CHE 132 Spring 2022 Final Exam Version A
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Stony Brook University *
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Course
132
Subject
Chemistry
Date
Jan 9, 2024
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CHE 132 Final Exam Form 1A Spring 2022 Directions 1.
Place all other items besides your exam out of sight and or beneath your seat.
2.
Place your ID (face up), pencils, erasers, and a scientific calculator at your desk now. You
may be asked to show proof of ID and sign an attendance sheet.
3.
On your scantron, starting from the left, write and bubble
–
in your (leave extra spaces
blank):
SBID NUMBER under IDENTIFICATION NUMBER
4.
Print your name clearly and sign.
5.
DO NOT OPEN YOUR EXAM until told to do so.
6.
All scantron forms must be submitted after 150
minutes. No extra time will be given.
7.
All answers must be entered on the Scantron answer sheet, which you must turn in.
8.
Answer sheets will not be returned, so record your answers next to each question on the
exam for comparison with the answers when they are posted on Blackboard. Use the
space between questions, the backs of pages, or the extra sheets for analysis and
calculations.
9.
There are 40
questions on this exam and 8
total pages.
10.
For multiple choice questions, bubble in the letter of the one choice that best completes
the statement or answers the question on the scantron form. No credit will be given for
multiple answers.
11.
Short Answer/Numerical –
Check to make sure that you
•
skip the multiple
–
choice scantron # when you have a short answer.
•
e
nter your answers on the right
–
hand side of the scantron in the appropriately
numbered line
.
•
Work will not be graded or afforded partial credit.
12.
Once you open the exam, be sure you have all the pages including a Periodic Table and
a sheet of equations/conversion factors.
General Chemistry Policies:
Possession of cell phones, smart watches, other communication devices, or any unauthorized
materials during the exam will result in a grade of 0 on this exam, a report to Academic
Judiciary, and a possible grade of F for the course.
If you do not have an ID, note that on the sign
–
in sheet, and you must see Dr. Amarante the
day after the exam and show her a valid ID.
Blank Page
CHE 132 Final Exam Equations D
H
soln
= -E
lattice
+
D
H
hydration
K = °
C + 273.15 °
F = 9/5 (
°
C) + 32
d=m/V
M
1
•V
1
=M
2
•V
2
∆࠵? = ࠵? + ࠵?
∆࠵? = ࠵?
!
࠵? = ࠵?࠵?
"!
∆࠵?
∆࠵?
#$%
°
= ∑ ࠵? ∆࠵?
’
°
(࠵?࠵?࠵?࠵?࠵?࠵?࠵?࠵?) − ∑ ࠵? ∆࠵?
’
°
(࠵?࠵?࠵?࠵?࠵?࠵?࠵?࠵?࠵?)
࠵? = ࠵?
(
࠵?࠵?࠵?
∆࠵? =
)
!"#
*
∆࠵?
"+##
= −
∆-
*
∆࠵?
.#/%"
=
∆-
$!%&’
*
$!%&’
∆࠵?
+%01
= ∆࠵?
"2"
+ ∆࠵?
"+##
∆࠵?
#$%
°
= ∑ ࠵? ࠵?
°
(࠵?࠵?࠵?࠵?࠵?࠵?࠵?࠵?) − ∑ ࠵? ࠵?
°
(࠵?࠵?࠵?࠵?࠵?࠵?࠵?࠵?࠵?)
D
G =
D
H-T
D
S
∆࠵?
#$%
°
= ∑ ࠵? ∆࠵?
’
°
(࠵?࠵?࠵?࠵?࠵?࠵?࠵?࠵?) − ∑ ࠵? ∆࠵?
’
°
(࠵?࠵?࠵?࠵?࠵?࠵?࠵?࠵?࠵?)
Rate=k[A]
m
[B]
n
࠵?
3/5
=
3
6[8]
(
࠵?
3/5
=
[8]
(
56
࠵?
3/5
=
:.<=>
6
ln[࠵?]
.
= −࠵?࠵? + ln[࠵?]
:
3
[8]
$
= ࠵?࠵? +
3
[8]
(
[
࠵?]
.
= −࠵?࠵? + [࠵?]
:
[࠵?]
.
= [࠵?]
:
࠵?
?6.
࠵? = ࠵?࠵?
?
)
%
*+
࠵?࠵?
6
,
6
-
= −
@
%
A
E
3
*
,
−
3
*
-
F
ln ࠵? = −
@
%
A
3
*
+ ln ࠵?
࠵?࠵?
!
= − log ࠵?
!
࠵?࠵? = − log[࠵?
"
࠵?
#
]
࠵?࠵? = ࠵?࠵?
$
+ ࠵?࠵?࠵?
[&
—
]
[(&]
࠵?࠵?࠵? = ࠵?࠵?
!
+ ࠵?࠵?࠵?
[)(
"
]
[)]
%࠵?࠵?࠵?࠵?࠵?࠵?࠵?࠵?࠵?࠵?࠵?࠵? =
[(
#
*
"
]
$%
[(&]
&’&(&)*
࠵? 100%
∆࠵?° = −࠵?࠵?࠵?࠵?࠵?
∆࠵?° = −࠵?࠵?࠵?
+,--
°
࠵?
+,--
°
=
/0
12
࠵?࠵?࠵?
࠵?
+,--
= ࠵?
+,--
°
—
/0
12
࠵?࠵?࠵?
࠵?
+,--
°
= ࠵?
+$3456,
°
− ࠵?
$156,
°
or
࠵?
+,--
°
= ࠵?
7,68+3951
°
+ ࠵?
5:96$3951
°
࠵?
+,--
°
=
;.;=>?@
1
࠵?࠵?࠵?࠵? ࠵?࠵? 298.15 ࠵?
࠵?
+,--
= ࠵?
+,--
°
—
;.;=>? @
1
Units Conversions 1 m
= 1.0936 yds
1 cm = 0.39370 inch
760 torr = 1 atm
1 inch
= 2.54 cm
1 km
= 0.621371 mile
760 mm Hg = 1 atm
1 Å
= 10
-10
m
1 mL = 1 cm
3
101,325 Pa = 1 atm
1 kg
= 2.2046 lbs
1 L
= 10
-3
m
3
1.013 bar = 1 atm
1 lb
= 453.59 g
1 gallon
= 3.7854 L
1 Pa = 1 N/m
2
1 lb
= 16 oz
1 quart
= 0.94633 L
4.184 J = 1 cal
1 ton
= 2000 lbs
1 mol = 6.022 X 10
23
4 qts = 1 gallon
C
=
k
H
P
P
solution
=
χ
solvent
P
solvent
o
Δ
T
=
K
b
m
solute
Δ
T
=
K
f
m
solute
Π
=
MRT
Δ
T
=
iK
f
m
solute
Π
=
iMRT
Δ
T
=
iK
b
m
solute
K
c
=
[
C
]
c
[
D
]
d
[
A
]
a
[
B
]
b
K
p
=
p
C
c
p
D
d
p
A
a
p
B
b
K
p
=
K
c
(
RT
)
Δ
n
a
x
=
x
⎡
⎣
⎤
⎦
°
c
a
i
=
p
i
p
ref
x
=
−
b
±
b
2
−
4
ac
2
a
K
w
= [H
3
O
+
][OH
—
] =[H
+
][OH
—
]= 1.0 x 10
-14
at 25 °C K
a
xK
b
=
K
w
pK
a
=
−
log
K
a
pOH
=
−
log[
OH
−
]
pH
+
pOH
=
14.00
=
pK
w
࠵?࠵?࠵?࠵?
at 298.15 K
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Table 1. K
a
and K
b
values for common acids and bases at 25.0°C. Table 2. Physical properties and constants for solvents. Solvent Boiling Point (°C) K
b
(°C/m)
Freezing Point (°C) K
f
(°C/m) Water 100.0 0.512 0.0 1.86 Ethanol 78.2 1.22 -114.1
12.5 Aniline 184.3 3.69 -5.96
5.87 Table 3. Henry’s Law constants for select gases at 25.0 °C. Gas k
H
(M/atm) O
2
1.3 x 10
-3
N
2
6.1 x 10
-4
CO
2
3.4 x 10
-2
NH
3
5.8 x 10
1
Name K
a
Name K
b
H
2
C
2
O
4
(Ka
1
) 5.9 x 10
-2
Ethylamine (CH
3
CH
2
NH
2
) 4.3 x 10
-4
HSO
4
–
1.2 x 10
-2
Methylamine (CH
3
NH
2
)
4.4 x 10
-4
H
2
SO
3 (Ka
1
) 1.2 x 10
-2
Dimethylamine [(CH
3
)
2
NH]
5.9 x 10
-4
Chlorous acid (HClO
2
)
1.1 x 10
-2
Diethylamine [(CH
3
CH
2
)
2
NH]
8.6 x 10
-4
Citric acid (H
3
C
6
H
5
O
7
, Ka
1
) 7.4 x 10
-3
Trimethylamine [(CH
3
)
3
N] 7.4 x 10
-5
H
3
PO
4
(Ka
1
) 7.2 x 10
-3
Ammonia (NH
3
)
1.76 x 10
-5
Hydrofluoric acid (HF) 6.8 x 10
-4
Pyridine (C
5
H
5
N)
1.7 x 10
-9
Nitrous acid (HNO
2
)
4.6 x 10
-4
Hydroxylamine (NH
2
OH) 6.6 x 10
-9
Cyanic acid (HCNO) 3.5 x 10
-4
Aniline (C
6
H
5
NH
2
)
4.0 x 10
-10
Formic acid (HCOO
H
) 1.8 x 10
-4
Lactic acid 1.4 x 10
-4
HC
2
O
4
–
6.4 x 10
-5
Benzoic acid (C
6
H
5
COO
H
) 6.3 x 10
-5
Hydrazoic acid (HN
3
) 1.9 x 10
-5
Acetic acid (CH
3
COO
H
)
1.8 x 10
-5
Propanoic acid (CH
3
CH
2
COO
H
)
1.3 x 10
-5
H
2
CO
3 (Ka
1
) 4.2 x 10
-7
H
2
PO
4
– (Ka
2
) 6.3 x 10
-8
HSO
3
–
6.2 x 10
-8
HClO 2.9 x 10
-8
HBrO 2.3 x 10
-9
HCO
3
— (Ka
2
) 4.8 x 10
-11
HPO
4
2— (Ka
3
) 4.2 x 10
-13
Table
4
. Solubility Product Constants at 25 °C.
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Table 5. Formation Constants at 25
o
C
Table
6
. Standard Reduction Potentials in Aqueous Solution at 25 °C.
Reduction Half-Reaction
E°(V)
F
2
(
g
) + 2 e
−
→
2 F
-
(
aq
)
+2.87
Au
+
(
aq
) + e
−
→
Au (
s
)
+1.83
→
Co
2+
(
aq
)
+1.82
Co
3+
(
aq
) + e
−
H
2
O
2
(l) + 2 H
+
(
aq
) + 2
e−
→
2 H
2
O (
l
)
+1.78
Au
3+
(
aq
) + 3 e
−
→
Au (
s
)
+1.52
MnO
4
−
(
aq
)
→
Mn
2+
(
aq
)
unbalanced under acidic conditions
+1.50
Cl
2
(
g
) + 2 e
−
→
2 Cl
−
(
aq
)
+1.36
MnO
2
(
s
)
→
Mn
2+
(
aq
)
unbalanced under acidic conditions
+1.23
Pt
2+
(
aq
) + 2 e
−
→
Pt (
s
)
+1.20
Ag
+
(
aq
) + e
−
→
Ag (s)
+0.7994
Fe
3+
(
aq
) + e
−
→
Fe
2+
(
aq
)
+0.771
MnO
4
−
(
aq
)
→
MnO
2
(
s
)
unbalanced under basic conditions
+0.59
I
2
(g) + 2 e
−
→
2 I
−
(
aq
)
+0.535
Cu
+
(
aq
) + e
−
→
Cu (
s
)
+0.521
Cu
2+
(
aq
) + 2 e
−
→
Cu (
s
)
+0.337
2 H
3
O
+
(aq) + 2 e
−
→
H
2
(
g
) + 2 H
2
O (
l
)
0.000
Fe
3+
(
aq
) + 3 e
−
→
Fe (
s
)
−0.036
Pb
2+
(
aq
) + 2 e
−
→
Pb (
s
)
−0.126
Sn
2+
(
aq
) + 2 e
−
→
Sn (
s
)
−0.1315
Ni
2+
(
aq
) + 2 e
−
→
Ni (
s
)
−0.25
Co
2+
(
aq
) + 2 e
−
→
Co (
s
)
−0.28
Cd
2+
(
aq
) + 2 e
−
→
Cd (
s
)
−0.40
Cr
3+
(
aq
) + e
−
→
Cr
2+
(
aq
)
−0.41
Fe
2+
(
aq
) + 2 e
−
→
Fe (
s
)
−0.44
Zn
2+
(
aq
) + 2 e
−
→
Zn (
s
)
−0.763
Cr
3+
(
aq
) + 3 e
−
→
Cr (
s
)
−0.74
Cr
2+
(
aq
) + 2 e
−
→
Cr (
s
)
−0.91
Mn
2+
(
aq
) + 2 e
−
→
Mn (
s
)
−1.18
V
2+
(
aq
) + 2 e
−
→
V (
s
)
−1.18
Zr
4+
(
aq
) + 4 e
−
→
Zr (
s
)
−1.53
Al
3+
(
aq
) + 3 e
−
→
Al (
s
)
−1.66
Mg
2+
(
aq
) + 2 e
−
→
Mg (
s
)
−2.356
Ba
2+
(
aq
) + 2 e
−
→
Ba (
s
)
−2.90
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�
1 2 3 4 5 6 Hydrogen 1 H 1.008 Key: Lithium Beryllium Element Name 3 4 Atomic number Li Be Symbol 6.94 9.0122 Atomic weiaht /mean relative mass) Sodium Magnesium 11 12 Na Mg 22.990 24.305 Potassium Calcium Scandium Titanium Vanadium Chromium 19 20 21 22 23 24 K Ca Sc Ti
V Cr 39.098 40.078(4) 44.956 47.867 50.942 51.996 Rubidium Strontium Yttrium Zirconium Niobium Molybdenum 37 38 39 40 41 42 Rb Sr y Zr Nb Mo 85.468 87.62 88.906 91.224/2) 92.906/2) 95.95 Caesium Barium Lutetium Hafnium Tantalum Tungsten 55 56 57-70
71 72 73 74 Cs Ba *
Lu Hf Ta
w 132.91 137.33 174.97 178.49(2) 180.95 183.84 Francium Radium Lawrencium Rutherfordium Dubnium Seaborgium 87 88 89-102
103 104 105 106 Fr Ra ** Lr Rf Db Sg 1223.021 1226.031 1262.111 1267.121 1270.131 1269.131 Lanthanum Cerium Praseodymiurr Neodymium 57 58 59 60 *lanthanoids
La Ce Pr Nd 138.91 140.12 140.91 144.24 Actinium Thorium Protactinium Uranium 89 90 91 92 **actinoids Ac Th Pa u
[227.031 232.04 231.04 238.03 The periodic table www.webelements.com 7 8 9 10 11 Manganese Iron Cobalt Nickel Copper 25 26 27 28 29 Mn Fe Co Ni Cu 54.938 55.845(2) 58.933 58.693 63.546(3) Technetium Ruthenium Rhodium Palladium Silver 43 44 45 46 47 Tc Ru Rh Pd Ag [98.9061 101.07/2) 102.91 106.42 107.87 Rhenium Osmium Iridium Platinum Gold 75 76 77 78 79 Re Os Ir Pt Au 186.21 190.23(2) 192.22 195.08 196.97 Bohrium Hassium Meitrlerium Darmstadtium Roentgenium 107 108 109 110 111 Bh Hs Mt Ds Rg 1270.131 1270.131 1278.161 1281.171 1281.171 Promethium Samarium Europium Gadolinium Terbium 61 62 63 64 65 Pm Sm Eu Gd Tb 1144.911 150.36/2) 151.96 157.25/3) 158.93 Neptunium Plutonium Americium Curium Berkelium 93 94 95 96 97 Np Pu Am Cm Bk [237.051 [244.061 [243.061 [247.071 [247.071 12 13 14 15 16 17 Boron Carbon Nitrogen Oxygen Fluorine 5 6 7 8 9 B C N 0 F 10.81 12.011 14.007 15.999 18.998 Aluminium Silicon Phosphorus Sulfur Chlorine 13 14 15 16 17 Al Si p s Cl 26.982 28.085 30.974 32.06 35.45 Zinc Gallium Germanium Arsenic Selenium Bromine 30 31 32 33 34 35 Zn Ga Ge As Se Br 65.38(2) 69.723 72.630(8) 74.922 78.971(8) 79.904 Cadmium Indium Tin Antimony Tellurium Iodine 48 49 50 51 52 53 Cd In Sn Sb Te I 112.41 114.82 118.71 121.76 127.60/3) 126.90 Mercury Thallium Lead Bismuth Polonium Astatine 80 81 82 83 84 85 Hg Tl
Pb Bi Po At 200.59 204.38 207.2 208.98 [208,981 [209,991 Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine 112 113 114 115 116 117 Cn Nh Fl Mc Lv Ts 1285.181 1286.181 1289.191 1289.191 1293.201 1293.211 Dysprosium Holmium Erbium Thulium Ytterbium 66 67 68 69 70 Dy Ho Er Tm Yb 162.50 164.93 167.26 168.93 173.05 Calfornium Einsteinium Fermium Mendelevium Nobelium 98 99 100 101 102 Cf Es Fm Md No [251.081 [252.081 [257.101 [258.101 [259.101 Symbols and names: the symbols and names of the elements, and their spellings are those recommended by the International Union of Pure and Applied Chemistry (IUPAC -
http://www.iupac.org/
). In some countries, the spellings alumlnum, cesium, and sulphur are usual. Group labels: the numeric system (1-18) used here is the current IUPAC convention. 18 Helium 2 He 4.0026 Neon 10 Ne 20.180 Argon 18 Ar 39.948 Krypton 36 Kr 83.798(2) Xenon 54 Xe 131.29 Radon 86 Rn 1222.021 Oganesson 118 Og 1294.211 Atomic weights (mean relative masses): these are the IUPAC 2013 values and given to 5 significant figures. The last significant figure of each value is considered reliable to ±1 except where a larger uncertainty is given in parentheses. IUPAC representative values are given for those elements having an atomic weight interval (H, Li, B, C, N, 0, Si, S, Cl, Tl). Elements for which the atomic weight is listed within square brackets have no stable nuclides and are represented by the element's longest lived isotope reported in the IUPAC 2013 values except Tc for which the value of Tc-99 given as that is the most commonly used isotope. ©2019 Prof Mark J Winter [WebElements Ltd and University of Sheffield]. All rights reserved. For updates to this table see https://www.webelements.com/nexus/printable-periodic-table/ (Version date: 22 June 2019)
.
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Final Exam, CHE 132 Spring 2022, Form 1A, Page 1
of 8
Final Exam CHE 132 Spring 2022 Form 1A
Short Answer Questions. Put the value in the blank for questions 1–10 with the appropriate number of significant figures.
1)
The osmotic pressure of a benzene solution containing 5.00 g of polystyrene per
liter was found to be 7.60 torr at 25°C. Estimate the average molecular weight
of the polystyrene in this sample.
2)
A concentration cell consists of two Al/Al
3+
electrodes. The electrolyte in
compartment A is 0.050 M Al(NO
3
)
3
and in compartment B is 1.25 M
Al(NO
3
)
3
. What is the voltage
(V)
of the cell at 25°C?
3)
What is the value of the (a) equilibrium constant and (b) Δ
G
° for the cell
reaction below
at 25°C?
2Cr(
s
) + 3Pb
2+
(
aq
) ⇄
3Pb(
s
) + 2Cr
3+
(
aq
) 4)
For the complex ion, Na
3
[CoCl
6
],
(a) What is the spin (low vs high)?
(b) The number of unpaired electrons.
(c) dia– vs paramagnetic?
5)
The coordination compound [V(OH
2
)
6
]Cl
3
has a
max
= 550 nm.
(a) Estimate the crystal field splitting energy (in kJ/mol).
(b) What color is being absorbed?
(c) What color would be an aqueous solution of the compound?
6)
What is the approximate pH of a weak acid–strong base titration if 25.0 mL of
aqueous 0.1900 M formic acid is titrated with 29.8 mL of 0.1594 M NaOH?
7)
Chromium crystallizes with a body–centered cubic unit cell. The radius of a
chromium atom is 128 pm. Calculate the density of solid crystalline chromium
in g/cm
3
.
8)
Calculate the ΔG°
rxn
using the following information at 298K.
2 HNO
3
(
aq
) + NO(
g
) → 3 NO
2
(
g
) + H
2
O(
l
) ΔG°
rxn
= ? ΔH
f ° (kJ/mol) –207.0 91.3 33.2 –285.8 S°(J/mol∙K) 146.0 210.8 240.1 70.0
Final Exam, CHE 132 Spring 2022, Form 1A, Page 2
of 8
9)
If you have 0.035 moles of SO
2
, 0.500 moles of SO
2
Cl
2
, and 0.080 moles of
Cl
2
are combined in an evacuated 5.00 L flask and heated to 100
o
C.
SO
2
Cl
2
(g) ⇄
SO
2
(g) + Cl
2
(g) K
c
= 0.078 at 100
o
C (a)
What is Q before the reaction begins?
(b) Which direction will the reaction proceed in order to establish
equilibrium?
10)
Ammonia will react with oxygen in the presence of a copper catalyst to form
nitrogen and water. From 164.5°C to 179.0°C, the rate constant increases by a
factor of 4.27. What is the activation energy
(kJ/mol)
of this oxidation reaction?
Multiple Choice Section
On the Scantron form, bubble in the letter of the one choice that best completes the statement or answers the question. No credit will be given for multiple answers. 11)
Determine the cell notation for the redox reaction given below.
3 Br
2
(
g
) + 2 Fe(
s
)
→
6 Br
⁻
(
aq
) + 2 Fe
3
+
(
aq
)
11) ______
A) Fe(
s
)
∣
Fe
3
+
(
aq
)
Br
2
(
g
)
∣
Br
⁻
(
aq
)
∣
Pt
B) Fe(
s
)
∣
Br
2
(
g
)
Fe
3
+
(
aq
)
∣
Br(
aq
)
∣
Pt
C) Fe
3
+
(
aq
)
∣
Fe(
s
)
Br
⁻
(
aq
)
∣
Br
2
(
g
)
∣
Pt
D)
Br
⁻
(
aq
)
∣
Br
2
(
g
)
∣
Pt
Fe
3
+
(
aq
)
∣
Fe(
s
)
E) Br
2
(
g
)
∣
Br
⁻
(
aq
)
∣
Pt
Fe(
s
)
∣
Fe
3
+
(
aq
)
12)
Give the major force in NaBr solution.
12) ______
A) dipole–dipole
B) nonpolar forces
C) ion–dipole
D) carbon bonding
E) ion–ion
13)
What is the pH of a buffer solution that is 0.192 M in lactic acid and 0.155 M in sodium lactate?
13) ______
A) 10.24
B) 3.94
C) 5.48
D) 14.09
E) 3.76
14)
What type of hybridization (according to valence bond theory) does Pt exhibit in the complex
ion, [Pt(NH
3
)
4
]
2
+
?
14) ______
A) sp
3
B) dsp
2
C) dsp
D) dsp
3
E) d
2
sp
2
Final Exam, CHE 132 Spring 2022, Form 1A, Page 3
of 8
15) Which of the following compounds can exhibit fac–mer isomerism?
15) ______ A) [Ni(H
2
O)
3
(CO)
3
]
3
+ B) [Cu(CO)
5
I]
+ C) [Cr(H
2
O)
4
Cl
2
]
+ D) [Co(CO)
5
NO
2
]
2
+ E) [Ti(NH
3
)
2
(H
2
O)
4
]
2
+ 16) Express the equilibrium constant for the following reaction.
4 NH
3
(
g
)
⇔
2 N
2
(
g
) + 6 H
2
(
g
)
16) ______ A) K =
B) K =
C) K =
D) K =
E) K =
17) Which of the following is the
strongest
reducing agent?
17) ______ A) Ca
2+
(
aq
)
B) Mg(
s
)
C) Ba(
s
)
D) K
+
(
aq
)
E) Li(
s
)
18) Determine the chemical formula for the compound, diamminetetraaquanickel(II) iodide.
18) ______ A) [Ni(NH
3
)
2
][(H
2
O)
4
I]
B) [Ni(NH
3
)
2
(H
2
O)
4
]I
2 C) [Ni(H
2
O)
4
][(NH
3
)
2
I]
D) [Ni(NH
3
)
2
(H
2
O)
4
]I
3 E) [Ni(NH
3
)
2
(H
2
O)
4
I]
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Final Exam, CHE 132 Spring 2022, Form 1A, Page 4
of 8
19) The equilibrium constant is given for one of the reactions below.
Determine the value of the
missing equilibrium constant.
Deuterium, D, is an isotope of hydrogen.
2 HD(
g
)
⇌
H
2
(
g
) + D
2
(
g
)
K
c
=
0.28
6 H
2
(
g
) + 6 D
2
(
g
)
⇌
12 HD(
g
)
K
c
= ?
19) ______ A) 1.2
B) 0.81
C) 2075
D) 0.00048
E) 0.78
20) What are the coefficients in front of NO
3
–
(
aq
) and Cu(
s
) when the following redox equation is
balanced in an acidic solution:
________ NO
3
–
(
aq
) + ________ Cu(
s
)
→
________ NO(
g
) + ________ Cu
2
+
(
aq
)?
20) ______ A) 2, 6
B) 3, 4
C) 3, 6
D) 2, 3
21) Given the following proposed mechanism, predict the rate law for the overall reaction.
A
2
+ 2 B
→
2 AB
(overall reaction)
Mechanism
A
2
⇌
2 A
fast
A + B
→
AB
slow
21) ______ A) Rate = k[A
2
]
1/2
[B]
1/2
B) Rate = k[A] [B]
2
C) Rate = k[A
2
]
2
D) Rate = k[A
2
][B]
E) Rate = k[A
2
]
1/2
[B]
22) Above what temperature does the following reaction become nonspontaneous?
FeO(
s
) + CO(
g
)
→
CO
2
(
g
) + Fe(
s
)
Δ
H = –11.0 kJ;
Δ
S = –17.4 J/K
22) ______ A) 632 K
B) 298 K
C) 191 K
D) This reaction is nonspontaneous at all temperatures.
E) This reaction is spontaneous at all temperatures.
Final Exam, CHE 132 Spring 2022, Form 1A, Page 5
of 8
23) A solution is prepared by adding 0.10 mol of ammonium nitrite, NH
4
NO
2
, to 1.00 L of water. Which statement about the solution is correct? 23) ______ A) The solution is neutral. B) The solution is acidic. C) The concentrations of potassium ions and acetate ions will be identical. D) The concentration of acetate ions will be greater than the concentration of potassium ions. E) The solution is basic. 24) Determine the partial pressure of oxygen necessary to form an aqueous solution that is 5.9 ×
10
–4 M O
2
at 25°C.
The Henry
ʹ
s law constant for oxygen in water at 25°C is 1.3 × 10
–3
M/atm.
24) ______ A) 2.2 atm
B) 0.78 atm
C) 0.45 atm
D) 1.3 atm
E) 0.61 atm
25) Which of the following pairs of coordination compounds or complex ions are examples of
coordination isomers?
25) ______ A) [Fe(NH
3
)
5
NO
2
]
2
+
and [Fe(NH
3
)
5
ONO]
2
+ B) [Fe(NH
3
)
2
(H
2
O)
4
]F
2 and [Fe(NH
3
)
2
(H
2
O)
4
]I
2 C) [MnCl
3
Br]
2
–
and [MnClBr
3
]
2
– D) [Cu(CO)
5
I]F and [Cu(CO)
5
F]I
E) [Fe(NH
3
)
2
(H
2
O)
4
]Br
2
and [Fe(NH
3
)
4
(H
2
O)
2
]Br
2 26) Use the standard half–cell potentials listed below to calculate the standard cell potential for the
following reaction occurring in an electrochemical cell at 25°C. (The equation is balanced.)
2 K(
s
) + I
2
(
s
)
→
2 K
⁺
(
aq
) + 2 I
⁻
(
aq
)
K
+
(
aq
) + e
⁻
→
K(
s
)
E° = –2.93 V
I
2
(
s
) + 2 e
⁻
→
2 I
⁻
(
aq
)
E° = +0.54 V
26) ______ A) +1.85 V
B) –5.32 V
C) +5.32 V
D) +6.40 V
E) +3.47 V
Final Exam, CHE 132 Spring 2022, Form 1A, Page 6
of 8
27) Use the tabulated half–cell potentials to calculate
Δ
G° for the following redox reaction.
2 Al(
s
) + 3 Mg
2
+
(
aq
)
→
2 Al
3
+
(
aq
) + 3 Mg(
s
)
Mg
2
+
(
aq
) + 2 e
⁻
→
Mg(
s
)
E° = –2.38 V
Al
3
+
(
aq
) + 3 e
⁻
→
Al(
s
)
E° = –1.66 V
27) ______ A) +4.1 × 10
2
kJ
B) +6.8 × 10
2
kJ
C) +1.4 × 10
2
kJ
D) –2.3 × 10
2
kJ
E) –7.8 × 10
2
kJ
28) Given the following balanced equation, determine the rate of reaction with respect to [SO
2
].
2 SO
2
(
g
) + O
2
(
g
)
→
2 SO
3
(
g
)
28) ______ A) Rate =
+
B) Rate =
C) Rate = –
D) Rate =
+
E) It is not possible to determine without more information.
29) A solution with a hydroxide ion concentration of 4.15 × 10
–
7
M is ________ and has a hydrogen
ion concentration of ________.
29) ______ A) acidic, 2.41 × 10
–
8
M
B) basic, 2.41 × 10
–
7
M
C) acidic, 2.41 × 10
–
7
M
D) acidic, 4.15 × 10
–
4
M
E) basic, 2.41 × 10
–
8
M
30) The ________ Law of Thermodynamics states the entropy of a perfect crystal at absolute zero is
zero.
30) ______ A) Zero
B) First
C) Second
D) Third
E) Fourth
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Final Exam, CHE 132 Spring 2022, Form 1A, Page 7
of 8
31) The following reaction is first order, C
2
H
6
→
2 CH
3
. If the rate constant is equal to 5.5 × 10
–4
s
–1
at 1000 K, how long will it take for 0.35 mol of C
2
H
6
in a 1.00 L container to decrease to 0.10
mol in the same container?
31) ______ A) 106 min
B) 38 min
C) 7.6 min
D) 131 min
32) Calculate the pH of a 0.200 M NaCH
3
CO
2
solution.
32) ______ A) 9.02
B) 4.98
C) 2.72
D) 11.28
33) The reaction below has a K
c
value of 1.0 × 10
12
.
What is the value of K
p
for this reaction at 400
K?
2 SO
2
(
g
) + O
2
(
g
)
⇌
2 SO
3
(
g
)
33) ______ A) 3.0 × 10
–14 B) 3.0 × 10
–
10 C) 3.0 × 10
10
D) 3.3 × 10
–11 E) 3.3 × 10
13
34) How many milliliters of 0.0991 M LiOH are required to titrate 25.0 mL of 0.0839 M HCl to the
equivalence point?
34) ______ A) 29.5
B) 4.58
C) 0.333
D) 0.208
E) 21.2
35) Identify the geometry of [PdCl
4
]
2
–
.
35) ______ A) trigonal bipyramidal
B) linear
C) tetrahedral
D) square planar
E) octahedral
36) Identify the location of oxidation in an electrochemical cell.
36) ______ A) the cathode
B) the electrode
C) the anode
D) the socket
E) the salt bridge
Final Exam, CHE 132 Spring 2022, Form 1A, Page 8
of 8
37) Choose the electron configuration for Ti
2+
.
37) ______ A) [Ne]3s
2
3p
6
B) [Ar]3d
2
C) [Ar]4s
2
D) [Ar]4s
2
3d
4
E) [Ar]4s
2
3d
2
38) Which of the following solutions is a good buffer system?
38) ______ A) a solution that is 0.10 M HCl and 0.10 M NH
4
+ B) a solution that is 0.10 M HF and 0.10 M LiC
2
H
3
O
2
C) a solution that is 0.10 M NaOH and 0.10 M KOH
D) a solution that is 0.10 M HC
2
H
3
O
2
and 0.10 M LiC
2
H
3
O
2
E) None of the above is a buffer system.
39) What is true if ln K is negative?
39) ______ A) Δ
G°
rxn
is negative and the reaction is spontaneous in the forward direction under
standard conditions.
B) Δ
G°
rxn
is positive and the reaction is spontaneous in the forward direction under
standard conditions.
C) Δ
G°
rxn
is zero and the reaction is at equilibrium under standard conditions.
D) Δ
G°
rxn
is negative and the reaction is spontaneous in the reverse direction under
standard conditions.
E) Δ
G°
rxn
is positive and the reaction is spontaneous in the reverse direction under
standard conditions.
40) Consider a reaction that has a positive
Δ
H and a positive
Δ
S.
Which of the following
statements is TRUE?
40) ______ A) This reaction will be spontaneous only at high temperatures.
B) This reaction will be nonspontaneous only at high temperatures.
C) This reaction will be nonspontaneous at all temperatures.
D) This reaction will be spontaneous at all temperatures.
E) It is not possible to determine without more information.
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