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CHAPTER 1.
STOICHIOMETRY
1. Convert 3.01 x 1023 atoms of rubidium to moles (0.500 mol Rb)
2. The bond energy for HF is 568 KJ/mol. How much energy in Joules is required to break a single HF bond? (9.44 x 10-19 J)
3. How many moles are present in 50.0 g of calcium chlorate? (0.242 mol)
4. How many formula units are present in 0.272 g of nickel(II) nitrate? (8.96 x 1020)
5. The density of liquid water is 0.997 g/mL at 25°C. How many moles of water are in 250.0 mL of water? (13.8 mol H2O)
6. Determine the molarity of a solution prepared by dissolving 141.6 g of citric acid, C3H5O(COOH)3 , in water and then diluting the
resulting solution to 3500.0 mL. (0.2106 M)
7. What mass of glucose, C6H12O6, would be required to prepare 5.000 x 103 L of a 0.215 M solution? (1.94 x 105 g)
8. Determine the mole fraction of a solution of 560 g of acetone, CH3COCH3, in 620 g of water. (0.219)
9. What volume of water would be added to 16.5 mL of a 0.0813 M solution of sodium borate in order to get a 0.0200 M solution?
(50.6 mL H2O)
10. A chemist wants to prepare a stock solution of H2SO4 so that samples of 20.00 mL will produce a solution with a concentration of
0.50 M when added to 100.0 mL of water.
a. What should the molarity of the stock solution be? (3.0 M)
b. If the chemist wants to prepare 5.00 L of the stock solution from concentrated H2SO4 , which is 18.0 M, what volume of
concentrated acid should be used? (0.83 L)
c. The density of 18.0 M H2SO4 is 1.84 g/mL. What mass of concentrated H2SO4 should be used to make the stock solution in
(b)? (1.5 x 103 g)
11. Maleic acid, which is used to manufacture artificial resin, has the empirical formula CHO. Its molar mass is 116.1 g/mol. What is its
molecular formula? (C4H4O4)
12. Hydrated salts are very common. If you heat 2.105 g of CoCl2xH2O, and find that 1.149 g of CoCl2 remains, what is the value of x?
(6)
13. Ammonia gas can be prepared by the following reaction: CaO(s)+2NH4Cl(s) 2NH3(g)+ H2O(g)+CaCl2(s). If you mix 112 g of CaO
and 224 g of NH4Cl, what is the theoretical yield of NH3? (68.1 g). What mass of excess is remaining? (10.3 g)
14. Disulfur dichloride can be made by allowing chlorine gas to react with molten sulfur: S8(l) + 4Cl2(g)  4S2Cl2(g) If you begin with
12.0 g of S8 and 13.03 g Cl2 and you isolate only 15.2 g of S2Cl2, what is the percentage yield of S2Cl2? (60.1%)
15. Styrene, the building block of polystyrene, is a hydrocarbon, a compound consisting only of C and H. If you burn 0.438 g of the
compound, and find that it produces 1.481 g of CO2 and 0.303 g of H2O, determine the empirical formula of the compound. (CH)
16. Aluminum bromide is a valuable laboratory chemical. If you use 25.0 mL of liquid bromine (D = 3.1023 g/mL) and excess
aluminum metal, what is the maximum theoretical yield of Al2Br6? (86.3 g)
17. In the photographic process silver bromide is dissolved by adding sodium thiosulfate, as shown in the balanced equation below. If
you want to dissolve 0.250 g of AgBr(molar mass = 187.8 g/mol), how many mL of 0.0138 M Na 2S2O3 should you add? (193 mL)
AgBr(s)+2Na2S2O3(aq) Na3Ag(S2O3)2(aq)+NaBr(aq)
18. A soft drink contains an unknown amount of citric acid, C6H8O7. If 100. mL of the soft drink required 33.51 mL of 0.0102 M NaOH
to neutralize completely the citric acid, how many grams of citric acid (MM = 192.13 g/mol) does the soft drink contain per 100
mL? The reaction of citric acid with NaOH is(all aq) shown below. (0.0219 g)
C6H8O7 + 3NaOH  Na3C6H5O7 + 3H2O(l)
19. NutraSweet is 57.14% C, 6.16% H, 9.52% N, and 27.18% O. Calculate the empirical formula of NutraSweet and find the molecular
formula. (The molar mass of NutraSweet is 294.30 g/mol). (C14H18N2O5)
20. Sodium thiosulfate, Na2S2O3, is used as a “fixer” in black and white photography. Assume you have a bottle of sodium thiosulfate
and want to determine its purity. You can titrate the thiosulfate ion with I 2 according to the equation below. If you use 40.21 mL of
0.246 M I2 in the titration, what is the weight percent of Na2S2O3 (MM = 158.12 g/mol) in a 3.232-g sample of impure Na2S2O3?
(96.7 %) 2S2O32- + I2  S4O62- + 2I21. Menthol, the substance we can smell in mentholated cough drops, is composed of C, H, and O. A 0.1005 g sample of menthol is
combusted, producing 0.2829 g of CO2 and 0.1159 g of H2O. What is the empirical formula for menthol? (C10H20O)
22. You are given a sample of a copper-containing alloy and asked to determine the mass percent of copper. After dissolving the metal
in acid, you add an excess of KI, and the Cu2+ and I- ions undergo the reaction:
2Cu2+ + 5I─  2CuI + I3─
The liberated I3- is titrated with sodium thiosulfate according to the equation:
I3─ + 2S2O32─  S4O62- + 3I─
If 26.32 mL of 0.101 M Na2S2O3 is required for titration to the equivalence point, what is the mass percent of Cu in 0.251 g of the alloy?
(67.2%)
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CHAPTER 2.
ATOM
A. Conceptual Structure
1. Explain all of the variables in Coulomb’s law as it applies to the atom.
2. Differentiate among energy levels, sublevels, and orbitals. Include what each communicates about region is space where
an electron is most likely to be found.
3. Define AufBau principle, Hund’s Rule, and Pauli Exclusion principle. Explain each based upon maximizing attractive
forces and minimizing repulsive forces in an atom.
4. Write the complete electron configurations for the following atoms.
a. Ruthenium
b. Radium
5. Write the complete electron configurations for the following ions.
a. Bismuth +3 ion
b. Iron +2 ion
c. Tellurium -2 ion
6. Write the noble gas configurations for the following atoms.
a. Nickel
b. Tungsten
7. Identify the element from the following photoelectron spectra. Indicate the sublevel and number of electrons
represented in each peak of the spectra. Justify the difference in energy values for the same sublevels in each based upon
atomic structure.
8. PES can be used to identify elements in a mixture. What elements are present in the mixture below?
Potassium
Silicon
Chlorine
Mixture
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B. Conceptual Periodic Trends
INTRO: You need to IDENTIFY and JUSTIFY the trends. You cannot use one trend (like electronegativity) to justify another
trend (such as ionization energy). You answer must include a comparison of attractive forces (proton-electron) and repulsive
forces (electron-electron). Distance from the nucleus, energy levels, conversion of energy levels, and shielding are critical
components of a justification. STUDY MY SUMMARY CHART CAREFULLY!
9. Which of the following groups of elements is arranged correctly in order of increasing first ionization energy?
A. Mg < C < N < F
B. N < Mg < C < F
C. Mg < N < C < F
D. F < C < Mg < N
10. Which of the following groups of elements is arranged correctly in order of decreasing atomic radius?
A. Mg < S < Al < Cl
B. Al < Mg < S < Cl
C. Mg < Al < S < Cl
D. Cl < S < Mg < Al
11. Which of the following elements would have the greatest difference between the first and the second ionization energy?
A. Li B. C
C. F
D. N
12. Which of the following groups of isoelectronic species show the elements arranged correctly in order of increasing size?
A. Na+ < O2- < FB. F- < Na+ < O2C. Na+ < F- < O2D. F- < O2- < Na+
13. An element having which of the following electronic configurations would have the greatest ionization energy?
A. [He]2s22p3
B. [He] 2s22p5
C. [Ne]3s23p3
D. [Ne]3s23p5
14. Periodic trends:
A. Which should be larger, the oxide ion, O2-, or the oxygen atom?
B. Which should have the largest difference between the 1st and 2nd ionization energy? O, S, or Se
15. Which of the following concerning second IE's is true?
A. That of Al is higher than that of Mg because Mg wants to lose the second electron, so it is easier to take the second
electron away.
B. That of Al is higher than that of Mg because the electrons are taken from the same energy level, but the Al atom
has one more proton.
C. That of Al is lower than that of Mg because Mg wants to lose the second electron, thus the energy change is
greater.
D. That of Al is lower than that of Mg because the second electron taken from Al is in a p orbital, thus it is easier to
take.
E. The second ionization energies are equal for Al and Mg.
Ionization Energies for element X (kJ mol¯1)
First
Second
Third
Fourth
Five
1086
2352
4619
6221
37820
16. The ionization energies for element X are listed in the table above. On the basis of the data, element X is most likely to be
A. Li
B. Be
C. B
D. C
E. P
17. 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?
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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?
18. The correct ranking of alkali metals from most reactive to least reactive is:
A. Be-Mg-Co-Sr-Ba
D. I-Br-Cl-F
B. Cs-Rb-K-Na-Li
E. Li-Na-K-Rb-Cs
C. F-Cl-Br-I12
C. Mathematical
19. Calculate the average atomic mass of chromium given the following isotopes with abundance.
Chromium-50
Chromium-52
Chromium-53
Chromium-54
Actual mass
49.945046
51.940519
52.940651
53.938882
Abundance
4.35
83.79
9.50
2.36
20. The mass spectrum below shows the relative % abundance for the isotopes of zinc.
a. Redraw the spectrum so that the true % abundance is on the y-axis. WATCH Y-AXIS CAREFULLY ! I AM NOT
SURE WHETHER AP WILL USE RELATIVE % OR ACTUAL %.
b. Calculate the average atomic mass of zinc.
21. Perform the following calculations regarding electromagnetic radiation.
a. Convert 4.398 x 10-19 J to frequency.
b. Convert 893 nm to energy
c. The energy of a photon is 8.22 x 10-18 J. What is the energy in KJ/mol?
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CHAPTER 3.
BONDING
1. Explain all of the variables in Coulomb’s law as it applies to bonding.
i. Ionic
ii. Covalent
iii. Metallic
2. Differentiate among non-polar covalent, polar covalent, and ionic BONDS.
3. What is/are the difference(s) between network covalent and molecular covalent bonding/compounds?
4. Define each of the terms in the acronym VSEPR. How is the model used to predict the structure of molecular covalent
substances?
5. Predict whether the following pairs of atoms are more likely to form an ionic, metallic or covalent bond using only a
periodic table. Justify your answer in each case.
a. Al & Ga
b. Zn & P
c. P & Br
6. Rank the polarity of the following bonds using only a periodic table. Justify your answer. P — N, P—S, P—Br, P—O
7. Define formal charge. Calculate the formal charge of each element in the molecule below. How is formal charge linked
to the selection of a possible structure of a molecule?
8. Complete (and memorize!) the following summary chart for molecular covalent substances
Bonding atoms
Non-bonded
“ABX”
Molecular structure
Bond angle(s)
(or group of
pairs of
atoms) on a
electrons on a
defined central
defined central
atom
atom
2
0
Hybridization
3
0
2
1
4
0
3
1
2
2
5
0
N/A
4
1
N/A
3
2
N/A
2
3
N/A
6
0
N/A
5
1
N/A
4
2
N/A
9. Explain why the bond angle decreases as the number of non-bonded pairs on the central atom increases.
10. Electromagnetic radiation is commonly used to study atoms and
molecules. What type of electromagnetic radiation is used to study
the following?
a. Electronic transitions
b. Ionization
c. Molecular Vibrations
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11. Electronic transitions are studied using the UV/Vis range of EMR. Which type of motions, vibrational or electronic,
requires the greatest energy? Justify your answer.
12. The structural formula for the amino acid, glycine is shown. Indicate the structure, hybridization, and bond angles for
all elements numbered. Justify your answer in each case.
13. What is Beer’s law and how is it commonly used to study molecules?
14. Molecules that have alternating (conjugated) double bonds absorb visible light (you should know this fact for test).
What ions are colored and form solutions that absorb visible light? Justify your answer.
a. Use the data below (excel or calculator) to make a Beer’s Law plot, write down the equation of the line, and
calculate the molarity of the unknown solution. (9.01e-5 M)
Molarity
Absorbance
7.817e-5
0.1
1.782e-4
0.188
2.7817e-4
0.272
3.7817e-4
0.43
4.7817e-4
0.469
5.7818e-4
0.566
unknown
0.111
b. Calculate the molar absorptivity () of the substance assuming the path length of the cuvette was 1 cm.
15. Which of the following pairs of bonded atoms would be expected to have the longest bond length?
D. C-N
B. C-S
C. C-B
D. C-F
16. Which of the descriptions below is the best representation of the energy change involved in the process of breaking
bonds in a molecule? (ignore any subsequent bond formation that may occur)
A. Always exothermic
B. Always endothermic
C. Net energy change is zero
D. Exothermic or endothermic depending on conditions.
17. How many sigma (σ) and pi(π) electron pairs are there in a carbon dioxide molecule?
A. Two sigma, zero pi
B. One sigma, one pi
C. Two sigma, two pi
D. Two sigma, one pi
18. Which of the following elements is most likely to form compound involving an expanded valence shell of electrons?
A. P B. Na C. O
D. N
19. Which of the following statements best describes the relationship between bond length and bond strength for a series
of compounds involving bonds between the same two atoms?
A. The greater the bond strength, the longer the bond.
B. The greater the bond strength, the shorter the bond.
C. Bond length and bond strength are not related
D. The relationship between bond length and bond strength depends on other factors.
20. Which of the following combinations of two elements is most likely to produce highly ionic bonds?
A. Nitrogen and oxygen
B. Nitrogen and fluorine
C. Boron and nitrogen
D. Lithium and fluorine
21. Which of the following combinations of two elements is most likely to produce covalent bonds?
A. nitrogen and oxygen
B. oxygen and calcium
C. sodium and nitrogen
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22.
23.
24.
25.
26.
27.
28.
29.
30.
31.
32.
33.
34.
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D. lithium and fluorine
Which of the following salts is expected to have the highest melting point?
A. NaF
C. NaI
B. NaCl
D. NaBr
Predict which compound in each of the following pairs should have the higher melting point.
(i)
NaCl or RbCl
(ii) NaCl or MgCl2
Draw Lewis dot structures for the following:
A. Ammonia
B. Hydrogen cyanide
C. N2O
D. BrF5
E. Ca3(PO4)2
Based on the VSEPR theory, what is the molecular shape of PCl5?
A. Linear
B. tetrahedral
C. Trigonal planar
D. Trigonal bipyramidal
Based on the VSEPR theory, which of the following corresponds most closely to the molecular shape of SCl2?
A. Linear
B. “T-shaped”
C. bent (bond angle 120o)
D. bent (bond angle 109.5o)
A certain molecule has five structural electron pairs and the molecular structure is linear. How many lone pairs are
present in this molecule?
A. None
C. Two
B. One
D. Three
A certain molecule has six structural electron pairs and the molecular structure is a square pyramid. How many lone
pairs are present in this molecule?
A. None
C. two
B. One
D. three
What is the approximate Cl-B-Cl angle in BCl3?
A. 90o
C. 120o
o
B. 109.5
D. 180o
What is the approximate I-I-I angle in I3─?
A. 90o
C. 120o
B. 109.5o
D. 180o
Which of the following best describes the variation of electronegativity of the elements with respect to their position
on the periodic table?
A. Increases across a period, increases down a group
B. Increases across a period, decreases down a group
C. Decreases across a period, increases down a group
D. Decreases across a period, decreases down a group
What is the formal charge on the O atoms in SO32─?
A. 0 B. +1 C. -1
D. +2
What is the formal charge of the S atom in SO3?
A. 0 B. +1 C. -1
D. +2
What is the average carbon-oxygen bond order in the formate ion?
A. One
C. Two
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B. 1 ½
D. 2 ½
35. What is the average sulfur-oxygen bond order in SO3?
A. One
C. Two
B. ½
D. 1 1/3
36. In which species is the carbon-oxygen bond longer?
A.
B.
37. Given the bond dissociation energies below, calculate the standard molar enthalpy of formation of NF 3 (in KJ/mol).
½ N2(g) + 3/2 F2(g)  NH3(g)
Bond Dissociation Energy (KJ/mol)
N≡N 946
F─F
159
N─F
272
A. 833
C. -104
B. 440.
D. -578
38. Which of the bonds below is least polar?
A. C─O
C. C─N
B. C─F
D. C─B
39. Which of the following molecules is polar?
A. NCl3
C. SF6
B. O2
D. CS2
40. Which of the following molecules is most likely to have a dipole moment?
A. CH4
C. SF6
B. BeF2
D. NF3
41. Cyanic Acid has the electron dot structure below (you must add non-bonding pairs of electrons):
H─O─C≡N
42.
43.
44.
45.
A. How many sigma (σ) bonds are there?
B. How many pi(π) bond are there?
C. What is the value of the C─O─H angle?
D. What is the value of the N─C─O angle?
Which of the following elements is most likely to display sp3d hybridization?
A. Oxygen
C. Phosphorus
B. Nitrogen
D. Carbon
What type of hybrid orbital set is used by the nitrogen atom in the molecule NH 3?
A. sp B. sp2 C. sp3 D. sp3d2
What type of hybrid orbital set is used by the xenon atom in the compound XeF4?
A. sp B. sp2 C. sp3 D. sp3d2
What hybrid orbital set is used by the nitrogen atom in the following molecule? You must add non-bonded pairs of
electrons.
H3C─N═C═O
A. sp
B. sp2
C. sp3
D. sp3d2
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CHAPTER 4.
BONDING FUNCTION PURE
1. Explain all of the variables in Coulomb’s law as it applies to intermolecular forces.
2. What is the difference between a “bond” and an “intermolecular force”? Do you think an AP reader would give you
credit if you used these words interchangeably? Why or why not?
3. Identify the force of attraction that must be broken for the following substances. If the force of attraction is an
intermolecular force, list all IMF’s involved. Please do not use “LD”, “HB” etc – spell it out! AP won’t accept acronyms
unless they are defined first.
Substance
Covalent, Ionic, Metallic, or
If IMF, which one(s)
IMF?
Brass
NH3
C6H12O6
HCl
SiO2
Si
CaCl2
PCl3
(an amino acid)
Na
NaCl
C 3H 6
4.
5.
6.
7.
Answer the questions below regarding hydrogen iodide and iodine.
a. Identify the intermolecular forces present in each of the pure substances.
b. The melting points are -50.80oC for HI and 113.7oC for I2. Explain the large difference using bonding and
intermolecular forces.
Below is a table of the vapor pressure at 20oC for a variety of substances.
Pvapor at 20oC (KPa)
Pentane
57.90
2-pentanone
3.6
1-pentanol
0.200
butane
203
a. Use intermolecular forces to explain the trend observed among the three 5-carbon substances. (You may
want to look up their structural formulas online)
b. Use intermolecular forces to explain the difference between butane and pentane.
Use molecular structures to model both hydrogen-bonding possibilities in water. Label each atom with partial
charges.
Predict which substance is likely to have the highest boiling point in each of the following pairs. Justify your answer
using Coulomb’s law.
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a. Na2O or CaO
c. KF or KCl
b. SrO or CaO
d. RbCl or SrBr2
CHAPTER 5.
BONDING FUNCTION MIXTURE
1. Dissolving and mixing always involves an increase in entropy. Despite this favorable change, not all substances mix or
dissolve due to unfavorable energetics. For each of the following pairs, indicate the forces of attraction that must be
broken and formed in order for the two to mix.
PAIR
Water and ethanol
BROKEN
FORMED
NET ENERGY PREDICTION
Water and benzene
Water and copper(II) sulfate
Benzene and cyclohexane
2.
3.
4.
5.
6.
Use coulombs law to explain why the aluminum ion, Al3+, has a greater attraction to water than the gallium ion, Ga3+.
Use structural formulas to show ethanol hydrogen bonding with water molecules. Make sure to show all possibilities.
Include partial charges on each atom.
Differentiate between P-type doping and N-type doping of silicon. Provide an example of each.
Differentiate between interstitial and substitutional alloys. Look up an example of each and list how properties
changed in the alloy compared to the pure metal.
How does the radius of a metal atom solute affect the hardness of the metal solvent in alloys?
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CHAPTER 6.
REACTIONS
1. What is characteristic of reactions that can be described by net ionic equations?
2. What classifications (decomposition, synthesis, combustion, double replacement, single replacement, neutralization,
redox) are often described by net ionic reactions?
3. When potassium chlorate is decomposed, a gas forms that is able to re-ignite a glowing splint. Write the balanced
reaction for the decomposition.
4. Write the balanced chemical reactions (including units) for the following:
a. Solutions of iron(III) perchlorate + calcium hydroxide.
b. A piece of copper is added to a solution of lead(II) nitrate.
c. Nitrous acid + aqueous potassium hydroxide
5. Write net ionic equations for the reactions in #3.
6. Write the balanced reaction equation for the combustion of the following substances
a. C3H6
b. C6H6
c. C5H9OH
7. Draw a particle diagram showing the synthesis reaction between nitrogen monoxide gas and oxygen gas to form
nitrogen dioxide. Use 5 molecules of oxygen in your drawing.
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CHAPTER 7.
KINETICS
1. List the three premises of collision theory.
2. What experimental factors increase the rate of a reaction? Explain why in terms of collision theory.
3. Define the following terms:
A. Mechanism
E. Rate constant expression
H. Order
B. Intermediate
F. Molecularity (unimolecular,
I. Initial rate
C. Catalyst
bimolecular, termolecular)
J. Instantaneous rate
D. Rate constant
G. Rate determining step.
K. Average rate
4. The following mechanism has been proposed for the formation of NOBr:
Step 1: NO(g) + Br2(g) ⇄ NOBr2(g)
(fast)
Step 2: NOBr2(g) + NO(g)  2NOBr(g) (slow)
A. Write the overall balanced equation for this reaction.
B. Derive the rate law expression predicted by this mechanism.
C. Why is this reaction not likely to occur in one step?
5. For the reaction that follows what is [NO2]/t with respect to [NO]/t? 2NO(g) + O2(g)  2NO2(g)
6. Given the initial data for the reaction A + B  C, determine the rate expression for the reaction. The reaction is catalyzed
by molecule “D”. Calculate “k” as well.
[A] (M) [B] (M) [D] (M) [C]/t (M/s)
0.10
0.20
0.10
2.58
0.10
0.10
0.10
1.29
0.24
0.10
0.20
2.58
0.10
0.20
0.20
5.16
7. Hydrogen peroxide decays into water and oxygen in a first-order process,
H2O2(aq)  H2O(l) + ½ O2(g)
Where the rate expression is −[H2O2]/ t = k[H2O2]. If we begin with 0.100 M H2O2 and find that after 3200 seconds the
peroxide concentration falls to 0.0825 M, what is the rate constant, k, at the temperature at which the experiment is
performed? (6.01 x 10−5 s−1)
8. The decomposition of SO2Cl2 is first order in SO2Cl2 with a rate constant of 0.17 hr─1. What is the half-life of SO2Cl2? (4.1
hours)
9. We know that the half-life for the first-order, radioactive decay of 222Rn is 3.82 days. If we have 5.80 micrograms of 222Rn,
how much will remain after 21.2 days?(1.27 x 10─7)
10. We are studying the reaction A  2B. A plot of 1/[A] vs. time as the reaction proceeds is linear with a slope of 0.187
M─1s─1. What is the rate expression for the reaction?
11. The reaction I─ + OCl─  IO─ + Cl─ is first order with respect to I─ and first order with respect to OCl─. The rate constant
is 6.1 x 10─2 L/mol-s. What is the rate of reaction when [I─] = 0.10 M and [OCl─] = 0.20 M? (1.2 x 10─3 M/s)
12. What fraction of a reactant remains after 3 half-lives of a first order reaction?
13. At what point on the potential energy diagram shown does the transition state
(activated complex) occur?
14. The rate of the chemical reaction between substances A and B is found to
follow the equation rate=k[A]2[B], where k is a constant. If the concentration
of A is halved, what should be done to the concentration of B to make the
reaction go to 75% of its former rate?
A. The concentration of B should be kept constant
B. The concentration of B should be doubled
C. The concentration of B should be tripled
D. The concentration of B should be halved
E. The concentration of B should be multiplied by 4/3.
15. Answer the following questions regarding the potential energy curve shown. Justify all answers.
A. How many steps are in the mechanism for this reaction?
B. Which step is the rate determining step?
C. Which letter(s) represent intermediates?
D. Which letter(s) represent activated complex(es)?
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Is the overall reaction endothermic or exothermic?
CHAPTER 8.
EQUILIBRIUM
A. General Equilibria
1. Explain the term “dynamic equilibrium”.
2. Define the reaction quotient, Q.
3. What is the difference between the reaction quotient “Q” and the equilibrium constant “K”.
4. Which of the following changes the value of an equilibrium constant. Justify your answer in each case.
Temperature
Concentration of
products
Catalyst
Concentration of
reactants
Nature of the reaction
5. Given the following two equilibria, calculate the equilibrium constant for the third (2.1)
PbI2(s) ⇄ Pb2+(aq) + 2I-(aq) K1 = 8.7 x 10-9
PbSO4(s) ⇄ Pb2+(aq) + SO42-(aq) K2 = 1.8 x 10-8
PbSO4(s) + 2I-(aq) ⇄ PbI2(s) + SO42-(aq) K3 = ??
6.
7.
Given the equilibrium constant, Kp, for the following reaction at 250C, calculate the equilibrium constant, Kc, for the
same reaction at the same temperature. (R = 0.0821) (6.1 x 10-3)
N2O4(g) ⇄ 2NO2(g) Kp = 0.15
Determine if the following system is at equilibrium, the reactant concentrations are too high, the product
concentrations are too high, or if one simply cannot determine with information given:
PCl5(g) ⇄ PCl3(g) + Cl2(g)
Kp = 11.5
𝑃𝑃𝐶𝑙5 = 1.15 𝑎𝑡𝑚
𝑃𝑃𝐶𝑙3 = 5.30 𝑎𝑡𝑚
𝑃𝐶𝑙2 = 2.80 𝑎𝑡𝑚
8. The equilibrium constant for the reaction: SO2(g) + NO2(g) ⇄ SO3(g) + NO(g) has a numerical value of 3.00 at a
given temperature. Equimolar amounts of SO2 and NO2 are reacted at this temperature and a total pressure of 3.00
atm. What percent of the SO2 is converted to product? (63.4 %)
9. For which of the reactions will an increase in pressure cause a decrease in product (temperature remaining constant)?
Justify your choices.
A. N2(g) + 3H2 (g) ⇄ 2NH3(g)
B. 3Fe(s) + 3H2O(g) ⇄ Fe2O3 (s) + 3H2(g)
C. PCl3(g) + Cl2(g) ⇄ PCl5(g)
D. HCl(g) + H2O(l) ⇄ H3O+ (aq) + Cl-(aq)
E. CaCO3(s) ⇄ CaO(s) + CO2(g)
10. Consider the equilibrium reaction: SO2(g) + O2(g) ⇄ SO3(g) What will be the effect of doubling the concentration of
SO3? Justify you answer(s)
A. SO2 and O2 increase equally
B. SO2 increases more than that of O2
C. O2 increases more than that of SO2
D. SO2 decreases more than that of O2
E. O2 decreases more than that of SO2
11. The synthesis of hydrogen sulfide gas is given by the following: H2 (g) + S (s) ⇄ H2S (g) + energy. List all of the
“stresses” that would increase the mass of sulfur. Justify your answer in each case.
12. For each of the following “stresses” on the equilibrium system, N2O4(g) ⇄ 2NO2(g), indicate how the stress alters the
forward and/or reverse rates (are they increased or decrease or unaffected?) ∆H = 57.2 KJ
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Forward
rate
Reverse
rate
Shift in response
to stress
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Justification
Temperature is
increased
N2O4 is
decreased
NO2 is
decreased
N2O4 is
increased
NO2 is increased
A catalyst is
added
The volume is
decreased
Helium is added
at constant
volume
13. What must be true of a reaction in order for Kc = Kp?
14. Indicating Drierite is a material used in the laboratory to remove water vapor from gases and as a dessicating agent.
When purchased, it is blue in color and it changes to pink upon absorbing moisture.
a. Based upon your observations in the LeChatelier lab, explain how Indicating Drierite works.
Blue Drierite + H2O ⇄ Pink Drierite + heat
b. Indicating Drierite can be regenerated and used over and over again. Propose a method of regenerating
Drierite quickly, efficiently, and at low cost. Explain how your method will work, using Le Chatelier’s
principle.
15. Nickel forms a green complex ion, [Ni(H2O)6]2+, in water. The addition of a small amount of ethylenediamine (en)
results in a light blue–colored solution. Ethelynediamine is a bidentate ligand, meaning it bonds to metals through
two atoms of the ligand. Further addition of ethylenediamine results in a royal blue solution while the addition of even
more ethylenediamine gives a violet solution. Further addition produces no additional color changes.
a. Propose a series of equilibrium reactions that would illustrate the observations.
b. Predict the effect of adding solid nickel(II) nitrate to the violet solution. Explain your answer using Le
Chatelier’s principle.
B. Solubility Equilibria
1. For slightly soluble salts, if Q < K is the solution saturated, unsaturated, or supersaturated? Justify your answer.
2. For slightly soluble salts, if Q > K, would the mass of the precipitate increase or decrease? Justify your answer.
3. Write the dissociation reactions and Ksp expressions for the following slightly soluble salts. Write Ksp in terms of the
solubility, S. The first one has been done for you as a model.
SALT
DISSOCIATION REACTION
EQUILIBRIUM CONSTANT
Ksp in terms of “S”
EXPRESSION
Cd(OH)2(s) ⇄ Cd2+(aq) + 2OH─(aq)
Ksp = [Cd2+][ OH─]2
Ksp = 4S3
Cd(OH)2
CoCO3
LaF3
Hg2S
Ba3(PO4)2
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4.
5.
6.
7.
8.
9.
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Calculate the molar solubility and g/L solubility values for Cd(OH)2 (Ksp = 2.5 x 10-14) and Ba3(PO4)2 (3.4 x 10-23).
In the above example: Will the addition of Na3PO4 increase or decrease the solubility of barium phosphate at a given
temperature? Justify your answer.
Lead(II) carbonate has a Ksp= 7.4 x 10-14 and lead(II) hydroxide has a Ksp = 1.2 x 10-15. Calculate the molar solubility
for each of these compounds. Can the magnitude of Ksp be used to predict solubility? Explain.
The solubility of copper (II) iodate in water is 3.27 x 10-3 M. What is the Ksp? (1.40 x 10-7)
A 10.0 mL solution is 0.10 M with respect to both calcium nitrate and magnesium nitrate. Which salt begins to
precipitate first if 0.10 M NaF is added by drops? (Ksp CaF2= 4.0 x 10-11 & Ksp MgF2 = 6.4 x 10-9) Justify your answer.
Which ions will remain in solution if excess magnesium nitrate is added to sodium fluoride? Rank the ions in order of
relative abundance in the solution.
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CHAPTER 9.
ACIDS AND BASES
1. Indicate whether each of the following pairs represent conjugate pairs (acid/CB 0r base/CA). Justify your answer in
each case.
A. H2C2O4 & C2O42B. H2PO4- & HPO42C. CH3NH2 & CH2NH2
D. H3O+ & OH2. This reaction is reversible. Identify all acids and bases. HSO3- + H2O ⇄ SO32- + H3O+
3. Answer the follow questions regarding the titration curve. Justify your answer in each case.
A. Which point is the best for determining whether the acid is strong or weak?
B. Why is the equivalence point basic? Which point(s) represent a region of buffering?
C. After which point is the acid primarily in its
deprotonated form?
D. At which point is the pH determined from the
amount of sodium hydroxide?
E. From which point can the Ka of the acid be
determined?
F. What two factors determine the starting pH in a
titration?
G. If the acid were HCN, which ions would be
present in the solution at the “X”? Draw a
particulate diagram to show the relative
amounts.
H. Use the data from the titration curve to
determine the molarity of the acid.
I. An indicator was used that changed color so that
the endpoint was slightly sooner than the
equivalence point, would the calculated molarity from the endpoint be too high or too low?
J. Sketch the curve for a strong acid of equal molarity. What are the three main differences in the curves?
K. Sketch the curve observed for the weak acid if the molarity of the sodium hydroxide is doubled.
4. Since ions are formed and consumed during a titration, it is sometimes
possible to measure conductivity instead of pH. This approach can be
used when titrating barium hydroxide with sulfuric acid
A. Write the reaction for the dissociation of barium
hydroxide.
B. Write the balanced, net ionic equation for the
reaction of barium hydroxide with sulfuric acid.
C. Why does the conductivity decrease at the beginning
of the titration?
D. Why does the conductivity increase after the
equivalence point?
5.
6.
You are given 4.554 g of a mixture of oxalic acid, H2C2O4, and sodium chloride. If 29.58 mL of 0.550 M NaOH are required
to titrate the 4.554 g sample to the equivalence point, what is the mass percent of oxalic acid (MM = 90.04 g/mol) in the
mixture? (16.1%)
H2C2O4 + 2NaOH  Na2C2O4 + 2H2O
Knowing that HF is a stronger acid than HC2H3O2, determine, if possible, in which direction the following equilibrium lies:
HF + C2H3O2─ ⇄ F─ + HC2H3O2
A. Left
7.
B. Right
C. Perfectly balanced
D. Cannot be determined
At 50oC the water ionization constant, Kw is 5.48 x 10-14. What is [H3O+] in neutral water at 50oC? (2.34 x 10-7 M)
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8.
9.
10.
11.
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We have a 0.00100 M solution of Sr(OH)2 solution at 25oC. What is [H3O+] in this solution? (5.00 x 10-12 M)
We have a 5.43 x 10-4 solution of HNO3 at 25oC. What is [OH─] in this solution? (1.84 x 10-11 M)
We have 500. mL of a solution that contains 0.0854 g of NaOH. What is the pH of this solution at 25 oC? (11.63)
For each solution below, tell if the pH is < 7, pH = 7 or if pH > 7. JUSTIFY
A. 0.10 M HNO3
B. 0.012 M KOH
C. 0.15 M acetic acid
D. 0.56 M Na2CO3
E. 0.45 M KBr
12. Predict the products of the following acid-base reaction: NH3+ HNO3
13. The following reactants are mixed in equimolar portions. Predict the resulting solution will be (A) acidic, (B) Basic,
(C) neutral (D) cannot be determined.
A. HCl + NaHCO3  ?
B. HCl + NaOH  ?
C. HF + KOH  ?
D. H2SO4 + KOH  ?
E. CH3COOH + NH3  ? (HINT: look up “K” values)
14. Indicate whether the following describes and acid, base, or both
A. Donates H+
B. Solution conducts electricity
C. Turns litmus paper red
D. Pink with phenolphthalein
E. Reacts with active metals
F. Tastes bitter
G. Feels slippery
15. Which of the following substances is amphoteric? JUSTIFY
A. Al(OH)3
D. Ca(OH)2
B. HCN
E. HI
C. CsBr
16. Which of the following substances is amphoteric? (Select all correct answers) JUSTIFY
A. SO42─
C. H2PO42─
B. H2O
D. H2SO4
17. In which of the following is the acid strength ranking INCORRECT? JUSTIFY your choice.
A. H2SO4 > H2SO3
B. HNO3 > HNO2
C. HClO4 > HClO3
D. HClO3 > HBrO3
E. H2SeO3 > H2SO3
18. Rank the hydrohalic acids from strongest to weakest: JUSTIFY
A. HF > HCl > HBr > HI
C. HCl > HBr > HI > HF
B. HI > HBr > HCl > HF
D. HF > HI > HBr > HCl
19. Rank the chlorine based acids from strongest to weakest. JUSTIFY
A. HClO > HClO2 > HClO3 > HClO4
B. HClO4 > HClO3 > HClO2 > HClO
20. Which of the following is the strongest acid? JUSTIFY
A. HClO3
C. HClO2
B. HBrO2
D. HBrO3
21. Which of the following is the strongest acid? JUSTIFY
A. H3AsO4
C. H3AsO3
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B. H3PO4
D. H3PO3
22. We have a solution of NH3. What effect will the addition of HCl have on the pH of the solution? JUSTIFY
A. Increase pH
B. Decrease pH
C. no effect
EQUILIBRIA, BUFFERS, TITRATIONS – YOU WILL NEED TO LOOK UP APPROPRIATE Ka AND Kb
23. What is the pH of a 0.155 M solution of H2S at 250C? (3.90)
24. What is the [OH─] in a 0.10 M solution of NaCN at 25oC? (1.57 x 10-3 M)
25. What is the pH of a 0.144 M solution of NaF at 25oC? (8.15)
26. If the pH of a 0.015 M solution of hypochlorous acid is 4.64, what is the concentration of the hypochlorite ion, OCl ─, in
solution? (2.29 x 10─5 M)
27. What is the value of Ka for hypochlorous acid from the previous question? (3.50 x 10─8)
28. We mix 50.0 mL of 0.050 M HNO3 and 25.0 mL of 0.10 M NaCH3COO. What is the pH of the resulting solution? (3.11)
29. We add 1.00 mL of 10.0 M NaOH to 50.0 mL of 0.20 M HNO2. What is the pH of the resulting solution? (8.32)
30. We have 100. mL of a 0.10 M solution of CH3COOH. How many grams of NaCH3COO must be added to make a buffer
solution of pH 5.00? Ka acetic acid = 1.8 x 10─5 (1.48 g)
31. A 0.10 M solution of HF is 8.1% ionized. What is the Ka? (7.1 x 10─4)
32. A 1.50 g sample of vitamin C is dissolved in 100.0 mL of water and titrated with 0.250 M NaOH to the methyl orange end
point. The volume of the base used is 34.1 mL. What is the molar mass of Vitamin C assuming one dissociable proton per
molecule? (176 g/mol)
33. A 25.00 mL sample of 0.100 M HCl is titrated with 0.100 M NaOH. What is the pH of the solution at the points where 25.1
and 25.5 mL of NaOH have been added? (10.30, 11.00)
34. A 25.00 mL sample of 0.100 M CH3CO2H is titrated with 0.100 M NaOH. What is the pH of the solution at the points where
24.5 and 25.5 mL of NaOH have been added? (K a = 1.8 x 10-5) (6.43, 11.00)
35. Which of the following mixtures will be a buffer when dissolved in a liter of water?
A. 1. 0.2 mol Ba(OH)2 and 0.3 mol HClO2
B. 2. 0.2 mol KNO3 and 0.2 mol HClO3
C. 3. 0.4 mol NH4Cl and 0.4 mol NaOH
D. 4. 0.2 mol HClO3 and 0.1 mol LiOH
E. 5. 0.4 mol HCOOH and 0.2 mol NaOH
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[LEGGETT--AP CHEMISTRY – MINIMAL FINAL REVIEW]
CHAPTER 10.
THERMOCHEMISTRY & THERMODYNAMICS
1. Define the following terms.
State Function
Internal Energy
Enthalpy
Entropy
Thermodynamically favorable
Heat capacity
Entropy driven
Enthalpy of formation
SATP
Endothermic
Surroundings
Calorimetry
2.
3.
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Work
Free energy
Specific heat capacity
Free Energy of formation
Exothermic
heat
spontaneous
Enthalpy driven
Standard “positional” entropy
System
Describe and explain what happens to the temperature of the surroundings during an exothermic process.
Describe and explain what happens to the temperature of the surroundings during an endothermic process.
ENDOTHERMIC
EXOTHERMIC
Verbally…
Energy is…_____________________
Energy is…_____________________
Observationally…
The temperature of the
surroundings…_____________________
The temperature of the
surroundings…_____________________
Symbolically…
∆H = _____________________
∆H = _____________________
In a Reaction…
_________+ A + B  C + D
A + B  C + D + __________
Graphically…
Predict the sign of the entropy change for each of the following processes: Remember that ∆ is always final - initial
(a) Fe(OH)3(s)  Fe3+(aq) + 3OH─(aq)
(b) O2(g) + 2H2(g)  2H2O(g)
(c) CO2(g) + H2O(l)  H2CO3(aq)
(d) Ag+(aq) + Cl-(aq)  AgCl(s)
LETS DO IT 1. Write the reactions for the enthalpy of formation for the following substances.
CO2(g)
Al2O3(s)
H2O(l)
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CALORIMETRY
•
The laboratory technique used to measure the heat released or absorbed during a chemical or physical change.
•
q - quantity of heat absorbed or released during a chemical or physical change
- is proportional to the change in temperature of the system being studied.
•
specific heat capacity (or specific heat) (c) - the quantity of heat needed to raise the temperature of 1 g
of a substance 1 K:
c = q/(m × ΔT ) or q
= cmΔT,--on equation sheet!!
where m is the mass of the substance.
units-- J/g(or mol).K.
Coffee Cup Calorimeter
Other ways of calculating ΔH reaction
1. Hess’ Law If several reactions add up to give an overall reaction, the ΔH of the reactions will add up to the overall ΔH reaction.
•
Standard Heats of Formation, ΔH f ° , are useful for this
purpose. This is the energy involved in making one mole of a
substance from its elements at 25°C and 1 atm pressure.
•
This law is often written as:
ΔH reaction.= ΣΔH f ° products - ΣΔH f ° reactants ON EQUATION SHEET!!!!!!
Note: ΔH f° for elements is 0.
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2. Using bond energy
•
Bond energy is the amount of energy needed to BREAK a
certain bond.
•
Breaking a bond is always endothermic
•
You can determine the approximate energy change in a
chemical reaction by taking the sum of the bond energies of the
reactants and subtracting from this the bond energies of the products.
Σ bond energies Reactants- Σ bond energies products
Note: this is opposite to the Hess’s Law (products -reactants) because bond energy involves breaking bonds
whereas ΔH f° involve forming bonds.
MC Practice
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g)
1.
For the reaction of methane represented above, H is -1,523 kJ mol-1. What is the value of H if the combustion produced liquid water
H2O(l), rather than water vapor, H2O(g)? (H for the phase change H2O(g)  H2O(l) is -44 kJ mol-1).
a. -1,479 kJ/mol
b. -1,567 kJ/mol
c. -1,323 kJ/mol
d. -1,611 kJ/mol
e.-1,411 kJ/mol
2. I2(g) + 3 Cl2(g)  2 ICl3(g)
According to the data in the table below, what is the value of Hof for the reaction represented above?
Bond
I-I
Cl-Cl
I-Cl
Average Bond Energy
(kilojoules/mol)
150
240
210
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a)
b)
c)
d)
e)
2015-2016
+450 kJ/mol
+1,260 kJ/mo
-870 kJ/mol
-390 kJ/mol
+180 kJ/mol
Entropy
•
Symbol is S (ΔS)
•
Unit J/K mol or J.mol-1K-1
•
The degree of disorder (randomness) in a system, or a refelction of the number of microstates a system has
•
Spontaneity tends toward + ΔS
•
The entropy of a pure crystalline substance at absolute zero is 0.
•
Larger and more complex molecules have greater entropies
•
gases have more entropy than liquids which have more than solids
•
solutions have more entropy than a pure solvent
FREE ENERGY
•
There is an upper limit to how much work can be derived
from any system. That maximum work is called the Gibb’s
Free Energy, ΔG.
•
The theoretical maximum work can be achieved if the steps
are so small that they can go either way, if they are
reversible.
•
A system that has a negative ΔG at a given temperature is spontaneous at that temperature
•
•
•
•
ΔG = ΔH – TΔS--On equation sheet!!!
This represents the theoretical maximum work that can be done by a system.
A reaction is spontaneous when ΔG is negative.
When ΔG = 0, the reaction is at equilibrium.
MAKE SURE THAT YOUR ΔH AND ΔS HAVE THE SAME UNITS (CONVERT J IN ΔS INTO kJ
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ΔH
+
+
ΔS
+
+
-
2015-2016
Spontaneous…
… at all temperatures
… at LOW temperatures
… at HIGH temperatures
..reverse rxn spontaneous
MC Practice
3. Which gives the thermodynamic parameters of the phase change in a system that begins as an open container of liquid water placed
in a constant temperature environment of 230 K?
I.
G < 0
II.
H < 0
III.
S < 0
a)
I only
b)
III only
c)
I and III only
d)
II and III only
e)
I, II and III
MC Practice
4.
H2(g) + O2(g)  H2O(l)
Ho = x
2 Na(s) + O2(g)  Na2O(s)
Ho= y
Na(s) + O2(g) + H2(g)  NaOH(s)
Ho= z
Based on the information above, what is the standard enthalpy change for the following reaction?
Na2O(s) + H2O(l)  2 NaOH(s)
a.
x+y+z
b.
x+y-z
c.
x + y - 2z
d.
2z - x - y
e.
z-x-y
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Thermochem Practice Free Response Problems
1. Lead iodide is a dense, golden yellow, slightly soluble solid. At 25C, lead iodide dissolves in water forming a system represented by the following
equation.
PbI2(s)  Pb2+ + 2 I-
H = +46.5 kilojoules
(a)
How does the entropy of the system PbI2(s) + H2O(l) change as PbI2(s) dissolves in water at 25C? Explain
(b)
If the temperature of the system were lowered from 25C to 15C, what would be the effect on the value of Ksp? Explain.
(c)
If additional solid PbI2 were added to the system at equilibrium, what would be the effect on the concentration of I- in the solution? Explain.
(d)
At equilibrium, G = 0. What is the initial effect on the value of G of adding a small amount of Pb(NO3)2 to the system at equilibrium?
Explain.
C2H2(g) + 2 H2(g)  C2H6(g)
2. Information about the substances involved in the reaction represented above is summarized in the following tables.
Substance
Hf (kJ/mol)
S (J/molK)
C2H2(g)
200.9
226.7
H2(g)
130.7
0
C2H6(g)
----
-84.7
Bond
Bond Energy (kJ/mol)
C-C
347
C=C
611
C-H
414
H-H
436
(a)
If the value of the standard entropy change, S, for the reaction is -232.7 joules per moleKelvin, calculate the standard molar entropy, S, of
C2H6 gas.
(b)
Calculate the value of the standard free-energy change, G, for the reaction. What does the sign of G indicate about the reaction above?
(c)
Calculate the value of the equilibrium constant, K, for the reaction at 298 K.
(d)
Calculate the value of the CC bond energy in C2H2 in kilojoules per mole.
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3.
CO(g) +
1
2
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) +
1
2
O2(g)  CO(g)
C(s) + O2(g)  CO2(g)
(b)
∆H˚298 = –110.5 kJ mol-1
∆H˚298 = –393.5 kJ mol-1
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.
S˚298
Substance
(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.
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The initial temperature of your water, the final temperature of your reaction are measured, and ΔT is calculated
Moles of reactants are calculated.
The specific heat of water is used (4.18 J/mol K)
q absorbed by water is calculated
qsurroundings = negative qsystem is assumed
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Getting from q to ΔH
Convert q Joules into KJ
divide by moles of reactant (from molarity(MxL) if solution and grams--> moles if solid )to get ΔH reaction!!!
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