pressure and boiling point.
1. Vapor pressure—
2. As temperature increases, more vapor forms, and vapor pressure
increases.
3. A liquid reaches its boiling point when its vapor pressure becomes
equal to the external pressure.
4. At higher altitudes, atmospheric pressures are lower and the
boiling point of water is lower than 100 °C.
5. Pressure cookers and autoclaves are sealed and increase the
pressure, which increases the boiling point of the liquid.
You are now able to do problems 11.30, 11.31 and 11.36 in your first
homework assignment.
VI. The combined gas law.
Read Section 11.6 in your eText! I will use the combined gas law to find the new pressure,
volume, or temperature of a gas when changes in two of these properties are given.
Do the Combined Gas Law Tutorial!
A. The Combined Gas Law.
1. The Combined Gas Law—
P1V1
PV
= 2 2
T1
T2
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Chapter 11 Notes from Basic Chemistry
Problem 6: A weather balloon is filled with 15.0 L of helium at a temperature of 25 °C and a pressure of 685
mmHg. What is the pressure, in mmHg, of the helium in the balloon in the upper atmosphere
when the temperature is –35 °C and the volume becomes 34.0 L?
You are now able to do problems 11.38 and 11.41 in your first homework assignment.
VII. Volume and moles (Avogadro’s Law).
Read Section 11.7 in your eText! I will use Avogadro’s law to describe the relationship
between the amount of a gas and its volume, and use this relationship in calculations.
Do the Volume and Moles Tutorial!
A. Avogadro’s Law.
1. Avogadro’s Law—
V1 V2
=
n1
n2
Problem 7: A sample containing 8.00 g of oxygen has a volume of 5.00 L. What is the volume (L) after 4.00 g
of oxygen is added to the 8.00 g of oxygen in the balloon, if the temperature and pressure do not
change?
B. STP and molar volume.
1. STP—
a. Standard temperature is exactly 0 °C (273 K).
b. Standard pressure is exactly 1 atm (760 mmHg).
2. Molar volume—
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Chapter 11 Notes from Basic Chemistry
1 mol gas (STP) = 22.4 L
22.4 L
1 mol gas (STP)
1 mol gas (STP)
22.4 L
Problem 8: What is the volume (L) of 5.10 g He at STP?
C. Density of a gas at STP.
1. Because of Avogadro’s Law, the volume of one mole of gas at STP is known to be 22.4 L.
2. The formula for the density of any gas at STP can be calculated using the mass of a mole (mass
found in the Blue Sheet) divided by the volume of a mole (22.4 L).
3. Calculating the density of carbon dioxide (44.00 g/22.4 L or 1.96 g/L) explains why balloons
filled with this gas settle on the floor and balloons of oxygen (32.00 g/22.4 L or 1.43 g/L) will
float just above them and helium balloons (4.00 g/22.4 L or 0.179 g/L) will rise rapidly because
the density of air is 1.29 g/L.
Problem 9: What is the density of hydrogen gas (H2) at STP?
You are now able to do problems 11.50 and Avogadro’s Law Simulation in your second homework
assignment.
VIII. The Ideal Gas Law.
Read Section 11.8 in your eText! I will use the ideal gas law to solve for P, V, T, or n of a gas
when given three of the four values in the ideal gas law. Calculate density, molar mass, or
volume of a gas in a chemical reaction.
Do the Introduction to the Ideal Gas Law Tutorial!
A. The Ideal Gas Law.
1. The Ideal Gas Law—
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PV = nRT
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Chapter 11 Notes from Basic Chemistry
PV
nT
(1 atm)(22.4 L)
0.0821 atm iL
For one mole at STP: R =
=
(1.00 mol)(273 K)
mol iK
If we rearrange the equation: R =
2. Universal gas constant, R,—
Problem 10: Chlorine gas, Cl2, is used to purify the water in swimming pools. How many moles of chlorine
gas are in a 7.00-L tank if the gas has a pressure of 865 mmHg and a temperature of 24 K?
Problem 11: What is the volume of 1.20 g of carbon monoxide at 8 °C if it has a pressure of 724 mmHg?
B. Molar mass of a gas.
1. The molar mass of a gas can be determined by determining the number of moles of a gas and
then dividing the mass of the gas by the number of moles.
Do the Introduction to the Ideal Gas Law Self Study!
Problem 12: What is the molar mass of an unknown gas in a 1.50-L container if 0.488 g of the gas has a
pressure of 0.0750 atm at 19.0 °C?
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Chapter 11 Notes from Basic Chemistry
You are now able to do problems 11.58 and 11.86 in your second homework assignment.
IX. Gas laws and chemical reactions.
Read Section 11.9 in your eText! I will determine the mass or volume of a gas that reacts or
forms in a chemical reaction.
Do the Ideal Gas Law and Stoichiometry Tutorial!
A. Gas laws and chemical reactions.
1. The Ideal Gas Law can be incorporated into a stoichiometry reaction.
a. Convert the given mass into moles.
b. Finding the moles of the wanted gas will set you up to use the Ideal Gas Law.
c. The moles of the wanted gas will become the “n” factor in the equation and volume can be
determined.
Guide to Reactions Involving Gases
STEP 1: Calculate the moles of given using molar mass or ideal gas law.
STEP 2: Determine the moles of needed using a mole-mole factor.
STEP 3: Convert the moles of needed to mass or volume using molar
mass or ideal gas law.
Problem 13: If 12.8 g of aluminum reacts with HCl, how many liters of H2 would be formed at 715 mmHg
and 19 °C?
2Al(s) + 6HCl(aq)  2AlCl3(aq) + 3H2(g)
You are now able to do problems 11.60, 11.62, 11.64, and 11.104 in your second homework assignment.
X. Partial Pressures (Dalton’s Law).
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Chapter 11 Notes from Basic Chemistry
Read Section 11.10 in your eText! I will use partial pressures to calculate the total pressure
of a mixture of gases.
Do the Mixtures of Gases Tutorial!
A. Partial Pressures (Dalton’s Law).
1. Partial pressure—
2. Dalton’s law—
Ptotal = P1 + P2 + P3 + …
B. Air is a gas mixture.
1. The air you breathe is approximately 78.2% nitrogen and 21.0% oxygen.
2. Carbon dioxide, argon, and water make up about 0.8%.
Guide to Solving for Partial Pressure
STEP 1: Write the equation for sum of partial pressures.
STEP 2: Solve for the unknown pressure.
STEP 3: Substitute known pressures and calculate the unknown.
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Chapter 11 Notes from Basic Chemistry
Problem 14: An anesthetic consists of a mixture of cyclopropane gas, C3H6, and oxygen gas, O2. If the
mixture has a total pressure of 1.09 atm, and the partial pressure of the cyclopropane is 73
mmHg, what is the partial pressure of the oxygen in the anesthetic?
C. Gases collected over water.
1. Collecting gas is hard to do because it can easily escape.
2. Bubbling a gas into an inverted container, filled with water can capture the gas the gas without
letting it escape.
3. The water vapor that occurs because of evaporation must be subtracted from the pressure reading
of the collected gas.
Guide to Gases Collected over Water
STEP 1: Obtain the vapor pressure of water.
STEP 2: Subtract vapor pressure from total P of gas mixture to give
partial pressure of needed gas.
STEP 3: Use ideal gas law to convert Pgas to moles or grams of gas
collected.
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Chapter 11 Notes from Basic Chemistry
Problem 15: A 456-mL sample of oxygen gas (O2) was collected over water at a pressure of 744 mmHg and a
temperature of 20. °C. How many grams of dry O2 were collected?
You are now able to do problems 11.66 in your second homework assignment.
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Chapter 11 Notes from Basic Chemistry