U.S. National Chemistry Olympiad: 1989 National
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1. Gallium reacts with element Y to form a compound with the formula GaY. If barium
were to combine with element Y, the formula of the product would most likely be
2. The number of significant figures one can assume in the atomic weight of 12C
3. A compound X2O3 contains 31.58% oxygen by weight. The atomic
weight of X is
(A) 34.66 g/mol.
(B) 45.01 g/mol.
(C) 52.00 g/mol.
(D) 104.0 g/mol.
4. The emission spectrum of the hydrogen atom
(A) is caused by the removal (ionization) of the electron.
(B) is continuous because the electron can emit any frequency of light during a
(C) is caused by the absorption of light at characteristic frequencies the electron to be
excited into higher energy levels.
(D) is a result of the excited electron undergoing transitions to lower energy levels and
emitting photons of light at specific frequencies.
5. Cathode rays are
(A) a beam of positively charged particles.
(B) nuclei of helium atoms.
(C) fast-moving neutrons.
(D) streams of electrons.
6. Which atom has the correct ground state electron configuration?
(A) Cl: [Ne]3s1 3p6
(B) Mo: [Kr]5s1 4d5
(C) Cu: [Ar]4s2 3d8
(D) As: [Ar]4s2 4d10 4p3
7. The correct number of unpaired electrons in the ground state of a neutral cobalt
8. Which electron configuration describes a neutral atom in an excited state?
(A) [Xe] 6s 4f 5d 6p
(C) [Ne] 3s 3p
9. The first ionization potential for S is lower than the first ionization potential
for P because
(A) Hund's rule is violated.
(B) ionization potentials decrease across a representative period.
(C) half-filled and filled sets of orbitals are more stable than configurations of one
more or fewer electrons.
(D) The statement above is false. The first ionization potential for S is higher than that
10. Which set of quantum numbers describe the most easily-removed electron from a
neutral iron atom?
(A) n = 4, l = 0, m = 0, s = -1/2
(B) n = 3, l = 2, m = 0, s = +1/2
(C) n = 3, l = 2, m = -1, s = -1/2
(D) n = 1, l = 0, m = 0, s = -1/2
11. The molecular ion XF2¯ has three pairs of non bonding electrons around
the central atom. The bond angle F-X-F will be closest to
12. Which species has a pyramidal structure?
13. Which molecule is polar?
14. Which of the following sets have correctly matched each molecule or ion and its
15. Which Lewis structure is correct?
16. The Lewis structure for dimethyformamide (a useful solvent) is shown below:
How many C and N atoms are sp3, sp2, and sp hybridized?
(A) 4 (sp3), 0 (sp2), 0 (sp)
(B) 3 (sp3), 0 (sp2), 1 (sp)
(C) 3 (sp3), 1 (sp2), 0 (sp)
(D) 2 (sp3), 2 (sp2), 0 (sp)
17. Consider the reaction 2 Al(s) + 3 Cl2(g) ---> 2 AlCl3(s).
What volume of chlorine at STP reacts with 324 g of aluminum?
(A) 121 L
(B) 134 L
(C) 260 L
(D) 403 L
18. How many liters of H2(g) at 273 °C and 200 torr can be obtained form
the reaction of 500 g of zinc with excess hydrochloric acid?
(A) 5.12 L
(B) 343 L
(C) 650 L
(D) 1300 L
19. What is the concentration of a solution prepared by dissolving 4.20 grams of NaF in
500 grams of water?
20. 10.24 mL of 0.568 M Al(NO3)3 is mixed with 3.12 mL, of 4.16 M
NaOH. How much solid Al(OH)3 could be formed?
(A) 0.055 g
(B) 0.075 g
(C) 0.111 g
(D) 0.166 g
21. The standard heat of formation of ammonia gas is -46.1 kJ mol¯1. How
much heat will be produced when 3.04 gram of N2 reacts with 6.04 g of H2?
(A) 5.00 kJ
(B) 10.00 kJ
(C) 92.1 kJ
(D) 185.3 kJ
22. At 25 °C and 1 atm, the densities of neon and krypton are 0.82 and 3.42
grams/liter, respectively. The density of argon at the same temperature and pressure is
estimated to be
(A) 2.12 g/L
(B) 2.60 g/L
(C) 2.80 g/L
(D) 4.17 g/L
23. When the hydrocarbon propane is burned in air, carbon dioxide and water are formed
as in the equation
C3H8(g) + 5 O2(g) ---> 3 CO2(g)
+ 4 H2O(l)
If 0.100 mol of CO2 is produced, how many drops of water will be formed,
assuming one drop is 0.05 cm3 and contains 1.70 x 1021 water
(A) 2.21 x 10¯1
24. According to the kinetic molecular theory of gases, the Kelvin temperature of a gas
is directly proportional to
(A) the average velocity of the gas molecules.
(B) the square of the mass of the gas molecules.
(C) the average kinetic energy of the gas molecules.
(D) the square root of the average velocity of the gas molecules
25. Which is the most viscous at 20 °C and 1.00 atm?
26. In the van der Waals equation given below,
[P + a (n/V)2] (V - nb) = nRT,
the a(n/V)2 and - nb terms represent, respectively, corrections for
(A) derivations in the pressure and the temperature.
(B) intermolecular attractive forces and inelastic collisions.
(C) intermolecular attractive forces and molecular volumes
(D) intermolecular repulsive forces and high temperatures.
27. The Bragg equation yields
(A) rates of flow of liquids.
(B) interatiomic distances in crystals
(C) trajectories of electron beams
(D) frequencies of spectral lines.
28. The empirical formula for a compound with a cubic close-packed arrangement of
anions and with cations occupying the octahedral sites is
29. The structure of ice is best described as
(A) a solid network of covalently bounded atoms.
(B) cubic-close packed with 74% of the available sites occupied.
(C) an open structure because of elaborate hydrogen bonding
(D) a flickering cluster of molecules in which protons are easily transferred through the
30. The structure of glass is best described as
31. Which solution has the lowest freezing point?
(A) 1.0 molal FeCl3
(B) 1.0 molal HCl
(C) 1.0 molal KCl
(D) 1.0 molal MgCl2
32. An aqueous solution is 1.00 molal in KI. Which change will cause the vapor pressure
of the solution to increase?
(A) addition of water
(B) addition of NaCl
(C) addition of Na2SO4
(D) addition of 1.00 molal KI.
33. Find the boiling point of a solution of 5.00g of naphthalene (C10H8)
in 100 g of benzene. Kb of benzene is 2.53 °C / m; the normal b.p. of
benzene = 80 °C.
(A) 0.99 °C
(B) 79 °C
(C) 81 °C
(D) 85 °C
34. What is the degree of ionization of the acid HB if a 0.120 molal solution of HB
freezes at -0.300 °C? Kf for water = 1.86 °C / m.
35. Given the standard enthalpies of formation : CO (g), -110.5 KJ/mol; CO2(g),
-393.5 KJ/mol. The enthalpy change for the reaction
2CO(g) + 2 CO2 (g) --> 4 C(s) + 3 O2 (g)
(A) -1008 kJ
(B) -566 kJ
(C) -504 kJ
(D) +1008 kJ
36. If DH° of vaporization for a liquid is 10.0 kJ / mol
and DS° of vaporization is 37.6 J / mol-K, which is the
nearest to the normal boiling point of the liquid?
(A) 240 K
(B) 266 K
(C) 300 K
(D) 310 K
37. A reaction has positive values of both DS° and DH°. From this you can deduce that the reaction
(A) must be spontaneous at any temperature.
(B) cannot be spontaneous at any temperature.
(C) will be spontaneous only at low temperatures.
(D) will be spontaneous only at high temperatures.
38. When NH4NO3 dissolves in water, the solution becomes cold.
From this you can deduce that the DS° for the process is
(D) undeterminable from the data.
39. The distribution coefficient, KD, for an organic compound between water
and methylene chloride is 3.40. An aqueous solution of the organic compound contains 0.500
g per 100 mL and is extracted with 50.0 mL of methylene chloride. What percentage of the
organic compound originally in the water is extracted?
40. Which compound is soluble in 8 M NaOH?
41. Which substance can be used to separate Cu2+ from Mg2+ in an
aqueous solution containing Cu(NO3)2 and Mg(NO3)2?
42. Magnesium fluoride is a slightly soluble salt whose solubility product constant is
Ksp = 3.7 x 10¯8. What is the approximate solubility of magnesium
(A) 9.2 x 10¯9 M
(B) 1.2 x 10¯8 M
(C) 1.4 x 10¯4 M
(D) 2.1 x 10¯3 M
43. Al(OH)3 is an insoluble solid whose Ksp = 1.9 x 10¯33.
What is the maximum concentration of OH¯ which can exist in 0.10 M AlCl3
solution without causing Al(OH)3 to precipitate?
(A) 2.7 x 10¯11 M
(B) 1.4 x 10¯10 M
(C) 8.7 x 10¯8 M
(D) 3.1 x 10¯7 M
44. MY and NY3, two nearly insoluble salts, have the same Ksp
values of 6.2 x 10¯13 at room temperature. Which statement would be true in
regard to MY and NY3?
(A) The molar solubilities of MY and NY3 in water are identical.
(B) The molar solubility of MY in water is less than that of NY3.
(C) The salts MY and NY3 are more soluble in 0.5 M KY than in pure water.
(D) The addition of the salt of KY to solution of MY and NY3 will have no
effect on their solubilities.
45. HCN is a weak acid (Ka = 4.0 x 10¯10). Which statement is
NOT true for an aqueous solution of hydrocyanic acid?
(A) The [H+] decreases as solution of HCN are made more dilute.
(B) Its percent dissociation increases as solutions are made more dilute.
(C) In a 0.100 M solution of HCN, the HCN is approximately 6.3% dissociated.
(D) The ionization constant, Ka, varies dramatically over a range of
concentrations of HCN.
46. Given a weak monoprotic acid such as HCN at equilibrium in aqueous solution, the
addition of a strong acid such as HBr to the solution would cause
(A) no change in the concentration of H+ or CN¯.
(B) the concentrations of both HCN and CN¯ concentration to increase.
(C) the HCN concentration to increase and the CN¯ concentration to decrease.
(D) the HCN concentration to decrease and the CN¯ concentration to increase.
47. Which solution has pH value above 7.00?
(A) 0.10 M KCI
(B) 0.10 M NH4NO3
(C) 0.10 M KCN
(D) 0.10 M HI
48. Which of the following would you expect to have the strongest conjugate acid?
49. If weak base has the ionization constant, Kb, Then the value of the
ionization constant, Ka, of its conjugate acid is given by
50. Which mixture is a buffer solution?
(A) 0.10 M HI + 0.10 M KI
(B) 0.10 M KCl + 0.10 M NaCl
(C) 0.10 M NaCN + 0.10 M HCN
(D) 0.10 M NaOH + 0.10 M KOH
51. Ethanol is CH3CH2OH. Which species is formed when ethanol
acts as a Bronsted base?
52. If molten CaCl2 is electrolyzed, the anode reaction is
(A) Ca2+ + 2e¯ ---> Ca
(B) 2Cl¯ ---> Cl2 + 2e¯
(C) CaCl2 ---> Ca + Cl2
(D) 2H2O ---> O2 + 4H+ + 4e¯
53. The permnganate ion is an excellent oxidizing agent in aqueous solutions. When the
MnO4¯ + H+ + e¯ ---> MnO2 + H2O
is balanced, the correct coefficients for the species involved are
54. A fuel cell uses H2(g) and Cl2(g) to form HCl(aq, 1M). Given
the standard electrode potential:
Cl2(g) + 2e¯ ---> 2Cl¯(aq) ; E° = 1.36 V,
What is the maximum electrical work per mole of HCl produced possible from this cell?
(A) 1.36 J
(B) 131,2 kJ
(C) 262,4 kJ
(D) 96,500 kJ
55. Given the standard electrode potentials:
Sn4+ (aq) + 2e¯ ---> Sn2+ (aq)
||E° = 0.154 V
|Br (aq) + 2e¯ ---> 2Br¯ (aq)
||E° = 1.087 V
The equilibrium constant for the process
2Br¯ (aq) + Sn4+ (aq) ---> Br2(aq) + Sn2+
(A) 1.19 x 10¯42
(B) 3.02 x 10¯32
(C) 3.31 x 1031
(D) 8.40 x 1041
56. The rate of the reaction 2NO + Cl2 ---> 2NOCl is given by the rate
equation rate = k[NO]2[Cl2]. The value of the rate constant can be
(A) increasing the concentration of the NO.
(B) increasing the concentration of the Cl2.
(C) increasing the temperature.
(D) doing all of these.
57. For a certain reaction the rate law is rate = k[C]3/2. If the rate of
the reaction is 0.020 mol L¯1 s¯1 when [C] = 1.0 M, what is the
rate when [C] = 0.60 M?
(A) 0.0093 mol L¯1 s¯1
(B) 0.012 mol L¯1 s¯1
(C) 0.033 mol L¯1 s¯1
(D) 0.040 mol L¯1 s¯1
58. The rate of the reaction 2A + B ---> Products is consistent with the rate
equation rate = k[A][B]. Which reaction mechanism is consistent with this information?
||A + B ---> AB
||AB + A ---> Products
A + A ---> A2
A2 + B ---> Products
||A + B ---> AB
A + A ---> A2
A2 + B ---> Products
59. Starting with three different kinds of amino acids, how many different kinds of
tripeptide molecules can be made assuming each amino acid can be used more then once in a
molecule if desired?
60. Metallic zinc will not reduce
61. What is the pH of the solution that results when 10.0 mL of 0.10 HF (Ka=
6.7 x 10¯4) and 10.0 mL of 0.040 M NaOH are mixed?
62. A saturated fat is
(A) a substance containing only carbon dioxide and hydrogen.
(B) a fat containing C=C bonds
(C) a substance in which the fat is at its maximum concentration in a solvent .
(D) the primary constitutent of proteins.
63. The principal used of hydrofluoric acid is
(A) in etching glass.
(B) as a bleaching agent.
(C) as an extremely stong oxidizing agent.
(D) in the preparation of organic fluorine compounds.
64. Important commerical extraction processes as welll as analytical techniques utilize
carbon dioxide above its critical temperature. Under theses conditions
(A) carbon dioxide is no longer fluid.
(B) one cannot condense the gas to a separate liquid state by applying pressure.
(C) carbon dioxide is known commonly as dry ice.
(D) carbon dioxide dissociates readily into C and O2.
65. The amount of copper in a 2.00 gram sample of the mineral cuprite was determined by
dissolving the sample in nitric acid (HNO3) and adding an excess of iodide
(I¯) solution to the resultant copper nitrate [Cu(NO3)2] solution.
The iodine (I2) liberated required 15.7 mL of 0.200 molar sodium thiosulfate
(Na2S2O3) solution to be titrated to an end point. What
is the percentage of Cu in the mineral? The essential reactions are:
2Cu2+ + 4I¯ ---> 2CuI + I2
I2 + 2S2O32¯ ---> S4O62¯
66. A mixture of n-hexane (C6H14) gas and O2 gas was
inside the vessel was 340 torr. After ignition by a spark, the mixture reacted completely
to form CO gas, CO2 gas, and steam. The total pressure exerted by the gaseous
products in the vessel was then 520 torr when the temperature was restored to 297 °C.
Which of the following equations describes the reaction that took place in the vessel?
(A) C6H14 + 8 O2 ---> 3 CO + 3 CO2 + 7 H2O
(B) C6H14 + 7 O2 ---> 5 CO + CO2 + 7 H2O
(C) 2 C6H14 + 17 O2 ---> 4 CO + 8 CO2 + 14
(D) 2 C6H14 + 15 O2 ---> 8 CO + 4 CO2 + 14
67. When conducting analyses of substances that are weak acids by titrating solutions
with a standardized strong base, the end-point indicator is chosen so that
(A) its color change occurs around the neutralization pH of 7.00.
(B) its color change occurs when the pH is about the same as the pKa of the weak acid.
(C) its color change occurs at a pH that is more basic than pH = 7.00.
(D) its color change occurs at a pH that is the same as that of the standardized base
68. Given the following standard enthalpies of formation in kJ mol¯1: CH4
(g), -74.8; C2H2(g), 226,9; C2H4(g), 52.6; C2H6(g),
-84.5; CO2(g), -393.5; H2O(l), -285.8. Which hydrocarbon, when
burned in an excess of oxygen will give the greatest heat per gram of fuel burned?
69. A solution may contain NaOH, Na2CO3, NaHCO3
individually or as any pair of the three analytes. A 25.00 mL aliquot of the solution is
treated with a standard solution of HCl. Given the following information, what analyte(s)
are present in the solution? Addition of phenolphthalein indicator (pink at pH = 10.0 and
colorless at pH =8.3) to the solution gives a pink color. The solution becomes colorless
after the methyl purple indictor (green at pH = 5.4 and purple at pH = 4.8) to this
solution gives a green color. Addition of an additional 15.30 mL of 0.1000 M HCl causes
the solution to become purple in color.
(A) NaOH + NaCO3
(B) NaCO3 + NaHCO3
70. A student standardizes a solution of HCl versus the primary standard Na2CO3.
He obtains the following values of the molarity of the HCl: 0.1055, 0.0998, 0.1032,
0.0995, 0.0990. What is the relative standard deviation (RSD) of his data expressed in
parts per thousand (ppt)?
(A) 5 ppt
(B) 12 ppt
(C) 24 ppt
(D) 27 ppt
1. (14pts.) A 1.529 gram of unknown ester is hydrolyzed in 50.00 ml of 0.236 M NaOH
solution according to the reaction below:
RCO2R'(l) + Na+(aq) + OH¯(aq) --->RCO2¯(aq)
+ Na+(aq) + R'OH(l)
When the reaction is complete, the excess OH¯ ion remaining in this solution is
titrated with a 0.115 M HCl solution. The volume of HCl required in this back-titration is
(a) How many moles of HCl are consumed during the back-titration?
(b) How many moles of NaOH reacted with the ester?
(c) Calculate the molar mass of the ester.
(d) If the molar mass of the alcohol is 74 g/mol, calculate the molar mass of the acid
which can be obtained by acidifying the final solution.
(e) Write structural formulas for all possible isomers of the alcohol.
2. (10pts.) Three samples of magnesium(II) sulfate hydrate were heated
and the following results were obtained.
(Sample No. 1)
(Sample No. 2)
(Sample No. 3)
|Mass of hydrate
|Mass heated salt
|Mass water lost
(A) Calculate the percentage of the water that was originally present in each sample.
(B) What percentage of water should be reprted for the magnesium sulfate hydrate? Why?
(C) How can you tell when you have heated a sample long enough?
(D) calculate the number of moles of water per mole of magnesium sulfate in the hydrate.
3. (14pts.) The pH of a 0.65 M solution of HF is 1.66.
(A) Calculate the ionization constant for HF.
(B) Calculate the pH of the solution which resultss when 15.00 mL of 0.42 M NaOH are added
to 25.00 mL of 0.65 M HF.
(C) Calculate the [Ca2+] required to initiate the precipitation of CaF2
in the solution in (b). (Ksp = 3.9 x 10¯11)
(D) Calculate the pH of the solution when 25.00 mL of 0.65 M HF is just neutralized with
0.15 M NaOH.
4. (12pts.) Mixtures of hydrazine (N2H4) and
dinitrogen tetroxide (N2O4) have been used as rocket fuels due to
their reaction to form N2 and H2O. The enthalpies of formation of
these substances are:
(A) Write a balanced chemical equation for the reaction of hydrazine and dinitrogen
tetroxide to form nitrogen and water.
(B) Calculate the enthalpy change for the reaction in (a).
(C) Calculate the heat which would be released if 0.0800 moles of N2H4
is reacted with a stoichiometric amount of N2O4.
(D) If 3.36 grams of N2 and 2.88 grams of H2O are produced in the
reaction in (c), calculate the temperature to which these gases would be heated if the
specific heats of N2 and H2O are 1.04 J / g-K and 1.87 J / g-K,
repectively. (Assume that the heat released contributes only to increasing the temperature
of the gaseous products and not of the container. Assume also that the specific heats of N2
and H2O do not vary with temperature.)
5. (16 pts.) Write equations for the following reactions.
(A) Concentrated sodium hydroxide solution is added to solid ammonium chloride.
(B) A solution of potassium bromide is added to a solution of Mercury(I) nitrate.
(C) Titanium(IV) chloride is added to excess water.
(D) Equimolar quantities of chlorine and toluene (methylbenzene) are mixed in the presence
of iron(III) chloride in the dark.
(E) Solutions of potassium permanganate and tin(II) chloride are mixed.
(F) Radon-220 decays with the loss of an alpha particle.
(G) Excess sodium hydroxide is added to a solution of aluminum nitrate.
(H) Carbon dioxide is bubbled through a solution of calcium hydroxide.
6. (12 pts.) Account for the following observations:
(A) Calcium exhibits an oxidation state of +2 in virtually all its compounds even
though the energy required to remove the second electron from a Ca atom is twice as great
as that required to remove the first.
(B) the melting point of MgO is 2800 °C while that of NaF is 993 °C even though the
interionic distances are within 10% of one another (2.10 x 10¯8 cm for MgO and
2.31 x 10¯8 cm for NaF).
(C) The F-E-F bond angles in the compounds ClF3, NF3, and BF3
are 90°, 101° and 120°, respectively.
(D) HF is a stronger acid than NH3 but a weaker acid than HCl even though the
electronegativity of fluorine (4.0) is greater than those of nitrogen and chlorine, which
are almost the same (3.1 and 2.9, respectively).
7. (12pts.) The standard reduction potentials for a series of technetium
TcO4¯ ---> TcO3
TcO4¯ ---> Tc2+
TcO3 ---> TcO2
TcO2 ---> Tc2+
Tc2+ ---> Tc
(A) Identify the reactions whose potentials would be affected by pH change and explain
(B) Compare the relative oxidizing ability of the TcO4¯ in its reaction to
form the Tc2+ ion with that of the MnO4¯ ion when it forms the Mn2+
ion (E° = 1.51 volts). Explain your answer.
(C) Identify the species which would be thermodynamically unstable toward
disproportionation and explain your reasoning.
(D) Predict the product formed when technetium metal is added to 1.0 M HCl and explain
8. (10pts.) Metal oxides are often referred to as bases and nonmetal oxides as acids.
(A) Explain why.
(B) Write chemical equations which illustrated the basicity of a metal oxide and the
acidity of a nonmetal oxide.
(C) Which oxide is more basic? Cr2O3 or CrO3? Why?
(D) Which oxide is more basic? Al2O3 or In2O3?