Download Chemistry Content Review Notes

STANDARDS OF LEARNING
CONTENT REVIEW NOTES
CHEMISTRY
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
OVERVIEW
Chemistry Content Review Notes are designed by the High School Science Steering Committee as a
resource for students and parents. Each nine weeks’ Standards of Learning (SOLs) have been
identified and a detailed explanation of the specific SOL is provided. Specific notes have also been
included in this document to assist students in understanding the concepts. A “
” section has
also been developed to provide students with the opportunity to check their understanding of the
content.
The document is a compilation of information found in the Virginia Department of Education
(VDOE) Curriculum Framework, Enhanced Scope and Sequence, and Released Test items. In
addition to VDOE information, Glencoe Textbook Series and resources have been used. Finally,
information from various websites is included. The websites are listed with the information as it
appears in the document.
The Chemistry Blueprint Summary Table is listed below as a snapshot of the reporting categories,
the number of questions per reporting category, and the corresponding SOLs.
It is the Chemistry Instructors’ desire that students and parents will use this document as a tool
toward the students’ success on the end-of-year assessment.
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
Standards of Learning
Notes
&
Activities
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
First Nine Weeks
LAB SAFETY EQUIPMENT
ATOMIC STRUCTURE
SOL 1C
SOL 2I
SOL 2A
SOL 2C
SOL 2B
SCIENTIFIC METHOD
SOL- 1D, E
REPORTING SCIENTIFIC
DATA
SOL- 1E, G
SOL 1G
SOL- 1A,E,F
SOL- 1F
MATTER AND ENERGY
SOL 2H
PERIODIC TABLE
SOL 3 A
SOL- 2D,E
SOL 2A
SOL 2F
SOL 2C
ELECTRON CONFIGURATION
SOL 2I
SOL-2 A, B,C, E,G
SOL-2 A, B, C, D, G
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
TOPIC: LAB SAFETY EQUIPMENT
CHEMISTRY
STANDARD CH. 1C
The student will investigate and understand that experiments in which
variables are measured, analyzed, and evaluated produce observations and
verifiable data. Key concepts include: c) proper response to emergency
situations
1. What is the correct procedure to neutralize an acid?
2. When diluting an acid solution, one should always do what and why?
3. Describe the sequence of actions you would take if you spill a strong acid on your
clothing in the lab.
4. Describe what you would do if your lab partner’s shirt-sleeve caught on fire.
5. When is it most important to wash your hands while in the laboratory?
6. Describe the procedure to smell a substance?
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
TOPIC: SCIENTIFIC METHOD
CHEMISTRY
STANDARD CH. 1D, E
The student will investigate and understand that experiments in which
variables are measured, analyzed, and evaluated produce observations
and verifiable data. Key concepts include:
d)
manipulation of multiple variables, using repeated trials; and
e)
accurate recording, organization, and analysis of data through
repeated trials.
Student A
Student B
Student C
Student D
Student measurement of Temperature in ° C
Reading 1
Reading 2
88.6
88.5
92.6
90.1
90.0
88.9
80.1
89.8
Reading 3
88.7
91.6
92.5
90.0
1. Use the above chart to answer the following question. Four students each took
temperature readings of a sample of water. The actual temperature of the sample was
90.0 ° C. Which student’s measurements were both accurate and precise?
2. A student was asked to determine the density of a mineral. Outline the steps in which to
accurately determine the density using the displacement method.
3. What is the first step that should be taken when a caustic chemical gets into a person’s
eye?
4. What is the mass of a substance whose density is 0.2396 g/mL and a volume of 20.0 ml?
5. What is the minimum number of trials needed for a measurement to be reliable?
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
TOPIC: REPORTING SCIENTIFIC METHOD
CHEMISTRY
STANDARD CH. 1E,G
The student will investigate and understand that experiments in which variables are
measured, analyzed, and evaluated produce observations and verifiable data. Key
concepts include:
e)
accurate recording, organization, and analysis of data through repeated trials.
g)
mathematical manipulations (SI units, scientific notation, linear equations,
graphing, ratio and proportion, significant digits, dimensional analysis).
1.
2.
3.
4.
5.
6.
Determine the number of significant figures in the following:
a. 5006
b. 1200
c. 0.001200
d. 5
Compute the following. Record your answer to the correct number of significant figures.
a. (4.33 x 103)(1.6 x 10 -2)
b. 2.36 x 102 x 4.2655 x 104
c. 9.65 x 105/7.2 x 10-3
d. 2.33 x 10-3 x 1.05 x 102/1.234
The SI base unit for mass is the ______________________.
The SI base unit for the volume of a liquid is the ________________.
Use the factor label method or dimensional analysis to solve the following problems.
a. Convert 950 m to km
b. Convert 2.56 kg to g
c. Convert 0.7526 mg to kg
d. Convert 4.85 L to mL
Convert the following temperatures.
a. 543K to Celsius
b. 23K to Celsius
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 1 A, E, F
The student will investigate and understand that experiments in which variables are
measured, analyzed, and evaluated produce observations and verifiable data. Key
concepts include:
a)
designated laboratory techniques
e)
accurate recording, organization, and analysis of data through repeated trials.
f)
mathematical and procedural error analysis
1. What piece of equipment would serve the purpose for each of the following?
a. Adding 5 ml of 5% HCl to 5 different solutions. ______________
b. To measure 97.0 ml of water. _______________
c. To mass a solid substance. ________________
d. To dry out a precipitate for measuring. _______________
2. Which of the following items should be used if you need approximately 100 ml of a
liquid? (Choose one)
a. Beaker
b. Erlenmeyer flask
c. Graduated cylinder
3. A student measures the mass of a piece of copper three times and records the results in
the following table:
Trial
Mass (grams)
1
15.5
2
15.6
3
15.7
The actual mass of the metal is 17.2 grams. Are the trial measurements accurate and/or precise?
Explain your answer.
4. A student measured the temperature of a solution and found it to be 46.0◦C at standard
pressure. The theoretical temperature of that boiling solution is 45.0◦C. What is the
percent error according to the student’s measurement?
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 1G
The student will investigate and understand that experiments in which
variables are measured, analyzed, and evaluated produce observations and
verifiable data. Key concepts include:
g) mathematical manipulations including SI units, scientific notation, linear
equations, graphing, ratio and proportions, significant digits and dimensional
analysis.
DIRECTIONS: Show all work using dimensional analysis
1. Convert 3.5 x 10-3 L to mL.
2. How many kilograms are there in 4223 g?
3. Using the rules for significant figures, what is the correct answer for 2.5 m times 44.8 m?
4. (3.5 x 106)(5.3 x 10-5) =
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 1 F
The student will investigate and understand that experiments in which
variables are measured, analyzed, and evaluated produce observations
and verifiable data. Key concepts include
f)
mathematical and procedural error analysis
DIRECTIONS: Determine the percent error in the following problems.
1. Experimental Value = 1.44 g
Accepted Value = 1.54 g
2. Experimental Value = 22.2 L
Accepted Value = 22.4 L
3. Experimental Value = 128.6 mg
Accepted Value =128.3 mg
4. Experimental Value =
9.98 x 10-3 g
Accepted Value =1.03 x 10-2 g
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
TOPIC: MATTER AND ENERGY
CHEMISTRY
STANDARD CH. 2H
The student will investigate and understand that the placement of elements on the
periodic table is a function of their atomic structure. The periodic table is a tool
used for the investigations of:
h)
chemical and physical properties.
Chemical and Physical Properties
The student will be given an object or substance and must identify the physical and chemical
properties of that object. (This assignment is to be given after lecture on chemical and physical
properties and how to correctly identify each.)
1.
Describe the chemical properties of this object below, citing any evidence you have
collected:
2.
Describe the physical properties of this object below, citing any evidence you have
collected:
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
TOPIC: ATOMIC STRUCTURE
CHEMISTRY
STANDARD CH. 2I
The student will investigate and understand that the placement of elements on the periodic
table is a function of their atomic structure. The periodic table is a tool used for the
investigations of:
i)
historical and quantum models.
Students will create a two page foldable for the models of an atom. Each outside flap will be the
names of the scientists (Democritus, Dalton, Rutherford, Bohr), and the inside picture will be
each scientist’s contribution to the historical model. A one page foldable will be created for the
quantum models with Heisenberg and Planck. Students will be creative using mnemonics and/or
graphics.
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 2A
The student will investigate and understand that the placement of elements on the
periodic table is a function of their atomic structure. The periodic table is a tool used
for the investigations of
a)
average atomic mass, mass number, and atomic number
Element
Ne
Al
Ge
P
Br
Ca
K
I
Sr
Sn
Average Atomic mass
Mass number
Atomic number
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 2C
The student will investigate and understand that the placement of
elements
on the periodic table is a function of their atomic structure.
.
The periodic table is a tool used for the investigations of:
c)
mass and charge characteristics of subatomic particles
Warm-up: Complete the following chart:
Subatomic Particle
Symbol
Charge
Relative mass
Location in atom
protons
neutrons
electrons
Atomic Structure
Element/Ion
3
1
Mass
Number
Protons
Neutrons
Electrons
H+
C
Li+
35
Cl-
39
24
Atomic
Mass
H
12
7
Atomic
Number
K
Mg2+
77
As3-
108
Ag
30 -2
S
238
U
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 2B
The student will investigate and understand that the placement of elements
on the periodic table is a function of their atomic structure. The periodic
table is a tool used for the investigations of:
b)
isotopes, half lives, and radioactive decay
DIRECTIONS: Answer the following questions
Average Atomic Mass
1. Rubidium has two common isotopes, Rb-85 and Rb-87. If the abundance of Rb-85 is
72.2% and the abundance of Rb-87 is 27.8%, what is the average atomic mass of
rubidium?
2. Uranium has 3 common isotopes. If the abundance of U-234 is 0.01%, the
abundance of U-235 is 0.71% and the abundance of U-238 is 99.28%, what is the
average atomic mass?
3. Titanium has five common isotopes: Ti-46 (8.0%), Ti-47 (7.8%), Ti-48 (73.4%), Ti-49
(5.5%), and Ti-50 (5.3%). What is the average atomic mass of titanium?
4. Explain why atoms have different isotopes. In other words, how is it that helium can
have three different sized atoms and they all are still the element Helium.
5. Draw and label 3 possible isotopes of hydrogen. Which isotope can you predict to be
the most abundant? Why? Correctly name the 3 isotopes using the symbol “H”?
6. Select an element on the Period Table. Create 4 isotopes for your element and form
an average atomic mass calculation to support the given atomic mass.
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
TOPIC: PERIODIC TABLE
CHEMISTRY
STANDARD CH. 3A
The student will investigate and understand how conservation of energy and
matter is expressed in chemical formulas and balanced equations. Key
concepts include
a)
nomenclature
Writing Binary Formulas
Cation
Anion
Na+
Formula
Cation
Anion
Cl-
Fe+2
O-2
Ba+2
F-
Fe+2
O-2
K+
S-2
Cr+2
S-2
Li+
Br-
Cr+3
S-2
Al+3
I-
Cu+
Cl-
Zn+2
S-2
Cu+2
Cl-
Ag+
O-2
Pb+2
O-2
Mg+2
P-3
Pb+4
O-2
Ni+2
O-2
Mn+2
Br-
Ni+3
O-2
Mn+4
Br-
Formula
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 2 D, E
The student will investigate and understand that the placement of elements on
the periodic table is a function of their atomic structure. The periodic table is a
tool used for the investigations of:
d)
families or groups;
e)
periods
Counting Valence Electrons
Element
Group Number
Number of Valence
Electrons
Ne
Al
Ge
P
Br
Ca
K
I
Sr
Sn
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 2 F
The student will investigate and understand that the placement of
elements on the periodic table is a function of their atomic structure.
The periodic table is a tool used for the investigations of:
f)
trends including atomic radii, electronegativity, shielding
effect, and ionization energy.
DIRECTIONS: State the periodic trend and explain.
PERIODIC TRENDS
LEFT TO
RIGHT
EXPLAINATION
TOP TO
BOTTOM
EXPLAINATION
ATOMIC RADII
ELECTRONEGATIVITY
SHIELDING EFFECT
IONIZATION ENERGY
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 2 C
The student will investigate and understand that the placement of
elements on the periodic table is a function of their atomic structure. The
periodic table is a tool used for the investigations of:
c)
mass and charge characteristics of subatomic particles.
DIRECTIONS: Complete the chart below.
Ions and Subatomic Particles
Ion Symbol
S
Protons
Electrons
Charge
2-
K1+
Ba2+
Fe3+
Fe2+
F1O2P3Sn4+
Sn2+
N3Br1Mg2+
Cu1+
Cu2+
U6+
Mn5+
Cl1Se2-
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
Second Nine Weeks
IONIC AND COVALENT BONDING AND COMPOUNDS
SOL- 3 A, C, D
MOLE/EMPIRICAL AND MOLECULAR FORMULA
SOL- 4A, B
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
TOPIC: IONIC COMPPOUNDS
CHEMISTRY
STANDARD CH. 3A,C, D
The student will investigate and understand how conservation of energy
and matter is expressed in chemical formulas and balanced equations.
Key concepts include
a) nomenclature
c) writing chemical formulas
d) bonding types
Multiple Choice
1. What is the ionic charge on the copper ion in the ionic compound that has the
formula CuSO4 ?
a.
b.
c.
d.
+1
-2
+3
+2
2. Which element when combined with Bromine would most likely form an ionic
compound?
a.
b.
c.
d.
potassium
carbon
phosphorus
iodine
3. An –ite or –ate ending on the name of a compound indicates that the compound
_________________.
a.
b.
c.
d.
is a binary ionic compound
is a binary molecular compound
contains a polyatomic ion
contains a polyatomic cation
4. In general, ionic compounds _____________.
a.
b.
c.
d.
are amorphous solids at room temperature
conduct electricity when in the solid state
conduct electricity when they are dissolved in water
all of these
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
5. Which of the following is the correct name for the compound MnF3?
a.
b.
c.
d.
Manganese fluoride (III)
Manganese (III) fluoride
Manganese (I) fluoride (III)
Manganese (III) fluoride (III)
6. According to the periodic table, Mg will most likely react with elements in which of
these groups?
a.
b.
c.
d.
1
3
17
18
7. Which of these compounds is most likely to contain an ionic bond?
a.
b.
c.
d.
H2
SO2
CH4
CaCl2
8. What is the name of NH4OH?
a.
b.
c.
d.
Ammonium hydroxide
Nitrogen oxygen hydride
Nitrogen hydroxide
Ammonium oxygen hydride
9. All of the following formulas are correct except:
a.
b.
c.
d.
NaOH
CaF2
H3PO4
AlSO4
10. In forming NaCl, energy is required to
a.
b.
c,
d.
change chlorine to a gas
add an electron to a chlorine atom
remove an electron from a sodium atom
bring together the sodium ions and chloride ions
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
Critical Thinking:
1. Why are there no rules for naming Group 18 ions?
2. Determine the ratios of cations to anions that are most likely in the formulas for ionic
substances of the following elements:
a. An alkali metal and a halogen
b. An alkaline earth metal and a halogen
c. An alkali metal and member of Group 16
d. An alkaline earth metal and a member of Group 16
3. Compound B has lower melting and boiling points than compound A does. At the same
temperature, compound B vaporizes faster and to a greater extent than compound A. If
only one of these compounds is ionic, which one would you expect it to be? Why?
TOPIC: COVALENT COMPOUNDS AND BONDING
CHEMISTRY
STANDARD CH. 3 A, C, D
The student will investigate and understand how conservation of energy and matter is
expressed in chemical formulas and balanced equations. Key concepts include
a) nomenclature
c) writing chemical formulas
d) bonding types
Multiple Choice
1. Which of the following is the name of the molecule PCl3?
a.
b.
c.
d.
Phosphorus trichloride
Phosphorus chloride
Potassium trichloride
Potassium chloride
2. The type of formula that shows the arrangements of atoms and bonds is calleda. Empirical
b. Chemical
c. Molecular
d. Structural
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
3. What shape does the molecule BF3 have?
a. Bent
b. Linear
c. Tetrahedral
d. Trigonal planar
4.
H
H
:Z–Z:
H
H
The figure above shows a compound containing hydrogen (H) and an unknown
element Z. To which group on the periodic table does element Z belong?
5. Bonding between two elements of equal electronegativity would bea. 100% covalent
b. Primarily ionic
c. 50% ionic
d. Metallic in character
6. The small attractive force between molecules of nitrogen gas, N2, is due toa.
b.
c.
d.
Hydrogen bonds
Van der Waals forces
Dipole-dipole attraction
Magnetism
7. Which of the following is a polar covalent bond?
a.
b.
c.
d.
NaCl
O2
Al
SO2
8. In chemical compounds, covalent bonds form whena.
b.
c.
d.
The electronegativity difference between two atoms is very large
Electrons are completely transferred between two metals
Pairs of electrons are shared between two nonmetal atoms
Two nonmetal ions are attracted to each other by opposite charges
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
9. The Lewis electron dot system represents electrons in thea.
b.
c.
d.
Outer energy level
Inner level
Middle level
Core level
10. An alien landed on Earth and created the periodic table shown. The astronaut was
trying to determine what type of bond would be present in several compounds. The
type of bond in a compound containing D and B would bea.
b.
c.
d.
A metallic bond
A nonmetallic bond
A covalent bond
An ionic bond
SOL Release Chemistry Test 2005
Critical Thinking
1.
Ionic compounds tend to have higher boiling points than covalent substances do. Both
ammonia, NH3, and methane, CH4, are covalent compounds, yet the boiling point of
ammonia is 130 C higher than that of methane. What might account for the large
difference?
2. Create a concept map using the following terms: valence electrons, non-polar, covalent
compounds, polar, dipoles, and Lewis structures.
3. Create a Venn Diagram to compare and contrast ionic and covalent bonding. Make sure
you have 5 similarities and 5 differences.
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
TOPIC: MOLE/EMPIRICAL AND MOLECULAR FORMULA
CHEMISTRY
STANDARD CH. 4 A, B
The student will investigate and understand that quantities in a
chemical reaction are based on molar relationships. Key concepts
include
a) Avogadro’s principle and molar volume
b) Stoichiometric relationships
Multiple Choice:
1. The molar mass (gram formula mass) for the compound sodium thiosulfate,
Na2S2O3 is
a. 71 grams
b. 153 grams
c. 158 grams
d. 254 grams
2. The empirical formula for C6H12 is ---a. C3H6
b. C2H4
c. CH3
d. CH2
3.
How many moles of copper are equivalent to 3.44 x 1023 atoms of copper?
a. 0.571 moles
b. 1.75 moles
c. 5.41 x 1021 moles
d. 5.71 x 1022 moles
4. What is the mass of one mole of CO2?
a.
b.
c.
d.
24 g
28 g
44 g
56 g
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
5. Which of the following shows the correct number of atoms of each element in the
formula Mg(NO3)2?
a.
b.
c.
d.
1 magnesium atom, 2 nitrogen atoms, and 6 oxygen atoms
1 magnesium atom, 2 nitrogen atoms, and 5 oxygen atoms
1 magnesium atom, 1 nitrogen atom, and 6 oxygen atoms
1 magnesium atom, 1 nitrogen atom, and 5 oxygen atoms
Practice:
1. Formula Mass: Determine the molar mass of the following. Show all work and units.
a. CuSO4
b. CaCO3
c. H3PO4
d. Al2(SO4)3
2. Determine the empirical formula for the following:
a. 80.0 g of carbon and 20.0 g of oxygen
b. 71.5% Ca and 28.5% O
3. Percent Composition. Determine the percent composition of the following
compounds. Show all work.
a. NO2
b. Al2(SO4)3
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
THIRD Nine Weeks
CHEMICAL REACTIONS
SOL 3E
STOICHIOMETRY
SOL- 1G, 3B, 4B
THERMOCHEMISTRY
SOL- 3E, 4B, 5D, 5E, 5F
CHEMICAL KINETICS
SOL- 3F
GAS LAWS
SOL- 4A, 5A, 5B, 5C
TOPIC: CHEMICAL REACTIONS
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 3e
The student will investigate and understand how conservation of energy and matter is
expressed in chemical formulas and balanced equations. Key concepts include:
e) reaction types
Directions: Determine the reaction type of each of the following reactions.
CO 
1.
Fe2O3 +
2.
Cr +
3.
Eu +
HF 
4.
NH4Cl
 HCl + NH3
5.
C12H22O11 + O2 
6.
Zn +
7.
SiO2 +
8.
Pb(NO3)2 +
9.
KClO3 
10.
C6H6 +
S8 
FeO +
CO2
Cr2S3
HCl 
C 
EuF3 +
CO2 +
ZnCl2 +
Si +
KCl +
H2O
H2
CO
H3AsO4 
O2 
H2
PbHAsO4 +
HNO3
O2
CO2 +
H2O
Directions: Complete the following table:
Reaction
Energy
Description
Exothermic
Endothermic
Sign of ΔH,
Positive or
Negative?
Temperature,
Increase or
Decrease?
Energy,
Product or
Reactant?
TOPIC: STOICHIOMETRY
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 1g
The student will investigate and understand that experiments in which
variables are measured, analyzed and evaluated produce observations and
verifiable data. Key concepts include:
g) mathematic manipulations (SI units, scientific notation, linear
equations, graphing, ratio and proportion, significant digits,
dimensional analysis)
1. Determine the number of significant figures in the following:
a. 602
b. 1200
c. 0.00345
d. 0.1040
e. 34.08
f. 3
g. 0.970
2. Record your answer to the correct number of sig. figs.
a. (7.502 x 102 )(5.43 x 104)
b. 9.01 x 106 / 1.22 x 105
3. Write the following numbers in scientific notation, or translate the numbers to
expanded notation.
c. 9.32 x 10-5
d. 7.68 x 107
a. 556,000,000,000
b. 0.00000751
4. Use conversion factors to solve the following: SHOW ALL WORK!!
a. Convert 250 m to km.
b. Convert 18.50 kg to g.
c. Convert 3.54 L to kL.
d. Convert 0.5420 kg to mg
5. Convert the following temperatures to Kelvin or to degree’s Celsius.
a. 100˚C
b. 45˚C
c. 273 K
d. 400 K
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY STANDARD
CH. 3B
The student will investigate and understand how conservation of
energy and matter is expressed in chemical formulas and balanced
equations. Key concepts include b) balancing chemical equations.
DIRECTIONS: Balance the following equations:
CO 
1. Fe2O3 +
2. Cr +
S8 
3. Eu +
HF 
FeO +
Cr2S3
EuF3 +
4. C12H22O11 + O2 
5. Zn +
6. SiO2 +
HCl 
C 
7. Pb(NO3)2 +
8. NaCl +
9. KClO3 
10.C6H6 +
CO2
H2
CO2 +
ZnCl2 +
Si +
Cl2 +
KCl +
O2 
H2
CO
H3AsO4 
H2O 
H2O
PbHAsO4 +
H2 +
HNO3
NaOH
O2
CO2 +
H2O
TOPIC: STOICHIOMETRY
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 4B
The student will investigate and understand that quantities in a chemical reaction
are based on molar relationships. Key concepts include: b) stoichiometric
relationships.
1. Fill in the blanks.
a. The amount of a substance equal to 6.02 x 1023 particles is called a
________________________.
b. 6.02 x 1023 is also known as __________________________________.
c. If you have 1 mole of copper, you have 6.02 x 1023 ________________________.
2. Molar (Formula) Mass. Determine the molar mass of the following and show all work
and units.
a. CuSO4
b. CaCO3
c. H3PO4
d. Al2(SO4)3
3. Mole Conversions. Solve the following problems. Show all work
a. How many grams are there in 1.55 x 1023 molecules of CO2?
b. How many atoms are there in 2.18 moles of nitrogen gas?
c. What is the volume, at STP, occupied by 4.30 moles of oxygen gas?
d. How many grams are there in 56.32 liters of carbon monoxide, at STP?
e. How many moles are in 8.67 x 1025 atoms of sulfur?
4. Percent Composition. Determine the percent composition of the following compounds.
Show all work.
a. NO2
b. Al2(SO4)3
c. What is the percent of Mg in MgCl2?
Solve the following problems and show all work.
5. Determine the empirical formula for 71.5% Ca and 28.5% O.
6. Determine the molecular formula for the empirical formula of C3H5O2, with a molar mass
of 146 g/mol.
7. 4Al + 3O2 → 2Al2O3 , How many moles of oxygen are required to react completely
with 0.84 mol of Al? How many moles of aluminum are needed to form 2.3 mol of
Al2O3?
8. The combustion of acetylene gas is represented by this equation: 2C2H2(g) + 5O2(g) →
4CO2(g) + 2H2O(g)
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
a. How many grams of oxygen are required to burn 13.0g of C2H2 ?
b. How many moles of C2H2 are needed to react completely with 98.0 g of water?
9. Calcium carbonate can be decomposed by heating. CaCO3(s) → CaO) (s) + CO2(g)
What is the percent yield of this reaction if 24.8 g of CaCO3 is heated to give 13.1 g of
CaO?
10. Copper reacts with sulfur to form copper (I) sulfide. 2Cu(s) + S(s) → Cu2S(s)
What is the limiting reactant when 80.0 g of Cu reacts with 25.0 g of S?
11. How many moles of CS2 form when 2.7 mol of C reacts? 5C + 2SO2 → CS2 +4CO
TOPIC: THERMOCHEMISTRY
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY STANDARD CH. 3E
The student will investigate and understand how conservation of energy and matter is
expressed in chemical formulas and balanced equations. Key concepts include:
e) reaction types
See Previous “Try It” in Chemical Reactions for CH.3E
CHEMISTRY
STANDARD CH. 4B
The student will investigate and understand that quantities in a chemical
reaction are based on molar relationships. Key concepts include:
b) stoichiometric relationships
See previous “Try It” in Stoichiometry for CH.4b
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 5E
The student will investigate and understand that the phases of matter
are explained by kinetic theory and forces of attraction between
particles. Key concepts include:
e) molar heats of fusion and vaporization
Solve the following problems. Show all work.
1. How much heat is released as a 75.0 g sample of ethanol gas at the boiling point condenses to
a liquid? Heat of vaporization is 879 J/g.
2. The heat of vaporization of water is 540. cal/g. How many calories would be needed to
convert 3 moles of water to vapor?
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 5F
The student will investigate and understand that the phases of
matter are explained by kinetic theory and forces of attraction
between particles. Key concepts include:
f) specific heat capacity
Solve the following problems. Show all work.
1. How much heat is lost as a 500. g cube of aluminum is cooled from 200°C to 25.0°C?
The specific heat for aluminum is 0.897 J/g°C
2. How much heat is required to melt 550.0 g of Cu that has already been heated to its
melting point? Heat of fusion is 205 J/g.
3. How many joules of heat are required to raise the temperature of 1.00 kg of water from
10.2˚C to 26.8˚C?
4. How much heat is released when 274 g of water cools from 85.2˚C to 38.4˚C?
TOPIC: CHEMICAL KINETICS
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 3F
The student will investigate and understand how conservation of energy and matter is
expressed in chemical formulas and balanced equations. Key concepts include:
f) reaction rates, kinetics and equilibrium
1. Define the following terms: catalyst, activated complex, activation energy, heating
curve, phase diagram, and heat capacity.
2. Draw labeled (reactants, products, activation energy, Δ H) reaction progress diagrams for
endothermic and exothermic reactions.
TOPIC: GAS LAWS
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 4A
The student will investigate and understand that quantities in a
chemical reaction are based on molar relationships. Key concepts
include:
a) Avogadro’s principle and molar volume.
Solve the following problems. Show all work.
1. What volume of oxygen gas at STP can be produced when 17.5 grams of potassium
chlorate is decomposed? 2KClO3 → 2KCl + 3O2
2. Chlorine gas will react with 3.45 L of hydrogen gas to yield what mass of hydrogen
chloride gas at STP? H2 + Cl2 → HCl
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 5B
The student will investigate and understand that quantities in a chemical
reaction are based on molar relationships. Key concepts include:
b) partial pressure and gas laws.
Solve the following problems. Show all work.
1. A mixture of He, Ne, Ar and Xe has a total pressure of 2.5 atmospheres. Calculate
the partial pressure of the He if the pressure due to Ne, Ar and Xe, respectively, is
0.3, 0.8 and 1.0 atmospheres.
2. A sample of carbon dioxide gas was collected over water at 24 °C. If the pressure of
the sample was 758 mm Hg and the vapor pressure of water is 20. mm Hg, what is the
pressure due to the carbon dioxide.
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 5B
The student will investigate and understand that quantities in a
chemical reaction are based on molar relationships. Key
concepts include:
b) partial pressure and gas laws.
1. If 4.50 g of methane gas (CH4) is introduced into an evacuated 2.00 L container at 35°C,
what is the pressure in the container, in atmospheres?
2. How many grams of iron would be produced from 52.5g of iron (III) oxide reacting with
aluminum? Al + Fe2O3 → Al2O3 + Fe
3. In the reaction between calcium hydride and water, the theoretical yield of hydrogen gas
from 75.0 g of calcium hydride is 7.18 g. In running this reaction the actual amount of
hydrogen produced was 6.94 g. What is the percent yield of this reaction? CaH2 + H2O
→ Ca(OH)2 + H2
4. Hydrogen gas can be produced in the laboratory by the reaction of magnesium metal with
hydrochloric acid. Mg + HCl → MgCl2 + H2 What is the limiting reactant, when 6.00 g
of HCl is added to 5.00 g of Mg to product hydrogen gas?
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 5B
The student will investigate and understand that the phases of matter are
explained by kinetic theory and forces of attraction between particles. Key
concepts include:
b) partial pressure and gas laws
1. A gas with a volume of 300 mL at 150°C is heated until its volume is 600 mL. What is
the new temperature of the gas if the pressure is constant?
2. Calculate the volume of a gas in liters at 1 atm if its volume at 900 mmHg is 1500 mL.
3. A 3.50 L gas sample at 20°C and a pressure of 650 mmHg is allowed to expand to a
volume of 8.00 L. What is the final temperature in degrees Celsius if the final pressure of
the gas is 425 mm Hg?
4. A gas cylinder contains nitrogen gas at 10 atm and temperature of 20°C. The cylinder is
left in the sun and the temperature of the gas increases to 50°C. What is the pressure in
the cylinder?
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 5C
The student will investigate and understand that the phases of matter are
explained by kinetic theory and forces of attraction between particles. Key
concepts include:
c) vapor pressure.
1.
How is the normal boiling point of a substance related to its vapor pressure?
2.
The atmospheric pressure of a mountain in South America is 400 mm Hg. Using the
vapor pressure curve of water from your textbook (page 401), determine the boiling
point of water at this altitude.
3.
Sketch a graph of the vapor pressure of a liquid. Show the dependency of pressure
on temperature.
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 5D
The student will investigate and understand that the phases of matter are
explained by kinetic theory and forces of attraction between particles. Key
concepts include:
d) phase changes.
Directions: Fill in the blanks with the correct response.
The particles in a ______________ phase are very close together in an orderly, fixed and usually
crystalline arrangement. _______________ is an endothermic change of state in which a solid
becomes a liquid. The temperature and pressure at which a solid becomes a liquid is its
____________________.
Because particles in the ____________ phase have enough kinetic energy to be able to move
past each other easily, they take the shape of their container. While many liquids flow readily,
many are resistant to flow, or are ________.
Because they are held close together, liquid particles are more affected by forces between
particles. They have attraction for each other, or ____________, as well as attraction for
particles of solid surfaces, called _______________. Liquids tend to form spherical drops
because of ____________, or the tendency to decrease their surface area to the smallest possible
size. Particles in a liquid can gain enough kinetic energy to leave the surface and become a gas
in a process called ______________.
Attractive forces between ____________ particles do not have a great effect, which makes the
particles independent of each other. The temperature and pressure at which the number of liquid
particles become gas particles is the same as the number of gas particles returning to the liquid
phase is called a substance’s _______________. Gas particles lose energy and become liquid
during ________________.
The process during which a liquid loses energy and becomes a sold is called ____________.
The temperature at which this change occurs is the _______________ of the substance.
When particles of solid become gas particles without first melting, the substance undergoes a
process called _____________. The reverse of this process, when a gas becomes a solid directly,
is called ________________.
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
TOPIC: ORGANIC CHEMISTRY AND BIOCHEMISTRY
CHEMISTRY
STANDARD CH. 6A, B
The student will investigate and understand how basic chemical properties relate to
organic chemistry and biochemistry. Key concepts include:
a) Unique properties of carbon that allow multi-carbon compounds
b) Uses in pharmaceuticals and genetics, petrochemicals, plastics and food
DIRECTIONS: Answer the questions below.
1.
List the 3 factors that make the bonding of carbon atoms unique.
2.
Draw Lewis structures, identify shape, and describe polarities of the following
molecules: CH4, C2H6, CH3CH2OH, C6H6
3.
List 6 natural, biological polymers.
4.
List 6 common pharmaceuticals that are organic compounds.
5.
List 6 man-made, synthetic polymers.
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
FOURTH Nine Weeks
SOLUTIONS
SOL- 3F, 4D, 5G
ACIDS AND BASES
SOL- 4D
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
TOPIC: SOLUTIONS
CHEMISTRY
STANDARD CH. 4C
The student will investigate and understand that quantities in a chemical
reaction are based on molar relationships. Key concepts include:
c) solution concentrations.
DIRECTIONS: Solve the following problems. Show all work.
1. What is the molality of a solution in which 3.0 moles of NaCl is dissolved in 1.5 kg of
water?
2. How many grams of I2 should be added to 750 g of CCl4 to prepare a 0.020 m solution?
3. Calculate the molarity of 5.85 g of NaCl in 2.00 L of solution.
4. How many moles are needed to make 2.0 L of 0.30M solution of Na2SO4?
5. How would you correctly prepare 125 mL of a 0.30M solution of copper(II) sulfate
(CuSO4) from a 2.00M solution of CuSO4?
6. What is the boiling point of a solution that contains 1.25 mol of CaCl2 in 1400 g of
water? Kb for water = 0.512 °C/m
7. What is the freezing point of Na2SO4 in 1750 g of water? Kf of water = 1.86 °C/m
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 3F
The student will investigate and understand that quantities in a
chemical reaction are based on molar relationships. Key concepts
include:
f) reaction rates, kinetics and equilibrium
DIRECTIONS: Answer the questions below.
1. What is a reversible chemical reaction? How does the format of the arrow differ from
a reaction that goes to completion?
2. What is Le Châtelier’s principle? How would the equilibrium in the following
reaction be impacted by:
CO (g) + O2 (g) ↔ CO2 (g)
a.
increasing [CO2]?
b.
decreasing pressure?
c.
decreasing [CO]?
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
TOPIC: ACIDS, BASES, and SOLUTIONS
CHEMISTRY
STANDARD CH. 4D
The student will investigate and understand that quantities in a chemical reaction are
based on molar relationships. Key concepts include:
d) acid/base theory: strong electrolytes, weak electrolytes, and nonelectrolytes;
dissociation and ionization, pH and pOH; and the titration process.
DIRECTIONS: Answer the questions below.
1. Define the following terms: conjugate acid, acid, base, conjugate base, Bronsted-Lowry
acid, Bronsted-Lowry base, indicator, neutralization and pH scale.
2. Solve the following problems.
a. Calculate the pH and pOH for each solution: [H+] = 5.0 x 10-6 and [OH-] = 4.5 x 10-11
b. Calculate the [H+] for each solution: pH = 5.0 and pH = 12.20
3. List the properties of an acid and a base.
4. Identify the conjugate acid and conjugate base in the following equation.
NH3 + H2O ↔ NH4+ + OH-
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
CHEMISTRY
STANDARD CH. 5G
The student will investigate and understand that the phases of
matter are explained by kinetic theory and forces of attraction
between particles. Key concepts include:
g) colligative properties.
DIRECTIONS: Answer the questions below.
1. What is the freezing point of a solution of 200 grams of ethylene glycol (HOCH2CH2OH)
mixed with 400 grams of water?
2. What is the boiling point elevation of a solution of 26.8 grams of NaCl dissolved in 500
grams of water?
3. Calculate the boiling point and freezing point of a solution of 220 g of sucrose
(C12H22O11) dissolved in 1.0 kg of water. How does this compare to the boiling point and
freezing point of pure water?
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
ANSWERS
FIRST NINE WEEKS
STANDARD CH. 1 F
1. Experimental Value = 1.44 g
Accepted Value = 1.54 g
6.5%
2. Experimental Value = 22.2 L
Accepted Value = 22.4 L
0.89%
3. Experimental Value = 128.6 mg
Accepted Value =128.3 mg
0.23%
4. Experimental Value = 9.98 x 10-3 g
Accepted Value =1.03 x 10-2 g
3.11%
STANDARD CH. 2H
Chemical and Physical Properties
The student will be given an object or substance and must identify the physical and chemical properties of that
object. (This assignment is to be given after lecture on chemical and physical properties and how to correctly
identify each.)
1.
Describe the chemical properties of this object below, citing any evidence you have
collected:
Answers will vary.
2.
Describe the physical properties of this object below, citing any evidence you have
Answers will vary.
collected:
STANDARD CH. 2I
.Students will create a two page foldable for the models of an atom. Each outside flap will be the names of the
scientists (Democritus, Dalton, Rutherford, Bohr), and the inside picture will be each scientist’s contribution to the
historical model. A one page foldable will be created for the quantum models with Heisenberg and Planck. Students
will be creative using mnemonics and/or graphics.
Answers will vary.
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
STANDARD CH. 2A
Element
Ne
Al
Ge
P
Br
Ca
K
I
Sr
Sn
Average Atomic mass
20.179
26.9812
72.59
30.97376
79.904
40.08
39.0983
126.904
87.62
118.69
Mass number
20
27
72
31
80
40
39
127
88
119
Atomic number
10
13
32
15
35
20
19
53
38
50
Atomic mass answers may slightly vary depending on the Periodic Table used.
STANDARD CH. 2C
Warm-up: Complete the following chart:
Subatomic Particle
Symbol
Charge
Relative mass
Location in atom
protons
P
+
1
In nucleus
neutrons
N
0
1
In nucleus
electrons
E
-
0
Outside nucleus or electron
cloud
Atomic Structure
Element/Ion
Atomic
Number
Atomic
Mass
Mass
Number
Protons
Neutrons
Electrons
3
H
1
1.00797
3
1
2
1
1
H+
1
1.00797
1
1
0
0
12
C
6
12.0115
12
6
6
6
7
Li+
3
6.941
7
3
4
2
17
35.453
35
17
18
18
39
K
19
39.0983
39
19
20
19
Mg2+
12
24.305
24
12
12
10
33
74.922
77
33
44
36
47
107.868
108
47
61
47
35
Cl-
24
77
As3-
108
Ag
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
30 -2
16
32.064
30
16
14
18
238
92
238.03
238
92
146
92
S
U
STANDARD CH. 2B
DIRECTIONS: Answer the following questions
Average Atomic Mass
1.
Rubidium has two common isotopes, Rb-85 and Rb-87. If the abundance of Rb-85 is
and the abundance of Rb-87 is 27.8%, what is the average atomic mass of rubidium? 85.556
Uranium has 3 common isotopes. If the abundance of
U-234 is 0.01%, the abundance of U-235 is 0.71% and the abundance of U-238 is
the average atomic mass? 237.9783
72.2%
2.
3.
Titanium has five common isotopes: Ti-46 (8.0%), Ti-47 (7.8%), Ti-48 (73.4%), Ti-49
and Ti-50 (5.3%). What is the average atomic mass of titanium? 47.923
99.28%, what is
(5.5%),
4.
Explain why atoms have different isotopes. In other words, how is it that helium can
have
three different sized atoms and they all are still the element Helium. Atoms of the same element will have the
same number of protons but different numbers of neutrons.
5.
Draw and label 3 possible isotopes of hydrogen. Which isotope can you predict to be
most abundant? Why? Correctly name the 3 isotopes using the symbol “H”?
Answers will vary.
the
6.
Select an element on the PTOE. Create 4 isotopes for your element and form an
average atomic mass calculation to support the given atomic mass.
Answers will vary.
STANDARD CH. 3A
Writing Binary Formulas
Cation
Anion
Formula
Cation
Anion
Formula
Na+
Cl-
NaCl
Fe+2
O-2
FeO
Ba+2
F-
BaF2
Fe+3
O-2
Fe2O3
K+
S-2
K2S
Cr+2
S-2
CrS
Li+
Br-
LiBr
Cr+3
S-2
Cr2S3
Al+3
I-
AlI3
Cu+
Cl-
CuCl
Zn+2
S-2
ZnS
Cu+2
Cl-
CuCl2
Ag+
O-2
Ag2O
Pb+2
O-2
PbO
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
Mg+2
P-3
Mg3P2
Pb+4
O-2
PbO2
Ni+2
O-2
NiO
Mn+2
Br-
MnBr2
Ni+3
O-2
Ni2O3
Mn+4
Br-
MnBr4
STANDARD CH. 2 D, E
Counting Valence Electrons
Element
Ne
Al
Ge
P
Br
Ca
K
I
Sr
Sn
Group Number
18
13
14
15
17
2
1
17
2
14
Number of Valence Electrons
8
3
4
5
7
2
1
7
2
4
STANDARD CH. 2 F
DIRECTIONS: State the periodic trend and explain.
ATOMIC RADII
LEFT TO
RIGHT
DECREASE
ELECTRONEGATIVITY
INCREASE
SHIELDING EFFECT
SAME
IONIZATION ENERGY
INCREASE
PERIODIC TRENDS
EXPLAINATION
STRONGER
NUCLEAR
ATTRACTION
MORE VALNCE
ELECTRONS
SAME ENERGY
LEVEL
MORE VALENCE
ELECTRONS
TOP TO
BOTTOM
INCREASE
EXPLAINATION
DECREASE
LARGER RADIUS
INCREASE
MORE ENERGY
LEVELS
MORE SHIELDING
DECREASE
MORE ENERGY
LEVELS
STANDARD CH. 2 C
Ions and Subatomic Particles
Ion Symbol
Protons
Electrons
Charge
2-
S
16
18
-2
K1+
19
18
+1
Ba2+
56
54
+2
Fe3+
26
23
+3
Fe2+
26
24
+2
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
F1-
9
10
-1
O2-
8
10
-2
P3-
15
18
-3
Sn4+
50
46
+4
Sn2+
50
48
+2
N3-
7
10
-3
Br1-
35
36
-1
Mg2+
12
10
+2
Cu1+
29
28
+1
Cu2+
29
27
+2
U6+
92
86
+6
Mn5+
25
20
+5
Cl1-
17
18
-1
Se2-
34
36
-2
SECOND NINE WEEKS
SOL 3C
1. D 2.A 3.C 4. C
5. B
6. C 7. D
8. A
9. D 10. D
SOL 3 A,C,D
1. A 2. D 3. B 4. Group 15 5. A 6. B 7. D 8. C 9. A
SOL 4A
1. C 2. C
3. A
4. C
5. A
THIRD NINE WEEKS
CHEMISTRY
STANDARD CH. 3e
11.
CO 
Fe2O3 +
S8 
12.
Cr +
13.
Eu +
HF 
14.
NH4Cl

15.
C12H22O11 +
16.
Zn +
17.
SiO2 +
FeO +
CO2 double replacement
Cr2S3 synthesis
EuF3 +
HCl +
NH3 decomposition
O2 
HCl 
C 
H2 single replacement
CO2 +
ZnCl2 +
Si +
H2O combustion
H2 single replacement
CO single replacement
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
18.
Pb(NO3)2 +
19.
KClO3 
20.
C6H6 +
H3AsO4 
KCl +
O2 
PbHAsO4 +
O2
CO2 +
HNO3 double replacement
decomposition
H2O
combustion
Directions: Complete the following table:
Reaction Energy
Description
Sign of ΔH, Positive
or Negative?
Exothermic
Endothermic
+
Temperature,
Increase or
Decrease?
increase
decrease
Energy,
Product or Reactant?
product
reactant
STANDARD CH. 1g
6.
Determine the number of significant figures in the following:
c.
d.
7.
602
1200
3
2
c. 0.00345 3
d. 0.1040 4
e. 34.08
f. 3
4
1
g. 0.970 3
Record your answer to the correct number of sig. figs.
c.
d.
8.
(7.502 x 102 )(5.43 x 104)
9.01 x 106 / 1.22 x 105
4.07 x 107
7.39 x 101
Write the following numbers in scientific notation, or translate the numbers to regular notation.
c.
d.
9.
556,000,000,000 5.56 x 1011
0.00000751 7.51 x 10-6
c. 9.32 x 10-5 0.0000932
d. 7.68 x 107 76,800,000
Use conversion factors to solve the following: SHOW ALL WORK!!
c.
d.
Convert 250 m to km. 0.25 km
Convert 18.50 kg to g. 1.85 x 104
c. Convert 3.54 L to kL. 3.54 x 10 -3
d. Convert 0.5420 kg to mg 5.42 x 10 5
10. Convert the following temperatures to Kelvin or to degree’s Celsius.
b.
100˚C 373K b. 45˚C 318K c. 273 K 0°C
d. 400 K 127°C
CH. 3B
21.
Fe2O3 +
CO 
2FeO +
22.
16Cr +
3S8 
8Cr2S3
23.
2Eu +
6HF 
2EuF3 +
CO2
3H2
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
24.
C12H22O11 + 12 O2 
25.
Zn +
26.
SiO2 +
27.
Pb(NO3)2 +
28.
2NaCl + 2 H2O 
29.
2 KClO3 
30.
2C6H6 +
2 HCl 
2C 
12 CO2 + 11H2O
ZnCl2 +
Si +
2CO
H3AsO4 
2KCl +
15O2 
H2
PbHAsO4 + 2 HNO3
Cl2 +
H2 + 2NaOH
3O2
12CO2 +
6 H2O
STANDARD CH. 4B
12. Fill in the blanks.
d. The amount of a substance equal to 6.02 x 1023 particles is called a _mole_______________________.
e. 6.02 x 1023 is also known as ____Avogadro’s number_____________.
f. If you have 1 mole of copper, you have 6.02 x 10 23 __atoms__________.
13. Molar (Formula) Mass. Determine the molar mass of the following and show all work and units.
b. CuSO4 159.58g/mol b. CaCO3 100.06g/mol c. H3PO4 97.96g/mol d. Al2(SO4)3 342.05g/mol
14. Mole Conversions. Solve the following problems. Show all work
f. How many grams are there in 1.55 x 1023 molecules of CO2? 11.3g
g. How many atoms are there in 2.18 moles of nitrogen gas? 1.31 x 1024atoms
h. What is the volume, at STP, occupied by 4.30 moles of oxygen gas?96.3L
i.
j.
How many grams are there in 56.32 liters of carbon monoxide, at STP?70.40g
How many moles are in 8.67 x 1025 atoms of sulfur?144 moles
15. Percent Composition. Determine the percent composition of the following compounds. Show all work.
b. NO2 30.5%N, 69.5%O b. Al2(SO4)3 15.8%Al, 28.1%S, 56.1%O c. What is the percent of Mg in
MgCl2?25.5%Mg
Solve the following problems and show all work.
16. Determine the empirical formula for 71.5% Ca and 28.5% O. CaO
17. Determine the molecular formula for the empirical formula of C3H5O2, with a molar mass of 146 g/mol.
C6H10O4
18. 4Al + 3O2 → 2Al2O3 , How many moles of oxygen are required to react completely with 0.84 mol of
Al? .63 mol O2 How many moles of aluminum are needed to form 2.3 mol of Al 2O3?4.6 mol Al
19. The combustion of acetylene gas is represented by this equation: 2C2H2(g) + 5O2(g) → 4CO2(g) + 2H2O(g)
a. How many grams of oxygen are required to burn 13.0g of C 2H2 ? 39.9 g O2 b. How many moles of C2H2
are needed to react completely with 98.0 g of water? 5.44 mol C2H2
20. Calcium carbonate can be decomposed by heating. CaCO3(s) → CaO(s) + CO2(g)
What is the percent yield of this reaction if 24.8 g of CaCO 3 is heated to give 13.1 g of CaO?94.2%
21. Copper reacts with sulfur to form copper (I) sulfide. 2Cu(s) + S(s) → Cu2S(s)
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
What is the limiting reactant when 80.0 g of Cu reacts with 25.0 g of S?LR=Cu
22. How many moles of CS2 form when 2.7 mol of C reacts? 5C + 2SO2 → CS2 +4CO
0.54 mol CS2
CHEMISTRY STANDARD CH. 3E
See Previous “Try It” in Chemical Reactions for CH.3E
STANDARD CH. 4B
See previous “Try It” in Stoichiometry for CH.4b
STANDARD CH. 5D
Solve the following problems. Show all work.
3.
How much heat is released as a 75.0 g sample of ethanol gas at the boiling point condenses to a liquid? Heat of
vaporization is 879 J/g.6.59 x 104J
4.
The heat of vaporization of water is 540. cal/g. How many calories would be needed to convert 3 moles of
water to vapor?3 x 104 cal
STANDARD CH. 5E
Solve the following problems. Show all work.
5.
How much heat is lost as a 500. g cube of aluminum is cooled from 200°C to 25.0°C? The specific heat for
aluminum is 0.897 J/g°C 8 x 104 J
6.
How much heat is required to melt 550.0 g of Cu that has already been heated to its melting point? Heat of
fusion is 205 J/g. 1.128 x 105 J
7.
How many joules of heat are required to raise the temperature of 1.00 kg of water from 10.2˚C to 26.8˚C?
6.95 x 104 J
8.
How much heat is released when 274 g of water cools from 85.2˚C to 38.4˚C?
-5.37 x 104 J
STANDARD CH. 3F
1.
Define the following terms: catalyst, activated complex, activation energy, heating curve, phase diagram,
and heat capacity.
a. Catalyst – a substance that changes the rate of a chemical reaction without being consumed.
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
b.
c.
d.
e.
f.
2.
Activated complex – molecule in an unstable state intermolecular to the reactant and products in
the chemical reaction.
Activation energy – the minimum amount of energy required to start a chemical reaction.
Heating curve – shows the change in temperature of a substance as heat is added and the substance
undergoes phase changes.
Phase diagram – graph of the relationship between the physical state of a substance, temperature,
and pressure of the substance.
Heat capacity – energy needed to increase the temperature of 1 mole (1 gram) of the substance by
1°C.
Draw labeled (reactants, products, activation energy, Δ H) reaction progress diagrams for endothermic and
exothermic reactions.
STANDARD CH. 4A
Solve the following problems. Show all work.
1.
What volume of oxygen gas at STP can be produced when 17.5 grams of potassium chlorate is
decomposed? 2KClO3 → 2KCl + 3O2 4.80L O2
2.
Chlorine gas will react with 3.45 L of hydrogen gas to yield what mass of hydrogen chloride gas
at STP? H2 + Cl2 → HCl 11.2g HCl
STANDARD CH. 4C
Solve the following problems. Show all work.
1.
A mixture of He, Ne, Ar and Xe has a total pressure of 2.5 atmospheres. Calculate the partial pressure
of the He if the pressure due to Ne, Ar and Xe, respectively, is 0.3, 0.8 and 1.0 atmospheres. 0.4 atm
2.
A sample of carbon dioxide gas was collected over water at 24 °C. If the pressure of the sample was
758 mm Hg and the vapor pressure of water is 20. mm Hg, what is the pressure due to the carbon
dioxide. 738 mmHg
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
STANDARD CH. 4D
1.
If 4.50 g of methane gas (CH4) is introduced into an evacuated 2.00 L container at 35°C,
what is the pressure in the container, in atmospheres? 3.5 atm
2.
How many grams of iron would be produced from 52.5g of iron (III) oxide reacting with aluminum? Al +
Fe2O3 → Al2O3 + Fe 36.7 g
3.
In the reaction between calcium hydride and water, the theoretical yield of hydrogen gas from 75.0 g of
calcium hydride is 7.18 g. In running this reaction the actual amount of hydrogen produced was 6.94 g.
What is the percent yield of this reaction? CaH2 + H2O → Ca(OH)2 + H2 96.4%
4.
Hydrogen gas can be produced in the laboratory by the reaction of magnesium metal with hydrochloric
acid. Mg + HCl → MgCl2 + H2 What is the limiting reactant, when 6.00 g of HCl is added to 5.00 g of Mg
to product hydrogen gas? LR=HCl
STANDARD CH. 5A
1.
A gas with a volume of 300 mL at 150°C is heated until its volume is 600 mL. What is
the new temperature of the gas if the pressure is constant? 800K
2.
Calculate the volume of a gas in liters at 1 atm if its volume at 900 mmHg is 1500 mL.
2L
A 3.50 L gas sample at 20°C and a pressure of 650 mmHg is allowed to expand to a volume of 8.00 L.
What is the final temperature in degrees Celsius if the final pressure of the gas is 425 mm Hg? 200°C
3.
4.
A gas cylinder contains nitrogen gas at 10 atm and temperature of 20°C. The cylinder is left in the sun and
the temperature of the gas increases to 50°C. What is the pressure in the cylinder? 10 atm
STANDARD CH. 5B
1.
2.
3.
How is the normal boiling point of a substance related to its vapor pressure?
When you increase the temperature of a system to the point at which the vapor pressure of a substance
is equal to the standard atmospheric pressure, you have reached the substance’s normal boiling point.
The atmospheric pressure of a mountain in South America is 400 mm Hg. Using the vapor pressure
cure of water from your textbook (page 401), determine the boiling point of water at this altitude.
84°C
Sketch a graph of the vapor pressure of a liquid. Show the dependency of pressure on temperature.
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
STANDARD CH. 5C
Directions: Fill in the blanks with the correct response.
The particles in a __solid________ phase are very close together in an orderly, fixed and usually crystalline
arrangement. __Melting_______ is an endothermic change of state in which a solid becomes a liquid. The
temperature and pressure at which a solid becomes a liquid is its _heat of fusion___.
Because particles in the __liquid______ phase have enough kinetic energy to be able to move past each other easily,
they take the shape of their container. While many liquids flow readily, many are resistant to flow, or are
viscous_____.
Because they are held close together, liquid particles are more affected by forces between particles. They have
attraction for each other, or cohesion__, as well as attraction for particles of solid surfaces, called adhesion_.
Liquids tend to form spherical drops because of surface tension_, or the tendency to decrease their surface area to
the smallest possible size. Particles in a liquid can gain enough kinetic energy to leave the surface and become a gas
in a process called _vaporization______.
Attractive forces between __gas____ particles do not have a great effect, which makes the particles independent of
each other. The temperature and pressure at which the number of liquid particles become gas particles is the same
as the number of gas particles returning to the liquid phase is called a substance’s __heat of vaporization___.
Gas particles lose energy and become liquid during _condensation____.
The process during which a liquid loses energy and becomes a sold is called freezing. The temperature at which this
change occurs is the _freezing point______ of the substance.
When particles of solid become gas particles without first melting, the substance undergoes a process called
sublimation___. The reverse of this process, when a gas becomes a solid directly, is called deposition____.
STANDARD CH.6A, B
1.
List the 3 factors that make the bonding of carbon atoms unique.
Strong single bond, not extremely reactive, different bonding arrangements
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
2.
Draw Lewis structures, identify shape, and describe polarities of the following molecules: CH 4, C2H6,
CH3CH2OH, C6H6
CH4 –tetrahedron, non-polar
C2H6,- tetrahedron at each carbon, non-polar
CH3CH2OH – tetrahedron at each carbon, polar due to the oxygen
C6H6 – hexagonal, planar, non-polar
3.
List 6 natural, biological polymers.
DNA, RNA, starch, cellulose, protein, poly-lipids, etc.
4.
List 6 common pharmaceuticals that are organic compounds.
aspirin, vitamins, insulin, Tylenol, ibuprofen, Ritalin, etc.
5.
List 6 man-made, synthetic polymers.
Nylon, polyethylene, polypropylene, PVC, polystyrene, Kevlar, etc.
FOURTH NINE WEEKS
STANDARD CH. 4E
1. What is the molality of a solution in which 3.0 moles of NaCl is dissolved in 1.5 kg of water? 2.0 m
2.
3.
How many grams of I2 should be added to 750 g of CCl4 to prepare a 0.020 m solution?
3.8 g
Calculate the molarity of 5.85 g of NaCl in 2.00 L of solution. 0.0500 M
4.
How many moles are needed to make 2.0 L of 0.30M solution of Na 2SO4? 0.60 mol
5.
How would you correctly prepare 125 mL of a 0.30M solution of copper(II) sulfate (CuSO 4) from a 2.00M
solution of CuSO4? 19 mL
6.
What is the boiling point of a solution that contains 1.25 mol of CaCl 2 in 1400 g of water? Kb for water =
0.512 °C/m 101.37°C = 101°C
7.
What is the freezing point of Na2SO4 in 1750 g of water? Kf of water = 1.86 °C/m
-3.19°C
STANDARD CH. 4F
DIRECTIONS: Answer the questions below.
3.
What is a reversible chemical reaction? A chemical reaction in which the products reform the original
reactants. How does the format of the arrow differ from a reaction that goes to completion? Arrows
that point in opposite directions.
4.
What is Le Châtelier’s principle? States that a system in equilibrium will oppose a change in a way that
helps eliminate the change. How would the equilibrium in the following reaction be impacted by:
CO (g) + O2 (g) ↔ CO2 (g)
d.
increasing [CO2]? Shift to the left
e.
decreasing pressure? Shift to the right
f.
decreasing [CO]? Shift to the left
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
STANDARD CH. 4G
DIRECTIONS: Answer the questions below.
1. Define the following terms: conjugate acid, acid, base, conjugate base, Bronsted-Lowry
acid, Bronsted-Lowry base, indicator, neutralization and pH scale.
Conjugate acid- compound with a hydrogen ion
Acid – donates a proton
Base – accepts a proton
Conjugate base – same compound of the acid but without the hydrogen ion
Bronsted-Lowery acid – substance that is capable of donating a proton
Bronsted-Lowery base – substance that is capable of accepting a proton
Indicator – dyes that can be added that will change color in the presence of an acid or base
Neutralization – adding an acid to a base to produce a salt and water
pH scale – a way of expressing the strength of acids and bases
2. Solve the following problems.
a. Calculate the pH for each solution: [H+] = 5.0 x 10-6 pH = 5.3and [OH-] = 4.5 x 10-11 pH = 3.7
b. Calculate the [H+] for each solution: pH = 5.0 [H] = 1x10-5 pH = 12.20 [H] = 6.3 x 10-13
1.
List the properties of an acid and a base.
Acid: produces hydrogen ions, taste sour, corrodes metal, electrolyte, pH<7, reacts with a base, litmus
paper turns red
Base: produces hydroxide ions, taste bitter, electrolytes, reacts with acids, pH>7, litmus paper turns
blue
4. Identify the conjugate acid and conjugate base in the following equation.
NH3 + H2O ↔ NH4+ + OHConjugate acid: NH4+ conjugate base: NH3
Conjugate acid: H2O, conjugate base: OH-
STANDARD CH. 5F
DIRECTIONS: Answer the questions below.
1.
What is the freezing point of a solution of 200 grams of ethylene glycol (HOCH 2CH2OH) mixed with 400
grams of water? -14.9°C
2.
What is the boiling point elevation of a solution of 26.8 grams of NaCl dissolved in 500 grams of water?
100.9°C
3.
Calculate the boiling point and freezing point of a solution of 220 g of sucrose (C 12H22O11) dissolved in 1.0
kg of water. Tb = 100.3°C Tf = -1.2°C How does this compare to the boiling point and freezing point of
pure water? Both are close to pure water.
CHEMISTRY STANDARD CH. 4G
See previous “Try It” in Solutions for CH.4G
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin
REFERENCES
www.222.nano.gov
www.chemmybear.com
www.chemistry.merlot.org
Holt Chemistry Textbook http://my.hrw.com/
www.nsta.org
Prentice Hall Textbook http://www.phschool.com/home.html
Virginia Department of Education Blueprint
http://www.doe.virginia.gov/testing/sol/blueprints/science_blueprints/blueprint_ch
emistry.pdf
Virginia Department of Education Framework
http://www.doe.virginia.gov/testing/sol/frameworks/science_framewks/framework
_science-chem.pdf
Virginia Department of Education Scope and Sequence
http://www.doe.virginia.gov/testing/sol/scope_sequence/science_scope_sequence/s
copeseq_science_chemistry.pdf
Virginia Department of Education Standards of Learning
http://www.doe.virginia.gov/testing/sol/standards_docs/science/courses/stds_chemi
stry.pdf
www.scitoys.com
www.stevespanglescience.com
Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver,
and Ashanta Ruffin