AP Homework for 2012-2013 Year
Topic 1: Stoichiometry
Problem Sets: 3, 4, 5, 19 (omit 1-4), and 20
Topic 2: Gas laws
Problem Sets: 15 and 16
Topic 3: Atomic Structure
Problem Sets 9, 10, and 11
Topic 4: Periodic Table and Trends
Problem Set 8
Topic 5: Thermochemistry
Problem Sets 6 and 7
Topic 6: Bonding
Problem Sets 12 and 13
Topic 7: Geometry
Problem Sets `14 and questions 1-4 of problem set 19
Topic 8: Condensed Phases
Problem Sets 17 and 18
Topic 9: Introduction to Equilibrium
Problem sets 24 and 25
Topic 10: Acids and Bases
Problem Sets 26 and 27
Topic 11: Applications of Equilibrium
Problem Sets 28 and 29
Topic 12: Thermodynamics
Problem Set 30
Topic 13: Kinetics
Problem Sets 21, 22, and 23
Topic 14: Electrochemistry
Problem Sets 31, 32, and 33
Problem Set #1
ALL work must be shown to receive full credit.
PS1.1. Diagram each of the following systems as viewed at the atomic level in the space provided.
PS1.2. A golden colored cube measures 2.00 cm on each side and has a mass of 154.4 g. Could the cube
be made of gold? Explain your answer using both words and a mathematical calculation.
PS1.3. Which temperature of each pair is lower: (Justify each answer with a calculation.)
PS1.4. The approximate volume of water contained in the oceans of the world is 3.3 x 108 mi3.
Determine the volume in km3.
PS1.5. A line of water molecules 2.5 inches long contains 2.5 x 108 molecules. Calculate the diameter of
a water molecule in centimeters, angstroms and nanometers.
PS1.6. The density of water is 1.00 g/mL. What is the weight, in pounds of 1.00 gallon of water?
PS1.7a During a flood an average of 1.00 x 103 ft3 of water flowed out of a dam's flood gate each
minute. What volume in ft3 and liters flowed out of the gate in one hour?
b) If the flow was constant for one entire day, what mass of water in kilograms was released from the
dam? What weight in pounds?
PS1.8. Using a standard linear measuring device (ruler) measure the length and width of this sheet of
paper in both inches and centimeters. (Be sure to report the appropriate number of significant figures).
Find the area of this page in square centimeters.
Problem Set #2
ALL work must be shown in Problems 1 - 10 full credit.
PS2.1. Sketch a neutral fluorine atom indicating the proper number and approximate
location of protons, neutrons and electrons.
PS2.2. How many electrons would be needed to equal the mass of a single proton?
PS2.3. Complete the following table
PS2.4. Copper exists in nature in two isotopic forms with masses 62.96 u and 64.96
u. The accepted atomic mass for copper is 63.55 u. Determine the percent
abundance for each isotope. How many grams of each isotope would be found
in a 10.0 g sample of pure copper metal?
PS2.5 Write a balanced nuclear equation for each of the following transformations.
a).
undergoes beta decay.
b) Radium-226 decays to a radon isotope.
c). Neptunium-233 undergoes alpha decay.
d). Two 21H atoms undergo a fusion reaction.
e). Uranium-235 absorbs a neutron and undergoes fission producing barium-141,
3 neutrons, and another nucleus.
PS2.6. Predict formulas for each of the following combinations. Write the name of
each compound formed. (Note: One of the compounds has two correct
formulas and four correct names!)
a) gallium and oxygen
e) iron and nitrate
b) magnesium and sulfate
c) aluminum and chlorine
d) calcium and phosphate
PS2.7a. The simple hydrocarbons listed in exercise 20e. have the same general
formula. Complete the formula where n is the number of carbon atoms.
CnH?
PS2.7b. For molecules having more than three carbon atoms, the straight chain
arrangement shown in exercise 20e. is not the only possible arrangement of
atoms. For example, the molecular formula C5H12 has three possible
structures. Write the condensed molecular formula or draw the structure for
each of these compounds and name each one.
PS2.8. Many familiar substances have common, unsystematic names. In each of the
following cass, give the correct systematic name:
PS2.9. Write the chemical formula of each substance mentioned in the following
word descriptions.
a) Zinc carbonate can be heated to form zinc oxide and carbon dioxide.
b) On treatment with hydrofloric acid, silicon dioxide forms silicon
tetrafluoride and water.
c) Sulfur dioxide reacts with water to form sulfurous acid.
d) The substance hydrogen phosphide is commonly called phosphine.
e) Perchloric acid reacts with cadmium to form cadmium (II) perchlorate.
f) Vanadium (III) bromide is a colored solid.
PS2.10. Using the CRC Handbook of Chemistry and Physics, find the density, melting
point, boiling point and the phase at room temperature for CF4, CCl4, CBr4
and CI4.
Problem Set #3
ALL work must be shown to receive full credit.
PS3.1. Predict the products for each of the following chemical reactions and balance the
chemical equation.
PS3.2. Determine the number of moles in each of the following.
a) 25.6 g of acetic acid (HC2H3O2)
b) 1.89 x 10-4 g of Ca3(PO4)2
PS3.3. Determine the number of oxygen atoms in each of the following.
a) 6.451 moles of C6H8O6 (vitamin C)
b) 1.89 x 10-4 g of Ca3(PO4)2
PS3.4. Determine the mass in grams in each of the following.
a) 0.0721 moles of H3PO4 (phosphoric acid)
b) 72 atoms of sulfur
PS3.5. Determine the percent composition of each element in Ni3(PO4)2.
PS3.6. Determine the empirical formula for a compound which is 26.6 % potassium,
35.4 % chromium and 38.1 % oxygen.
PS3.7. When a solid compound which is 27.62 % Ca, 22.06 % S, 49.62 % O and
0.700 % H is heated, a solid product is isolated which is 29.45 % Ca, 23.52 %
S and 47.03 % O. What is the other product formed? Write the chemical
equation which describes the reaction which has occurred.
PS3.8. Determine the simplest formula for gold chloride if a 0.303 g sample of gold
chloride reacts with AgNO3 to produce gold nitrate and 0.430 g of AgCl.
HINT: Given that 0.430 g of AgCl are formed in the reaction one must
assume all the chloride in 0.430 g of AgCl comes from the gold chloride.
Problem Set #4
ALL work must be shown to receive full credit.
PS4.1. Gallium metal reacts with hydrogen chloride according to the equation
a) How many grams of hydrogen chloride are required to completely react with
24.4 g of gallium metal?
b) b) How many grams of gallium chloride are formed?
PS4.2. How many grams of gallium chloride are formed when 24.4 g of gallium metal
reacts with 40.0 g of hydrogen chloride?
PS4.3. In the formation reaction
Calculate the number of moles of C2H6 formed when
a) 2.0 moles of C are reacted with 5.0 moles of H2
4b. 6.0 moles of H2 are reacted with 4.0 moles C
c) 9.0 moles of H2 are reacted with 5.0 moles of C
d) 0.0812 moles of C reacted with 0.125 moles of H2
PS4.4. In the reaction
20.0 g of Hg are reacted with 5.00 g of O2. Calculate the maximum amount, in
grams, of HgO formed.
PS4.5. In the reaction
30.0 g of NH3 are reacted with 30.0 g of O2. Calculate the mass of H2O and
NO formed. What mass of the excess reactant remains unreacted?
PS4.6. A common laboratory method for determination of arsenic is described in the
reaction,
a) If 2.50 g of As2O3, 4.50 g of I2 and 4.00 g of H2O are mixed, and the
reaction proceeds to completion, which reactant is the limiting reagent?
6b) Calculate the mass of As2O5 which is theoretically possible.
c) If only 1.80 g of As2O
reaction. 5 is actually produced, determine the percent yield in the
PS4.7. An industrial method for the production of elemental zinc from its ore is to
'roast' the crude ore in oxygen and then react it with CO. The following
reactions describe the chemical process,
A 10.0 kg sample of an ore containing ZnS was chemically treated according to
the reactions described above producing 2.85 kg of Zn. Determine the
percentage of ZnS contained in the original sample of the ore.
Problem Set #5
ALL work must be shown to receive full credit.
PS5.1. Calculate the molarity of a solution containing,
a) 23.9 g of NaCl dissolved in enough water to have 300 mL of solution
b) 1.69 x 10-3 moles HCl in 12.0 mL of water
PS5.2. Describe how to prepare the following solutions.
a) 450 mL of 0.500 M NaOH, from solid NaOH
b) 900 mL of 1.00 M HCl, from 12.0 M HCl (concentrated HCl)
PS5.3. In the reaction,
How many grams of magnesium will react with 250. mL of 0.500 M HCl?
PS5.4. How many grams of H2 will be produced when 1.00 g of Mg is added to a
250. mL sample of 0.500 M HCl? (Use the chemical reaction in 5.3.)
PS5.5. Calculate the volume of 0.360 M NaOH needed to neutralize 27.0 mL of 0.820
M HCl. (Note: Write the chemical equation.)
PS5.6. Calculate the concentration of phosphoric acid produced when 2.00 g of P2O5
is added to 500 mL of distilled water. (Assume the final volume of the solution
is 500 mL.) The important reaction is,
How many moles of water remain unreacted?
PS5.7. Potassium dichromate in acidic solution is frequently used to determine the
concentration of Fe(II) in solution. The equation which describes the reaction
is:
A solution of Cr2O72- is prepared by dissolving 6.425 g of K2Cr2O7 in 800 mL of water. (Assume no
significant change in volume when the solution is prepared.) A total of 21.35 mL of this solution is
required to reach the end-point in a titration of a 250.0 mL sample containing Fe(II). Determine the
concentration of Fe(II) in the solution.
Problem Set #6
ALL work must be shown to receive full credit.
PS6.1. Rank the following substances from lowest specific heat to highest specific heat.
C(s), Fe(s), N2(g), H2O(l), Hg(l)
PS6.2. Rank the following substances from lowest heat capacity to highest heat capacity.
1 kg H2O(l), 10 kg Cu(s) cube, 5 kg Fe(s) ball, 5 kg C(s) rod
PS6.3. Below is a table of specific heats for several different metals
Calculate the molar heat capacity
for each of the metals. Based on the results
of your calculations, estimate the specific heats of iron and silver.
PS6.4a. How much heat must be absorbed by a 75.0 g sample of water for the temperature to change
from 24.5 ºC to 63.8 ºC?
b. Calculate the temperature change if a 75.0 g sample of zinc absorbed the same amount of heat. The
specific heat of zinc is listed in problem 6.3.
PS6.5. A 28.4 g sample of an unknown metal was heated to 110.0 ºC and plunged into a 100 g sample
of water initially at a temperature of 24.60 ºC. The final temperature of the mixture was 25.34 ºC.
Calculate the specific heat of the metal. Identify the metal.
PS6.6. When 1.000 g of KNO3 is dissolved in 120.0 g of water initially at 24.25 ºC in a coffee-cup
calorimeter, the final temperature is found to be 23.44 ºC. Calculate the heat absorbed per gram and per
mole when KNO3 dissolves in water. (Assume the heat capacity of the calorimeter is zero.)
PS6.7. Calculate the heat produced per mole of aspirin when 2.216 g of C9H8O4 are reacted with excess
oxygen in a bomb calorimeter containing 4.40 kg of water. The temperature change measured is 2.32 ºC.
The heat capacity of the calorimeter is 2340 J/ ºC) .
Problem Set #7
ALL work must be shown to receive full credit.
is -683 kJ. Calculate the heat produced when 2.57 g of CaBr2 are formed.
PS7.3. For which of the following reactions is
PS7.4. Given the enthalpy change for the two reactions
PS7.5. Given the following equations;
reasoning in each case.
PS7.6. Using a table of Standard Enthalpies of Formation in your text or another reference book,
calculate the enthalpy of reaction for each of the following;
PS7.7. The molar heat of combustion of nitroethane, C2H5NO2(l), to CO2(g), H2O(l) and N2(g) at 25 ºC is
-1348 (kJ/mol). Determine the
for nitroethane.
PS7.8a. Given the following thermodynamic data,
Calculate
exothermic?
for 1 mol of CaCl2 dissolving in water. Is the dissolution of CaCl2 endothermic or
b. If 18.0 g of CaCl2 is added to 100. g of water initially at 23.5 ºC, calculate the temperature of the
solution after the CaCl2 dissolves. (Assume no heat is lost to the container or the surroundings and the
specific heat of the solution is the same as that of water.)
PS7.9. Determine the standard enthalpy of vaporization (transition from liquid to gas) for CCl4(l).
Problem Set #8
ALL work must be shown in all problems for full credit.
PS8.1. Consult the periodic table and arrange the following pure solid elements in order
of increasing electrical conductivity at room temperature: Ca, Sn, S, and Si.
Explain the reason for the order you chose.
PS8.2. a) Arrange the following series of element in order of increasing metallic
character: Si, Sn, C, Ge, Pb.
b) Arrange the folllowing series in order of inreasing nonmetallic character:
As, P, Bi, Sb, N
PS8.3. Write the formula for the compound formed by the combination of each of the
following pairs of elements.
a) magnesium and iodine
b) gallium and oxygen
c) potassium and sulfur
PS8.4. Which of the following compounds are ionic and which are molecular:
N2O, Na2 O, CaO, CO, P2O5 , Cl2O7 , Fe2O3 ? Explain your reasoning.
PS8.5. Predict the products and write balanced chemical equations for each of the
following pairs of reactants.
a) Hydrochloric acid and nickel
b) sulfuric acid and iron
c) hydrobromic acid and zinc
d) acetic acid (HC2H3O2) and magnesium
PS8.6. Based on the acitivitiy series, predict the products of each proposed reaction.
Write a balanced chemical equation for each reaction that occurs.
PS8.7. a) Use the following reactions to prepare an acitvity series for the halogens:
PS8.7b. Relate the positions of the halogens in the eriodic table with their locations
in this activity series.
c) Predict whether the following reactions will occur.
PS8.8. Write balanced chemical equations for the reactions that occur when the
following substance are mixed.
a) sodium oxide and water
b) copper (I) oxide and nitric acid
c) sulfur trioxide and water
d) Selenium odioxide and sodium hydroxide
PS8.9. Write the symbol including the charge for each of the following ions:
a) selinide
b) astinide
c) phosphide
d) cesium ion
PS8.10. From the following list of elements select the one that best fits each description:
oxygen
sulfur
lithium
silver
iodine
germanium
aluminum
a) An acive metal that forms a cation with a +1 charge.
b) A grayish-black solid that readily forms a purple vapor.
c) A metal that dissolves in NaOH solution.
d) A metalloid.
e) A yellow nonmetal.
Problem Set #9
ALL work must be shown in all problems for full credit.
PS9.1. In the wave form shown below label the wavelength and amplitude. Describe how the frequency
of the wave can be determined.
PS9.2. The double line in the sodium emission spectrum have wavelengths of 5.896 x 10-5 cm and 5.890
x 10-5 cm (589.6 and 589 nm). Calculate the frequency of each line.
PS9.3. The wavelengths of the sodium doublet can be accurately measured despite the fact that they are
very short. The IUPAC standard for the meter is, in fact, based on the wavelength of a particular line in
the emission spectrum of the element krypton. According to the standard, one meter is defined as
1,650,763.73 wavelengths. What color is this line?
PS9.4. Briefly describe the difference between the continuous model of matter and quantized view of
matter.
PS9.5. Which color of light is higher energy, blue of red? Explain.
PS9.6. An argon laser emits light with a wavelength of 488 nm. Calculate the energy of a photon of light
emitted from the laser.
PS9.7. The wavelength of a photon of light capable of breaking a carbon-carbon single bond is 344 nm.
What area of the electromagnetic spectrum is light of this wavelength located? Calculate the energy
required to break 1 mol of C-C single bonds.
PS9.8. Calculate the energy and the wavelength of light emitted when an electron in an excited hydrogen
atom falls from the n = 5 level to the n = 1 level.
PS9.9. An electron initially in the n = 2 energy level in a hydrogen atom absorbs a photon of light with a
frequency of 6.167 x 1014 s-1. Calculate the new energy level the electron will occupy.
PS9.10. Will a photon of light of wavelength 480 nm excite an electron in the hydrogen atom from the n
= 1 level to the n = 2 level? Explain.
Problem Set #10
ALL work must be shown in all problems for full credit.
PS10.1. How do electrons shells, subshells and orbitals differ?
PS10.2. What are the possible values of l, and ml for the n = 2 shell?
PS10.3. What are the set of quantum numbers for an electron in a 2s orbital? A 3d sublevel? A 6f
orbital?
PS10.4. Which of the following sets of quantum numbers describe an electron in an excited state in a
hydrogen atom? Which describe a ground state? Which are not allowed?
a) n = 3; l = 3; mi = 0
b) n = 2; l = 1; mi = -1
c) n =1 ; l = 0; mi = 0
d) n = 2; l = 0; mi = +1
e) n = 5; l = 2; mi = +2
PS10.5. How many orbitals are available in each of the following subshells or shells?
a) n = 3 shell
b) 2p subshell
c) 4d subshell
d) n = 6 shell
PS10.6. Sketch a 1s subshell and a 2s subshell. What are the major similarities and differences between
the two subshells.
Problem Set 11
ALL work must be shown in all problems for full credit.
PS11.1. Of the set of four quantum numbers, n, l, ml, ms , which pair determines the energy of an
electron in an orbital in a many electron atom?
PS11.2. Order the following orbitals from highest to lowest energy.
4s, 2p, 4d, 3p, 4f
PS11.3. Explain the term degenerate, as it relates to orbitals, and explain how adding electrons to a set
of degenerate orbitals differs from adding electrons to orbitals which are not degenerate.
PS11.4. List the possible values of the four quantum numbers for each electron of the valence electrons
in Br.
PS11.5. Draw the orbital diagrams for S and Mo.
PS11.6. Write the electron configuration for Mg, Kr, Rh, Pt and Bi. Use the periodic table.
PS11.7. Briefly explain each of the following statements.
a) The atomic radius of Mg is smaller than Ba.
b) The ionization energy of S is lower than P.
PS11.7c) The energy required to remove the fourth electron in Al is significantly larger than the third
electron.
d) The energy change associated with gaining an electron, the electron affinity, is positive for some
atoms and negative for others.
PS11.8. Predict the products of the following reactions.
a) A sample of potassium is added to water at 25 ºC
b) Mg(s) + S(s)
c) Sodium solid is dropped into a sample of bromine
d) Ba(s) + O2(g)
Problem Set #12
ALL work must be shown in all problems for full credit.
PS12.1. Write the electron configurations for each of the following atoms or ions.
a) Na
b) Mn2+
c) Brd) Kr
e) Pb2+
PS12.2. Predict the formula of the ionic compound formed between the following pairs of elements.
a) Na and Br2
b) Al and O2
c) Ba and S
d) Fe and Cl2
PS12.3. Explain why Mg2+ is smaller than S2-. Explain why Mg is larger than S.
PS12.4. Which of the following species form an isoelectronic group?
N3-, Cl-, Ne, Mg2+, Se2-, H+
PS12.5. Which of the following salts has the largest lattice energy? Explain.
LiF, LiCl, LiBr, LiI
PS12.6. Write the equation which describes the reaction which is associated with the lattice energy of an
ionic compound such as MgO.
PS12.7. Predict whether the following compounds are ionic or covalent.
SiCl4, MgBr2, PH3, NH4Cl, HCl, Al2O3
PS12.8. Arrange the following elements from smallest to largest electronegativity.
O, Al, Ga, I, H, Na
PS12.9. Write the Lewis structures for the following ions or molecules.
a) HBr
b) PCl3
c) SO32-
9d. ClO3-
e) C2H4
f) CH2Cl2
g) Cl2CO
h) HCN
Problem Set #13
ALL work must be shown in all problems for full credit.
PS13.1. Write the Lewis structures for the following ions or molecules. (If a molecule cannot be
adequately represented by a single diagram, include all resonance structures.)
a) CO32-
b) NO2-
c) HCO2-
d) N2O
PS13.2. Use bond dissociation energies to estimate the
for HCl(g).
PS13.3. Use bond dissociation energies to estimate the enthalpy of the reaction for
PS13.4. Use bond dissociation energies to estimate the enthalpy of the reaction for
PS13.5. Use bond dissociation energies to estimate the enthalpy of the reaction for
PS13.6. Determine the oxidation state of the boldfaced elements in each of the following:
a) Ca3P2
b) SO42c) K2Cr2O7
d) Na2O2
e) FePO4
PS13.7. Give the name or chemical formula for each of the following substances:
a) lead(II) nitrate
b) dinitrogen pentoxide
c) Na2CO3
d) FeCl3
e) HCl
f) P4O6
g) chromium(III) sulfate
PS13.8. Predict the products of the following reactions.
Problem Set #14
ALL work must be shown in all problems for full credit.
PS14.1. Complete the following table
PS14.2. Use simple structure and bonding models to account for each of the following.
a). The H-N-H bond angle is 107.5 º in NH3.
b). The I3- ion is linear.
PS14.3. Which of the molecules listed in Problem #14.1 above are polar and which are nonpolar?
PS14.4. Consider theLewis structure for glycine, the simplest amino acid:
a). What are the approximate bond angles about each of the two carbon atoms, and what are the
hybredizations of the orbitals on each of them?
b). What are the hybridizations of the orbitals on the two oxygens and the nitrogen atom, and what are
the approximate bond angles about the nitrogen?
c). What is the total number of s bonds in the entire molecule? Sigma () bonds?
PS14.5. Indicate the hybridzation on the central atom for each of molecules in Problem 14.1 above.
PS14.6. Indicate the atomic orbitals on each atom in the following molecules which are involved in
forming the covalent bond.
a). H-F
b). F2
c).
PS14.7a. Draw two resonance structures for the nitrite ion, NO2-.
b). What is the hybridization around the nitrogen atom in this polyatomic ion?
c). Describe the pi-bonding in the polyatomic ion.
d). Label both the sigma and the pi bonds as either localized or delocalized.
PS14.8. Use simple structure and bonding models to account for the following. The bond lengths in
CO32- are all identical and are shorter than a carbon- oxygen single bond.
Problem Set #15
AP Chemistry by Satellite
ALL work must be shown in all problems for full credit.
PS15.1. The pressure of the mercury in the gas (vapor) phase above the liquid mercury in a barometer is
2.0 x 10-3 mmHg. Calculate the pressure in units of atmospheres.
PS15.2. A sample of SF6 occupies a container of variable volume. If the the sample originally occupies a
volume of 50.0 mL at 755 mmHg, calculate the pressure which must be exerted to lower the volume to
25.0 mL.
PS15.3. If the pressure of a sample of an ideal gas, initially at 1.25 atm, is tripled, by what factor will
the volume of the gas change?
PS15.4. A 545 mL sample of nitrogen gas initially at -220 ºC is heated to 100 ºC. Calculate the new
volume, assuming the pressure does not change.
PS15.5. Calculate the volume of 0.390 moles of an ideal gas at 750 mmHg and 23 ºC.
PS15.6. Calculate the density of SF6 at 1.00 atm and 0.00 ºC.
PS15.7. Calculate the volume of a sample of helium at -33.0 ºC and 1.23 atm if it occupies a volume of
2.34 L at 54.5 ºC and 1026 mmHg.
PS15.8. A 0.751 mol sample of an ideal gas occupies a 10.0 liter flask at 27.0 ºC and 1.85 atm. If 0.257
mol of the gas are removed from the container, calculate the new pressure. (Assume the temperature
remains constant.)
PS15.9. Which of the following contains the largest number of particles?
a) 5.00 g of He at 1.00 atm and 0 ºC
b) 225 g of Au
c) 34.5 L of an ideal gas at -5.0 ºC and 2000 mmHg
PS15.10. Which of the following samples has the greater mass?
a) 278 L of Ar at 25 ºC and 300 mmHg
b) 225 mL of CH4 at 300 ºC and 5.34 atm
c) 34.5 L of Cl2
Problem Set #16
AP Chemistry by Satellite
ALL work must be shown in all problems for full credit.
PS16.1. The first laboratory preparation of O2, performed by Joseph Priestly, required the heating of
mercury (II) oxide,
What volume of O2 at 22.5 ºC and 753 mmHg will be produced when 15.3 g of HgO are completely
decomposed?
PS16.2. Hydrogen, H2, can be prepared by passing steam through a hollow air tube which has been
heated to high temperature. The reaction is,
Calculate the volume of H2 formed at 0.98 atm and a temperature of 450 ºC when 98.3 g of H2O are
passed through an iron tube.
PS16.3. Calculate the volume of SO2(g) produced when 2.00 liters of O2 react with excess sulfur at
constant temperature and pressure.
PS16.4. A gaseous mixture contains 3.00 g of N2, 0.430 moles of Ar and 2.15 x 1023
molecules of CH4. If the total pressure of the mixture is 3.00 atm, calculate the partial pressure of each
component.
PS16.5. Experimentally oxygen gas is frequently collected over water. When a sample of oxygen,
produced from a chemical reaction, is collected over water at 24 ºC, the total pressure was 765 mmHg.
Calculate the pressure due only to oxygen.
PS16.6. A sample of oxygen collected over water at 29 ºC exerts a total pressure of 759 mmHg. If the
volume of the container is 125 mL calculate the mass of oxygen present.
PS16.7. Can the speed of a molecule of a gas be doubled if the temperature of the gas sample is held
constant, yes or no? Briefly explain your answer.
PS16.8. Explain, in terms of the kinetic molecular model, why increasing the volume of a sample of an
ideal gas decreases the pressure of the gas at constant temperature.
PS16.9. Rank the following gases in order of rate of diffusion. Explain your order.
Ne, CH4, SF6, CO2
PS16.10. Describe the two factors responsible for the deviation of real gases from ideal behavior as
assumed in the ideal gas equation.
Problem Set #17
ALL work must be shown in all problems for full credit.
PS17.1. a) How much heat is produced when 75.0 g of steam at 135 ºC is converted to water at 20.0 ºC?
b) How much heat is required to convert 30.0 g of ice at -10.0 ºC to steam at 105.0 ºC?
PS17.2. For many years drinking water has been cooled in hot climates by the evaporation of water from
the surface of canvas bags or porous clay pots. How many grams of water can be cooled from 35 ºC to
20 ºC by the evaporation of 1.0 g of water?
PS17.3. Ethyl alcohol melts at -114 ºC and boils at 78 ºC. The enthalpy of vaporization for ethyl alcohol
at 78 ºC is 870 (J/g) and the enthalpy of fusion is 109 (J/g) . If the heat capacity of solid ethyl alcohol is
taken to be 0.97 (J/g.ºC), and that for the liquid 2.3 (J/g.ºC), how much heat is required to convert 10.0 g
of ethyl alcohol at -120 ºC to the vapor phase at 78 ºC?
PS17.4. A sample of water vapor in a flask of constant volume exerts a pressure of 530 mm Hg at 100
ºC. The flask is slowly cooled.
a) Assuming no condensation, use the Ideal Gas Law to calculate the pressure of the vapor at 90 ºC; at
80 ºC.
b) Compare your answers to the equilibrium vapor pressure of water at 90 ºC and at 80 ºC. Will
condensation occur at 90 ºC; 80 ºC?
c) On the basis of your answers in a) and b), predict the pressure exerted by the water vapor at 90 ºC; at
80 ºC.
PS17.5. Consider the following data for the vapor pressure, Pv, of acetic acid, CH3COOH(l) (use graph
paper):
a) Plot ln (Pv) vs. 1/T and use your graph to estimate the heat of vaporization of acetic acid.
For parts b) and c) use the Clausius-Clapeyron equation.
b) Determine the temperature of a sample of acetic acid when the vapor pressure is 42.4 mmHg.
c) Determine the vapor pressure of acetic acid of a sample of liquid acetic acid at 13.5 ºC.
PS17.6. The vapor pressure of bromine at 9 ºC is 113 mm Hg; at 20 ºC it is 184 mm Hg. Using the
Clausius-Clapeyron equation estimate the
a) heat of vaporization of bromine;
b) vapor pressure of bromine at 50 ºC.
PS17.7. The vapor pressure of mercury is 17.3 mm Hg at 200 ºC; its heat of vaporization is 59.4 kJ/mol.
Use the Clausius-Clapeyron equation
a) the vapor pressure of mercury at 340 ºC.
b) the normal boiling point of mercury.
Problem Set #18
ALL work must be shown in all problems for full credit.
PS18.1. Indicate what change, if any, should occur in each of the following properties as a result of an
increase in the strength of intermolecular forces:
a) vapor pressure;
b) normal boiling point;
c) normal melting point
d) surface tension;
e) viscosity;
f) heat of fusion;
g) heat of vaporization;
h) molecular weight.
PS18.2. Indicate all the various types of intermolecular attractive forces that may operate in each of the
following:
a) CH3OH(l);
b) Xe(l);
c) H2S(l);
d) ClF(l)
e) Ca(NO3)2(s)
PS18.3. What are the required structural features for a substantial hydrogen-bonding contribution to the
intermolecular attractive forces?
PS18.4. List the following compounds in the expected order of increasing energy of the hydrogenbonding interaction between molecules:
H2S; CH3NH2; C6H5OH (phenol).
PS18.5. For each of the following pairs of substances predict which will have the higher melting point
and indicate why:
(a) CuBr2 > Br2
(b) CO2 < SiO2
(c) S < Cr
(d) CsBr < CaF2
PS18.6. Indicate the type of crystal (atomic, molecular, metallic, covalent, or ionic) each of the
following would form upon solidification:
(a) O2
(b) H2S
(c) Ag
(d) KCl
(e) Si
(f) Al2(SO4)3
(g) Ne
(h) SiO2
(i) NH3
(j) MgO
(k) NaOH
(l) CH4 .
PS18.7. Sketch the layer structure arrangement for the packing arrangements; simple cubic, bodycentered cubic, hexagonal closest-packed and cubic closest- packed.
PS18.8. Sketch a simple cubic unit cell and explain why only one atom is contained in a simple cubic
unit cell.
PS18.9. Using the phase diagram in problem 18.7., determine the physical state of water at
a) 900 mmHg and 40 ºC
b) 500 mmHg and 30 ºC
c) 300 mmHg and 90 ºC
PS18.10. In the phase diagram for water shown below;
a) At 400 mmHg what is the approximate temperature needed to convert water from a solid to a liquid?
b) What is the approximate pressure at which water changes from a liquid to a gas at 80 ºC?
Problem Set #19
ALL work must be shown in all problems for full credit.
PS19.1. Predict whether in the following pairs (solute:solvent) a solution is formed. In each case explain
why or why not.
a) HCl(g):H2O(l)
b) CH3COOH(l):H2O(l)
c) CH3OH(l):C7H8 (toluene)(l)
d) C6H14(l):H2O(l)
e) I2(s):C7H8 (toluene)(l)
f) Br2(l):H2O(l)
g) NaNO3(s):H2O(l)
h) NaClO4(s):CCl4(l)
PS19.2. Predict whether the following compounds are soluble or insoluble in water.
a) Na2SO4(s)
b) C12H22O11(s)
c) CoCl2(s)
d) CH3OH(l)
e) C7H8(l)
f) HF(l)
g) NH3(g)
h) (CH3)2CO(l)
PS19.3. For those compounds in 19.2 that are soluble write the equation which describes the
compound's behavior when added to water.
PS19.4. Describe the attractive forces present when KI(s), NH3(g) and CH3CH2OH(l) dissolve in water.
Prepare sketches depicting at the atomic level how each of these substances interact with water
molecules.
PS19.5. If 52.6 g of AgNO3 are added to 684 mL of water, determine;
a) how many grams of AgNO3 (solute).
b) how many grams of H2O (solvent). (Density of pure water is 1.00 g/mL )
c) how many grams of solution.
d) how many moles of AgNO3.
e) how many moles of H2O.
f) how many moles of solution.
g) how many milliliters of solution, if the density is 1.106 g/mL.
h) the weight percent AgNO3.
i) the mole fraction of AgNO3.
j) the molality of the solution.
k) the molarity of the solution.
PS19.6. Given that an aqueous solution, which weighs 367 g, is 40.0 % (ethylene glycol) C2H6O2 by
mass and has a density of 1.05 g/mL, determine;
a) the mole fraction of ethylene glycol.
b) the molality of the solution.
c) the molarity of the solution.
PS19.7. An aqueous solution of cesium chloride is 5.94 molal and has a density of 1.58 g/mL. Calculate
the
a) weight percent cesium chloride.
b) mole fraction of cesium chloride.
c) molarity of the solution.
PS19.8. Describe how you would prepare;
a) 520 mL of a 0.760 M NaCl solution.
b) 37 g of a 15.2 % (by weight) solution of NaCl.
c) 210. g (grams of solution) of a 0.185 molal NaCl solution.
PS19.9. A 10.00 % (by weight) solution of K2SO4 in water was found to be 0.574 M by titration at 25
ºC. Calculate
a) the molality of the solution
b) the density of the solution
Problem Set #20
ALL work must be shown in all problems for full credit.
PS20.1. Calculate the vapor pressure for each of the following solutions at 25 ºC; a) 16.8 g of urea
(NH2)2CO dissolved in 108 g of water.
b) 8.54 g of MgCl2 dissolved in 108 g of water.
PS20.2. To what temperature (ºC) would a solution containing 150 g of glycerol, C3H5(OH)3 (assume to
be a nonvolatile solute), in 100 g of water have to be heated to have a vapor pressure of 91.1 mmHg?
PS20.3. Determine the freezing point of a solution which is 0.50 molal urea (a nonelectrolyte).
Determine the boiling point of this solution.
PS20.4. A solution containing a nonelectrolyte dissolved in water has a freezing point of -1.62 ºC.
Calculate the boiling point of the same solution.
PS20.5. Given the following data;
a) If each of the solutions is prepared by adding 1 mole of compound to 1 kg of water why does each
have a different Tf ?
b) Predict the ideal Tf for the above compounds.
b) Predict the ideal Tf for the above compounds.
c) Why does the ideal Tf differ from the experimental Tf?
d) Classify each compound as a strong, weak or nonelectrolyte.
PS20.6. Determine the ideal freezing point of a solution prepared by mixing 3.53 g of MgCl2 in 430 g of
water.
PS20.7. Calculate the "molecular formula" of solid sulfur if the freezing point of carbon tetrachloride is
lowered by 0.28 ºC when 0.24 g of sulfur is added to 100 g of carbon tetrachloride. (Note: Kf(CCl4) =
29.8 ºC/m .)
PS20.8. When 7.20 g of HBr are added to 100 g of water, the freezing point of the solution is lowered by
3.30 ºC. How can you explain this temperature change? Is HBr a strong, weak or nonelectrolyte?
PS20.9. 5.76 g of an ionic compound with the formula K3X is dissolved in 350 g of water. The freezing
point of the solution was lowered by 0.370 ºC. Determine the formula weight of X.
Problem Set #21
ALL work must be shown in all problems for full credit.
PS21.1. The following data was collected for the reaction
a) Plot the data for Exp. #1 and graphically estimate
i) the initial rate
ii) the instantaneous rate at 100 sec? 750 sec? 1350 sec?
iii) the time it takes for half of the HI to react
b) Repeat a) for Exp #2
i) the initial rate
ii) the instantaneous rate at 100 sec? 750 sec? 1350 sec?
iii) the time it takes for half of the HI to react
c) By what factor did the initial concentration change in going from Exp #1 to Exp #2?
d) By what factor did the initial rate change in going from Exp #1 to Exp #2?
PS21.1. (Continued)
e) What is the order of the reaction with respect to HI?
f) How did the half-life change for the two experiments?
g) Determine the rate constant for the reaction including units.
h) What would the initial rate be if the concentration of HI is 0.654 M? 1.25 x 10-2 M?
PS21.2. The following initial rate data were collected for the reaction
at 25 ºC.
a) Determine the reaction order for A2 and B.
b) Determine the overall order of the reaction.
c) Write the specific rate law for the reaction.
d) Determine the rate constant for the reaction (include units).
PS21.3. The following initial rate data were collected for the reaction
a) Determine the reaction order for NO2 and F2.
b) Determine the overall order of the reaction.
c) Write the specific rate law for the reaction.
d) Determine the rate constant for the reaction (include units).
Problem Set #22
AP Chemistry by Satellite Name___________________________________
ALL work must be shown in all problems for full credit.
PS22.1. The reaction
NO2(g)
N2O4(g)
follows simple second order kinetics. If the [NO2]0 is 0.156 M,
a) calculate the rate constant for the reaction if it takes 1.00 x 10-3 s for the concentration of NO2 to fall
to 0.147 M.
b) calculate the half-life for the reaction. (When the [NO2]0 = 0.156 M.)
c) how long will it take for the [NO2] to fall to 5.00 x 10-2 M?
d) what is the [NO2] after 1.00 s? (When [NO2]0 = 0.156 M.)
PS22.2. The reaction
follows simple first order kinetics with a half-life of 12.4 s.
a) Calculate the rate constant for the reaction.
PS22.2. (Continued)
b) How long will it take for the [H2O2] to fall from 0.300 M to 0.0452 M?
c) What is the [H2O2] after 30 minutes if [H2O2]0 = 1.25 M?
d) How long will it take for the [H2O2] to decrease by a factor of 6?
PS22.3. C4H8 decomposes according to the following equation;
the rate constant for the decomposition is 6.07 x 10-10 sec-1 at 25 ºC.
a) What is the order of the reaction?
b) How long would it take for 1.00 % of a sample of C4H8 to decompose at 25 ºC and 1 atm?
PS22.3b. (Continued)
c) What is the half-life of the reaction?
d) How long would it take for 1.00 % of a sample of C 4H8 to decompose at 25 ºC and 10 atm?
PS22.4. The second-order decomposition of nitrous oxide, N2O, has a half-life of 75.0 min at 900 K
when the initial concentration of N2O is 2.00 x 10-2 M.
a) What is the concentration of nitrous oxide after 150 minutes?
b) How long will it take for 40.0 % of the sample to decompose?
PS22.5. The first-order rate constant for the reaction
is 4.00 x 10-4 sec-1 at 573 K.
a) What will be the concentration of CH3N2CH3 after 600 seconds, given that the initial concentration is
1.03 x 10-2 M?
b) What is the half-life of the reaction for this initial concentration?
PS22.6. In the reaction
the [N2O] was followed with time and the data shown below was obtained.
Determine the order of the reaction and its half-life.
Problem Set #23
AP Chemistry by Satellite Name___________________________________
ALL work must be shown in all problems for full credit.
PS23.1. The activation energy for the decomposition of N2O5 is 102 (kJ/mol). The rate constant for the
reaction at 45 ºC is 5.00 x 10-4 M-1·sec-1. What is the value of the rate constant at 65 ºC?
PS23.2. Using the data in PS23.1, calculate the temperature at which the rate constant is 3.00 x 10-6 M1
·sec-1.
PS23.3. Using the data in PS23.1, calculate the rate constant at 0 ºC.
PS23.4. A chemist was able to determine that the rate of a particular reaction at 100 ºC was four times
faster than at 30 ºC. Calculate the approximate energy of activation for such a reaction.
PS23.5. Data for the reaction
was collected and is shown below. Plot the data and determine the activation energy for the reaction.
PS23.6. Consider the simple reaction,
Determine what the order of the reaction must be for each statement to be true. a) The initial
concentration of A is doubled and the initial rate increase by a factor of four.
b) The half-life for the disappearance of A is inversely proportional to the initial concentration of A.
PS23.7. Given the following reaction mechanism
What is the overall reaction? Write the rate law for the reaction.
PS23.8. If the reaction
occurred via a one step mechanism, draw a picture of activated complex.
Discuss how likely a one-step mechanism would be for this reaction.
PS23.9. Suggest a two step mechanism for the reaction
if the experimental rate law is rate = k[NO][Cl2].
PS23.10. The suggested mechanism for the reaction between peroxide and iodide ion is,
Identify a specie(s), if any, which is acting as a catalyst and a specie(s) which is acting as an
intermediate.
Problem Set #24
AP Chemistry by Satellite Name___________________________________
ALL work must be shown in all problems for full credit.
PS24.1. Given the reaction
Initially (before any reaction occurs) a 1.00 liter reaction vessel at 400 ºC contains 0.502 moles of O2
and 0.791 moles of NH3 and no water or nitrogen. Consider the following:
a) If 0.0873 moles of O2 react, how many moles of NH3 must react and how many moles of H2O and N2
are formed? How many moles of O2, NH3, H2O and N2 remain after completion of the reaction?
b) If 0.234 moles of NH3 react, how many moles of O2 must react and how many moles of H2O and N2
are formed? How many moles of O2, NH3, H2O and N2 remain after completion of the reaction?
c) If '3x' moles of O2 react, how many moles of NH3 must react and how many moles of H2O and N2 are
formed(in terms of 'x')? How many moles of O2, NH3, H2O and N2 remain after completion of the
reaction?
d) If 0.875 moles of H2O are formed, how many moles of N2 are formed and how many moles of O2 and
NH3 must react? How many moles of O2, NH3, H2O and N2 remain after completion of the reaction?
PS24.2. Write the equilibrium expression for each of the following chemical equations;
PS24.3. Equilibrium constants for the following reactions have been determined at 550 ºC:
Calculate K (at the same temperature) for the commercially important water gas shift reaction
PS24.4. Calculate Kc for the reaction
if Kc for the reaction
is 1.3 x 104.
PS24.5. A 1.00 liter container initially holds 0.257 moles of NOBr at a given temperature. The reaction
which occurs is:
At equilibrium analysis shows 0.240 moles of NO and 0.120 moles of Br2.
a) Which direction did the reaction proceed to establish (reach) equilibrium?
b) How many moles of NOBr reacted in order to form 0.240 moles of NO and 0.120 moles of Br2?
c) How many moles of NOBr remain after equilibrium was established?
d) What is the magnitude of Kc?
PS24.6. In a container, the partial pressure of NOCl is initially 0.340 atm at a given temperature. The
reaction which occurs is:
At equilibrium analysis shows the partial pressure of NO is 0.0916 atm.
a) Which direction did the reaction proceed to establish (reach) equilibrium?
b) What is the partial pressure of NOCl which reacted in order for the partial pressure of NO to be
0.0916 atm?
c) What is the partial pressure of Cl2 at equilibrium?
d) What is the partial pressure of NOCl at equilibrium?
e) What is the magnitude of Kp?
PS24.7. A 1.00 liter container holds 1.06 moles of H2 and 1.57 moles of CO at a temperature of 162 ºC.
At this temperature, the following reaction occurs,
After equilibrium is established, analysis shows 0.200 moles of CH3OH in the container. Calculate the
[CO]eq, [H2]eq and Kc.
PS24.8. The following reaction,
occurs at 298K. If 2.00 mol of HI are placed into a 1.00 liter container and permitted to react, at
equilibrium it is found that 20.0 % of the HI has decomposed. Calculate Kc and Kp.
PS24.9 The equation which describes the preparation of ammonia is:
A 3.000 L reaction vessel initially contains 0.3000 moles N2 and 0.4500 moles H2. When the reaction is
allowed to attain equilibrium at a given temperature, analysis determines 0.09992 M N2. Calculate Kc
for the reaction.
Problem Set #25
AP Chemistry by Satellite Name___________________________________
ALL work must be shown in all problems for full credit.
PS25.1. A 0.622 gram quantity of COBr2 is sealed in a glass bulb of 0.100 L volume and heated to a
temperature of 73 ºC. At 73 ºC the COBr2 partially decomposes according to the equation
for which Kc = 0.190. Calculate the concentration of each species at 73 ºC.
PS25.2. The equilibrium constant, Kp, for the reaction
is 7.31. Calculate the partial pressure of all species at equilibrium for each of the following original
mixtures:
a) 1.0 atm of CO and 1.0 atm of H2O.
b) 1.0 atm of CO, 1.0 atm of H2O and 1.00 atm of H2.
c) 1.0 atm of H2 and 1.0 atm of CO2.
PS25.3. At 1000 K the equilibrium constant, Kc, for the reaction
is 0.833. Calculate the concentrations of all species at equilibrium when 0.200 moles of NO2 are placed
in a 2.00 L container at 1000 K.
PS25.4. At 25 ºC, 0.560 mol of O2 and 0.20 mol of N2O were placed in a 1.00 liter vessel and allowed to
react according to the equation
When the system reached equilibrium, the concentration of NO2 was found to
be 0.020 M.
a) What were the equilibrium concentrations of N2O and O2?
PS25.4. (Continued)
b) What is the value of Kc for this reaction at 25 ºC?
PS25.5. The reaction
has been carefully studied at 350 ºC and the Kc is 0.079. Which direction (left-to-right or right-to-left)
will the reaction proceed to establish equilibrium under each of the following initial conditions?
PS25.6. Consider the reaction
for which Hrxn = -1036 kJ. Predict how the [SO2] will change when the equilibrium is disturbed by;
a) Addition of O2
b) Addition of H2O
c) Addition of a catalyst
d) Increase in temperature
e) Decrease in the volume of the reaction container
PS25.7. The equilibrium constant, Kp, for the reaction
is 6.25 at 25 ºC and Hº = 34.4 kJ. Calculate the magnitude of the equilibrium constant at 50 ºC.
PS25.8. Given the reaction
A 10.0 liter vessel at 298 K initially contains a sample of XeF4 at 0.750 atm. After the reaction achieves
equilibrium, the total pressure in the vessel is 1.95 atm. Calculate Kp from this data.
PS25.9. The equilibrium constant, Kc, for the reaction
is 2.50 x 10-6 at a particular temperature. If the [N2]o = 2.00 M, the [O2]o = 1.00 M and the [NO]o = 0 M,
calculate the equilibrium concentration of all species.
PS25.4. A 0.383 gram quantity of crystalline PCl5 is sealed in a glass bulb of 0.100 L volume and heated
to a temperature of 250 ºC, at which temperature all of the PCl5 has vaporized. At 250 ºC the PCl5
partially decomposes according to the equation
If the [PCl3]eq = 1.4 x 10-2 M, calculate Kc at 250 ºC.
Problem Set #26
AP Chemistry by Satellite Name___________________________________
ALL work must be shown in all problems for full credit.
PS26.1. For aqueous solutions of the following substances, write the dissociation reaction and indicate
whether the substance behaves as an Arrhenius acid or base.
a) HF(aq)
b) HC6H5O(aq)
c) Ba(OH)2(aq)
d) LiOH(aq)
e) H2O(aq)
f) H2CO3(aq)
PS26.2. Calculate the pH and pOH in each of the following aqueous solutions. In each case, indicate
whether the solution is acidic or basic.
PS26.3. Calculate the [H+] and [OH-] in each of the following aqueous solutions.
a) pH = 7.41
b) pH = 11.0
c) pOH = 0.230
d) pOH = 7.00
e) pH = 14.9
f) pH = -0.543
PS26.4. For each of the following acids, write the formula for the conjugate base.
PS26.5. For each of the following bases, write the formula for the conjugate acid.
PS26.6. For the following compounds, write the reaction with water and indicate the Bro/nsted acids,
bases, the conjugate acid and conjugate base.
a) HCl(g)
b) NH3(g)
c) HCN(g)
d) HC6H5O(s)
e) CH3CH2NH2(l)
PS26.7. Determine the equilibrium constant for the following solutions. (Show your work clearly!)
a) 0.100 M HC2H3O2 whose pH = 2.87.
PS26.7. (Continued)
b) 0.812 M NH3 whose pH = 11.58.
c) 0.500 M B whose pH = 10.67.
d) 0.0751 M HA whose pH = 4.00.
PS26.8. Given the following substances and their initial concentration:
Answer the following,
i) identify each as an acid, base or neutral substance.
ii) list the Ka value for each acid and Kb value for each base.
iii) identify each substance as strong or weak.
iv) calculate the [H+] and the pH of each of the solutions.
v) rank all substances from strongest acid...weakest acid...neutrals..
...weakest base...strongest base.
a) 0.100 M HNO3
b) 55.5 M H2O
c) 0.100 M NaOH
d) 0.100 M C2H5NH2
e) 0.100 M HF
f) 0.100 M HNO2
g) 0.100 M CH3NH2
h) 0.100 M C5H5N
i) 0.100 M HC6H5O
j) 0.100 M Ba(OH)2
k) 0.00491 M HF
l) 0.100 M HOCl
PS26.8. (Continued)
Problem Set #27
AP Chemistry by Satellite Name___________________________________
ALL work must be shown in all problems for full credit..
PS27.1. Ascorbic acid, H2C6H6O6, is a diprotic acid.
a) The equilibrium constant for the first dissociation is Ka1 = 8.0 x 10-5. Assuming the initial
concentration of H2C6H6O6 is 0.100 M, calculate [H+] assuming only the first dissociation occurs.
b) Now consider the second dissociation equation for which Ka2 = 1.6 x 10-12. What is the initial
concentration of [HC6H6O6-]? What is the initial concentration of [H+]? Calculate the final [H+]
assuming the second dissociation occurs.
PS27.2. Calculate the pH of a 0.100 M H2S. Calculate the [S2-] in the solution.
PS27.3. Predict the products of the following neutralization reactions.
a) HCl(aq) + NaOH(aq)
b) HNO3(aq) + Ba(OH)2(aq)
c) NaOH(aq) + H2CO3(aq)
d) NH3(aq) + H2SO4(aq)
e) HC6H5O(aq) + NaOH(aq)
f) HCN(aq) + KOH(aq)
PS27.4. Given a solution containing the following ions (neglect the counter-ion for the moment), write a
reaction (with water) and indicate whether the ion acts as an acid or as a base.
a) F-(aq)
b) ClO2-(aq)
c) NO2-(aq)
d) NH4+(aq)
e) CH3NH3+(aq)
f) C2H5NH3+(aq)
PS27.5. Can you make any generalizations about the acid-base character of the ions in Problem #27.4?
If so, state them.
PS27.6. Indicate an acid and a base which could react, in a neutralization reaction, to form each of the
following salts. In some cases water will be present as another product.
a) KC2H3O2(aq)
b) KClO(aq)
c) C2H5NH3NO3(aq)
d) NH4Cl(aq)
e) KCl(aq)
f) (NH4)2SO4(aq)
PS27.7. If each salt in Problem 27.6 is added to water, indicate whether the resulting solution is acidic,
basic or neutral.
PS27.8. Calculate the pH of the following salt solutions
a) 0.355 M KClO
b) 0.777 M NH4Cl
c) 0.0345 M KCl
d) 0.411 M KC2H3O2
PS27.8. (Continued)
e) 1.00 M NaHSO4
PS27.9. In the series of oxyacids, XOH, OXOH, and O2XOH, list the acids in order of increasing acid
strength. Justify your answer.
Problem Set #28
AP Chemistry by Satellite Name___________________________________
ALL work must be shown in all problems for full credit.
PS28.1. Determine the pH for a solution containing the following substances+ .
a) 0.600 M HC3H5O3 and 0.600 M NaC3H5O3
b) 0.200 M HC3H5O3 and 0.200 M NaC3H5O3
c) 0.400 M NH4Cl and 0.811 M NH3
+ Important equilibrium constants are located in Appendix E.1 and E.2 on pages 979 and 980 of Brown and LeMay.
d) 0.200 M HBr and 1.00 M HC3H5O2
e) 0.835 M C6H5NH3NO3 and 0.255 M C6H5NH2
PS28.2. Determine the magnitude of the equilibrium constant for the following reactions
(Note: write the net ionic equation)
PS28.3. A titration is performed by adding 0.400 M KOH to 40.0 mL of 0.300 M HCl.
a) Calculate the pH before addition of any KOH.
b) Calculate the pH after the addition of 5.0, 20.0 and 29.5 mL of the base.
c) Calculate the volume of base needed to reach the equivalence point.
d) Calculate the pH at the equivalence point.
e) Calculate the pH after adding 5.00 mL of KOH past the endpoint.
f) Plot pH (y axis) versus volume of KOH added (x axis) for each calculation above. Sketch the titration
curve.
PS28.4. A titration is performed by adding 0.200 M NaOH to 24 mL of 0.350 M HC2H3O2.
a) Calculate the pH before addition of any NaOH.
b) Calculate the pH after the addition of 5.0, 25.0, and 40.0 mL of the base.
c) Calculate the volume of base needed to reach the equivalence point.
d) Calculate the pH at the equivalence point.
e) Calculate the pH after adding 5.00 mL of NaOH past the endpoint.
f) Plot pH (y axis) verses volume of NaOH added (x axis) for each calculation above. Sketch the
titration curve.
PS28.5. Calculate the pH at the equivalence point when 22.0 mL of 0.200 M hydroxylamine, HONH2, is
titrated with 0.150 M HCl.
PS28.6. Calculate the pH of a solution prepared by mixing
a) 25.0 mL of 0.212 M NaOH and 36.0 mL of 0.187 M HCl
b) 125 mL of 0.345 M KOH and 87.0 mL of 0.400 M HC3H5O2
c) 400 mL of 0.100 M NH3 and 250 mL of 0.120 M HCl
Problem Set #29
AP Chemistry by Satellite Name___________________________________
ALL work must be shown in all problems for full credit.
PS29.1. Calculate the pH of the following buffer solutions.
a) 0.250 M HC2H3O2 and 0.250 M NaC2H3O2
b) 0.00250 M HC2H3O2 and 0.00250 M NaC2H3O2
c) 0.470 M NH4Cl and 0.470 M NH3
d) 0.260 M HCN and 0.340 M NaCN
PS29.2. Specify the reagents (an acid and its conjugate base or a base and its conjugate acid) and the
concentration of each reagent needed to prepare buffer solutions having the listed pH values.
NOTE: The optimum buffer solution is one with equal concentrations of the weak acid (weak base) and
its conjugate base (conjugate acid). Under these conditions, the pH of the solution is equal to the pKa
(pKb). So the best reagent for each of the solutions below is one whose pK is equal to the pH. Since the
tables in the appendix list K values, each of the pH's must be converted to their corresponding [H+] and
compared to an equilibrium constant in Appendix E1 or E2.
a) 4.74
b) 10.81
c) 5.23
PS29.3. Determine the pH of a buffer prepared by mixing 0.475 moles of HC2H3O2 and 0.525 moles of
NaC2H3O2 in enough water to give 1.00 liter of solution.
Calculate the pH when
a) 0.0500 mol of HCl is added
b) 0.0500 mol of NaOH is added
c) 1 liter of water is added
d) 5.00 mL of a 1.00 M HNO3 solution is added
PS29.4. Calculate the pH change produced when 0.100 mol of solid NaOH is added to each of the
following buffer solutions.
a) 500 mL of 0.900 M HC2H3O2 and 0.900 M NaC2H3O2
b)500 mL of 0.550 M HC2H3O2 and 0.550 M NaC2H3O2
PS29.4.
c) 500 mL of 0.200 M HC2H3O2 and 0.800 M NaC2H3O2
d) 500 mL of 0.100 M HC2H3O2 and 0.900 M NaC2H3O2
PS29.5. Calculate the pH change produced when 0.100 mol of gaseous HCl is added to each of the
following buffer solutions.
a) 500 mL of 0.900 M NH3 and 0.900 M NH4Cl
b) 500 mL of 0.200 M NH3 and 0.800 M NH4Cl
c) 500 mL of 0.100 M NH3 and 0.900 M NH4Cl
PS29.6. Complete and balance the following reactions. Identify all products phases as either (g)as,
(l)iquid, (s)olid or (aq)ueous. If no reaction occurs, write NR.
PS29.7. Calculate Ksp for the following salts using the information provided. a) The concentration of
CrO42-(aq) in a saturated solution of Ag2CrO4 is 6.50 x 10-5 M.
b) The solubility of AgBrO3 in water is 7.2 x 10-2 g/L.
c) A sample of a saturated solution of PbSO4 contains .0262 g/L of Pb2+.
PS29.8. Calculate the solubility of the following compounds in water. (Use a table of solubility product
constants in your text or some other reference book.)
a) BaCO3
b) AuCl
c) AuCl3
d) Cu3(PO4)2
PS29.9. Calculate the solubility of;
a) BaCO3 in 0.500 M Ba(NO3)2
b) PbCl2 in 0.0250 M CaCl2
c) Cu3(PO4)2 in 0.200 M Cu(NO3)2
PS29.10. A 45 mL sample of 0.015 M calcium chloride is added to 55 mL of 0.010 M sodium sulfate. Is
a precipitate expected? Explain. (Your answer must include a calculation!)
Problem Set #30
AP Chemistry by Satellite Name___________________________________
ALL work must be shown in all problems for full credit.
PS30.1. Using the table of thermodynamic values found on page 17 - 24, calculate the Hºrxn (standard
enthalpy change) for each of the following reactions:
PS30.2. For each of the following pairs, indicate which substance you would expect to possess the larger
standard entropy:
a) 1 mol H2(g) at 298 K and 1 atm or 1 mol H2(g) at 298 K and at 10 atm.
b) 1 mol H2O(s) at 5 oC or 1 mol H2O(l) at 5 oC
c) 1 mol Br2(l) at 58.8 oC and 1 atm or 1 mol Br2(g) at 58.8 oC and 1 atm.
d) 1 mol KNO3(aq) at 30 oC or 1 mol KNO3(s) at 30 oC.
PS30.3. Predict whether the entropy change in the system is positive or negative for each of the
following processes:
PS30.4. For each reaction below, use the table of thermodynamic values on page 17 - 24 of the
Lectureguide to determine the values of Hº and Sº.
PS30.4. (Continued)
PS30.4. (Continued)
PS30.5. a) Calculate Gº for each of the reactions in problem 30.4.
PS30.5. (Continued)
b) Which of the reactions in 30.4 are spontaneous at 298 K?
c) For each of the reactions listed in b), find the temperature above or below which the reaction becomes
nonspontaneous.
PS30.5. (Continued)
d) Which of the reactions in 30.4 are nonspontaneous at 298 K?
e) For each of the reactions listed in d), find the temperature above or below which the reaction becomes
spontaneous.
PS30.6. Consider the reaction
a) The concentration equilibrium constant, Kc, for this reaction is 1.12 x 10-3 at 750 K. Would you
expect this value to increase or decrease at 300 K? Explain your answer using words, no calculations
yet!
b) Calculate Hº and Sº for the reaction from the table of thermodynamic values in your text or from
the table on page 17 - 24 of the Lectureguide.
PS30.6.(Continued)
c) Using the values obtained in b), calculate Gº for the reaction at 300 K and at 750 K.
d) Calculate Kc for the reaction at 300 K. Compare your value with your prediction in a).
PS30.7. a) If 4.00 moles each of Cl2, H2O and O2 are placed into a 2.00 L vessel at 750 K, what will be
the equilibrium concentration of HCl? (See PS30.6b.)
b) If the experiment described in a) is repeated at 300 K, what will the concentration of HCl be?
PS30.8. Under what conditions do enthalpy, entropy and free energy take on values of zero?
PS30.9. a) Given Kb for ammonia is 1.8 x 10-5 at 298 K, calculate Gº for the reaction
b) What is the value of G at equilibrium?
c) What is the value of G when the [NH3] = 0.100 M, [NH4+] = 0.100 M and [OH-] = 0.0500 M?
PS30.10. The enthalpy of combustion, Hºcomb, for oxalic acid, C2H2O4(s), is
a) Write the balanced chemical equation which describes the combustion of one mole of oxalic acid.
PS30.10. (Continued)
b) Write the balanced chemical equation which describes the standard formation of oxalic acid.
c) Using the information given above and the equations in a) and b), calculate Hºf for oxalic acid.
d) Calculate Sºf for oxalic acid and Sºrxn for the combustion of one mole of oxalic acid.
e) Calculate Gºf for oxalic acid and Gºrxn for the combustion of one mole of oxalic acid.
f) Is the formation of oxalic acid from its elements spontaneous? Is the combustion of oxalic acid at 25
ºC spontaneous?
Problem Set #31
AP Chemistry by Satellite Name___________________________________
ALL work must be shown in all problems for full credit.
PS31.1. Balance the following oxidation-reduction reactions using the half-reaction method.
PS31.2. Draw a diagram of the cells in which the following reactions occur. In each case, label the
anode and cathode, the anode and cathode electrode material, the half-reaction at each electrode, the ions
in the anode and cathode compartments and salt bridge, the direction of electron flow, and the direction
of ion movement.
PS31.3. Write the balanced chemical equation for the overall cell reaction represented in each of the
following cell notations.
(a) Sn(s)|Sn2+(aq)||Ag+(aq)|Ag(s)
(b) Zn(s)|Zn2+(aq)||(Pt(s))Cl2(g)|Cl-(aq)
(c) Zn(s)|Zn2+(aq)||Cr2O72-(aq), H+(aq), Cr3+(aq)|Pt(s)
PS31.4. Which of the following species are reduced by Ag?
Cl2, Fe3+, Pb2+, I2, NO3- (in H+)
PS31.5. Which of the following species are oxidized by nitrate in acidic solution?
Cl-, Fe2+, Cu, Au, H+
PS31.6. Select a suitable species for each of the following
(a) an oxidizing agent able to convert Sn to Sn2+, but not Cu to Cu2+.
(b) a reducing agent capable of converting H+ to H2, but not Zn2+ to Zn.
(c) an oxidizing agent capable of converting Cl- to Cl2, but not F- to F2.
Problem Set #32
AP Chemistry by Satellite Name___________________________________
ALL work must be shown in all problems for full credit.
PS32.1. Calculate Eº and determine which of the following reactions will occur in the forward direction
under standard conditions? Balance the equations in acid solution.
PS32.2. Use standard reduction potentials to predict the spontaneous reaction, if any, that occurs
between the following. If no spontaneous reaction occurs, write NR.
PS32.2. (Continued)
PS32.3. Calculate Gº and K for the spontaneous reactions in problem 32.2.
PS32.3. (Continued)
PS32.4. Calculate the solubility product constant for the iron(II) hydroxide, given the following
information.
PS32.5. Calculate Ecell of the following cells at 25 ºC.
(a) Cr(s)|Cr3+(1.0 x 10-2 M)||Ni2+(2.0 M)|Ni(s)
(b) Zn(s)|Zn2+(1.0 x 10-2 M)||Zn2+(1.0 M)|Zn(s)
(c) Fe(s)|Fe2+(2.0 M)||Cr2O72-(1.0 M), H+(3.0 x 10-3 M), Cr3+(0.500 M)|Zn(s)
PS32.6. A galvanic cell consisting of a Cu versus a hydrogen electrode was used to determine the pH of
an unknown solution. The unknown was placed in the hydrogen electrode compartment and the pressure
of the hydrogen gas was controlled at 1 atm. The concentration of Cu2+ was 1 M and the E(not standard
E) of the cell at 25 ºC was determined to be +0.48 V. Calculate the pH of the solution.
PS32.7. Consider the following half-reaction
Eº = +0.96 volts. Calculate Ered when NO3- is 1 M and NO is 1 atm and the pH is;
(a) 0
(b) 2
(c) 7
PS32.8. An iron ore sample weighting 0.8500 g is dissolved in HCl(aq), treated with a reducing agent to
convert all the iron to Fe2+(aq), and then titrated with exactly 28.28 mL of 0.0652 M K2Cr2O7. What is
the percent Fe, by mass,
in the ore sample?
Problem Set #33
AP Chemistry by Satellite Name___________________________________
ALL work must be shown in all problems for full credit.
PS33.1. When a sample of molten MgCl2 and a sample of aqueous MgCl2 are electrolyzed with inert
electrodes different products are obtained. Determine the product of the electrolysis of each sample and
explain why there is a difference.
PS33.2a) A sample of molten CdCl2 and a sample of aqueous CdCl2 are electrolyzed with inert
electrodes. Determine the products of the electrolysis of each sample.
b) If the sample of aqueous CdCl2 is electrolyzed using a cadmium metal anode and an iron metal
cathode what are the expected results at each electrode?
PS33.3. Sketch a cell which depicts the electrolysis of a 1.00 M aqueous solution of HCl. Assume both
the electrodes are copper. Indicate the ions in the solution and the half-reaction at each electrode.
Calculate the minimum voltage required for the electrolysis to occur.
PS33.4. In the electrolysis of KCl(aq), how many liters of Cl2(g) are formed, measured at STP, by a
current of 10.0 amps for 1.5 hours? How many moles of KOH are formed in the same period?
PS33.5. If an aqueous solution containing Pb2+ ions is electrolyzed using a current of 0.900 amps,
calculate the mass of Pb(s) plated out after 50 minutes.
PS33.6. A certain quantity of electricity is passed through a series of solutions containing AgNO3,
CrCl3, ZnSO4 and CuSO4.
a) If 1.00 g of Ag(s) was deposited in the first solution, how many grams of metal were deposited in each
of the remaining solutions?
b) How much electricity (in coulombs) was used?
c) If the process took place in 3.00 hours, what was the current produced by the power source?
PS33.7. Can you electrolyze water into H2(g) and O2(g) using the following power sources
(theoretically)?
a) standard car battery
b) transistor radio battery
c) lantern battery
d) 'D' cell flashlight battery
e) Zn(s)|Zn2+(aq)||Cu2+(aq)|Cu(s)
PS33.8. An iron object must be coated with another metal to protect it from corrosion. Rank the
following metals from most protective to least protective to iron.
Al, Mg, K, Cu, Ag, Sn, Zn
Which metal would be the most appropriate choice (consider reactivity and cost) for an iron tool that
will be exposed to the atmosphere?