FAQs about HONORS CHEMISTRY EXAM 5: GASES
1. What does the test cover?
All the material in the Unit 5 study guide and homework, plus review from earlier units (see practice
test).
2. How long is the test?
There are 25 multiple choice questions and 5 free response questions (4 calculation questions and 1
question on Lewis structures). The test is worth 100 points.
3. What can I use during the test?
You many use a calculator for the free-response part of the test. I will provide the ideal gas equation
(both forms), value of R and standard molar volume (see the practice test). Of course you can use a
periodic table.
4. Are there any gas law calculations?
Are you kidding? Review homework questions 1-9 & 23 for basic gas properties (like pressure
readings and pressure and temperature conversions), and 4, 10-16 for gas law problems. Don’t forget
to change temperatures into Kelvins, and match units pressure and volume units within the problem.
5. What about stoichometry?
Of course! Look at homework questions 17-22. You need to know what STP means (including
values), and how to use it in stoichiometry problems.
6. So the test is mostly just calculations, right?
Not entirely. There are some conceptual questions about gas behavior, and of course the review is
mostly atoms, molecules, periodic table, balancing equations. Do the ALL of the practice test.
7. Do I have to memorize much?
You do need to know that 1 atm = 760 mm Hg, that one mole of any gas = 22.4 L at STP, and the
values for temperature and pressure at STP. And how to change into Kelvins.
8. Any hints?
It’s very helpful to be able to visualize molecules in constant motion to explain gas properties such as
pressure and volume, and to see temperature as a measure of molecular motion. Oh, and don't forget
to change temperatures into Kelvins for all gas law calculations!!
PRACTICE TEST FOR EXAM 5: GASES
For full credit, you must show all setups with units and round off the answers appropriately.
P1V1 P2V2
Standard molar volume = 22.4!L mol
PV = nRT
=
n1T1 n2T2
atm!L
1L = 1000mL
1!atm = 760!mmHg
R = 0.08206
mol!K
1. a. Explain each of these in molecular terms:
i. gas pressure
ii. gas temperature
iii. why gas pressure increases when the container size is decreased
iv. why gas pressure increases when the temperature is increased
v. why a gas always fill its entire container
b. O2 gas is collected by displacing water from a collection tube. The total pressure in the tube is 742
mm Hg; the water vapor pressure in the experiment is 23 mm Hg. What is the pressure of the O2?
c. Two identical containers are at the same temperature and pressure; one contains Ne gas and the
other, Xe gas. Compare the kinetic energies, masses, velocities, and number of particles in the
two containers.
d. For each container, the difference between the mercury levels is given below the container. Draw
in the mercury in the U-tube. If the atmospheric pressure is 760 mm Hg, what is the pressure of
the gas in each of these containers:
left 35 mm above right
levels same
left 40 mm below right
e. If containers of the following gases (all at the same temperature) were opened at the same time,
which would reach you first, and why?
a. Cl2 (71 g/mol) b. H2S (34 g/mol) c. NH3 (17 g/mol) d. all of them at the same time
2. Gas law problems:
a. A balloon filled with 635 mL of oxygen gas at 23 °C is placed in a freezer, where it cools to –10
°C. What is the volume of the cold balloon? The pressure and amount of gas remain constant.
b. A small gas cylinder contains 3.22 L of argon at 11.7 atm pressure. What is the volume of the gas
at 1.05 atm? Assume temperature and amount of gas remain constant.
c. 325 mL of air at room pressure (765 mm Hg) are compressed with a piston to a volume of 42 mL.
What is the pressure of the compressed air? The temperature and amount of air remain constant.
d. A He weather balloon has a volume of 18 L at the earth's surface, where the temperature is 25 °C
and the pressure is 750 mm Hg. What is the volume of the balloon when it reaches an altitude
where the temperature is –22 °C and the pressure is 375 mm Hg? The amount of He is constant.
e. What is the volume of 0.085 mol hydrogen gas at 0.97 atm and 21 °C?
f. A canister of gas at 2.8 atm and 75 °C is cooled to 0 °C. If the volume and amount of gas are
constant, what is the pressure in the cold bottle?
g. A sample of nitrogen gas in a 275 mL container at 0.82 atm pressure and 26 °C is transferred to
larger container with volume 750 mL . The temperature in the larger container is now 17 °C.
What is the pressure of the gas in the larger container? The amount of nitrogen is constant.
h. How many moles of chlorine gas are in a 345 L tank at 7800 mm Hg and 27 °C?
3. a. What does STP stand for? What are the STP conditions? Include units!
b. What is the volume of 0.045 mole H2 gas at STP?
c. How many moles are in 25.4 mL of O2 gas at STP?
4. Stoichiometry problems
a. How many mL of CO2 gas at STP can react with 0.18 g LiOH in this reaction:
2 LiOH (s) + CO2 (g) → Li2CO3 (s) + H2O (l)
b. A car airbag inflates when sodium azide decomposes explosively to produce nitrogen gas:
2 NaN3 (s) → 2 Na (l) + 3 N2 (g)
How many grams of NaN3 must decompose to produce 41 L of N2 at STP?
c. How many L of CO2 gas at STP will form when 50.0 g of propane (C3H8) burn?
C3H8 (g) + 5 O2 (g) → 3 CO2 (g) + 4 H2O (l)
ANSWERS to PRACTICE TEST FOR EXAM 5
1. a. i. Gas pressure is caused by collisions of gas molecules with container wall.
ii. Gas temperature is proportional to molecular kinetic energy.
iii. Molecules have less space to bounce around, so they collide with container walls more often.
iv. Molecules are moving faster, so they strike container walls more often and with greater force.
v. Molecules are in constant random motion, so they eventually distribute throughout entire
container, no matter how large it is.
b. 719 mm Hg
c. Both have same kinetic energy (same temperature); Xe atom is heavier, Ne atom is faster; same
number of particles in both containers (Avogadro’s hypothesis).
d. First bulb has gas P = 795 mm Hg; second bulb, 760 mm Hg; third bulb, 720 mm Hg
e. All would reach you eventually (because the molecules are in constant random motion), but NH3
would reach you first because it has the smallest molar mass and therefore the fastest molecules,
given that they are all at the same temperature (same kinetic energy).
2. a. 564 mL
e. 2.1 L
b. 35.9 L
f. 2.2 atm
c. 5920 mm Hg
g. 0.29 atm
d. 30. L
h. 144 mol
3. a. STP is Standard Temperature and Pressure: 0 °C or 273 K and 1 atm or 760 mm Hg
b. 1.0 L
c. 0.00113 mol
4. a. 84 mL
b. 79 g
c. 76.2 L
REVIEW PRACTICE TEST
1. Fill in this table:
Symbol Protons Neutrons Electrons Atomic number Mass number
131
I
47
61
9
19
2. On a blank periodic table, fill in the period numbers, group numbers, and staircase; label the alkali
metal, alkaline earth, halogen, and noble gas families; indicate the nonmetal, metalloid, metal, and
transition metal regions; add the trends in atomic size and ionization energy.
3. Use words or phrases from the table at right to fill in the missing parts of this
description of chemical bonding. Some may be used more than once.
When forming a chemical bond, atoms will _____________ valence electrons in
order to achieve the same number of electrons as _______________ . That number
of electrons is ___, except for small atoms like H that try to achieve just ___
electrons in their valence level. This bonding theory is called the _________ rule.
This rule predicts that in general metals tend to _____ valence electrons and
eight
gain
lose
lose, gain, or share
octet
share
the nearest noble gas
two
nonmetals tend to _____ or _____ valence electrons when forming chemical bonds. Thus, an atom of Mg will
_____ __ electrons when forming a bond, and an atom of O will _____ __ electrons when forming a bond.
4. Show the number of valence electrons in each atom by drawing its dot structure:
a. Ca
b. Si
c. O
5. Compare & contrast covalent and ionic bonds. How would you recognize which type of bonding is
present from the formula of the substance?
6. Draw Lewis dot structures for each of these molecules. Include the number of valence electrons.
a. CCl4
b. COS
c. H2S
d. NF3
e. CH2O
7. How many particles are in a mole? What is this number called?
8. Balance these equations
a. Na (s) + H2O (l) → NaOH (aq) + H2 (g)
b. C5H12 (l) + O2 (g) → CO2 (g) + H2O (g)
c. Ag (s) + H2S (g) + O2 (g → Ag2S (s) + H2O (l)
d. AgNO3 (aq) + Cu (s) → Cu(NO3)2 (aq) + Ag (s)
e. C2H5OH (l) + O2 (g) → CO2 (g) + H2O (g)
9. Define the terms “exothermic” and “endothermic.” How can you tell whether a reaction is exothermic
or endothermic?
REVIEW ANSWERS
1.
Symbol Protons Neutrons Electrons Atomic number Mass number
131
I
53
78
53
53
131
108
Ag
47
61
47
47
108
19
F
9
10
9
9
19
2. See pg 162-63, 171, & 174 in your text
3. When forming a chemical bond, atoms will lose, gain, or share valence electrons in order to
achieve the same number of electrons as the nearest noble gas . That number of electrons is 8,
except for small atoms like H that try to achieve just 2 electrons in their valence level.
This bonding theory is called the octet rule. This rule predicts that in general metals tend to
lose valence electrons and nonmetals tend to gain or share valence electrons when forming
chemical bonds. Thus, an atom of Mg will lose 2 electrons when forming a bond, and an atom of O will gain
2 electrons when forming a bond.
4.
Si
Ca
O
5. Ionic compounds are composed of a repeating pattern of positive and negative ions, held together by
the attraction of their opposite charges. The ions form when metal atoms donate their valence
electrons to non-metal atoms, forming positive metal ions (cations) and negative nonmetal ions
(anions). Ionic compounds (salts) are made of metals combined with nonmetals.
Molecular compounds are composed of molecules; the atoms in the molecules are held together by
covalent bonds, in which the atoms share valence electrons. Each atom stays with its own electrons.
Molecular compounds are made of nonmetals combined with other nonmetals (remember H is a
nonmetal).
6. a. CCl4 (32 e–)
b. COS (16 e–)
c. H2S (8 e–)
d. NF3 (26 e–)
e. CH2O (12 e–)
Cl
Cl
C
Cl
Cl
H
S
C
O
S
F
H
N
F
F
O
H
C
H
7. 6.02 x 1023 (Avogadro’s number)
8. a.
b.
c.
d.
e.
2 Na (s) + 2 H2O (l) → 2 NaOH (aq) + H2 (g)
C5H12 (l) + 8 O2 (g) → 5 CO2 (g) + 6 H2O (g)
4 Ag (s) + 2 H2S (g) + O2 (g → 2 Ag2S (s) + 2 H2O (l)
2 AgNO3 (aq) + Cu (s) → Cu(NO3)2 (aq) + 2 Ag (s)
C2H5OH (l) + 3 O2 (g) → 2 CO2 (g) + 3 H2O (g)
9. Exothermic reactions release energy; the energy is shown on the right side of the chemical equation.
Endothermic reactions absorb energy; the energy is written on the left side of the chemical equation.