ODD & WEIRD: Descriptive, Coordination, Nuclear, and Organic
2002
(A) PbSO4
(B) CuO
6.
Is purple in aqueous solution
7.
Is white and very soluble in water
(C) KMnO4
(D) KCl
(E) FeCl3
(A) H2SeO4(aq) + 2 Cl–(aq) + 2 H+(aq)  H2SeO3(aq) + Cl2(g) + H2O(1)
(B) S8(s) + 8 O2(g)  8 SO2(g)
(C) 3 Br2(aq) + 6 OH–(aq)  5 Br–(aq) + BrO3–(aq) + 3 H2O(l)
(D) Ca2+(aq) + SO42–(aq)  CaSO4(s)
(E) PtCl4(s) + 2 Cl–(aq)  PtCl62–(aq)
11.
12.
13.
14.
A precipitation reaction
A reaction that produces a coordination complex
A reaction in which the same reactant undergoes both oxidation and reduction
A combustion reaction
235
92 U
1
+0n 
141
55
1
Cs + 3 0 n + X
23. Neutron bombardment of uranium can induce the reaction represented above. Nuclide X is which of the
following?
(A)
92
35
Br
(B)
94
35
Br
(C)
91
37
Rb
(D)
92
37
Rb
94
(E) 37 Rb
29. The best explanation for the fact that diamond is extremely hard is that diamond crystals
(A) are made up of atoms that are intrinsically hard because of their electronic structures
(B) consist of positive and negative ions that are strongly attracted to each other
(C) are giant molecules in which each atom forms strong covalent bonds with all of its neighboring
atoms
(D) are formed under extreme conditions of temperature and pressure
(E) contain orbitals or bands of delocalized electrons that belong not to single atoms but to each crystal
as a whole
32. Which of the following oxides is a gas at 25˚C and 1 atm?
(A) Rb2O
(B) N2O
(C) Na2O2
(D) SiO2
59. All of the halogens in their elemental form at 25˚C and 1 atm are
(A) conductors of electricity
(B) diatomic molecules
(C) odorless
(D) colorless
(E) gases
(E) La2O3
H H H
| | |
H—C—C—C—Br
| | |
H H H
62. Which of the following structural formulas represents an isomer of the compound that has the structural
formula represented above?
H H H
H H H
|
| |
| | |
(A) Br—C—C—C—H
(B) H—C—C—C—H
|
| |
| | |
H H H
H Br H
H H H
| | |
(C) H—C—C—C—H
| | |
H H Br
H H H
|
| |
(D) Br—C—C—C—Br
|
| |
H H H
H H H H
| | | |
(E) H—C—C—C—C—Br
| | | |
H H H H
65. Which of the following substances is LEAST soluble in water?
(A) (NH4)2SO4
(B) KMnO4
(C) BaCO3
(D) Zn(NO3)2
(E) Na3PO4
72. A colorless solution is divided into three samples. The following tests were performed on samples of the
solution.
Sample
Test
Observation
+
1
Add H (aq)
No change
2
Add NH3(aq)
No change
3
Add SO42–(aq)
No change
Which of the following ions could be present in the solution at a concentration of 0.10 M?
(A) Ni2+(aq)
(B) Al3+(aq)
(C) Ba2+(aq)
(D) Na+(aq)
(E) CO32–(aq)
ANSWERS: 6 C,7 D,11-14 D/E/C/B, 23 D,29 C,32 B,59 B,62 B,65 C,72 D
2007 part A, form B, question #2
Answer the following problems about gases.
(a) The average atomic mass of naturally occurring neon is 20.18 amu. There are two common isotopes of
naturally occurring neon as indicated in the table below.
Isotope
Mass (amu)
Ne-20
19.99
Ne-22
21.99
(i) Using the information above, calculate the percent abundance of each isotope.
(ii) Calculate the number of Ne-22 atoms in a 12.55 g sample of naturally occurring neon.
(b) A major line in the emission spectrum of neon corresponds to a frequency of 4.341014 s-1. Calculate the
wavelength, in nanometers, of light that corresponds to this line.
(c) In the upper atmosphere, ozone molecules decompose as they absorb ultraviolet (UV) radiation, as shown by
the equation below. Ozone serves to block harmful ultraviolet radiation that comes from the Sun.
O3(g) UV
 O2(g) + O(g)
A molecule of O3(g) absorbs a photon with a frequency of 1.001015 s-1.
(i)
How much energy, in joules, does the O3(g) molecule absorb per photon?
(ii) The minimum energy needed to break an oxygen-oxygen bond in ozone is 387 kJ mol-1. Does a photon with
a frequency of 1.001015 s-1 have enough energy to break this bond? Support your answer with a calculation.
1989 D
The carbon isotope of mass 12 is stable. The carbon isotopes of mass 11 and mass 14 are unstable. However, the type
of radioactivity decay is different for these two isotopes. Carbon-12 is not produced in either case.
(a) Identify a type of decay expected for carbon-11 and write the balanced nuclear reaction for that decay process.
(b) Identify the type of decay expected for carbon-14 and write the balanced nuclear reaction for that decay process.
(c) Gamma rays are observed during the radioactive decay of carbon-11. Why is it unnecessary to include the gamma
rays in the radioactive decay equation of (a)?
(d) Explain how the amount of carbon-14 in a piece of wood can be used to determine when the tree died.
1991 D
Explain each of the following in terms of nuclear models.
(a) The mass of an atom of 4He is less than the sum of the masses of 2 protons, 2 neutrons, and 2 electrons.
(b) Alpha radiation penetrates a much shorter distance into a piece of material than does beta radiation of the same
energy.
(c) Products from a nuclear fission of a uranium atom such as 90Sr and 137Ce are highly radioactive and decay by
emission of beta particles.
(d) Nuclear fusion requires large amounts of energy and to get started, whereas nuclear fission can occur
spontaneously, although both processes release energy.
1997 D
Answer each of the following questions regarding radioactivity.
(a) Write the nuclear equation for decay of 234
94 Pu by alpha emission.
(b) Account for the fact that the total mass of the products of the reaction in part (a) is slightly less than that of the
original 234
94 Pu .
(c) Describe how , , and  rays each behave when they pass through an electric field. Use the diagram below to
illustrate your answer.
(d) Why is it not possible to eliminate the hazard of nuclear waste by the process of incineration?
2002 B
Consider the hydrocarbon pentane, C5H12 (molar mass 72.15 g).
(a) Write the balanced equation for the combustion of pentane to yield carbon dioxide and water.
(b) What volume of dry carbon dioxide, measured at 25˚C and 785 mm Hg, will result from the complete combustion
of 2.50 g of pentane?
(c) The complete combustion of 5.00 g of pentane releases 243 kJ of heat. On the basis of this information, calculate
the value of ∆H for the complete combustion of one mole of pentane.
(d) Under identical conditions, a sample of an unknown gas effuses into a vacuum at twice the rate that a sample of
pentane gas effuses. Calculate the molar mass of the unknown gas.
(e) The structural formula of one isomer of pentane is shown below. Draw the structural formulas for the other two
isomers of pentane. Be sure to include all atoms of hydrogen and carbon in your structures.
2003 D
Compound
Name
Compound
Formula
∆H˚vap
(kJ mol-1)
Propane
CH3CH2CH3
19.0
Propanone
CH3COCH3
32.0
1-propanol
CH3CH2CH2OH
47.3
Using the information in the table above, answer the following questions about organic compounds.
(a) For propanone,
(i)
draw the complete structural formula (showing all atoms and bonds);
(ii) predict the approximate carbon-to-carbon-to-carbon bond angle.
(b) For each pair of compounds below, explain why they do not have the same value for their standard heat of
vaporization, ∆H˚vap. (You must include specific information about both compounds in each pair.)
(i)
Propane and propanone
(ii) Propanone and 1-propanol
(c) Draw the complete structural formula for an isomer of the molecule you drew in, part (a) (i).
(d) Given the structural formula for propyne below,
H
| 
H—C—CC—H
|
H
(i)
indicate the hybridization of the carbon atom indicated by the arrow in the structure above;
(ii) indicate the total number of sigma () bands and the total number of pi (π) bonds in the molecule
2006 B
Answer the following questions that relate to the analysis of chemical compounds.
(a) A compound containing the elements C, H, N, and O is analyzed. When a 1.2359 g sample is burned in excess
oxygen, 2.241 g of CO2(g) is formed. The combustion analysis also showed that the sample contained 0.0648 g of
H.
(i)
Determine the mass, in grams, of C in the 1.2359 g sample of the compound.
(ii) When the compound is analyzed for N content only, the mass percent of N is found to be 28.84 percent.
Determine the mass, in grams, of N in the original 1.2359 g sample of the compound.
(iii) Determine the mass, in grams, of O in the original 1.2359 g sample of the compound.
(iv) Determine the empirical formula of the compound.
(b) A different compound, which has the empirical formula CH2Br, has a vapor density of 6.00 g L-1 at 375 K and
0.983 atm. Using these data, determine the following.
(i) The molar mass of the compound
(ii) The molecular formula of the compound
1988 D
12
X
10
X
8
pH
X
X
6
X
X
X
4X
2
0
0
5
10
15
20
millilitres of NaOH
25
30
A 30.00 millilitre sample of a weak monoprotic acid was titrated with a standardized solution of NaOH. A pH meter
was used to measure the pH after each increment of NaOH was added, and the curve above was constructed.
(a) Explain how this curve could be used to determine the molarity of the acid.
(b) Explain how this curve could be used to determine the dissociation constant Ka of the weak monoprotic acid.
(c) If you were to repeat the titration using a indicator in the acid to signal the endpoint, which of the following
indicators should you select? Give the reason for your choice.
Methyl red
Ka = 110–5
Cresol red
Ka = 110–8
Alizarin yellow
Ka = 110–11
(d) Sketch the titration curve that would result if the weak monoprotic acid were replaced by a strong monoprotic
acid, such as HCl of the same molarity. Identify differences between this titration curve and the curve shown
above.
1998 D (Required)
An approximately 0.1–molar solution of NaOH is to be standardized by titration. Assume that the following materials
are available.
• Clean, dry 50 mL buret
• 250 mL Erlenmeyer flask
•
•
•
•
Wash bottle filled with distilled water
Analytical balance
Phenolphthalein indicator solution
Potassium hydrogen phthalate, KHP, a pure solid monoprotic acid (to be used as the primary standard)
(a) Briefly describe the steps you would take, using the materials listed above, to standardize the NaOH solution.
(b) Describe (i.e., set up) the calculations necessary to determine the concentration of the NaOH solution.
(c) After the NaOH solution has been standardized, it is used to titrate a weak monoprotic acid, HX. The equivalence
point is reached when 25.0 mL of NaOH solution has been added. In the space provided at the right, sketch the
titration curve, showing the pH changes that occur as the volume of NaOH solution added increases from 0 to
35.0 mL. Clearly label the equivalence point on the curve.
(d) Describe how the value of the acid–dissociation constant, Ka, for the weak acid HX could be determined from the
titration curve in part (c).
(e) The graph below shows the results obtained by titrating a different weak acid, H2Y, with the standardized NaOH
solution. Identify the negative ion that is present in the highest concentration at the point in the titration
represented by the letter A on the curve.