Chemistry
Final Exam Problems Review
1) For the following
reaction, determine how many moles of oxygen gas (O2), are needed to
react with each of the given moles of glucose (C6H12O6
) listed below:
C6H12O6
(s) +
6 O2(g) à 6 CO2(g) + 6
H20(l)
2) A mass of 46.0
grams of solid sodium reacts with 38.0 g of fluorine gas to produce 84.0 g of
sodium fluoride [NaF(s)].
a. Write a balanced equation for
this reaction.
b. Calculate the number of moles
of each substance.
3) If 17.5 g of zinc
metal are reacted with phosphoric acid [H3PO4(aq)], then
zinc (II) phosphate [Zn3(PO4)2 ]
and hydrogen gas are produced.
a.
How many moles of phosphoric acid are required?
b.
If the phosphoric acid is a 3.00 molar solution, how
many liters are needed?
c.
What mass of zinc phosphate will be produced?
d.
How many moles of hydrogen gas will be given off?
4) What volume of
0.60 M copper(II) sulfate [CuSO4(aq)] will react with 45 mL of 1.5 M
sodium hydroxide [NaOH(aq)] to form copper(II) hydroxide [Cu(OH)2(s)]
and sodium sulfate [Na2SO4(aq)]?
5) Explain why the
percent yield of a product in a reaction is usually less than 100.
6) If 5.45 g of
potassium chlorate [KClO3(s)] are decomposed to form potassium choride
[KCl(s)], then 1.95 g of oxygen gas (O2) are also given off.
a.
Calculate the theoretical yield of oxygen.
b.
Calculate the percent yield of oxygen.
c.
Explain why the percent yield of oxygen is less than
100.
7) a.
How many grams of FeI2(s) can be formed when 25.7 g of Fe(s)
reacts with 105 g of I2(s)?
b. Which
reactant is the limiter?
d.
Which reactant is the excess? How much excess (in grams)?
8) Using a molecular
model, explain why a gas can be easily compressed, while a liquid and a solid
cannot.
9) The volume of a
gas is often referred to as one of the properties that can be measured. Is this volume simply the sum of all of the
individual molecular volumes? Explain
your answer.
10) Explain why the air pressure is greater on
11) Carbon dioxide
does not exist in the liquid state unless the pressure is at least 5.1
atm. Convert this pressure into units
of: a. torr
b. kPa
12) If air pressure
is reduced from normal sea level values (this happens at higher elevations),
the boiling point of a liquid falls. For instance, water boils at only 95
degrees Celsius if the atmospheric pressure is 634 mm Hg. Explain this in terms of intermolecular
forces and vapor pressure. Also, convert
this pressure into units of:
a.
atm
b.
13) Hydrogen gas is
collect by bubbling it through water.
Calculate the partial pressure of the hydrogen gas if:
a.
the total pressure is 94,000 Pa, and the partial
pressure of water is 1200
b.
the total pressure is 100.3 kPa, and the partial
pressure of water is 2600.
c.
Whose gas law helps to solve this problem?
14) In a flask that has a volume of 273 dm3, you have a
sample of two noble gases: neon and xenon.
The partial pressure of the neon is 96,950 Pa, and the partial pressure
of the xenon is 1.025 atm. What is the
total pressure (in kPa) exerted by these
two gases?
15) Explain why you
must convert the temperature to an absolute temperature scale (such as a Kelvin
scale) when you wish to use the direct proportionality of the
temperature-volume law (Charle’s Law).
16) Change the
following volumes of gases from the conditions given to the new
conditions. Assume that the pressure and
amount of gas is constant.
a.
85 cm3 at 61 degrees C to 35 degrees C
b.
7.3 dm3 form 228 degrees C to –48 degrees
Celsius
17) An
anaesthesiologist is about to administer gas to a patient. The gas has a temperature of 22.4 degrees
C. When the gas enters the patient’s
body, it is warmed to a temperature of 37.2 degrees C. Assuming that the gas does not undergo a
change in pressure, what percentage increase in volume does the gas experience
as it reaches the new temperature?
18) A sample of
carbon dioxide has a volume of 2.0 dm3 at a temperature of –10 degrees C. What volume will this sample have when the
temperature is increased to 110 degrees C?
Assume that the pressure does not change nd tht no carbon dioxide leaks from
the sample.
19) Change the following volumes of gases from the conditions
given to the new conditions. Assume that
the temperature and amount of gas is constant.
a.
1.15 E3 cm3 at 75.2 kPa to 14.0 kPa
b.
94.7 dm3 at 1.00 kPa to 100.0 kPa
20) A 12.7-L sample
of gas is under a pressure of 9.3 kPa.
What will be the pressure of the gas when the volume increases to 20.1 L
(assume the temperature is held constant).
21) A sample of nitrogen gas having a volume of 2.44 dm3 is
collected at a pressure of 67.4 kPa.
What volume will the gas occupy when its pressure is changed to 145.1
kPa if the temperature is held constant?
22) What did Avogadro
assume to be true about equal volumes of gases if they are held at the same
temperature and pressure?
23) Which of the
following sample of gases occupies the largest volume, assuming that each
sample is at the same temperature and pressure—50.0 g of neo, 50.0 g of argon,
or 50.0 g of xenon?
24) What volume of carbon dioxide gas contains the same
number of oxygen atoms as 250.0 cm3 of carbon monoxide gas, if each gas sample
is measured at the same temperature and pressure?
25) Mathematically,
the ideal gas law may be expressed as PV = nRT.
What do each of the letters in this equation represent? Identify the variables and constants.
26) What is the volume in dm3 of 12.0 g of nitrogen gas if the
gas is measured at a pressure of 125 kPa and a temperature of 45 degrees C?
27) What mass of carbon dioxide will occupy a volume of 5.5
L at a temperature of 5 degrees C and a pressure of 75 kPa?
28) What is the temperature of a 0.00893 mol sample of neon
gas that has a volume of 302 mL and a pressure of 715 torr?
29) A balloon is filled with helium to a volume of 12.5
L. If the temperature of the gas is 25
degrees C and the pressure is 101 kPa, How many moles of helium are in the
balloon? How many helium atoms are in
it?
30) What is the value
of the standard temperature expressed in degrees C and in Kelvin?
31) What is the value
of standard pressure in atm and kPa?
32) What volume will
25.0 g of ammonia (NH3) occupy at STP?
33) Cooks sometimes use the fermentation of glucose to
produce the gas required to make bread rise, as shown by the following chemical
equation:
C6H12O6
(s) +
2 O2(g) à 2 CH3COOH(l) + CO2(g) + 2 H20(l)
What mass of glucose is required to produce 150. cm3
of carbon dioxide gs measured at STP?
34) When coke (almost
pure carbon) is burned in th presence of air, the product is carbon dioxide in
the following equation:
C(s) + O2(g) à CO2(g)
How many liters of carbon dioxide reproduced from burning
750. g of coke with an excess supply of oxygen?
Assume that the carbon dioxide is measured at STP.
35) What is meant by the term ideal gas? What is a real gas?
36) Gases tend to
follow the ideal gas law more closely when their pressure is low than when it
is high. Considering what happens to gas
volume under high pressure, why do you suppose this trend is true?
37) According to the kinetic molecular theory, what is
different about a sample of xenon gas at 25 degrees C and another sample at 100
degrees C.
38) According to the kinetic molecular theory, what is the
relationship between the average kinetic energy of a sample of gas molecules
and their temperature? Does it matter
what scale you use to express this temperature?
39) Suppose you have two balloons filled with air. The smaller balloon has a volume of 500. mL,
while the larger balloon has a volume of 1500. mL. Both balloons have the same temperature and
pressure. In which balloon (if either)
are the nitrogen molecules moving faster, the larger balloon or the
smaller? How do you know?
40) In an experiment, it takes an unknown gas 1.5 times
longer to diffuse than the same amount of oxygen gas (O2). Find the molar mass of the unknown gas.
41) Suppose a gas
diffuses 1.41 times as fast as sulfur dioxide at the same temperature and
pressure. What is the molar mass of the
unknown gas?
42) Two beams of
ions—germanium-74 and selenium-74—pass between a pair of charged plates. The ions have the same charge, but the
velocity of the germanium is twice as large as the velocity of the selenium
ions. Identify which beam will be
deflected more and explain your reasoning?
43) Compare and contrast the subatomic particles contained
in the following atoms or ions
(identify number of protons, neutrons, and electrons of
each).
a.
a barium-137 atom and a barium-137 ion with a 2+
charge
b.
a molybdenum-96 atom and a molybdenum-96 ion with a 6+
charge
c.
a phosphorus-31 atom and a phosphorus-31 ion with a 3-
charge
44) List which of the three types of radiation—alpha, beta,
or gamma—each of the following describes.
a.
is not deflected by a magnet
b.
has a negative charge
c.
moves with the greatest speed
d.
consists of ions
e.
is similar to light rays
f.
consists of the same particles as cathode rays
45) Describe the problem with Rutherford’s model of the
atom.
46) A proton beam, an electron beam, and a neutron beam pass
between two charged plates. Assuming the
velocities to be the same, draw a diagram showing the paths of each type of
subatomic particle.
47) How are
chlorine-37 and calcium-40 similar?
48) List the number
of protons, neutrons, and electrons in each of the following ions.
a. 7Li1+
b.
24Mg2+
c.
27Al3+
d.
31P1-
e.
127I1-
49) A very bright
line in the bright0ine spectrum of sodium has a wavelength of 5.90 E –7 m. What is the frequency of this line?
50) p. 357 11 A radio
station broadcasts at a frequency of 105.4 MHz.
What is the wavelength of this electromagnetic wave?
51) What must be done
to a hydrogen atom to change its 2s electron to a 3s electron? What happens when a hydrogen atom with a 3s
electron becomes a hydrogen atom with a 2s electron?
52) What orbital will the sixth electron of ground-state
nitrogen occupy?
53) What are the three quantum numbers that describe an
orbital?
54) Which of the following statements about orbitals is
false?
a.
Orbitals are distributed in space round the nucleus.
b.
Orbitals are regions of space in which electrons are
likely to be found.
c.
Orbitals show the path of the electron.
d.
Orbitals are part of one model for atomic structure.
55) Write the
electron configurations for the following elements: arsenic, krypton, bromine,
and phosphorus.
56) Write the electron configurations for the following
ions: K+, O2-, Br-
57) Name the elements
that correspond to each of the following electron configurations. (Assume all
are neutral atoms in their ground states.)
a.
1s22s22p1
b.
1s22s2
c.
1s22s22p63s23p2
58) Write the symbol for the elements found in
the following locations.
a.
period 3, Group 3A
b.
period 1, Group 8A
c.
period 4, Group 2B
d.
period 6, Group 5A
59) At room
temperature, nitrogen is a nonmetallic gas, and bismuth is a sloid metal. Why are they both in Group 5A?
60) Name the group of
elements in the periodic table that has the following outer electron
configurations.
a.
s2
b.
s2p5
c.
electrons filling the d orbital
d.
s2p6
61) Use the periodic table to determine the electron
configuration of the vlaence electrons for the following representative
elements. (The principal quantum number
is the same as the period number, and the total number of s and p
electrons is the same as the Group number.)
62) Predict the common ions formed when atoms of the
elements listed gain or lose electrons.
Then name the noble gas with which the ion is isoelectronic.
63) For which of the following elements do yo expect there
to be a very large increase, going from the 2nd to the 3rd
ionization energy (IE): Na, Mg, Al?
64) What trend is
observed for the atomic radius of atoms
a.
going down a group?
b.
Going left to right across a period?
65) p. 390 34 How does core electron shielding affect the
attraction of the nucleus for the valence electrons?
66) What is ionization energy, and how is it determined?
67) Sodium reacts
with water to form a basic solution, as the following equation shows.
2
Na(s) +
2 H20(l) à 2 NaOH(aq)
+ H2(g)
Write the equation for the reactions of the following metals
with water:
a.
K
b.
Li
c.
Ca
68) Use the
electron-sea model to explain the high electrical conductivity of metals.
69) Explain why fire involving an alkali metal cannot be put
out with water.
70) What is the most
reactive element?
71) Briefly explain
how the potential energy changes as two hydrogen atoms approach each other.
72) How does an ionic
bond form?
73) How does a
covalent bond form?
74) What is a crystal lattice?
75) What types of elements usually form an ionic bond?
76) What types of elements usually form a covalent bond?
77) Classify bonds in
each of the following substances as nonpolar covalent, polar covalent, or
ionic.
a.
K2O
b.
BeO
c.
NH3
d.
KCl
e.
CBr4
f.
N2
g.
CS2
78) Show the partial
distribution in the following bonds:
a.
H—Br
b.
N—O
c.
Be—H
79) Describe the arrangement of the electron pairs for each
of the following molecules and predict
their shape.
a.
CCl4
b.
SiH4
c.
AsH3
d.
H2Se
e.
CS2
80) Show the bond dipoles and predict the
molecular shape of each of the molecules in question 79.
81) On a molecular level, describe
dipole-dipole intermolecular forces.
82) Explain the origin of dispersion forces
between two molecules of nitrogen, N2.
83) Describe the relationship between the
polarity of individual molecules and the nature and strength of intermolecular
forces.
84) How do dipole-dipole forces
differ from H bonds?
85) Place the following in order
of strength from least strength to greatest strength: all chemical bonds, dipole-dipole, H bond,
London dispersion forces
86) Of the following types of
solids, order these from least structured to most structured: ionic, covalent molecular, covalent network,
amorphous, metallic.
87) What effect do intermolecular
forces have on the boiling point of a liquid?
88) What is the relationship
between boiling point, external pressure, and vapor pressure?
89) For a particular substance, why is the molar
heat of fusion less than the molar heat of vaporization?
90) Why is the temperature of a substance
constant during melting?
91) What is the relationship between kinetic
energy and temperature?
92) Why does increasing the
temperature increase the vapor pressure of a liquid?
93) Explain why 12 g of steam at
100 degrees C can melt more ice than 12 g of liquid water at 100 degrees C.
94) What is the triple point of a
substance?
95) What is an indicator?
96) What does it mean for an acid
or base to be neutralized?
97) On a heating curve, why do
some areas increase in temperature and other areas plateau?
98) What does pH measure?
99) What are the products of a
neutralization?
100) What types of ionizing
radiation exist? Know the
characteristics of each.
101) Complete the following
nuclear equation: 204
Pb à 200
Hg +
____ What type of decay
is this?
102) Complete the following
nuclear equation: 234Th à 234
Pa +
____ What type of decay
is this?
103) Complete the following
nuclear equation: 8 B à 8
Be +
____ What type of decay is
this?
104) What is the difference between fission and
fusion?
105) Describe C-14 dating.
106) Describe the following
nuclear equation:
106
Ag +
0 e à
106 Pd