Pre-AP Chemistry Review – May 2006
Measurement – Ch. 2
1. How many sig figs are in the following numbers?
a. 2.35
b. 34,000
c. 89.70
d. 0.0052
2. Osmium is the densest element with a density of 22.57 g/cm3. Find the mass of a 56.2 cm3 sample of osmium.
3. Perform the following SI prefix conversions.
a. 65.2 mm = ______dm
b. 2.3 kg = ______g
c. 65,000 mL = ______mL
d. 0.502 km = ______cm
4. How many milliliters are in a 2.0 quart jug of milk?
5. Mrs. J. spent last weekend grading lab notebooks. If she spent 5.5 min on each notebook, how many hours did it take her to grade all 95 notebooks?
Matter – Ch. 1
6. Classify the following substances as solid, liquid, gas, or plasma based on their properties.
a. flexible volume, high KE, particles can disperse freely.
b. flexible volume, very high KE, particles are charged.
c. fixed volume, very low KE, orderly particles.
d. fixed volume, low KE, particles can move past each other.
7. Compare and contrast a solution, colloid, and suspension.
8. Classify the following as element, compound, heterogeneous mixture, or solution.
a. graphite (carbon)
b. grape juice
c. table salt (NaCl)
d. pepper
9. Classify the following as chemical or physical changes.
a. cutting wire
b. ripening tomato
c. apple slices turning brown
d. compressing a gas
Atomic Structure – Ch. 3
10. Complete the table for the following isotopes.
Symbol
/ ZnAtomic # / 20
Mass # / 65 / 74 / 40
# of protons / 34
# of neutrons / 21
# of electrons / 18
11. What are isotopes?
Nuclear Chemistry – Ch. 22
12. Match each description with the appropriate type of radiation – alpha, beta, positron, or gamma.
a. A negatively charged electron.
b. Blocked only by several feet of concrete.
c. A positively charged particle stopped by lead.
d. Blocked by paper or clothing.
e. Radiation energy with no electrical charge.
13. Write an equation for the decay of polonium-218 by alpha (a) emission.
14. Write an equation for the decay of carbon-14 by beta (b-) emission.
15. Carbon-14 has a half-life of 5,730 years. If a plant contained 2.0 g of 14C when it died, how much is left after 11,460 years?
16. Define fission and fusion.
Electrons in Atoms – Ch. 4 & 5
17. Describe how Bohr’s model explains the bright lines (red, green, violet, violet) in the emission spectrum of hydrogen.
18. What is the primary difference between the modern model of the atom and Bohr’s model?
19. Draw orbital diagrams for the following elements.
Symbol / Atomic # /Orbital Diagram
FV
20. Give the shorthand electron configuration for the following.
Symbol / # e- /Shorthand e- Configuration
PdAt
21. Predict the ions that will form from the following atoms and give the shorthand configuration of the ion.
Atom / Ion /Noble Gas
/Shorthand e- Configuration
RbTe
22. What is a photon?
Periodic Table – Ch. 5
23. How did Mendeleev and Mosely arrange the elements in the periodic table?
24. Circle the atom with the larger radius.
a. Ra N b. Ne Xe
25. Circle the atom with the HIGHER first ionization energy.
a. Li Cs b. Ba As
26. Why are the Noble Gases inert?
Chemical Bonding – Ch. 6 & 7
27. Based on their electronegativities (p.151), are the bonds in the following substances IONIC, POLAR, or NONPOLAR?
a. MgO c. LiCl
b. H2O d. Br2
28. Are the following properties characteristics of ionic, covalent, or metallic bonding?
a. These bonds are formed by delocalized electrons in an “electron sea.”
b. These bonds involve a transfer of electrons.
c. Substances containing these bonds are malleable and have very high melting points.
d. Substances containing these bonds do not conduct electricity and have low melting points.
e. Compounds containing these bonds have a crystal lattice structure.
f. These bonds are formed by sharing electrons.
29. Explain the relationship between potential energy and stability.
30. Write formulas for the following compounds (HINT: First determine ionic/acid/covalent).
a. calcium bromide d. silicon dioxide
b. iron(III) sulfate e. dinitrogen tetroxide
c. hydrofluoric acid f. sulfurous acid
31. Write names for the following compounds (HINT: First determine ionic/acid/covalent).
a. CrCl3 d. MgSO4
b. CaO e. P4O6
c. AsCl5 f. HClO3
Molecular Structure – Ch. 6
32. For each of the following molecules, draw the Lewis electron dot diagram, give the shape and bond angle(s), and state whether the molecule is polar or nonpolar. Show your work in the spaces provided for counting valence e- and e- pairs.
Number of valence e- / Lewis Diagram / Countinge- pairs / Molecular Shape
& Bond Angle(s) / Molecular Polarity
AsBr3
SiO2
33. Draw the dipole moments for each bond in the following molecules and circle whether the molecule is polar or nonpolar.
TeCl2 bent / BCl3 trigonal planar / CH2Cl2 tetrahedralpolar nonpolar / polar nonpolar / polar nonpolar
The Mole – Ch. 3 & 7
34. How many magnesium sulfate molecules are in 25.0 g?
35. Find the molarity of a 750 mL solution containing 346 g of potassium nitrate.
36. Find the % composition of copper(II) chloride.
37. The percent composition of a compound is 40.0% C, 6.7% H, and 53.7% O. The molecular mass of the compound is 180.0 g/mol. Find its empirical and molecular formulas.
Chemical Reactions – Ch. 8
38. Write and balance the following word equation using chemical formulas, physical states, and energy. – When solid sodium chlorate absorbs energy, it produces solid sodium chloride and oxygen gas.
39. Predict the products and balance. Write N.R. if no reaction will occur. Include physical states for extra credit.
a. Cu(s) + MgSO4(aq) ®
b. C5H12(l) + O2(g) ®
c. NH4Cl(aq) + Pb(NO3)2(aq) ®
d. Fe2O3(s) ®
40. For each of the reactions in #39, specify whether it is combustion, synthesis, decomposition, single replacement, or double replacement.
a.
b.
c.
d.
41. Name four ways to increase the rate of a reaction.
42. Define endothermic and exothermic reactions.
Stoichiometry – Ch. 9
43. How many grams of copper would be produced from 49.48 g of chromium? ___Cr + ___CuSO4 ® ___Cu + ___Cr2(SO4)3
44. How many grams of chromium are required to react with 125 mL of 0.75M CuSO4. (same reaction as #43)
45. How many grams of ZnS are required to react with 12.6 L of oxygen gas at STP? ___ZnS + ___O2 ® ___ZnO + ___SO2
46. 6.45 g of ZnS reacts with 9.20 g of oxygen gas to produce zinc oxide. How many grams of ZnO are formed?
47. What are the limiting and excess reactants in #46?
48. The actual yield of the reaction in #46 is 4.42 g. What is the percent yield of this reaction?
Gases – Ch. 10 & 11
49. Identify the gas law that explains each situation – Boyle’s Law, Charles’ Law, or Gay-Lussac’s Law
a. Do not store aerosol cans at temperatures above 120°F. Danger of explosion.
b. A balloon pops after floating high into the atmosphere.
c. A balloon pops in a hot car on a summer day.
For the following problems, Identify the gas law (Boyle’s, Charles’, Gay-Lussac’s, combined, ideal) and calculate the answer.
50. A jar is tightly sealed at 22°C and 772 torr. What is the pressure inside the jar after it has been heated to 178°C?
51. 300.0 mL of gas has a pressure 75.0 kPa. When the volume is decreased to 125.0 mL, what is its pressure?
52. 50.0 L of gas has a temperature of 75°C. What is the temp in Celsius when the volume changes to 110 L?
53. A gas occupies 325 L at 25°C and 98.0 kPa. What is its volume at 70.0 kPa and 15°C?
54. What is the volume of a container that holds 48.0 g of helium at a pressure of 4.0 atm and temperature of 52°C?
Liquids & Solids – Ch. 12
55. Identify which intermolecular force is being described – dispersion forces, dipole-dipole forces, or hydrogen bonding.
a. Attraction between any two polar molecules.
b. Very weak force that increases with molar mass.
c. Attraction between two momentary dipoles.
d. Very strong attractive force between molecules with N-H, O-H, or F-H bonds.
56. Identify the type(s) of intermolecular forces present in the following molecules.
a. CH4
b. SCl2
c. F2
d. NH3
57. Identify which type of solid is being described – ionic, metallic, covalent molecular, covalent network, or amorphous.
a. Every atom is covalently bonded to another atom.
b. Atoms are surrounded by a sea of electrons.
c. Particles are connected only by IMF.
d. There is no geometric pattern in the structure.
e. Charged particles in a geometric pattern.
58. Explain the relationship between the strength of intermolecular forces and the boiling point of a substance.
59. Use the vapor pressure graph to answer the following questions.
a. What is the vapor pressure of CCl4 at 70°C?
b. What is the boiling point of water at 20 kPa?
c. What is the normal boiling point of CHCl3?
60. Assign the correct letter(s) from the heating curve below to each of the following descriptions.
a. liquid is boiling
b. solid is warming
c. solid is melting
d. potential energy is increasing
e. kinetic energy is increasing
Pre-AP Chemistry Review – May 2006
ANSWER KEY
9. a. physical, b. chemical, c. chemical, d. physical
10.
/Symbol
/Zn
/ Ca / Se / ArAtomic # / 30 / 20 / 34 / 18
Mass # / 65 / 41 / 74 / 40
# of protons / 30 / 20 / 34 / 18
# of neutrons / 35 / 21 / 40 / 22
# of electrons / 30 / 20 / 34 / 18
13.
14. /
15. 0.50 g
20. / Pd / 46 / [Kr] 5s24d8At / 85 / [Xe] 6s24f145d106p5
30. a. CaBr2, b. Fe2(SO4)3, c. HF, d. SiO2, e. N2O4, f. H2SO3
31. a. chromium(III) chloride, b. copper(I) carbonate, c. arsenic pentachloride, d. magnesium sulfate,
e. tetraphosphorous hexoxide, f. chloric acid.
e- pairs / Molecular Shape
& Bond Angle(s) / Molecular Polarity
AsBr3 / 1(5)+3(7)=26 / / 3B, 1L / trigonal pyramidal, 107° / polar
SiO2 / 1(4)+2(6)=16 / / 2B, 0L / linear, 180° / nonpolar
34. 1.25 ´ 1023 molecules MgSO4
35. 4.6M KNO3
36. OMIT
37. OMIT
38. 2NaClO3(s) 2NaCl(s) + 3O2(g)
39. OMIT – review balancing equations
40. a. single replacement, b. combustion, c. double replacement, d. decomposition
41. increase the surface area by grinding or dissolving the solid in water, increase the concentration of the reactants, increase the temperature of the reactants, use a catalyst
42. Exothermic reactions release energy because products have a lower PE than reactants. Endothermic reactions absorb energy because products have a higher PE than reactants.
43. 2Cr + 3CuSO4 ® 3Cu + Cr2(SO4)3, 90.71 g Cu
44. 3.3 g Cr
45. 2ZnS + 3O2 ® 2ZnO + 2SO2, 36.5 g ZnS
46. OMIT
47. OMIT
48. OMIT
49. a. Gay-Lussac’s Law (T P), b. Boyle’s Law (P¯ V), c. Charles’ Law (T V)
50. Gay-Lussac, 1180 torr
51. Boyle, 180. kPa
52. Charles, 490°C
53. Combined, 440. L
54. Ideal, 80. L
55. a. dipole-dipole, b. dispersion forces, c. dispersion forces, d. hydrogen bonding
56. a. dispersion; b. dispersion, dipole-dipole; c. dispersion; d. dispersion, dipole-dipole, hydrogen bonding
57. OMIT
58. Stronger IMF result in a higher boiling point because it takes more energy to break the strong attraction between molecules.
59. a. 80 kPa, b. 60°C, c. 62°C
60. a. liquid is boiling=d, b. solid is warming=a; c. solid is melting=b; d. PE is increasing=b,d; e. KE is increasing=a,c,e