AP ChemistryName:
End of Year ReviewPeriod:
Unit 2:
1971
Permanganate ion, MnO4-, oxidizes sulfite ions to sulfate ion. The manganese product depends upon the pH of the reaction mixture. The mole ratio of oxidizing to reducing agent is two to five at pH 1, and is two to one at pH 13. For each of these cases, write a balanced equation for the reaction, and indicate the oxidation state of the manganese in the product containing manganese.
1981 B
A 1.2516 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)Write the balanced equation for the titration reaction, shown unbalanced below.
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.
Unit 3
1996 D (Required)
Represented above are five identical balloons, each filled to the same volume at 25C and 1.0 atmosphere pressure with the pure gases indicated.
(a)Which balloon contains the greatest mass of gas? Explain.
(b)Compare the average kinetic energies of the gas molecules in the balloons. Explain.
(c)Which balloon contains the gas that would be expected to deviate most from the behavior of an ideal gas? Explain.
(d)Twelve hours after being filled, all the balloons have decreased in size. Predict which balloon will be the smallest. Explain your reasoning.
2003 B
A rigid 5.00 L cylinder contains 24.5 g of N2(g) and 28.0 g of O2(g)
(a)Calculate the total pressure, in atm, of the gas mixture in the cylinder at 298 K.
(b)The temperature of the gas mixture in the cylinder is decreased to 280 K. Calculate each of the following.
(i)The mole fraction of N2(g) in the cylinder.
(ii)The partial pressure, in atm, of N2(g) in the cylinder.
(c)If the cylinder develops a pinhole-sized leak and some of the gaseous mixture escapes, would the ratio in the cylinder increase, decrease, or remain the same? Justify your answer.
A different rigid 5.00 L cylinder contains 0.176 mol of NO(g) at 298 K. A 0.176 mol sample of O2(g) is added to the cylinder, where a reaction occurs to produce NO2(g).
(d)Write the balanced equation for the reaction.
(e)Calculate the total pressure, in atm, in the cylinder at 298 K after the reaction is complete.
Unit 4
1988 D
An experiment is to be performed to determine the standard molar enthalpy of neutralization of a strong acid by a strong base. Standard school laboratory equipment and a supply of standardized 1.00 molar HCl and standardized 1.00 molar NaOH are available.
(a)What equipment would be needed?
(b)What measurements should be taken?
(c)Without performing calculations, describe how the resulting data should be used to obtain the standard molar enthalpy of neutralization.
(d)When a class of students performed this experiment, the average of the results was -55.0 kilojoules per mole. The accepted value for the standard molar enthalpy of neutralization of a strong acid by a strong base is -57.7 kilojoules per mole. Propose two likely sources of experimental error that could account for the result obtained by the class.
1995 B
Propane, C3H8, is a hydrocarbon that is commonly used as fuel for cooking.
(a)Write a balanced equation for the complete combustion of propane gas, which yields CO2(g) and H2O(l).
(b)Calculate the volume of air at 30C and 1.00 atmosphere that is needed to burn completely 10.0 grams of propane. Assume that air is 21.0 percent O2 by volume.
(c)The heat of combustion of propane is -2,220.1 kJ/mol. Calculate the heat of formation, Hf, of propane given that Hf of H2O(l) = -285.3 kJ/mol and Hf of CO2(g) = -393.5 kJ/mol.
(d)Assuming that all of the heat evolved in burning 30.0 grams of propane is transferred to 8.00 kilograms of water (specific heat = 4.18 J/g.K), calculate the increase in temperature of water.
Unit 5
2007 Part A
1.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 4.341014 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) O2(g) + O(g)
A molecule of O3(g) absorbs a photon with a frequency of 1.001015 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 1.001015 s-1 have enough energy to break this bond? Support your answer with a calculation.
2006 D
2.Suppose that a stable element with atomic number 119, symbol Q, has been discovered.
(a)Write the ground-state electron configuration for Q, showing only the valence-shell electrons.
(b)Would Q be a metal or a nonmetal? Explain in terms of electron configuration.
(c)On the basis of periodic trends, would Q have the largest atomic radius in its group or would it have the smallest? Explain in terms of electronic structure.
(d)What would be the most likely charge of the Q ion in stable ionic compounds?
(e)Write a balanced equation that would represent the reaction of Q with water.
(f)Assume that Q reacts to form a carbonate compound.
(i)Write the formula for the compound formed between Q and the carbonate ion, CO32–.
(ii)Predict whether or not the compound would be soluble in water. Explain your reasoning.
1997 D
3.Answer each of the following questions regarding radioactivity.
(a)Write the nuclear equation for decay ofby 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 .
(c)Describe, or trace, 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?
Unit 6
2000 D
- Answer the following questions about the element selenium, Se (atomic number 34).
a)Samples of natural selenium contain six stable isotopes. In terms of atomic structure, explain what these isotopes have in common, and how they differ.
b)Write the complete electron configuration (e.g., 1s2 2s2... etc.) for a selenium atom in the ground state. Indicate the number of unpaired electrons in the ground-state atom, and explain your reasoning.
c)In terms of atomic structure, explain why the first ionization energy of selenium is
iless than that of bromine (atomic number 35), and
iigreater than that of tellurium (atomic number 52).
Selenium reacts with fluorine to form SeF4. Draw the complete Lewis electron-dot structure for SeF4 and sketch the 1997 D (Required)
- Consider the molecules PF3 and PF5.
a)Draw the Lewis electron-dot structures for PF3 and PF5 and predict the molecular geometry of each.
b)Is the PF3 molecule polar, or is it nonpolar? Explain.
c)On the basis of bonding principles, predict whether each of the following compounds exists. In each case, explain your prediction.
(i)NF5
(ii)AsF5
d)molecular structure. Indicate whether the molecule is polar or nonpolar, and justify your answer.
Unit 8
2005 B
1.Answer the following questions related to the kinetics of chemical reactions.
I–(aq) + ClO–(aq) IO–(aq) + Cl–(aq)
Iodide ion, I–, is oxidized to hypoiodite ion, IO–, by hypochlorite, ClO–, in basic solution according to the equation above. Three initial-rate experiments were conducted; the results shown in the following table.
Experiment / [I–](mol L–1) / [ClO–]
(mol L–1) / Initial Rate of Formation of IO– (mol L–1 s–1)
1 / 0.017 / 0.015 / 0.156
2 / 0.052 / 0.015 / 0.476
3 / 0.016 / 0.061 / 0.596
(a)Determine the order of the reaction with respect to each reactant listed below. Show your work.
(i)I–(aq)
(ii)ClO–(aq)
(b)For the reaction,
(i)write the rate law that is consistent with the calculations in part (a);
(ii)calculate the value of the specific rate constant, k, and specify units.
The catalyzed decomposition of hydrogen peroxide, H2O2(aq), is represented by the following equation.
2 H2O2(aq) 2 H2O(l) + O2(g)
The kinetics of the decomposition reaction were studied and the analysis of the results show that it is a first-order reaction. Some of the experimental data are shown in the table below.
[H2O2](mol L–1) / Time (minutes)
1.00 / 0.0
0.78 / 5.0
0.61 / 10.0
(c)During the analysis of the data, the graph below was produced.
AP ChemistryName:
End of Year ReviewPeriod:
(i)Label the vertical axis of the graph
(ii)What are the units of the rate constant, k, for the decomposition of H2O2(aq) ?
(iii)On the graph, draw the line that represents the plot of the uncatalyzed first-order decomposition of 1.00 M H2O2(aq).
AP ChemistryName:
End of Year ReviewPeriod:
2007part B, question #6
2.The reaction between SO2(g) and O2(g) to form SO3(g) is represented below.
2 SO2(g) + O2(g) 2 SO3(g)
The reaction is exothermic. The reaction is slow at 25˚C; however, a catalyst will cause the reaction to proceed faster.
(a)Using the axes provided, draw the complete potential-energy diagram for both the catalyzed and uncatalyzed reactions. Clearly label the curve that represents the catalyzed reaction.
(b)Predict how the ratio of the equilibrium pressures, , would change when the temperature of the uncatalyzed reaction mixture is increased. Justify your prediction.
(c)How would the presence of a catalyst affect the change in the ratio described in part (b)? Explain.
Unit 9
2000 A
1.2 H2S(g) 2 H2(g) + S2(g)
When heated, hydrogen sulfide gas decomposes according to the equation above. A 3.40 g sample of H2S(g) is introduced into an evacuated rigid 1.25 L container. The sealed container is heated to 483 K, and 3.7210–2 mol of S2(g) is present at equilibrium.
(a)Write the expression for the equilibrium constant, Kc, for the decomposition reaction represented above.
(b)Calculate the equilibrium concentration, in molL-1, of the following gases in the container at 483 K.
(i)H2(g)
(ii)H2S(g)
(c)Calculate the value of the equilibrium constant, Kc, for the decomposition reaction at 483 K.
(d)Calculate the partial pressure of S2(g) in the container at equilibrium at 483 K.
(e)For the reaction H2(g) + S2(g) H2S(g) at 483 K, calculate the value of the equilibrium constant, Kc.
1998 D
C(s) + H2O(g) CO(g) + H2(g)Hº = +131kJ
A rigid container holds a mixture of graphite pellets (C(s)), H2O(g), CO(g), and H2(g) at equilibrium. State whether the number of moles of CO(g) in the container will increase,decrease, or remain the same after each of the following disturbances is applied to the original mixture. For each case, assume that all other variables remain constant except for the given disturbance. Explain each answer with a short statement.
(a)Additional H2(g) is added to the equilibrium mixture at constant volume.
(b)The temperature of the equilibrium mixture is increased at constant volume.
(c)The volume of the container is decreased at constant temperature.
(d)The graphite pellets are pulverized.
Unit 10
1986 D
H2SO3HSO3–HClO4HClO3H3BO3
Oxyacids, such as those above, contain an atom bonded to one or more oxygen atoms; one or more of these oxygen atoms may also be bonded to hydrogen.
(a)Discuss the factors that are often used to predict correctly the strengths of the oxyacids listed above.
(b)Arrange the examples above in the order of increasing acid strength.
2005 A
HC3H5O2(aq) ↔ C3H5O2–(aq) + H+(aq)Ka= 1.3410–5
Propanoic acid, HC3H5O2, ionizes in water according to the equation above.
(a)Write the equilibrium constant expression for the reaction.
(b)Calculate the pH of a 0.265 M solution of propanoic acid.
The methanoate ion, HCO2–(aq) reacts with water to form methanoic acid and hydroxide ion, as shown in the following equation.
HCO2–(aq) + H2O (l) ↔ H2CO2(aq) + OH–(aq)
(d)Given that [OH–] is 4.1810–6M in a 0.309 M solution of sodium methanoate, calculate each of the following.
(i)The value of Kb for the methanoate ion, HCO2–(aq)
(ii)The value of Ka for methanoic acid, HCO2H
(e)Which acid is stronger, propanoic acid or methanoic acid? Justify your answer.
Unit 11
1994 A
MgF2(s) Mg2+(aq) + 2 F-(aq)
In a saturated solution of MgF2 at 18ºC, the concentration of Mg2+ is 1.2110-3 molar. The equilibrium is represented by the equation above.
(a)Write the expression for the solubility-product constant, Ksp, and calculate its value at 18ºC.
(b)Calculate the equilibrium concentration of Mg2+ in 1.000 liter of saturated MgF2 solution at 18ºC 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 MgF2will form when 100.0 milliliters of a 3.0010-3-molar Mg(NO3)2 solution is mixed with 200.0 milliliters of a 2.00l0-3-molar NaF solution at 18ºC. Calculations to support your prediction must be shown.
(d)At 27ºC the concentration of Mg2+ in a saturated solution of MgF2is 1.1710-3 molar. Is the dissolving of MgF2in water an endothermic or an exothermic process? Give an explanation to support your conclusion.
1985 A
At 25ºC the solubility product constant, Ksp, for strontium sulfate, SrSO4, is 7.610-7. The solubility product constant for strontium fluoride, SrF2, is 7.910-10.
(a)What is the molar solubility of SrSO4 in pure water at 25ºC?
(b)What is the molar solubility of SrF2 in pure water at 25ºC?
(c)An aqueous solution of Sr(NO3)2 is added slowly to 1.0 litre of a well-stirred solution containing 0.020 mole F- and 0.10 mole SO42- at 25ºC. (You may assume that the added Sr(NO3)2 solution does not materially affect the total volume of the system.)
1.Which salt precipitates first?
2.What is the concentration of strontium ion, Sr2+, in the solution when the first precipitate begins to form?
(d)As more Sr(NO3)2 is added to the mixture in (c) a second precipitate begins to form. At that stage, what percent of the anion of the first precipitate remains in solution?
Unit 12
2006 B
CO(g) + O2(g) CO2(g)
The combustion of carbon monoxide is represented by the equation above.
(a)Determine the value of the standard enthalpy change, ∆H˚rxn for the combustion of CO(g) at 298 K using the following information.
C(s) + O2(g) CO(g)∆H˚298 = –110.5 kJ mol-1
C(s) + O2(g) CO2(g)∆H˚298 = –393.5 kJ mol-1
(b)Determine the value of the standard entropy change, ∆S˚rxn, for the combustion of CO(g) at 298 K using the information in the following table.
Substance / S˚298(J mol-1 K-1)
CO(g) / 197.7
CO2(g) / 213.7
O2(g) / 205.1
(c)Determine the standard free energy change, ∆G˚rxn, for the reaction at 298 K. Include units with your answer.
(d)Is the reaction spontaneous under standard conditions at 298 K? Justify your answer.
(e)Calculate the value of the equilibrium constant, Keq, for the reaction at 298 K.
2007 part A, form B, question #1
A sample of solid U308 is placed in a rigid 1.500 L flask. Chlorine gas, Cl2(g), is added, and the flask is heated to 862˚C. The equation for the reaction that takes place and the equilibrium-constant expression for the reaction are given below.
U308(s) + 3 Cl2(g), 3 UO2Cl2(g) + O2(g)
When the system is at equilibrium, the partial pressure of Cl2(g) is 1.007 atm and the partial pressure of UO2Cl2(g) is 9.73410-4 atm
(a)Calculate the partial pressure of O2(g) at equilibrium at 862˚C.
(b)Calculate the value of the equilibrium constant, KP, for the system at 862˚C.
(c)Calculate the Gibbs free-energy change, ∆G˚, for the reaction at 862˚C.
(d)State whether the entropy change, ∆S˚ for the reaction at 862˚C is positive, negative, or zero. Justify your answer.
(e)State whether the enthalpy change, ∆H˚, for the reaction at 862˚C is positive, negative, or zero. Justify your answer.
(f)After a certain period of time, 1.000 mol of O2(g) is added to the mixture in the flask. Does the mass of U308(s) in the flask increase, decrease, or remain the same? Justify your answer.
Unit 13
1998 D
Answer the following questions regarding the electrochemical cell shown.
(a)Write the balanced net-ionic equation for the spontaneous reaction that occurs as the cell operates, and determine the cell voltage.
(b)In which direction do anions flow in the salt bridge as the cell operates? Justify your answer.
(c)If 10.0 mL of 3.0-molar AgNO3 solution is added to the half-cell on the right, what will happen to the cell voltage? Explain.
(d)If 1.0 gram of solid NaCl is added to each half-cell, what will happen to the cell voltage? Explain.
(e)If 20.0 mL of distilled water is added to both half-cells, the cell voltage decreases. Explain.
2004 D Required
An electrochemical cell is constructed with an open switch, as shown in the diagram above. A strip of Sn and a strip of unknown metal, X are used as electrodes. When the switch is closed, the mass of the Sn electrode increases. The half-reactions are shown below.
Sn2+(aq) + 2 e– Sn(s)E˚ = –0.14 V
X3+(aq) + 3 e– X(s)E˚ = ?
(a)In the diagram above, label the electrode that is the cathode. Justify your answer.
(b)In the diagram above, draw an arrow indicating the direction of electron flow in the external circuit when the switch is closed.
(c)If the standard cell potential E˚cell is +0.60 V, what is the standard potential, in volts for the X3+/X electrode?
(d)Identify metal X.
(e)Write balanced net-ionic equation for the overall chemical reaction occurring in the cell.
(f)In the cell, the concentration of Sn2+ is changed from 1.0 M to 0.50 M, and the concentration of X3+ is changed from 1.0 M to 0.10 M.
(i)Substitute all appropriate values for determining the cell potential, Ecell, into the Nernst equation. (Do not do any calculations.)
(ii)On the basis of your response in (f) (i), will the cell potential be greater than, less than, or equal to E˚cell? Justify your answer.
Answers:
Unit 2:
1971
Answer:
2 MnO4- + 5 SO32- + 6 H+ 2 Mn2+ + 3 H2O + 5 SO42- (oxidation state Mn = +2)
2 MnO4- + SO32- + 2 OH- 2 MnO42- + H2O + SO42- (oxidation state Mn = +6)
1981 B
Answer:
(a)2 MnO4- + 5 H2C2O4 + 6 H+ 2 Mn2+ + 10 CO2 + 8 H2O
oxidizing agent: MnO4- , reducing agent: H2C2O4
(b)
(c)moles of H2C2O4 = moles CaCO3, therefore, 9.7210-3 mol H2C2O4 = 9.7210-3 mol CaCO3
(d)
Unit 3:
1996 D (Required)
Answer:
(a)CO2; according to Avogadro’s Hypothesis, they all contain the same number of particles, therefore, the heaviest molecule, CO2 (molar mass = 44), will have the greatest mass.
(b)all the same; at the same temperature all gases have the same kinetic energy.
(c)CO2; since they are all essentially non-polar, the largest intermolecular (London) force would be greatest in the molecule/atom with the largest number of electrons.
(d)He; it has the smallest size and has the greatest particulate speed and, therefore, it’s the easiest to penetrate the wall and effuse.
2003 B
Answer:
(a)24.5 g N2 x = 0.875 mol N2
28.0 g O2 x = 0.875 mol O2
P = =
= 8.56 atm
(b)(i) = 0.500 mole fraction N2
(ii)
= 8.05 atm x mole fraction = 8.05 atm x 0.500 = 4.02 atm N2
(c)decrease; since N2 molecules are lighter than O2 they have a higher velocity and will escape more frequently (Graham’s Law), decreasing the amount of N2 relative to O2
(d)2 NO + O2→ 2 NO2
(e)all 0.176 mol of NO will react to produce 0.176 mol of NO2, only 1/2 of that amount of O2 will react, leaving 0.088 mol of O2, therefore, 0.176 + 0.088 = 0.264 mol of gas is in the container.