Download Chemistry EOC Review Name

Chemistry EOC Review
Name __________________________________
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.
Unit 1 (Chapters 1 & 2, 6):
1. What is organic chemistry?
2. List the characteristics of the four states of matter.
3. How are physical properties different from chemical properties? List two examples of each type of
property.
4. Classify the following as homogeneous mixture, heterogeneous mixture, element, or compound:
a. air
b. salt water
c. aluminum foil
d. table salt
5. Distinguish between periods and families (Group).
6. Which family (group) on the periodic table is:
a. the most active metals
b. the most active nonmetals
7. Distinguish between metals and nonmetals.
8. How are the noble gases different from other families?
Unit 2 (Chapter 3):
9. Classify the following as having good or poor accuracy and good or poor precision:
a. 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 10 8 m/sec.
b. 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.
10. Write the following in scientific notation:
a. 0.000 85
b. 1, 000, 000
c. 0.4565
d. 112
11. List the SI units for the following quantities:
a. Mass
b. Temperature
c. Volume
d. Length
e. Energy
f. Amount of Substance
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?
Unit 3 (Chapter 4) (Lots of good questions may come from this section):
15. Explain what the early Greeks believed about the atom (Democritus and Aristotle)
16. What did Dalton believe about the atom?
17. Describe Millikan’s experiment and what he determined.
18. What did Rutherford determine about the atom?
19. Describe Rutherford’s experiment.
20. Complete the following chart:
Particle
Proton
Neutron
Electron
Location
Charge
Mass
Discover
21. Distinguish between the mass number and the atomic number:
Atomic
Number
Mass Number Number of
Protons
9
Number of
Neutrons
Number of
Electrons
10
16
108
8
47
35
80
22. All isotopes have the same number of _____________ while all isotopes differ in their number of
________________.
Unit 5 (Chapter 5):
23. Compare and contrast the following four atomic models:
a. Thomson
b. Rutherford
c. Bohr
d. Quantum Mechanical
24. How are frequency and wavelength related?
25. Calculate the wavelength of a yellow light by a sodium lamp if the frequency of the radiation is 3.34 x 10 14 Hz.
26. What is the energy associated with the photon in problem 37?
27. High energy electrons are found _________________ while low energy electrons are found ____________.
28. The Heisenberg Uncertainity Principle states __________________________________________________.
29. List the four atomic orbitals and the maximum number of electrons each orbital can hold.
30. Write electron configurations for the following:
a. He
b. Na
c. P
d. Fe
e. Br
31. Which blocks (s, p, d, or f) will the following elements fall into:
a. U
b. Fe
c. K
d. I
f. Une
Unit 6 (Chapter 6):
32. The modern periodic table is arranged according to ______________________________.
33. Who is considered to be the father of the periodic table?
34. Describe the location of the following groups on the periodic table:
a. metals
b. transition metals
c. halogens
d. metalloids
e. nonmetals
35. Metals form _________ (type of ions) while nonmetals form ______________ (type of ions).
36. How does the atomic radius change moving across the table?
37. List he following in terms of increasing atomic radius: Cs, F, O, Ir
38. What is ionization energy?
39. What is electronegativity?
40. Arrange in terms of increasing ionization energy: H, K, Si, O
41. How is the name given to the following groups: a. Group I
b. Group II
c. Group 17
Units 8, 9, &11 (Chapters 7, 8,9):
42. How is an ionic bond different from a covalent bond?
43. How is a polar covalent different from a nonpolar bond?
44. Classify each bond as ionic, polar, or nonpolar:
a. H – Cl
b. C – O
c. Mg – O
d. F - F
45. List the seven diatomic molecules.
46. What are valence electrons?
47. What is the Octet Rule?
48. Describe how single, double, and triple bonds differ from each other? Provide examples of each other.
49. Write dot structures for the following: a. Si
b. O
c. Ne
d. Cl e. Na
50. Write electron configurations for the following: a. Mg2+
b. Clc. N351. What is a metallic bond?
52. How are alloys made?
53. List three properties of ionic compounds.
54. What is the VSEPR theory?
55. Use the VSEPR theory to predict the shape of the following molecules:
a. CH4
b. O2
c. NH3
d. H2O
56. What is an ion?
57. 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
d. Al
e. Ba
f. He
58. How is the charge on a transition metal indicated?
59. Write formulas for the following polyatomic ions:
a. sulfate
b. carbonate
c. hydroxide
d. ammonium e. phosphate
f. phosphite
60. Write formulas for the following:
a. sodium iodide
b. barium hydroxide
c. iron (III) oxide
d. ammonium phosphate
e. cobalt (III) chloride
61. Name the following compounds:
a. Na2SO4 b. CuOH
c. BaCl2
d. AgCO3
e. FeO
62. Name the following compounds:
a. N2O5
b. CO
c. CO
d. SO2
e. SF6
63. Write formulas for the following:
a. dihydrogen monoxide
b. tetrasulfur hexachloride
c. nitrogen dioxide
64. List two differences between ionic and covalent compounds.
Unit 10 (Chapter 11):
65. What do the following symbols mean:
a. 
b. (aq)
c. (g)
d. (s)
e. 
66. Why must equations be balanced?
67. Balanced the following equations:
a. CH4 + O2  CO2 + H2O
b. Na + I2  NaI
c. N2 + H2  NH3
d. CaSO4 + AlBr3  CaBr2 + Al2(SO4)3
68. 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
69. 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
70. What are the products of a combustion reaction?
71. Determine if the following compounds are soluble or insoluble: a) NaCl b) CO2 c) Mg(OH)2
72. Write net ionic equations for the following:
a) NaCl (aq) + Pb(NO3)2 (aq) 
b) MgI2 (aq) + KOH (aq) 
Unit 11 (Chapter 10) – The Mole
81. What is a mole?
82. How many particles are in the following:
a. 0.4 mole Ca b. 12 mol Cu
83. What is the molar mass (gfm) of the following: a. H2O
b. CuSO4
c. CoCl36H2O
84. How many moles are in the following (see previous problem for help):
a. 12.3 g H2O
b. 0.0885 mg CuSO4
c. 1.23 kg CoCl36H2O
85. Convert the following to volume at STP:
a. 6 moles H2
b. 3.4 g H2O
c. 2.3 x 1024 molecules CO2
86. How many particles are in 23.9 g Ar?
87. Determine the percent composition of each element in the following compounds:
a. H2O
b. Cu3PO4
c. NaOH
88. What is the relationship between an empirical and molecular formula?
89. What is the empirical formula of a compound that is 25.9% nitrogen and 74.1% oxygen.
90. 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?
91. What is the empirical formula of C8H16?
92. Calculate the percent of water in CoCl32H2O.
Unit 12 (Chapter 12 Stoichiometry):
93. What information does a balanced equation provide?
94. How many moles of Al are needed to form 3.7 mol Al 2O3 in the following reaction: Al + O2  Al2O3
95. 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
96. How many grams of N2 are needed to produce 30.6 g NH3 from the reaction in #95?
97. How many molecules of O2 are produced when 29.2 g of H2O decomposes?
Units 14 (Physical Characteristics of Gases):
98. What are the four assumptions of the Kinetic Molecular Theory?
99. What is gas pressure?
100. List four possible units for gas pressure.
101. Convert the following units for pressure:
a. 5 kPa to atm
b. 2334 torr to kPa
102. Describe two devices used to measure pressure.
103. What are the values for STP?
104. Identify the name of the following phase changes (If you forget on EOC, consider water and its changes)
105. a. solid to liquid
b. gas to liquid
c. liquid to solid
d. solid to gas
106. How are the pressure and volume of a gas related?
107. 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?
108. State Charles’s Law (*Remember
that temperature in Charles’s Law must
be in Kelvin)
109. 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?
110. 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?
111. What is Dalton’s Law of Partial Pressure?
112. What is the value for “R” in the Ideal Gas Law?
113. How many moles of oxygen will occupy a volume of 2.5 liters at 1.2 atm and 25C?
114. What pressure will be exerted by 25 g CO2 at a temperature of 25C and a volume of 500 mL?
115. How many oxygen molecules are in 5.56 liters of O2 at STP?
116. What is a triple point?
Unit 15 (Chapter 16 -Solutions):
117. What is an aqueous solution?
118. Distinguish between solvent and solute.
119. How are electrolytes different from nonelectrolytes?
120. Name three factors that increase the rate of solvation
121. What is meant by solubility?
122. What is the rule for determining if substances will soluble in each other?
123. Explain how saturated, unsaturated, and supersaturated solutions are different from each other.
124. 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 16.4 on page 474 to answer the following questions:
a. What is the solubility of KBr at 70 C?
b. How much KNO3 will dissolve in 100 g of water at 30?
c. How much NaCl can be dissolved in 300 grams of H2O at 50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.
Unit 19: ( Chapter 19-Acids and Bases):
133. Name the following acids or bases:
a. HNO3
b. KOH
c. H2SO4
d. HCl
134. All acids contain __________ while all bases produce __________ in solutions.
135. Describe the formation of the hydronium ion (H3O+).
136. List three properties of acids and bases
137. 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
138. Complete the following chart:
pH
[H+]
[OH-]
pOH
Acidic, Basic,
Or Neutral
e. Mg(OH)2
5.9
5 x 10-5
8.6
1 x 10-11
139. What is an indicator?
140. Where are the stronger acids located on the pH scale?
141. Where are the weaker bases located on the pH scale?
142. What are the products of a neutralization reaction?
143. Write balanced equations for the following:
a. HCl + KOH 
b. H3PO4 + Ca(OH)2 
144. What is the purpose of a titration?
Unit (Chapter 17):
145. A 500 g sample of water is heated from 45 C to 60 C. How much energy is required? (See Specific Heat
Formula and Table 17-1)
146. Distinguish between exothermic and endothermic reactions?
147. Sketch graphs of endothermic and exothermic reactions.
148. What is enthalpy?
149. What is entropy?
150. What does it mean when entropy has a negative value?
151. Indicate if the following will have a positive or negative value for S:
a. the melting of ice
b. increase in pressure
c. the reaction of H2 (g) and O2 (g) to form liquid H2O
152. Explain how the collision theory relates to the creation of chemical bonds.
153. What roles does a catalyst play in a chemical reaction?
154. What is activation energy?
155. Draw an energy diagram for an endothermic reaction. Label the following points: relative energy of reactants,
relative energy of products, activation energy, activated complex, different in energy between products and
reactants, a catalyst (with a dotted line), and an inhibitor (with a dotted line).
156. How does a catalyst increase the rate of reaction?
157. Describe how the following affect reaction rate:
a. temperature
b. concentration
c. particle size
Unit 20: (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)
Unit 25 (Chapter 25-Nuclear Chemistry):
162. Provide the symbols for the following radioactive particles:
a. alpha
b. beta
c. gamma
163. 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?
164. Complete the following nuclear reactions:
a. 238 92 U  234 90 Th + ____________
b. 218 84 Po  218 85 At + ____________
165. Distinguish between nuclear fission and fusion.
*Be familiar with the following pieces of lab equipment: balance, graduated cylinder, pipet, beaker, buret,
thermometer, crucible and LAB SAFETY RULES!