These are questions from a previous year's exam. They may not all apply to what is expected this year 2000-2001
Part I: Multiple Choice: Pick the best answer to the
following questions and fill in the appropriate space on the “scantron”
sheet. Make sure your name is on the scantron.
1.
Which of the following gas particles would have the
greatest velocity if each had the same kinetic energy?
a) bromine b) chlorine
c) ammonia d) hydrogen
e) argon
2.
How many significant figures does 0.00045 have?
a.
one; b. two; c. three; d.
five; e. six.
3.
Aluminum is an example of a class of elements known as:
a. metals; b. non-metals; c.
metalloids ; d. noble gases.
4.
How
many electrons can a d sublevel hold?
a. 10; b. 2; c. 5;
d. 7
5.
The
state in which the electron in an atom has the lowest possible energy is the
a. ground state; b.
excited state; c. inert state; d.
radiation-emitting state.
6.
The
Bohr model of the atom: a. was
referred to as the “plum pudding atom.”
b. described a nuclear atom in which electrons surround a dense nucleus.
c. described the path of electrons as orbits around the nucleus.
d. was the quantum mechanical model.
7.
What
is the charge of iron in FeSO4?
a. +1; b. +2; c.
+3; d. none of the above.
FeSO4 is a covalent compound.
8.
The
outermost electrons of elements in groups IIIA - 0 are filling what kind of
sublevel?
a. s; b. p ; c. d;
d.
f.
9.
How
many electrons are needed to fill the 4th energy level?
a. 18; b. 2 ; c.
32; d. 40.
10.
The energy required to remove an electron from an atom in the gas phase
is the
a. ionic energy; b.
ionization energy; c.
electronegativity; d.
electron affinity.
11.
What is the correct electron configuration of phosphorus?
a. 1s2 2s2 2p6 3s2 3p6 4s2
3d10 4p3
b. [Ar] 4s2 3d10 4p3
c. both a and b are correct
d. neither a nor b is correct
12.
What is the correct orbital diagram for nitrogen?
a.
b.
c.
d.
13.
Which
statement about the periodic table is FALSE.
a. Elements of the same chemical families have similar properties.
b. The periodic table is arranged according to increasing atomic number.
c. Elements of the same period have similar chemical and physical
properties.
d. The periodic table is arranged according to increasing number of
protons.
14.
Which one of
the following sets of elements contains an element that is NOT in the same
family as the others?
a.
tin, lead, carbon
b.
oxygen, carbon, nitrogen
c.
argon, krypton, helium
d.
fluorine, chlorine, iodine
15.
As one
proceeds from sodium to chlorine in the Periodic Table
a.
the atomic radius increases while the electronegativity decreases
b.
the atomic radius decreases while the electronegativity increases
c.
the atomic radius and the electronegativity both decrease.
d.
the atomic radius and the electronegativity both increase.
16.
The periodic
law states that the physical and chemical properties of the elements are
periodic functions of their:
a. wavelengths; b.
combining weights; c. atomic
weights; d. atomic numbers.
17.
The mass of
6.02 x 1023 molecules of chlorine gas is:
a. 1.00 g; b. 35.5 g; c.
71.0 g; d. 6.022 x 1023
g.
18.
What is the
percent composition of nitrogen in the compound (NH4)2S?
a. 21%; b. 36%; c.
41%; d. 0.36%;
e. 0.41%.
19.
What is the
proper name of CuSO4?
a. Copper (I) sulfate; b.
Copper (II) sulfate; c.
Copper sulfur tetroxide; d.
Copper (II) sulfide.
20.
What is the
correct formula of iron (III) phosphate?
a. FePO3
b. FePO4
c. Fe3(PO3)3
d. Fe2(PO4)3
21.
The polyatomic ion, NH4+, is called:
a. nitrate; b.
ammonia; c. ammonium;
d. nitride.
22.
What is the
molecular formula of a compound which has an empirical formula of CH2O
and has a molecular mass of 180.00 g.
a. CH2O; b.
C2H4O2;
c. C5H10O5;
d. C6H12O6
23.
.....Al2(C2O4)3(s)
---->......Al2O3(s) + ......CO(g) + ......CO2(g)
According to the equation for the reaction represented above, what is the
mole ratio of CO to
CO2 that is produced by the decomposition of aluminum oxalate
(Al2(C2O4)3)?
a. 1 mole CO: 1 mole CO2
b.
1 mole CO: 2 moles CO2
c. 1 mole CO: 3 moles CO2
d. 2 moles CO: 1 mole CO2
24.
......LiAlH4(s) + ......BCl3(l) ----->......B2H6
+ ......LiCl(s) + ......AlCl3(s)
When the
equation for the reaction represented above is balanced and all coefficients are
reduced to the lowest whole number terms, the coefficient of B2H6
is: a.1; b.2; c.3;
d.4; e.6.
25.
2Al(s) + Fe2O3 --> Al2O3 +
2Fe(s)
According to the equation for the reaction represented above which of the
following statements is true.
a. If 2 moles of Al is used, 1 mole of Fe is produced.
b. If 1 mole of Al is used, 0.5 mole of Fe2O3 is
consumed.
c. If one mole of Al is used, 1 mole of Al2O3 is
produced.
d. If 0.5 mole of Al is used, 1mole of Al2O3 is
produced.
e. If 0.5 mole of Al is used, 0.5 mole of Fe2O3 is
consumed.
26.
Which of the following combinations of particles represents an ion net
charge of -1 and of a mass number
80
a.
44 neutrons, 35 protons, 36 electrons
b.
44 neutrons, 36 protons, 35 electrons
c.
44 neutrons, 36 protons, 36 electrons
d.
45 neutrons, 35 protons, 35 electrons
e.
45 neutrons, 35 protons, 36 electrons
27.
C2H4(g) + 3O2(g) ----> 2CO2(g) +
2H2O(l)
When 100.0
grams of O2 are allowed to react completely with 1.00 mole of C2H4
according to the equation above, which of the following results?
a. Some C2H4 remains unreacted.
b. Some O2 remains unreacted.
c. Only CO2 and H2O are present when the reaction
has run to completion.
d. Less than 2 moles of CO2 is formed.
e. The partial pressure of O2 falls to zero.
28.
.....Ag+(aq) + ......S2-(aq) ----> ........(s)
When the
equation for the reaction represented above is completed and balanced by the use
of lowest whole number coefficients, the coefficient for Ag+ is: a.1;
b.2; c.3;
d.4; e.6.
29.
According to quantum mechanics, correct statements concerning the
electron in the hydrogen atom include which of the following?
I.It moves in a definite circular orbit around the nucleus.
II. It is associated with definite energy levels.
III. It occupies a fixed position in space with reference to the nucleus.
a.
I only
b.
II only
c.
I and III only
d.
II and III only
e.
I, II, and III
30.
When the external pressure is 505 kPa, what is the
vapor pressure of water at its boiling point?
a) 0 kPa b) 101 kPa
c) 505 kPa d) 1010 kPa
e) not enough information
31. How many kilocalories of heat are
required to raise the temperature of 225 g of aluminum from 20oC to
100oC? (specific heat of aluminum = 0.21 cal/(g xoC))
a)
0.59 kcal
b) 3.8 kcal c) 85 kcal
d) None of the above
32.
For a
particular compound, which of the following pairs can represent the empirical
and the molecular formula, respectively?
a.
CH3 and C3H6;
b. CH2 and C2H2;
c. CH2 and C3H9;
d. CH and C6H6;
e. CH and CH4
33.
H2O(l) --> H2O(g)
Correct
statements concerning the process above occurring at 100o C include
which of the following?
I.
The vaporization process is endothermic.
II.
The system is absorbing heat.
III. The average potential energy of the vapor
molecules is greater than that of the liquid molecules.
a.
I only; b. II only;
c. I and III only; d.
II and III only; e. I, II,
and III.
34.
Which of the following statements about the halogens (X) is true?
a.
They all form X- ions.
b.
They have the lowest ionization (potential) energies of the elements in their
respective periods.
c.
The are all solids.
d.
They have the largest atomic radii of their period.
e.
They occur as single atoms in their elemental form.
Diagram for # 35
35.
Both of the
gas samples represented above are at the same temperature and pressure. The mass
of H2 in the 1 liter container is 0.20 gram. The mass of X in the two
liter container is 8.0 grams. The molecular weight of X is:
a.10; b.20;
c.40; d.60;
e.80.
36.
At 23oC,
200 milliliters of an ideal gas exerts a pressure of 750 mm of Hg. The volume of
the gas at 0oC and 760 mm of Hg found from which of the following
expressions?
a.
200 x (760/750) x (273/296) ml
b.
200 x (750/760) x (0/23) ml
c.
200 x (760/750) x (23/0) ml
d.
200 x (760/750) x (296/273) ml
e.
200 x (750/760) x (273/296) ml
Table for #37:
37.
The
ionization energies (potentials) for the removal of different electrons from an
atom of an element in the gas phase are shown above. An atom of this element is
most likely to form an ion that has the charge of:
a. +1; b. +2;
c. +3; d. +4;
e. +5.
38.
Of the
following ground state electron configurations, the one that represents the
element of lowest first ionization energy (potential) is:
a.1s22s22p5
b.1s22s22p6
c.1s22s22p63s1
d.1s22s22p63s2
e.1s22s22p63s23p1
39.
A 250 gram sample of a hydrated salt was heated at 110o C
until all water was driven off. The
remaining solid weighed 160 grams. From these data, the percent of water
by weight in the
original sample can be correctly calculated as:
a.
(160/340) x 100
b.
(90/340) x 340
c.
(160/250) x 100
d.
(90/250) x 100
e.
(90/160) x 100
40.
Cl2(g) + 2KBr (aq) ------>
When 1 mole
of chlorine gas reacts completely with the excess KBr solution, as shown above,
the products obtained are:
a.
1 mole of Cl- ions and 1 mole of Br
b.
1 mole of Cl- ions and 2 moles of Br
c.
1 mole of Cl- ions and 1 mole Br2
d.
2 moles of Cl- ions and 1 mole of Br2
e.
2 moles of Cl- ions and 2 moles of Br2
41.
Some solid crystalline compounds slowly change to a gaseous state when
left at room temperature in an open container. Which of the following is true
about this phenomenon?
a.
It requires the release of heat to the surroundings.
b.
It is accompanied by an absorption of heat by the solid.
c.
It is the result of a chemical reaction with air.
d.
It is best described as fusion.
e.
It is observed only with ice.
42.
A compound composed of three different types of elements is:
a. table salt; b. water; c.
potassium perchlorate; d.
carbon dioxide.
43.
Which of the following formulas does NOT represent a molecule?
a. NH3; b. H2O;
c. NaCl; d. CO2.
44.
The formula for dinitrogen tetrafluoride would be:
a. NF; b. N2F;
c. N2F3;
d. N2F4.
45. What happens to the average
kinetic energy of the particles in a sample of matter as the temperature of the
sample increased? a) It decreases.
b) It increases. c) It does not
change.
46.
When a gas is heated:
a) all of the absorbed
thermal energy is converted to kinetic energy.
b) some
of the absorbed thermal energy is converted to the internal energy of the gas
particles, and some is converted to kinetic energy.
c) all of the absorbed
thermal energy is converted to potential energy.
d)
one-half
the absorbed thermal energy is converted to potential energy and the other half
is converted kinetic energy.
47.
What is the volume occupied by 71.0 g of chlorine gas at STP?
a) 22.4 L
b) 44.8 L
c) 56.0 L
d) 67.2 L
e) 78.4 L
48.
What is the number of moles of gas in 56.0 L of oxygen at STP?
a) 0.50 mol
b) 1.00 mol
c) 1.5 mol
d) 2.00 mol
e) 2.50 mol
49. Compared
with 1 mole of chlorine gas at STP, 1 mole of hydrogen gas at STP occupies:
a) more volume. b) less
volume. c) the same volume.
50. The first particles to evaporate
from a liquid are: a) those with
the lowest kinetic energy. b) those with the highest kinetic energy.
c) those farthest from the surface of the liquid.
51. What happens to the temperature
of a liquid as it evaporates? a) It
increases. b) It decreases. c) It does not change.
52. If a liquid is sealed in a
container and kept at constant temperature, how does its vapor pressure change
over time? a) It rises
continuously. b) It rises at first,
then remains constant. c) It rises
at first, then falls.
53.
Why does a liquid’s evaporation rate increase when the liquid is
heated?
a) because
more surface molecules have enough energy to overcome the attractive forces
holding them in the liquid.
b) because the average
kinetic energy of the liquid decreases.
c) because the surface
area of the area is reduced.
d) because the
potential energy of the liquid increases.
54. In an exothermic reaction, the
energy stored in the chemical bonds of the reactants is:
a) equal to the energy stored in the bonds of the products. b) greater than the energy stored in the bonds of the
products. c) less than the energy
stored in the bonds of the product. d)
less than the heat released. e)
less than the heat absorbed.
55. If the heat involved in a
chemical reaction has a negative sign, a)
heat is lost to the surroundings. b)
heat is gained from the surroundings. c)
no heat is exchanged in the process.
56. Calculate
the energy required to produce 6.0 mol Cl2O7 on the basis
of the following balanced equation: 2Cl2(g)
+ 7O2(g)
+ 130 kcal
--> 2Cl2O7(g)
a) 12 kcal
b) 21 kcal
c) 43 kcal
d) 130 kcal
e) 390 kcal
57.
If you were to touch the flask in which an endothermic reaction were
occurring
a) the flask would
probably feel cooler than before the reaction started.
b) the flask would
probably feel warmer than before the reaction started.
c) the flask would
feel the same as before the reaction started.
d) None of the above
58.
Calculate the energy released when 24.8 g Na2O reacts in the
following reaction.
Na2O(s) +
2HI(g) -->
2NaI(s) +
H2O(l)
H = -120.00 kcal
a) 0.207 kcal
b) 2.42 kcal
c) 48.0 kcal
d) 3.00 x 102
kcal
e) 2980 kcal
59.
What is the standard heat of reaction for this reaction:
Zn(s) +
Cu2+ (aq) --> Zn2+ (aq) +
Cu(s)
(DHof for Cu2+
= +64.4
kJ/mol; DHof for Zn2+
= -152.4 kJ/mol)
a) 216.8 kJ created
per mole
b) 88.0 kJ created per
mole
c) 88.0 kJ absorbed
per mole d)
216.8 kJ absorbed per mole
60. DHof for the formation of rust (Fe2O3)
is -826 kJ/mol. How much energy is
involved in the formation of 5 grams of rust?
a) -25.9 kJ b) -25.9 J
c) -66 kJ d) 66 J
61.
Why does the pressure inside a container of gas increase if more gas is
added to the container?
a) Because
there is a corresponding increase in the number of particles striking an area of
the wall of the container per unit time.
b) Because there is a
corresponding increase in temperature.
c) Because there is a
corresponding decrease in volume.
d) Because
there is a corresponding increase in the force of the collisions between the
particles and the walls of the container.
62. Increasing
the volume of a given amount of gas at constant temperature causes the pressure
to decrease because: a) the
molecules are striking a larger area with the same force.
b) there are fewer molecules. c)
the molecules are moving more slowly. d)
there are more molecules.
63.
Which of these changes would NOT cause an increase in the pressure of a
gaseous system?
a) The
container is made larger.
b) Additional amounts
of the same gas are added to the container.
c) The temperature is
increased.
d) Another gas is
added to the container.
64. A gas occupies a volume of 0.20 L
at 10.1 kPa. What volume will the
gas occupy at 101 kPa?
a) 38 L b) 20. L
c) 2.0 L d) 0.020 L
65. A sample of gas occupies 40.0 mL
at –123oC. What volume
does the sample occupy at 27oC?
a) 182 mL b) 8.80 mL c)
80.0 mL d) 20.0 mL
66. At a certain temperature and
pressure, 0.20 mol of CO2 has a volume of 3.1 L.
A 3.1-L sample of hydrogen at the same temperature and pressure: a) has the same mass. b)
contains the same number of atoms. c)
has a higher density. d) contains
the same number of molecules.
67. If a balloon containing 1000. L
of gas at 50oC and 101 kPa rises to an altitude where the pressure is
50.5 kPa and the temperature is 10 C, the volume of the balloon under these new
conditions would be
a) 1000 L
x 101 kPa x 10oC
50.5
kPa 50oC
b) 1000 L
x 323 K x 50.5
kPa
283 K
101 kPa
c) 1000 L
x 101 kPa x 283
K
50.5 kPa 323 K
d) 1000 L
x 50oC
x 50.5 kPa
10oC
101 kPa
Part
II: Matching:
(1 point each) Match the letter of the correct definition or example with the
numbered vocabulary. some letters will not be used.
There is only one appropriate answer for each number.
Place your answer on the scantron sheet starting with number 68.
68.
isotope
a. his experiments led to
the discovery of the electron.
69.
allotrope
b. his experiments led to
the discovery of the nucleus.
70.
atomic number
c.
the number of atoms found in 12.0 grams of carbon-12
71.
Rutherford
d. is a cute brown furry
thing that lives in the ground
72.
mass number
e.
carbon-14 and carbon-12 are examples
73.
Charles
ab. graphite and diamond are
examples.
74.
atomic mass
ac. H2, N2
and O2 are examples.
75.
Hund’s Rule
ad. developed a gas law
relating pressure and temperature
76.
Boyle
ae. the ability to attract
and acquire electrons; this process releases energy
77.
Avogadro’s number
bc. ability of an atom to
attract electrons when bonded to another atom
78.
Pauli Exclusion Princ.
bd. developed a gas law
relating pressure and volume
79.
Dalton
be. developed a gas law
relating volume and temperature
80.
electronegativity
cd. developed a gas law
relating moles and pressure
81.
transition elements
ce. orbitals of the same
energy level will contain one electron each with the
same
spin until they have to contain two
82.
Gay-Lussac
de. This number is in common
with all isotopes of the same element.
abc. This number indicates
the number of protons and neutrons in an atom.
abd. relative, weighted
average of all known types of atoms of the element
abe. developed a
relationship between rates of effusion and molar masses
acd. copper, silver and gold
are examples of these
ace. Most actinides fall in
this category.
ade. provided the first
principles of the modern atomic theory
bcd. No two electrons can
have the same four quantum numbers in the
same
atom.
Part II: Problems and Short
Answers:
Answer the following problems
in the space given. BE SURE TO SHOW YOUR
WORK FOR FULL CREDIT. For
each question answer all the
parts of the question in complete sentences.
1.
Determine the empirical formula of a certain copper sulfide ore if a sample of
the compound contains 79.8% by mass of copper (Cu).
2.
Write balanced equations for each of the following:
a. the formation of
crystalline mercury (I) oxide from the elements.
(Mercury is Hg)
b. the combustion of ethane
gas (C2H6)
c.
Al(s) +
CuCl2(aq) à
d.
Fe2O3(s) à
3.
Calculate the volume of carbon dioxide produced when 250.0 g of pentane,
C5H12, burn. Assume
the carbon dioxide is cooled to STP when measured.
(Hint: You need an equation to do stoichiometry.)
4.
What mass of silver can be produced
from 3.00 mol of copper and 3.85 mol of silver nitrate?
(a)
Write
a balanced equation for the reaction.
(b)
Determine
which reactant is limiting if copper(II) nitrate is one product.
(c)
How
much of the excess reactant is left over?
5.
What is the actual amount of magnesium oxide that can be produced when
carbon dioxide reacts with 42.8 g of magnesium metal?
The percent yield for this reaction is 81.7%.
2 Mg +
CO2 ----> 2
MgO +
C
6.
Calculate the amount of heat produced when 34.8 g of methane, CH4,
burns in an excess of air, according to the following equation:
CH4 +
2 O2 ----> CO2
+ 2 H2O
+ 890.2 kJ
7.
Calculate the change in enthalpy (DHrxn)
for the following (unbalanced)
reaction. Assume all the reactants
and products are gases. [DHof
(NH3) = -46.2 kJ, DHof
(NO) = 90.4 kJ,
DHof
(H2O) = -241.8 kJ]
NH3 +
O2 ---->
NO +
H2O
8.
Explain the relationship between the periodic trends, ionization energy
and atomic size in terms of the fundamental forces of the atom.
9.
One of the cylinders in an automobile engine is heated and the piston
moves, allowing the gas inside to expand. The
original pressure was 187 kPa, while its original volume was 175 mL, measured at
18oC. The final measured
pressure was 87 kPa and the temperature was measured at 382oC. Calculate the final volume of the cylinder.
10.
Calculate the mass of 275 mL of nitrogen dioxide at a pressure of 240.0
kPa and 28oC.
(R = 8.31 (kPa x L)/(mol x K))
11.
Discuss three
significant differences between the three common phases or states of matter.
12.
Discuss the difference between an ideal
and a real gas. Also tell why and
under what conditions the assumptions about an ideal gas are good
approximations.
13.
Discuss Rutherford’s experiment regarding the structure of the atom.
What did he discover about the atom’s structure?
How did he come to this conclusion?
What question remained to be answered by his student, Neils Bohr?
Extra Credit:
Balance the following equation.
NaI +
MnO2 +
H2SO4 à
Na2SO4 +
MnSO4 +
H2O +
I2
