Download Chemistry Final Exam Study Guide_S2014

Chemistry Final Exam Study Guide
STRATEGY: Start by reading through your notes to refresh your memory on these topics. Then, use this review
sheet as a starting point to identify the areas on which you need to spend more study time. For those areas, go back to
homework assignments, quizzes, and reviews to practice more problems. Don’t just answer the questions you know! If
you don’t know an answer, look it up in your notes, book or online. Also, these are not the questions that will be on
the test, you need to be able to understand WHY an answer is correct, and not just WHAT the answer is…copying
won’t help you on the Final Exam!
Introduction Unit (Ch. 1-4)
1. In a lab, the average measured density for Pre-1982 pennies was 7.98 g/cm3. Given that the literature value for
the density is 8.92 g/cm3, calculate the percent error.
2. Convert the following numbers into or out of scientific notation.
a. 548,000
b. 0.0000770 c. 1.200 × 10-3 d. 9.25 × 107
3. Osmium is the densest element with a density of 22.57 g/cm3. Find the mass of a 56.2 cm3 sample of osmium.
4. Perform the following SI prefix conversions.
a. 65.2 mm = ? cm b. 2.3 kg = ?g c. 65,000 ml = ? L
d. 0.502 km = ? cm
5. How many milliliters are in a 2.0 quart jug of milk? (1L = 1.06qts)
6. Ms. Townsend spent last weekend grading tests. If she spent 4 min on each test, how many hours did it take
her to grade all 74 tests?
7. Calculate the density from the slope of a "Mass vs. Volume" graph.
8. Write the isotope symbol, including atomic number & mass number, for the following isotopes.
a. carbon-14
b. chromium-53
c.
nickel-63
d. zirconium-92
9. Complete the table for the following isotopes.
Symbol
Atomic #
Mass#
Protons
Neutrons
electrons
Zn
65
20
21
74
34
40
18
10. Calculate the average atomic mass of copper if 69.17% of the copper atoms occurring in nature are 63Cu and
30.83% are 65Cu.
Atoms/Electrons (Ch. 5)
11. Draw an orbital diagram, complete electron configuration and noble gas notation for:
a. Na
b. C
c. Mo
d. Se
12. How does an electron become excited? What does it do when it returns to the ground state?
13. What is a photon? Quantum?
14. Describe the relationship between wavelength and frequency.
15. What type of electromagnetic radiation is emitted when an electron moves from n=4 to n=1?
16. Where are the s,p,d,f blocks located on the periodic table? How many orbitals are associated with each? How
many electrons can be in each orbital?
Periodic Table (Ch. 6)
17. DEFINE each of the following trends, describe WHAT the trend is (increases up, etc), explain WHY each
trend occurs
a. Atomic Radius
b. Ionic Radius
c. Ionization Energy
d. Electronegativity
18. Circle the atom with the LARGER radius.
a. Be
N
b. Ne
Xe
19. Circle the particle with the LARGER radius.
a. Cl
Cl –
b. Mg
Mg2+
20. Circle the atom with the HIGHER first ionization energy.
a. Li
Cs
b. Ca
As
21. Circle the atom with the HIGHER electronegativity.
a. Cl
Si
b. O
Po
Bonds/Names (Ch. 8&9)
22. Based on their electronegativities (p. 169), are the bonds in the following substances IONIC, POLAR, or
NONPOLAR?
a. MgO
b. LiCl
c. H2O
d. Br2
23. 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.
24. Use Lewis Diagrams to show the formation of MgO.
25. Use Lewis Diagrams to show the formation of H2O.
26. Explain the relationship between potential energy and stability.
27. 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.
lithium phosphate
f.
ammonium carbonate
28. Write names for the following compounds (HINT: First determine ionic/acid/covalent).
a.
CrCl3
d.
MgSO4
b.
Cu2CO3
e.
P4O6
c.
AsCl5
f.
NaClO3
29. Explain the difference between nonpolar covalent, polar covalent, and ionic bonds in terms of sharing of
electrons and electric charge.
Molecular Geometry (Ch. 9&13.2)
30. Explain the main idea of the VSEPR Theory
31. For each of the following molecules, draw the Lewis structure, name the number of lone and bond pairs, and
give the name of the shape.
a. BeCl2
b. AlCl3
c. CF4
d. H2O
e. NH3
32. For each of the molecules in the previous question, draw the dipole moments, state whether the molecule is
polar or nonpolar, then list any intermolecular forces that each molecule will have.
33. What type of IMF is indicated by each statement? Hint: IMF’s are: hydrogen bonding, dipole-dipole and
London dispersion
a. Attraction between any two polar molecules
b. Very weak force that increases with molar mass
c. Attraction between two momentary (or temporary) dipoles
d. Very strong attractive force between molecules with N-H, O-H or F-H bonds.
Moles (Ch. 11)
34. How many magnesium sulfate molecules are in 25g?
35. Find the percent composition of each element in copper (II) chloride
36. The percent composition of a compound is 40%C, 6.7% H and 53.7%O. The molecular mass of the
compound is 180.0g/mol. Find its empirical and molecular formulas.
37. How many moles are in 5g of H2SO4?
38. How many grams of H2 is in 34L of H2?
39. 2.3 moles of Ca(OH)2 is equal to how many grams?
40. How many liters will 3.4 mol of O2 occupy?
Reactions (Ch. 10)
41. Rewrite and balance the following word equation using chemical formulas, states, and energy. When solid
sodium chlorate absorbs energy, it produces solid sodium chloride and oxygen gas.
42. Predict the products of the following single replacement reactions using the Activity Series in your Reference
Tables. If no reaction occurs, write NR. If a reaction occurs be sure to write the correct formulas and balance.
c. Fe + Na3PO4
a. K + ZnCl2
b. Cl2 + HF
43. Write balanced, complete and net ionic equations for the following pairs of compounds. Use the Solubility
Rules in your Reference Tables.
a. Potassium iodide and silver nitrate
b. Ammonium phosphate and sodium sulfate
44. Balance the following equations. Also classify as synthesis, decomposition, single replacement, double
replacement or combustion
a. Cu + MgSO4  CuSO4 + Mg
d. Ca(NO3)2 + Na2SO4  CaSO4 +
NaNO3
b. C5H12 + O2  CO2 + H2O
e. H2 + O2  H2O
c. Fe2O3  Fe + O2
45. Identify each as endothermic or exothermic:
a. PE of products is lower than PE of reactants
b. PE of products is higher than PE of reactants
c. When substances are mixed, the test tube feels cold
d. In your car’s engine, fuel is burned to produce energy
46. List 3 conditions required for a successful collision according to the kinetic molecular theory.
Stoichiometry (Ch. 12)
47.
48.
49.
50.
How many grams of copper would be produced from 49.48g of chromium? Cr + CuSO4  Cu + Cr2(SO4)3
How many grams of chromium are required to react with 5 mole of CuSO4?
How many grams of ZnS are required to react with 12.6L of oxygen gas at STP? ZnS + O2  ZnO + SO2
6.45g of lithium reacts with 9.2g of oxygen gas to produce lithium oxide.
a. How many grams of Li2O are formed? Hint: write the balanced chemical equation first.
b. The actual yield is 12.5g. What is the percent yield of this reaction?
Gas Laws (Ch. 14)
51. Explain why each of the examples occurs (explain what variable changes (pressure, temperature and
volume) and what the result of that change is.)
a. A balloon pops after floating high into the atmosphere
b. A balloon pops in a hot car on a summer day.
c. An aerosol can label that reads: “Do not store aerosol cans at temperatures above 120F. Danger of
explosion”
52. A jar is tightly sealed at 22C and 772mmHg. What is the pressure inside the jar after it has been heated to
178C?
53. What is the volume of 5 mole of He at 50C and 1.2 atm?
54. H2 and He gas are collected in a vessel and have a total pressure of 5.9atm. If the partial pressure of He is
1.8atm, what is the partial pressure of H2?
55. 3 mol of gas occupies a volume of 50L. If 1.3 mole of gas is added at constant pressure and temperature,
what is the new volume?
56. 300mL of a gas has a pressure of 75kPa. When the volume is decreased to 125mL, what is its pressure?
57. 50L of gas has a temperature of 75C. What is the temp in Celsius when the volume changes to 110L?
58. A gas occupies 325L at 25C and 98kPa. What is its volume at 70kPa and 15C?
Solutions, Acids/Bases (Ch. 15&19)
59. Explain the effect of adding more solute to unsaturated, saturated and supersaturated solutions.
60. Explain how temperature and pressure affect solubility of gases dissolved in water and of solids dissolved in
water.
61. State whether each pair is soluble or insoluble.
a. KCl in water
c. Wax in benzene
b. Ammonia in oil
d. CH4 in water
62. How many grams of AlCl3 are required to make a 2.25M solution in 3L of water?
63. What volume of 12M HCl is needed to prepare 250mL of 0.20M HCl?
64. State whether each characteristic indicates an acid, a base or both.
a. Have a sour taste
c. Are corrosive
e. Turn blue litmus
b. React with metals
d. Feel slippery
paper red
65. Define acids and bases according to Arrhenius and Bronsted-Lowry.
66. Identify each substance as an acid, a base, a conjugate acid and conjugate base: H2S + H2O  HS- + H3O+
67. What is the pH and pOH of each solution, then classify as acidic or basic:
a. [H+] = 1x10-13
b. [OH-] = 1x10-6
c. [H+] = 1x10-6
68. When a neutralization reaction between a strong acid and a weak base reaches the equivalence point, will the
solution be acidic, basic or neutral?
69. If 43.5mL of 0.15M HBr is required to neutralize 25mL of Ca(OH)2, what is the molarity of Ca(OH)2?
Nuclear Chemistry (Ch. 25)
70. Match each description with the appropriate type of radiation (alpha, beta, positron, gamma):
a. A negatively charged electron
d. Blocked by paper
b. Blocked only by several feet of
e. Radiation energy with no electrical
concrete
charge
c. A positively charged particle stopped
by lead
71. Write equations for the nuclear decay of the following:
a. Decay of polonium-218 by alpha
c. Decay of carbon-14 by beta emission
emission
d. Decay of chlorine-32 by position
b. Decay of sodium-22 by electron
emission
capture
72. Carbon-14 has a half-life of 5730 years. If a plant contained 2.0g Carbon-14 when it died, how much is left
after 34380 years?
73. Compare and contrast nuclear fission and nuclear fusion.
Energy and Thermodynamics (Ch.16&17)
74.
75.
76.
77.
78.
79.
80.
81.
82.
83.
84.
85.
86.
87.
What is the relationship between temperature and kinetic energy?
What is specific heat? How does specific heat affect an object?
What is energy?
Compare/contrast endothermic and exothermic reactions. Include the following: definition, sign of ΔH,
temperature change, where heat is written in a chemical equation, enthalpy of products and reactants,
reaction pathway diagram.
What is the universe? System? Surroundings?
What is entropy? How can you tell if a substance has high or low entropy?
From each pair, pick the one with the greater entropy:
a. NH3(g) and NH3(l)
c. CO2(s) and CO2(g)
b. NH3(g) and NH3(aq)
d. KBr(s) and KBr(l)
State whether the reaction shown indicates an increase or a decrease in entropy:
a. C2H5OH(s)  C2H5OH(l)
f. PCl5(g)PCl3(g) +Cl2(g)
b. H2O(g)  H2O(l)
g. 2CO(g) + O2  2CO2(g)
c. CH3OH(l)  CH3OH(g)
h. C6H12O6(s)2C2H5OH(l) + 2CO2(g)
d. NH3(l) NH3(g)
i. H2SO4(l)H2O(l) + SO3(g)
e. 2SO3(g) + CO2(g)  CS2(g) +4O2(g)
State whether the following processes are exothermic or endothermic:
a. C2H5OH (l)  C2H5OH(g)
d. NaCl(s)NaCl(l)
b. NH3(g)  NH3(l)
e. C5H12(g) + O2(g)CO2(g) + H2O(g)
c. Br2(l)Br2(s)
If the temperature of 34.4g of ethanol increases from 25C to 78.8C, how much heat has been absorbed by
the ethanol?
Calculate the heat required to melt 25.7g of solid water.
How much energy is released when 25g of water vapor condense to form a liquid?
Draw a heating curve. Label the following: solid, liquid, gas, melting point, boiling point, kinetic energy
increasing, kinetic energy constant, potential energy increasing, and temperature increasing.
Draw a potential energy diagram for an endothermic reaction and one for an exothermic reaction. Label the
reactants, products, activated complex, activation energy, and draw a dotted line to represent how the
reaction pathway would be different if a catalyst is used.