AP Chemistry Spring Semester Midterm Review
Please write your responses to the following questions on your own paper. Try to use only the
periodic table and equation sheet as you work out your solutions.
1981
1) Ammonia hydrogen sulfide is a crystalline solid that decomposes as follows:
NH4HS(s) -----> NH3(g) + H2S(g)
(a) Some solid NH4HS is placed in an evacuated vessel at 25 °C. After equilibrium is attained , the total pressure inside the
vessel is found to be 0.659 atmosphere. Some solid NH 4HS remains in the vessel at equilibrium. For this decomposition,
write the expression for Kp and calculate its numerical value at 25 °C.
(b) Some extra NH3 gas is injected into the vessel containing the sample described in part (a). When equilibrium is
reestablished at 25 °C, the partial pressure of NH 3 is twice the partial pressure of H2S. Calculate the numerical value of the
partial pressure of NH3 and the partial pressure of H2S in the vessel after the NH3 has been added and equilibrium has been
reestablished.
(c) In a different experiment, NH3 gas and H2S gas are introduced into an empty 1.00-liter vessel at 25 °C. The initial partial
pressure of each gas is 0.500 atmosphere. Calculate the number of moles of solid NH 4HS that is present when equilibrium is
established.
3) A 1.2156-gram sample of a mixture of CaCO3 and Na2SO4 was analyzed by dissolving the sample and completely
precipitating the Ca2+ as CaC2O4. The CaC2O4 was dissolved in sulfuric acid and the resulting H2C2O4 was titrated with a
standard KMnO4 solution.
(a) On a page of your answer booklet, write the balanced equation for the titration reaction, shown balanced below.
MnO4¯ + H2C2O4 + H+ ----> Mn2+ + CO2 + H2O
Indicate which substance is the oxidizing agent and which substance is the reducing agent.
(b) The titration of the H2C2O4 obtained required 35.62 milliliters of 0.1092-molar MnO4¯ solution. Calculate the number of
moles of H2C2O4 that reacted with the MnO4¯.
(c) Calculate the number of moles of CaCO3 in the original sample.
(d) Calculate the percentage by weight of CaCO3 in the original sample.
6) Assume that you have two different gases that you know are not cyclic (i.e., not ring) compounds, each with the following
elementary analysis: C, 85.7 per cent; H, 14.3 per cent. Each gas has a molecular weight
of 56 ± 1 g/mol.
(a) What is the molecular formula for the compounds?
(b) Draw the structural formulas for the four possible noncyclic isomers with this molecular formula.
(c) In the presence of an appropriate catalyst, both gases add hydrogen. The hydrogenated products are identical, their
molecular weight is 58. Which of the structures you drew to answer (b) can definitely be eliminated on the basis of this
additional information?
8) The emission spectrum of hydrogen consists of several series of sharp emission lines in the ultraviolet (Lyman series). in the
visible (Balmer series). and in the infrared (Paschen series, Brackett series, etc,) regions of the spectrum.
(a) What feature of the electronic energies of the hydrogen atom explains why the emission spectrum consists of discrete
wavelengths rather than a continuum of wavelengths?
(b) Account for the existence of several series of lines in the spectrum. What quantity distinguishes one series of lines from
another?
(c) Draw an electronic energy level diagram for the hydrogen atom and indicate on it the transition corresponding to the line
of lowest frequency in the Balmer series.
1982 -- 5
(a) From the standpoint of the kinetic-molecular theory, discuss briefly the properties of gas molecules that cause deviations from
ideal behavior.
(b) At 25 °C and 1 atmosphere pressure, which of the following gases shows the greatest deviation from ideal behavior? Give
two reasons for your choice.
CH4
SO2
O2
H2
(c) Real gases approach ideality at low pressure, high temperature, or both. Explain these observations
6) The values of the first three ionization energies (I1, I2, I3) for magnesium and argon [in kJ/mole] are as follows:
I1
I2
I3
Mg
735
1443
7730
Ar
1525
2665
3945
(a) Give the electronic configuration of Mg and Ar.
(b) In terms of these configurations, explain why the values of the first and second ionization energies of Mg are significantly
lower than the values for Ar, whereas the third ionization energy of Mg is much larger than the third ionization energy Ar.
(c) If a sample of Ar in one container and a sample of Mg in another container are each heated and chlorine is passed in to
each container, what compounds, if any , will be formed? Explain in terms of the electronic configuration given in part (a).
(d) Element Q has the following first three ionization energies [in kJ/mole]:
Q
I1
496
I2
4568
I3
6920
What is the formula for the most likely compound of element Q with chlorine? Explain the choice of formula on the basis of
the ionization energies.
1983
1) Sulfuryl chloride, SO2Cl2, is a highly reactive gaseous compound. When heated, it decomposes as follows.
SO2Cl2(g) <===> SO2(g) + Cl2(g)
This decomposition is endothermic. A sample of 3.509 grams of SO 2Cl2 is placed in an evacuated 1.00-liter bulb and the
temperature is raised to 375 K.
(a) What would be the pressure in atmospheres in the bulb if no dissociation of the SO 2Cl2(g) occurred?
(b) When the system has come to equilibrium at 375 K, the total pressure in the bulb is found to be 1.43 atmospheres.
Calculate the partial pressures of SO2, Cl2, and SO2Cl2 at equilibrium at 375 K.
(c) Give the expression for the equilibrium constant (either K p or Kc) for the decomposition of SO2Cl2(g) at 375 K. Calculate
the value of the equilibrium constant you have given, and specify its units.
(d) If the temperature were raised to 500 K, what effect would this have on the equilibrium constant? Explain briefly.
3) The molecular weight of a monoprotic acid HX was to be determined. A sample of 15.126 grams of HX was dissolved in
distilled water and the volume brought to exactly 250.00 milliliters in a volumetric flask. Several 50.00-milliliter portions of
this solution were titrated against NaOH solution, requiring an average of 38.21 milliliters of NaOH. The NaOH solution was
standardized against oxalic acid dihydrate, H2C2O4 . 2 H2O (molecular weight: 126.066 gram mol¯ 1). The volume of NaOH
solution required to neutralize 1.2596 grams of oxalic acid dihydrate was 41.24 milliliters.
(a) Calculate the molarity of the NaOH solution.
(b) Calculate the number of moles of HX in a 50.00-milliliter portion used for titration.
(c) Calculate the molecular weight of HX.
(d) Discuss the effect on the calculated molecular weight of HX if the sample of oxalic acid dihydrate contained a nonacidic
impurity.
1984
4) Give a scientific explanation for the following observations. Use equations or diagrams if they are relevant.
(a) It takes longer to cook an egg until it is hard-boiled in Denver (altitude 1 mile above sea level) than it does in New York
City (near sea level).
(b) Burning coal containing a significant amount of sulfur leads to "acid rain."
(c) Perspiring is a mechanism for cooling the body.
(d) The addition of antifreeze to water in a radiator decreases the likelihood that the liquid in the radiator will either freeze or
boil.
1985
3) The formula and the molecular weight of an unknown hydrocarbon compound are to be determined by elemental analysis
and the freezing-point depression method.
(a) The hydrocarbon is found to contain 93.46 percent carbon and 6.54 percent hydrogen. Calculate the empirical formula of
the unknown hydrocarbon.
(b) A solution is prepared by dissolving 2.53 grams of p-dichlorobenzene (molecular weight 147.0) in 25.86 grams of
naphthalene (molecular weight 128.2). Calculate the molality of the p-dichlorobenzene solution.
(c) The freezing point of pure naphthalene is determined to be 80.2 °C. The solution prepared in (b) is found to have an initial
freezing point of 75.7 °C. Calculate the molal freezing-point depression constant of naphthalene.
(d) A solution of 2.43 grams of the unknown hydrocarbon dissolved in 26.7 grams of naphthalene is found to freeze initially
at 76.2 °C. Calculate the apparent molecular weight of the unknown hydrocarbon on the basis of the freezing-point
depression experiment above.
(e)What is the molecular formula of the unknown hydrocarbon?
5) Properties of the chemical elements often show regular variations with respect to their positions in the periodic table.
(a) Describe the general trend in acid-base character of the oxides of the elements in the third period (Na to Ar). Give
examples of one acidic oxide and one basic oxide and show with equations how these oxides react with water.
(b) How does the oxidizing strength of the halogen elements vary down the group? Account for this trend.
(c) How does the reducing strength of the alkali metals vary down the group? Account for this trend.
7) Describe a separate laboratory procedure for preparing each of the following.
(a) Pure barium sulfate from an aqueous solution of barium chloride.
(b) A pure aqueous solution of copper(II) nitrate from solid copper(II) carbonate.
(c) A pure aqueous solution of calcium chloride from an aqueous solution of calcium bromide.
9)
Substance
H2
C3H8
HF
CsI
LiF
SiC
Melting Point, °C
-129
-190
-92
621
870
>2,000
(a) Discuss how the trend in the melting points of the substances tabulated above can be explained in terms of the types of
attractive forces and/or bonds in these substances.
(b) For any pairs of substances that have the same kind(s) of attractive forces and/or bonds, discuss the factors that cause
variation in the strengths of the forces and/or bonds.
1986
3) Three volatile compounds X, Y, and Z each contain element Q. The percent by weight of element Q in each compound was
determined. Some of the data obtained are given below.
Percent by Weight
Compound
Molecular Weight
of Element Q
X
64.8%
?
Y
73.0%
104.
Z
59.3%
64.0
(a) The vapor density of compound X at 27 °C and 750. mm Hg was determined to be 3.53 grams per liter. Calculate the
molecular weight of compound X.
(b) Determine the mass of element Q contained in 1.00 mole of each of the three compounds.
(c) Calculate the most probable value of the atomic weight of element Q.
(d) Compound Z contains carbon, hydrogen, and element Q. When 1.00 gram of compound Z is oxidized and all of the
carbon and hydrogen are converted to oxides, 1.37 grams of CO 2 and 0.281 gram of water are produced. Determine the most
probable molecular formula of compound Z.
1987
5) Use the details of modern atomic theory to explain each of the following experimental observations.
(a) Within a family such as the alkali metals, the ionic radius increases as the atomic number increases.
(b) The radius of the chlorine atom is smaller than the radius of the chloride ion, Cl¯. (Radii: Cl atom = 0.99 Å; Cl= ion =
1.81 Å)
(c) The first ionization energy of aluminum is lower than the first ionization energy of magnesium. (First ionization energies:
12Mg = 7.6 ev, 13Al = 6.0 ev)
(d) For magnesium, the difference between the second and third ionization energies is much larger than the difference
between the first and second ionization energies. (Ionization energies, in electron-volts, for Mg: 1st = 7.6, 2nd = 14, 3rd = 80)
9) Two important concepts that relate to the behavior of electrons in atomic system are the Heisenberg uncertainty principle
and the wave-particle duality of matter.
a) State the Heisenberg uncertainty principle as it relates to determining the position and momentum of an object.
b) What aspect of the Bohr theory of the atom is considered unsatisfactory as a result of the Heisenberg uncertainty principle?
c) Explain why the uncertainty principle or the wave nature of particles is not significant when describing the behavior of
macroscopic objects, but is very significant when describing the behavior of electrons.
1989
5)
CF4
XeF4
ClF3
(a) Draw a Lewis electron-dot structure for each of the molecules above and identify the shape of each.
(b) Use the valence shell electron-pair repulsion (VSEPR) model to explain the geometry of each of these molecules.