Leader: Grady
Course: Chem 167
Instructor: Houk
Final Review
Supplemental Instruction
Iowa State University
Date: Week of 4/26-5/1
1. Calculate the following to the correct number of significant figures. Assume that all
these numbers are measurements.
(a) 13.0217/17.10
√(0.0034)2 +4(1.000)(6.3𝑥10−4 )
(b) .0034+
2𝑥1.000
2. The “Western Stone” in Jerusalem is one of the largest stone building blocks ever to
have been used. It has a mass of 517 metric tons, and measures 13.6 m long, 3.00 m
high, and 3.30 m wide. What is the density of this rock in g/cm3? (1 metric ton = 1000
kg)
3. Rubidium has two naturally occurring isotopes, 85Rb (relative mass 84.9118 amu) and
87
Rb (relative mass 86.9092 amu). If rubidium has an average atomic mass of 85.47
amu, what is the abundance of each isotope (in percent)?
4. The elemental analysis of methanol is 37.5% C, 12.6% H, and 49.9% O. What is the
empirical formula of methanol?
5. Name/give the formula the following substances:
a. Cu(NO3)2
b. N2O7
c. Carbon disulfide
d. PCl5
e. Dibromine pentoxide
6. Give the molecular formula for the following substance:
7. Describe covalent bonds and ionic bonds.
8. Balance the following equation and then write the net ionic equation.
a. Zns) + HCl(aq)  H2(g) + ZnCl2(aq)
9. When barium nitrate and sodium phosphate are mixed together in water a precipitate
is formed. What is the most likely chemical formula of the precipitate?
10. Which of the following is a weak electrolyte? (Weak acid/base)
a. HCl
b. CH3COOH
c. NaOH
d. HNO3
11. If 21 g of H2S is mixed with 38 g of O2 and 31 g of SO2 form, what is the percentage
yield? 2H2S + 3O2  2SO2 + 2H2O
12. What is the final molarity of sodium ions if 50.0 mL of 0.25 M Na2SO4 is combined
with 50.0 mL of 1.00 M NaCl? Assume the total solution volume is exactly 100.0
mL.
13. Small quantities of hydrogen gas can be prepared by the following reaction:
Zn(s) + H2SO4(aq)  ZnSO4(aq) + H2(g) How many grams of H2 can be prepared
from 25.0 mL of 6.00M H2SO4 and excess zinc?
14. What volume of oxygen at 24 oC and 0.88 atm is needed to completely react via
combustion with 45 g of methane gas?
15. Which of the following is not a postulate of the kinetic molecular theory of gases?
a. Gas molecules are in constant random motion
b. Gas molecules are attracted to each other
c. Gas molecules are infinitely small
d. Gas molecules have elastic collisions
16. Which of the following conditions might we expect a gas not to behave like an ideal
gas?
a. High temperature
b. Low temperature
c. High pressure
d. Both (a) and (c)
e. All (a), (b), and (c)
17. Find the wavelength of light that would be emitted when a hydrogen atom undergoes
the transition from the state n=4 to the state n=2. Express your answer in nm. (hint:
the wavelength for the transition 41 is 97.534 nm, 21 121.556 nm, and 31
1102.583 nm)
18. Which of the following sets of quantum numbers is not possible?
a. n = 3, l =2, ml = -1, ms = +1/2
b. n = 2, l = 0, ml = 0, ms = +1/2
c. n = 4, l = 3, ml = 4, ms = -1/2
d. n = 1, l = 0, ml = 0, ms = +1/2
19. Using the shortcut method, give the electron configuration and orbital drawing for the
following:
a. S
b. Ni
c. Mg2+
d. O2-
20. Describe the trends for ionization energy, electron affinity, and atomic radius.
21. Give the correct order for atomic radius for Mg, Na, P, Si, and Ar.
22. What is the frequency of yellow light with wavelength of 595 nm?
23. Draw the Lewis structures for the following molecules. Determine the shapes, bond
angles, hybridization of the central atom, polarity of the molecule, and how many
sigma and pi bonds are present.
a. FNO
b. HCO3-
c. ICl5
d. H2NCN
e. TeCl4
f. C2F4
g. SNF3
24. Barium crystallizes with an edge of 5.025x10-8 cm in a body-centered cubic unit cell.
What is the density of barium?
25. Explain using a diagram how doping will increase the conductivity of a
semiconductor. Arsenic would be which type of doping for germanium?
26. Identify the kinds of intermolecular forces (London dispersion, dipole-dipole, or
hydrogen bonding) that are the most important in each of the following substances.
a. Methane, CH4
b. Methanol, CH3OH
c. Chloroform, CHCl3
d. Benzene, C6H6
e. Ammonia, NH3
f. Sulfur dioxide, SO2
27. Referring to the previous question, which molecule has the highest boiling point,
highest vapor pressure, and lowest surface tension?
28. Which of the following is an example of an alternating copolymer?
a. ABABABABABABABABAB
b. AABAAABBBAAABAABAB
c. AAABBBAAABBBAAABBB
d. BBAAAAAAABBBABBBAA
29. Describe the process that creates a polymer (describe the three steps). Also, give the
typical products of a polymer condensation reaction.
30. A 1.000 g sample of octane (C8H18) is burned in a calorimeter containing 1200 g of
water at an initial temperature of 25.00 oC. After the reaction, the final temperature
of the water is 33.20 oC. The heat capacity of the calorimeter is 837 J/oC. The specific
heat of water is 4.184 J/g oC. Calculate the heat of combustion of octane in kJ/mol.
31. Calculate the heat of combustion of diborane.
Δ𝐻𝑓𝑜
B2H6(g)
H2O(g)
B2O3(s)
36.0 kJ/mol
-241.8 kJ/mol
-1273.4 kJ/mol
𝑜
32. Calculate Δ𝐻𝑟𝑥𝑛
for: 2H2O(l)  2H2(g) + O2(g)
H2O(l)  H2O(g)
2H2(g) + O2(g)  2H2O(g)
Δ𝐻 𝑜 = 44.01 𝑘𝐽
Δ𝐻 𝑜 = −483.64 𝑘𝐽
33. Given the thermochemical equation:
N2(g) + O2(g)  2NO(g)
Δ𝐻 = 180.6 𝑘𝐽
If 558 kJ of energy are supplied, what mass of NO can be made?
34. Describe the meaning of positive/negative values for enthalpy, entropy, and Gibb’s
free energy.
35. When 0.0157 kg of a compound with a heat of combustion of -37.6 kJ/mol is burned
in a calorimeter, 18.5 kJ of heat are released. What is the molar mass of the
compound?
What is the rate law expression for this reaction?
Determine the value of the rate constant k in the correct units.
37. The decomposition of N2O5 in solution is a first-order reaction:
2 N2O5  4NO2 + O2
The rate constant at a give temperature is found to be 5.25𝑥10−4 𝑠 −1 If the initial
concentration of N2O5 is 0.200 M, what is its concentration after 10 minutes have
passed?
38. Determine the rate law (including the rate constant) using the data provided.
Q + X  products
[Q] (M)
0.12
0.24
0.12
[X] (M)
0.10
0.10
0.20
Rate = Δ[products]/ Δt (M/min)
1.5 x 10-3
3.0 x 10-3
1.2 x 10-2
39. The activation energy for the reaction in which CO2 decomposes to CO and a free radical
oxygen atom, O, has an activation energy of 460 kJ/mol. The frequency factor is
2𝑥1011 𝑠 −1 What is the rate constant for this reaction at 298 K?
40. The reaction of NO2(g) and CO(g) is thought to occur in two steps:
Step 1: NO2(g) + NO2(g)  NO(g) + NO3(g)
(slow)
Step 2: NO3(g) + CO(g)  NO2(g) + CO2(g)
(fast)
a. Show that the elementary steps add up to give the overall stoichiometric equation.
b. What is the molecularity of each step?
c. For this mechanism to be consistent with kinetic data, what must be the
experimental rate equation?
d. Identify any intermediates in this reaction.
41. A system of 0.100 moles of oxygen gas is placed in a closed 1 L container and is brought
to equilibrium at 600 K.
O2(g) <--> 2O(g)
𝐾 = 2.8𝑥10−39
What are the equilibrium concentrations of O and O2?
42. The pH of a 0.25 M HCOOH formic acid solution is 2.17 at 25 oC. Calculate the Ka for
the acid.
HCOOH(aq) <--> H+(aq) + HCOO-(aq)
43. Calculate the value of Kc for the reaction: 2N2O(g) + 3O2(g) <--> 2N2O4(g) using the
following information.
2N2(g) + O2(g) <--> 2N2O(g)
N2O4(g) <--> 2NO2(g)
½ N2(g) + O2(g) < NO2(g)
Kc = 1.2𝑥10−35
Kc = 4.6𝑥10−3
Kc = 4.1𝑥10−9