Download Resource for Final Exam Prep

Dr. Jesudoss Kingston
Fall 2015
This exam consists of 70
questions
You have 110 minutes to
complete the exam
Exam booklet serial number:
____________ Yellow or
Name___________________________________________
Chem 177 AE
FINAL Exam
Dec. 16, 2015
Recitation TA Name:
_______________________________________________
Recitation Section number:
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Gray
DO NOT WRITE ANYTHING IN THE EXAM BOOKLET! Use this sheet as well as scratch
papers for doing calculations and other scratch work. Record your answers on the
computer scan sheet (the bubble sheet). Turn in the computer scan sheet and the exam
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Select the one best answer for each question. Place your answer on the computer scan sheet
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to go to the correct room for your recitation section.
Room
Troxel 1001
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TA
Beatrise Berzina
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James Shi
Section(s)
3, 6
11, 15
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53, 55
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12
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5, 7
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61, 63
8, 10
4, 17
56, 59
13
52
General Instructions:
 All electronic devices except non-programmable or non-graphing
calculators, should be turned off and put away.
 Bring your ISU ID.
 DO NOT write anything on the ACS booklet (Do your work on the white
cover sheet with pasge-1 instruction on them).
 Write your name, ACS booklet # on the half sheet and remember to
turn it in at the end of the exam.
 Hat/Cap should be removed while taking the exam.
 Do not look around while taking the exam. You will be suspected of
cheating and severe disciplinary action will be taken and reported to
Dean of students.
As mentioned in the class, only limited equations and constants will be provided
on the ACS exam and you are required to know all other stuff (see below).
Things that are provided:
A Periodic table with element symbol, atomic number & atomic mass
NA = 6.022  1023
h = 6.626  10-34 J.s
c = 2.998  108 m
R = 0.0821 L.atm.K-1.mol-1 = 8.314 J.K-1.mol-1
Things that you should know: (These are just a summary of each chapter.
You are required to read relevant stuff from your textbook, because I might
have missed few here and there)
Chapter-1:
Precision, Accuracy, atoms, elements, compounds, mixture, pure substance, extensive,
intensive, physical-chemical properties,
Density = mass/volume
1 cm3 = 1 mL 1 m3 = 1000L and other simple unit conversions including nm, mm, cm,
m, km, mg, g, kg, etc
Chapter-2:
Average atomic mass = (mass of isoptope-1) (fraction of isotope-1 abundance) + (mass
of isoptope-2) (fraction of isotope-2 abundance) + ………
OR = ∑(fractional isotope abundance)(mass of isotope)
Naming of ionic and molecular compounds, writing chemical formula etc. Know all of the
polyatomic ions. Know the rules for naming.
Chapter-3:
Balancing chemical equation.
Mole = mass/molar mass or n = m/MM
Molarity = moles of solute/Liters of solution or M = n/L
Chapter-4:
Complete ionic and net-ionic equations, weak and strong electrolytes, dilution, ion
concentration in a mixture of solutions, solution stoichiometry etc.
Solubility Rules: Know that ionic compounds containing acetates (C2H3O2-) or Nitrates
(NO3-) or ammonium (NH4+) or group-1 metal cations would be soluble in water.
Concept of redox reaction. Assigning oxidation number. Remember LEO the lion says
GER. Oxidation is Lose of Electron and Gain of Electron is Reduction. So, in an oxidation,
the oxidation number increases and in the Reduction, the oxidation number decreases.
Reducing agent  is the one that get oxidized and Oxidizing agent  is the one get
reduced!
Chapter-5:
E = q + w [where q is heat and w is work and they are NOT state functions]
Sign of q and w and their meanings. Energy profile diagrams for an endo and exothermic
reactions
Recognize the chemical equation for standard enthalpy of formation, Hof
Calculation of H using standard enthalpy of formation, Hof data:
H = ∑ mH𝑓0 (products) − nH𝑓0 (reactants)
Calculation of H using bond enthalpy data:
H = Energy Input – Energy output or
= [sum of energy needed to break all the bonds] – [sum of energy released to form all the
bonds]
q = m.s. T
qrxn = -qsoln
H = q at constant pressure (open system such as coffee-cup calorimeter)
E = q at constant volume (closed system such as bomb calorimeter)
Hess’s Law of manipulating thermochemical equations
Different energy units, fuel density, Food calorie etc
Chapter-6:
E = h
and n = c/
So, E = hc/
1 nm = 1  10-9 m
Photoelectric effect  kinetic energy of ejected electron
Uncertainty principle  you can estimate accurately the position and the momentum of
an electron at the same time
Quantum numbers (n, l, ml, ms), what do they each define? Value of l for s, p, d and f
orbitals
Energies of different level depends on the n+l values [for example, 2p is slightly higher
in energy than 2s]
Pauli-Exclusion principle, Hund’s rule, Electron configuration (condensed and full
version), Exceptions such as Cr, Cu, Mo and Ag (half-filled or fully-filled orbitals are
more stable). Ground state vs excited state electron configuration, degenerate orbitals
Nodes, nodal planes, shapes of s and p orbitals
Chapter-7:
Periodic properties: Trends of Size/atomic radii, Effective nuclear charge, ionization
energy, electron affinity, compare sizes of isoelectronic cations and anions.
Classification of elements (metals, non-metals, metalloids), special group names,
physical states of elements (gas, liquid, solid, monoatomic/diatomic etc).
Know equations for general reactions such as the following:
For active metals: metal + water  metal hydroxide + H2(g)
Metal oxide + water  metal hydroxide (aq) [basic solution]
metal oxide + acid  salt + H2O
Nonmetal oxide + water  acid (aq) [an acidic solution]
Non-metal oxide + base  salt + H2O
Carbonate or bicarbonate + acid  salt + CO2 + H2O
Chapter-8:
Concept of chemical bonding  ionic, covalent, polar covalent, metallic bonding
Electronegativity, its trend, Lattice energy, Valence electrons, Lewis structures, formal
charges, octet rule, resonance structures, bond polarity, molecular polarity, dipole
moment, relationship between bond strength and lengths of single, double and triple
bonds.
Chapter-9:
Theories of Bonding: Valence bond theory uses hybridization, sp, sp2, sp3 to explain
shapes, bond angles etc, electronic/molecular geometry, sigma and pi-bonds, head-on
and side-way overlaps.
Molecular orbital theory uses linear combination of atomic orbitals to form “bonding”
and “Antibonding” orbitals. Know the shapes of these orbitals for the combination of s
and p atomic orbitals.
MO energy level diagrams for simple diatomic molecules such as H2, He2 etc.
Bond order = ½[# of electrons in the bonding orbital - # of electrons in the antibonding
orbitals]
Presence of unpaired electrons  paramagnetic
Absence of unpaired electrons  diamagnetic
Chapter-10:
All of the gas laws, PV = nRT
Boyle’s Law: P1V1 = P2V2
Charles’s Law: V1/T1 = V2/T2
Gay-Lussac’s Law: P1/T1 = P2/T2
Avogadro’s Law: V1/n1 = V2/n2
Combined Gas Law: P1V1/T1 = P2V2/T2
Temperature
should be in
Kelvin
Molar volume: 22.4 L at STP (0oC, 1 atm, 1 mol)
Ideal behavior: at high temperature and low pressure
Dalton’s law of partial pressure:
Total pressure = partial pressure of gas -A + partial pressure of gas-B + etc
Partial pressure of gas-A = (mole fraction of gas-A)(Total pressure) or PA = A.PT
Rate of diffusion/effusion depends on the molar mass  smaller molecular can
diffuse/effuse much faster than a bigger molecule (see Figure 10.14).
𝑟1
𝑀2
= √
𝑟2
𝑀1
Chapter-11:
Intermolecular forces: London-dispersion (due to temporary dipole, this is the only
force for non-polar molecules), dipole-dipole (polar molecules), ion-dipole (ions and
polar molecule), hydrogen bonding (molecule should be H attached to F, O or N and
there should be at least one lone pair of electrons in F, O or N)
Stronger the intermolecular forces  high melting point, high boiling point, high
viscosity, high surface tension, low vapor pressure
During phase changes (such as melting, boiling, vaporization), one need to overcome
these intermolecular forces. Usually, the heat needed to do vaporization (Hvap) is
higher than heat needed to do melting or fusion (Hfus).
Understand heating curve: good review of heat calculation
Understand phase diagram: Existence of physical states at different pressure and
temperature conditions. Triple point => All three phases will be co-existing, critical
temperature => beyond this temperature, gas cannot be liquefied by applying pressure.
Critical pressure => the pressure that is needed to liquefy a gas at its critical
temperature.
Know how vapor pressure is related to boiling point. The boiling point of a liquid refers
to the temperature at which its vapor pressure will be equal to the atmospheric
pressure. If the atmospheric pressure is 1 atm, then it is called normal boiling point.
Here are some questions for you to practice. More practice problem sets are posted
in Mastering Chemistry. Review all of our exams, quizzes, worksheets, textbook
problems etc for additional help. Good luck.
1.
When the following operation is performed, the answer to the correct number of significant figures
(assuming no number shown is exact) is
(1.0 + 12.01) / (12 – 1.2)
1) 1.2
2.
3) 1.205
4) 1.2046
All of the following are properties of sodium. Which one is a physical property of sodium?
1)
2)
3)
4)
3.
Its surface turns black when first exposed to air.
It is a solid at 25 ˚C and changes to a liquid when heated to 98 ˚C.
When placed in water it sizzles and a gas is formed.
When placed in contact with chlorine it forms a compound that melts at 801 ˚C.
The active ingredient in many antacids is CaCO3. Suppose that the average dose of CaCO3 in an antacid
tablet is 500.0 mg, and that a bottle of antacid tablets contains 250 tablets. What mass of Ca is present
in 4 bottles of antacid tablets? CaCO3 is 40.0% Ca by mass.
1) 2.00 g
4.
2) 40.0 g
3) 120. g
4) 200. g
An object will sink in a liquid if the density of the object is greater than that of the liquid. The mass of a
sphere is 9.83 g. If the volume of this sphere is less than _____ cm3, then the sphere will sink in liquid
mercury (density = 13.6 g/cm3).
1) 0.723
5.
2) 1.38
3) 134
4) 7.48 5) none of these
Atomic number refers to
1)
2)
3)
4)
5)
6.
the number of protons in a nucleus.
the coefficient preceding a substance in a balanced equation.
the subscript following an element’s symbol in a formula.
the number of nucleons in an atom.
the number of atoms in an isomer.
What is the molar mass of (NH4)2CO3?
1) 96 g
7.
2) 87 g
3) 78 g
4) 69 g
The hypothetical element studentium (Sd) has two isotopes: Sd-88; natural abundance 66.7% with
atomic mass 88.1 amu; and Sd-91; natural abundance 33.3% with atomic mass 91.1 amu. What is the
average atomic mass of Sd?
1) 90.1 amu
89.6 amu
8.
2) 1.20
2) 89.1 amu
3) 91.1 amu
4) 88.1 amu
5)
Two students count the grains of uncooked rice in a small cup. This measurement was repeated four times
by both students with the following results: Mike – 256, 263, 262, 266 and Ike – 250, 242, 270, 278. The
actual number of grains is 260. Which student is more accurate, which is more precise? [Remember:
Accuracy is how close the avg measurement to the expected value and Precision is how close each
measurement with the avg of the measured value (standard deviation)]
1.
2.
3.
4.
Mike is more accurate; Ike is more precise
Mike is inaccurate; Ike is more precise
Mike is more precise; Ike is not precise
Mike is more precise; Ike is more accurate
9.
What is the name of CuCl2?
1) cupric bichlorine
2) copper(II) chloride
10.
3) copper(I) chloride
4) copper dichloride
The name of PCl3 is
1) potassium chloride
2) phosphorus trichloride
3) phosphorus(III) chloride
11.
4) phosphorus chloride
5) trichloro potassium
How many oxygen atoms are there in 52.06 g of carbon dioxide?
1) 1.424  1024 2) 6.022  1023 3) 1.204  1024 4) 5.088  1023 5) 1.018  1024
12.
What is the formula of chromium(III) sulfate?
1) Cr3SO4
13.
2) Cr2(SO4)3
4) mass
5) all are extensive properties
The formula for magnesium nitrate is
1) MgNO3
15.
4) Cr2(SO3)3
Which of the following is an extensive property?
1) density
2) boiling point
3) color
14.
3) Cr3S2
2) Mg(NO3)2
3) Mg2NO3
4) Mg2NO4
5) MgNO
Balance the equation using the smallest set of whole numbers. What is the coefficient for H2O?
_____PCl3() + _____H2O()  _____H3PO3(aq) + _____HCl(aq)
1) 1
16.
2) 2
3) 1.5  102 L
4) 1.5  104 L
how close a measured number is to other measured numbers
how close a measured number is to the true value
how close a measured number is to the calculated value
how close a measured number is to zero
how close a measured number is to infinity
To the correct number of significant figures, what is the temperature reading on the
Celsius thermometer at the right?
1)
2)
3)
4)
19.
5) none of these
Precision refers to ___________.
1)
2)
3)
4)
5)
18.
4) 5
Convert 15 m3 to liters.
1) 1.5  10–2 L 2) 1.5 L
17.
3) 3
21 ˚C
21.7 ˚C
21.70 ˚C
22 ˚C
An element cannot
1) be part of a heterogeneous mixture.
3) be separated into other substances by chemical means.
4) interact with other elements to form compounds.
2) be part of a homogeneous mixture.
5) be a pure substance.
Questions 20 and 21 refer to the following diagram.
a)
20.
2) drawing b)
3) drawing c)
4) drawing d)
If shaded and unshaded spheres represent atoms of different elements, which of the above drawings
most likely represents a molecular compound at room temperature and a pressure of 1 atm?
1) drawing a)
22.
d)
c)
If shaded and unshaded spheres represent atoms of different elements, which of the drawings below
most likely represents an ionic compound at room temperature and a pressure of 1 atm?
1) drawing a)
21.
b)
2) drawing b)
3) drawing c)
4) drawing d)
Which one of the following is not one of the postulates of Dalton’s atomic theory?
1) Atoms are composed of protons, neutrons, and electrons.
2) All atoms of a given element are identical; the atoms of different elements are different and have
different properties.
3) Atoms of an element are not changed into different types of atoms by chemical reactions; atoms are
neither created nor destroyed in chemical reactions.
4) Compounds are formed when atoms of more than one element combine; a given compound always
has the same relative number and kind of atoms.
5) Each element is composed of extremely small particles called atoms.
23.
When 0.387 g of Cr is heated in an atmosphere of Cl2 gas, a combination reaction occurs and 1.178 g of a
solid compound is formed. Assuming that all of the chromium reacts, what is the mass of chlorine that
reacted?
1) 0.387 g
24.
2) sulfurous acid
5) sulfur hydroxide
4) 0.396 g
5) 1.565 g
3) hydrosulfuric acid
Which is a chemical property of chlorine?
1) It is yellowish green.
3) It has a density of 3.2 g•L–1 at STP.
5) It boils at -34 ˚C.
26.
3) 0.791 g
The correct name for H2SO3 is ________.
1) sulfuric acid
4) hydrosulfic acid
25.
2) 1.178 g
2)
4)
It burns in sodium vapor.
It dissolves in carbon tetrachloride.
Aluminum oxide, Al2O3, is used as a filler for paints and varnishes as well as in the manufacture of
electrical insulators. Calculate the number of moles in 47.51 g of Al2O3.
1) 2.38 mol
2) 2.15 mol
3) 1.10 mol
4) 0.466 mol
5) 0.421 mol
27.
A molecule of water contains hydrogen and oxygen in a 1:8 ratio by mass. This is a statement of _____.
28.
1) the law of multiple proportions
2) the law of constant composition
3) the law of conservation of mass
4) the law of conservation of energy
5) none of the above
What is the formula of the compound formed between strontium ions and nitrogen ions?
1) SrN
2) Sr3N2
3) Sr2N3
4) SrN2
5) SrN3
29.
What is the mass of 1 mole of Na2S?
1) 55.1 g
30.
2) 55.1 amu
2) 1.5
3) 2.8
2) 2
3) 3
c) Mg(OH)2  MgO + H2O
d) 2 Na + Cl2  2 NaCl
4) 4
5) zero
Which of the following is the balanced net ionic equation for the reaction between Al(OH)3(s) and
HNO3(aq)? HNO3 is a strong acid.
1)
2)
3)
4)
Al(OH)3(s) + 3HNO3(aq)  Al(NO3)3(aq) + 3H2O()
Al(OH)3(s) + 3H+(aq)  Al3+(aq) + 3H2O()
Al3+(aq) + 3OH–(aq) + 3H+(aq) + 3NO3–(aq)  Al3+(aq) + 3NO3–(aq) + 3H2O()
H+(aq) + OH–(aq)  H2O()
How many grams of KOH are required to prepare 250.0 mL of 2.0 M KOH solution?
1) 0.5 g
5.
5) 184
How many of the following are oxidation-reduction reactions?
1) 1
4.
4) 0.36
CuCl2 will precipitate, and Ba2+ and SO42– are spectator ions.
CuSO4 will precipitate, and Ba2+ and Cl– are spectator ions.
BaSO4 will precipitate, and Cu2+ and Cl– are spectator ions.
BaCl2 will precipitate, and Cu2+ and SO42– are spectator ions.
No precipitate will form.
a) NaOH + HCl  NaCl + H2O
b) Cu + 2 AgNO3  2 Ag + Cu(NO3)2
3.
5) 6.022  1023
In accordance with the solubility rules, which of the following is true when solutions of CuSO 4(aq) and
BaCl2(aq) are mixed?
1)
2)
3)
4)
5)
2.
4) 78.1 amu
How many grams of hydrogen are in 46 g of CH4O?
1) 5.8
1.
3) 78.1 g
2) 14 g
3) 28 g
4) 56 g
5) 112 g
When 48.0 g of nitrogen gas reacts with an excess amount of hydrogen gas to give an actual yield of 5.90
g of ammonia gas, NH3(g), what is the percentage yield for this reaction?
N2(g) + 3 H2(g)  2 NH3(g)
1) 82.3%
6.
2) 60.7%
3) 58.3%
4) 20.2%
5) 10.1%
The amount of chloride ion in a water supply can be determined by reaction with aqueous silver nitrate.
AgNO3(aq) + Cl–(aq)  AgCl(aq) + NO3–(aq)
What mass of chloride ion is present in a 100.0 g sample of aqueous solution if 40.8 mL of
0.100 M AgNO3 is required to react with all the chloride ion in the sample?
1) 4.08  10–3 g 2) 0.144 g
3) 0.585 g
4) 0.693 g
5) 4.08 g
The following statement and diagrams refer to problems 7 and 8.
Consider a representation of a small volume of a reaction flask before the start of a reaction between nitrogen
gas and oxygen gas to give dinitrogen pentoxide,
.
____N2(g) + ____O2(g) 
N2 =
N2O5(g)
O2 =
10
?
10
initial
7.
after
Which of the following represents this same small volume after the reaction occurs?
4
2
10
4
6
8
1)
8.
4)
2) N2
3) O2
4) both N2 and O2
5) none of these
When 25.0 mL of H2SO4 solution was completely neutralized in a titration with 0.0500 M NaOH solution,
18.3 mL of the base was required. What was the molarity of the H2SO4 solution? The reaction was
H2SO4 + 2 NaOH  Na2SO4 + 2 H2O
1) 0.0366 M
10.
3)
What is the limiting reagent in the above problem?
1) N2O5
9.
2)
2) 0.0100 M
3) 0.0148 M
4) 0.200 M
5) 0.0183 M
Write a balanced net ionic equation for the reaction of NiBr2(aq) with (NH4)2S(aq).
1) NiBr2(aq) + (NH4)2S(aq)  NiS(s) + 2 NH4Br(aq)
2) Ni2+(aq) + 2 Br – (aq) + 2 NH4+(aq) + S2–(aq)  NiS(s) + 2 NH4+(aq) + 2 Br – (aq)
3) Ni2+(aq) + 2 Br – (aq) + 2 NH4+(aq) + S2–(aq)  NiS(s) + 2 NH4Br(s)
4) Ni2+(aq) + S2–(aq)  NiS(s)
11.
What type of reaction will take place when NH3(aq) and H2SO4(aq) are mixed?
1) precipitation
2) oxidation-reduction
12.
Which of the following compounds is a weak electrolyte?
1) HCl
2) CH3COOH (acetic acid)
3) C6H12O6 (glucose)
13.
3) acid-base
4) NH3(aq) and H2SO4(aq) do not react.
4) O2
5) NaCl
Identify the major ionic species present in an aqueous solution of Na 2CO3
1) Na2+, CO32– 2) Na2+, C2–, O3 3) Na+, C4+, O32– 4) Na+, C+, O2–
5) Na+, CO32–
14.
What element is reduced in the following chemical reaction?
Cu + 2 H2SO4  CuSO4 + SO2 + 2 H2O
1) Cu
15.
2) H
2) +3
1) a and d
2) a, b, c, and d
4) b, c, and d
5) c and d
A2Y
A2Z
2) 1.85  10–3
3) 1.85
4) 3.52
5) 0.104
The molarity of a solution prepared by diluting 43.72 mL of 5.005 M aqueous K2Cr2O7 to 500 mL is
2) 0.0044
3) 0.438
4) 0.0879
5) 0.870
3) HF
4) HBr
5) HCl
Which of the following is a weak acid?
1) H2SO4
2) HNO3
The concentration of iodide ions in a 0.193 M solution of barium iodide, BaI2(aq), is
1) 0.193 M
22.
3) a, c, and d
The molarity (M) of an aqueous solution containing 22.5 g of sucrose (C 12H22O11) (MW = 342) in 35.5 mL of
solution is
1) 57.2
21.
5) none of these
A2X is the strongest electrolyte and A2Y is the weakest electrolyte.
A2Y is the strongest electrolyte and A2X is a nonelectrolyte.
A2Y is the strongest electrolyte and A2Z is a nonelectrolyte.
A2Z is the strongest electrolyte and A2Y is the weakest electrolyte.
1) 0.0657
20.
4) +1
Three different substances, A2X, A2Y, A2Z, were dissolved in water with the following results. (Water
molecules are omitted for clarity.) Which of the substances is the strongest electrolyte, and which is the
weakest?
1)
2)
3)
4)
19.
3) +2
CH4(g) + O2(g)  CO2(g) + H2O()
CaO(s) + CO2(g)  CaCO3(s)
PbCO3(s)  PbO(s) + CO2(g)
CH3OH() + O2(g)  CO2(g) + H2O()
A2X
18.
5) H2O
Which of the following are combustion reactions?
a)
b)
c)
d)
17.
4) O
What is the oxidation state of the C atom in C2O42–?
1) +4
16.
3) S
2) 0.386 M
3) 0.0965 M
4) 0.579 M
5) 0.0643 M
Ethylene glycol, used as a coolant in automotive engines, has a specific heat capacity of 2.42
J/(g•K). Calculate q when 3.65 kg of ethylene glycol is cooled from 132 ˚C to 85 ˚C.
1) -1900 kJ
2) -420 kJ
3) -99 kJ
4) -0.42 kJ
5) -4.2  10–6 kJ
23.
A 275-g sample of nickel at 100.0 ˚C is placed in 100.0 mL of water at 22.0 ˚C. What is the
final temperature of the water? Assume that no heat is lost to or gained from the surroundings.
Specific heat capacity of nickel = 0.444 J/(g•K)
1) 39.6 ˚C
24.
2) 40.8 ˚C
3) 61.0 ˚C
4) 79.2 ˚C
5) 82.4 ˚C
A sample of baking soda dissolves in water. The initial temperature of the water is 23 ˚C. The
temperature of the resultant solution is 18 ˚C.
H2O
NaHCO3(s)
Na+(aq) + HCO3–(aq)
For the dissolving process described above, the reaction shown is _____ and heat is _____.
1) exothermic, given off by the reaction.
2) exothermic, taken in (absorbed) by the reaction.
3) endothermic, given off by the reaction.
4) endothermic, taken in (absorbed) by the reaction
5) cannot tell without knowing the mass of water and the specific heat of the resultant
solution
25.
A 13.8 mL aliquot of 0.176 M H3PO4(aq) is to be titrated with 0.110 M NaOH(aq). What volume (mL) of
base will it take to reach the equivalence point?
H3PO4(aq) + 3 NaOH(aq)  Na3PO4(aq) + 3 H2O()
1) 7.29
26.
2) 22.1
3) 199
4) 66.2
5) 20.9
Lithium and nitrogen react in a combination reaction to produce lithium nitride.
6 Li(s) + N2(g)  2 Li3N(g)
In a particular experiment, 3.50 g samples of each reagent are reacted. The theoretical yield of lithium
nitride is _____ g.
1) 3.52
27.
5) 8.70
2) H+ and OH–
3) K+ and NO3– 4) H+ and NO3– 5) OH– only
2) ∆E does not change
3) ∆E decreases
4) ∆E = q
A nitrogen oxide is 63.65% by mass nitrogen. The molecular formula could be _______.
1) NO
30.
4) 5.85
A gas expands against an external pressure while in thermal isolation from the surroundings. For this
expansion
1) ∆E increases
29.
3) 17.6
What are the spectator ions in the reaction between KOH(aq) and HNO3(aq)?
1) K+ and H+
28.
2) 2.93
2) NO2 3) N2O 4) N2O4 5) either NO2 or N2O4
Which one of the following conditions would always result in an increase in the internal energy of a
system?
1)
2)
3)
4)
5)
The system loses heat and does work on the surroundings.
The system gains heat and does work on the surroundings.
The system loses heat and has work done on it by the surroundings.
The system gains heat and has work done on it by the surroundings.
None of the above is correct.
1.
Which sketch represents an orbital that can have a principal quantum number of 2?
II
I
1) I
2.
2) II & III
IV
3) I, II, & III
V
4) IV & V
5) I, II, III, IV, & V
The wavelength of light associated with an electronic transition from n = 3 to n = 2 in a hydrogen atom is
1) 306 nm
3.
III
2) 361 nm
3) 434 nm
4) 656 nm
The electronic configuration of cadmium (Z = 48) is
1) 1s22s22p62d103s23p63d104s25d10
2) 1s22s22p63s23p64s23d104p65s24d10
3) 1s22s22p62d103s23p63d104s24p64d2
4) 1s21p62s22p62d103s23p63d104s24p2
5) 1s22s22p63s23p63d104s24p65s25p65d4
4.
Which is isoelectronic with N3–?
1) C4+
5.
2) F
2) 1s22s22p23s2 3) 1s22s22p1
2) t
4) v
4)
K+
3) O– is larger than O2–
is larger than Cl–
4) directly, inversely
5) indirectly, not
Which of the following is the correct electron configuration for the ground state of a Mn atom?
1) 1s22s22p63s33p63d7
2) 1s22s22p63s23p64s24d5
10.
3) u
The energy of a photon of light is _____ proportional to its frequency and _____ proportional to its
wavelength.
1) directly, directly
2) inversely, inversely
3) inversely, directly
9.
5) [He]2s22p4
Which of the following is true?
1) Fe3+ is larger than Fe2+
2) Na is larger than Na+
8.
4) 1s22s22p4
Elements with the outermost electron configuration ns2np3 are found in which portion of the periodic
table?
1) s
7.
4) Ne
Which one of the following configurations depicts an excited oxygen atom?
1) 1s22s22p2
6.
3) Na
3) 1s21p62s22p63s23p63d5
4) 1s22s22p63s23p64s23d5
What is the frequency of blue light that has a wavelength of 400 nm?
1) 2.65  10–31 Hz
2) 1.3  10–6 Hz 3) 7.5  105 Hz 4) 2.5  106 Hz 5) 7.5  1014 Hz
11.
Which of the following is a correct orbital diagram for the valence electrons in a S atom?
1)
2)
3s
3p
3s
3p
3)
4)
3s
3p
3s
3p
Use the following diagram to answer Q-12. Two electromagnetic waves are represented below.
12.
Wave (a) has the
1)
2)
3)
4)
13.
longer wavelength and higher frequency than wave (b).
longer wavelength and lower frequency than wave (b).
shorter wavelength and higher frequency than wave (b).
shorter wavelength and lower frequency than wave (b).
What is the energy in joules of a mole of photons associated with red light of wavelength 7.00  102 nm?
1) 256 kJ
14.
3) 4.72  10–43 J
2) Rb
3) Ge
2) The Rutherford gold foil experiment
4) The discovery of atomic line spectra
2) 1
3) 2
4) 3
5) 4
The uncertainty principle states that
1)
2)
3)
4)
5)
18.
4) Br
There are _____ unpaired electrons in a ground state phosphorus atom.
1) 0
17.
5) 2.12  1042 J
Which experiment caused Bohr to invoke an explanation that quantized the energy levels of electrons in
atoms?
1) The discovery of isotopes by mass spectrometry
3) The Mulliken oil-drop experiment
5) The placing of Mentos in 2 L bottles of Diet Coke.
16.
4) 12.4 kJ
Which has the highest Zeff for its valence electrons?
1) K
15.
2) 1.71  105 J
matter and energy are really the same thing.
it is impossible to know anything with certainty.
it is impossible to know both the exact position and momentum of an electron.
there can only be one uncertain digit in a reported number.
it is impossible to know how many electrons there are in an atom.
Of the following, which gives the correct increasing order for atomic radius for Mg, Na, P, Si and Ar? [start
with small one first]
1) Mg < Na < P < Si < Ar
2) Si < P < Ar < Na < Mg
3) Ar < Si < P < Na < Mg
4) Ar < P < Si < Mg < Na
5) Na < Mg < Si < P < Ar
19.
The spheres below represent atoms of Sb, As, P, and N (not necessarily in that order).
A, r = 75 pm
B, r = 110 pm
C, r = 120 pm
D, r = 140 pm
Which one of these spheres represents an atom of Sb?
1) sphere A
20.
For the fourth-shell orbital shown below, what are the principal quantum
numbers, n, and the angular momentum quantum number, ?
2) n = 4 and  = 1
2) [Ar]3d4
3) [Ar]3d6
2) temperature, pressure
3) volume, pressure
4) pressure, volume
When 50.0 mL of 0.400 M Ca(NO3)2 is added to 50.0 mL of 0.800 M NaF, CaF2 precipitates, as shown in
the net ionic equation below. The initial temperature of both solutions is 21.0 ˚C. Assuming that the
reaction goes to completion, and that the resulting solution has a mass of 100.00 g and a specific heat of
4.18 J/(g•˚C), calculate the final temperature of the solution.
Ca2+(aq) + 2 F–(aq)  CaF2(s)
1) 20.45 ˚C
25.
4) [Ar]3d104s1 5) [Ar]3d3
Heat transferred in a chemical reaction or physical change is expressed as ∆E when the heat is
transferred under constant _____and is expressed as ∆H when heat is transferred under constant _____
conditions.
1) pressure, temperature
24.
3) n = 4 and  = 2
Fe+2 ions are represented by __________.
1) [Ar]3d1
23.
4) sphere D
higher the orbital energy and the higher Zeff for the electron.
higher the orbital energy and the lower Zeff for the electron.
lower the orbital energy and the higher Zeff for the electron.
lower the orbital energy and the lower Z eff for the electron.
1) n = 4 and  = 0
4) n = 4 and  = 3
22.
3) sphere C
Within a given shell of a multielectron atom, the lower  for an orbital, the
1)
2)
3)
4)
1.
2) sphere B
2) 21.55 ˚C
∆Ho = –11.5 kJ
3) 22.10 ˚C
4) 22.65 ˚C
o
Given the data in the table, ∆H rxn for the reaction
Substance
2 CO(g) + O2(g)  2 CO2(g)
is __________ kJ.
1) –566.4

2) –283.2
o
5) The ∆H f of O2(g) is needed for the calculation

CO(g)
CO2(g)
CaCO3(s)
3) 283.2

4) –677.0
o
∆H f (kJ/mol)
–110.5
–393.7
–1207.0
26.
Hydrogen peroxide can decompose to water and oxygen by the following reaction:
2 H2O2(l)  2 H2O(l) + O2(g)
H = 196 kJ
Calculate the value of q when 5.00 g of H2O2(l) decomposes at constant pressure.
[Molar mass of H2O2 = 34.0 g]
1) 14.4 kJ
27.
4) +98 kJ
5) 28.8 kJ
2) 2, 2, 1, –1/2
3) 1, 0, 1, +1/2
4) 2, 1, +2, +1/2
5) 1, 1, 0, –1/2
Which of the following have the same number of valence electrons?
1) K, As, Br
29.
3) 98 kJ
Which of the following is a valid set of four quantum numbers? (n, , m, ms)
1) 2, 1, 0, +1/2
28.
2) 490 kJ
2) B, Si, As
3) N, As, Bi
4) He, Ne, F
Which atom in each group (I and II) has the smallest atomic radius? (I) K, Zn, Br (II) Al, Ga, In
1) K; Al
2) K; In
3) Br; Al
4) Br; In
30.
Write all the quantum numbers for the last electron in Scandium atom
1.
What is the formal charge on the oxygen atom in OCN- ion. The correct Lewis structure is shown.
OC–N
1) +2
2.
3) 0
4) –1
5) –2
2) O=PCl
3) O=P=Cl
4) PCl=O
Based on valence bond theory, the bond order of the N–N bond in the N2 molecule is _____.
1) 0
4.
_
Which one of the following is the best Lewis electron-dot formula for POCl?
1) P=OCl
3.
2) +1
1
2) 1
3) 2
4) 3
5) 5
Consider the following properties of an element:
(i)
It is solid at room temperature.
(ii) It easily forms an oxide when exposed to air.
(iii) When it reacts with water, hydrogen gas evolves.
(iv) It must be stored submerged in oil.
Which element fits the above description the best?
1) sulfur
5.
6.
2) copper
3) mercury
4) sodium
5) magnesium
Which is not generally considered to be a chemical reaction of the alkaline earth metal calcium?
1) Ca(s) + Cl2(g)  CaCl2(s)
2) 2 Ca(s) + 2 H2O()  2 Ca(OH)2(aq) + H2(g)
3) 6 Ca(s) + 2 N2(g)  2 Ca3N2(s)
4) Ca(s) + O2(g)  CaO2(s)
What is the general trend in ionization energy and electron affinity values?
1)
2)
3)
4)
Both decrease as one traverses a period from left to right and both decrease as one descends a group.
Both decrease as one traverses a period from left to right and both increase as one descends a group.
Both increase as one traverses a period from left to right and both decrease as one descends a group.
Both increase as one traverses a period from left to right and both increase as one descends a group.
7.
Select the best Lewis structure for P2I4.
1)
I
8.
I
I
P
P
I
I
P
2) bent
P
I
I
I
I
P
P
I
4) Si
5) Na
4) tetrahedral
5) see-saw
3) Br2
4) S8
5) Cr
Which one of the following equations correctly represents the process involved in the electron affinity of X?
X(g)
X+(g)
X+(g) + e–
X(g) + e–
X+(g) + Y–(g)





What is the molecular shape of
X+(g) + e–
X+(aq)
X(g)
X–(g)
XY(s)
SiF62–
2
F
as predicted by the VSEPR theory?
F
4) see-saw
5) octahedral
F
Si
F
For the SF4, the number of electron pairs (regions of electron density) around S is
molecular geometry is described as _______.
F
F
, and its
2) 5, seesaw
3) 4, seesaw
5) 5, trigonal bipyramidal
Which one of the following is a polar molecule?
1) PBr5
15.
P
3) trigonal planar
2) Si
1) 4, tetrahedral
4) cannot be determined
14.
I
3) Al
1) trigonal bibyramidal
2) hexagonal
3) tetrahedral
13.
I
I
In which of these substances are the atoms held together by metallic bonding?
1)
2)
3)
4)
5)
12.
P
I
According to VSEPR theory, a molecule with the general formula AX2E2 will have a _____ molecular shape.
1) CO2
11.
I
2) Mg
1) linear
10.
I
4)
An element M reacts with chlorine to form MCl2, with oxygen to form MO, and with nitrogen to form
M3N2. The most likely candidate for the element is.
1) Li
9.
3)
2)
2) CCl4
3) BrF5
4) XeF2
5) XeF4
Choose the incorrect statement about molecular geometries (shapes).
1) PCl5 is trigonal bipyramidal.
2) SO32– is tetrahedral.
3) XeF2 is linear.
4) The O–C–O angle in
CH3COH is about 120˚ (sp2)
O
16.
For which molecule can the bonding be described in terms of sp3 hybrid orbitals of the central atom?
1) SF6
17.
2) BF3
3) PCl5
4) NH3
5) BeH2
Which of the following is a correct Lewis structure for SO2?
1)
O
S
O
2) O
S
O
3)
O
S
O
4)
O
S
O
_
18.
Which bond in the Lewis structure shown is the most polar?
H
O
C
N
The electronegativities of the elements are: H = 2.1; O = 3.5; C = 2.5; N = 3.0
1) H–O
2) O–C
3) CN
4) They are equally polar.
Questions 19 and 20 refer to the Lewis structure for ozone that is shown.
O
19.
2) bent
4) trigonal pyramidal
2) bent
3) trigonal planar
4) trigonal pyramidal
Element A has an electronegativity of 2.8 and element B has an electronegativity of 3.5. Which
statement best describes the bonding in A3B?
1)
2)
3)
4)
22.
3) trigonal planar
Based on the VSEPR model, what is the molecular geometry for ozone?
1) linear
21.
O
Based on the VSEPR model, what is the electronic geometry for ozone?
1) linear
20.
O
The AB bond is largely covalent with a – on A.
The AB bond is largely covalent with + on A.
The compound is largely ionic with A as the cation.
The compound is largely ionic with A as the anion.
What are the bond angles in the molecule model of PF5 at the right?
1) some less than 90˚ and one less than 180˚
2) 90˚ and 180˚
3) some less than 90˚ and some less than 120˚ but greater than 90˚, one less than
180˚ but greater than 120˚
4) 90˚, 120˚, and 180˚
23.
Which of the following molecules or ions will exhibit delocalized bonding?
SO2
1) SO2, SO3, and SO32–
4) SO3 and SO32–
24.
2) SO32– only
3) SO2 and SO3
5) None of the above will exhibit delocalized bonding.
2) 1
3) 2
4) 3
5) 4
According to modern bonding theory the number of sigma () and pi (π) bonds in the ethylene molecule
H2C=CH2.
1) 4 π and 1  2) 5 π and 1 
26.
SO32–
How many lone pairs of electrons are on the iodine atom of IF 4+?
1) 0
25.
SO3
3) 4  and 1 π
4) 5  and 1 π
5) 4  and 2 π
Which compound contains the longest carbon-to-nitrogen bond?
1) H
H
H
C
N
H
H
2)
H
C
N
3) H
H
H
C
C
H
N
H
4) All are equal.
27.
Which of the following has the bonds correctly arranged in order of increasing polarity?
1) Be–F, Mg–F, N–F, O–F
4) N–F, Be–F, Mg–F, O–F
28.
2) O–F, N–F, Be–F, Mg–F
5) Mg–F, Be–F, N–F, O–F
3) O–F, Be–F, Mg–F, N–F
Using the table of bond dissociation energies, the ∆H for the following reaction is ____ kJ.
Bond
H–Cl
F–F
H–F
Cl–Cl
2 HCl(g) + F2(g)  2 HF(g) + Cl2(g)
1) -359
29.
4) 223
5) 208
A reactive element with a relatively low electronegativity would be expected to have a relatively
1)
2)
3)
4)
30.
2) -223 3) 359
D (kJ/mol)
431
155
567
242
small negative electron affinity and a relatively low ionization energy.
small negative electron affinity and a relatively high ionization energy.
large negative electron affinity and a relatively low ionization energy.
large negative electron affinity and a relatively high ionization energy.
Which of the following are allowed resonance forms of NCS–?
1) only I
_
I
N
C
S
II
N
C
S
III
N
C
S
_
_
2) only II
_
and
N
C
S
and
N
C
S
and
S
C
N
_
_
3) only III
4) I and III
A19. How many nitrogen atoms are there in 60.0 g of ammonium nitrate sample?
B14. Using standard enthalpy of formation data given in Appendix-C (at the end of the book), calculate H for
the combustion of Octane.
B15. How many molecules of oxygen is needed to completely burn 275 mL of Octane
[density = 0.703 g/cm3]?
B18. Based on the picture giving in B17, if you have 8.0 moles of A and 6.0 moles of B2, how many moles of
A3B4 will be formed?
(A) 5.3 mol (B) 6.8 mol
(C) 2.7 mol
(D) 1.8 mol