Download Science SOL CH

Science Curriculum Matrix
Chemistry
August 1, 2009
The Science Vertical Team has revised the Chemistry Curriculum Matrix for 2009-2010. In addition to the necessary correlation
to the Virginia Science Standards of Learning, the Chemistry content is organized by both concepts and topics. We encourage
you to utilize this document while planning for instruction. A more dynamic version of this matrix is available on our wiki site at
http://acpsscience.pbworks.com/. We anticipate making additional updates to this document as the school year progresses.
Please contact Tony Borash with your comments and suggestions at [email protected].
In addition to this document, we recommend that you review the Chemistry Curriculum Framework for additional clarification
regarding the Chemistry Science SOL and the Chemistry Enhanced Scope and Sequence for unit and lesson planning resources.
Thanks,
The Science Vertical Team
Scientific Investigation, Reasoning, and Logic
The Science Vertical Team will work to develop a recommended integration for the science processes and skills generally found
in standard CH.1 into the more content specific standards CH.2 through CH.5. In the meantime, Virginia Science SOL CH.1 and
the associated essential skills and processes are presented below. Please incorporate these specific science processes and skills
into your daily science instruction as much as possible and practical.
CH.1 The student will investigate and understand that experiments in which variables are measured, analyzed, and evaluated,
produce observations and verifiable data.
a.
b.
c.
d.
e.
f.
g.
h.
i.
designated laboratory techniques;
safe use of chemicals and equipment;
proper response to emergency situations;
manipulation of multiple variables with repeated trials;
accurate recording, organizing, and analysis of data through repeated trials;
mathematical and procedural error analysis;
mathematical manipulations (SI units, scientific notation, linear equations, graphing, ratio and proportion, significant
digits, dimensional analysis);
the use of appropriate technology including computers, graphing calculators, and probeware for gathering data and
communicating results; and
construction and defense of a scientific viewpoint (the nature of science).
Essential Skills and Processes
In order to meet this standard, it is expected that students should know and be able to:
For CH.1 a,b, and c:
•
•
•
•
•
•
Make the following measurements, using the specified equipment:
- volume: graduated cylinder, pipette, volumetric flask, buret
- mass: electronic or dial-a-gram
- temperature: thermometer and/or temperature probe
- pressure: barometer or pressure probe.
Identify, locate, and know how to use laboratory safety equipment, including aprons, goggles, gloves, fire extinguishers,
fire blanket, safety shower, eye wash, broken glass container, and fume hood.
Demonstrate the following basic lab techniques: filtering, decanting, using chromatography, and lighting a gas burner.
Identify the following basic lab equipment: beaker, flask, graduated cylinder, test tube, test tube rack, test tube holder,
ring stand, wire gauze, clay triangle, crucible with lid, evaporation dish, watch glass, wash bottle, and dropping pipette.
Understand Material Safety Data Sheet (MSDS) warnings, including handling chemicals, lethal dose (LD), hazards,
disposal, and chemical spill cleanup.
Demonstrate safe laboratory practices, procedures, and techniques.
For CH.1 d and e:
•
•
Design and perform experiments to test predictions.
Identify variables.
•
•
•
•
Predict outcome(s) when a variable is changed.
Record data, using the significant digits of the measuring equipment.
Demonstrate precision (reproducibility) in measurement.
Recognize accuracy in terms of closeness to the true value of a measurement.
For CH.1 f and g:
• Discover and eliminate procedural errors.
• Know most frequently used SI prefixes and their values (milli-, centi-, deci-, kilo-).
• Demonstrate the use of scientific notation, using the correct number of significant digits with powers of ten notation for
the decimal place.
• Correctly utilize the following when graphing data.
- dependent variable (vertical axis)
- independent variable (horizontal axis)
- scale and units of a graph
- regression line (best fit curve).
• Calculate mole ratios, percent composition, conversions, and relative atomic mass.
• Use the rules for performing operations with significant digits.
• Utilize dimensional analysis.
• Use graphing calculators correctly.
• Read a measurement from a graduated scale, stating measured digits plus the estimated digit.
• Use data collected to calculate percent error.
• Determine the mean of a set of measurements.
For CH.1 h and i:
•
•
•
Use appropriate technology for data collection and analysis, including probeware interfaced to a graphing calculator and/or
computer.
Use probeware to gather data.
Explain the emergence of modern theories based on historical development. For example, students should be able to
explain the origin of the atomic theory beginning with the Greek atomists and continuing through the most modern
Quantum models.
Physical Science: Matter: Structure
COURSE: Chemistry
CONCEPT: SCALE: Properties
ENDURING UNDERSTANDING: Properties characterize objects, organisms, and substances.
Essential Understandings
Assessment Samples – SOL/Blooms
Students should understand:
Knowledge/Comprehension Level
• The periodic table is
arranged in order of
increasing atomic numbers.
• The atomic number of an
element is the same as the
number of protons.
• In a neutral atom, the
number of electrons is the
same as the number of
protons.
• All atoms of an element have
the same number of protons.
• The atomic mass for each
element is the weighted
average of that element’s
naturally occurring isotopes.
•
Students use colored counters that represent protons, electrons, and
neutrons to form different “atoms.” They should learn that if the number of
protons is changed, the identity of the element is changed. If the number
of neutrons is changed, the mass number is changed. If the number of
electrons is changed, the charge is changed. (See Appendix A)
Application/Analysis Level
•
Students use lima beans as a model of different isotopes. Students are
given small beans and large beans. Students can calculate the average
mass of each type of bean. Students are then given a sealed bean sample
that contains some small and some large beans. Based on the average
mass of their sample, they can calculate the percentage of each type of
bean, using algebra. (See Appendix B)
Synthesis/Evaluation Level
•
Vocabulary
atom
protons
electrons
neutrons
charge
atomic number
average atomic mass
mass number
isotope
periodic table
half-life
radioactive decay
alpha radiation
beta radiation
gamma radiation
Students can do research on radioactive isotopes to explore the following
questions:
o What factors influence the stability of a nuclide?
o What are the different types of nuclear decay?
o What happens when an atom undergoes nuclear decay?
o Is there a way to predict what type of decay a particular
radioisotope will undergo?
o How can I use half-life data to calculate the amount of radioactive
material present after a certain time period?
SOL: CH.2 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;
(b) isotopes, half lives, and radioactive decay;
(c) characteristics of subatomic particles as to mass and charge.
August 1, 2009
SOL CH.2a,b,c – Appendix A
PROTONS, NEUTRONS, AND ELECTRONS
Name
Part One – Basic Information about the Atom
The ATOMIC NUMBER of Helium is 2.
Atomic Mass
Element Symbol
This means that every atom of helium contains
exactly 2
Atomic Number
.
Element Name
The ATOMIC NUMBER can be used to identify
an element. There are at least 110 different
elements on the periodic table. Each element can be uniquely identified by its atomic number.
Look at your periodic table. What is the ATOMIC NUMBER of Silver (Ag)?
This means that each atom of silver contains
PROTONS have a
ELECTRONS have a
.
.
charge, and they are located in the
of the atom.
charge, and they are located
.
The ATOMIC MASS of Helium is 4.00260. This number represents the average mass of all of the
different isotopes of Helium that exist in nature. We will be discussing more information about
isotopes later.
For now, let us assume that we have an atom of helium that has a MASS NUMBER of 4.
You should know that protons are much, much heavier than electrons. In fact, a proton is about 2000
times heavier than an electron. For this reason, scientists have assigned the following mass numbers
to a proton and an electron.
PROTON
ELECTRON
Mass Number
1
0
The ATOMIC NUMBER of a Helium atom is 2, which corresponds to 2 protons.
If the MASS NUMBER of a Helium atom is 4, where are the other 2 mass units coming from?
The answer is that this helium atom also contains 2
.
Fill in the following table, based on the information for each of the three subatomic particles.
Mass Number Charge
PROTON
ELECTRON
NEUTRON
Location in the Atom
An atom is neutral and has no overall charge. This means that in the atom there should be the same
number of
and
.
The MASS NUMBER of an atom is equal to the sum of the
and
.
Check with your teacher before proceeding to the next step.
Part Two – Subatomic Particles
Each of you should have a partner, and each of you should have a cup that contains the following
items:
PROTONS (brown beans)
NEUTRONS (black beans)
ELECTRONS (white beans)
Look in your cup and count all of your beans.
Write the numbers of each type of particle for your cup and for your partner’s cup in the table below.
Your Cup
PROTONS NEUTRONS ELECTRONS
Your Partner’s Cup
PROTONS NEUTRONS ELECTRONS
Now we will review the symbol that chemists use to represent an atom. Let’s use helium-4 as an
example:
Mass Number (protons + neutrons)
Atomic Number (protons)
!
2
He
Based on the number of subatomic particles in your cup, write the symbol for the atom that is
represented by these particles. Do the same thing for your partner’s cup.
Your Atomic Symbol
Your Partner’s Atomic
Symbol
Check with your teacher before proceeding to the next step.
Part Three – Creating New Elements
If you are the person with Fluorine, take ONE PROTON and ONE ELECTRON from your pile of
beans and give them to your partner. Now fill in the table below with the new count of subatomic
particles.
Your Atom
PROTONS NEUTRONS ELECTRONS
Your Partner’s Atom
PROTONS NEUTRONS ELECTRONS
Did the chemical identity of your elements change when you made this change?
If you said yes, then what are the NEW elements that you have now?
If you said no, then explain why the chemical identity did NOT change.
When an atom gains or loses a PROTON, what property of the atom is going to change?
Based on the number of subatomic particles that you have now, write the symbol for the atom that
you have. Write your partner’s symbol as well..
Your Atomic Symbol
Your Partner’s Atomic
Symbol
Check with your teacher before proceeding to the next step.
Part Four – Exploring Isotopes
If you are the person with Oxygen, take TWO NEUTRONS from your pile and give them to your
partner. Now fill in the table below with the new count of subatomic particles.
Your Bag
PROTONS NEUTRONS ELECTRONS
Your Partner’s Bag
PROTONS NEUTRONS ELECTRONS
Did the chemical identity of your elements change when you exchanged NEUTRONS?
If you said yes, then what are the NEW elements that you have now?
If you said no, then explain why the chemical identity did NOT change.
When an atom gains or loses a NEUTRON, what property of the atom is going to change?
Based on the number of subatomic particles that you have now, write the symbol for the atom that
you have. Write your partner’s symbol as well..
Your Atomic Symbol
Your Partner’s Atomic
Symbol
Two atoms that have the SAME number of PROTONS and a DIFFERENT number of NEUTRONS
are called
.
Are the atoms shown below are ISOTOPES of each other?
Explain why or why not.
Are the atoms shown below are ISOTOPES of each other?
Explain why or why not.
A chemical property is defined as “the ability of a substance to undergo a change that transforms it
into a new substance.” For example, the ability of iron to rust is a chemical property. If you have 2
different isotopes of the same element, should they have the same chemical properties or different
chemical properties? Explain why you think so.
Check with your teacher before proceeding to the next step.
Part Five – Charges and Ions
If you are the person with Magnesium, take TWO ELECTRONS from your pile and give them to your
partner. Now fill in the table below with the new count of subatomic particles.
Your Bag
PROTONS NEUTRONS ELECTRONS
Your Partner’s Bag
PROTONS NEUTRONS ELECTRONS
Did the chemical identity of your elements change when you exchanged ELECTRONS?
If you said yes, then what are the NEW elements that you have now?
If you said no, then explain why the chemical identity did NOT change.
When an atom gains or loses an ELECTRON, what property of the atom is going to change?
A atom is neutral when the number of
is equal to the number of
An ION is defined as an atom that has a positive or negative charge.
If you compare the number of protons and electrons, how can you determine the overall charge on an
ION?
If an atom LOSES ELECTRONS, then it will become
charged.
If an atom GAINS ELECTRONS, then it will become
charged.
The charge on an ion can be calculated as follows: (PROTONS) – (ELECTRONS) = charge on the
ion
Based on the number of subatomic particles that you have now, write the symbol for the atom that
you have. Write your partner’s symbol as well..
Your Ion Symbol
Your Partner’s Ion Symbol
Check with your teacher to find out what to do next
PLEASE MAKE SURE THAT THE CLEAR PLASTIC CUP HAS
9 BROWN BEANS, 9 WHITE BEANS, AND 10.BLACK BEANS.
PLEASE MAKE SURE THAT THE RED PLASTIC CUP HAS
11 BROWN BEANS, 11 WHITE BEANS, AND 12.BLACK BEANS.
PROTONS, NEUTRONS, AND ELECTRONS
Name
Summary Questions
1. When you change the number of PROTONS in an atom, you will change the
2. When you change the number of NEUTRONS in an atom, you will change the
3. When you change the number of ELECTRONS in an atom, you will change the
4. ISOTOPES are defined as two atoms that have the same number of
but they
each have a different number of
5. An ION is defined as an atom that has
6. True or False: “For a neutral atom, the number of protons should be equal to the number of electrons.”
If you answered True, then explain WHY this statement is true.
If you answered False, then write the complete symbol for a neutral atom in which the number of protons is NOT equal to
the number of electrons.
7. Choose a neutral atom in which the number of protons is NOT equal to the number of neutrons. (Circle your answer)
+
-2
8. A student saw the following list of subatomic particles: 22P, 26N, 20E
The student then wrote the atomic symbol as
Is the student correct?
-2
If yes, explain why. If not, then write the correct symbol
9. A student saw the following list of subatomic particles: 50P, 69N, 50E
The student then wrote the atomic symbol as
Is the student correct?
If yes, explain why. If not, then write the correct symbol
SOL CH.2a,b,c – Appendix B
Exploring Isotopes of “Beanium”
Name
PRE-LAB QUESTIONS
63
65
1. Two different isotopes of copper ( Cu and Cu) are being compared to each other.
For each property listed below, answer if you think that these two isotopes would have the SAME property or
if you think that they would have DIFFERENT properties.
(a) Number of PROTONS in the atom
(c) Mass of the atom
(b) Number of NEUTRONS in the atom
(d) Chemical Reactivity
2. Suppose that your grades in Math class were based on Tests and Homework, and you received the following
scores:
Tests = 90
Homework = 80
(a) Your teacher tells you that he will count Tests as 50% of your grade and Homework as 50% of your grade.
Calculate your Math Grade, based on these percentages
(b) Suppose that your teacher changes his mind and decides that he will count Tests as 80% of your grade and
Homework as 20% of your grade. Now calculate your new Math Grade, based on these new percentages.
3. At the end of the semester, a student’s grades were reported as follows:
Nine Weeks Grade
90
Exam Grade
70
Semester Grade
85
(a) Based on the results shown above, do you think that the Nine Weeks Grade and the Exam Grade were each
worth 50% of the Semester Grade? Answer Yes or No, and then explain why you think so.
(b) Which grade (Nine Weeks or Exam) do you think counted for a higher percentage when determining the
student’s Semester Grade? Explain why you think so.
10
11
4. There are two possible isotopes of Boron: B and B. Based on the average atomic mass found on the periodic
table, which isotope of boron is more abundant in nature? Explain.
Exploring Isotopes of “Beanium”
Name
Part One – Determining the Average Mass of a Single Atom of Beanium
You should two cups of beans. One is labeled “Isotope A” and the other is “Isotope B.”
1. Place the small plastic cup labeled “A” on the electronic scale and press the zero button.
The scale should read “0.00 g” with the cup on the pan.
2. Add the Beanium sample that is labeled “Isotope A” to the plastic cup.
Record the mass of Isotope A in Data Table 1.
3. Count the individual “atoms” of Isotope A as you put them back in the larger cup. (There should be 20 “atoms.”)
Record this number of atoms in Data Table 1.
4. Repeat Steps 1 through 3, using the small plastic cup labeled “B” and “Isotope B.”
5. Perform the calculations necessary to complete Data Table 1.
(Hint: Average Mass of One Atom = Mass of Isotope divided by Number of Atoms)
Data Table 1
Mass of Isotope A
Number of Atoms in
Isotope A
Average Mass of
One Atom of
Isotope A
Mass of Isotope B
Number of Atoms in
Isotope B
Average Mass of
One Atom of
Isotope B
Part Two – Determining the Average Mass for Different Samples of Beanium
1. Place TWO small plastic cups (one marked “A” and the other marked “B”) on the electronic scale and press the
zero button. The scale should read “0.00 g” with the two cups on the pan.
2. Place 2 atoms of Isotope A into the A cup. Place 18 atoms of Isotope B into the B cup.
3. Record the Total Mass of the Beanium Sample in Data Table 2..
4. Now adjust the number of atoms so that you have 5 atoms of A in the A cup, and 15 atoms of B in the B cup.
Record the Total Mass of the Beanium Sample.
5. Now adjust the number of atoms so that you have 8 atoms of A in the A cup, and 12 atoms of B in the B cup.
Record the Total Mass of the Beanium Sample.
6. Now adjust the number of atoms so that you have 12 atoms of A in the A cup, and 8 atoms of B in the B cup.
Record the Total Mass of the Beanium Sample.
7. Now adjust the number of atoms so that you have 15 atoms of A in the A cup, and 5 atoms of B in the B cup.
Record the Total Mass of the Beanium Sample.
8. Now adjust the number of atoms so that you have 18 atoms of A in the A cup, and 2 atoms of B in the B cup.
Record the Total Mass of the Beanium Sample.
Data Table 2
Number of
Atoms of
Isotope A
Number of
Atoms of
Isotope B
2
18
5
15
8
12
12
8
15
5
18
2
Total Number
of Atoms in
the Beanium
Sample
Percentage
of Isotope A
Percentage
of Isotope B
Helpful Hints for Calculations
Number of Atoms of Isotope A
Percentage of Isotope A = -------------------------------------------- x 100%
20
Number of Atoms of Isotope B
Percentage of Isotope B = ------------------------------------------ x 100%
20
Average Mass of One Atom of Beanium = Total Mass of Beanium Sample ÷ 20
Total Mass of
Beanium
Sample
Average Mass
of One Atom
of Beanium
Exploring Isotopes of “Beanium”
Name
ANALYSIS QUESTIONS
1. “Beanium” exists in two isotopes, which are called Isotope A and Isotope B.
What is the main difference between Isotope A and Isotope B?
2. Chlorine exists in two isotopes, which are called chlorine-35 and chlorine-37.
What is the main difference between chlorine-35 and chlorine-37?
3. Copy your data from Data Table 1 into the blanks below:
Average Mass of One
Atom of Isotope A:
Average Mass of One
Atom of Isotope B:
4. Suppose that you are given a sample of Beanium. You know that it contains 18 atoms of Isotope A and 2 atoms
of Isotope B. You are asked to calculate the average mass of one atom of your sample. Should you use the
following formula? Answer yes or no and explain why you think so.
(Mass of Isotope A) + (Mass of Isotope B)
-------------------------------------------------------- = Average Mass of Beanium
2
5. Calculate the average mass of Beanium for a sample that contains 90% Isotope A and 10% Isotope B. Use the
numbers that you wrote down in Question 3. Show your work below, by filling in the blanks with the appropriate
numbers.
x
Percent Abundance of A
as a decimal
=
Average mass of Isotope A
x
Percent Abundance of B
as a decimal
=
add these two
values
together
Average mass of Isotope B
Average mass of Beanium sample =
6. Compare your answer from Question 5 with the value from Data Table 2.
(Look for the box that is heavily outlined in the lower right corner of Data Table 2.)
Are these two values similar?
Why?
35
37
7. There are two possible isotopes of Chlorine: Cl and Cl. Based on the average atomic mass found on the
periodic table, which isotope of chlorine is more abundant in nature? Explain why you think so.
63
65
8. The element copper is present in two isotopes: Cu and Cu. The percent abundance of each isotope is listed
below. Use these percentages to calculate the average atomic mass of copper.
Isotope
63
Cu
65
Cu
% Abundance
73%
27%
Physical Science: Matter: Periodicity
COURSE: Chemistry
CONCEPT: SYSTEMS
ENDURING UNDERSTANDING: A system is an organized collection of parts, which display observable and predictable
patterns.
Essential Understandings
Assessment Samples – SOL/Blooms
Vocabulary
Students should understand:
• Periodicity is regularly
repeating patterns or trends
in the chemical and physical
properties of the elements
arranged in the periodic
table.
• Horizontal rows called
periods have predictable
properties based on an
increasing number of
electrons in the outer
orbitals.
• Vertical columns called groups
have similar properties
because of their similar
valence electron
configurations.
Knowledge/Comprehension Level
• Students are given a periodic table and they should be able to identify the
following features: periods, groups, alkali metal family, alkaline earth
family, transition metals, halogens, and noble gases.
Application/Analysis Level
•
Students should understand that if the formula of one compound is known,
then the formulas of similar compounds are predictable. For example, if the
formula of water is H2O, then the formulas of the other Group 16 hydrides
should be H2S, H2Se, and H2Te.
Synthesis/Evaluation Level
•
Students are given a set of cards that represent fictional elements. They
must arrange them in order of increasing atomic mass (just as Dimitri
Mendeleev did back in 1869). Students should look for similarities among
the elements so they can be classified into groups/families. The cards will
be set up so that there are two missing elements. These should create gaps
in their table. Just like Mendeleev, students should be able to predict
several properties of the missing elements, based on the existing trends.
(See Appendix A)
periodic law
periods
groups
alkali metals
alkaline earth metals
transition metals
halogens
noble gases
metals
nonmetals
metalloids (semimetals)
atomic radius
ionization energy
electronegativity
shielding effect
electron configuration
valence electrons
SOL: CH.2 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) series and periods;
(f) trends including atomic radii, electronegativity, shielding effect, and ionization energy.
August 1, 2009
SOL CH.2def – Appendix A
Physical Science: Matter: Periodicity
COURSE: Chemistry
CONCEPT: SYSTEMS
ENDURING UNDERSTANDING: A system is an organized collection of parts, which display observable and predictable
patterns.
Essential Understandings
Students should understand:
• Electron configuration is the
arrangement of electrons
around the nucleus of an
atom based on their energy
level.
• Atoms can gain, lose, or
share electrons within the
outer energy level.
Assessment Samples – SOL/Blooms
Vocabulary
Knowledge/Comprehension Level
• Students should be able to write the full electron configuration AND the
noble gas abbreviated configuration for any element from hydrogen (Z=1)
through krypton (Z=36).
electron configuration
orbitals (s, p, d, f)
quantum numbers
energy levels
Aufbau Principle
Pauli Exclusion
Principle
Hund’s Rule
cation
anion
oxidation state
Example: Titanium = 1s22s22p63s23p63d24s2 or [Ar] 3d24s2
Application/Analysis Level
•
Students should be able to recognize if a given electron energy diagram is
correct or incorrect. If the diagram is incorrect, they should be able to
identify which rule or principle (aufbau, Pauli, Hund) has been violated.
(See Appendix A)
Synthesis/Evaluation Level
•
Students should be able to give examples of several ions that are
isoelectronic with a given noble gas. They should also be able to identify an
element from its charge and configuration.
Example: Name six ions that are isoelectronic with neon
(N-3, O-2, F–, Na+, Mg+2, and Al+3)
Example: A certain element, M, has the following formula for its
oxide: M2O3. The M ion is isoelectronic with neon.
What is the identify of element M? (Aluminum)
SOL: CH.2 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
(g) electron configurations, valence electrons, and oxidation numbers.
August 1, 2009
SOL CH.2g – Appendix A
Identify the rule (aufbau principle, Hund’s rule, or Pauli exclusion principle), if any, that has been violated in each of the following electron energy
diagrams. If no rule has been violated, then state that it is correct as written.
3p
3s
3p
3s
2p
3p
3s
2p
3p
3s
2p
2p
2s
2s
2s
2s
1s
1s
1s
1s
Physical Science: Matter: Structure
COURSE: Chemistry
CONCEPT: SCALE: Properties
ENDURING UNDERSTANDING: Properties characterize objects, organisms, and substances.
Essential Understandings
Students should understand:
• Matter is classified by its
chemical and physical
properties.
• Physical properties refer to
the condition or quality of a
substance that can be
observed or measured
without changing the
substance’s composition.
• Chemical properties refer
to the ability of a
substance to undergo
chemical reaction and form
a new substance.
Assessment Samples – SOL/Blooms
Vocabulary
Knowledge/Comprehension Level
• Students should be able to classify a change as a physical change (e.g., ripping
paper, melting ice) or a chemical change (e.g., burning paper, electrolysis of
water) and explain their reasoning. Physical changes do not alter the chemical
identity of the substances involved. Chemical changes involve a change in
chemical identity and the formation of a new chemical substance. (See
Appendix A)
matter
mixture
pure substance
homogeneous
heterogeneous
compound
element
physical property
chemical property
density
conductivity
melting point
boiling point
malleability
ductility
reactivity
Application/Analysis Level
•
Students should be able to classify examples of matter as mixture vs. pure
substance. If something is a mixture, students should be able to identify it as
homogeneous or heterogeneous. If something is a pure substance, students
should identify it as an element or a compound. If a mixture is separated or a
compound is decomposed, students should know if that would be a physical or
chemical change. (See Appendix B)
Synthesis/Evaluation Level
•
Before discussing the concept of mixture and pure substance, students will
discover the defining characteristics of these concepts on their own, through
the concept-attainment model. Examples and non-examples are revealed to
the students one at a time.
o Examples (mixtures): coffee, tea, milk, gasoline, soda, steel, saltwater,
blood, air
o Non-examples (pure substances): gold, silver, iron, oxygen, nitrogen,
water, ethanol, sucrose, salt, carbon dioxide
Students will brainstorm to come up with ideas about what all the examples
have in common. Students’ ideas will be evaluated continually as additional
examples and non-examples are revealed.
SOL: CH.2 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.
August 1, 2009
SOL CH.2h – Appendix A
Laboratory Activity – Physical and Chemical Changes
Part 1 – Signs of a Chemical Change
There are FOUR basic observations that may suggest that a chemical reaction may have occurred.
•
•
•
•
Change in color (out of nowhere)
Formation of a gas
Change in energy (heat, light, sound)
Formation of a precipitate
Add a few drops of phenolphthalein to the beaker that contains sodium hydroxide.
1.
What do you observe?
Now add a few drops of hydrochloric acid to the beaker until you see a color change.
2.
What do you observe?
Pour the contents of the beaker in the sink and rinse out the beaker with water.
Add a piece of magnesium metal ribbon to the test tube that contains hydrochloric acid.
Cover the test tube with another empty test tube (upside-down) to trap the gas.
Hold the upside-down test tube in place.
When all of the fizzing has stopped, call your teacher to put a flame to the gas.
3.
What did you observe when the magnesium was added to the acid?
4.
What did you observe when the gas in the upside-down test tube was lit?
5.
Feel the outside of the test tube with the acid in it. How does it feel?
Add about 3 drops of lead(II) nitrate to a clean, empty beaker.
Now add about 3 drops of potassium iodide to the same beaker.
6.
What do you observe?
7.
All of the changes described in this section are classified as
new substance was produced.
changes because a
Pour the contents of the beaker in the sink and rinse out the beaker with water.
Part 2 – Phase changes
1.
There are six phase changes that you need to memorize. They are listed below. Write the name of the phase
change in each blank.
solid ! liquid
liquid ! solid
liquid ! gas
gas ! liquid
solid ! gas
gas ! solid
2.
Look at the beaker with ice cubes and water in it. What do you observe on the outside of the beaker?
3.
How do you suppose those drops of water got there?
Now place the beaker on the hot plate.
4.
What will happen to the ice cubes over time?
5.
Once all of the ice has melted, what will happen to the water over time ?
Add a piece of dry ice to a flask. Stretch a balloon over the flask. Place the flask on the hot plate.
6.
What do you observe happening to the balloon?
Your teacher will demonstrate phase changes involving iodine crystals.
7.
When the iodine crystals change from a solid to a gas, this process is called
8.
When the iodine vapor is changed from a gas to a solid, this process is called
9.
All of the changes described in this section are phase changes. Phase changes are all classified as
changes, because the chemical identity of the substance is
.
Part 3 – Physical or Chemical?
For each of the following activities, you should check the boxes that apply, and then decide if the change is
classified as PHYSICAL or CHEMICAL.
Activity
Take a piece of aluminum
foil and rip it into smaller
pieces.
Add the aluminum pieces to
a beaker, and then add some
copper chloride solution to it.
Squirt some methanol on the
lab counter and wait for it to
disappear.
Squirt some methanol on the
lab counter and call your
teacher to put a flame near
the liquid.
Place an antacid tablet in
the mortar and crush it with
the pestle.
Add some water to the
powder in the mortar.
Add about 3 drops of silver
nitrate to a clean beaker.
Then add about 3 drops of
copper chloride to the beaker.
TEACHER DEMO:
Solid potassium chlorate is
heated in a test tube until it
turns into a liquid.
TEACHER DEMO:
A piece of candy is added to
the molten potassium
chlorate in the test tube.
Did you see
a color change?
Did you see the
formation
of a gas?
Did you
notice a
change
in energy?
Did you
see the
formation
of a
precipitate?
Did you
see a
phase
change?
Is this a
PHYSICAL
or a
CHEMICAL
change?
SOL CH.2h – Appendix B
Laboratory Activity – Matter
Part 1 – Classifying Matter
At each lab station, you will see three samples of matter. Your job is to classify each sample as either
a MIXTURE, a COMPOUND, or an ELEMENT. Each lab station has one of each type of matter represented.
Station #1
Station #7
vinegar solution =
copper sulfate
salt (sodium chloride) =
hydrogen peroxide solution
copper =
nitrogen
Station #2
Station #8
sugar (sucrose) =
oil and water
aluminum =
ammonium nitrate
sedimentary rock =
carbon
Station #3
Station #9
zinc =
hexane (C6H14)
carbon dioxide =
air
bronze =
helium
Station #4
Station #10
potassium iodide =
soda and spaghetti pieces
pepper and water =
oxygen
hydrogen =
cobalt chloride
Station #5
Station #11
water =
calcium carbonate
calcium =
green ink
Gatorade =
magnesium
Station #6
Station #12
iodine crystals =
copper
baking soda (sodium bicarbonate) =
methane (CH4)
iron and sulfur =
chicken noodle soup
1.
A MIXTURE is defined as a sample of matter that represents a physical combination of more than one substance.
Fill in the blanks below with the names of the samples of matter from each lab station that you classified as a
MIXTURE:
#1
#5
#9
#2
#6
#10
#3
#7
#11
#4
#8
# 12
Mixtures are classified as either HOMOGENEOUS or HETEROGENEOUS. What’s the difference between these terms?
A HOMOGENEOUS MIXTURE appears to be uniform throughout. In other words, it appears to look like just one phase or
one substance. On the other hand, in a HETEROGENEOUS MIXTURE you can see and identify more than one
component or phase.
2.
Sort all 12 mixtures into one of the following categories:
HOMOGENEOUS MIXTURES
3.
HETEROGENEOUS MIXTURES
A COMPOUND is defined as a sample of matter that is formed from a chemical combination of more than one
element. Fill in the blanks below with the names of the samples of matter from each lab station that you classified
as a COMPOUND:
#1
#5
#9
#2
#6
#10
#3
#7
#11
#4
#8
# 12
It is very important that you understand the difference between a COMPOUND and a MIXTURE.
A COMPOUND has elements that are chemically bonded together.
A MIXTURE has components that are only physically combined together.
4.
If you wanted to separate a mixture into simpler substances, you could do it with a
change. On the other hand, if you wanted to break down a compound into simpler substances, you would need to
use a
change.
5.
An ELEMENT is defined as a sample of matter that CANNOT be broken down or decomposed by either a
physical or a chemical change. Fill in the blanks below with the names of the samples of matter from each lab
station that you classified as an ELEMENT:
#1
#5
#9
#2
#6
#10
#3
#7
#11
#4
#8
# 12
6.
You should realize that all of the substances listed above can be found on the
Part 2 – Separation Techniques
1.
Paper Chromatography
You are going to be given a piece of paper with a line of green ink on it. You should lower the filter paper into the
beaker of liquid so that only the bottom edge of the paper is touching the water. Over time, the water will rise up
through the paper and separate the green ink into smears of blue and yellow.
2.
a.
Is the green ink from the marker classified as a compound or a mixture?
b.
In this experiment, are we producing a new substance that wasn’t there before?
c.
Is paper chromatography classified as a physical change or a chemical change?
Filtration
You should have a plastic cup with some water and pepper in it. You are going to use a piece of filter paper to
separate the pepper from the water.
a.
Is pure water (H2O) classified as a compound or a mixture?
b.
Is the sample of water and pepper classified as a compound or a mixture?
c.
If you said “mixture” for question 2b., do think that water and pepper represents a homogeneous mixture or a
heterogeneous mixture?
How can you tell?
Now you are going to pour the water and pepper mixture through a funnel lined with filter paper.
d.
Describe the appearance of the filter paper after all of the water has passed through it.
e.
Do you think the liquid that came through the filter paper should be classified as a compound or a mixture?
Explain why.
f.
Filtration is a separation technique that is based on particle size. Do you think filtration should be classified
as a physical change or a chemical change?
3
Distillation
Your teacher will perform this experiment as a demonstration.
a.
Is solid copper sulfate (CuSO4) classified as a compound or a mixture?
b.
Is the combination of water and copper sulfate classified as a homogeneous mixture or a heterogeneous
mixture?
c.
How can you tell?
In distillation, the mixture is heated until it begins to boil. Which substance, water or copper sulfate, is going
to appear in the collection flask first?
d.
Distillation is a technique used to separate two substances that have different
e.
When we use distillation to separate copper sulfate from water, are we producing a new substance that
wasn’t there before?
f.
4.
Is distillation classified as a physical change or a chemical change?
Electrolysis of water
Your teacher will perform this experiment as a demonstration.
a.
Is pure water (H2O) classified as a compound or a mixture?
b.
During electrolysis of water, bubbles of gas are formed in
two separate tubes. What two gases are being collected?
and
c.
When we use electrolysis to break water down into
its elements, are we producing new substances that
weren’t there before?
d.
Is electrolysis classified as a physical change
or a chemical change?
e.
The diagram at right shows a set-up for the electrolysis of water.
The tube on the LEFT contains
gas, and
the tube on the RIGHT contains
gas.
Part 3 – Summary Questions
1.
2.
Classify each of the following as COMPOUND, ELEMENT, HETEROGENEOUS MIXTURE, or HOMOGENEOUS
MIXTURE
green ink
copper
water
carbon dioxide
hydrogen
bronze
sand & water
oxygen
air
saltwater
sugar
salt & sand
Fill in the following “MATTER” chart with the words COMPOUND, ELEMENT, HETEROGENEOUS MIXTURE,
HOMOGENEOUS MIXTURE, MATTER, and PURE SUBSTANCE.
MIXTURE
3.
Green ink can be separated easily by using
4.
Sand & water can be separated easily by using
5.
Saltwater can be separated easily by using
6.
Water can be decomposed into hydrogen and oxygen by using
7.
MIXTURES can be separated by
changes.
8.
COMPOUNDS can be separated by
changes.
9.
CANNOT be separated by either physical or chemical means.
Physical Science: Matter: Structure
COURSE: Chemistry
CONCEPT: COMMUNICATION: Theory
ENDURING UNDERSTANDING: Theories explain why natural phenomena occur, and they evolve to incorporate new knowledge.
Essential Understandings
Students should understand:
• Discoveries and insights
related to the atom’s
structure have changed the
model of the atom over
time.
• The modern atomic theory is
called the Quantum
Mechanical Model.
Assessment Samples – SOL/Blooms
Vocabulary
Knowledge/Comprehension Level
• Students can prepare a timeline that shows the important scientists who
contributed to atomic theory, including a few details about their specific
contribution. Scientists should include Democritus, John Dalton, J. J.
Thomson, Ernest Rutherford, Robert Millikan, Niels Bohr, Dimitri
Mendeleev, Henry Moseley, Max Planck, Werner Heisenberg, and Louis de
Broglie.
atom
electron
cathode ray tube
nucleus
planetary model
periodic table
quantum theory
uncertainty principle
particle-wave duality
Application/Analysis Level
•
Students will observe a “think tube” which has various ropes extending
from the sides. When a single rope is pulled as part of the teacher’s
demonstration, students must draw a diagram of what they think the
inside of the tube looks like. As more ropes are pulled, students are
challenged to explain a few seemingly impossible results and then modify
their diagrams accordingly. Students then can try to construct their own
think tube that will behave similarly to the teacher’s model.
Synthesis/Evaluation Level
•
Students can design their own mystery box, which contains a simple
geometric shape (e.g., triangle, square, circle) that is sandwiched between
two larger pieces of wood. The only way that the identity of the hidden
shape can be revealed is that other students can shoot marbles at the
mystery shape from different angles. Based on the deflection pattern of the
marbles, students try to guess what shape is inside the box. This activity is
somewhat parallel the Rutherford gold foil experiment.
SOL: CH.2 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.
August 1, 2009
Physical Science: Matter: Reactions
COURSE: Chemistry
CONCEPT: COMMUNICATION: Model
ENDURING UNDERSTANDING: Models vary in complexity and facilitate understanding through the use of familiar concepts.
Essential Understandings
Students should understand:
• Conservation of matter is
represented in balanced
chemical equations.
• Chemical formulas are used
to represent compounds.
• Subscripts represent the
relative number of each type
of atom in a molecule or
formula unit.
• A coefficient is a quantity
that precedes a reactant or
product symbol or formula in
a chemical equation and
indicates the relative number
of particles involved in the
reaction.
• Bonds form between atoms
to achieve stability.
Assessment Samples – SOL/Blooms
Knowledge/Comprehension Level
• Students should know the formula and charge of monoatomic ions (e.g.,
Na+, Mg+2, F–, O-2) based on their location on the periodic table. (See
Appendix A)
• Students should know the formula and charge of the following polyatomic
ions: ammonium (NH4+), hydroxide (OH–), nitrate (NO3–), carbonate
(CO32-), sulfate (SO42-), and phosphate (PO43-).
Application/Analysis Level
•
Students should be able to write correct chemical formulas for molecular
compounds, ionic compounds, and acids if they are given the name (and
vice versa). They should also be able to identify the oxidation state of
certain elements within a chemical formula. (See Appendix B)
Synthesis/Evaluation Level
•
•
•
Students are given several chemical formulas that refer to a mystery
element. They are also given information about the electron configuration
of the mystery element when it becomes an ion. They must identify the
element (See Appendix C)
Students are given an incorrect Lewis dot structure. They must identify
what is wrong with the structure and draw a correct Lewis structure. (See
Appendix C)
Students are given a molecular formula that contains a mystery element
and the total number of valence electrons for the formula. They must draw
the Lewis dot structure, identify the molecular geometry, and identify the
mystery element. (See Appendix C)
Vocabulary
Law of Multiple
Proportions
empirical formula
molecular formula
structural formula
covalent bond
ionic bond
ionization energy
electronegativity
polar
nonpolar
binary covalent
binary ionic
monoatomic ions
polyatomic ions
Lewis dot diagrams
bonding pair
nonbonding pair
(lone pair)
single bond
double bond
triple bond
tetrahedral
pyramidal
bent
trigonal planar
linear
SOL: CH.3 The student will investigate and understand how conservation of energy and matter is expressed in chemical formulas and
balanced equations.
(a) nomenclature;
(b) balancing chemical equations;
(c) writing chemical formulas (molecular, structural, empirical, and Lewis diagrams);
(d) bonding types (ionic, covalent).
August 1, 2009
SOL CH.3abcd – Appendix A
At the top of each column, write the charge on the ion (oxidation state) that would be formed when
these elements undergo chemical reactions. While the value for oxidation states can certainly vary,
consider what the element would have to do to achieve a noble gas configuration.
Your choices are +1, -1 +2, -2, +3, and -3
Common Oxidation States
+4
H
Li
Na
K
Rb
Cs
Be
Mg
Ca
Sr
Ba
B
Al
Ga
In
Tl
C
Si
Ge
Sn
Pb
0
N
P
As
Sb
O
S
Se
Te
F
Cl
Br
I
He
Ne
Ar
Kr
Xe
Rn
SOL CH.3abcd – Appendix B
1. Binary Ionic Compounds
sodium oxide
NiBr3
magnesium chloride
Ba3P2
iron(III) sulfide
Al2O3
potassium nitride
CoF2
calcium phosphide
SrI2
2. Ionic Compounds with Polyatomic Ions (You need to memorize those polyatomic ions!)
ammonium acetate
Na2CO3
potassium chlorate
Li2CrO4
magnesium hydrogen carbonate
FeSO4
aluminum hydroxide
Cu3(PO4)2
calcium nitrate
Ba(CN)2
3. Binary Molecular Compounds
carbon diselenide
PBr5
sulfur trioxide
SeF6
silicon tetrachloride
ICl7
4. Acids
hydrofluoric acid
H2SO4
nitric acid
HC2H3O2
5. Oxidation Numbers – Write the oxidation number of the element listed in bold type
H2SO4
F2
N H3
SCl2
ClO3–
PCl3
HBr
PF5
HOCl
Cr2O7-2
SOL CH.3abcd – Appendix C
A certain element, X, forms the following ionic compounds: XCl2, XO, and X3N2
When an atom of element X becomes a stable ion, it has the electron configuration 1s22s22p63s23p6.
a) Based on the given information, is element X classified as a metal or a nonmetal?
b) When element X becomes a stable ion, what is the charge?
(Don’t just say “positive” or “negative.” I want the sign and the number, like “+3” or “-2”
c) What is the identity of element X?
A student tried to write a Lewis dot structure for the molecule C3H3N.
The student wrote a skeleton structure that was correct, but the arrangement of electrons was
incorrect. The incorrect Lewis dot structure is shown on the left.
a) How many valence electrons should be used for the molecule C3H3N?
b) In the left box, circle all of the atoms (other than hydrogen) that do NOT have access to a full
octet of electrons.
c) In the right box, draw the correct Lewis dot structure for C3H3N.
You should keep the same skeleton arrangement of atoms as the structure on the left.
The number of electrons in your structure should be the same as the number you mentioned in
part a). Make sure that all atoms (except hydrogen) have access to a full octet.
INCORRECT Lewis Dot Structure
CORRECT Lewis Dot Structure
Consider the following information, as summarized in the table below. Note that the letters A, D,
G, and J are used to represent elements from the periodic table. The letters O and F represent
atoms of oxygen and fluorine.
Molecule
Total Number of
Valence
Electrons
Additional Information
AO2
16
none
DO2
18
none
GF3
24
JF3
26
The Lewis structure
contains ALL single
bonds
The Lewis structure
contains ALL single
bonds
(a) For each molecule listed in the table above, identify the molecular geometry/shape.
You might want to sketch the Lewis electron dot structure on your own paper.
AO2 =
DO2 =
GF3 =
JF3 =
(b) The molecules AO2, DO2, GF3, and JF3 represent covalent compounds.
Based on the information in the table above, identify the elements A, D, G, and J.
(More than one correct answer may be possible for each letter, but you only have to give ONE
possible answer for each letter.)
A=
D=
G=
J=
Physical Science: Matter: Reactions
COURSE: Chemistry
CONCEPT: CHANGE AND CONSTANCY: Cause and Effect
ENDURING UNDERSTANDING: Observable changes occur in nature, and inferences can be made to explain their causes.
Essential Understandings
Students should understand:
• Elements and compounds
react in different ways.
• Spontaneous reactions may
be fast or slow.
• Randomness (entropy), heat
content (enthalpy), and
temperature affect
spontaneity.
• Reaction rates/kinetics are
affected by activation
energy, catalysis, and the
degree of randomness
(entropy).
Assessment Samples – SOL/Blooms
Knowledge/Comprehension Level
• Students should be able to balance an equation and identify the type of
chemical reaction (See Appendix A)
Application/Analysis Level
•
Students are given an equilibrium system. They should apply LeChatelier’s
principle to predict the direction (if any) that the system will shift as a
result of various changes to the system. For example: Consider an
exothermic reaction. In which direction will the equilibrium position shift
when the temperature of the system is increased? (Answer: to the left.)
Synthesis/Evaluation Level
•
If students are given a sentence describing the chemicals that react
together, they should be able to identify the type of reaction AND write a
complete balanced chemical equation for that process. (See Appendix A)
Vocabulary
synthesis
decomposition
single replacement
double replacement
neutralization
oxidation
reduction
endothermic
exothermic
equilibrium
Le Chatelier’s
Principle
activation energy
catalyst
rate of reaction
SOL: CH.3 The student will investigate and understand how conservation of energy and matter is expressed in chemical formulas and
balanced equations.
(e) reaction types (synthesis, decomposition, single and double replacement, oxidation-reduction, neutralization, exothermic and
endothermic);
(f) reaction rates and kinetics (activation energy, catalysis, degree of randomness).
August 1, 2009
SOL CH.3ef – Appendix A
Identify the TYPE of reaction that is represented in each diagram. Use the following choices:
SYNTH
DECOMP
synthesis
decomposition
SR
DR
single replacement
double replacement
!
+
1)
H
+
X
M
+
!
+
B
OH
C
+
M
X
C
B
+
!
+
5)
A
H
!
+
4)
6)
OH
combustion
neutralization
+
!
2)
3)
COMB
NEUTRAL
D
!
A
+
D
BALANCE each reaction using the lowest possible whole numbers. Identify the TYPE of reaction. Use the following
choices.
SYNTH
DECOMP
synthesis
decomposition
SR
DR
single replacement
double replacement
7)
8)
9)
Fe(OH)3
Ca(C2H3O2)2
+
!
(NH4)2CO3 !
COMB
NEUTRAL
combustion
neutralization
Fe2O3
+
H2 O
CaCO3
+
NH4C2H3O2
Ni
+
Cl2
!
NiCl3
10)
C3 H6 O
+
O2
!
CO2
+
H2 O
11)
F2
+
ScBr3
!
ScF3
+
Br2
+
H2SO4 !
Al2(SO4)3
+
H2 O
12)
Al(OH)3
WRITE THE CORRECT FORMULAS for each substance listed. BALANCE each reaction using the lowest possible whole
numbers. Identify the TYPE of reaction. Use the following choices.
SYNTH
DECOMP
synthesis
decomposition
SR
DR
single replacement
double replacement
13)
octane (C8H18)
+
14)
diphosphorus pentoxide
15)
hydrochloric acid
16)
aluminum
17)
potassium chlorate
18)
iron(III) bromide
+
+
oxygen
+
!
+
carbon dioxide
water
magnesium hydroxide
lead(II) nitrate
!
!
COMB
NEUTRAL
!
!
water
phosphoric acid (H3PO4)
!
lead
potassium chloride
sodium sulfate
+
combustion
neutralization
magnesium chloride
+
+
water
aluminum nitrate
+
iron(III) sulfate
oxygen
+
sodium bromide
Physical Science: Matter: Structure
COURSE: Chemistry
CONCEPT: SCALE: Measurement
ENDURING UNDERSTANDING: Measurement represents properties on a numerical scale.
Essential Understandings
Students should understand:
• Atoms and molecules are too
small to count by usual
means.
• A mole is a way of counting
any type of particle (atoms,
molecules, and formula
units).
• Stoichiometry involves
quantitative relationships.
• Stoichiometric relationships
are based on mole quantities
in a balanced equation.
Assessment Samples – SOL/Blooms
Vocabulary
Knowledge/Comprehension Level
• Students should know that there are 6.02 x 1023 particles in one mole.
They should know that 1 mole of gas at standard temperature and
pressure (STP) has a volume of 22.4 L.
• Students should know how to use the periodic table to calculate the molar
mass of a substance by using the atomic masses and adding up the total
mass of all of the atoms that are found in the chemical formula.
Avogadro’s number
mole
average atomic mass
molar mass
molar volume
STP
reactants
products
limiting reactant
percent yield
Application/Analysis Level
•
If given a balanced chemical reaction, students should be able to do the
following conversions: mole-mole, mole-mass, mass-mass, and volumevolume. (See Appendix A)
Synthesis/Evaluation Level
•
Students are given the amounts of each reactant in a chemical equation,
and they are to do calculations to determine the limiting reactant,
theoretical yield, and the percent yield (if given the actual yield). (See
Appendix B)
SOL: CH.4 The student will investigate and understand that quantities in a chemical reaction are based on molar relationships.
(a) Avogadro's principle and molar volume;
(b) stoichiometric relationships.
August 1, 2009
SOL CH.4a,b – Appendix A
Moles to Moles:
KClO3
1
!
KCl
+
O2
(a) Balance the equation above.
(b) If you started with 5.17 moles of potassium chlorate, how many moles of oxygen gas
would be produced?
Grams to Moles:
Mg
2
+
H3PO4
!
Mg3(PO4)2
+
H2
(a) Balance the equation above.
(b) How many moles of magnesium metal would be needed to react completely with
37.8 grams of phosphoric acid?
Grams to Grams:
C4H10
3
+
O2
!
CO2
+
H 2O
(a) Balance the equation above.
(b) How many grams of oxygen gas are required to react with 74.2 grams of butane (C4H10)?
SOL CH.4a,b – Appendix B
Limiting Reactant:
K3PO4
+
Ca(OH)2
!
KOH
+
Ca3(PO4)2
1. Balance the equation above.
2. Suppose that you reacted 41.5 grams of potassium phosphate with
32.4 grams of calcium hydroxide.
(a) Determine which reactant is the limiting reactant, based on the product calcium phosphate.
(b) Calculate the theoretical yield (in grams) of calcium phosphate.
(c) When this experiment was performed (as described above), the mass of recovered
calcium phosphate was equal to 24.7 grams. Calculate the percent yield for the reaction.
Physical Science: Matter: Structure
COURSE: Chemistry
CONCEPT: SYSTEMS
ENDURING UNDERSTANDING: A system is an organized collection of parts, which display observable and predictable patterns
Essential Understandings
Students should understand:
•
•
•
•
•
•
•
•
Gases have mass and occupy
space.
Gas particles are in constant,
rapid, random motion and exert
pressure as they collide with the
walls of their containers.
Gas molecules with the lightest
mass travel fastest.
Relatively large distances
separate gas particles from
each other.
An Ideal Gas does not exist, but
this concept is used to model
gas behavior.
A Real Gas exists, has
intermolecular forces and
particle volume, and can change
states.
Equal volumes of gases at the
same temperature and pressure
contain an equal number of
particles.
Solutions can be a variety of
solute/solvent combinations:
gas/gas, gas/liquid,
liquid/liquid, solid/liquid,
gas/solid, liquid/solid, or
solid/solid.
Assessment Samples – SOL/Blooms
Knowledge/Comprehension Level
• Students should be able to write the basic postulates of the kinetic
molecular theory.
• Students should know the gas laws (Boyle, Charles, Gay-Lussac, ideal) and
understand how to use them to solve for a particular variable.
Application/Analysis Level
•
Students should be able to explain why each of the following events occur,
in terms of the KMT.
o A balloon will expand when heated.
o Pressure inside a sealed can will increase when the can is heated.
o A balloon will expand when taken to a higher elevation (at constant
temp.)
o Pressure inside a sealed can will increase when the can is squashed.
Synthesis/Evaluation Level
•
•
Students can perform an experiment that involves the calculation of the
value of the ideal gas constant, R. (See Appendix A)
Students can collect a sample of butane gas by water displacement. Based
on the data collected, students can calculate the molar mass of the butane
gas.
Vocabulary
pressure
volume
temperature
kinetic-molecular
theory
Boyle’s Law
Charles’ Law
Gay-Lussac’s Law
Dalton’s Law of
partial pressures
Ideal Gas Law
ideal gases
real gases
solution
concentration
solute
solvent
molarity
SOL: CH.4 The student will investigate and understand that quantities in a chemical reaction are based on molar relationships.
(c)
(d)
(e)
(f)
partial pressure;
gas laws;
solution concentrations;
chemical equilibrium.
August 1, 2009
SOL CH.4c,d,e,f – Appendix A
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Physical Science: Matter: Structure
COURSE: Chemistry
CONCEPT: SCALE: Properties
ENDURING UNDERSTANDING: Properties characterize objects, organisms, and substances.
Essential Understandings
Students should understand:
• Two important classes of
compounds are acids and
bases.
• Acids and bases are defined
by several theories.
• Acids and bases dissociate in
varying degrees.
Assessment Samples – SOL/Blooms
Knowledge/Comprehension Level
• Students can estimate the pH of various solutions (e.g., water, vinegar,
ammonia, lemon juice, drain opener, etc.) by using colorful acid-base
indicators and/or pH paper.
Application/Analysis Level
•
Students can write balanced chemical equations for neutralization
reactions. For example, hydrochloric acid in the stomach is neutralized with
milk of magnesia (magnesium hydroxide).
2 HCl + Mg(OH)2 ! MgCl2 + 2 H2O
Synthesis/Evaluation Level
•
Students can determine the molarity of household vinegar by performing a
titration experiment with a standardized NaOH solution. (See Appendix A)
Vocabulary
Arrhenius definition
Bronsted-Lowry
definition
acid
base
pH
pOH
molarity
strong electrolyte
weak electrolyte
dissociation
titration
SOL: CH.4 The student will investigate and understand that quantities in a chemical reaction are based on molar relationships.
(g) acid/base theory: strong electrolytes, weak electrolytes, and nonelectrolytes; dissociation and ionization; pH and pOH; and the
titration process.
August 1, 2009
SOL CH.4g – Appendix A
Standardization of NaOH with KHP (potassium hydrogen phthalate, KHC8H4O4)
Procedure:
1. Measure out approximately 1.00 gram of the solid powder KHP. Record the mass of KHP in
the data table below.
2. Transfer the KHP into a 500 mL flask. Rinse the plastic weighing tray with distilled water to
make sure that it is completely transferred.
3. Add approximately 100 mL of distilled water to the flask.
4. Add about 5 drops of phenolphthalein solution to the flask.
5. Make sure the buret is filled with NaOH solution so that the volume is in between 0 and 5 mL.
6. Record the exact volume of the buret to two decimal places in the data table below.
7. Place the flask directly under the buret. Slowly add the NaOH and swirl to mix it.
8. As the pink color starts to persist, try to add the NaOH drop by drop.
9. When one drop of NaOH results in a pink color that does not fade, you have reached the
equivalence point in the titration.
10. Record the final volume on the buret to two decimal places.
11. Rinse out your flask several times with with water. You don’t have to dry it, but it should be
clean.
12. Repeat Steps 1-11. You may need to add more NaOH to the buret in step 5.
Trial 1
Trial 2
mass of KHP
(3 sig figs)
g
molar mass of KHP (KHC8H4O4)
g/mol
moles of KHP
(3 sig figs)
mol
mol
initial volume of NaOH in buret
(two decimal places)
mL
mL
final volume of NaOH in buret
(two decimal places)
mL
mL
volume of NaOH used
(two decimal places)
mL
mL
moles of NaOH used
mol
mol
liters of NaOH used
L
L
molarity of NaOH
mol/L
mol/L
Average value for Molarity of NaOH solution
g
mol/L
DETERMINING THE MOLARITY OF ACETIC ACID IN VINEGAR
In this experiment you will be titrating a vinegar solution with the sodium hydroxide that you have
standardized. Vinegar contains acetic acid (CH3CO2H). Sodium hydroxide is a base. You will be using
phenolphthalein indicator to help you visualize the endpoint. When all of the acid has been
neutralized by the base, the color of the solution will turn pink. You will do 2 trials. Based on your
results, you can calculate the molarity of the acetic acid.
Write a balanced equation for the neutralization of acetic acid with sodium hydroxide:
Trial 1
volume of vinegar
(3 sig figs)
10.0
Trial 2
mL
10.0
mL
initial volume of NaOH in buret
(two decimal places)
mL
mL
final volume of NaOH in buret
(two decimal places)
mL
mL
volume of NaOH used
(two decimal places)
mL
mL
liters of NaOH used
L
L
Molarity of NaOH
(from earlier experiment)
M
M
moles of NaOH
mol
mol
moles of vinegar
L
L
molarity of vinegar
mol/L
mol/L
Average value for Molarity of vinegar solution
mol/L
What is the percentage of acetic acid by mass in your vinegar sample? Let’s assume that the density
of vinegar is close to 1 g/mL. Plug in your molarity into the calculations below and find out.
mol CH3CO2H x
L
60 g CH3CO2H x
1L
x
1 mol CH3CO2H
1000 mL
1 mL
1g
x 100% =
Physical Science: Matter: Structure
COURSE: Chemistry
CONCEPT: COMMUNICATION: Theory
ENDURING UNDERSTANDING: Theories explain why natural phenomena occur, and they evolve to incorporate new knowledge.
Essential Understandings
Students should understand:
• Atoms and molecules are in
constant motion.
• The Kinetic Molecular Theory
is a model for predicting and
explaining gas behavior.
• Forces of attraction between
molecules determine the
physical changes of state.
• Vapor pressure is a property
of a substance determined
by intermolecular forces.
Assessment Samples – SOL/Blooms
Knowledge/Comprehension Level
• If students are given a vapor pressure diagram for a certain liquid, they
should be able to determine the normal boiling point of the liquid.
Application/Analysis Level
•
•
If students are given a vapor pressure diagram for several liquids, they
should be able to predict which liquid has the strongest/weakest
intermolecular forces. (See Appendix A)
If students are given a vapor pressure diagram, they should be able to
identify the boiling point for conditions when the pressure is reduced below
atmospheric pressure. (See Appendix A)
Vocabulary
kinetic-molecular
theory
pressure
temperature
volume
vapor pressure
intermolecular forces
London dispersion
forces
dipole-dipole forces
hydrogen bonding
Synthesis/Evaluation Level
•
Students should be able to analyze the structural formulas of various
substances and identify the type of intermolecular forces that each
substance has. They should predict which substance in a given pair has a
higher boiling point. (See Appendix B)
SOL: CH.5 The student will investigate and understand that the phases of matter are explained by kinetic theory and forces of attraction
between particles.
(a) pressure, temperature, and volume;
(b) vapor pressure.
August 1, 2009
SOL CH.5ab – Appendix A
Why is zinc a good conductor of electricity, but solid zinc chloride is not a good conductor of electricity. Explain your
answer by discussing the properties of each solid at the particle level.
Match these substances with their boiling points using your knowledge of intermolecular forces.
o
magnesium fluoride
A. –188 C
fluorine
B.
hydrogen fluoride
C. 1400 C
o
20 C
o
The diagram below shows how the temperature of a substance changes as it is being heated.
Label each portion of the graph with either a phase of matter or a phase change
Identify the letter that matches each of the following:
1.00 atm
B
D
C
Pressur
e
E
solid
liquid
vapor
A
triple point
F
G
critical point
normal MP
Temperatur
e
normal BP
Based on the phase diagram below, answer the following questions:
2.00
atm
1.00
atm
`
o
a.
At –10 C and 0.50 atm, what phase
of matter would this substance be in?
b.
If the temperature were to then increase to
o
30 C at constant pressure, what phase
change would occur?
c. If the pressure were to then increase
to 1.5 atm at constant temperature,
what phase change would occur?
0.50
atm
d.
If the temperature were to then
o
decrease to 0 C at constant pressure,
what phase change would occur?
e.
From the graph, estimate the normal
BP of this substance:
-20 -10 0 10 20 30 40 50 60 70 80 90
Temp (oC)
Based on the vapor pressure curves above, answer the following questions:
a. Which liquid has the highest vapor pressure at room temperature?
b. Which liquid has the highest boiling point?
c. Which liquid experiences the weakest intermolecular forces?
d. If only one of these liquids experiences hydrogen bonding, which one would it be?
o
e. If liquid B is at a temperature of 40 C, what would the equilibrium vapor pressure be?
f. What is the normal boiling point of liquid C?
SOL CH.5ab – Appendix B
For each substance below, indicate whether the substance is polar or nonpolar. Then indicate the type of intermolecular
forces
that is most appropriate. Use the following abbreviations:
LDF = London dispersion forces
Substance
Polar or
Nonpolar?
H2
HCl
H 2O
HF
Type
of IMF
DD = dipole-dipole forces
Substance
Polar or
Nonpolar?
CO
CO2
CH3OH
CH3F
Type
of IMF
HB = hydrogen bonding
Substance
Ne
NH3
PH3
C 3H 8
For each pair of substances below, circle the one that should have the higher boiling point:
a.)
Ne
b.)
CH4
c.)
N2
Xe
C4H10
CO
d.)
HF
e.)
NH3
HBr
PH3
Polar or
Nonpolar?
Type
of IMF
Physical Science: Force, Motion, and Energy: Energy
COURSE: Chemistry
CONCEPT: SCALE: Properties
ENDURING UNDERSTANDING: Properties characterize objects, organisms, and substances.
Essential Understandings
Students should understand:
• Solid, liquid, and gas phases
of a substance have different
energy content.
• Specific amounts of energy
are absorbed or released
during phase changes.
• Specific heat capacity is a
property of a substance.
• Polar substances dissolve
ionic or polar substances;
nonpolar substances dissolve
nonpolar substances.
• The number of solute
particles changes the
freezing point and boiling
point of a pure substance.
Assessment Samples – SOL/Blooms
Vocabulary
Knowledge/Comprehension Level
• Students should know that likes dissolve likes. For example, they should
know that water is polar and that polar solutes will dissolve in water,
whereas nonpolar solutes do not dissolve in water. Likewise, two nonpolar
substances (e.g., turpentine and gasoline) should be very miscible with
each other.
solid
liquid
gas
specific heat capacity
polar
nonpolar
freezing point
boiling point
heat of fusion
heat of vaporization
phase diagram
vapor pressure
triple point
heating curve
solute
solvent
Application/Analysis Level
•
•
•
Students can do calorimetric calculations to calculate the amount of heat
required to raise the temperature of a certain mass of water by a certain
number of degrees Celsius. (See Appendix A)
If given the molar heat of fusion (or molar heat of vaporization) for a
certain substance, they can calculate the amount of heat required to melt
(or evaporate) a certain quantity of that substance. (See Appendix A)
Students can calculate the molar heat of fusion for ice (See Appendix B)
Synthesis/Evaluation Level
•
Students can record the mass of water in a Styrofoam cup. Then they can
dissolve a known mass of an ionic solid into the water and record the
change in temperature as the solute dissolves completely. Based on the
temperature change, the mass of the water, and the mass of the solute,
they can calculate the molar enthalpy of solution.
SOL: CH.5 The student will investigate and understand that the phases of matter are explained by kinetic theory and forces of attraction
between particles.
(c) phase changes;
(d) molar heats of fusion and vaporization;
(e) specific heat capacity;
(f) colligative properties.
August 1, 2009
SOL CH.5cdef – Appendix A
38. The molar heat of fusion for ice is 6.02 kJ/mol. Use this information to answer the following questions.
a.
How much heat energy is required to melt 125 grams of ice?
b.
How many grams of ice at 0 °C can be melted by the addition of 1500 J of heat?
SOL CH.5cdef – Appendix B
CALCULATING THE MOLAR HEAT OF FUSION OF ICE
Introduction
Water freezes at 0oC. Ice melts at the same temperature, 0oC. The amount of heat energy required
to melt one mole of ice at its melting point is known as the molar heat of fusion. The units for
the heat of fusion are normally given as kilojoules per mole (kJ/mol).
Most students are more familiar with the calorie (cal) than they are with the joule (J). One calorie is
defined as the amount of heat required to raise the temperature of 1 gram of water by 1 degree
Celsius. The conversion factor between calories and joules is as follows.
1 cal = 4.18 J
In this experiment you will melt a small quantity of ice in a Styrofoam cup that contains warm water.
You will then be able to use your data to calculate the molar heat of fusion of ice.
Laboratory Procedure
1. There should be a beaker filled with water at your lab station, as well as two Styrofoam cups
nested together. Record the mass of the clean, dry cups to the nearest 0.01 g.
2. Carefully pour water into the Styrofoam cup until the cup is approximately one-third full.
Record the mass of the cups with the water to the nearest 0.01 g.
3. Place the Styrofoam cup in a 400 mL beaker. This is done so that the cup does not tip
over. Record the initial temperature of the warm water to the nearest 0.1oC.
4. Obtain a cup of ice cubes and place them on a paper towel. Blot the excess water from
them so they are dry. IMMEDIATELY after you have recorded the water temperature,
carefully add two or three ice cubes to the water. The total mass is important, so do not allow
the water to splash or spill out of the cup.
5. Gently stir the ice water with the thermometer. Do not touch the bottom of your cup
with the thermometer. Your goal is to reach a temperature of exactly 0oC.
6. Once the water temperature reaches 0oC, place the thermometer on the counter so that it
doesn’t roll off the table. Carefully remove any extra pieces of ice with the plastic spoon. Try to
remove ice only (not water).
7. Record the mass of the Styrofoam cups with water and melted ice to the nearest 0.01 g.
8. Pour the water into the sink and place the cups upside-down on a towel to dry. Return any
unused ice to your teacher. Dry off the countertop.
9. If time permits, you can repeat the entire experiment to get data for a second trial.
Recorded Data
Trial 1
Trial 2
Mass of two clean, dry Styrofoam cups (to nearest 0.01 g)
g
g
Mass of cups + water (to nearest 0.01 g)
g
g
Initial temperature of water (to nearest 0.1 oC)
o
Mass of cups + water + melted ice (to nearest 0.01 g)
C
g
o
C
g
Calculations (Use the data for either Trial 1 or Trial 2)
1. Calculate the initial mass of water used in this experiment. Show your calculations below.
2. Calculate the amount of heat (in joules) which was transferred from the water to the ice in
this experiment. Use the following equation, and show your work below.
heat = (mass of water, in grams) x
4.18 J
g oC
x (change in temperature, in oC)
3. Calculate the mass of ice that melted in this experiment. Show your calculations below.
4. Calculate the moles of ice that melted in this experiment. Show your calculations below.
5. Calculate the heat of fusion of ice, in units of kilojoules per mole.
(Round answer to 2 sig figs.) Show your calculations below.
6. The accepted value for the molar heat of fusion of ice is 6.0 kJ/mol.
Calculate your percent error, based on your answer in #5.