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Gas Practice

Stoichiometry and the Ideal Gas Law
1. What is the density of nitrogen gas (N2N2) at 248.0 Torr and 18º C?
(see answer for Example 10.5.1)
2. As temperature increases, density of a gas ___________. As pressure increases,
density of a gas _____________.
a. increases; increases
b. increases; decreases
c. decreases; increases
d. decreases; decreases
ANSWER: (un-highlight ) The answer is C. Use equation (10.5.1) and (10.5.2)
from the website to answer
3. What volume of carbon dioxide gas is produced at STP by the decomposition of
0.150 g CaCO3 via the equation: CaCO3(s)→CaO(s)+CO2(g)
(see answer for Example 10.5.2A)
4. Sulfuric acid, the industrial chemical produced in greatest quantity (almost 45 million
tons per year in the United States alone), is prepared by the combustion of sulfur in air to
give SO2, followed by the reaction of SO2 with O2 in the presence of a catalyst to give SO3,
which reacts with water to give H2SO4. The overall chemical equation is as follows:
What volume of O2 (in liters) at 22°C and 745 mmHg pressure is required to produce 1.00
ton (907.18 kg) of H2SO4?
(see answer for Example 10.5.2C)
5. Charles used a balloon containing approximately 31,150 L of H 2 for his initial flight in
1783. The hydrogen gas was produced by the reaction of metallic iron with dilute
hydrochloric acid according to the following balanced chemical equation:
How much iron (in kilograms) was needed to produce this volume of H2 if the temperature
was 30°C and the atmospheric pressure was 745 mmHg?
(see answer for Exercise 10.5.2)
6. Sodium azide (NaN3) decomposes to form sodium metal and nitrogen gas according
to the following balanced chemical equation:
This reaction is used to inflate the air bags that cushion passengers during automobile
collisions. The reaction is initiated in air bags by an electrical impulse and results in the
rapid evolution of gas. If the N2 gas that results from the decomposition of a 5.00 g sample
of NaN3 could be collected by displacing water from an inverted flask, what volume of gas
would be produced at 21°C and 762 mmHg?
(see answer for Example 10.5.3)
7. A 1.00 g sample of zinc metal is added to a solution of dilute hydrochloric acid. It
dissolves to produce H2 gas according to the equation
The resulting H2 gas is collected in a water-filled bottle at 30°C and an atmospheric
pressure of 760 mmHg. What volume does it occupy?
(see answer for Exercise 10.5.3)
Ideal Gas Law and Combined Gas Law
1. The balloon that Charles used for his initial flight in 1783 was destroyed, but we can
estimate that its volume was 31,150 L (1100 ft3), given the dimensions recorded at the
time. If the temperature at ground level was 86°F (30°C) and the atmospheric pressure
was 745 mmHg, how many moles of hydrogen gas were needed to fill the balloon?
(see answer for Example 10.4.1)
2. Suppose that Charles had changed his plans and carried out his initial flight not in
August but on a cold day in January, when the temperature at ground level was −10°C
(14°F). How large a balloon would he have needed to contain the same amount of
hydrogen gas at the same pressure as in Example 10.4.1?
(see answer for Example 10.4.2)
3. At a laboratory party, a helium-filled balloon with a volume of 2.00 L at 22°C is
dropped into a large container of liquid nitrogen (T = −196°C). What is the final volume
of the gas in the balloon?
(see answer for Exercise 10.4.2)
4. Aerosol cans are prominently labeled with a warning such as “Do not incinerate this
container when empty.” Assume that you did not notice this warning and tossed an
“empty” aerosol can (0.025 mol in 0.406 L, initially at 25°C and 1.5 atm internal
pressure) into a fire at 750°C. What would be the pressure inside the can (if it did not
(see answer for Example 10.4.3)
5. Suppose that a fire extinguisher, filled with CO2 to a pressure of 20.0 atm at 21°C at
the factory, is accidentally left in the sun in a closed automobile in Tucson, Arizona, in
July. The interior temperature of the car rises to 160°F (71.1°C). What is the internal
pressure in the fire extinguisher?
(see answer for Exercise 10.4.3)
6. We saw in Example that Charles used a balloon with a volume of 31,150
L for his initial ascent and that the balloon contained 1.23 × 103 mol of H2 gas initially at
30°C and 745 mmHg. Suppose that Gay-Lussac had also used this balloon for his
record-breaking ascent to 23,000 ft and that the pressure and temperature at that
altitude were 312 mmHg and −30°C, respectively. To what volume would the balloon
have had to expand to hold the same amount of hydrogen gas at the higher altitude?
(see answer for Example 10.4.4)
7. A steel cylinder of compressed argon with a volume of 0.400 L was filled to a
pressure of 145 atm at 10°C. At 1.00 atm pressure and 25°C, how many
15.0 mL incandescent light bulbs could be filled from this cylinder? (Hint: find the
number of moles of argon in each container.)
(see answer for Exercise 10.4.3)