Chemistry 400: General Chemistry
Miller Fall 2015
Exam III
December 9, 2015
Name: _____________________________________
Approximately 150 points
Please answer each of the following questions to the best of your ability. If you wish to receive partial credit, please show your work.
For multiple choice, there is no partial credit (unless otherwise noted) and there is only one correct answer. For multiple choice,
please clearly mark the correct answer. For all ionic species, please show the charge on each ion to receive full credit.
Element
F
O
Cl
N
S
Br
C
H
Electronegativity
4.0
3.5
3.0
3.0
2.8
2.8
2.5
2.1
I. Multiple choice are 4 points each.
1. Atoms having equal electronegativities are expected to form
a. no bonds
b. polar covalent bonds
c. nonpolar covalent bonds
d. ionic bonds
e.
2. In what may have been the greatest lecture ever given on the day before Thanksgiving (LOL), I mentioned that intermolecular
forces are important to understand. Which of the following correctly describes the influence of intermolecular forces on the properties
of matter?
1. Intermolecular forces determine the solubility of gases, liquids, and solids in various solvents.
2. Intermolecular forces are directly related to the energy required to accomplish phase changes.
3. Intermolecular forces influence the shapes of molecules like DNA and proteins.
a.
1 only
b.
2 only
c.
3 only
d.
1 and 2
3. Hydrogen bonding is present in all of the following molecular solids EXCEPT ____.
a. H2SO4
b. CH3OH
c. CH3OCH3
d. HF
e. CH3CO2H
4. Equilibrium is established between a liquid and its vapor when
a. the rate of evaporation equals the rate of condensation.
b. equal masses exist in the liquid and gas phases.
c. equal concentrations (in molarity) exist in the liquid and gas phases.
d. all the liquid has evaporated.
e. the liquid ceases to evaporate and the gas ceases to condense.
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e.
1, 2 and 3
5. Three possible structures of C2H2Cl2 are shown below. Which of these molecules are polar?
a.
1 only
b.
2 only
c.
3 only
d.
1 and 3
e.
2 and 3
II. Drawing Section
1. Draw the orbital overlap diagram for acetylene C2H2. Your diagram should include the type of each atomic orbital (s, p, d, sp2, etc.,)
and whether the overlap forms a sigma or pi bond. (10 points)
2
2. For each of the following chemical species:
A. NH3
(i) Draw the best Lewis structure. Label all
atoms with nonzero formal charge.
(4 points)
(ii) Draw the correct electron geometry
and (iii) Draw dipoles. (6 points)
(iv) What is the hybridization on N?
(2 points)
(v) What is the bond angle (including the
correction for electron pairs, if necessary)?
(2 points)
B. NO3–
(i) Draw the best Lewis structure. Label all
atoms with nonzero formal charge.
(4 points)
(ii) Draw the correct electron geometry
and (iii) Draw dipoles. (6 points)
(iv) Draw any resonance structures (if
there are any). (4 points)
3. Draw all of the hydrogen bonds between one molecule of NH3 and as many other water molecules as necessary. (6 points)
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III. Free Response
1. For each of the following pairs of molecules, list the dominant IMF for each molecule (2 points each) and circle the one with the
higher boiling point (1 point each, only awarded if your answer is correct and both dominant IMFs are correct):
A.
H
B. HCN
C
C. NaCl
NaBr
D. HF
F2
C
H
2. Four non-equivalent Lewis structures for SCO are given below.
S
C
A
O
S
C
O
S
C
B
C
O
S
C
O
D
A. One of these structures is clearly worst. Choose one and explain why it is the worst. (6 points)
B. One of these structures is clearly the best. Choose one and explain why it is the best. (6 points)
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3. Consider O2(g).
A. Draw the molecular orbital energy diagram for O2 (showing σ2s, σ*2s, etc.,). (10 points)
B. What is the bond order for O2? (4 points)
C. Is O2 paramagnetic or diamagnetic? (4 points)
D. According to molecular orbital theory, which should be more stable: O2 or O22–? Why? (6 points)
E. Calculate the bond energy for making the bond in O2 and making the bond in O22- out of oxygen atoms based on their
different bond dissociation energies. Based on these bond making energies, which bond is more stable? Why? (10 points)
2 O(g) + 2 e– → O22=(g)
(ignore e in using bond dissociation energies)
2 O(g) → O2(g)
-
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4A. Draw a heating curve for water. Label the axes, phases, and phase transitions. (10 points)
B. Tell me as much as possible about the final state of the mixture when three 25.0 gram ice cubes at –10.0°C are placed in a glass
with 50.0 mL of water at 25.0°C. Does the ice dissolve completely? If partially, how much ice is left? What is the final temperature?.
(10 points)
6
Chemistry 400
Conversions and Equations
1 L = 1.057 qt
453.6 g = 1 lb
ºF = 1.8 º C + 32
1 m = 39.37 in
1 yd = 36 in = 3 ft
Na = 6.02 x 1023
1 gal = 4 qt
1 lb = 16 oz
1 calorie = 4.184 J
specific heat of water = 4.184 J/g °C
q = Energy = (mass) (Csp) (ΔT)
average atomic mass =
PO2 = (% O2) PT
λ=
€
E = hν
2
% yield = actual/theoretical × 100%
PT = P1 + P2 + P3 + …
λ=
$ 1 1'
E = −2.18 ×10 J Z && 2 − 2 ))
€
% n2 n1€
(
−18
€
(mass isotope 1)(%) + (mass isotope 2)(%) + (mass isotope 3)(%)
100%
P1 V1 / T1 = P2 V2 / T2
c
ν
1 qt = 4 cups
1 mile = 5280 feet
1 atm = 101.3 kPa = 1.013 bar = 14.7 psi
For gases: standard T = 273.15 K, P = 1 atm
h
KE = ½ mv2
mv
h = 6.626×10–34 J•s
c = 3.00×108 m/s€
1 L•atm = 101.325 J
PV=nRT
ΔxmΔv =
w = -P ΔV
h
4π
Mass of electron =
9.1×10–31 kg
Kw = [H3O+] [OH–] = 1.0×10–14
[H+] [OH–] = 1.0×10–14
C 1V 1 = C 2V 2
R=0.08206 L•atm/mol•K = 8.314 J/mol•K
Specific heat of ice: 2.09 J/g•°C
Specific heat of water: 4.184 J/g•°C
Specific heat of steam: 2.03 J/g•°C
Heat of fusion of H2O = Δ Hfus = 6.02 kJ/mol
Heat of vaporization of H2O = Δ Hvap = 40.7 kJ/mol
€
ΔT = m i K
π =iMRT
€
€
" P % −ΔHvap " 1 1 %
ln$ 2 ' =
$ − '
R # T2 T1 &
# P1 &
€
€
1
1 ppm = 1 x 10-6 g/mL
€
1 ppb = 1 x 10-9 g/mL
Chemistry 400
Conversions and Equations
H
436
C
414
347
Average Single Bond Dissociation Energy (in kJ/mol)
N
O
F
Si
P
S
Cl
Br
I
389 464 565 323 322 368 431 364 297
305 360 485 301
272 339 276 213
163 222 272
200 243 159
142 190 452 335
203 201 234
159 565 490 327 253 237 278
226
293 464 310 234
201
326
184
266 253 218
243 216 208
193 175
151
H
C
N
O
F
Si
P
S
Cl
Br
I
Comparison of Average Single, Double and Triple Bond Energies (in kJ/mol)
Bond Type Single Bond Double Bond Triple Bond
C–C
347
611
837
N–N
163
418
946
O–O
142
498
C–N
305
615
891
C–O
360
736*
1072
C–Cl
339
651
*For CO2, the C=O bond is 799 kJ/mol
Boiling Point Elevation and Freezing Point Depression
Constants
Solvent
Formula
Kb (°C/m) Kf (°C/m)
Water
H 2O
0.512
–1.86
Ethanol
CH3CH2OH
1.22
–1.99
Chloroform
CHCl3
3.63
–4.70
Benzene
C 6H 6
2.53
–5.12
Diethyl ether CH3CH2OCH2CH3
2.02
–1.79
Element
F
O
Cl
N
S
Br
C
H
2
Electronegativity
4.0
3.5
3.0
3.0
2.8
2.8
2.5
2.1
Chemistry 400
Conversions and Equations
Material
Ag(s)
Ag+(aq)
Al(s)€
Al3+(aq)
Al2O3(s)
AlCl3(aq)
AlCl3(s)
Br(g)
Br2(g)
Br2(l)
C(g)
C(s, dia)
C(s, gr)
C2H4(g)
C2H4O(g)
C2H5OH(l)
C6H12O6(s)
C3H8(g)
CH3CH2CH2CH3(l)
Ca(g)
Ca(OH)2(aq)
Ca(OH)2(s)
Ca(s)
Ca2+(aq)
Ca2+(g)
CaCl2(s)
CaCO3(s)
CaF2(s)
CaO(s)
CH3OH(g)
CH4(g)
CHCl3(l)
Cl–(aq)
Cl(g)
Cl2(g)
CO(g)
CO2(g)
Cu(s)
Cu2+(aq)
Fe2O3(s)
Fe3O4(s)
H(g)
H+(aq)
H2(g)
H2O(g)
H2O(l)
H2O(s)
H2O2(aq)
Δ Hf°(kJ/mol)
0
105.79
0
–538.4
–1675.7
–1039.7
–704.2
111.9
30.9
0
716.7
1.88
0
52.4
–166.2
–277.6
–1273.3
–107.85
–147.3
177.8
–1003
–985.2
0
–542.8
1934.1
–795.4
–1207.6
–1228.0
–634.9
–201.0
–74.6
–134
–167.1
121.3
0
–110.5
–393.5
0
64.9
–824
-1118
218
0
0
–241.8
–285.8
–291.8
–191.2
3
Material
H2O2(l)
H3O+(aq)
HBr(g)
€
HCl(aq)
HCl(g)
H2SO4(l)
I2(g)
I2(s)
Mg2+(aq)
MgCl2(aq)
N(g)
N2(g)
N2H4(l)
N2O(g)
N2O4(g)
Na(s)
Na+(aq)
Na2SO4(s)
NaCl(aq)
NaCl(s)
NaOH(aq)
NH3(aq)
NH3(g)
NH4+(aq)
NH4Cl(aq)
NH4Cl(s)
NH4NO3(aq)
NH4NO3(s)
NI3(s)
NO(g)
NO2(g)
O(g)
O2(g)
O3(g)
OH–(aq)
SO2Cl2(g)
SO2(g)
SO42–(aq)
Zn(s)
Zn2+(aq)
HgO(s)
Hg(l)
Δ Hf°(kJ/mol)
–187.78
–285.8
–36.3
–167.2
–92.3
–814
62.42
0
–467.0
–801.2
472.7
0
50.6
81.6
11.1
0
–240.34
–1387.1
–407.2
–411.2
–470.1
–80.29
–45.9
–133.26
–299.66
–314.43
–339.9
–365.6
192
91.3
33.2
249.2
0
142.7
–230.02
–364
–296.8
–909.3
0
–153.39
–90.8
0
Chemistry 400
Conversions and Equations
Unit of Concentration
Symbol
Formula
Mass/mass percent
% (w/w)
%(m/m) = grams of solute x 100
grams of solution
Mass/volume percent
% (w/v)
% (m/v) = grams of solute x 100
mL of solution
Volume/volume percent
% (v/v)
% (v/v) =
Molarity
M
M = moles of solute
L of solution
parts per million
ppm
ppm = grams of solute x 106
grams of solution
mL of solute x 100
mL of solution
4