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Frequent Response Questions
Chapter 4: Representing Aqueous Reactions
1.)
Give the formulas to show the reactants and the products for FIVE of the following
chemical reactions. Each of the reactions occurs in aqueous solution unless otherwise
indicated. Represent substances in solution as ions if the substance is extensively
ionized. Omit formulas for any ions or molecules that are unchanged by the reaction.
In all cases a reaction occurs. You need not balance. (Look up solubility guidelines!)
Example: A strip of magnesium is added to a solution of silver nitrate.
Mg + Ag+ ---> Mg2+ + Ag
(a) Excess sodium cyanide is added to a solution of silver nitrate.
(b) Solutions of manganese(II) sulfate and ammonium sulfide are mixed.
(c) Phosphorous(V) oxide powder is sprinkled over distilled water.
(d) Solid ammonium carbonate is heated.
(e) Carbon dioxide gas is bubbled through a concentrated solution of potassium
hydroxide.
(f) A concentrated solution of hydrochloric acid is added to solid potassium
permanganate.
(g) A small piece of sodium metal is added to distilled water.
(h) A solution of potassium dichromate is added to an acidified solution of iron(II)
chloride.
2.)
Give the formulas to show the reactants and the products for FIVE of the following
chemical reactions. Each of the reactions occurs in aqueous solution unless otherwise
indicated. Represent substances in solutions as ions if the substance is extensively
ionized. Omit formulas for any ions or molecules that are unchanged by the reaction.
In all cases a reaction occurs. You need not balance. Example: A strip of magnesium
is added to a solution of silver nitrate.
(a) Solutions of zinc sulfate and sodium phosphate are mixed.
(b) Solutions of silver nitrate and lithium bromide are mixed.
(c) A stream of chlorine gas is passed through a solution of cold, dilute sodium
hydroxide.
(d) Excess hydrochloric acid solution is added to a solution of potassium sulfite.
(e) A solution of tin(II) chloride is added to an acidified solution of potassium
permanganate.
(f) A solution of ammonium thiocyanate is added to a solution of iron(III) chloride.
(g) Samples of boron trichloride gas and ammonia gas are mixed.
(h) Carbon disulfide vapor is burned in excess oxygen.
3.)
Give a scientific explanation for the following observations. Use equations of
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. ((d) Not Testable on New AP
Exam)
Chapter 5: Gas’s
1.)
A student collected a sample of hydrogen gas by the displacement of water as shown
by the diagram above. The relevant data are given in the following table.
(a) Calculate the number of moles of hydrogen gas collected.
(b) Calculate the number of molecules of water vapor in the sample of gas.
(c) Calculate the ratio of the average speed of the hydrogen molecules to the average
speed of the water vapor molecules in the sample.
(d) Which of the two gases, H2 or H2O, deviates more from ideal behavior? Explain
your answer.
Chapter 6: Thermochemistry
1.)
Chapter 8: Periodic Properties of Elements
1.)
Use principle of atomic structure and/or chemical bonding to answer of each of the
following.
(a) The radius of the Ca atom is 0.197 nanometer; the radius of the Ca2+ ion is 0.099
nanometer. Account for this difference.
(b) The lattice energy of CaO(s) is -3,460 kilojoules per mole; the lattice energy for
K2O(s) is -2,240 kilojoules per mole. Account for this difference.
(c) Explain the difference between Ca and K in regard to:
(i) their first ionization energies.
(ii) their second ionization energies.
(d) The first ionization energy of Mg is 738 kilojoules per mole and that of Al is 578
kilojoules per mole. Account for this difference.
2.)
The melting points of the alkali metals decrease from Li to Cs. In contrast, the melting
of the halogens increase from F2to I2.
(a) Using bonding principles, account for the decrease in the melting point of the
alkali metals.
(b) Using bonding principles, account for the increase in the melting points of the
halogens.
(c) What is the expected trend in the melting points of the compounds LiF, NaCl,
KBr, and CsI? Explain this trend using bonding principles.
Chapter 10: Chemical Bonding II: Molecular Shapes, Valence Bond Theory, and MO Theory
1.)
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.
Chapter 11: Liquids, Solids, and IMFs
1.)
For each of the following, use appropriate chemical principles to explain the
observation.
(a) Sodium chloride may be spread on an icy sidewalk in order to melt the ice;
equimolar amounts of calcium chloride are even more effective.  As of now,
this question will not be asked on the AP exam due to the fact that your knowledge
of colligative properties will not be tested (However, colligative properties will still
be tested on the regents exam)
(b) At room temperature, NH3 is a gas and H2O is a liquid, even though NH3 has a
molar mass of 17 grams and H2O has a molar mass of 18 grams.
(c) C (graphite) is used as a lubricant, whereas C (diamond) is used as an abrasive.
(d) Pouring vinegar onto the white reside in a kettle used for boiling water results in
fizzing/bubbling phenomenon.
2.)
Given solid samples of KI and of (NH4)2CO3. Briefly describe four simple laboratory
tests by which these two compounds can be distinguished. For each test, report the
expected result for each compound.
3.)
Discuss some differences in physical and chemical properties of metals and
nonmetals. What characteristic of the electronic configurations of atoms distinguishes
metals from nonmetals? On the basis of this characteristic, explain why there are
many more metals than nonmetals.
Chapter 13: Reaction Rates
1.)
2 NO(g) + 2 H2(g) ---> N2(g) + 2 H2O(g)
Experiments conducted to study the rate of the reaction represented by the equation
above. Initial concentrations and rates of reaction are given in the table below.
(a)
(i) Determine the order for each of the reactants, NO and H 2, from the data given and
show your reasoning.
(ii) Write the overall rate law for the reaction.
(b) Calculate the value of the rate constant, k, for the reaction. Include units.
(c) For experiment 2, calculate the concentration of NO remaining when exactly onehalf of the original amount of H2 has been consumed.
(d) The following sequence of elementary steps is a proposed mechanism for the
reaction.
I. NO + NO <===> N2O2
II. N2O2 + H2 ---> H2O + N2O
III. N2O + H2 ---> N2 + H2O
Based on the data present, which of the above is the rate-determining step? Show that
the mechanism is consistent with:
(i)
the observed rate law for the reaction, and
(ii)
the overall stoichiometry of the reaction.
2.)
C2H4 (g) + H2 (g) ---> C2H6 (g)
For the above reaction, H° = -137 kJ
Account for the following observations regarding the exothermic reaction represented
by the equation above.
(a) An increase in the pressure of the reactants causes an increase rate.
(b) A small increase in temperature causes a large increase in the reaction rate.
(c) The presence of metallic nickel causes an increase in reaction rate.
(d) The presence of powdered nickel causes a larger increase in reaction rate than does
the presence of a single pieces of nickel of the same mass.
3.)
For a hypothetical chemical reaction that has the stoichiometry 2 X + Y ---> Z, the
following initial rate data were obtained. All the measurements were made at the same
temperature.
Initial Rate of
Initial [X]o, Initial [Y]o,
Formation of Z,
mol L¯1
mol L¯1
mol L¯1 sec¯1
7.0 x 10¯4
0.20
0.10
3
1.4 x 10¯
0.40
0.20
2.8 x 10¯3
0.40
0.40
3
4.2 x 10¯
0.60
0.60
(a) Give the rate law for this reaction from the data above
(b) Calculate the specific rate constant for this reaction and specify its units.
(c) How long must the reaction proceed to produce a concentration of Z equal to 0.20
molar, if the initial reaction concentrations are [X]o = 0.80 molar, [Y]o = 0.60
molar, and [Z]o = 0 molar?
(d) Select from the mechanisms below the one most consistent with the observed data,
and explain your choice. In these mechanisms M and N are reaction intermediates.
(NB - the words fast and slow are placed above the reaction arrow on the original
test.)
(1)
X + Y ----> M (slow)
X + M ----> Z (fast)
(2)
X + X <===> M (fast)
Y + M----> Z (slow)
(3)
Y ---> M (slow)
M + X ---> N (fast)
N + X ----> Z (fast)
4.)
Chapter 14: Chemical Equilibrium
1.)
Discuss the following phenomena in terms of the chemical and physical properties of
the substances involved and general principles of chemical and physical change.
(a) As the system shown above approaches equilibrium, what change occurs to the
volume of water in beaker A ? What happens to the concentration of the sugar
solution in beaker B ? Explain why these changes occur.
(b) A bell jar connected to a vacuum pump is above. As the air pressure under the
bell jar decreases, what behavior of water in the beaker will be observed?
Explain why this occurs.
(c) What will be observed on the surfaces of zinc and silver strips shortly after they
are placed in separate solutions of CuSO4, as shown above. Account for these
observations.
(d) A water solution of I2 is shaken with an equal volume of a nonpolar solvent
such as TTE (trichlorotrifluoroethane). Describe the appearance of this system
after shaking. (A diagram may be helpful.) Account for this observation.
Chapter 15: Acids and Bases
1.)
A chemical reaction occurs when 100. milliliters of 0.2000-molar HCl is added
dropwise to 100. milliliters of 0.100-molar Na3PO4 solution.
(a) Write the two net ionic equations for the formation of the major species.
(b) Identify the species that acts as both a Brönsted acid and as a Brönsted base in the
equations in (a). Draw the Lewis electron-dot diagram for this species.
(c) Sketch a graph using the axis provided, showing the shape of the titration curve
that results when 100. milliliters of the HCl solution is added slowly from a buret
to the Na3PO4 solution. Account for the shape of the curve.
(d) Write the equation for the reaction that occurs if a few additional milliliters of HCl
solution are added to the solution resulting from the titration in (c).
2.)
In an experiment to determine the molecular weight and the ionization constant for
ascorbic acid (vitamin C), a student dissolved 1.3717 grams of the acid in water to
make 50.00 milliliters of solution. The entire solution was titrated with a 0.2211molar NaOH solution. The pH was monitored throughout the titration. The
equivalence point was reached when 35.23 milliliters of the base had been added.
Under the conditions of this experiment, ascorbic acid acts as a monoprotic acid that
can be represented as HA.
(a) From the information above, calculate the molecular weight of ascorbic acid.
(b) When 20.00 milliliters of NaOH had been added during the titration, the pH of the
solution was 4.23. Calculate the acid ionization constant for ascorbic acid.
(c) Calculate the equilibrium constant for the reaction of the ascorbate ion, A¯, with
water.
(d) Calculate the pH of the solution at the equivalence point of the titration.
3.)
1) Sodium benzoate, C6H5COONa, is a salt of the weak acid, benzoic acid,
C6H5COOH. A 0.10-molar solution of sodium benzoate has a pH of 8.60 at room
temperature.
(a) Calculate the [OH¯] in the sodium benzoate solution described above.
(b) Calculate the value for the equilibrium constant for the reaction
C6H5COO¯ + H2O <===> C6H5COOH + OH¯
(c) Calculate the value of Ka, the acid dissociation constant for benzoic acid.
(d) A saturated solution of benzoic acid is prepared by adding excess solid benzoic
acid to pure water at room temperature. Since this saturated solution has a pH of
2.88, calculate the molar solubility of benzoic acid at room temperature.
4.)
Discuss the roles of indicators in the titration of acids and bases. Explain the basis of
their operation and the factors to be considered in selecting an appropriate indicator
for a particular titration.
5.)
Chapter 16: Aqueous Ionic Equilibrium
1.)
MgF2(s) <===> Mg2+(aq) + 2 F¯(aq)
In a saturated solution of MgF2 at 18° C, the concentration of Mg2+ is 1.21 x
10¯3 molar. The equilibrium is represented by the equation above.
(a) Write the expression for the solubility-product constant, Ksp, and calculate its
value at 18° C.
(b) Calculate the equilibrium concentration of Mg2+ in 1.000 liter of saturated
MgF2 solution at 18°C to which 0.100 mole of solid KF has been added. The KF
dissolves completely. Assume the volume change is negligible.
(c) Predict whether a precipitate of MgF2 will form when 100.0 milliliters of a 3.00 x
10¯3 molar Mg(NO3)2 solution is mixed with 200.0 milliliters of a 2.00 x
10¯3 molar NaF solution at 18°C. Calculations to support your prediction must be
shown.
(d) At 27°C the concentration of Mg2+ in a saturated solution of MgF2 is 1.17 x
10¯3 molar. Is the dissolving of MgF2 in water an endothermic or an exothermic
process? Give an explanation to support your conclusion.
2.)
Answer the following questions about the solubility of some fluoride salts of alkaline
earth metals.
(a) A student prepares 100. mL of a saturated solution of MgF2 by adding 0.50 g of
solid MgF2 to 100. mL of distilled water at 25’ C and stirring until no more
solid dissolves. (Assume that the volume of the undissolved MgF2 is negligibly
-
small.) The saturated solution is analyzed, and it is determined that [F ] in the
solution is 2.4x10-3 M.
(i) Write the chemical equation for the dissolving of solid MgF2 in water.
(ii) Calculate the number of moles of MgF2 that dissolved.
(iii) Determine the value of the solubility-product constant, Ksp , for MgF2 at
25oC.
(b) A beaker contains 500. mL of a solution in which both Ca2+(aq) and Ba2+(aq)
are present at a concentration of 0.10 M at 25oC. A student intends to separate
the ions by adding 0.20 M NaF solution one drop at a time from a buret. At
25oC the value of Ksp for CaF2 is 3.5x10-11; the value of Ksp for BaF2 is
1.8x10-6.
(i) Which salt will precipitate first, CaF2 or BaF2? Justify your answer.
For parts (b)(ii) and (b)(iii) below, assume that the addition of the NaF solution
does not significantly affect the total volume of the liquid in the beaker.
-
(ii) Calculate the minimum concentration of F (aq) necessary to initiate
precipitation of the salt selected in part (b)(i).
(iii) Calculate the minimum volume of 0.20 M NaF that must be added to the
beaker to initiate precipitation of the salt selected in part (b)(i).
(c) There are several ways to dissolve salts that have limited solubility. Describe
one procedure to re-dissolve the precipitate formed in part (b).
Chapter 17: Free Energy and Thermodynamics
1.)
2 H2S(g) + SO2(g) <===> 3 S(s) + 2 H2O(g)
At 298 K, the standard enthalpy change, H°, for the reaction represented above is 145 kilojoules.
(a) Predict the sign of the standard entropy change, S°, for the reaction. Explain the
basis for your prediction.
(b) At 298 K, the forward reaction (i.e., toward the right) is spontaneous. What
change, if any, would occur in the value ofG° for this reaction as the temperature
is increased? Explain your reasoning using thermodynamic principles.
(c) What change, if any, would occur in the value of the equilibrium constant, K eq, for
the situation described in (b)? Explain your reasoning.
(d) The absolute temperature at which the forward reaction becomes nonspontaneous
can be predicted. Write the equation that is used to make the prediction. Why does
this equation predict only an approximate value for the temperature?
2.)
Br2 (l) ---> Br2 (g)
At 25 °C the equilibrium constant, Kp, for the reaction above is 0.281 atmosphere.
(a) What is G°298 for this reaction?
(b) It takes 193 joules to vaporize 1.00 gram of Br2(l) at 25 °C and 1.00 atmosphere
pressure. What are the values of H°298and of S°298 for this reaction?
(c) Calculate the normal boiling point of bromine. Assume that H° and S° remain
constant as the temperature is changed.
(d) What is the equilibrium vapor pressure of bromine at 25 °C?
3.)
Substance
C(s)
CO2(g)
H2(g)
H2O(l)
O2(g)
C3H7COOH(l)
Standard Heat of
Formation, Hf°,
in kJ mol¯1
0.00
-393.5
0.00
-285.85
0.00
?
Absolute Entropy
S°, in J mol¯1K¯1
5.69
213.6
130.6
69.96
205.0
226.3
The enthalpy change for the combustion of butyric acid at 25 °C, H°comb, is -2,183.5
kilojoules per mole. The combustion reaction is
C3H7COOH(l) + 5 O2(g) <===> 4 CO2(g) + 4 H2O(l)
(a) From the data above, calculate the standard heat of formation, Hf°, for butyric
acid at 25 °C.
(b) Write a correctly balanced equation for the formation of butyric acid from its
elements.
(c) Calculate the standard entropy change, Sf°, for the formation of butyric acid at 25
°C. The entropy change, S°, for the combustion reaction above is -117.1 J K¯1 at 25
°C.
(d) Calculate the standard free energy of formation, Gf°, for butyric acid at 25 °C.
Chapter 18: Electrochemistry
The electrolysis of an aqueous solution of potassium iodide, KI, results in the
formation of hydrogen gas at the cathode and iodine at the anode. A sample of 80.00
milliliters of a 0.150-molar solution of KI was electrolyzed for 3.00 minutes, using a
constant current. At the end of this time, the I2 produced was titrated against a 0.225molar solution of sodium thiosulfate, which reacts with iodine according to the
equation below. The end point of the titration was reached when 37.2 milliliters of the
Na2S2O3 solution had been added.
I2 + 2 S2O32¯ ---> 2 I¯ + S4O62¯
(a) How many moles of I2 was produced during the electrolysis?
(b) The hydrogen gas produced at the cathode during the electrolysis was collected
over water at 25 °C at a total pressure of 752 millimeters of mercury. Determine
the volume of hydrogen collected. (The vapor pressure of water at 25 °C is 24
millimeters of mercury.)
(c) Write the equations for the half reaction that occurs at the anode during the
electrolysis.
(d) Calculate the current used during the electrolysis.
2.)
Consider three unlabeled bottles, each containing small pieces of one of the following
metals.
- Magnesium
- Sodium
- Silver
The following reagents are used for identifying the metals.
- Pure water
- A solution of 1.0-molar HCl
- A solution of concentrated HNO3
(a) Which metal can be easily identified because it is much softer than the other two?
Describe a chemical test that distinguishes this metal from the other two, using
only one of the reagents above. Write a balanced chemical equation for the
reaction that occurs.
(b) One of the other two metals reacts readily with the HCl solution. Identify the metal
and write the balanced chemical equation for the reaction that occurs when this
metal is added to the HCl solution. Use a table of standard reduction potentials to
account for the fact that this metal reacts with HCl while the other does not.
(c) The one remaining metal reacts with the concentrated HNO3 solution. Write a
balanced chemical equation for the reaction that occurs.
(d) The solution obtained in (c) is diluted and a few drops of 1 M HCl is added.
Describe what would be observed. Write a balanced chemical equation for the
reaction that occurs.
3.)