Download February Homework Packet

UNIT 1: INTRODUCTION TO MATTER
Key Points:
 Matter is anything that has mass and takes
up space
 Matter can be classified as pure
substances or mixtures
 Elements and compounds are pure
substances
 Mixtures can be either homogenous or
heterogeneous
 Mixtures can be separated by physical
means using one of the following
techniques: distillation, crystallization,
filtration, chromatography
 An atom is the simplest form of matter
 Matter can be a in the solid, liquid, or gas
phase
 Chemistry is the study of the change in
matter
 Physical changes are changes in matter in
which the appearance of a substance
changes but the identity of the compound
remains the same
 Chemical changes are changes in matter in
which the identity is changed
 Physical properties include color, smell,
freezing point, boiling point, melting point,
density, etc.
 Chemical properties include reactivity
 In all chemical and physical changes, mass,
energy, and charge must always be
conserved
Questions:
1. How do you distinguish between an element,
compound, and mixture looking at just the
chemical formula?
2. Which type of matter can be separated by
PHYSICAL MEANS?
3. Which type of matter can be separated by
CHEMICAL MEANS?
4. What are three examples of physical properties?
5. What are three examples of chemical
properties?
6. What are two types of pure substances?
7. What are the physical properties that is used to separate mixtures in a) distillation, b) crystallization,
c) filtration, and d) chromatography?
8. How do you know if a substance is a solid, liquid, gas, or mixture based on the chemical formula?
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UNIT 1 REGENTS PRACTICE:
1. Which substance represents a compound?
(1) C (s)
(2) Co (s)
(3) CO (g)
(4) O2 (g)
2. Two substances, A and Z, are to be
identified. Substance A can not be broken
down by a chemical change. Substance Z can
be broken down by a chemical change.
What can be concluded about theses
substances?
(1) Both substances are elements.
(2) Both substances are compounds.
(3) Substance A is an element and substance
Z is a compound.
(4) Substance A is a compound and
substance Z is an element.
3. Which particle diagram represents one pure
substance, only?
6. Which statement describes a chemical
property of hydrogen gas?
(1) Hydrogen gas burns in air.
(2) Hydrogen gas is colorless.
(3) Hydrogen gas has a density of 0.00009
g/cm3 at STP.
(4) Hydrogen gas has a boiling point of 20. K
at standard pressure.
7. Which substance can be broken down by a
chemical change?
(1) antimony
(2) carbon
(3) hexane
(4) sulfur
8. Which statement describes a chemical
property that can be used to distinguish
between compound A and compound B?
(1) A is a blue solid, and B is a white solid.
(2) A does not corrode in acid, and B does
corrode in acid.
(3) A has a high melting point, and B has a
low melting point.
(4) A dissolves in water, and B does not
dissolve in water.
9. An aqueous solution of sodium chloride is
best classified as a
(1) homogeneous compound
(2) homogeneous mixture
(3) heterogeneous compound
(4) heterogeneous mixture
4. Which grouping of the three phases of
bromine is listed in order from left to right
for increasing distance between bromine
molecules?
(1) gas, liquid, solid
(2) liquid, solid, gas
(3) solid, gas, liquid
(4) solid, liquid, gas
5. Which process represents a chemical
change?
(1) melting of ice
(2) corrosion of copper
(3) evaporation of water
(4) crystallization of sugar
10. Which process is a chemical change?
(1) melting of ice
(2) boiling of water
(3) subliming of ice
(4) decomposing of water
11. Which two substances can not be broken
down by chemical change?
(1) C and CuO
(2) C and Cu
(3) CO2 and CuO
(4) CO2 and Cu
2
12. Which statement best describes the shame
and volume of an aluminum cylinder at
STP?
(1) It has a definite shape and a definite
volume.
(2) It has a definite shape and no definite
volume.
(3) It has no definite shape and a definite
volume.
(4) It has no definite shape and no definite
volume.
14. Given the particle diagram representing
four molecules of a substance:
Which particle diagram best represents this
same substance after a physical change has
taken place?
13. Which particle diagram represents a
mixture of element X and element Z, only?
15. Which statement describes a chemical
property of oxygen?
(1) Oxygen has a melting point of 55 K.
(2) Oxygen can combine with a metal to
produce a compound.
(3) Oxygen gas is slightly soluble in water.
(4) Oxygen gas can be compressed.
16. Which diagram best represents a gas in a
closed container?
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Base your answers to questions 17-19 on the diagram below concerning the classification of matter.
17. What type of mixture is represented by X? [1]
18. What type of substance is represented by Z? [1]
19. Given a mixture of sand and water, sand one process that can be used to separate the water from the
sand. [1]
Base you answers to questions 20 through 22 on the particle diagrams below.
20. Explain, in terms of composition, why sample A represents a pure substance. [1]
21. Explain why sample C could represent a mixture of fluorine and hydrogen chloride. [1]
22. Contrast sample A and sample B, in terms of compounds and mixtures. Include both sample A and
sample B in your answer. [1]
4
Base your answers to questions 24-26 on the diagram of a molecule of nitrogen shown below:
24. Draw a particle model that shows at least six molecules of nitrogen gas. [1]
25. Draw a particle model that shows at least six molecules of liquid nitrogen. [1]
26. Describe, in terms of particle arrangement, the difference between nitrogen gas and liquid nitrogen.
[1]
Base questions 27 and 28 on the information below.
In an investigation, a dripless wax candle is massed and then lighted. As the candle burns, a small
amount of liquid wax forms near the flame. After 10 minutes, the candle’s flame is extinguished and the
candle is allowd to cool. The cooled candle is massed.
27. Identify one physical change that takes place in this investigation. [1]
28. State one observation that indicates a chemical change has occurred in this investigation. [1]
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UNIT 3: ATOMIC CONCEPTS
Key Points:
 The plum pudding model was the first
model of the atom and presented the atom
as a hard sphere with charges evenly
distributed
 Rutherford’s gold foil experiment
concluded that the atom had a positively
charged nucleus and that the atom is
mostly empty space
 The Bohr model suggests that electrons
travel in circular orbits
 The wave-mechanical model of the atom
claims that electrons exist in orbitals,
regions with high probability of electron
location
 The atom is made up of two parts: the
nucleus and the electron cloud
 There are three subatomic particles that
make up the atom: protons, electrons,
neutrons
 Every element is defined solely by the
number of protons
 The atomic number and atomic mass can
be found on the periodic table
 The atomic number is the number of
protons in an atom and the atomic mass is
the weighted average of the masses of the
isotopes (atomic mass also equals the
number protons plus the number of
neutrons).
 Isotopes are atoms of the same element
that have different number of neutrons
 The electron configuration of an atom tells
us how many electrons exist on each
energy level of the atom
 A ground state atom and an excited state
atom can be determined by looking at its
electron configuration
 Valence electrons are the electrons that
exist on the outer most shell of an atom.
They can be represented via a Lewis dot
diagram
Questions:
1. Draw a model of the atom and label the proton,
neutron, electron, and the nucleus.
2. What is the difference between an electron orbit
and an orbital?
3. What is the mass of 1 proton?
4. What is the mass of 1 neutron?
5. What is the mass of 1 electron?
6. How do you count the number of protons in an
atom?
7. How do you count the number of neutrons in an
atom?
8. How do you count the number of electrons in an
atom?
9. How is an atom different from an ion?
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10. What does the atomic mass tell us?
11. What is an isotope?
12. Where in the atom do electrons have the lowest energy? Highest energy?
13. How do you find the number of valence electrons in an atom?
14. Why is the number of valence electrons so important in chemistry?
15. How do you excite an electron?
16. What is produced when an excited electron comes back to the ground state?
17. How do you know the difference between a ground state electron configuration and an excited state
electron configuration?
18. Why is the atomic mass given in the reference table a decimal?
19. Explain the steps required to calculate the average atomic mass of an element?
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UNIT 3 REGENTS PRACTICE:
1. Which subatomic particle has no charge?
(1) neutron
(2) electron
(3) alpha particle
(4) beta particle
2. The atomic number of an atom always equal to
the number of its
(1) protons, only
(2) protons plus electrons
(3) neutrons, only
(4) protons plus neutrons
3. Which two particles each have a mass
approximately equal to one atomic mass unit?
(1) electron and neutron
(2) proton and electron
(3) proton and neutron
(4) electron and positron
4. According to the wave-mechanical model of
the atom, electrons in an atom
(1) are most likely found in the excited state
(2) have a positive charge
(3) are located in orbitals outside the nucleus
(4) travel in defined circles
5. What is the total charge of the nucleus of a
carbon atom?
(1) 0
(2) +12
(3) -6
(4) +6
6. The nucleus of an atom of K-42 contains
(1) 20 protons and 19 neutrons
(2) 23 protons and 19 neutrons
(3) 19 protons and 23 neutrons
(4) 19 protons and 42 neutrons
7. What is the total number of protons in an atom
with the electron configuration 2-8-18-32-181?
(1) 69
(2) 118
(3) 197
(4) 79
8. How do the energy and the most probable
location of an electron in the third shell of an
atom compare to the energy and the most
probable location of an electron in the first
shell of the same atom?
(1) In the third shell, an electron has more
energy and is closer to the nucleus.
(2) In the third shell, an electron has less
energy and is closer to the nucleus.
(3) In the third shell, an electron has less
energy and is farther from the nucleus.
(4) In the third shell, an electron has more
energy and is farther from the nucleus.
9. An atom of oxygen is in an excited state. When
an electron in this atom moves from the third
shell to the second shell, energy is
(1) absorbed by the nucleus
(2) absorbed by the electron
(3) emitted by the nucleus
(4) emitted by the electron
10. Which statement best describes the nucleus of
an aluminum atom?
(1) It has a charge of -13 and is surrounded by
a total of 10 electrons.
(2) It has a charge of -13 and is surrounded by
a total of 13 electrons.
(3) It has a charge of +13 and is surrounded by
a total of 10 electrons.
(4) It has a charge of +13 and is surrounded by
a total of 13 electrons.
11. Which list consists of elements that have the
most similar chemical properties?
(1) Mg, Ca, and Ba
(2) K, Al, and Ni
(3) K, Ca, and Ga
(4) Mg, Al, and Si
12. Which two elements have the most similar
chemical properties?
(1) Be and Mg
(2) Cl and Ar
(3) Na and P
(4) Ca and Br
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20. Chlorine-37 can be represented as
13. Which electron configuration represents an
atom of aluminum in an excited state?
(1) 2-7-4
(2) 2-7-7
(3) 2-8-6
(4) 2-8-3
14. Which two notations represent different
isotopes of the same element?
15. Which atom in the ground state has a partially
filled second electron shell?
(1) hydrogen atom
(2) lithium atom
(3) potassium atom
(4) sodium atom
16. Which quantity identifies an element?
(1) atomic number
(2) mass number
(3) total number of neutrons in an atom of the
element
(4) total number of valence electrons in an
atom of the element
17. Which statement describes a chemical
property of the element magnesium?
(1) magnesium reacts with an acid
(2) magnesium has a high boiling point
(3) magnesium is malleable
(4) magnesium conducts electricity
18. Matter that is composed of two or more
different elements chemically combined in a
fixed proportion is classified as
(1) a mixture
(3) a compound
(2) an isotope
(4) a solution
19. An atom in the ground state contains a total of
5 electrons, 5 neutrons, and 5 protons. Which
Lewis electron-dot diagram represents this
atom?
21. The atomic mass of an element is calculated
using the
(1) atomic number and the ratios of its
naturally occurring isotopes
(2) masses and the half-lives of each of its
isotopes
(3) atomic number and half-lives of each of its
isotopes
(4) masses and the ratios of its naturally
occurring isotopes
22. Which of these terms refers to matter that
could be heterogeneous?
(1) element
(2) mixture
(3) solution
(4) compound
23. Which substance can be decomposed by
chemical means?
(1) ammonia
(2) phosphorous
(3) silicon
(4) oxygen
24. What is represented by the dots in a Lewis
electron-dot diagram of an atom of an element
in Period 2 of the Periodic Table?
(1) the number of neutrons in the atom
(2) the number of valence electrons in the
atom
(3) the total number of electrons in the atom
(4) the number of protons in the atom
25. What is the total number of neutrons in an
atom of
(1) 26
(2) 31
(3) 57
(4) 83
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26.
a) Determine the total number of neutrons in an atom of Si-29.
b) In the space below, show a correct numerical setup for calculating the atomic mass of Si.
c) A scientist calculated the percent natural abundance of Si-30 in a sample to be 3.29%. Determine
the percent error for this value.
27.
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a) Identify one piece of information shown in the electron-shell diagrams that is not shown in the
Lewis electron-dot diagrams.
b) Determine the mass number of the magnesium atom represented by the electron-shell diagram.
28. Write an electron configuration for an atom of aluminum-27 in the excited state.
29. What is the total number of neutrons in an atom of aluminum-27?
30. Draw the Lewis electron-dot diagram for Selenium in ground state.
31. Explain, in terms of subatomic particles, why K-37 and K-42 are isotopes of potassium.
11
UNIT 9: NUCLEAR CHEMISTRY (this part is due end of March)
Key Points:
 There are three main types of nuclear
radiation: alpha, beta, and gamma
radiation
 Spontaneous decay is when a radioactive
isotope decays by releasing a radioactive
particle
 Spontaneous decay reactions can be
represented using a chemical equation
 Each radioactive isotope has a specific
mode and rate of decay
 Half-life is the amount of time it takes to
reduce the amount of a radioactive isotope
by half
 There are natural and artificial
transmutation reactions
 Fusion and fission reactions are types of
artificial transmutation reactions that
produce tons of energy
 Radioisotopes are commonly used for
various procedures including medical
procedures and radioactive dating.
Questions:
1. Fill in the following table:
Particle
Symbol
Mass
Charge
Alpha
Beta
Positron
Gamma
2. What causes certain isotopes to spontaneously
decay?
3. Write a decay equation for C-14.
4. What is half-life?
5. What is the difference between fusion and fission? How do you identify them?
6. How is the difference between natural transmutation and artificial transmutation? How do you
identify them?
7. How do nuclear reactions differ from chemical reactions?
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UNIT 4 REGENTS QUESTIONS:
1. Which element has chemical properties that
are most similar to those of calcium?
(1) Co
(2) Sr
(3) K
(4) N
2. What was concluded about the structure of the
atom as a result of the gold foil experiment?
(1) A negatively charged nucleus is
surrounded by positively charged particles.
(2) A negatively charged nucleus is
surrounded by mostly empty space.
(3) A positively charged nucleus is surrounded
by positively charged particles.
(4) A positively charged nucleus is surrounded
by mostly empty space.
3. Which particles are found outside of an atom?
(1) electrons, only
(2) neutrons, only
(3) protons and electrons
(4) protons and neutrons
9. Which type of reaction converts one element
to another element?
(1) neutralization
(2) polymerization
(3) substitution
(4) transmutation
10. Which is an electron configuration for an atom
of phosphorous in the excited state?
(1) 2-8-5
(2) 2-8-8
(3) 2-8-4-1
(4) 2-8-7-1
11. What is the half-life and decay mode of Au198?
(1) 2.69 days and alpha decay
(2) 2.69 days and beta decay
(3) 3.82 days and alpha decay
(4) 3.82 days and beta decay
12. Which equation represents a transmutation
reaction?
4. What is the total number of valence electrons
in an atom of oxygen in the ground state?
(1) 2
(3) 6
(2) 4
(4) 7
5. Which nuclear emission has the greatest mass?
(1) 
(3) 
(2) +
(4) -
13. Which equation represents the radioactive
decay of
6. What is the decay mode of 37K?
(1) 
(3) 
+
(2) 
(4)
7. What is the mass number of an beta particle?
(1) 1
(3) 2
(2) 0
(4) 4
8. Which nuclear emission has the weakest
penetrating power?
(1) alpha particle
(2) beta particle
(3) gamma radiation
(4) positron
14. Which equation represents positron decay?
13
15. Which equation represents a fusion reaction?
16. The amount of energy released from a fission
reaction is much greater than the energy
released from a chemical reaction because in a
fission reaction
(1) mass is converted into energy
(2) energy is converted into mass
(3) ionic bonds are broken
(4) covalent bonds are broken
17. How many days are required for 200. grams of
radon-222 to decay to 50.0 grams?
(1) 1.91 days
(2) 3.82 days
(3) 7.64 days
(4) 11.5 days
18. Which fraction of an original 20.00-gram
sample of nitrogen-16 remains unchanged
after 36.0 seconds?
(1) 1/5
(2) 1/8
(3) 1/16
(4) 1/32
19. Which statement describes a chemical
property of hydrogen gas?
(1) Hydrogen gas burns in air.
(2) Hydrogen gas is colorless.
(3) Hydrogen gas has a density of 0.00009
g/cm3 at STP
(4) Hydrogen gas has a boiling point of 20. K at
standard pressure.
21. Which equation is an example of artificial
transmutation?
22. Which nuclide is paired with a specific use of
that nuclide?
(1) iodine-131, treatment of thyroid disorders
(2) uranium-238, dating of once-living
organisms
(3) carbon-14, dating of rock formations
(4) cobalt-60, treatment of cancer
23. One benefit of nuclear fission reactions is
(1) nuclear reactor meltdowns
(2) biological exposure
(3) production of energy
(4) storage of waste materials
24. Given the nuclear equation:
25.
20. The nucleus of a radium-226 atom is unstable,
which causes the nucleus to spontaneously
(1) decay
(2) oxidize
(3) absorb electrons
(4) absorb protons
14
15
Base your answers to questions part a-d on the information below, which relates the numbers of neutrons
and protons for specific nuclides of C, N, Ne, and S.
a) Using the point plotted on the graph for neon, complete the table below:
b) Explain, in terms of atomic particles, why S-32 is a stable nuclide.
c) Using the point plotted on the graph for nitrogen, what is the neutron-to-proton ratio of this
nuclide?
d) Complete the decay equation for N-16.
16
1. A U-238 atom decays to a Pb-206 atom through a series of steps. Each point on the graph below
represents a nuclide and each arrow represents a nuclear decay mode.
a) Based on this graph, what particle is emitted during the nuclear decay of a Po-218 atom?
b) Explain why the U-238 disintegration series ends with the nuclide Pb-206.
2. The fossilized remains of a plan were found at a construction site. The fossilized remains contain
1/16 the amount of carbon-14 that is present in a living plant.
a) Determine the approximate age of these fossilized remains.
b) Write out the nuclear equation for the decay of C-14.
17
3. The [Marie and Pierre] Curies set out to study radioactivity in 1898. Their first accomplishment was to
show that radioactivity was a property of atoms themselves. Scientifically, that was the most important
of their findings, because it helped other researchers refine their understanding of atomic structure.
More famous was their discovery of polonium and radium. Radium was the most radioactive
substance the Curies had encountered. Its radioactivity is due to the large size of the atom, which
makes the nucleus unstable and prone to decay, usually to radon and then lead, by emitting particles
and energy as it seeks a more stable configuration.
Marie Curie struggled to purify radium for medical uses, including early radiation treatment
for tumors. But radium’s bluish glow caught people’s fancy, and companies in the United States began
mining it and selling it as a novelty: for glow-in-the-dark light pulls, for instance, and bogus cure-all
patent medicines that actually killed people.
What makes radium so dangerous is that it forms chemical bonds in the same way as calcium,
and the body can mistake it for calcium and absorb it into the bones. Then, it can bombard cells with
radiation at close range, which may cause bone tumors or bone-morrow damage that can give rise to
anemia or leukemia.
-- Denise Grady, The New York Times, October 6, 1998
a) State one risk associated with the use of radium.
b) Write the decay reaction for radium-226
c) If a scientist purifies 1.0 gram of radium-226, how many years must pass before only 0.50 gram of
the original radium-226 sample remain unchanged?
4. Some radioisotopes used as tracers make it possible for doctors to see the images of internal body parts
and observe their functions. The table below lists information about three radioisotopes and the body
part each radioisotope is used to study.
a) Write the nuclear decay equation for the radioisotope used to study red blood cells.
b) It could take up to 60. hours for a radioisotope to be delivered to the hospital from the laboratory
where it is produced. What fraction of an original sample of Na-24 remains unchanged after 60.
hours.
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UNIT 4: THE PERIODIC
TABLE
Key Points:
 The periodic table is arranged in order of
increasing atomic number
 There are three main classification of
elements: metals, non-metals, and
metalloids
 A row of the periodic table is called a
period and a column is called a group
 Elements within the same group are
chemically similar to each other because
they have the same number of valence
electrons
 Atomic radius is the distance between the
nuclei of two atoms of the same elements
in the solid state
 Atomic radius increases as we go across a
period and decreases as we go down a
group
 Electronegativity is an atom’s attraction
for electrons in a chemical bond
 Ionization energy is the energy required to
remove a valence electron from an atom
 Electronegativity and ionization energy
increases as we move across a period
UNIT 5 REGENTS QUESTIONS:
1. Which subatomic particle has a negative
charge?
(1) electrons
(2) protons
(3) neutrons
(4) positrons
2. The elements in the Periodic Table are
arranged in order of increasing
(1) atomic radius
(2) atomic number
(3) mass number
(4) neutron number
3. Which list of elements from Group 2 on the
Periodic Table is arranged in order of
increasing atomic radius?
(1) Be, Mg, Ca
(2) Ca, Mg, Be
(3) Ba, Ra, Sr
(4) Sr, Ra, Ba
4. The elements located in the lower left corner
of the Periodic Table are classified as
(1) metals
(2) nonmetals
(3) metalloids
(4) noble gases
5. Germanium is classified as a
(1) Metal
(2) Metalloid
(3) Nonmetal
(4) Noble gas
6. Which list of elements contains a metal, a
metalloid and a nonmetal.
(1) Zn, Ga, Ge
(2) Si, Ge, Sn
(3) Cd, Sb, I
(4) F, Cl, Br
7. Which element has chemical properties that
are most similar to the chemical properties of
sodium?
(1) Mg
(2) K
(3) Se
(4) Cl
8. From which of these atoms in the ground state
can a valence electron be removed using the
least amount of energy?
(1) nitrogen
(2) oxygen
(3) carbon
(4) chlorine
9. Which of these elements has the lowest boiling
point?
(1) Li
(2) K
(3) Na
(4) Rb
10. Which element has both metallic and nonmetallic properties?
(1) Rb
(2) Rn
(3) Si
19
(4) Sr
11. Atoms of which element have the greatest
tendency to gain electrons?
(1) bromine
(2) chlorine
(3) fluorine
(4) iodine
12. Element X is a solid that is brittle, lacks luster,
and has six valence electrons. In which group
on the Periodic Table would Element X be
found?
(1) 1
(2) 2
(3) 15
(4) 16
13. Which statement describes a chemical
property of the element magnesium?
(1) Magnesium is malleable.
(2) Magnesium conducts electricity.
(3) Magnesium reacts with an acid.
(4) Magnesium has a high boiling point.
14. Which statement best describes Group 2
elements as they are considered in order from
top to bottom of the Periodic Table?
(1) The number of principal energy levels
increases, the number of valence electrons
increases.
(2) The number of principal energy levels
increases, the number of valence electrons
remains the same.
(3) The number of principal energy levels
remains the same, and the number of
valence electrons increases.
(4) The number of principal energy levels
remains the same, and the number of
valence electrons remains the same.
(4) low ionization energy and good electrical
conductivity
16. The data table below shows elements Xx, Yy,
and Zz from the same group on the Periodic
Table.
Atomic Mass
Atomic
Element
(atomic mass
Radius
unit)
(pm)
Xx
Yy
69.7
114.8
141
?
Zz
204.4
171
What is the most likely atomic radius of
element Yy?
(1) 103 pm
(2) 127 pm
(3) 166 pm
(4) 185 pm
17. Which term indicates how strongly an atom
attracts the electrons in a chemical bond?
(1) alkalinity
(2) atomic mass
(3) electronegativity
(4) activation energy
15. What are two properties of most nonmetals?
(1) high ionization energy and poor electrical
conductivity
(2) high ionization energy and good electrical
conductivity
(3) low ionization energy and poor electrical
conductivity
20
Base your answers to questions 18-20 on the Reference Tables for Chemistry
18. Complete the data table provided below for the following Group 18 elements: He, Ne, Ar, Kr, Xe
Atomic
Number
Element
2
10
18
36
54
He
Ne
Ar
Kr
Xe
First Ionization Energy
(kJ/mol)
19. Using information from your data table in question 7, construct a line graph following the directions
below. Circle each point and connect the points.
20. Based on your graph in question 7, describe the trend in first ionization energy of Group 18 elements
as the atomic number increases.
21
UNIT 1/2: CHEMICAL BONDING
Key Points:
 Noble gases are stable because they have eight valence electrons
 Atoms will gain or lose electrons to achieve a full octet
 Full octet = stability
 Ions are charged atoms
 Ions form by gaining or losing electrons to achieve a full octet
 Ions are positively charged if the atom loses electrons and negatively charged if the atom gains
electrons
 Polyatomic ions are listed in Table E and are ions that are made up of two or more elements
 Ionic bonds form between metals and non-metals
 An ionic bond require the transfer of electrons between two atoms so that both atoms can achieve a
full octet
 The criss-cross rule determines the molecular formula of an ionic salt
 The cation is always named first by the name of the element
 The anion is named second by dropping the ending of the element of the element and replacing it
with –ide
 Covalent bonding is the sharing of electrons between two non-metals so that each atom fulfills the
octet rule
 Lewis electron dot diagrams can be used to show the bonding of atoms in a covalently bonded
molecule
 Numerical prefixes must be used to indicate the number of atoms of each element present
 Polarity is the difference in electronegativity between the atoms that make up a compound
 Polarity is the unequal sharing of electrons
 Non-polar molecules are either symmetrical or are diatomic
 Molecules with a electronegativity difference of 0.6 are usually polar; less than 0.6 are usually nonpolar
 Intermolecular forces refer to forces that attract one molecule to another molecular
 Dipole-dipole interactions occur between polar molecules
 Hydrogen bonding occurs between a partially positive hydrogen and a partially negative atom like O,
N, or F
Questions:
1. What is the difference between cations and anions?
2. How do you know if you have a polyatomic ion?
3. What are the three types of chemical bonds?
4. Explain, in terms of electrons, how ionic bonds are different from covalent bonds?
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5. For which type of bonding do you use the criss-cross rule?
6. For which type of bonding do you use numerical prefixes?
7. The roman numeral (ex: Iron (III) oxide) in a compound name is used when the metal is what kind of
metal?
8. Write the chemical formula for Lithium oxide.
9. What is the name of CO2?
10. List three polar molecules and three nonpolar molecules.
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UNIT 6 REGENTS QUESTIONS
1. Which statement best describes electrons?
(1) They are positive subatomic particles and
are found in the nucleus.
(2) They are positive subatomic particles and
are found surrounding the nucleus.
(3) They are negative subatomic particles
and are found in the nucleus
(4) They are negative subatomic particles
and are found surrounding the nucleus.
2. Which Group of the Periodic Table contains
atoms with a stable outer electron
configuration?
(1) 1
(3) 8
(2) 16
(4) 18
3. An atom of carbon-12 and an atom of carbon14 differ in
(1) atomic number
(2) mass number
(3) nuclear charge
(4) number of electrons
4. What is the chemical formula for iron (III)
oxide?
(1) FeO
(3)Fe2O3
(2) Fe3O
(4) Fe3O2
5. Given a formula for oxygen:
What is the total number of electrons shared
between the atoms represented in this
formula?
(1) 1
(3) 2
(2) 8
(4) 4
6. Which formulas represent two polar
molecules?
(1) CO2 and HCl
(2) CO2 and CH4
(3) H2O and HCl
(4) H2O and CH4
7. Magnesium nitrate contains chemical bonds
that are
(1) covalent, only
(2) ionic, only
(3) both covalent and ionic
(4) neither covalent nor ionic
8. When sodium and fluorine combine to
produce the compound NaF, the ions formed
have the same electron configuration as
atoms of
(1) argon, only
(2) neon, only
(3) both argon and neon
(4) neither argon nor neon
9. Which formula represents a nonpolar
molecule?
(1) H2S
(2) HCl
(3) CH4
(4) NH3
10. What occurs when an atom loses an electron?
(1) The atom’s radius decreases and the atom
becomes a negative ion.
(2) The atom’s radius decreases and the atom
becomes a positive ion.
(3) The atom’s radius increases and the atom
becomes a negative ion.
(4) The atom’s radius increases and the atom
becomes a positive ion.
11. Which list of elements that have the most
similar chemical properties?
(1) Mg, Al, and Si
(2) Mg, Ca, and Ba
(3) K, Al, and Ni
(4) K, Ca, and Ga
12. Radioactive cobalt-60 is used in radiation
therapy treatment. Cobalt-60 undergoes beta
decay. This type of nuclear reaction is called
(1) natural transmutation
(2) artificial transmutation
(3) nuclear fusion
(4) nuclear fission
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13. Which Lewis electron-dot diagram is correct
for S2- ion?
19. Which electron configuration represents an
atom of aluminum in an excited state?
(1) 2-7-4
(2) 2-7-7
(3) 2-8-3
(4) 2-8-6
20. Which formula represents a hydronium ion?
(1) H3O+
(2) NH4+
(3) OH(4) HCO3-
14. Which formula represents an ionic
compound?
(1) H2
(2) CH4
(3) CH3OH
(4) NH4Cl
15. An ion of which element has a larger radius
than an atom of the same element?
(1) aluminum
(2) chlorine
(3) magnesium
(4) sodium
16. What is the total number of different
elements present in NH4NO3?
(1) 7
(2) 9
(3) 3
(4) 4
17. The atomic mass of an element is the
weighted average of the masses of
(1) its two most abundant isotopes
(2) its two least abundant isotopes
(3) all of its naturally occurring isotopes
(4) all of its radioactive isotopes
18. What is the chemical formula for sodium
sulfate?
(1) Na2SO3
(2) Na2SO4
(3) NaSO3
(4) NaSO4
21. Which group on the Periodic Table of the
Elements contains elements that react with
oxygen to form compounds with the general
formula X2O?
(1) Group 1
(2) Group 2
(3) Group 14
(4) Group 18
22. Given the formula of a substance:
What is the total number of shared electrons
in a molecule of this substance?
(1) 22
(2) 11
(3) 9
(4) 6
23. Which polyatomic ion contains the greatest
number of oxygen atoms?
(1) acetate
(2) carbonate
(3) hydroxide
(4) peroxide
24. Which isotopic notation represents an atom
of carbon-14?
25
25. Which compound has hydrogen bonding
between its molecules?
(1) CH4
(2) CaH2
(3) KH
(4) NH3
26. Which symbol represents a particle with a
total of 10 electrons?
(1) N
(2) N3+
(3) Al
(4) Al3+
27. Element X reacts with iron to form two
different compounds with the formulas FeX
and Fe2X3. To which group on the Periodic
Table does element X belong?
(1) Group 8
(2) Group 2
(3) Group 13
(4) Group 16
28. Which type of bond is found in sodium
bromide?
(1) covalent
(2) hydrogen
(3) ionic
(4) metal
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29. Draw a Lewis electron-dot diagram of Cl2.
30. Draw a Lewis electron-dot diagram of phosphorous trichloride, PCl3.
31. Explain, in terms of electronegativity, why a P-Cl bond in a molecule of PCl5 is more polar than a P-S
bond in a molecule of P2S5.
32. Explain, in terms of molecular structure or distribution of charge, why methane is nonpolar.
33. Write the IUPAC name for Fe2O3.
34. Using the key below, draw at least two water molecules in the box, showing the correct orientation of
each water molecule when it is near the Cl – ion in the aqueous solution.
27
UNIT 3: CHEMICAL REACTIONS
Key Points:
 The two kinds of chemical formulas:
empirical formulas and molecular formulas
 Empirical formula is the simplest wholenumber ratio of atoms of the elements in a
compound
 The molecular formula is the actual ratio of
atoms of the elements that combine to form a
compound
 Gram formula mass is the calculated mass of
a compound per 1 mole of that compound
 The molecular formula is found by dividing
the gram formula mass of the unknown by
the gram formula mass of the given empirical
formula to produce an integer value
 The integer value will be the number we
multiply the subscripts of the empirical
formula by to obtain the molecular formula
 The percent composition tells us what
percent of a compound is made up of a
certain element based on its mass
 Percent composition is calculated by finding
the gram-formula mass of an element and
then dividing it by the gram formula mass of
the compound
 In all chemical rxns, there is a conservation of
mass, energy, and charge
 The starting material is called the reactant
and is on the left side of the arrow
 The final product is called the product and is
on the right side of the arrow
 A coefficient tells us the number of molecules
we have
 When the law of conservation of mass is
obeyed, then the equation is balanced
 We can change the coefficients to make sure
all chemical reactions are balanced
 Table J lists elements from reactive to less
reactive
 Elements listed higher on Table J are more
reactive and will replace a lower placed
element
 A mole is a measurement of quantity
 We find mass from moles by multiplying the
number of moles by the atomic mass
Questions:
1. What is the difference between molecular
and empirical formula?
2. How do you calculate for the gram-formula
mass of a compound?
3. How do you solve for the percent
composition of an element in a compound?
4. How do you perform a:
a) Mole-to-mole conversion?
b) Mole-to-gram conversion?
c) Gram-to-mole conversion?
5. What are the types of chemical reactions?
6. What is table J used for? What type of reaction does table J work for?
28
UNIT 7 REGENTS QUESTIONS
1. Which phrase best describes an atom?
(1) a positive nucleus surrounded by a hard
negative shell
(2) a positive nucleus surrounded by a cloud
of negative charges
(3) a hard sphere with positive particles
uniformly embedded
(4) a hard sphere with negative particles
uniformly embedded
2. All chemical reactions have a conservation of
(1) mass, only
(2) mass and charge, only
(3) charge and energy, only
(4) mass, charge, and energy
3. Which two particles each have a mass
approximately equal to one atomic mass
unit?
(1) electron and neutron
(2) electron and positron
(3) proton and electron
(4) proton and neutron
4. Which equation shows conservation of
atoms?
(1) H2 + O2  H2O
(2) H2 + O2  2H2O
(3) 2H2 + O2  2H2O
(4) 2H2 + 2O2  2H2O
5. Compared to an electron in the first electron
shell of an atom, an electron in the third shell
of the same atom has
(1) less mass
(2) less energy
(3) more mass
(4) more energy
6. What is the gram-formula mass of Ca3(PO4)2?
(1) 248 g/mol
(2) 263 g/mol
(3) 279 g/mol
(4) 310 g/mol
7. What is the total number of neutrons in an
atom of an element that has a mass number
of 19 and an atomic number of 9?
(1) 9
(2) 10
(3) 19
(4) 28
8. Given the balanced equation:
2 C + 3 H2  C2H6
What is the total number of moles of C that
must completely react to produce 2.0 moles
of C2H6?
(1) 1.0 mol
(2) 2.0 mol
(3) 3.0 mol
(4) 4.0 mol
9. Which is an empirical formula?
(1) P2O5
(2) P4O6
(3) C2H4
(4) C3H6
10. Which metal reacts spontaneously with a
solution containing zinc ions?
(1) magnesium
(2) nickel
(3) copper
(4) silver
11. Which pair consists of a molecular formula
and its corresponding empirical formula?
(1) C2H2 and CH3CH3
(2) C6H6 and C2H2
(3) P4O10 and P2O5
(4) SO2 and SO3
12. Which equation represents a double
replacement reaction?
(1) 2 Na + 2 H2O  2 NaOH + H2
(2) CaCO3  CaO + CO2
(3) LiOH + HCl  LiCl + H2O
(4) CH4 + 2 O2  CO2 + 2 H2O
29
13. Given the balanced equation representing the
reaction between propane and oxygen:
C3H8 + 5 O2  3 CO2 + 4 H2O
According to this equation, which ratio of
oxygen to propane is correct?
14. What type of change must occur to form a
compound?
(1) chemical
(2) physical
(3) nuclear
(4) phase
15. What is the percent composition by mass of
aluminum in Al2(SO4)3 (gram formula mass =
342 grams/mole)?
(1) 7.89%
(2) 15.8 %
(3) 20.8%
(4) 36.0 %
16. Which equation shows a conservation of
mass?
(1) Na + Cl2  NaCl
(2) Al + Br2  AlBr3
(3) H2O  H2 + O2
(4) PCl5  PCl3 + Cl2
18. A compound has a molar mass of 90. grams
per mole and the empirical formula CH2O.
What is the molecular formula of this
compound?
(1) CH2O
(2) C2H4O2
(3) C3H6O3
(4) C4H8O4
19. Which formula represents lead (II)
chromate?
(1) PbCrO4
(2) Pb(CrO4)2
(3) Pb2CrO4
(4) Pb2(CrO4)3
20. Given the balanced equation:
2 C4H10 (g) + 13 O2 (g)  8 CO2 (g) + 10 H2O
(g)
What is the total number of moles of O2 (g)
that must react completely with 5.00 moles of
C4H10(g)?
(1) 10.0
(2) 20.0
(3) 26.5
(4) 32.5
17. Given the balanced equation representing a
reaction:
2 CO (g) + O2 (g)  2 CO2 (g)
What is the mole ratio of CO (g) to CO2 (g) in
this reaction?
(1) 1:1
(2) 1:2
(3) 2:1
(4) 3:2
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21. “Hand Blasters” is a toy that consists of a set of two ceramic balls, each coated with a mixture of
sulfur and potassium chlorate, KClO3. When the two balls are struck together, a loud popping noise is
produced as sulfur and potassium chlorate react with each other.
a) Balance the equation for the “Hand Blaster” reaction, using the smallest whole-number
coefficients.
________S(s) + ________ KClO3(s)  _______SO2(g) + _______KCl(s) + energy
22. During a laboratory activity, a student reacted a piece of zine with 0.1 M HCl (aq).
a) Complete the equation below by writing the formula of the missing product.
Zn (s) + 2 HCl (aq)  ________ (aq) + H2 (g)
b) What type of reaction is this?
c) Based on Reference Table J, identify one metal that not spontaneously react with HCl.
23. The particle diagrams below represent the reaction between two nonmetals, A2 and Q2.
a) Using the symbols A and Q, write the chemical formula of the product.
b) Identify the type of chemical bond between an atom of element A and an atom of element Q.
c) Compare the total mass of the reactants to the total mass of the product.
31
24. What is the total number of moles in 80.0 grams of C2H5Cl (gram-formula mass= 64.5 grams/mole)?
25. What is the mass of 4.76 moles of Na3PO4 (gram-formula mass = 164 grams/ mole)?
26. Air bags are an important safety feature in modern automobiles. An air bag is inflated in milliseconds
by the explosive decomposition of NaN3 (s). The decomposition reaction produces N2(g), as well as
Nas), according to the unbalanced equation below.
NaN3(s)  Na(s) + N2 (g)
a) Write the balanced equation below.
b) How many moles of N2 is produced with 20 grams of NaN3 is reacted?
27.
Archimedes (287-212 BC), a Greek inventor and mathematician, made several discoveres
important to science today. According to a legend, Hiero, the king of Syracuse, commanded
Archimedes to find out if the royal crown was made of gold, only. The king suspected that the crown
consisted of a mixture of gold, tin, and copper.
Archimedes measured the mass of the crown and the total amount of water displaced by the
crown when it was completely submerged. He repeated the procedure using individual samples, one
of gold, one of tin, and one of copper. Archimedes was able to determine that the crown was not
made entirely of gold without damaging it.
a) Identify one physical property that Archimedes used in his comparison of the metal samples.
b) Determine the volume of a 75-gram sample of gold at STP.
32
28. Base your answer to 28a and 28b on the balanced equation below.
Fe(s) + 2HNO3(aq)  Fe(NO3)2(aq) + H2(g)
a) What is the total number of oxygen atoms represented in the formula of the iron compound
produced?
b) Explain, using information from Reference Table J, why this reaction is spontaneous.
And of course, because you should always review your nuclear chemistry…
29. In living organisms, the ratio of the naturally occurring isotopes of carbon, C-12 to C-13 to C-14, is
fairly consistent. When an organism such as a wooly mammoth died, it stopped taking in carbon. And
the amount of C-14 present in the mammoth began to decrease. For example, one fossil of a woolly
mammoth is found to have 1/32 of the amount of C-14 found in a living organism.
a) Identify the type of nuclear reaction that caused the amount of C-14 in the wooly mammoth to
decrease after the organism died.
b) Determine the total time that has elapsed since this woolly mammoth died.
c) State, in terms of subatomic particles, how an atom of C-14 differs from an atom of C-12
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PRACTICE REGENTS
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Please answer questions in the answer book provided on page 44.
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Answer Book for Constructed Response Questions
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