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Chemistry I - REVIEW
**DO NOT DEPEND ENTIRELY ON THIS PRACTICE TEST FOR YOUR PREPARATION**
Chapters 1 and 2 – Science of Chemistry / Matter and Energy
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Define the following:
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Physical Change -
Chemical Change -
Physical Property-
Chemical Property-
1.
Chemistry is defined as the study of the composition and structure of materials and:
a. the categories of matter.
c. the electrical currents in matter.
b. the changes in matter.
d. molecules in living things.
2.
All of the following are chemical changes except:
a. zinc metal dissolves in acid
c. the rusting of an iron gate
a. magnesium metal burns in a flame
d. water freezing to become ice
3.
Which of the following represents a physical property?
a. corrosiveness
b. reactivity with oxygen
c. state of matter
d. toxicity
4. A mixture that contains substances that are not evenly distributed:
a. heterogeneous
b. extensive
c. homogeneous
5. Which of the following involves a chemical change?
a. sharpening a pencil
b. baking of bread
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c. melting of wax
d. phase
d. tasting a fruit
6. The difference between intensive and extensive properties is that:
a. intensive properties deal with quantity or amounts of matter while extensive do not
b. extensive properties deal with quantity or amounts of matter while intensive do not
c. intensive properties take in heat, extensive give off heat
d. intensive properties create new substances while extensive do not
7. Of the four common states of matter, ______ is the highest energy:
a. solid
b. liquid
c. gas
d. plasma
8. A mixture that contains substances that are uniformly distributed:
a. heterogeneous
b. extensive
c. homogeneous
9.
All of the following are examples of units except
a. weight
b. kilometer.
c. gram.
10. The reason for organizing, analyzing, and classifying data is
a. so that computers can be used.
b. to prove a law.
c. to find relationships among the data.
d. to separate qualitative and quantitative data.
d. phase
d. teaspoon.
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11. Poor precision in scientific measurement may arise from
a. the standard being too strict.
b. human error
c. limitations of the measuring instrument.
d. both human error and the limitations of the measuring instrument.
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12. A logical explanation of a body of observed natural phenomena is a scientific:
a. principle.
b. experiment
c law.
d. theory.
13. Electricity can convert oxygen into ozone. The ozone created by this process is
a. a new substance.
c. a different state of oxygen.
b. an isotope of oxygen.
d. an allotrope of oxygen.
List the observable indications that a chemical change has occurred.
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Describe the motion of particles in each of the three common state of matter.
Solid –
Liquid –
Gas –
Plasma –
Reporting Data: re-familiarize yourself with: Density
14. A substance has a volume of 23.0 ml and a mass of 45.0 g. what is its density?
a. 1.95 g
b. 1.95 ml
c. .51 g/ml
d. 1.95 g/ml
15. A substance has a density of 5.50 g/ml. What is the volume of the substance if its mass is 34.0 g?
m
a. 187 ml
b. 6.18 ml
c. .16 ml
d. 187 g
D =
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16. A substance has a density of 78.00 g/ml. What is the mass of the substance if its volume is 123.0 ml?
a. .63 g
b. 1.57 g
c. 9594 g
d. 78 g
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17. A substance has a volume of 28 cm3 and a mass of 26 g. What is its density?
a. .92 g/cm3
b. 1.07 g/cm3
c. 728 g/cm3
d. 123 g/cm3
Scientific Notation: Review your guidelines
18. Convert 1234000.00 to scientific notation.
a. 1.23400000 x 106
b. 1.23 x 10-6
c. 1.23 x 106
d. 1 x 109
19. 3.5 x 103 would be expressed as:
a. 3500
b. 35000
c. .00035
d. .000035
20. Convert .000097 to scientific notation.
a. 9.7 x 106
b. 9.7 x 10-6
c. 9.7 x 10-5
d. .97 x 105
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21. 3.5 x 10-4would be expressed as:
a. .0035
b. 35000
c. .00035
d. 3500
Signficant Figures: Review your rules
Determine the number of significant figures in each of the following measurements;
503 m ______3____
0.0015 mg _____2_____
635.526 _____6_____
630 ml _____2_____
601L _____3_____
10001 m _____5_____
2407 g ____4______
500 ml ___1_______
1.0000006 µg _____8_____
22. Multiply (2.3 x 103) and (3.43 x 104)
a. 7.889 x 107
b. 8 x 101
23. Divide (4.5 x 108 )
a. 2.64 x 105
c. 7.89 x 107
by (1.7 x 103)
b. 3 x 105
c. 2.6 x 105
d. 2.647 x 105
24. Add (3.75 x 104) and (6.1 x 105)
a. 64.75 x 104
b. 6 x 105
c. 6.475 x 105
25. Subtract (2.20 x 107) from (4.4 x 107)
a. –2.2 x 107
b. 6.6 x 107
c. 2.2 x 107
26. 23.334 + 13.00 + 0.00234
a. 36.34
b. 36.3
27. 23.098 x
a. 70.0
3.033
c. 36.336
x 1.0
b. 70.1
28. 31.1 – 5.456 – .00981 – 4.0
a. 21
b. 22
29. 23.045 ÷ 1.34
a. 143.3
÷ .12
b. 140
d. 7.9 x 107
d. 6.5 x 105
d. 2 x 107
d. 36.33634
c. 70.056
d. 70.
c. 21.6
d. 21.634
c. 143.31
d. 143
Why are significant figures important when reporting measurements?
Explain why textbook pictures of atoms and molecules are “models”.
Chapter 3 -Atoms and Moles
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You should, by now be familiar with locations of common individual elements on the table, as well as
the element groupings.
Remember - elements (atoms) on the TABLE are electrically Neutral - same # of Protons and Electrons
You should be familiar with basic atomic structure. Identify everything you can about the following:
ProtonsNeutrons-
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Atomic Mass Atomic Number-
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Electrons-
Isotopes-
Law of Multiple Proportions-
Law of Conservation of Mass-
Law of Definite Proportions-
Atomic Theory-
Periods-
Groups-
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Behavior of electrons
Understand what happens when electrons become "excited" and jump to higher energy levels, then
return to their ground states.
Complete the table for the following ISOTOPES.
Symbol
Atomic Number
Atomic Mass
Number
Number of protons
Number of
neutrons
Number of
electrons
Zn
30
65
30
35
30
Ca
20
41
20
21
20
Se
34
79
34
45
34
Ar
18
40
18
22
18
Sample Problems:
30. Which of the following is not an element?
a. potassium
b. californium
c. hydronium
d. radon
31. How many protons are there in a neutral atom of Calcium (Ca)?
a. 40
b. 20
c. 60
d. 10
32. How many electrons are present in a neutral atom of Zinc (Zn)?
a. 65
b. 35
c. 30
d. 95
33. How many neutrons are in a neutral atom of Iridium (Ir)?
a. 115
b. 192
c. 77
d. 269
34. An element that is said to be electrically neutral has the same numbers of:
a. neutrons and protons
b. protons and atoms
c. electrons and neutrons
d. protons and electrons
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35. Sodium (Na) has an atomic number of 11 and an atomic mass of 23. How many neutrons does it have?
a. 11
b. 23
c. 12
d. 34
For some of the following, look at figure - XYZ
1
1
H
18
2
13
3
Mg
3
4
5
6
7
K
8
Fe
5
6
15
C
2
4
14
9
10
11
17
O
F
P
12
Cu
16
Ga
D
He
B
C
Kr
A
Cs
7
36. When in its ground state, how many energy shells will Magnesium's electrons use?
a. 2
b. 3
c. 12
d. 24
37. When in its ground state, how many outer (valence) electrons does Gallium have?
a. 3
b. 13
c. 4
38. Which area is classified as metals?
a. C
b. D
c. B
d. A
39. Which area is classified as the Noble gases?
a. C
b. D
c. B
d. A
40. Phosphorous (P) would have ___ outer electrons when in its ground state.
a. 15
b. 30
c. cannot be determined
-d. 70
d. 5
41. Silicon (Si) has an atomic number of 14 and an atomic mass of 28. How many protons does it have?
a. 28
b. 14
c. 42
d. 7
42. Xenon (Xe) has an atomic mass of 131 and an atomic number of 54. How many electrons does it
have?
a. 131
b. 77
c. 27
d. 54
43. What period is Cesium (Cs) located in?
a. 1
b. 55
c. 133
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d. 6
44. When in its ground state, what energy level will Oxygen's (O) valence electrons be found in?
a. 16
b. 2
c. 8
d. none of these
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45. When in its ground state, how many outer electrons would Krypton (Kr) have?
a. 8
b. 18
c. 36
d. 47
46. Written in this form, 35Cl, Chlorine has an atomic number of:
17
a. 35
b. 17
c. 18
d. 52
47. For the following, 80$. The $ represents what element?
35
a. Mercury
b. Rhodium
c. Bromine
d. Moneydemum
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48. When electrons return from the excited state to the (unexcited) ground state they release:
a. light energy (photons)
b. heat energy
c. chemical energy
d. hydrogen gas
Review the basics of ELECTRON CONFIGURATION
Determine the FULL electron configuration and draw the orbital diagram for:
Sc 1s22s22p63s23p64s23d1
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑
1s 2s
2p
3s
3p
4s
3d
1s22s22p63s23p64s23d104p65s24d105p1
In
↑↓ ↑↓ ↑↓↑↓↑↓ ↑↓ ↑↓↑↓↑↓ ↑↓ ↑↓↑↓↑↓↑↓↑↓ ↑↓↑↓↑↓ ↑↓ ↑↓↑↓↑↓↑↓↑↓
1s
2s
2p
3s
3p
4s
3d
4p
5s
4d
↑
5p
Cl 1s22s22p63s23p5
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑
1s 2s
2p
3s
3p
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49.The element with electron configuration 1s2 2s2 2p6 3s2 3p2 is
a. Mg
b. C c. S
50.A single orbital in the 4d level can hold ____ electrons.
a. 10 b. 2 c. 3
51.How many total orbitals can exist at the third main energy level?
a. 3 b. 6 c. 9 d. Si
d. 6
d. 18
52. With the quantum model of the atom, scientists have come to believe that determining an electron's
exact location around the nucleus is:
a. impossible
b. easily determined with a complex equation
c. only possible under certain conditions
d. easily determined with specialized equipment
53. According to the Bohr model of the atom, the single electron of a hydrogen atom circles the nucleus:
a. at any of an infinite number of distances, depending on its energy.
b. in many fixed orbits within an electron cloud.
c. Counterclockwise
d. in a specific, allowed or “quantized” energy level
54. How many electrons are needed to completely fill the fourth energy level?
a. 8 b. 18 c. 32 d. 40
You should be familiar with the following terms
Aufbau Principle
Pauli Exclusion Principle
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Hund’s Rule
Heisenberg Uncertainty Principle
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Give the Shorthand electron configuration for the following:
Te
Rf
[Kr] 4d105s25p4
[Rn] 5f146d25s2
55. Which of the following is NOT part of Dalton’s atomic theory?
a. Atoms cannot be divided, created, or destroyed.
b. The number of protons in an atom is its atomic number.
c. In chemical reactions, atoms are combined, separated, or rearranged.
d. All matter is composed of extremely small particles called atoms.
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ATOMS/ELEMENTS - The Mole / Avogadro's Number - know the following:
What is the Mole? 1 mol = 6.022 x 1023 atoms
Know the related conversion factors
1 mol “X”
6.022 x 1023 atoms “X”
Molar mass of element “X” g
1 mol “X”
or
When speaking of individual elements - what is a Molar Mass? -
56. What is the mass of 8.00 mol of Na?
a. 184 g
b. 2.80 g
57. What is the mass of 5.50 mol of Cl?
a. 93 g
b. 193 g
c. .347 g
c. 6.36 g
d. 880. g
d. .139 g
58. Which of the following has the largest number of moles?
a. 3.0 g K
b. 3.0 g Fe
c. 3.0 g C
d. 3.0 g I2
59. How many atoms are in 2.00 moles of Carbon?
a. 1.20 x 1024
b. 3.01 x 1023
c. 1.00 x 1010
d. 3.32 x 10-23
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60. How many moles of Nickel are in 459 g of Nickel?
a. 27081 mol
b. 0.129 mol
c. 7.78 mol
d. 400 mol
61. How many moles are in 4.02 x 1022 atoms of hydrogen?
a. .0667
b. 6.022 x 1023
c. 14.97
d. 2.42 x 1046
62. How many atoms are present in 13 mol of Al?
a. 6.02 x 1023
b. 4.6 x 1022
c. 7.83 x 1024
d. 3.01 x 1030
63. What is the molar mass of Sulfur?
a. 16.00 g
b. 22.98 g
c. 48.00 g
d. 32.06 g
64. How many moles are in 2.45 x 1012 atoms of carbon?
a. 6.02 x 1023
b. 4.07 x 10-12
c. 7.83 x 1024
d. 2.45 x 1011
65. What is the molar mass of Lithium?
a. 6.94 g
b. 3.00g
d. 6.022 x 1023
c. 10.25 g
Chapter 4 - The Periodic Table -----------------------------------
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Be Able To Define/Describe the following:
Metals-
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NonmetalsMetalloidsNoble gasesPeriodic LawValence ElectronsBe sure to become familiar with:
The Family Groups of elements: What are these, where are these on the table? How reactive are they?
Alkali MetalsAlkaline Earth MetalsTransition MetalsMixed Group ElementsNoble GasesHydrogen-
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Refamiliarize yourself with the contributions of:
Dmitri Mendeleyev –
Henry Mosley –
Sample Problems: look back at Figure XYZ - (The Periodic Table)
66. Which of the following is the most reactive element?
a. Ga
b. Fe
c. Mg
d. K
67. Which of the following is the least reactive element:
a. Kr
b. Cs
c. O
d. Cu
68. Which element would be classified as an Alkaline Earth metal?
a. K
b. C
c. Fe
d. Mg
69. Which of the following would be classified as a transition metal?
a. Ga
b. H
c. Fe
d. Kr
70. Which of the following would not be classified as a Main block element?
a. Mg
b. Cu
c. Kr
d. P
71. In the modern periodic table, elements are ordered according to:
a. decreasing atomic mass
b. Mendeleev’s original design
c. increasing atomic number
d. when they were discovered
Periodic Trends - All aspects of the trends and how they INCREASE / DECREASE
What is Electronegativity? -
What is Electron Affinity? –
What is Ionization Energy –
What is the Atomic radius? –
What is the Shielding Effect? –
72. Based on the information from the periodic table, the shielding effect increases with increasing atomic
number within a(n):
a. period
b. group
c. orbital
d. configuration
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73. Within a group, the nucleus has a stronger ability to pull on the outermost electrons in elements of
____________ atomic numbers.
a. lower
b. even
c. greater
d. odd
74. The difference between an electron affinity value of –47 and one of –329 is that with a value of –47:
a. the more difficult it is for an atom to take on an extra electron
b. the more easily the atom can take on an extra electron.
c. the greater effect there is on the atomic radii
d. the more electrons the atom gives away
75. Generally, ionization energy values:
a. increase as you move from right to left in a period
b. decrease as you move from the bottom to the top of a group
c. increases as you move from left to right in a period
d. increases as you move from ion to ion in a compound
76. The energy needed to remove electrons from atoms is know as:
a. isotope energy
b. electronegativity
c. electron affinity
77. The most electronegative elements when forming compounds are:
a. alkali metals
b. transition metals
c. main block elements
d. ionization energy
d. halogens
78. As the number of electrons in the outer shells of atoms within a period decreases, the atoms (in
general) have:
a. increasing atomic radii, decreasing ionization energies and electronegativity values
b. decreasing atomic radii, ionization energies and electronegativity values
c. decreasing atomic radii, increasing ionization energy, decreasing electronegativity values
d. increasing atomic radii, increasing ionization energy and electronegativity values
79. The periodic law states that:
a. The physical/chemical properties of the elements are periodic functions of their atomic number
b. no two electrons with the same spin can be found in the same place at the same time.
c. Electrons exhibit properties of both particles and waves.
d. The chemical properties of elements can be grouped according to periodicity
80. The energy needed to remove an electron from an atom ___ as you move left to right across a period.
a. generally increases
b. remains constant
c. generally decreases
d. varies unpredictably
81. An element with a high ionization energy is:
a. most likely an isotope of the element
b. more likely to gain a proton
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c. more likely to lose an electron
d. more likely to gain an electron
82. Which of the following statements is FALSE:
a. Ionization energies decrease as you move down a group
b. Atomic radius generally increases as you move from right to left within a period
c. Shielding effect is constant as you move from left to right across a period
d. Electron affinities decrease steadily from left to right across the periodic table
e. All are TRUE
83. The tendency for an atom to attract the electrons to itself when it is combined with another atom:
a. generally increases as you move from left to right from boron to fluorine
b. generally increases as you move from right to left from phosphorous to aluminum
c. eventually levels off as you move from left to right across a period
d. varies unpredictably and cannot be determined under normal conditions
84. According to the information of the periodic table, Magnesium…..
a. has a stable octet in its outermost energy level
b. will accept an extra valence electron from Fluorine to complete its outer shell
c. will form a cation when it loses its valence electrons
d. ions always combine readily with any available cations to complete their outer shells
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85. According to information on the periodic table, an atom of Xenon:
a. has a stable octet in its outermost energy level
b. will accept an extra electron to complete its outermost energy level
c. forms a cation when it completes its outer energy level
d. always combine readily with any available anions to complete their outer shells
86. Which of the following lists of elements is in order of increasing atomic radii.
a. Sr, Rb, Sn, I, In d. In, Sb, I, Sr, Rb
b. Rb, Sr, In, Sb, I e. Sb, I, In, Sr, Rb
c. I, Sb, In, Rb, Sr
f. I, Sb, In, Sr, Rb
87. Which of the following lists of elements is in order of increasing electronegativity.
a. At, Bi, Cl, F, I
d. Bi, At, I, Cl, F
b. Bi, At, Cl, F, I
e. F, Cl, Bi, I, At
c. F, Cl, I, At, Bi
f. At, Bi, I, Cl, F
Chapter 5 Ionic Bonds ---------------------------------------------------
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What are ions? -
-Understand role the Periodic Trends play in the formation of Ionic Bonds.
Know the charges for your groups and WHY those charges occur.
The difference between cations and anions
Understand that metals combine with nonmetals to form ionic bonds.
Understand the reactivity of metals and which groups of nonmetals they would most likely combine
with to form ionic bonds.
88. Negative ions:
a. contain a single proton
c. contain more protons than electrons
b. contain more electrons than protons
d. always have a 1- charge
89. When a neutral atom of magnesium loses electrons to obtain a stable octet, the charge of the resulting
ion is:
a. 2+
b. 2c. 3+
d. 1+
90. An atom or group of atoms that carry a positive charge is called a:
a. anion
b. polyatomic
c. cation
d. neutron
91. When a neutral atom of phosphorous becomes an ion to obtain a stable octet, its charge would be:
a. 2b. 3+
c. 3d. 292. Neutral atoms that are likely to gain three electrons (to obtain stable octets) when they react can be
found in group:
a. 2
b. 18
c. 15
d. 17
93. Group 1 Metals would be most likely to combine with non-metals in group __ to form ionic bonds.
a. 17
b. 18
c. 15
d. 2
94. Ions form when:
a. protons are lost or gained
c. electrons are lost or gained
b. neutrons are gained by atoms
d. ionic bonds are formed
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95.
Noble gases have a low reactivity because:
a. their outermost electron levels can only accept a single electron.
b. their potential energy is much too high.
c. their electrons are only exchanged with other noble gases.
d. their outermost energy levels are completely filled.
96. Which of the following atoms does not naturally form an ion?
a. Carbon
b. Lead
c. Uranium
d. Argon
97. Which of the following will form an ion with a 3+ charge
a. Beryllium
b. Helium
c. Aluminum
d. Potassium
98.
Positive ions:
a. contain a single proton
c. contain more protons than electrons
99.
When an electron is added to an atom:
a. energy input is required
c. energy is released
b. contain more electrons than protons
d. always have a 1+ charge
b. stable octets always form
d. ionic bonds are broken
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Ionic Bonds - Bonds
You will have to remember the charges present for each ion to correctly write its formula.
That includes knowing or at least knowing how to identify charges for the “variable” transition
metals.
100. Identify the compound that would form between Lithium and Sulfur
a. Li2S
b. LI2S2
c. LiS2
d. SLi2
101. The name of the compound formed from lithium and sulfur would be:
a. lithium sulfur
b. lithate sulfate
c. lithium sulfide
d. sulfur lithite
102. Binary compounds
a. are two or more bonded ions
b. are named with the anion first
c. consist of two monatomic ions
d. generally have a charge of 1+ or 2+.
103. An Ionic compound is defined as:
a. an atom that carries a positive or negative charge
b. the three dimensional arrangement of atoms or ions in a crystal
c. a compound composed of cations and anions combined so that (+) and (-) charges equal zero
d. compound composed of atoms that bond by sharing pairs of electrons in their outer shells.
104. Atoms form bonds in order to:
a. attain the stability common to the noble gases
b. to give themselves a more stable future
c. to provide a stable alternative source of energy
d. to become a polyatomic ion
105. What would the formula be for the formation of an ionic compound between Iron (III) and Oxygen
a. FeO2
b. Fe2O
c. Fe2O3
d. IrO2N
106. The charge on lead (IV) would be:
a. 4+
b. 4-
c. 5+
107. Which of the following is improperly named:
a. K2SO4: Potassium sulfate
b. Al(OH)3 : Aluminum Hydroxide
c. HgBr2: Mercury (II) bromide
d. CuCl2 : Copper Chloride
d. there is no charge on transition metals
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Write names for the following Ionic Compounds:
FeO
_____Iron (II) Oxide_________
Al2(SO4)3
_____Aluminum Sulfate________
MnCl5
___Manganese (V) Chloride____
RbI
______Rubidium Iodide________
K2CO3
___Potassium Carbonate_______
AgC2H3O2 ______Silver (I) Acetate________
Pb3(PO4)2 ____Lead (II) Phosphate_______
FeBr2
_____Iron (II) Bromide_________
CoCl2 .3H2O ______Cobalt (II) Chloride TriHydrate_________
MgCO3.5H2O ______Magnesium Carbonate PentaHydrate____
Write proper formulas for the following Ionic compounds
Copper (II) Sulfate
_____ CuSO4______
Ammonium Iodide
_____ NH4I______
Magnesium Phosphate
___ Mg3(PO4) 2_____
Calcium Hydroxide
___ Ca(OH)2_____
Titanium (IV) Nitride
_____ Ti3N4_____
Copper (I) Oxide
_____ Cu2O______
Manganese (II) Chlorate ____Mn(ClO3)2_____
Lithium Bicarbonate
____LiHCO3_______
Magnesium Sulfate Heptahydrate _________ MgSO4 . 7H2O __________
Sodium Tetraborate Decahydrate _________ Na2B4O7 . 10H2O ________
Describe what is wrong with the following formulas/names for binary compounds
a.) Mg2S2
b.) ZnClO32
c.) monostrontium dioxide
d.) aluminum (III) bromide
e.) (NO3)LI
What is a Metallic Bond?
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How can Metallic Bonds be used to describe the behavior of Metals?
Predict the IONS that will form from the following - give the shorthand configuration of the ION.
Atom
ION
Noble Gas
Shorthand Electron Configuration
Rb
Rb1+
Kr
[Ar]4s23d104p6
Te
Te2-
Kr
[Kr]5s24d105p6
Chapter 6 – Covalent Bonds -----------------------------------------------------
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Understand the role electronegativity plays in the formation of covalent bonds
You’ll need to be good at Lewis Symbols and Lewis Structures
Draw Lewis symbols for the following elements: Potassium, Silicon, Chlorine, Xenon
K.
:Si:
.
:Cl:
..
..
:Xe:
..
Draw Lewis Structures for the following covalently bonded compounds: Carbon Tetraiodide; Silicon
Dioxide; Dihydrogen monoxide; Oxygen Dichloride
..
.. : ..I : ..
: ..I :C:
.. ..I :
: ..I :
..
..
O
..
.. ::Si:: O
..
H :O:
.. H
.. .. ..
:Cl
.. Cl:
.. :O:
..
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108. To draw a Lewis structure, it is NOT necessary to know:
a.which atoms are in the molecule
c. number of valence electrons
b.bond length
d. the number of atoms in the molecule
109. A Lewis structure does not show
a. molecular shape
b. valence electrons
c. atoms
d. bonds
110. When more than one “correct” Lewis structure for a molecule can be drawn it is said to be a:
a. oxidation number
b. resonance hybrid
c. empirical formula
d. covalent bond
111. These electrons are found in the outermost energy level
a. resonance electrons
b. oxidation electrons
c. molecular electrons
d. valence electrons
112. A diagram showing the arrangements of outer electrons among atoms in a molecule
a. Bohr model
b. molecule
c. lewis structure
d. empirical formula
113. For the following, which one represents the proper Lewis structure for HCl
a. Cl
H
b. H Cl
c. H Cl
d. H Cl
114. Which of the following represents the correct Lewis structure for carbon tetraiodide
.. .. .. .. ..
a. : C : I : I : I : I :
.. .. .. .. ..
c.
.. .. ..
:..I ..I C..:
: ..I :
b.
..
: ..I :
..
..
: ..I : C : ..I :
..
: ..I :
.. .. .. .. ..
d. : I : I : C : I : I :
.. .. .. .. ..
115.What is wrong with the following Lewis structure? ::Cl :: S :: Cl::
a. Sulfur has too many electrons
b. Chlorine should be in the center of the structure
c. There are too many total electrons - a triple bond is needed
d. There are not enough electrons shown – no double bonds are needed
Review Figure 7 pg 195
Bonds and Electronegativity
116. A bond with an electronegativity difference of less than 0.5 would be categorized as:
a. ionic
b. polar covalent
c. non-polar covalent
d. james
117. A bond with an electronegativity difference of between 0.5-2.1 would be categorized as:
a. ionic
b. polar covalent
c. covalent
d. dura
118. A bond with an electronegativity difference of greater than 2.1 would be categorized as:
a . ionic
b. polar covalent
c. covalent
d. glue
Identify the following as either Polar Covalent, Non-Polar Covalent or Ionic Bonds
MgO ___________________I
H2O __________________C
Cl2 ___________________C
LiCl __________________I
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Covalent Bonds – Naming and Writing
Remember that covalent bonds are Nonmetal to Nonmetal or Metalloid to Nonmetal - Groups 13-17
They are written differently from ionic compounds – Using: The prefix system
Formulas for covalent bonds can also be represented by empirical formulas.
Make sure you review the procedure to find an empirical formula:
% to Grams then Grams to Moles then DIVIDE all by the smallest moles for the ratio. Get Whole
numbers in the end.
Make sure you review the procedure to find the molecular formula from an empirical formula
What is a Molecular formula –
Write names for the following Covalent Compounds:
NF3 ___nitrogen trifluoride__________
S2Cl2 ___disulfur dichloride__________
N2O ___dinitrogen momoxide________
I2O5 ___diiodine pentoxide__________
SiH4 ___silicon tetrahydride__________
C2H6 ___dicarbon hexahydride_______
As2O5 ___diarsenic pentoxide__________
P4O10 ___tetraphosphorous decaoxide__
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Write proper formulas for the following Covalent compounds
Dinitrogen Tetroxide
______ N2O4______
Dihydrogen Monosulfide
____ H2S _____
Sulfur Hexafluoride
______ SF6_______
Nitrogen Triiodide
____ NI3______
Tricarbon Octahydride
______ C3H8______
Iodine Heptachloride
____ ICl7______
Carbon monoxide
______ CO_______
Bromine
____ Br2______
119. The proper formula for the compound dintrogen pentoxide using the prefix system
a. N2O4
b. NO5
c. 5NO2
d. N2O5
120.Covalent bonds are called molecular bonds because:
a. their crystal lattice structures often form common molecular shapes
b. they only occur between molecules of the same type to form a crystal lattice
c. the atoms involved are only bonded to each other and not to other molecules
d. they show actual depictions of molecular shapes and numbers of atoms
121. A Non-Polar covalent bond is unlikely when two different atoms join because the atoms are likely to
differ in:
a. density
b. electronegativity
c. state of matter
d. polarity
Are the following properties characteristics of ionic, covalent, or metallic bonding?
a.
These bonds are formed by delocalized electrons in an “electron sea.” ____________________M
b.
These bonds involve a transfer of electrons. ____________________ I
c.
Substances that are malleable and have high melting points. ____________________ M
d.
Substances that do not conduct electricity and have low melting points. ___________________ C
e.
Compounds that have a rigid crystal lattice structure. ____________________ I
f.
Have bonds are formed by sharing electrons. ____________________ C
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Chapter 7 – The Mole and Chemical Composition --------------------------Know the related conversion factors
1 mol “X”
6.022 x 1023 atoms “X”
Molar mass of compound “X” g
1 mol “X”
or
When speaking of compounds (ionic and covalent) what is Molar Mass? -
What is and Empirical Formula? –
What is a Molecular Formula? –
122. What is the mass of 7.54 mol of NaCl?
a. 58 g
b. 7.69 g
c. 437 g
123. What is the mass of 3.00 mol of H2SO4?
a. 2.94 g
b. .030 g
c. 32.67 g
d. 294 g
124. What is the molar mass of Al2(SO4)3?
a. 278 g
b. 314 g
c. 102 g
d. 342 g
d. .13 g
125. How many moles are in 123 g of Magnesium Oxide?
a. 4920
b. 6.02 x 1023
c. 3.08
--
d. .325
Mass relationships and % composition (mass %) of compounds
Remember that these involve individual atom molar masses and total molar masses of compounds.
126. Determine the mass % of sodium in the compound: Na2CO3.
a. 43.3 % Na
b. 21.6 % Na
c. 23% Na
d. 46% Na
127. Which of the following has the highest mass % of Br?
a. MgBr2
b. KBr
c. FeBr2
d. CaBr2
128. Determine the mass % of each element in the compound : NaOH
a. 57.5% Na, 39.9% O, 2.5% H
b. 39.9%, Na, 57.5% O, 2.5% H
c. 33.3 % Na, 33.3%, 33.3 % H
d. 50 % Na, 50% OH
129.Given what you know about % composition, how many grams of Magnesium can be obtained from a
1200. g sample of magnesium sulfate?
a. 2517.6 g
b. 242.4 g
c. 436.6 g
d. 345.88 g
130.What is the percent composition of water in Na2HPO4 · 12H2O
a. 39.6%
b. 35.4%
c. 60.37%
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d. 64.5%
131.A substance was found to have the following percentages; Zn 23%, S 11%, O 22%, H2O 44%.
What is the empirical formula of this compound?
a. ZnSO4 . 1 H2O
b. ZnSO4 . 7 H2O
c. ZnSO4 . 2 H2O
d. ZnSO4 . 6 H2O
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132. What is the proper name for CoCl4 . 3 H2O
a. Cobalt (II) Chloride TriHydroxide
c. Cobalt (II) Chloride TriHydrate
b. Cobalt (II) Chloride TriHydride
d. Cobalt (II) TriDihydrogenMonoxide
133. Determine the empirical formula for a compound that is 37.51% C, 4.20 % H and 58.29 % O.
a. C6H8O7
b. CHO
c. C3H4O7
d. C12H8O14
134. Determine the empirical formula for a compound that is 29.71% C, 6.22% H and 64.07% Pb.
a. C8H10Pb
b. C8H20Pb
c. CHPb
d. C4H20Pb
135. What is the molecular formula for C3ClH2 if it has a molar mass of 147.00 g/mol?
a. CClH
b. C3Cl2H4
c. C3ClH2
d. C6Cl2H4
136. If 12.94 g of Boron combine with 3.62 g of Hydrogen to produce a compound with a molecular
formula mass of 27.68 g, what is the molecular formula of this compound?
a. BH
b. B2H6
c. BH3
d. B2H
137. A compound is known to have 39.99% Carbon, 6.73% Hydrogen and 53.28% Oxygen. If the actual
molecular formula of the compound is 180.18 g/mol, what is the molecular formula of the compound?
a. CHO
b. C4H8O4
c. C6H12O6
d. C2H4O2
138. Of the following, the only empirical formula is
a. N2F2
b. H2C2
c. N2F4
d. HNF2
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Chapter 8 – Chemical Equations and Reactions -----------------------------139. The reaction type of
H2 (g) + Cl2 (g)
a. synthesis
b. single displacement
→
2HCl (g)
c. decomposition
d. double displacement
140. The reaction type of
2AlCl3 (aq) + 3Pb(NO3)2 (aq)
→
3PbCl2 (s) + 2Al(NO3)3 (aq)
a. combustion
b. single displacement
c. decomposition
d. double displacement
141. The reaction type of
C3H8 (g) + 5O2 (g)
a. synthesis
b. single displacement
142. The reaction type of 2KClO3 (s)
a. combustion
b. single displacement
→
→ 3CO2 (g) + 4H2O(l)
c. combustion
d. double displacement
2KCl (s) + 3O2 (g)
c. decomposition
d. double displacement
143. When CuSO4 (aq) + Fe (s)
→
Fe2(SO4)3 (aq) + Cu (s) is balanced,
the coefficient in front of Cu will be:
a. no coefficient
b. 3
c. 4
d. 8
144. When CO2 (g) + H2O (g)
→
the coefficient in front of O2 will be:
a. no coefficient
b. 3
C6H12O6 (aq)
+
O2 (g)
c. 4
145. In the combustion of propane, the reactants are propane and:
a. water
b. an oxide
c. hydrogen
is balanced,
d. 6
d. oxygen
146. Use the activity series to determine which of the following elements will replace
Copper in Copper (II) Sulfate.
a. Fe
b. Hg
c. Ag
d. Au
147. Use the activity series to determine which of the following elements will replace
Hydrogen in HCl.
a. Cu
b. Ag
c. Au
d. Mg
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148. The state of matter for a reactant or a product in a chemical equation is indicated by
a. coefficient before the formula.
c. symbol after the formula.
b. subscript after the formula.
d. superscript after the formula.
149. An insoluble solid produced by a chemical reaction in solution is called
a. a precipitate.
c. a molecule.
b. a reactant.
d. the mass of the product.
150. How would oxygen be represented in the formula equation for the reaction of methane and oxygen to
yield carbon dioxide and water?
a. oxygen
b. O
c. O2
d. O3
151. The products of the reaction, C2H5OH + 3O2 → 2CO2 + 3H2O have the same _____ as the reactants.
a. atoms.
b. coefficients.
c. molecules
d. subscripts.
152. Which equation is NOT balanced?
a. 2H2 + O2 → 2H2O
c. H2 + H2 + O2 → H2O + H2O
b. 4H2 + 2O2 → 4H2O
d. 2H2 + O2 → H2O
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153. Complete and balance the following reactions. Include all reaction conditions
Solutions of aluminum bromide and magnesium hydroxide are mixed
2AlBr3 (aq)
+
3Mg(OH)2 (aq) →
2Al(OH)3 (s)↓
+
3MgBr2 (aq)
Solutions of iron (III) sulfate and potassium hydroxide are mixed
Fe2(SO4)3 (aq) + 6KOH( aq) → 3K2SO4 (aq) + 2Fe(OH)3 (s)↓
Solutions of calcium sulfate and ammonium phosphate are mixed
3CaSO4 (aq) +
2(NH4)3PO4 (aq)
→
Ca3(PO4)2 (s)↓ +
3(NH4)2SO4 (aq)
Solutions of lithium carbonate and calcium chloride are mixed
Li2CO3 (aq)
+
CaCl2 (aq)
→
2LiCl (aq)
+
CaCO3 (s)↓
Solutions of potassium sulfate and barium nitrate are mixed
K2SO4
(aq)
+ Ba(NO3)2 (aq)
→ 2KNO3 (aq)
+
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Solutions of siver (I) nitrate and nickel (II) sulfate are mixed
2AgNO3(aq) + NiSO4(aq)
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→ Ag2SO4 (s)↓ + Ni(NO3)2 (aq)
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Ionic Compounds in Solutions: Precipitate lab ; Given the following rules:
General Rules for Solubility of Ionic Compounds (salts) in water at 25°C
1. Salts containing Group 1 elements are soluble (Li+, Na+, K+, Cs+, Rb+). Exceptions to this rule are
rare. Salts containing the ammonium ion (NH4+) are also soluble.
2. Salts containing nitrate ion (NO3-) are generally soluble.
3. Salts containing Cl -, Br -, I - are generally soluble. Important exceptions to this rule are salts of Ag
+, Pb2+, and (Hg )2+. Thus, AgCl, PbBr , and Hg Cl are all insoluble.
2
2
2 2
4. Most silver salts are insoluble. AgNO3 and Ag(C2H3O2) are common soluble salts of silver;
virtually anything else is insoluble.
5. Most sulfate salts are soluble. Important exceptions to this rule include BaSO4, PbSO4, Ag2SO4,
and CaSO4.
6. Most hydroxide salts are only slightly soluble. Hydroxide salts of Group 1 elements are soluble.
Hydroxide salts of Group 2 elements (Ca, Sr, and Ba) are slightly soluble*. Hydroxide salts of
transition metals and Al3+ are insoluble. Thus, Fe(OH)3, Al(OH)3, Co(OH)2 are not soluble.
7. Most sulfides of transition metals are highly insoluble. Thus, CdS, FeS, ZnS, Ag2S are all
insoluble. Arsenic, antimony, bismuth, and lead sulfides are also insoluble.
8. Carbonates are frequently insoluble. Group 2 carbonates (CaCO3, SrCO3, and BaCO3) are
insoluble. Some other insoluble carbonates include FeCO3 and PbCO3.
9. Chromates, Phosphates and Fluorides are insoluble. Examples: PbCrO4, Ca3(PO4)2, PbF2.
10. Phosphates are frequently insoluble. Examples: Ca3(PO4)2, Ag3PO4
11. Fluorides are frequently insoluble. Examples: BaF2, MgF2 PbF2.
**The term slightly soluble really means that not enough dissolves for you to notice - it is
therefore insoluble (does not dissolve)
The following solutions are available for a student to mix. Use the RULES FOR SOLUBILITY above.
A) 0.1 M Na3PO4
D) 0.1 M BaCl2
B) 0.1 M K2CO3
E) 0.1 M MgSO4
C) 0.1 M NaNo3
F) 0.1 M NH4NO3
154. Which of the following combination of solutions should produce a precipitate?
a. A + F
b. B + F
c. C + D
d. A + D
155. Which of the following combinations will also produce a precipitate?
a. C + D
b. A + C
c. E + B
156.
d. E + C
Which of the following combinations of solutions will NOT form a precipitate?
a. B + D
b. E + D
c. B + C
d. A + D
157. Use the activity series to predict whether or not the following reactions would occur. If they
occur, finish the reaction and balance – if not – write no reaction Note: Use Zinc (II) and Copper (II)
ZnCl2 (aq) + Cu (s)
Zn (s)
+
CuCl2 (aq) →
Pb (s)
+
Mg(NO3)2 (aq) →
No Reaction
Cu (s)
+
AgNO3 (aq) →
Cu(NO3)2 (aq) + Ag (s)
I2 (s)
+
CuCl2 (aq)
→
No Reaction
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158. A chemical reaction has NOT occurred if the products have:
a. the same mass as the reactants
b. less total bond energy than the reactants
c. more total bond energy than the reactants
d. the same chemical properties as the reactants
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159. An endothermic chemical reaction occurs when:
a. plants take in energy from the sun to make food
b. a match ignites
c. copper chloride reacts with water in the “in the bag” lab
d. sugar in plants combines with oxygen to give off energy
160. In the chemical equation:
3CuSO4 (aq) + 2Fe (s)
→
Fe2(SO4)3 (aq)
One (1) mole of Fe will yield:
a. 1.0 mol of Fe2(SO4)3 and 3 mol of Cu
b. 1.0 mol of Fe2(SO4)3 and 1.5 mol of Cu
c. 0.5 mol of Fe2(SO4)3 and 1.5 mol of Cu
d. 0.33 mol of Fe2(SO4)3 and 1.0 mol of Cu
Chapter 9 - Stoichiometry -
+ 3Cu(s)
Problems will include:
-Converting MOLES to MOLES using ratios
-Converting MOLES to MASS and MASS to MASS
-Predicting the amounts of products/reactants needed/produced – given an amount and an excess
-Identifying a limiting reactant. Using limiting reactant to predict the amount of product obtained
-Calculating the theoretical yield in a chemical reaction
-Calculating the actual % yield in a chemical reaction
161. You need at least 1.03 x 10–3 mol of O2 every minute. If all this oxygen is used for the cellular
respiration process that breaks down glucose C6H12O6 - Given the following equation:
C6H12O6 (s)
+
6 O2 (g)
→ 6 CO2 (g)
+
6 H2O (g)
How many grams of glucose does your body consume each minute?
1.03 x 10–3 mol of O2
1 mol C6H12O6
6 mol O2
180.18 g C6H12O6
1 mol C6H12O6
=
.031 g C6H12O6
=
.0453 g CO2
=
.0816 g H2O
How many grams of Carbon Dioxide are produced in the reaction above?
1.03 x 10–3 mol of O2
6 mol CO2
6 mol O2
44.01 g CO2
1 mol CO2
How many grams of Water are produced in the above reaction?
1.03 x 10–3 mol of O2
6 mol H2O
6 mol O2
18.02 g H2O
1 mol H2O
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Given the following unbalanced equation:
2 BF3 + 3 H2O
→
B2O3 +
6 HF
162. What is the mole ratio between Dihydrogen monoxide and Hydrogen Monofluoride
a. 1 : 4
b. 3 : 6
c. 6 : 3
d. 2 : 3
Consider the equation:
Cu(s) + HNO3(aq) → Cu(NO3)2(aq) + NO(g) + H2O(l)
163. What is the mole ratio between Copper (Cu) and Nitrogen Monoxide (NO)
a. 1 : 1
b. 3 : 2
Consider the following for 164-167
c. 1: 2
d. 2 : 3
C3H8 (g) + 5O2 (g)
3CO2 (g) + 4H2O(l)
→
164. If given 19.00 moles of C3H8 - how many moles of CO2 would be produced?
a. 19.00 moles
b. 3.00moles
c. 9.50 moles
d. 57.00 moles
165. How many grams of CO2 would be produced in the previous question?
a. 57.00g
b. 1.295g
c. 2508 g
d. 44.01g
166. If given 165.6g of Oxygen gas, how many grams of H2O would be produced?
a. 74.60 g
b. 18.02g
c. 5.175g
d. 4.140g
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167. If given 12.5 moles of Oxygen gas, how many grams of CO2 would be produced?
a. 32.00g
b. 44.01g
c. 330.1g
d. 7.500g
168. If you were given 50.0 g of each reactant, which would be the limiting reactant?
2AlCl3 (aq) + 3Pb(NO3)2 (aq) → 3PbCl2 (s) + 2Al(NO3)3 (aq)
a. AlCl3
b. Pb(NO3)2
c. PbCl2
d. Al(NO3)3
169. Given 20.0 g of Al and 30.0 g O2 , the limiting reactant would be……….
4Al(s) + 3O2(g) → 2Al2O3(s)
a. Al
b. O2
c. PbCl2
d. Al2O3
170. In the question above, How many grams of Al2O3 would be produced?
a. 3.63 x 10-3 g
b. 37.82 g
c. 27.81 g
171. Consider the following:
3CuSO4 (aq) + 2Fe (s)
→
Fe2(SO4)3 (aq)
d. 151 g
+ 3Cu(s)
If you were given 50.0 g of Copper (II) Sulfate, predict the amount Iron (III) Sulfate that could be
produced with an excess of Iron?
a. 41.6 g
b. 2.60 x 10-4 g
c. 0.104 g
d. 21.6 g
172. Consider the following:
3AgNO3 (aq) + Na3PO4 (aq)
→
Ag3PO4 (s) + 3NaNO3 (aq)
Given an excess of Sodium Phosphate and 120.5 g of Silver Nitrate, how much Silver Phosphate is
produced?
a. 5.638 x 10-4 g
b. .001163 g
c. 98.97g
d. 890.3 g
173. Consider the following: 2AlCl3 (aq) + 3Pb(NO3)2 (aq)
→ 3PbCl2 (s) +
2Al(NO3)3
Given 92.0g Aluminum Chloride and 92.0 g of Lead (II) Nitrate, how much Lead (II) Chloride
would be produced?
a. 289.22g
b. .278g
c. 77.31 g
d. 278.10g
174. An experiment was designed to produce 345g of NaCl. The experiment only produced 235 g. What is
the % yield of the reaction?
a. 146 %
b. 68%
c. 50%
d. 48%
175. 400.0 g of Hydrogen gas are added to an excess of Nitrogen gas. The reaction produces 2040.0 g of
NH3. What is the % yield of the reaction?
a. 100%
b. 50%
c. 91%
d. cannot be determined
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