Name_________________
You must check out a Chemistry book from the textbook room for this summer assignment.
You may check out books June2-4 from the textbook room.
There will be 2 summer study sessions for Chemistry on July 15 and 29th. Call or e-mail Penny
Howard for details. Penny.howard@marion.k12.fl.us or 352-239-5427
AICE Chemistry Summer Review
1. Distinguish between quantitative and qualitative measurements. Give an example of
each.
2. Rewrite the following in scientific notation: .000007345, 4,304,000, 0.0230
3. Give the number of significant figures in the following: 1603, 3.28 x 10^-3, 1500, .00240
4. Round off the measurement 417.01g to three significant figures. _________
5. Round off the measurement 0.0030955 to three significant figures. _________
6. What is the sum of 6.210 L and 4 L expressed in the correct number of significant
figures?
7. A rock has a mass of 127 g and displaces 32.4 ml of water. What is the density of the
rock?
8. A piece of copper has a volume of 28.6 cm^3. The density of copper is 8.92 g/cm^3.
What is the mass of copper?
9. The accepted value for the density of sodium chloride is 2.165 g/ml. A student
measured the density as 2.075 g/ml. Calculate the percent error of the measurement.
10. What is the volume of an object with a density of 5.20 g/ml and a mass of 600 g?
11. What is the temperature 195 K expressed in degrees Celsius?
12. How many joules are there in 165.2 calories?
13. Make the following conversions:
a) 382.2mg to g
b) 5.0 x 10^3 mm to meters
c)
.12 L to ml
d) 13.5 km to meters
e) 18.7 x 10^3 mg to decigrams
14. Define isotope.
15. What is the charge and the location of the following subatomic particles in a neutral
atom?
a) Proton-
b) Neutron-
c) electron-
16. What is the charge of the nucleus of a neutral atom? What is the charge of a neutral
atom?
17. How can you find the number of neutrons an atom has?
18. The atomic number is always equal to the number of ___ in a atom.
19. What are the two different ways to express an isotope? _________________
20. How do the isotopes protium, deuterium, and tritrium differ?
21. Average relative atomic masses are measured in _____
22. There are four naturally occurring isotopes of the element chromium. The relative
abundance of each is:
Chromium-50 4.31%
chromium-52 83.76%
Chromium-53 9.55%
Chromium -54 2.38%
Calculate the average atomic mass of these isotopes
23. Explain how the atoms of one element differ from those of another element.
24. If the atomic number of an element is 8, then how many neutrons, protons, and
electrons are there in a neutral atom?
25. Define quantum.
26. Define orbital.
27. Who discovered the electron and how? The proton and How?
28. What are the shapes of the s and p orbitals?
29. What is the maximum number of electrons that each of the following orbitals can hold?
a. s
b. p
c. d
d. f
30. How many orbitals are in the s orbital, p orbital, d orbital, and f orbital.
31. If three electrons are available to fill three empty 2p atomic orbitals, how will the
electrons be distributed in three orbitals?
32. It the spin of one electron in an orbital is clockwise, what is the spin of the other
electron in that orbital? Why?
33. What is the next atomic orbital in the series 1s, 2s, 2p, 3s, 3p, 4s?
34. Write the electron configurations for the following elements and their ions.
a. K K+
b. P Pc. O O-2
d. Ca Ca+2
35. Define stable electron configuration.
36. Which of the following are represents elements:
Mg, Li, Fe, Se, K, Pb, Ni, Br, Na, S, O
37. Which of the following are transition elements:
Sr, Ru, As, W, Ag, Al, Ni, Mg, Zn, Sn, Sb,
38. Which of the following are metals:
Ta, Nd, Se, F, C, Br, Ba, Na, Cl, Ag, Sr
39. Which elements constitute the halogens?
40. Which elements constitute the noble gases?
41. Which elements constitute the alkali metals?
42. Which elements constitutes the alkaline earth metals?
43. List 4 different Things you can find out about lithium by using the periodic table.
44. What are the typical charges for the following monatomic ions?
F__, O__, Al__, Mg__, Cl__, P__, Be__, Na__, Ba__
45. Name all of the diatomic elements
46. What are the formulas for the following polyatomic ions.
a. Acetate
d. Nitrate
b. Phosphate
e. Sulfate
c. Carbonate
f. Sulfite
47. Distinguish between ionic compounds and molecular compounds.
48. Define empirical formula.
49. Write the chemical formulas for each of the following:
a. Magnesium chloride
e. tin IV oxide
i. diphosphorus trioxide
b. Potassium nitrate
f. calcium hydroxide
j. carbon tetrachloride
c. Carbonic acid
g. hydrochloric acid
k. sulfurous acid
d. Ammonium chloride
h. dinitrogen pentaoxide
l. iron III acetate
50. State the number of electrons lost or gained in forming each of these ions.
a. Mg+2
b. Br-1
c. Ag+1
d. O-2
e. Al+3
51. Determine the empirical formula of the compound with the percent composition of
29.1 % Na, 40.5 % S and 30.4% O.
52. How many kilograms of Iron can be recovered from 643 kilograms of the ore Fe2O3?
53. What is the percent composition of each of the following? a.) Cr2O3
b.) C6H12O6
54. Avogadro’s number of particles is equal to one ____________.
55. The representative particles of an element is an _________________.
56. The representative particles of a molecular compound is an _________________.
57. The representative particle of an ionic compound is an ____________________.
58. Convert 50.4 grams of CaBr2 to moles.
59. How many grams are there in 1.26 moles NaI?
60. How many males are there in 86.2 C2H4?
61. How many molecules are there in 0.943 moles of H2O?
62. How many grams are there in 2.54 liters of O2 gas?
Complete and write the balanced equations for each of the following:
63. Calcium chloride + nitric acid 
64. Aluminum + iron III oxide 
65. Mg (OH)2 + HNO3 
66. PbO2 
67. Pb(NO3)2 + H2SO4 
68. Water and diphosphorous pentaoxide yields phosphoric acid
69. Hydrogen and chlorine yields hydrochloric acid.
70. Calcium oxide and water yields calcium hydroxide.
71. What are the products of a combustion reaction?
72. What are the five different types of chemical reactions? Give an example of each.
73. Write a balance equation for the following:
a. Water decomposes into hydrogen and oxygen gas when electricity is added.
74. How many moles of oxygen are produced if you have 2.5 moles of water?
75. How many liters of hydrogen gas are produced if 7.2 moles of water is used?
76. How many grams of oxygen will be produced if 10.0 g of water is used?
77. Distinguish between limiting and excess reagents.
78. The percentage yield of the reaction 2Al + 2H3PO4 2AlPO4 + 3H2 is 73.7%. How many
grams of Al must be used to react with an excess of H3PO4 to give 30.0g of AlPO 4?
79. If a 200.0g sample of Al is reacted with 175.0L of O2 is STP according to the reaction 4Al
+ 3O2  2Al2O3, what is the limiting reagent?
80. Balance the following equation: CuCl (aq) + H2S (g)  Cu2S (s) + HCl (aq)
a. How many moles of Cu2S could be produced from 9.5 moles of CuCl reacting
with an excess of H2S gas?
b. How many grams of Cu2S could be produced from 9.90g of CuCl reacting with an
excess of H2S gas?
c. How many liters of gas would be used to produce 2.50 moles of HCl?
d. How many grams of CuCl would you need to produce 1.75 moles of Cu 2S?
81. From the following empirical formulas and the formula weight for each compound,
determine the correct molecular formulas.
a. Vitamin C, C3H4O3, molar mass = 176.0 g/mol
b. Nicotine, C5H7N, molar mass = 81.0 g/mol
82. A Freon gas sample with a molecular weight of 121.0 g/mol was found to be 9.92%
carbon, 58.68% chlorine, and 31.40% Fluorine. What is the molecular formula of the
Freon gas?
83. Write the formulas for the following compounds: calcium hydroxide, barium sulfate,
calcium selenide, iron(II) oxide, zinc oxide, sodium peroxide, copper II sulfate, sulfurous
acid, nitric acid
84. Write the outer electron configuration for the following: K, Na, Cl, Br, N, O, I-, Ca+2, H
85. A diatomic molecule with a triple covalent bond is ___________.
86. A diatomic molecule with a double covalent bond is ___________.
87. How many unshared pairs of electrons does the nitrogen atom in ammonia possess?
88. How many unshared pairs of electrons are there in a molecule of hydrogen iodide?
89. Calculate the number of moles of oxygen in a 12.5 liter tank at a pressure of 25,325 kPa,
measured at 220C
90. The volume of a balloon is 340 ml at a pressure of 142 kPa and a temperature of 250C.
Calculate the new volume if the temperature is raised to 850C and the pressure changes
to 182 kPa.
91. A volume of 3.0L of air is warmed from 500C to 1000C. What is the new volume if the
pressure remains constant?
92. What is the temperature of the gas inside a 750 ml balloon filled with 0.030g H2 gas?
The pressure of the balloon is 120 kPa
93. Nitrogen gas in a steel cylinder is under a pressure of 15,200 kPa at 27 0C. What will the
pressure in the tank be if the tank is left in the sun and the internal temperature rises to
550C?
94. A sample of gas occupies a volume of 80 ml at a pressure of 0.50 atm and a temperature
of 00C. what will be its volume at a pressure of 1.50 atm and a temperature of 500C?
95. Explain why a magnesium atom is smaller than atoms of both sodium and calcium.
96. Would you expect a Cl- ion to be larger or smaller than an Mg+2 ion? Explain.
97. Which family of elements is characterized by an s2p3 configuration?
98. Name the element that matches the following description:
a. One that has 5 outer electrons in the third row of the periodic table.
b. One with a 4s2 4p5 electron configuration
c. The noble gas on the 5th row
d. The 4th period group 6A element
99. Compare the ionization energy of sodium to that of potassium.
100.
Will the electronegativity of barium be larger or smaller than that of strontium?