AP CHEMISTRY FALL SEMESTER EXAM REVIEW
90 minute MC Section + 20 minute FRQ
Unit 2: Reactions and Aqueous Solutions
1. Complete combustion of a sample of a hydrocarbon in excess oxygen produces equimolar quantities of carbon
dioxide and water. What could be the molecular formula of the compound?
2. If the mass percent of an element in a pure sample is given, be able to determine why another given sample is at
a lower mass percent of the element.
3. If you have two beakers of the same volume with different compounds and you combine them into one beaker,
be able to predict what concentrations of each compound would give a precipitate.
4. When given different compounds, be able to identify which one would be an isomer, dissolves in acidic/basic
solutions, and solubility in water.
5. Calculate molarity of one ion if another ion’s concentration is given when dissolved in water.
6. Know weak/strong acids/bases Identify oxidizing and reducing agents.
7. Be able to calculate the amount of moles produced from combing two molecules.
8. Write a net ionic equation
9. Solubility rules
10. Calculate mole fraction
11. Calculate empirical formula
12. What do you rinse a buret with before you do a titration?
13. Know how to calculate the amount of water needed to make a dilution.
14. Calculate percent yield.
15. Predict solubility of salts in water.
16. What piece of laboratory glassware is needed for preparing a certain volume of an aqueous solution from a solid.
Unit 3: Gas Laws
17. Draw a particulate showing the formation of water vapor from hydrogen gas and oxygen gas in a rigid container.
18. If given the molar mass, temperature and pressure of a gas, be able to:
a. Calculate the average kinetic energy of a gas molecule
b. Calculate the density of a gas
c. If the pressure of each gas is increased at constant temperature until condensation occurs, which gas
will condense at the lowest pressure? Highest pressure?
19. If different containers of different gases have the same volume and temperature, be able to find the total pressure
once they are all combined if the ratio of atoms is given.
20. If a reaction were to go to completion, know how to calculate the total pressure in the container.
21. Based on temperature and pressure of different molecules, be able to determine which one behaves most like an
ideal gas and lowest/highest root mean square.
22. Gas Law calculations (Boyles, Charles, Gay-Lussac, Ideal)
23. Calculate number of moles produced in a gas collection tube.
24. Gas Lab: what measurements were needed to be able to calculate the molar mass of the gas?
25. Be able to calculate vapor pressure of a substance when given the mole fraction and vapor pressure of water.
Unit 4: Atomic Structure and Periodicity
26. When given ionic radius, be able to predict relative strength of the ion’s attraction to water and why.
27. If there is a mixture of two elements with similar average atomic masses, be able to know what would happen if
you ran them through a mass spectrometer. How would you determine if a different sample was purely one
element?
28. What would explain why one molecule can absorb UV but not visible light and another molecule could absorb
both?
29. When comparing two different element’s PES, be able to explain why one peak comes before another.
30. Based on periodic trends, be able to predict the atomic radius and first ionization of an element when given the
same information for other elements.
31. Know periodic trends for electronegativity, first-ionization energy, atomic radius.
Unit 5: Bonding
32. Lewis electron-dot diagrams, molecular geometry, bond angle and polarity
33. Predict the hybridization of a compound.
Unit 6: States of Matter, IMF, Solutions
34. Be able to rank liquids based on IMF when given vapor pressure
35. Be able to identify hydrogen bonding in water.
36. Explain why two molecules of different elements would have different boiling points (ie. type of solid, type of
bond, polarity, IMFs, electronegativity)
37. The dissolution of an ionic solute in a polar solvent can be imagined as occurring in three steps. In step 1, the
separation between ions in the solute is greatly increased, just as will occur when the solute dissolves in the polar
solvent. In step 2, the polar solvent is expanded to make spaces that the ions will occupy. In the last step, the ions
are inserted into the spaces in the polar solvent. Be able to describe the enthalpy change, ∆H, for each step.
38. Be able to relate IMF, frequency of collisions, number of molecules and speed to the pressure of a container
before and after equilibrium.
39. What are the covalent network solids and why would conductivity increase if another element is placed within
the network?
40. Know properties of solids: covalent network, molecular, ionic
41. Why would one element be stronger when alloyed with one element than a different element?
42. Be able to explain why different compounds would have different boiling points based on IMF, bonding, ionic
radius, or energy.
43. Know what phases exist where on a phase diagram.
44. Be able to explain why one substance is a liquid whereas another is a gas at the same temperature and pressure.
PRACTICE PROBLEMS
1. What is the mass of H2SO4 (molecular weight 98.1) in 50.0 milliliters of a 6.00-molar solution?
2. How many milliliters of 11.6-molar HCl must be diluted to obtain 1.0 liter of 3.0-molar HCl?
3. A 40.0 mL sample of 0.25 M KOH is added to 60.0 mL of 0.15 M Ba(OH)2. What is the molar concentration of OH¯(aq)
in the resulting solution? (Assume that the volumes are additive)
4. The volume of distilled water that should be added to 10.0 mL of 6.00 M HCl(aq) in order to prepare a 0.500 M HCl(aq)
solution is approximately ___?
5. What is the final concentration of barium ions, [Ba2+], in solution when 100. mL of 0.10 M BaCl2(aq) is mixed with
100. mL of 0.050 M H2SO4(aq)?
6. What number of moles of O2 is needed to produce 142 grams of P4O10 from P? (Molecular weight P4O10 = 284)
7. The alkenes are compounds of carbon and hydrogen with the general formula CnH2n. If 3.50 gram of any alkene is
burned in excess oxygen, what number of moles of H2O is formed?
8. 3 Ag(s) + 4 HNO3 → 3 AgNO3 + NO(g) + 2 H2O
The reaction of silver metal and dilute nitric acid proceeds according to the equation above. If 0.10 mole of powdered
silver is added to 10. milliliters of 6.0-molar nitric acid, what is the number of moles of NO gas that can be formed?
9. When 70. milliliter of 3.0-molar Na2CO3 is added to 30. milliliters of 1.0-molar NaHCO3, what is the resulting
concentration of Na+ ?
10. A 20.0-milliliter sample of 0.200-molar K2CO3 solution is added to 30.0 milliliters of 0.400-molar Ba(NO3)2 solution.
Barium carbonate precipitates. What is the concentration of barium ion, Ba2+, in solution after the reaction?
11. A 27.0-gram sample of an unknown hydrocarbon was burned in excess oxygen to form 88.0 grams of carbon dioxide
and 27.0 grams of water. What is a possible molecular formula of the hydrocarbon?
12. How many grams of calcium nitrate, Ca(NO3)2, contains 24 grams of oxygen atoms?
13. The mass of element Q found in 1.00 mole of each of four different compounds is 38.0 grams, 57.0 grams, 76.0 grams,
and 114 grams, respectively. A possible atomic weight of Q is
19. The simplest formula for an oxide of element X (MM = 76.0) that is 24.0 percent oxygen by weight is___.
14. BrO3¯ + 5 Br¯ + 6 H+  3 Br2 + 3 H2O
If 25.0 milliliters of 0.200-molar BrO3¯ is mixed with 30.0 milliliters of 0.450-molar Br¯ solution that
contains a large excess of H+, what is the amount of Br2 formed, according to the equation above?
15. How many moles of solid Ba(NO3)2 should be added to 300. milliliters of 0.20-molar Fe(NO3)3 to increase
the concentration of the NO3¯ ion to 1.0-molar? (Assume that the volume of the solution remains constant.)
16. Samples of F2 gas and Xe gas are mixed in a container of fixed volume. The initial partial pressure of the F2
gas is 8.0 atmospheres and that of the Xe gas is 1.7 atmospheres. When all of the Xe gas reacted, forming a
solid compound, the pressure of the unreacted F2 gas was 4.6 atmospheres. The temperature remained
constant. What is the formula of the compound?
17. It is suggested that SO2 (molar mass 64 grams), which contributes to acid rain, could be removed from a
stream of waste gases by bubbling the gases through 0.25-molar KOH, thereby producing K2SO3. What is
the maximum mass of SO2 that could be removed by 1,000. liters of the KOH solution?
18. When a 1.25-gram sample of limestone was dissolved in acid, 0.44 gram of CO2 was generated. If the rock
contained no carbonate other than CaCO3, what was the percent of CaCO3 by mass in the limestone?
19. ... Fe(OH)2 + ... O2 + ... H2O  ... Fe(OH)3
If 1 mole of O2 oxidizes Fe(OH)2 according to the reaction represented above, how many moles of Fe(OH)3
can be formed?
20. What mass of Au is produced when 0.0500 mol of Au2S3 is reduced completely with excess H2?
21. A 1.0 L sample of an aqueous solution contains 0.10 mol of NaCl and 0.10 mol of CaCl2. What is the
minimum number of moles of AgNO3 that must be added to the solution in order to precipitate all of the Cl¯
as AgCl(s) ? (Assume that AgCl is insoluble.)
22. When hafnium metal is heated in an atmosphere of chlorine gas, the product of the reaction is found to
contain 62.2 percent Hf by mass and 37.4 percent Cl by mass. What is the empirical formula for this
compound?
40. What is the ground state electron configuration for the Mn3+ ion?
41. Ca, V, Co, Zn, As: Gaseous atoms of which of the elements above are paramagnetic?
23. Two flexible containers for gases are at the same temperature and pressure. One holds 0.50 gram of hydrogen and the
other holds 8.0 grams of oxygen. Which of the following statements regarding these gas samples is FALSE?
(A) The volume of the hydrogen container is the same as the volume of the oxygen container.
(B) The number of molecules in the hydrogen container is the same as the number of molecules in the
oxygen container.
(C) The density of the hydrogen sample is less than that of the oxygen sample.
(D) The average kinetic energy of the hydrogen molecules is the same as the average kinetic energy of the
oxygen molecules.
(E) The average speed of the hydrogen molecules is the same as the average speed of the oxygen molecules.
24. Which of the following aqueous solutions has the highest boiling point?
(A) 0.10 M potassium sulfate, K2SO4
(B) 0.10 M hydrochloric acid, HCl
(C) 0.10 M ammonium nitrate, NH4NO3
(D) 0.10 M magnesium sulfate, MgSO4
(E) 0.20 M sucrose, C12H22O11
25. Which of the following does NOT behave as an electrolyte when it is dissolved in water?
(A) CH3OH
(B) K2CO3
(C) NH4Br
(D) HI
(E) Sodium acetate, CH3COONa
26. Which of the following has the lowest conductivity?
(A) 0.1 M CuS04
(B) 0.1 M KOH
(C) 0.1 M BaCl2
(D) 0.1 M HF
(E) 0.1 M HNO3
30. The elements in which of the following have most nearly the same atomic radius?
a. Be, B, C, N
b. Ne, Ar, Kr, Xe
c. Mg, Ca, Sr, Ba
d. C, P, Se, I
e. Cr, Mn, Fe, Co
36. All of the following statements concerning the characteristics of the halogens are true EXCEPT:
a. The first ionization energies (potentials) decrease as the atomic numbers of the halogens increase.
b. Fluorine is the best oxidizing agent.
c. Fluorine atoms have the smallest radii.
d. Iodine liberates free bromine from a solution of bromide ion.
e. Fluorine is the most electronegative of the halogens.
Next 2 questions
(A) Lithium
(B) Nickel
(C) Bromine
(D) Uranium
(E) Fluorine
42. Is a gas in its standard state at 298 K
43. Reacts with water to form a strong base
Ionization Energies for element X (kJ mol¯1)
First
Second
Third
Fourth
Five
580
1815
2740
11600
14800
44. 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. Na
b. Mg
c. Al
d. Si
e. P
45. Which of the following represents a pair of isotopes?
Atomic
Number
Mass
Number
(A)
I.
II.
6
7
14
14
(B)
I.
II.
6
14
7
14
(C)
I.
II.
(D)
I.
II.
7
7
13
14
(E)
I.
II.
8
16
16
20
6
14
14
28
Next 3 questions
a. 1s2 2s22p5 3s23p5
b. 1s2 2s22p6 3s23p6
c. 1s2 2s22p62d10 3s23p6
d. 1s2 2s22p6 3s23p63d5
e. 1s2 2s22p6 3s23p63d3 4s2
46. The ground-state configuration for the atoms of a transition element
47. An impossible electronic configuration
48. The ground-state configuration of a common ion of an alkaline earth element
Next 4 questions
a. Heisenberg uncertainty principle
b. Pauli exclusion principle
c. Hund's rule (principle of maximum multiplicity)
d. Shielding effect
e. Wave nature of matter
40.
41.
42.
43.
Can be used to predict that a gaseous carbon atom in its ground state is paramagnetic
Explains the experimental phenomenon of electron diffraction
Indicates that an atomic orbital can hold no more than two electrons
Predicts that it is impossible to determine simultaneously the exact position and the exact velocity of
an electron
Next 4 questions.
32.
33.
34.
35.
Represents an atom that is chemically unreactive
Represents an atom in an excited state
Represents an atom that has four valence electrons.
Represents an atom of a transition metal.
Next 3 questions
(A) F
(B) S
(C) Mg
(D) Ar
(E) Mn
36. Forms monatomic ions with 2¯ charge in solutions
37. Forms a compound having the formula KXO4
38. Forms oxides that are common air pollutants and that yield acidic solution in water
Next 3 questions
(A) O
(B) La
(C) Rb
(D) Mg
(E) N
38. What is the most electronegative element of the above?
39. Which element exhibits the greatest number of different oxidation states?
40. Which of the elements above has the smallest ionic radius for its most commonly found ion?
39. Given that a solution is 5 percent sucrose by mass, what additional information is necessary to calculate the molarity
of the solution?
I. The density of water
II. The density of the solution
III. The molar mass of sucrose
40. Given that a solution is 5 percent sucrose by mass, what additional information is necessary to calculate the molarity
of the solution?
I. The density of water
II. The density of the solution
III. The molar mass of sucrose
AP CHEMISTRY FREE-RESPONSE QUESTION POSSIBILITIES
20 minutes
1. Account for each of the following in terms of principles of atom structure, including the number, properties, and
arrangements of subatomic particles.
a. The second ionization energy of sodium is about three times greater than the second ionization energy of
magnesium.
b. The difference between the atomic radii of Na and K is relatively large compared to the difference
between the atomic radii of Rb and Cs.
c. A sample of nickel chloride is attracted into a magnetic field, whereas a sample of solid zinc chloride is
not.
d. Phosphorus forms the fluorides PF3 and PF5, whereas nitrogen forms only NF3. [actually a bonding
moment]
2. Methanamide, CH3NO, is a liquid at 25°C.
a. The complete Lewis electron-dot diagram for methanamide is shown below.
i. In the molecule, angle x is not 180°. Estimate the observed angle. Justify your answer.
ii. In the molecule, angle y is not 90°. Explain why in terms of electron domains (VSEPR
model).
b. Consider a molecule with the formula CH2O2 . The structure of this molecule has a geometry
around the carbon atom similar to the geometry around carbon in methanamide. In the box
provided below, draw the complete Lewis electron-dot diagram for the molecule.
3. The structures of a water molecule and a crystal of LiCl(s) are represented above. A student prepares a
1.0 M solution by dissolving 4.2 g of LiCl(s) in enough water to make 100 mL of solution.
a. In the space provided below, show the interactions of the components of LiCl(aq) by making a
drawing that represents the different particles present in the solution. Base the particles in your
drawing on the particles shown in the representations above. Include only one formula unit of
LiCl and no more than ten molecules of water. Your drawing must include the following details.
i.
identity of ions (symbol and charge)
ii. the arrangement and proper orientation of the particles in the solution
4. Ethene, C2H4(g) (molar mass 28.1 g/mol), may be prepared by the dehydration of ethanol, C2H5OH(g) (molar
mass 46.1 g/mol), using a solid catalyst. A setup for the lab synthesis is shown in the diagram above. The
equation for the dehydration reaction is given below.
C2H5OH(g)
ethanol
catalyst
C2H4(g) + H2O(g)
ethane
water
A student added a 0.200 g sample of C2H5OH(l) to a test tube using the setup shown above. The student
heated the test tube gently with a Bunsen burner until all of the C2H5OH(l) evaporated and gas
generation stopped. When the reaction stopped, the volume of collected gas was 0.854 L at 0.822 atm
and 305 K. (The vapor pressure of water at 305 K is 35.7 torr).
a. Calculate the number of moles of C2H4(g)
i.
That are actually produced in the experiment and measured in the gas collection tube
ii.
That would be produced if the dehydration reaction went to completion.
b. Calculate the percent yield of C2H4(g) in the experiment.
c. The Lewis electron-dot diagram for C2H4 is shown in the box on the left. In the box on the right, complete
the Lewis electron-dot diagram for C2H5OH by drawing in all of the electron pairs.
d. What is the approximate value of the C-O-H bond angle in the ethanol molecule?
e. During the dehydration experiment, C2H4(g) and unreacted C2H5OH(g) passed through the tube into the
water. The C2H4 was quantitatively collected as a gas, but the unreacted C2H5OH was not. Explain this
observation in terms of the intermolecular forces between water and each of the two gases.
Mg(s) + 2 H+(aq) --- Mg2+(aq) + H2(g)
5. A student performs an experiment to determine the volume of hydrogen gas produced when a given mass of
magnesium reacts with excess HCl(aq) , as represented by the net ionic equation above. The student begins with
a 0.0360 g sample of pure magnesium and a solution of 2.0 M HCl(aq) .
a. Calculate the number of moles of magnesium in the 0.0360 g sample.
b. Calculate the number of moles of HCl(aq) needed to react completely with the sample of magnesium.
As the magnesium reacts, the hydrogen gas produced is collected by water displacement at 23.0�C. The pressure
of the gas in the collection tube is measured to be 749 torr.
c. Given that the equilibrium vapor pressure of water is 21 torr at 23.0�C, calculate the pressure that the
H2(g) produced in the reaction would have if it were dry.
d. Calculate the volume, in liters measured at the conditions in the laboratory, that the H2(g) produced in
the reaction would have if it were dry.
e. The laboratory procedure specified that the concentration of the HCl solution be 2.0 M, but only
12.3 M HCl solution was available. Describe the steps for safely preparing 50.0 mL of 2.0 M HCl(aq)
using 12.3 M HCl solution and materials selected from the list below. Show any necessary calculation(s).
10.0 mL graduated cylinder
250 mL beakers
50.00 mL volumetric flask
Distilled water
Balance
Dropper