Download Chemistry EOC Review Spring 2013

Chemistry EOC Review
Spring 2013
Directions: The following is an End-Of-Course Review Guide designed to assist you as prepare for your EOC.
It is imperative that you complete this guide to the best of your ability. This will help you to achieve a higher
average on your second quarter grade.
Unit 1 & 2 Lab Safety, Introduction to Chemistry, & Measurement (Chapters 1 & 3):
1. What is chemistry?
2. Distinguish between hypothesis, theory, and law.
3. Classify the following as having good or poor accuracy and good or poor precision:
4. A scientist experimentally determines the speed of light to be 2.98 x 108 m/sec. In a second experiment,
she determines the speed to be 2.99 x 108 m/sec.
5. The actual concentration of a solution is found to be 1.5 M. A scientist finds the concentrations to be
5.0 M and 5.2 M in two separate experiments.
6. Write the following in scientific notation:
a. 0.000 85
b. 1, 000, 000
c. 0.4565
7. How many significant figures are in the following:
a. 0.98
b. 1, 000
c. 0.3450
d. 112
d, 1,98 x 102
8. Report you answer to the correct number of significant figures:
a. 0.56 + 2.345
b. 9.45 – 8.4
c. 98.7 x 18.34
d. 1.23 / 3
9. Explain how the measurement 45 cm is different from 45.0 cm.
10. List the SI units for the following quantities:
a. Mass
b. Temperature
c. Volume
11. How many meters are in the following:
a. 1 km
b. 1 cm
d. 1 mm
d. Length
12. Calculate the density of a 5.0 g object that has a volume of 2.0 cm3.
13. Convert the following temperatures: a. 34C to K b. 50 K to C
14. How will the graphs of a direct and indirect (inverse) proportion differ from each other?
15. Distinguish between periods and families.
16. Which family on the periodic table is:
a. the most active metals
b. the most active nonmetals
17. Distinguish between metals and nonmetals.
18. Which metalloid is used in computer chips?
19. How are the noble gases different from other families?
Unit 3 Matter & Atomic Structure (Chapter 2 & 4) (Lots of good questions may come from this section):
20. List the characteristics of the four states of matter.
21. How are physical properties different from chemical properties? List two examples of each type of
property.
22. What are intensive or extensive properties? List two example of each type of property.
23. Classify the following as homogeneous mixture, heterogeneous mixture, element, compound, intrinsic
or extensive.
a. air
b. salt water
c. aluminum foil
d. table salt
e. dirt
f. Amt. of Substance
g. color
24. Explain what the early Greeks believed about the atom (Democritus and Aristotle)
25. What did Dalton believe about the atom?
26. Describe Millikan’s experiment and what he determined.
27. What did Rutherford determine about the atom?
28. Describe Rutherford’s experiment.
29. Complete the following chart:
Particle
Proton
Neutron
Location
Charge
Mass
Discover
Electron
30. Distinguish between the mass number and the atomic number:
Atomic
Mass Number Number of
Number of
Number of
Number
Protons
Neutrons
Electrons
9
10
16
8
108
47
35
45
31. All isotopes have the same number of _____________ while all isotopes differ in their number of
________________.
32. The basis for the atomic mass unit (amu) is ___________________________.
Unit 5 Electrons Configuration, Quantum Theory, Periodic Table (Chapter 5 & 6)
33. Compare and contrast the following four atomic models:
a. Thomson
b. Rutherford
c. Bohr
d. Quantum Mechanical
34. How are frequency and wavelength related?
35. Calculate the wavelength of a yellow light by a sodium lamp if the frequency of the radiation is 3.34 x 1014 Hz.
36. What is the energy associated with the photon in problem above?
37. High energy electrons are found _________________ while low energy electrons are found ____________.
38. The Heisenberg Uncertainity Principle states __________________________________________________.
39. List the four atomic orbitals and the maximum number of electrons each orbital can hold.
40. Write electron configurations for the following:
a. He
b. Na
e. Br
c. P
41. Which blocks (s, p, d, or f) will the following elements fall into:
a. U
b. Fe
c. K
d. I
d. Fe
f. Mt
Unit 4 Periodic Table & Trends, Bonds (Chapter 7 & 8):
42. The modern periodic table is arranged according to ______________________________.
43. Who is considered to be the father of the periodic table?
44. Describe the location of the following groups on the periodic table:
a. metals
b. transition metals
d. metalloids
e. nonmetals
c. halogens
45. Metals form _________ (type of ions) while nonmetals form ______________ (type of ions).
46. How does the atomic radius change moving across the table?
47. List he following in terms of increasing atomic radius: Cs, F, O, Ir
48. What is ionization energy?
49. What is electronegativity?
50. Arrange in terms of increasing ionization energy: H, K, Si, O
51. How is the name given to the following groups:
a. Group I
b. Group II
c. Group 17
52. How is an ionic bond different from a covalent bond?
53. How is a polar covalent different from a nonpolar bond?
54. Classify each bond as ionic, polar, or nonpolar:
a. H – Cl
b. C – O
c. Mg – O
d. F – F
55. List the seven diatomic molecules.
56. What are valence electrons?
57. What is the Octet Rule?
58. Describe how single, double, and triple bonds differ from each other? Provide examples of each other.
59. Write dot structures for the following:
a. Si b. O
60. Write electron configurations for the following:
61. What is a metallic bond?
a. Mg2+
c. Ne
d. Cl e. Na
b. Cl-
c. N3-
62. How are alloys made
63. List three properties of ionic compounds
64. If the wavelength of a certain light is 6.5X10-7m what is the frequency?
65. The frequency of a wave is found to be 9.0X1014Hz. What is the wavelength?
66. The energy of one photon of light is 4.9X10-19J. What is the frequency of this light?
67. The frequency of a wave is 4.0X1014Hz. Calculate the energy.
68. Electrons travel as waves within the atom. Calculate the wavelength of a wave if the energy is 6.9X1019
J.
Unit 6 Molecular Geometry & Naming/Writing Formulas (Chapters 8.3 & 9):
69. What is the VSEPR theory?
70. Use the VSEPR theory to predict the shape of the following molecules:
a. CH4
b. O2
c. NH3
d. H2O
71. Predict the hybridization of the molecules in question OMIT
72. What is an ion?
73. What is the oxidation number for each of the following elements in ionic compounds? (oxidation number =
charge atoms take in ionic compounds)?
a. Na
b. Mg
c. N
74. How is the charge on a transition metal indicated?
d. Al
e. Ba
f. He
75. Write formulas for the following polyatomic ions:
a. sulfate
b. carbonate
c. hydroxide
d. ammonium
e. phosphate
f. phosphite
76. Write formulas for the following:
a. sodium iodide
b. barium hydroxide
d. ammonium phosphate
77. Name the following compounds:
a. Na2SO4
d. AgCO3
78. Name the following compounds:
a. N2O5
e. SF6
79. Write formulas for the following:
a. dihydrogen monoxide
c. iron (III) oxide
e. cobalt (III) chloride
b. CuOH
c. BaCl2
e. FeO
b. CO
c. CO2
d. SO2
b. tetrasulfur hexachloride
c. nitrogen dioxide
80. List two differences between ionic and covalent compounds.
Unit 7 Chemical Reactions & Equations (Chapter 8.4 & 11):
81. What do the following symbols mean:
a. 
b. (aq)
c. (g)
d. (s)
e. 
82. Why must equations be balanced?
83. Balanced the following equations:
a. CH4 + O2  CO2 + H2O
c. N2 + H2  NH3
b. Na + I2  NaI
d. CaSO4 + AlBr3  CaBr2 + Al2(SO4)3
84. Write balanced equations for the following:
a. Iron plus lead (II) sulfate reacts to form iron (II) sulfate plus lead
b. Ammonium carbonate and magnesium sulfate react to yield ammonium sulfate and magnesium carbonate
85. Classify the following reactions as synthesis, decomposition, single replacement, double replacement, or
combustion reactions:
a.
2 KClO3  2 KCl + 3 O2
b. HCl + NaOH  NaCl + H2O
c. Mg + 2 HCl  MgCl2 + H2
d. CH4 + 2 O2  CO2 + 2 H2O
e. SO3 + H2O  H2SO4
86. What are the products of a combustion reaction?
87. Determine if the following compounds are soluble or insoluble: a) NaCl
b) CO2
c) Mg(OH)2
88. Write net ionic equations for the following:
a) NaCl (aq) + Pb(NO3)2 (aq) 
b) MgI2 (aq) + KOH (aq) 
Unit (Chapter 10) – The Mole:
87. What is a mole?
88. How many particles are in the following:
a. 0.4 mole Ca
b. 12 mol Cu
89. What is the molar mass (gfm) of the following:
a. H2O
b. CuSO4
c. CoCl36H2O
90. How many moles are in the following (see previous problem for help):
a. 12.3 g H2O
b. 0.0885 mg CuSO4
91. Convert the following to volume at STP:
a. 6 moles H2
b. 3.4 g H2O
c. 1.23 kg CoCl36H2O
c. 2.3 x 1024 molecules CO2
92. How many particles are in 23.9 g Ar?
93. Determine the percent composition of each element in the following compounds:
a. H2O
b. Cu3PO4
c. NaOH
94. What is the relationship between an empirical and molecular formula?
95. What is the empirical formula of a compound that is 25.9% nitrogen and 74.1% oxygen.
96. You find that 7.39 g of a compound has decomposed to give 6.93 g of oxygen. The rest of the compound is
hydrogen. If the molecule mass of the compound is 34 g/mol, what is its molecular formula?
97. What is the empirical formula of C8H16?
Unit 9 Stoichiometry (Chapter 12):
97. What information does a balanced equation provide?
98. How many moles of Al are needed to form 3.7 mol Al2O3 in the following reaction: Al + O2  Al2O3
99. Find the number of grams of NH3 produced if 5.40 g of H2 reacts with excess N2 in the following reaction:
N2 + H2  NH3
100.
How many grams of N2 are needed to produce 30.6 g NH3 from the reaction in #95?
101.
How many molecules of O2 are produced when 29.2 g of H2O decomposes?
Unit 10 Thermochemistry (Chapter 17):
88. What is a triple point?
89. What is a phase diagram?
90. An endothermic process _____________ heat from the surroundings.
91. An exothermic process ______________ heat to its surroundings.
92. What is enthalpy? What is the symbol?
93. List the formulas to calculate heat. (3 formulas)
94. What is the heat of fusion and vaporization?
95. Explain in your own words the law of conservation of energy.
96. Calculate the quantity of heat gained gained or lost in the following changes.
a. 3.50 mol of water freezes at 0C.
b. 0.44 mol of steam condenses at 100C.
Units 11 Behavior of Gases (Chapter 14):
97. What are the four assumptions of the Kinetic Molecular Theory?
98. What is gas pressure?
99. List four possible units for gas pressure.
100.
Convert the following units for pressure:
101.
Describe two devices used to measure pressure.
102.
What are the values for STP?
a. 5 kPa to atm
b. 2334 torr to kPa
103.
Identify the name of the following phase changes (If you forget on EOC, consider water and its change
a. solid to liquid
b. gas to liquid
c. liquid to solid
d. solid to gas
104.
How are the pressure and volume of a gas related?
105.
A gas is originally at a volume of 6 mL and a pressure of 1 atm. If the pressure is increased to 2 atm,
what is the new volume of the gas?
106.
State Charles’s Law (*Remember that temperature in Charles’s Law must be in Kelvin)
107.
Oxygen gas is at a temperature of 40C when it occupies a volume of 2.3 liters. To what temperature
should it be raised to occupy a volume of 6.5 liters?
108.
A gas initially has a pressure of 1.5 atm and is at 20C. It has a volume of 3.0 L. If the pressure is
increased to 2.5 atm and temperature is increased to 30C, what new volume will the gas occupy?
109.
What is Dalton’s Law of Partial Pressure?
110.
What is the value for “R” in the Ideal Gas Law?
111.
How many moles of oxygen will occupy a volume of 2.5 liters at 1.2 atm and 25C?
112.
What pressure will be exerted by 25 g CO2 at a temperature of 25C and a volume of 500 mL?
113.
How many oxygen molecules are in 5.56 liters of O2 at STP?
Unit 12 Solutions & Equilibrium (Chapter 16 & 18):
114. What is an aqueous solution?
115.
Distinguish between solvent and solute.
116.
Name three factors that increase the rate of solvation.
117.
What is meant by solubility?
118.
What is the rule for determining if substances will soluble in each other?
119.
Explain how saturated, unsaturated, and supersaturated solutions are different from each other.
120.
Generally, an increase in temperature causes the solubility of most substances to __________ with the
exception of gases. A gas’s solubility will _______________ with an increase with temperature.
130. Use figure from homework packect to answer the following questions:
a. What is the solubility of NH4Cl at 50C?
b. How much KI will dissolve in 100 g of water at 80C?
c. How much NaCl can be dissolved at 10C?
d. Which salt shows the least change in solubility over increasing temperature?
e. Which substance is most soluble at 50C in Figure 16.4?
131.
What is the formula for molarity? (Formula provided on EOC).
132.
Calculate the molarity of a solution that contains 85 g BaCl2 in 500 mL of water.
133.
What is entropy?
134.
What does it mean when entropy has a negative value?
135.
a.
b.
c.
Indicate if the following will have a positive or negative value for S:
the melting of ice
increase in pressure
the reaction of H2 (g) and O2 (g) to form liquid H2O
136.
Explain how the collision theory relates to the creation of chemical bonds.
137.
What roles does a catalyst play in a chemical reaction?
138.
What is activation energy?
139.
Draw an energy diagram for an endothermic reaction. Label the following points: activation energy,
activated complex, different in energy between products and reactants, and a catalyst (with a dotted
line).
140.
How does a catalyst increase the rate of reaction?
141.
Describe how the following affect reaction rate:
a. temperature
b. concentration
c. particle size
Unit 13: ( Chapter 19-Acids and Bases):
142. Name the following acids or bases:
a. HNO3
b. KOH
c. H2SO4
d. HCl
143.
All acids contain __________ while all bases produce __________ in solutions.
144.
Describe the formation of the hydronium ion (H3O+).
145.
List three properties of acids and bases
146. Classify each of the following as being acidic, basic, or neutral:
a. [H+] = 3 x 10-3 M
b. [OH-] = 8 x 10-9 M
e. Mg(OH)2
147.
Complete the following chart:
pH
[H+]
[OH-]
pOH
Acidic, Basic,
Or Neutral
5.9
5 x 10-5
8.635
1 x 10-11
148.
What is an indicator?
149.
Where are the stronger acids located on the pH scale?
150.
Where are the weaker bases located on the pH scale?
151.
What are the products of a neutralization reaction?
152. Write balanced equations for the following:
a. HCl + KOH 
b. H3PO4 + Ca(OH)2 
153.
What is the purpose of a titration?
Unit 14 (Chapter 25-Nuclear Chemistry):
154.
Provide the symbols for the following radioactive particles:
a. alpha
b. beta
c. gamma
155. A atom has a half life of 9.0 days. If 100 grams of the atom are initially present, how many grams will remain
after 36 days?
156.
Complete the following nuclear reactions:
238
234
a.
92 U 
90 Th + ____________
b.
157.
218
84
Po  218 85 At + ____________
Distinguish between nuclear fission and fusion.
Unit 15: (Chapter 20 -Oxidation-Reduction Reactions):
158.
What is meant by oxidation?
159.
Define reduction.
Remember LEO says GER, Loss of Electrons Oxidation (LEO) says Gain of Electrons Reduction (GER).
160. Determine the oxidation number for each element in the following:
a. NaOH
b. CO2
c. PCl5
161. Determine which of the following substances is oxidized and reduced
a. Na(s) + Ag+ (aq)  Na+ (aq) + Ag (s)
b. MnO4- (aq) + H2O2 (l)  O2 (g) + Mn2+ (aq)
*Be familiar with the following pieces of lab equipment: balance, graduated cylinder, pipet, beaker, buret,
thermometer, crucible and LAB SAFETY RULES!