Year-Long Chemistry First Semester Exam Review
Short Answer
1. The following length measurements were taken by students using several different measuring devices. Find
the average of the measurements. Make sure that your answer has the correct number of significant figures.
10.05 cm, 10.1 cm, 9.741 cm, 10.6 cm, 10.5 cm
2. Round off the measurement 0.003 095 5 m to three significant figures.
3. What is the sum of 6.210 L and 3 L expressed in the correct number of significant digits?
4. What is the product of the number 1000 and the measurement 0.003 57 m expressed in the correct number of
significant digits?
5. The mass of the electron is 9.109 39
figures.
10
6. Perform the following operation: 3.43 cm
significant figures.
kg. Express the mass of the electron to 1, 2, 3, and 4 significant
5.2 cm. Make sure that your answer has the correct number of
7. What is the temperature 128 K expressed in degrees Celsius?
8. Express 0.05 grams in kilograms, using the correct abbreviations.
9. Express 0.06 liters in cubic meters, using the correct abbreviations.
10. What is the density of an object having a mass of 4.0 g and a volume of 39.0 cubic centimeters?
11. A cube of a gold-colored metal with a volume of 59 cm has a mass of 980 g. The density of pure gold is 19.3
g/cm . Is the metal pure gold? Show calculations to justify your answer.
12. The density of osmium, which is the densest metal, is 22.57 g/cm . What is the mass of a block of osmium
that measures 1.00 cm 4.00 cm 2.50 cm?
13. List the number of protons, neutrons, and electrons in
C.
14. Consider an element Z that has two naturally occurring isotopes with the following percent abundances: the
isotope with a mass number of 19.0 is 55.0% abundant; the isotope with a mass number of 21.0 is 45.0%
abundant. What is the average atomic mass for element Z?
15. Give the electron configuration for a neutral atom of beryllium.
16. Give the electron configuration for a neutral atom of selenium.
17. Write the electron configuration for chromium.
18. What is the frequency of ultraviolet light with wavelength 2.94
10
m? (c = 3.00
10 m/s)
19. What is the wavelength of gamma ray electromagnetic radiation with a frequency of 2.73
3.00 10 m/s)
20. What is the frequency of a radar photon with energy 6.80
10
J? (h = 6.63
21. Which group of elements in the periodic table is known as the alkali metals?
22. Which group in the periodic table is known as the noble gases?
10
J s)
10
Hz? (c =
23. An element has an atomic number of 80. How many protons and electrons are in an atom of the element?
24. What is the electron configuration of oxygen?
25. What is the electron configuration of sulfur?
26. In which group in the periodic table do the elements have the highest electronegativity values?
27. Write the formula for the compound rubidium phosphide.
28. Write the formula for the compound boron chloride.
29. How many representative particles are in 1.45 g of a molecular compound with a molar mass of 237 g?
30. Find the number of moles of argon in 607 g of argon.
31. Find the mass, in grams, of 1.40 10 molecules of N .
Numeric Response
32. What is the melting point of water in degrees Celsius?
33. In how many physical states does water commonly exist?
34. What is the boiling point of water in degrees Celsius?
35. How many grams of liquid water are produced when 60 grams of ice melt?
36. What is the relative charge carried by an electron?
37. What is the relative charge of a proton?
38. About how many more times massive is a proton than an electron?
39. Use the periodic table to determine the number of electrons in a neutral atom of lithium.
40. Use the periodic table to determine the number of protons in an atom of barium.
41. How many protons are present in an atom of Be-9?
42. What is the total number of subatomic particles in the nucleus of an atom of
Bi?
43. Determine the number of electrons in an atom of iridium.
44. What is the atomic number for an element with 41 neutrons and a mass number of 80?
45. How many electrons are in an atom of gold?
46. What is the mass number for an oxygen atom that has 10 neutrons in its nucleus?
47. How many protons are present in the nuclei of the three known isotopes of hydrogen?
48. Use the periodic table to determine the number of neutrons in nitrogen-14.
49. How many neutrons are present in an atom of the isotope
50. Calculate the number of neutrons in
U?
Pb.
51. How many electrons are in the highest occupied energy level of a neutral strontium atom?
52. How many electrons are present in the d sublevel of a neutral atom of nickel?
53. What is the usual charge on an ion from Group 7A?
54. How many electrons does the ion Ca
contain?
55. How many valence electrons are in rubidium?
56. What is the charge of a particle having 9 protons and 10 electrons?
57. How many valence electrons does an iodine atom have?
58. How many electrons are shared in a single covalent bond?
59. How many electrons are shared in a double covalent bond?
60. How many covalent bonds are in a covalently bonded molecule containing 1 phosphorus atom and 3 chlorine
atoms?
61. How many unshared pairs of electrons are in a molecule of hydrogen iodide?
62. What is the charge on the cation in CuSO ?
63. What is the ionic charge on the zirconium ion in the ionic compound zirconium oxide, ZrO ?
64. How many iron(II) ions combine with oxygen to form iron(II) oxide?
65. Determine the subscript for ammonium in the chemical formula for ammonium dichromate.
Essay
66. Name five areas of chemistry research and give examples of the types of problems chemists are trying to
address within each area.
67. How did Lavoisier transform the field of chemistry in the late 1700s?
68. What is the difference between a solution and a heterogeneous mixture? Give an example of each.
69. Give five examples of heterogeneous mixtures and five examples of homogeneous mixtures.
70. Give an example of a two-phase mixture and describe how you would separate the substances.
71. Define element and compound. Explain the difference between an element and a compound.
72. What does the chemical formula of a compound indicate? Use an example to illustrate your points.
73. Explain the difference between chemical properties and physical properties. Give an example of each.
74. Discuss the difference between physical changes and chemical changes. Use an example to illustrate the
difference.
75. State the law of conservation of mass. Then apply the law to this question: What would be the total mass of
the products of a reaction in which 10 grams of water decomposes into the elements hydrogen and oxygen?
76. Explain the difference between precision and accuracy. Suppose you made three different mass measurements
of a sugar sample you knew to have a mass of 1 g. How would you know whether or not the measurements
were accurate? How would you know whether or not they were precise? Could the three measurements be
precise, but not accurate? Explain.
77. Describe the rules that are used to determine the number of significant figures in the results of addition,
subtraction, multiplication, and division.
78. You are given a jar containing a mixture of three different solid substances. The particles of each substance
are larger than 0.5 cm in diameter. All you know about the substances is that each is a pure metal. Describe a
method for determining the identities of the metals that make up the mixture. Assume that each metal is a
different color, so you can tell one from the others.
79. Explain how Dalton improved upon atomic theory more than 2000 years after Democritus’s hypotheses about
atoms.
80. What observations by Rutherford led to the hypothesis that atoms are mostly empty space, and that almost all
of the mass of the atom is contained in an atomic nucleus?
81. In what way are two isotopes of the same element different? Explain why isotopes of the same element have
the same chemical behavior.
82. Describe the different principles that govern the building of an electron configuration.
83. What is the quantum mechanical model?
84. Explain what is meant by the Heisenberg uncertainty principle.
85. Describe the trends in the atomic size of elements within groups and across periods in the periodic table.
Provide examples.
86. Explain how ions form. Provide examples.
87. Describe the trends in first ionization energy within groups and across periods in the periodic table. Provide
examples.
88. Describe the trends in electronegativity within groups and across periods in the periodic table. Provide
examples.
89. Explain how atoms (ions) are held together in an ionic bond. Give an example of an ionic compound.
90. Explain what is meant by VSEPR theory. Give an example of how VSEPR theory can be applied to predict
the shape of a molecule.
91. Explain what a polar molecule is. Provide an example.
92. What determines the degree of polarity in a bond? Distinguish between nonpolar covalent, polar covalent, and
ionic bonds in terms of relative polarity.
93. What are dispersion forces? How is the strength of dispersion forces related to the number of electrons in a
molecule? Give an example of molecules that are attracted to each other by dispersion forces.
94. Compare the compositions of binary ionic and binary molecular compounds.
95. Given the name of a binary molecular compound, describe how to write its formula. Use carbon tetrachloride
as an example.
96. Explain how to write an ionic formula, given an anion and a cation. As an example, use the phosphate anion.
Write formulas for the compounds produced by the combinations of these ions. Name the compounds for
which you have written formulas.
97. Name the compounds CuBr , SCl , and BaF . Explain the use or omission of the Roman numeral (II) and the
prefix di-.
98. Explain what a saturated solution is. Give a specific example.
99. Discuss the phenomenon of supersaturation. Indicate how crystallization can be initiated in a supersaturated
solution.
100. Discuss the different factors that can affect the solubility of a substance. Include specific examples in your
discussion.
Year-Long Chemistry First Semester Exam Review
Answer Section
SHORT ANSWER
1. Average = (10.05 + 10.1 + 9.741 + 10.6 + 10.5) / 5 = 10.2 cm
2. 0.003 10 m
3. 9 L
4. 3.57 m
5. one: 9
10
6. 3.43 cm
7.
kg; two: 9.1
10
kg; three: 9.11
10
kg; four: 9.109
10
5.2 cm = 18 cm
C = K – 273 = 128 – 273 = –145 C
8. 0.05 g
1 kg/1000 g = 0.000 05 kg
9. 0.06 L
1 m /1000 L = 0.000 06 m
10. Density = mass/volume = 4.0 g/39.0 cm = 0.10 g/cm
11. No; mass/volume = 980 g/59 cm = 17 g/cm
12. Mass = density
volume = 22.57 g/cm
1.00 cm
13. protons: 6; neutrons: 7; electrons: 6
14. 19.9 amu
15. 1s 2s
16. 1s 2s 2p 3s 3p 3d 4s 4p
17. 1s 2s 2p 3s 3p 3d 4s
18.
= c/ =
1.02
10
= 1.02
s
= 1.02
10
10
s
Hz
4.00 cm
2.50 cm = 226 g
kg
19.
= c/ =
20.
= E/h
= 1.10
10
m
=
= 1.03
= 1.03
10 s
10 Hz
21. 1A
22. 8A
23. 80 protons, 80 electrons
24. 1s 2s 2p
25. 1s 2s 2p 3s 3p
26. 7A
27. Rb P
28. BCl
29. 1.45 g 1.00 mol/237 g 6.02
= 3.68 10 molecules
10
molecules/1.00 mol
30. 607 g Ar
1 mol Ar/39.9 g Ar = 15.2 mol Ar
31. 1.40 10
= 6.51 g
molecules N
NUMERIC RESPONSE
32.
33.
34.
35.
36.
37.
38.
39.
40.
0
3
100
60
1–
1+
1840
3
56
(1.00 mol N /6.02
10
molecules N )
(28.0 g N /1 mol N )
41.
42.
43.
44.
45.
46.
47.
48.
49.
50.
51.
52.
53.
54.
55.
56.
57.
58.
59.
60.
61.
62.
63.
64.
65.
4
209
77
39
79
18
1
7
143
128
2
8
1
18
1
1
7
2
4
3
3
2
4
1
2
ESSAY
66. For each area, students should provide at least one example that reflects an overall goal. For example,
materials must meet a specific need such as increased strength and reduced weight. The demand for energy
requires conservation and production, including the development of alternate energy sources. Chemists
contribute to the health and well-being of humans through the development of medicines and materials for
replacement body parts. Chemists help to develop more productive crops and more effective, safer ways to
protect crops from pests and diseases. Chemists help to identify pollutants and prevent pollution. Chemists
analyze materials from locations other than Earth directly or indirectly.
67. Lavoisier helped transform chemistry from a science of observation to a science of measurement. Toward this
end, he developed a balance that could measure small differences in mass. He was part of a general movement
toward basing conclusions on experimental evidence.
68. A solution is a mixture that has a uniform composition. The composition of a heterogeneous mixture is not
uniform throughout. All solutions are mixtures, but not all mixtures are solutions. Salt water is a solution;
sand in water is a heterogeneous mixture.
69. The examples should follow this rule: A homogeneous mixture is uniform in composition, while a
heterogeneous mixture is not. Examples of a homogeneous mixture include air and stainless steel. Examples
of a heterogeneous mixture include chicken noodle soup and the mixture of oil and vinegar.
70. One example would be a solution of salt in water. The salt and water can be separated by distillation. Another
example would be a mixture of iron and sulfur, which can be separated by using a magnet to attract the iron.
71. An element is the simplest form of matter that has a unique set of properties. A compound is a substance that
contains two or more elements chemically combined in a fixed proportion. An element cannot be broken
down into simpler components through chemical reactions. Compounds are substances that can be broken
down into simpler substances through chemical reactions.
72. The chemical formula shows the elements that are in the compound and gives the proportions in which
elements are found in the compound. The formula for sodium chloride (NaCl), for instance, indicates that the
elements in the compound are sodium (Na) and chlorine (Cl). The formula also indicates that the elements are
present in a proportion of 1:1 in the compound.
73. ANS:
Chemical properties describe the ability of a substance to undergo particular chemical changes. Chemical
properties can only be observed when a substance undergoes a change in composition. Physical properties can
be observed without altering the composition of a substance. The ability of iron to rust is a chemical property
of iron. The fact that iron is a solid at room temperature is a physical property of iron.
74. Physical changes do not alter the composition of a substance, but chemical changes do. When water is boiled,
the resulting gas is still water. However, when water is broken down into hydrogen and oxygen in a chemical
reaction, the water no longer exists. The latter is a chemical change; the former a physical change.
75. The law of conservation of mass states that in any physical change or chemical reaction, mass is neither
created nor destroyed. According to this law, the total mass the products in the decomposition of water would
be 10 grams. The total mass of the products would be the same mass as the total mass of the reactants.
76. Precision is the reproducibility of a measurement made under the same conditions; accuracy is the closeness
of a measurement to the true value of the measurement. The three measurements would be precise if they
were very close to each other in value; they would be accurate if they were close to the actual 1-g mass of the
sample. If the measurements are very close to each other, they are precise, regardless of how close they are to
the real value. Therefore, the measurements could be precise, but not accurate.
77. The answer of an addition or subtraction can have no more digits to the right of the decimal point than are
contained in the measurement with the least number of digits to the right of the decimal point. The answer of
a multiplication or division can have no more significant figures than the measurement having the least
number of significant figures. For multiplication and division, the position of the decimal point has nothing to
do with the number of significant figures.
78. Each substance is a pure metal, so they should be identifiable by determining their densities. Separate the
solids by color into three separate batches. Determine the mass of each metal using a balance. Determine the
volume of each metal as follows. Immerse each metal in a measured volume of a suitable liquid in a
graduated cylinder. Measure the volume of the liquid plus the metal. Calculate the difference between the
volume of the liquid plus the metal and the liquid by itself. To calculate the density, divide the mass of each
metal by its volume. Compare the density and color of each metal to the density and color listed in reference
books.
79. John Dalton used experimental methods to hypothesize that atoms of the same element are identical, and that
they differ from atoms of other elements. Dalton also noted that atoms of different elements can form
compounds in which the elements combine in whole-number ratios. He noted that individual atoms are not
transformed into different atoms as a result of a chemical reaction. What Dalton did not contribute was the
fact that individual atoms are divisible into subatomic particles, which came from the discovery of the
electron and the atomic nucleus by Thomson and Rutherford, respectively.
80. Rutherford’s gold-foil experiment led to this hypothesis. Alpha particles were observed to mostly pass
through a gold foil, which suggests that the volume of individual gold atoms consists mainly of empty space.
The observation that some alpha particles were scattered at large angles led to the suggestion that the gold
atom has a central core, or nucleus, composed of a concentrated mass capable of deflecting the alpha
particles.
81. Isotopes of the same element have different numbers of neutrons, and therefore, different mass numbers and
different atomic masses. Isotopes of the same element have the same number of protons and electrons. The
electrons, not the neutrons, are responsible for an element's chemical behavior.
82. The aufbau principle states that electrons enter the orbitals of lowest energy first. The Pauli exclusion
principle states that each orbital can hold only two electrons. Hund's rule states that electrons first enter
separate orbitals of the same energy, with each electron having the same spin, before pairing with electrons
that have opposite spins.
83. It is a model that describes subatomic particles and atoms as waves. Schrodinger applied a mathematical
model of the wave/particle nature of matter to hydrogen. Solutions to the Schrodinger equation determine the
energies an electron can have and how likely it is to find the electron in various locations.
84. The measurement of the speed or position of a moving particle necessarily involves an interaction with the
particle. Therefore, the position or the speed of the particle is changed as a result of the measurement. As a
consequence, accurate measurements of both these variables cannot be made at the same time.
85. Atomic size increases with increasing atomic number within a group. For example, sodium atoms are larger
than lithium atoms, and potassium atoms are larger than sodium atoms. Atomic size decreases with increasing
atomic number across a period. For example, lithium atoms are larger than beryllium atoms, and beryllium
atoms are larger than boron atoms.
86. Ions form when electrons are transferred among atoms. For example, a group 1A element, such as potassium,
tends to transfer one electron to other atoms, causing it to form a net positive charge, in this case K . Such
ions with net positive charges are called cations. Ions with net negative charges are anions. Nonmetal atoms,
such as chlorine, tend to accept electrons from other atoms. Chlorine tends to gain a single electron, forming
the anion Cl .
87. First ionization energies decrease from top to bottom within a group and increase across a period from left to
right. For example, the first ionization energy of rubidium is less than that of lithium. The first ionization
energy of iodine is much greater than that of lithium.
88. Electronegativity values decrease from top to bottom within a group, and from right to left across a period.
For example, rubidium is less electronegative than lithium. Lithium is less electronegative than fluorine.
89. In an ionic bond, oppositely charged ions are held together by the electronic force of attraction that exists
between oppositely charged particles. In the ionic compound, anions and cations are present in a ratio that
causes the total charge on the compound to be zero. Sodium phosphide, Na P, has three sodium ions for each
phosphide ion. This ratio insures a zero total charge given the charges on the two individual ions (Na = 1 , P
= 3–).
90. VSEPR (valence-shell electron-pair repulsion) theory states that because electron pairs repel, molecules adjust
their shapes so that the valence-electron pairs, both bonding and non-bonding, are as far apart as possible.
Methane, CH , for example, has four bonding electron pairs and no unshared pairs. The bonding pairs are
farthest apart when the angle between the central carbon and each of its attached hydrogens is 109.5 . This is
the angle that is observed experimentally.
91. A polar molecule is one in which one end of the molecule has a slightly negative electric charge and the other
end has a slightly positive electric charge. An example of a polar molecule is water. The oxygen atom in
water develops a slightly negative charge and the hydrogen atoms develop slightly positive charges because
of the difference in electronegativity between the oxygen and hydrogen atoms.
92. The relative electronegativity of the two bonded atoms determines the polarity of a bond. If the difference in
electronegativities between the two atoms is less than 0.4, the bond is nonpolar covalent. If the difference in
electronegativities between the two atoms is 0.4 to 1.0, the bond is moderately polar covalent. If the
difference in electronegativities between the two atoms is 1.0 to 2.0, the bond is highly polar covalent. If the
difference in electronegativities between the two atoms is more than 2.0, the bond is ionic.
93. Dispersion forces are the weakest of all molecular interactions, and are thought to be caused by the motion of
electrons. Generally, the strength of dispersion forces increases as the number of electrons in a molecule
increases. Diatomic molecules of halogen elements are an example of molecules whose attraction for one
another is caused by dispersion forces.
94. Binary ionic and binary molecular compounds are composed of two elements. Ionic compounds consist of
oppositely charged ions. Molecular compounds consist of molecules. Ionic compounds are composed of one
metallic cation and a nonmetallic anion. Binary molecular compounds are composed of two nonmetallic
elements.
95. Begin by writing the chemical symbol for the first element in the compound name on the left, and then write
the chemical symbol for the next element on the right. Prefixes that describe elements in the compound names
determine the subscripts for those elements. For example, to write the formula for carbon tetrachloride, first
write the symbols for the chemical elements as indicated in the name “CCl.” Next, the “tetra” in front of
“chloride” indicates the presence of 4 chlorine atoms, which should be indicated by the subscript “4.”
Therefore, the correct formula is CCl .
96. The cation is written first, the anion second. The resulting formula must indicate an electrically neutral
substance. The charge of the anion becomes the subscript of the cation, and the charge of the cation becomes
the subscript of the anion. When the charges of the two ions are the same, no subscripts are written.
Example 1: Cu , PO
Example 2: Cu , PO
combine as copper(I) phosphate Cu PO
combine as copper(II) phosphate Cu (PO )
97. CuBr is copper(II) bromide. The name must include a Roman numeral because copper is a transition element
that can form ions with more than one charge. SCl 2 is sulfur dichloride. The compound is named with
prefixes because sulfur and chlorine are both nonmetals and thus form a molecular compound. BaF is barium
fluoride. A Roman numeral is not needed in this name because barium is a Group A metal and forms only the
2+ ion. Prefixes are not used in ionic compounds.
98. A saturated solution contains the maximum amount of solute for a given amount of solvent at a constant
temperature. For example, no more than 36.2 g of sodium chloride will dissolve in 100 g of water. In a
saturated solution, a dynamic equilibrium exists between the solid and its dissolved ions. In this equilibrium,
as many ions are going out of the solution as are going in per unit time.
99. A solution that contains more solute than it can theoretically hold at a given temperature is a supersaturated
solution. Crystallization can be initiated by adding a seed crystal to the solution or by exposing the solution to
a rough surface. The latter can be done by scratching the inside of the container holding the solution.
100. The factors are temperature, pressure, and the nature of the solute and solvent. Specific examples include the
following. Potassium nitrate is more soluble in water at high temperature than at low temperature. Gases are
less soluble at high temperatures than at low temperatures. The solubility of a particular gas increases as the
partial pressure of that gas above the solution increases. Sodium nitrate is much more soluble in water than is
barium sulfate, regardless of temperature, because the attractive forces between the ions in BaSO are
stronger than the attractive force of the solvent molecules.