Download 1 Chemistry 400: General Chemistry Name: Miller Fall 2015 Final

Chemistry 400: General Chemistry
Miller Fall 2015
Final Exam Part Deux
December 14, 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 all ionic species, please show the
charge on each ion to receive full credit. Good luck!
All Lewis Structures must include all valence electrons.
Element Electronegativity
F
4.0
O
3.5
Cl
3.0
N
3.0
S
2.8
Br
2.8
C
2.5
H
2.1
I. Nomenclature
(4 points each, spelling counts) If the name is given, please give the formula. If the formula is given, please give the name.
A. NaNO2
B. CH3COOH
C. Fe2S3
D. N2O3
E. CuSO4•5H2O
II. Drawing
1. For CF4:
(i) Draw the correct Lewis structure.
(4 points)
(ii) Draw the correct electron geometry
and (iii) Draw dipoles. (6 points)
(iv) What is the bond angle around the
oxygen atom? Use a <, >, or = symbol as
part of your answer. (2 points)
(v) Is this molecule polar or nonpolar?
(4 points)
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2. If MgSO4·xH2O, is heated to 250°C, all the waters of hydration evaporate.
On heating a 0.4000 g sample of the hydrate MgSO4·xH2O, 0.1953 g of MgSO4 remains. How many molecules of water occur per
formula unit of MgSO4? In other words, solve for x.
For this problem, your show your work if you'd like to receive any credit. (10 points)
3. Will the pH of 0.100 M HCl or 0.100 M CH3COOH be lower? Explain your choice with specific reference to the types of acids
involved and their percents of ionization. (8 points)
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4. There is a relationship between intermolecular forces, temperature, and the temperatures at which the solid, liquid, and gas
phases exist. Each blank will be filled in with one of the words in bold. You may use each word more than once. (1 point per blank)
A. ______________________ is proportional to the average kinetic energy.
B. At low temperatures, materials are in the ______________________ phase. In this phase, ______________________ are
strong enough to keep the particles in the material (atoms, ions, or molecules) close enough such that the particles have 3dimensional order.
C. At high temperatures, materials are in the ______________________ phase. In this phase, no matter how strong the
______________________ are, the energy supplied by the high temperatures is enough to overcome the
______________________ .
D. At intermediate temperatures, materials are in the ______________________ phase. The exact temperatures of the phase
transitions depend upon the strength of ______________________ .
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5. Calculate the change in the enthalpy of reaction, ΔHrxn, for the following combustion reaction two ways:
C2H6(g) + 3.5 O2(g) → 2 CO2(g) + 3 H2O(g)
A. Using bond dissociation energies. (14 points)
B. Using standard changes in the enthalpy of formation, ΔHf°. (6 points)
ΔHf°(C2H6) = -84 kJ/mol
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6. In the experiment below, the electron shooter shoots one electron at a time through the double slits at the screen. Draw the pattern
that the electrons make on the screen after a long period of time in which many electrons have been shot at the screen but the electrons
have been shot at the screen one at a time.
A. No laser (no observer) to determine which slit the electron goes through (4 points)
electron shooter: shoots
one electron at a time
double slit
screen
B. A laser (an observer) is present to determine which slit the electron goes through (4 points)
electron shooter: shoots
one electron at a time
double slit
screen
7. Define an orbital. Make sure to use the words volume and probability in your definition (6 points)
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8. Write a paragraph about the relationship between the 3d and 4s orbitals in terms of energy. For what elements is 3d lower in
energy? For what elements is 4s lower in energy? To answer this question, use the electron configurations (with or without noble gas
configurations) of a potassium atom, a potassium ion, a titanium atom, and a doubly charged titanium ion. (12 points)
electron configurations:
K atom
K+ ion
Ti atom
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Ti2+ ion
9. A. Draw the best Lewis Structure for NH3 and the best Lewis structure for ClH3, named chlorine trihydride. All atoms will have
zero formal charge. For each Lewis structure, calculate the difference in electronegativity for each type of bond and draw dipoles for
each bond (if they exist). (8 points)
B. Determine the dominant type of intermolecular force for each molecules above and suggest a reason why they are
different (even though the differences in electronegativity of each type of bond might suggest otherwise). (8 points)
10. The nonvolatile, nonelectrolyte aspirin, C9H8O4, is soluble in diethyl ether CH3CH2OCH2CH3. How many grams of aspirin are
needed to generate an osmotic pressure of 7.72 atm when dissolved in 190 ml of a diethyl ether solution at 298 K. (8 points)
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11. Shown below is the distribution of velocities for hydrogen (H) atoms at 300 K. Place a line on this graph that approximately
represents the distribution of velocities for helium (He) atoms at 300 K. (4 points)
Fraction of Atoms x 10^5 100 90 80 70 60 50 40 30 20 10 0 0 1000 2000 3000 4000 5000 Velocity (m/s) 8
6000 12. From calculations in class, it was clear that the wavelength of an electron was approximately the same size as an atom of
hydrogen, r = 0.05 nm.
A. Calculate the wavelength of an electron traveling at 2.2 x 106 m/s. (6 points)
B. Are there any other particles or atoms for which the wavelength is comparable to the size of the atom? If so, please show
through calculations. If not, then please explain why. (10 points)
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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
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