Download Scientific Measurement

Scientific Measurement
Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence.
_____1. I can calculate percent
error using the equation on
Table T.
A researcher measures the mass of a sample to be 5.51 g. The actual mass of
the sample is known to be 5.80 g. Calculate the error of the measurement.
% error = 5.51-5.80 x 100 = 5% (can make the answer positive)
5.80
How many joules are in 9.3 kJ? 9,300 J
_____2. I can convert between
different metric units.
How many grams are in 39,983 mg? 39.983 g
How many milliliters are in 452 L? 452,000 mL
What is the volume of 54.6 g of beryllium (density = 1.85 g/cm3)?
_____3. I can solve for any
variable of the density equation.
Volume = 54.6g / 1.85 g/cm3 = 29.5 cm3
What is the density of a 27.8 g substance with a volume of 2.3 mL?
Density = 27.8 g / 2.3 mL = 12 g/mL
_____4. I can determine the
number of significant figures in
a measurement.
_____5. I can determine the
answer to a math problem to the
correct number of significant
figures.
How many significant figures are there in 30.50 cm? 4
How many significant figures are there in 400 sec? 1
To the correct number of significant figures, what is the answer to
5.93 mL + 4.6 mL? 10.5 mL
To the correct number of significant figures, what is the answer to
5.93 cm * 4.6 cm? 27 cm2
_____6. I can read the meniscus
on a graduated cylinder to the
correct number of significant
figures.
The volume is 75.7 mL. (only 1 decimal)
(The 7 can be any one estimated number)
_____7. I can convert numbers
into scientific notation from
standard notation.
Convert 87,394,000,000,000 to scientific notation. 8.7394 x 1013
_____8. I can convert numbers
into standard notation from
scientific notation.
Convert 5.8 x 10 to standard notation. 5,800,000,000
Convert 0.0000040934 to scientific notation. 4.0934 x 10-6
9
-5
Convert 4.3 x 10 to standard notation. 0.000043
Unit 1: Matter & Energy
Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence.
Definitions:
Atom: the smallest particle of matter that retains the properties of an
element.
_____1. I can define the terms
atom, element, compound, and
mixture.
Element: a substance that cannot be broken down; made of one type of
atom.
Compound: atoms of two or more different elements chemically
combined in a fixed (definite) ratio.
Mixture: two or more substances physically combined in a variable ratio.
Put each of the following examples into the correct column.
_____2. I can classify substances
as a pure substance (element or
compound) or a mixture.
_____3. I can list the 7 diatomic
elements.
_____4. I can identify a binary
compound.
_____5. I can define
homogeneous mixture and
heterogeneous mixture.
Examples: C12H22O11, NaCl, Fe, salt water, air, CO2, H2, Ar, soda
Element
Compound
Mixture
Fe, H2, Ar
C12H22O11, NaCl, CO2
The seven diatomic elements are:
Br2 I2 N2 Cl2 H2 O2 F2
Which of the following is a binary compound?
A) NaOH
B) H3PO4
C) H2O
D) MgSO4
Definitions:
Homogeneous Mixture: two or more substances physically combined with
uniform composition; no visible differences.
Heterogeneous Mixture: two or more substances physically combined
with a non-uniform composition; visibly different parts.
Two examples of Homogeneous Mixtures:
A) salt water (any solution – aq)
_____6. I can give examples of
homogeneous and
heterogeneous mixtures.
salt water, air, soda
B) Kool Aid
Two examples of Heterogeneous Mixtures:
A) soil
B) sand in water
_____7. I can determine how
matter will be separated using
filtration.
_____8. I can describe how
matter can be separated using
distillation.
When a mixture of sand, salt, sugar, and water is filtered, what passes
through the filter? salt, sugar, water
Which physical property makes it possible to separate the components of
crude oil by means of distillation? difference in boiling points
To separate a mixture of salt and water, use evaporation or distillation.
_____9. I can state which
separation process is best for a
given situation.
To separate a mixture of ethanol and water, use distillation.
To separate a mixture of sand and water, use filtration.
Monoatomic Element
Diatomic Element
(same colors)
____10. I can draw particle
Compound
diagrams to represent a
monoatomic element, a diatomic
element, a compound, a mixture
of 2 compounds, and a mixture
(different colors)
of an element and a compound.
Mixture of an element and compound
Mixture of 2 compounds
All chemical reactions have conservation of
_____11. I can answer questions
about the Law of Conservation
of Matter (Mass).
A) mass only
C) mass and charge only
B) charge and energy only
D) mass, charge, and energy
Given the reaction: C6H12O6 + 6 O2  6 CO2 + 6 H2O, determine the mass
of CO2 produced when 9 grams of glucose completely reacts with 9.6 grams
of oxygen to produce 5.4 grams of water.
13.2 g (total mass of reactants = total mass of products)
Write “P” for physical or “C” for chemical on the line provided.
P Copper (II) sulfate is blue.
C Copper reacts with oxygen.
_____12. I can classify a
P Copper can be made into wire.
property as physical or chemical.
P Copper has a density of 8.96 g/cm3.
P Copper melts at 1358 K.
P Copper doesn’t dissolve in water.
Write “P” for physical or “C” for chemical on the line provided.
P Copper (II) sulfate dissolves in water.
C Copper reacts with oxygen to form solid copper (I) oxide.
_____13. I can classify a change
as physical or chemical.
P Solid copper is melted.
P A chunk of copper is pounded flat.
P Copper and zinc are mixed to form brass.
P A large piece of copper is chopped in half.
C Copper reacts with bromine to form copper (II) bromide.
Circle the particle diagram that best represents Substance A after a physical
change has occurred.
_____14. In a particle diagram, I
can distinguish between a
physical change and a chemical
change.
Substance A
Draw a particle diagram to represent atoms of Li in each phase.
Solid
Liquid
Gas
____15. I can use particle
diagrams to show the
arrangement and spacing of
atoms/molecules in different
phases.
_____16. I can compare solids,
liquids, and gases in terms of
their shape, volume, particle
distance, and attraction between
particles.
Shape
Volume
Particle
Distance
Attraction
Between
Particles
Solid
Definite
Definite
Close together
Liquid
Indefinite
Definite
Random, but
close together
Gas
Indefinite
Indefinite
Random and
far apart
Strong
Weak
Very Strong
*Gases fill the container most completely
Fusion: change from solid to liquid.
Solidification: change from liquid to solid.
Condensation: change from gas to liquid.
Vaporization: change from liquid to gas.
_____17. I can state the change
of phase occurring during each
phase change.
Melting: change from solid to liquid.
Boiling: change from liquid to gas.
Sublimation: change from solid to gas.
Deposition: change from gas to solid.
Freezing: change from liquid to solid.
_____18. I can answer questions
about the Law of Conservation
of Energy.
As electrical energy is converted into heat energy, the total amount of energy
A) decreases
B) increases
C) remains the same
Definitions:
Kinetic Energy: energy a substance has due to its motion (related to
temperature).
Potential Energy: energy a substance has stored in its bonds (related to
distance between particles).
_____19. I can define kinetic
energy, potential energy,
temperature, heat, endothermic,
and exothermic.
Temperature: a measure of the average kinetic energy of a substance.
Heat: energy transferred from an object of high temperature to an object
of low temperature.
Endothermic: a process in which energy is absorbed.
Exothermic: a process in which energy is released.
Indicate whether the change is exothermic or endothermic:
fusion/melting endothermic
solidification/freezing exothermic
_____20. I can indicate if a
phase change is exothermic or
endothermic.
condensation exothermic
vaporization/boiling endothermic
sublimation endothermic
deposition exothermic
Which sample of water contains particles having the highest kinetic energy?
_____21. I can answer questions
about average kinetic energy.
A) 25 mL of water at 95˚C
B) 45 mL of water at 75˚C
C) 75 mL of water at 75˚C
D) 95 mL of water at 25˚C
Object A at 40˚C and object B at 80˚C are placed in contact with each other.
Which statement describes the heat flow between objects?
_____22. I can state the direction A) Heat flows from object A to object B
B) Heat flows from object B to object A
of heat flow.
C) Heat flows in both directions between the objects
D) No heat flow occurs between the objects
_____23. I can convert between
Celsius and Kelvin.
_____24. Given a
heating/cooling curve, I can
identify the process occurring
during each segment.
What Kelvin temperature is equal to 200˚C? 473 K
What Celsius temperature is equal to 200 K? -73˚C
AB = heating or cooling a solid
BC = melting or freezing
CD = heating or cooling a liquid
DE = vaporizing/boiling or condensing
EF = heating or cooling a gas
On the graph, circle the sections that show a change in potential energy.
_____25. Given a
heating/cooling curve, I can
determine which sections of the
curve show changes in potential
energy.
On the graph, circle the sections that show a change in kinetic energy.
_____26. Given a
heating/cooling curve, I can
determine which sections of the
curve show changes in kinetic
energy.
On the graph, circle the sections that show a phase change.
_____27. Given a
heating/cooling curve, I can
determine which sections of the
curve show phase changes.
113oC
_____28. Given a
heating/cooling curve, I can
determine the temperature at
which a substance melts/freezes
or vaporizes/condenses.
53oC
What is the freezing point of this substance? 53˚C
What is the boiling point of this substance? 113˚C
_____29. I can state the
temperature at which water
freezes in ˚C and K.
_____30. I can state the
temperature at which water
melts in ˚C and K.
_____31. I can state the
temperature at which water
boils in ˚C and K.
_____32. I can state the
temperature at which water
condenses in ˚C and K.
_____33. I can state the unit
used to measure energy.
What is the freezing point of water in ˚C and K? 0˚C or 273 K
What is the melting point of water in ˚C and K? 0˚C or 273 K
What is the boiling point of water in ˚C and K? 100˚C or 373 K
What is the condensing point of water in ˚C and K? 100˚C or 373 K
Energy is measured in joules.
Definitions:
Specific Heat Capacity: amount of heat required to increase the
temperature of 1 gram of a substance 1˚C (or 1 K).
_____34. I can define specific
heat capacity, heat of fusion,
and heat of vaporization.
Heat of Fusion: amount of heat that needs to be absorbed to melt 1 gram
of a substance (or released to freeze 1 gram of a substance).
Heat of Vaporization: amount of heat that needs to be absorbed to
vaporize 1 gram of a substance (or released to condense 1 gram of a
substance).
How much heat is needed to vaporize 10 g of water at 100˚C?
q = mHV
q = 10(2,260) = 22,600 J
(Hv on Table B, don’t plug in temp.)
How much heat is needed to change the temperature of 100.0 g of water
from 20˚C to 35˚C?
q = mC∆T
q = 100(4.18)(15) = 6,270 J
(C on Table B)
_____35. I can use the heat
equations to solve for any
variable.
How much heat is needed to melt 20 g of ice at 0˚C?
q = mHf
q = 20(334) = 6,680 J
(Hf on Table B, don’t plug in temp.)
How many grams of water can be heated by 15˚C using 13,500 J of heat?
q = mC∆T
13,500 = m(4.18)(15) = 215 g
It takes 5,210 J of heat to melt 50 g of ethanol at its melting point. What is
the heat of fusion of ethanol?
q = mHf
5,210 = 50(Hf) = 104.2 J/g
(solving for Hf, substance is not water so you cannot use Table B value)
The five parts of the Kinetic Molecular Theory are:
A. Particles of an ideal gas move in constant, random, straight line
motion.
_____36. I can state the 5 parts
of the Kinetic Molecular
Theory.
B. Particles of an ideal gas have elastic collisions (energy is not lost).
C. The size of the particles is so small compared to the space between the
particles that the volume of the actual gas particles is negligible.
D. Particles of an ideal gas have no forces of attraction between them.
E. Particles of an ideal gas collide with container walls causing pressure.
_____37. I can define an ideal
gas.
Definition:
Ideal Gas: follows all of the parts of the Kinetic Molecular Theory.
_____38. I can state why real
gases do not behave like ideal
gases.
Real gases do not behave like ideal gases because real gases have some
_____39. I can state the two
elements that act ideally most of
the time.
_____40. I can state the
conditions under which real
gases behave most like ideal
gases.
The two elements that act ideally most of the time are hydrogen &
volume and some attraction.
helium.
A real gas will act most ideally under the conditions of low pressure
and high temperature. (PLIGHT)
STP stands for Standard Temperature and Pressure.
_____41. I can state the
meaning of “STP” and the
Table on which it can be found.
The values can be found on Table A.
Standard Pressure = 1 atm or 101.3 kPa
Standard Temperature = 0˚C or 273 K
Avogadro’s Hypothesis says that samples of a gas that are at the same
temperature, pressure, and volume will contain the same number of
molecules. The actual identity of the substances does not matter.
_____42. I can state Avogadro’s
Hypothesis (Law) and answer
questions based on this concept.
Which cylinder contains the same number of molecules at STP as a 2 L
cylinder of H2 (g) at STP?
A) 1 L cylinder of O2 (g)
C) 2 L cylinder of CH4 (g)
B) 1.5 L cylinder of NH3 (g)
D) 4 L cylinder of He (g)
_____43. I can state the
relationship between pressure
and volume for gases.
At constant temperature, as the pressure on a gas increases, the volume
_____44. I can state the
relationship between
temperature and volume for
gases.
At constant pressure, as the temperature of a gas increases, the volume
_____45. I can state the
relationship between
temperature and pressure for
gases.
In a fixed container (has constant volume), as the temperature of a gas
decreases.
increases.
increases, the pressure increases.
Draw a curve for each of the graphs below:
_____46. I can recognize the
relationships stated above in
graph form.
P
V
V
P
T
T
A rigid cylinder with a movable piston contains a 2 L sample of neon gas at
STP. What is the volume when its temperature is increased to 30˚C while its
pressure is decreased to 90 kPa?
P1V1 = P2V2
101.3(2) = 90V2
V2 = 2.5 L
T1
T2
273
303
_____47. I can solve problems
using the combined gas law.
* STP values on Table A
* Choose the pressure with the same units as the final pressure
* Temperature needs to be in Kelvin
* Use the temperature conversion equation on Table T
Definitions:
Vapor Pressure: the force a gas exerts above a liquid.
_____48. I can define vapor
pressure, boiling point, and
volatile.
Boiling Point: the temperature at which the vapor pressure is equal to the
atmospheric pressure.
Volatile: easily turns into a vapor (easily evaporates).
_____49. I can state the
conditions of pressure that are
used for normal boiling points.
The normal boiling point of a substance occurs at a pressure of
1 atm or 101.3 kPa. This is known as standard pressure and can be found
on Table A.
_____50. I can state the
As the temperature of a substance increases, its vapor pressure increases.
relationship between
temperature and vapor pressure.
_____51. I can state the
As the atmospheric pressure increases, the boiling point increases.
relationship between
atmospheric pressure and boiling
point.
_____52. I can state the
As the attraction between molecules increases, the boiling point increases.
relationship between molecular
attraction and boiling point.
_____53. I can determine
strength of attraction between
molecules using Table H.
Based on Table H, which substance has the weakest intermolecular forces?
Propanone, because at any temperature it has the highest vapor pressure.
This means it is easily converted into a vapor (weak forces).
_____54. I can determine the
vapor pressure of ethanol,
ethanoic acid, propanone, or
water at a given temperature.
What is the vapor pressure of ethanol at 56˚C? 39 kPa
_____55. I can determine the
temperature of ethanol, ethanoic
acid, propanone, or water at a
given vapor pressure.
When the vapor pressure of water is 30 kPa, what is the temperature of the
water? 70˚C
Unit 2: The Atom
Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence.
_____1. I can describe John
Dalton’s contribution to our
understanding of the atom.
Dalton’s Model: the atom is the basic unit of matter; atoms of the same
element have the same properties, atoms of different elements have
different properties.
Diagram:
Thomson’s Model: “Plum Pudding” Model – small, negative electrons
embedded in a large positive sphere.
_____2. I can describe JJ
Thomson’s contribution to our
understanding of the atom.
Diagram:
Rutherford’s Experiment: directed positive alpha particles at a thin sheet
of gold foil and observed their path.
Experiment Results: most alpha particles went straight through, a few
were slightly deflected, some bounced back.
_____3. I can describe Ernest
Rutherford’s contribution to our
understanding of the atom.
Experiment Conclusions: the atom is mostly empty space, with the mass
concentrated in the small, dense, positively charged nucleus.
Diagram:
Bohr’s Model: “Planetary Model” – electrons are in definite orbits
(circular paths) around the nucleus.
_____4. I can describe Niels
Bohr’s contribution to our
understanding of the atom.
Diagram:
Where are electrons likely to be found according to the modern model?
Orbitals – regions of probable location
_____5. I can describe the
modern model (wave mechanical
model) of the atom.
Diagram:
Particle #1
Particle #2
Particle #3
Name
Proton
Neutron
Electron
Charge
+1
0
-1
Mass
1 amu
1 amu
0 amu
Location in
Atom
Nucleus
Nucleus
Outside
Nucleus
_____6. I can state the three
subatomic particles, their
location in an atom, and the
magnitude of their charges and
their masses (in amu).
_____7. I can identify the nuclear
charge of an atom.
_____8. I can explain why atoms
are electrically neutral.
_____9. I can define atomic
number and mass number.
What is the charge of an oxygen-18 nucleus? +8
Atoms are electrically neutral because the number of protons is
equal to the number of electrons.
Definitions:
Atomic Number: the number of protons in the nucleus; defines the
identity of the element.
Mass Number: the sum of the protons and neutrons.
_____10. I can use the Periodic
Table to determine the atomic
number of an element.
How many protons are in an atom of selenium? 34
_____11. I can determine the
mass number of an atom.
What is the mass number of an atom that has 20 protons, 19 neutrons, and
20 electrons? 20 + 19 = 39
_____12. I can represent an atom
using different notations.
How many protons are in an atom of silicon? 14
What are three other notations that can be used to represent 80Br?
80
Br
Bromine-80
Br-80
35
_____13. Given the mass
number, I can determine the
number of protons, neutrons,
and electrons in an atom.
In an atom of 212Po, how many protons are present? 84
In an atom of 212Po, how many electrons are present? 84
In an atom of 212Po, how many neutrons are present? 128
Definitions:
Ion: charged particle due to the loss or gain of electrons.
_____14. I can define ion,
cation, and anion.
Cation: a positive ion formed when an atom loses one or more electrons.
Anion: a negative ion formed when an atom gains one or more electrons.
How many protons are in 18O-2? 8
How many neutrons are in 18O-2? 10
_____15. Given the mass number
and the charge, I can determine
the number of protons, neutrons,
and electrons in an ion.
How many electrons are in 18O-2? 10
What is this ion’s electron configuration? 2-8
How many electrons are in 27Al+3? 10
What is this ion’s electron configuration? 2-8
_____16. I can define isotope.
_____17. I can identify isotopes
based on their symbols.
_____18. I can calculate average
atomic mass given the masses of
the naturally occurring isotopes
and the percent abundances.
Definition:
Isotope: atoms of the same element (same atomic #, same # of protons)
with a different mass # (different # of neutrons).
Which symbols represent atoms that are isotopes?
A) 14C and 14N B) 16O and 18O C) 131I and 131I D) 222Rn and 222Ra
Element Q has two isotopes. If 77% of the element has an isotopic mass of
83.7 amu and 23% of the element has an isotopic mass of 89.3 amu, what is
the average atomic mass of the element?
1st % (1st mass #) + 2nd % (2nd mass #) + …
100
= 77(83.7) + 23(89.3)
100
= 84.988 amu
(answer should be in between the two given masses, but closer to the one
with the larger percent abundance)
Definitions:
Principal Energy Level: the main energy level or shell of an atom.
Orbital: the region where an electron is most likely to be found.
_____19. I can define principal
energy level, orbital, ground
state, excited state, electron
configuration, and bright line
spectrum.
Ground State: electrons are in the lowest energy levels possible (closest
to the nucleus); the Periodic Table shows ground state configurations.
Excited State: a temporary state that exists when an electron jumps to a
higher energy level (farther from the nucleus).
Electron Configuration: shows the number of electrons in each shell.
Bright Line Spectrum: narrow lines of color given off when an electron in
the excited state releases energy and returns to the ground state.
The 1st shell holds a maximum of 2 electrons.
_____20. I can state the
maximum number of electrons
that will fit into each of the first
four shells.
The 2nd shell holds a maximum of 8 electrons.
The 3rd shell holds a maximum of 18 electrons.
The 4th shell holds a maximum of 32 electrons.
_____21. I can define valence
electrons.
Definition:
Valence Electrons: electrons in the last shell; it is the last number written
in the electron configuration on the Periodic Table.
_____22. I can define a stable
octet.
Definition:
Stable Octet: having a full outer shell of 8 valence electrons; it makes the
atom stable and unreactive.
How many valence electrons does rubidium have in the ground state? 1
_____23. I can locate and
interpret an element’s electron
configuration on the Periodic
Table.
How many principal energy levels contain electrons in an atom of iodine in
the ground state? 5
How many electrons does an aluminum atom have in its 2nd shell? 8
_____24. I can state the
relationship between distance
from the nucleus and energy of
an electron.
_____25. I can state the
relationship between the number
of the principal energy level and
the distance to the atom’s
nucleus.
As the distance between the nucleus and the electron increases, the energy
_____26. I can identify an
electron configuration that
shows an atom in the excited
state.
Which electron configuration represents an atom of potassium in the
excited state?
_____27. I can explain, in terms
of subatomic particles and energy
states, how a bright line
spectrum is created.
A bright-line spectrum is created when an electron in the excited state (a
higher energy state) releases energy in the form of color as it returns to
the ground state (a lower energy state).
of the electron increases.
As the number of the principal energy level increases, the distance to the
nucleus increases.
A) 2-8-7-1
B) 2-8-8-1
C) 2-8-7-2
D) 2-8-8-2
_____28. I can identify the
elements shown in a bright line
spectrum.
Which element(s) is/are present in the mixture? D and E
Draw the Lewis electron dot diagram for the following atoms:
_____29. I can draw Lewis
electron dot diagrams for a given
element.
Unit 3: Nuclear Chemistry
Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence.
_____1. I can explain what
makes a nucleus stable or
unstable.
The stability of the nucleus is dependent on the neutron to
proton ratio.
Type
Symbol
Notation
alpha
_____2. I can compare types of
radiation.
Charge
Mass
Penetrating
Power
+2
4 amu
Low
beta
β-
-1
0 amu
Moderate
positron
β+
+1
0 amu
Moderate
gamma
γ
0
0 amu
High
The difference between natural transmutation and artificial transmutation is
that in natural transmutation an unstable nucleus breaks apart on its own
_____3. I can explain the
difference between natural
transmutation and artificial
transmutation.
becoming a different element and in artificial transmutation a stable
nucleus is made unstable by hitting it with a high energy particle (such as a
proton, neutron, or gamma radiation). Once hit, the nucleus changes into a
nucleus of a different element.
Two synonyms for spontaneous decay are: natural transmutation
_____4. I can state two synonyms
for spontaneous decay.
and natural decay.
Which equation represents a natural decay?
_____5. I can identify a natural
decay reaction from a list of
reactions.
Which equation represents artificial transmutation?
_____6. I can identify an
artificial transmutation reaction
from a list of reactions.
Complete the following nuclear equations:
_____7. I can show how mass
number and electrical charge are
conserved in any nuclear
reaction.
30
15
P
92
36
_____8. I can define half-life.
Kr
Definition:
Half-life: the amount of time it takes for half of the mass of a radioactive
sample to decay.
Based on Reference Table N, what fraction of a radioactive sample of Au198 will remain unchanged after 10.78 days?
1 half-life = 2.695 days, so 10.78 days is 4 half-lives = 1/16 of the sample
(1 HL = 1/2, 2 HL = 1/4, 3 HL = 1/8, 4 HL = 1/16, 5 HL = 1/32)
_____9. Given the length of the
half-life and the amount of time
that has passed, I can determine
the amount of radioactive sample
that will remain or the original
amount.
What was the original mass of a radioactive sample of K-37 if the sample
decayed to 25.0 g after 4.92 seconds?
H
A
T
0
400 g
0
1
200 g
1.23 s
2
100 g
2.46 s
3
50 g
3.69 s
4
25 g
4.92 s
(half-life on Table N, add 1 half-life down T column, multiply by 2 to get
to the original amount in between the two zeros)
_____10. Given the length of the
half-life and the amount of
radioactive sample remaining, I
can determine the amount of
time that has passed.
A 100.0 g sample of Co-60 decays until only 12.5 g of it remains. Given
that the half-life of Co-60 is 5.271 years, how long did the decay take?
H
A
T
0
100 g
0
1
50 g
5.271 y
2
25 g
10.542 y
3
12.5 g 15.813 y
(divide by 2 to move down the A column, then add 1 half-life)
_____11. Given the amount of
time that has passed and the
amount of radioactive sample
remaining, I can determine the
length of the half-life.
What is the half-life of a radioisotope if 25.0 g of an original 200.0 g sample
remains unchanged after 11.46 days?
H
A
T
0
200 g
0
1
100 g
3.82 d
2
50 g
7.64 d
3
25 g
11.46 d
11.46 d / 3
(divide the last time by the number of half-lives that passed)
Compared to K-37, the isotope K-42 has
_____12. Using Table N, I can
determine the length of half-life
and/or decay mode for a specific
radioactive isotope.
A) shorter half-life and the same decay mode
B) shorter half-life and a different decay mode
C) longer half-life and the same decay mode
D) longer half-life and a different decay mode
_____13. I can explain why all
nuclear reactions release A LOT
more energy than chemical
reactions do.
Nuclear reactions release A LOT more energy than chemical reactions do
because some of the mass is converted into energy.
Definitions:
Fission: when an atom is hit with a neutron, producing lighter nuclei,
more neutrons, and energy (BIG  SMALL). Fission equations usually
contain uranium and neutrons are on both sides of the arrow. The
neutrons produced allow a chain reaction to occur.
_____14. I can define fission and
fusion.
Fusion: when two light nuclei combine to form a heavier nucleus
(SMALL  BIG). Fusion reactions have two hydrogen atoms, two
helium atoms, or one hydrogen atom and one helium atom on the left
side of the arrow.
Which equation represents fission?
_____15. I can identify a fission
reaction from a list of reactions.
Which equation represents fusion?
_____16. I can identify a fusion
reaction from a list of reactions.
_____17. I can state the
conditions of temperature and
pressure that are needed for a
fusion reaction to happen.
The temperature and pressure conditions needed for fusion to happen are:
high temperature and high pressure.
Which of the following equations represent NUCLEAR reactions?
_____18. Given a list of
reactions, I can differentiate a
nuclear reaction from a chemical
reaction.
Five beneficial uses for radioactive isotopes are:
A. Tracing chemical and biological processes
_____19. I can state 5 beneficial
uses for radioactive isotopes.
B. Detection and treatment of disease (must have short half-lives)
C. Food storage
D. Radioactive dating
E. Nuclear power
Tc-99 and Co-60 are used for treating cancer.
_____20. I can state the scientific
use of specific radioactive
isotopes.
I-131 is used for treating thyroid disorders.
C-14 is used for dating once living organisms (wood, bone, animal skin).
U-238 is used for dating geological formations.
Three risks associated with radioactivity and radioactive isotopes are:
A. Biological exposure and damage
_____21. I can state three risks
associated with radioactivity and
radioactive isotopes.
B. Long term storage and disposal (waste stays radioactive for long
periods of time because they have very long half-lives)
C. Nuclear accidents and pollution
Unit 4: Periodic Table
Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence.
_____1. I can state how the
elements on the Periodic Table
are arranged.
The elements on the Periodic Table are arranged by increasing
atomic number.
Group 1 is called the alkali metals.
Group 2 is called the alkaline earth metals.
_____2. I can state the group
names for elements in groups 1,
2, 3-12, 17, and 18.
Groups 3-12 are called the transition metals or transition elements.
Group 17 is called the halogens.
Group 18 is called the noble gases.
_____3. I can explain why
elements in the same group have
similar chemical properties.
Elements in the same group have similar chemical properties because they
have the same number of valence electrons (the last number in their
electron configuration is the same).
_____4. I can explain what
elements in the same period have
in common.
Elements in the same period have the same number of occupied principal
energy levels (shells) (their electron configurations consist of the same
amount of numbers).
Definitions:
Electronegativity: the strength of attraction between the nucleus and a
bonded electron; the higher the EN, the stronger the attraction (Table S).
_____5. I can define
electronegativity, ionization
energy, atomic radius, and ionic
radius.
Ionization Energy: the amount of energy needed to remove the outermost
electron (Table S).
Atomic Radius: the distance from the nucleus to the outermost electron
(the size of the atom) (Table S).
Ionic Radius: the size of an ion once the electron(s) have been lost or
gained.
Classify each of the following elements as metals (M), nonmetals (NM), or
metalloids (MTLD) (have properties of both metals and nonmetals).
_____6. I can classify elements as
metals, nonmetals, or metalloids
based on their placement on the
Periodic Table.
MTLD B
M
K
M
Li
NM
C
NM
Ar
MTLD Sb
NM
H
M
Fe
M
Au
NM
S
MTLD
Si
M
Fr
NM
He
NM
Rn
M
Al
MTLD As
M
Bi
NM
I
NM
F
MTLD Ge
_____7. I can state the
names/symbols for the two
elements on the Periodic Table
that are liquids at STP.
The two elements that are liquids at STP are:
Mercury (Hg) and Bromine (Br).
The 11 elements that are gases at STP are:
_____8. I can state the
names/symbols of the 11
elements that are gases at STP.
Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe),
Radon (Rn), Oxygen (O), Nitrogen (N), Hydrogen (H), Fluorine (F), and
Chlorine (Cl).
The properties of metals are:
Located to the left of the “staircase,” magnetic, malleable, ductile, have
_____9. I can list properties of
metals.
luster, solids (except Mercury), good conductors of heat and electricity,
high melting and boiling points, tend to lose electrons and form positive
ions because they have large radii, low EN values, and low IE values.
The properties of nonmetals are:
Located to the right of the “staircase,” exist in all 3 phases (Group 17
contains all 3 phases), brittle, dull, poor conductors of heat and
_____10. I can list properties of
nonmetals.
electricity (insulators), low melting and boiling points, tend to gain
electrons and form negative ions because they have small radii, high EN
values, and high IE values.
As one reads down a group from top to bottom, the activity/reactivity of
_____11. I can state the trend for
activity/reactivity for METALS
as one reads down a group.
_____12. I can state the trend for
activity/reactivity for
NONMETALS as one reads
down a group.
METALS increases.
The most reactive metal is Francium (Fr).
As one reads down a group from top to bottom, the activity/reactivity of
NONMETALS decreases.
The most reactive nonmetal is Fluorine (F).
Metals lose electrons and become positive ions that are smaller
_____13. I can explain how
losing or gaining electrons
affects the radius of an ion.
than the original atom because the number of shells decreases.
Nonmetals gain electrons and become negative ions that are larger
than the original atom because of more electron repulsion (more e -).
_____14. I can explain why the
elements in Group 18 don’t
usually react with other
elements.
Elements in Group 18 don’t usually react with other elements because they
have a stable octet of 8 valence electrons.
Definition:
Allotrope: forms of the same element in the same state that have
different structures and different properties.
_____15. I can define allotrope
and give examples.
Ex: Oxygen (O2 (g)) and Ozone (O 3 (g))
Ex: Carbon – diamond and graphite
As one reads down a group from top to bottom, electronegativity
decreases because the number of shells increases (shielding effect – inner
_____16. I can state the periodic
trend for electronegativity and
explain why it occurs.
shells block/weaken attraction between nucleus and valence electrons).
As one reads across a period from left to right, electronegativity
increases because the number of protons (nuclear charge) increases (more
protons pull electrons in closer and strengthen the attraction).
As one reads down a group from top to bottom, ionization energy
decreases because the number of shells increases (shielding effect – inner
shells block/weaken attraction between nucleus and valence electrons so
_____17. I can state the periodic
trend for ionization energy and
explain why it occurs.
it is easier for them to be removed).
As one reads across a period from left to right, , first ionization energy
increases because the number of protons (nuclear charge) increases (more
protons pull electrons in closer and strengthen the attraction so it is
harder for them to be removed).
As one reads down a group from top to bottom, atomic radius
increases because the number of shells increases (more shells take up
_____18. I can state the periodic
trend for atomic radius and
explain why it occurs.
more space).
As one reads across a period from left to right, atomic radius
decreases because the number of protons (nuclear charge) increases (more
protons pull electrons in closer, making the atom smaller).
Unit 5: Bonding
Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence.
Definitions:
Intramolecular Forces: forces of attraction within the same molecule.
Examples: ionic bonds, covalent bonds, metallic bonds
_____1. I can define
intramolecular forces and
intermolecular forces and give
examples of each.
Intermolecular Forces: forces of attraction between different molecules.
Examples: Van der Waals (dipole-dipole forces, dispersion forces),
hydrogen bonding, molecule-ion attraction.
BARF stands for Break Absorb Release Form
This means that when a bond is FORMED, energy is released
and when a bond is BROKEN, energy is absorbed.
_____ 2. I can explain and apply
the meaning of BARF as it
applies to chemical bonding.
Given the balanced equation: N + N  N2
Which statement describes the process represented by this equation?
A) A bond is formed as energy is absorbed.
B) A bond is formed as energy is released.
C) A bond is broken as energy is absorbed.
D) A bond is broken as energy is released.
_____3. I can state the number of
valence electrons that an atom
attains to be most stable.
_____4. I can state the three
types of chemical bonds.
Atoms are most stable when they have 8 valence electrons.
The three types of chemical bonds are ionic, covalent, and metallic.
Definitions:
Ionic Bond: form between a metal and a nonmetal. Electrons are
transferred from the metal to the nonmetal. The metal becomes a
positive ion and the nonmetal becomes a negative ion and the two
oppositely charged ions attract.
_____5. I can define ionic bond,
covalent bond, and metallic bond
in terms of the types of elements
(metals, nonmetals) from which
they are formed.
Covalent Bond: form between two or more nonmetals. Electrons are
shared between the atoms.
Metallic Bond: form between the same metal. The metals lose their
electrons and the positive ions that form are attracted to the “sea of
mobile electrons.”
In an ionic bond, the valence electrons of the metals are transferred to the
nonmetal so that each atom attains a stable octet (like noble gases). In these
_____6. I can define ionic,
covalent, and metallic bonds
based on what happens to the
valence electrons.
compounds, the electronegativity difference is large (usually greater than
1.7). In a covalent bond, the valence electrons of the two nonmetals are
shared so that each atom attains a stable octet. In these compounds, the
electronegativity difference is small (usually less than 1.7). In a metallic
bond, positive ions are immersed in a sea of mobile electrons.
_____7. I can determine the
compound that has the greatest
ionic character.
Which of the following compounds has the greatest ionic character?
A) LiCl
B) CaCl2
C) BCl3
D) RbCl
Draw Lewis dot diagrams for the following ionic compounds:
NaCl
_____8. I can draw a Lewis dot
diagram to represent an ionic
compound.
CaCl2
Physical properties of ionic substances are: they have hard, crystalline
_____9. I can state the properties
of ionic substances.
structures as a solid (crystal lattice – geometric pattern), they have high
melting and boiling points, they conduct electricity as a liquid and in
solution (aq) (mobile charges), they are soluble (dissolve in water).
A solid substance was tested in the lab. The results are shown below.
*dissolves in water
_____10. I can identify a
substance as ionic based on its
properties.
*is an electrolyte
* has a high melting point
Based on these results, the solid substance could be
A) Hg
B) AuCl
C) CH4
D) C12H22O11
Based on bond type, which compound has the highest melting point?
A) CH4
B) C12H22O11
C) NaCl
D) C5H12
In a single covalent bond, 2 electrons or 1 pair is shared.
_____ 11. I can state the number
of electrons that are shared in
single and multiple covalent
bonds.
In a double covalent bond, 4 electrons or 2 pairs are shared.
In a triple covalent bond, 6 electrons or 3 pairs are shared.
Draw Lewis dot diagrams for the following molecular substances and
identify their shapes. (Use the CNOF table)
H2O - bent
_____12. I can draw a Lewis dot
diagram to represent a molecular
(covalently bonded) substance
and identify its shape as linear,
bent, trigonal pyramidal, or
tetrahedral.
I2 - linear
NH3 – trigonal pyramidal
CO2 - linear
CH4 - tetrahedral
CHCl3 – tetrahedral
Physical properties of molecular substances are: they can exist in all 3
_____13. I can state the
properties of molecular
substances.
phases, they are usually not soluble in water, solids are soft, they are
poor conductors in any phase (no mobile charges), they have low melting
and boiling points.
_____14. I can identify a
substance as “molecular” based
on its properties.
_____15. I can state the type of
bonding that occurs in
compounds containing
polyatomic ions.
_____16. Given the chemical
formula for a compound, I can
determine the type(s) of bonding
in the compound.
Compounds containing polyatomic ions have both ionic and covalent
bonding.
State the type(s) of bonding in the following compounds:
NaCl - ionic
CO - covalent
Hg - metallic
Na3PO4 - ionic and covalent
Explain, in terms of valence electrons, why the bonding in methane (CH4) is
similar to the bonding in water (H2O).
_____17. In terms of valence
electrons, I can find similarities
and differences between the
bonding in several substances.
In both molecules, the valence electrons are shared between the two
atoms to form a covalent bond.
Explain, in terms of valence electrons, why the bonding in HCl is different
than the bonding in NaCl.
In HCl the valence electrons are shared to form a covalent bond, while in
NaCl the valence electrons are transferred to form an ionic bond.
Polar covalent bonds are formed when different nonmetals share electrons
_____18. I can explain the
difference between a polar
covalent bond and a nonpolar
covalent bond.
unequally. The electronegativity difference is between 0 and 1.7.
Nonpolar covalent bonds are formed when the same nonmetals share
electrons equally. The electronegativity difference is zero.
_____19. I can explain how to
determine the degree of polarity
of a covalent bond.
_____20. I can explain why one
covalent bond is more or less
polar than another covalent
bond.
The degree of polarity of a covalent bond is determined by the
electronegativity difference between the elements.
Explain, in terms of electronegativity difference, why the bond between
carbon and oxygen in a carbon dioxide molecule is less polar than the bond
between hydrogen and oxygen in a water molecule.
The difference in electronegativity between carbon and oxygen is less
than the difference in electronegativity between hydrogen and oxygen.
Draw a water molecule and include slight charges. Explain why you
assigned each atom its charge.
_____21. I can determine which
end of a polar covalent bond is
slightly negative and which end
is slightly positive and explain
why.
Oxygen is slightly negative because it has a higher electronegativity and
hydrogen is slightly positive because it has a lower electronegativity.
_____22. I can define
symmetrical and asymmetrical.
Definitions:
Symmetrical: a molecule in which all slight charges cancel (looks the
same when the molecule is turned upside down). There is an equal
distribution of charge in a symmetrical molecule.
Asymmetrical: a molecule in which all slight charges do not cancel (looks
different when the molecule is turned upside down). There is an unequal
distribution of charge in an asymmetrical molecule.
SNAP means symmetrical nonpolar asymmetrical polar.
Why is a molecule of CH4 nonpolar even though the bonds between the
carbon and hydrogen are polar?
_____23. I can explain and apply
the meaning of SNAP as it
relates to determining molecule
polarity.
A) The shape of the CH4 molecule is symmetrical.
B) The shape of the CH4 molecule is asymmetrical.
C) The CH4 molecule has an excess of electrons.
D) The CH4 molecule has a deficiency of electrons.
Explain, in terms of charge distribution, why a molecule of water is polar.
Water is asymmetrical so the slight charges that exist do not cancel out
(the molecule has asymmetrical distribution of charge).
Determine which molecules are polar and which are nonpolar.
Justify your answer.
_____24. I can determine if a
molecular is polar or nonpolar.
H2O – polar, asymmetrical
CO2 – nonpolar, symmetrical
I2 – nonpolar, symmetrical
(nonpolar bonds)
CH4 – nonpolar, symmetrical
NH3 – polar, asymmetrical
CHCl3 – polar, asymmetrical
Definitions:
Dipole-Dipole: a type of van der Waal’s force that exists between polar
molecules. The slight negative end of one polar molecule is attracted to
the slight positive end of another polar molecule.
_____25. I can define the
different types of intermolecular
forces.
Dispersion: a type of van der Waal’s force that exists between nonpolar
molecules. The attraction comes from the temporary slight charges that
exist when the electrons move around the nucleus. Wherever the
electron is, it brings a slight negative charge with it.
Hydrogen Bonding: occurs between the hydrogen atom in one molecule
and a nitrogen, oxygen, or fluorine atom in another molecule (H-NOF).
Hydrogen bonds are very strong and explain the unusually high boiling
point of water.
Molecule-Ion Attraction: exists between ions of an ionic compound and
the molecules of a polar liquid. This attraction is present in all solutions
(aq) and explains how a substance dissolves.
_____26. I can state the
relationship between polarity
and intermolecular force
strength.
_____27. I can state the
relationship between size of the
molecule and intermolecular
force strength.
_____28. I can state the
relationship between the strength
of the intermolecular forces and
vapor pressure.
_____29. Given the physical state
of some substances, I can
compare the relative strength of
the intermolecular forces.
_____30. Given the boiling
points (or freezing points) of
some substances, I can compare
the relative strength of the
intermolecular forces.
As the polarity of the molecule increases, the strength of the forces
increases.
As the size of the molecule increases, the strength of the forces increases.
As the strength of the forces increases, vapor pressure decreases.
At STP, iodine (I2) is a crystal and fluorine (F2) is a gas. Compare the
strength of the intermolecular forces in a sample of I2 at STP to the strength
of the intermolecular forces in a sample of F2 at STP.
Since I2 is a solid, it must have stronger intermolecular forces of
attraction compared to fluorine, which is a gas. The stronger the forces
of attraction, the closer the particles are arranged.
At STP, CF4 boils at -127.8˚C and NH3 boils at -33.3˚C. Which substance
has stronger intermolecular forces? Justify your answer.
NH3 has stronger intermolecular forces of attraction because it has a
higher boiling point. The stronger the forces of attraction, the higher the
boiling point is.
Which compound has hydrogen bonding between its molecules?
A) CH4
_____31. I can answer questions
about hydrogen bonding.
B) CaH2
C) KNO3
D) H2O
The relatively high boiling point of water is due to water having
A) hydrogen bonding
C) nonpolar covalent bonding
B) metallic bonding
D) strong ionic bonding
In which system do molecule-ion attractions exist?
A) NaCl (aq)
B) NaCl (s)
C) C6H12O6 (aq)
D) C6H12O6 (s)
Which diagram best illustrates the ion-molecule attractions that occur when
the ions of NaCl (s) are added to water?
_____32. I can answer questions
about molecule-ion attraction.
Unit 6: Chemical Reactions
Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence.
Write the name for the following compounds:
CrO chromium (II) oxide
_____1. Given the chemical
formula, I can write the name for
binary compounds.
MgI2 magnesium iodide
Ni2O3 nickel (III) oxide
FeCl2 iron (II) chloride
SnS tin (II) sulfide
Write the chemical formula for the following compounds:
sodium bromide NaBr
_____2. Given the name, I can
write the chemical formula for
binary compounds.
lithium iodide LiI
iron (III) fluoride FeF3
vanadium (V) oxide V2O5
Write the name for the following compounds:
_____3. Given the chemical
formula, I can write the name for
ternary compounds.
CuSO4 copper (II) sulfate
(NH4)3PO4 ammonium phosphate
Fe3(PO4)2 iron (II) phosphate
Write the chemical formula for the following compounds:
_____4. Given the name, I can
write the chemical formula for
ternary compounds.
calcium oxalate CaC2O4
nickel (II) thiosulfate NiS2O3
Write the name for the following compounds:
_____5. Given the chemical
formula, I can write the name for
covalent molecules.
N2O dinitrogen monoxide
SF6 sulfur hexafluoride
N2O3 dinitrogen trioxide
Write the chemical formula for the following compounds:
_____6. Given the name, I can
write the chemical formula for
covalent molecules.
diphosphorus pentasulfide P2S5
carbon monoxide CO
dinitrogen tetroxide N2O4
How many atoms are in N2? 2
_____7.Given the chemical
symbol/formula, I can determine
how many atoms are present.
What is the total number of atoms in Pb(C2H3O2)2? 15
How many atoms of carbon are in Pb(C2H3O2)2? 4
Using the symbols shown below, complete the equation below to illustrate
conservation of mass.
_____8. I can use particle
diagrams to show conservation
of mass in a chemical equation.
= Al
= Br
2 Al
+
3 Br2

2 AlBr3
_____9. I can balance a chemical
equation showing conservation
of mass using the lowest whole
number coefficients.
Balance the following equation using the lowest whole number coefficients.
_____10. Given a partially
balanced equation, I can predict
the missing reactant or product.
Use the law of conservation of mass to predict the missing product.
_____11. Given a list of chemical
reactions, I can classify them as
being a synthesis reaction,
decomposition reaction, single
replacement reaction, or double
replacement reaction.
_____12. From a list of given list
of elements, I can determine
which element is most active.
Al2(SO4)3 + 3 Ca(OH)2  2 Al(OH)3 + 3 CaSO4
2 NH4Cl + CaO  2 NH3 + H2O + CaCl2
Classify the following reactions as synthesis, decomposition, single
replacement, or double replacement.
SR
S
D
DR
Which of the following elements is most likely to react? (Table J)
A) Cu
B) Al
D) Mg
C) Li
Which reaction occurs spontaneously?
_____13. I can predict if a single
replacement reaction will occur.
_____14. I can predict the
products in a double replacement
reaction.
A) Cl2 + 2 NaBr  Br2 + 2 NaCl
C) I2 + 2 NaBr  Br2 + 2 NaI
B) Cl2 + 2 NaF  F2 + 2 NaCl
D) I2 + 2 NaF  F2 + 2 NaI
Complete and balance the following reaction:
MgCl2 +
2 AgNO3 
Mg(NO3)2
+ 2 AgCl
Unit 7: Moles
Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence.
_____1. I can define gram
formula mass (molar mass).
The gram formula mass of a substance is the mass of 1 mole.
_____2. I can define a mole as it
pertains to chemistry.
Definition:
Mole: a unit of measurement for a substance;
1 mole = 6.02 x 1023 particles, mass is equal to the gram formula mass.
1 mole of a substance will occupy 22.4 L.
What is the gfm for Ni2O3? Ni: 58.7 x 2 = 117.4
O: 16.0 x 3 = 48.0
gfm = 165.4 g
_____3. I can determine the
gram-formula mass for any
compound or hydrate.
What is the gfm for Pb(C2H3O2)2? Pb: 207.2 x 1 = 207.2
C: 12.0 x 4 = 48.0
H: 1.0 x 6 = 6.0
O: 16.0 x 4 = 64.0
gfm = 325.2 g
What is the gfm for CuSO4 • 5 H2O? Cu: 63.5 x 1 = 63.5
S: 32.0 x 1 = 32.0
O: 16.0 x 4 = 64.0
H2O: 18.0 x 5 = 90.0
gfm = 249.5 g
_____4. I can determine the
percent composition of an
element in a compound.
What is the percent by mass of Mg in Mg(NO3)2?
Mg: 24.3 x 1 = 24.3
% comp = part x 100
N: 14.0 x 2 = 28.0
whole
O: 16.0 x 6 = 96.0
gfm = 148.3 g
= 24.3 x 100 = 16.4 %
148.3
_____5. I can define hydrate.
Definition:
Hydrate: crystals with water inside; an anhydrous substance does not
contain water and is formed when a hydrate is heated.
_____6. I can determine the
percent composition of water in
a hydrate.
What is the percent by mass of H2O in Na2SO4 • 10 H2O?
Na: 23.0 x 2 = 46.0
% comp = part x 100
S: 32.0 x 1 = 32.0
whole
O: 16.0 x 4 = 64.0
H2O: 18.0 x 10 = 180.0
= 180.0 x 100 = 55.9 %
gfm = 322.0 g
322.0
A student heated a 9.1 g sample of a hydrate to a constant mass of 5.41 g.
What percent by mass of water did the hydrate contain?
Mass of H2O = hydrate – anhydrous = 9.1 – 5.41 = 3.69
% comp = part x 100
whole
= 3.69 x 100 = 40.5 %
9.1
_____7. I can define empirical
formula and molecular formula.
_____8. Given the molecular
formula, I can determine the
empirical formula of a
compound.
_____9. Given the percent
composition, I can determine the
empirical formula of a
compound.
_____10. Given the empirical
formula and the molar mass, I
can determine the molecular
formula of a compound.
Definitions:
Empirical Formula: the simplest ratio in which atoms combine; the
subscripts are reduced.
Molecular Formula: the actual ratios of atoms that combine; the
subscripts are not reduced and are a multiple of the empirical.
Which pair consists of a molecular formula and its corresponding empirical
formula?
A) C2H2 and CH3CH3
C) P4O10 and P2O5
B) C6H6 and C2H2
D) SO2 and SO3
A compound consists of 25.9% nitrogen and 74.1% oxygen by mass. What
is the empirical formula of the compound?
N = 25.9 g
O = 74.1 g
14.0
16.0
= 1.85
= 4.63
N2O5
1.85
1.85
= 1x 2
= 2.5 x 2
What is the molecular formula of a compound that has the empirical
formula of CH2O and a molar mass of 180 g/mol?
C: 12.0 x 1 = 12.0
H: 1.0 x 2 = 2.0
C6H12O6
O: 16.0 x 1 = 16.0
gfm = 30.0 g
molecular mass = 180 = 6
empirical mass
30
_____11. I can find the number of
moles of a substance if I am
given the mass and formula for
the substance.
How many moles of NaCl are equivalent to 94.3 g?
Na: 23.0 x 1 = 23.0
moles = mass x = 94.3 = 1.6 moles
Cl: 35.5 x 1 = 35.5
gfm
58.5
gfm = 58.5 g
_____12. I can find the mass of a
substance if I am given the moles
and formula for the substance.
How many grams of LiBr would 3.5 moles of LiBr be?
Li: 6.9 x 1 = 6.9
moles = mass
3.5 = x
Br: 79.9 x 1 = 79.9
gfm
86.8
gfm = 86.8 g
How many moles of O2 contain 3.01 x 1023 molecules?
_____13. I can convert between
moles and numbers of particles
using Avogadro’s number.
1 mole
=
x
6.02 x 1023
3.01 x 1023
x = 0.5 moles
How many molecules are in 0.25 moles of H2?
1 mole
=
6.02 x 1023
0.25
x
x = 1.5 x 1023 molecules
x = 303.8 g
How many liters does 4.6 moles of N2 occupy?
1 mole
22.4 L
_____14. I can convert between
moles and volume.
_____16. I can define
stoichiometry.
4.6
x
x = 103.04 L
How many moles of CO2 are present in an 11.2 L sample?
1 mole
22.4 L
_____15. I can answer questions
using Avogadro’s Hypothesis.
=
=
x
11.2
x = 0.5 moles
Which gas sample at STP has the same total number of molecules as 2 L of
CO2 (g) at STP?
A) 5 L of CO2 (g) B) 2 L of Cl2 (g) C) 3 L of H2S (g) D) 6 L of He (g)
Definition:
Stoichiometry: a proportional relationship between two or more
substances in a reaction.
Given the following balanced equation, state the mole ratios between the
following substances.
C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O ()
_____17. Given a balanced
equation, I can state the mole
ratios between any of the
reactants and/or products.
The mole ratio between C3H8 and O2 is 1 to 5.
The mole ratio between C3H8 and CO2 is 1 to 3.
The mole ratio between C3H8 and H2O is 1 to 4.
The mole ratio between CO2 and O2 is 3 to 5.
The mole ratio between H2O and CO2 is 4 to 3.
Using the equation from question #17, determine how many moles of O 2
are needed to completely react with 7.0 moles of C3H8.
_____18. Given the number of
moles of one of the reactants or
products, I can determine the
number of moles of another
reactant or product that is
needed to react.
5 O2
=
1 C3H8
x
7
x = 35 moles of O2
Using the equation from question #17, determine how many moles of CO2
are produced when 8.0 moles of H2O completely react.
3 CO2
4 H2O
=
x
8
x = 6 moles of CO2
Unit 8: Solutions
Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence.
Definitions:
Solute: the substance being dissolved; usually present in a smaller
amount. Ex – salt in a salt water solution.
_____1. I can define solute,
solvent, solution, and solubility.
Solvent: the substance that does the dissolving; usually present in a
greater amount. Ex – water in a salt water solution.
Solution: a homogeneous mixture of solute and solvent.
Solubility: the amount of solute that will dissolve in a certain amount of
solvent at a given temperature.
“Like dissolves like” means the polarity of the solute and the polarity of
the solvent must be the same in order for a substance to dissolve.
Nonpolar solutes dissolve in nonpolar solvents. Polar or ionic solutes
dissolve in polar solvents.
_____2. I can explain and apply
the expression “like dissolves
like.”
Explain, in terms of molecular polarity, why ammonia is more soluble than
methane in water at 20˚C at standard pressure.
Ammonia is a polar molecule and will dissolve in water which is also a
polar molecule. Methane is a nonpolar molecule and will not dissolve in
polar water.
_____3. I can describe the trend
in solubility for solids and liquids
as the temperature changes.
_____4. I can describe the trend
in solubility for gases as the
temperature changes.
_____5. I can describe the trend
in solubility for gases as the
pressure changes.
_____6. I can describe the
conditions in which gases are
most soluble.
As the temperature increases, the solubility of solids and liquids increases.
As the temperature increases, the solubility of a gas decreases.
As the pressure increases, the solubility of a gas increases.
Gases are most soluble in low temperatures and high pressures.
Write “S” for soluble and “NS” for not soluble.
_____7. I can use Table F to
determine if a substance will be
soluble in water.
_____8. I can use Table G to
determine how much solute to
add at a given temperature to
make a saturated solution.
S potassium chlorate
NS silver bromide
S lithium carbonate
NS calcium carbonate
How many grams of KClO3 must be dissolved in 100 grams of water at
20˚C to make a saturated solution?
8 grams (must be on the KClO 3 line to be saturated)
_____9. I can use Table G to
determine if a solution is
saturated, unsaturated, or
supersaturated.
If 20.0 g of NaNO3 are dissolved in 100.0 g of water at 25.0˚C, will the
resulting solution be saturated, unsaturated, or supersaturated?
Unsaturated (point below the line)
Saturated solutions are on the line, supersaturated solutions are above.
Definitions:
Concentration: amount of dissolved solute in a given amount of solvent.
_____10. I can define
concentration, dilute, and
concentrated.
Dilute: a solution that contains a small amount of dissolved solute.
Concentrated: a solution that contains a large amount of dissolved
solute.
What is the molarity of 3.5 moles of NaBr dissolved in 500 mL of water?
Molarity = moles
L
_____11. I can calculate the
concentration of a solution in
molarity.
x = 3.5
.5
x=7M
What is the molarity of 1.5 L of a solution containing 52 g of LiF?
moles = mass
gfm
Molarity = moles
L
x = 52
25.9
x=2
1.5
x = 2 moles
x = 1.3 M
Which solution is most concentrated?
_____12. I can determine which
solution is the most concentrated
or the most dilute.
_____13. I can calculate the
concentration of a solution in
ppm.
A) 1 mole of solute dissolved in 1 liter of solution
B) 2 moles of solute dissolved in 3 liters of solution
C) 6 moles of solute dissolved in 4 liters of solution
D) 4 moles of solute dissolved in 8 liters of solution
What is the concentration, in ppm, of a 2,600 g solution containing 0.015 g
of CO2?
ppm = mass of solute x 1,000,000 ppm = 0.015 x 1,000,000 = 5.77 ppm
mass of solution
2,600
_____14. I can explain how
adding a solute to a pure solvent
affects the freezing point of the
solvent.
When a solute is added to a solvent, the freezing point of the solvent
____15. I can explain how
adding a solute to a pure solvent
affects the boiling point of the
solvent.
When a solute is added to a solvent, the boiling point of the solvent
____16. I can determine which
solution will have the lowest
freezing point/highest boiling
point.
Which solution has the highest boiling point?
decreases. The more solute that is added, the lower the freezing point gets.
increases. The more solute that is added, the higher the boiling point gets.
A) 1.0 M KNO3 B) 2.0 M KNO3 C) 1.0 M Ca(NO3)2 D) 2.0 M Ca(NO3)2
Unit 9: Acids & Bases
Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence.
_____1. I can use two different
systems to define acids and
bases.
_____2. I can determine which
substance is an acid or base
according to the alternative
theory given a reaction.
Arrhenius Theory
Alternative Theory
(Bronsted-Lowry)
Acid
Produces H+ ions (H3O+)
Donates H+ ions
Base
Produces OH- ions
Accepts H+ ions
Given the reaction: HI + H2O  H3O+ + I-, identify the acid and base.
Acid: HI
Base: H2O
List the chemical names of three common acids and three common bases.
_____3. I can give examples of
the chemical names of common
acids and bases.
Acids (Table K)
Bases (Table L)
Hydrochloric Acid
Sodium Hydroxide
Nitric Acid
Potassium Hydroxide
Sulfuric Acid
Calcium Hydroxide
List the chemical formulas of three common acids and three common bases.
_____4. I can give examples of
the chemical formulas of
common acids and bases.
_____5. I can list properties of
common acids and bases.
Acids (Table K)
Bases (Table L)
HCl
NaOH
HNO3
KOH
H2SO4
Ca(OH)2
Acids
Taste sour, electrolytes, react
with bases to form water and salt
in a neutralization reaction, react
with active metals (higher than H2
on Table J) to produce hydrogen
gas and salt, cause indicators to
change color
Bases
Taste bitter, feel slippery,
electrolytes, react with acids to
form water and salt in a
neutralization reaction, do not
react with metals, cause
indicators to change color
Definitions:
Electrolyte: any substance that conducts electricity when dissolved in
water; produces mobile charges.
_____6. I can define electrolyte,
nonelectrolyte, hydronium ion,
and hydroxide ion.
Nonelectrolyte: any substance that does not conduct electricity when
dissolved in water; does not produce mobile charges.
Hydronium Ion: H3O+, also known as a hydrogen ion; is produced by
acids and is present in larger amounts in acidic solutions.
Hydroxide Ion: OH-; is produced by bases and is present in larger
amounts in basic solutions.
_____7. I can state another name
for the hydronium ion.
_____8. I can state the two types
of substances that are
nonelectrolytes.
_____9. I can state the three
types of substances that are
electrolytes.
_____10. I can identify
nonelectrolytes.
_____11. I can identify
electrolytes.
The hydronium ion is also known as the hydrogen ion (H+).
Sugars and alcohols are two classes of compounds that are nonelectrolytes.
Acids, bases, and salts are three classes of compounds that are electrolytes.
Circle all of the substances that are nonelectrolytes.
A) C2H5OH
B) C6H12O6
C) C12H22O11
D) CH3COOH E) LiOH
Which substance, when dissolved in water, forms a solution that conducts
an electric current?
A) C2H5OH
B) C6H12O6
C) C12H22O11
D) CH3COOH
Predict the products in the following reactions, then balance.
Zn + 2 HCl  H2 + ZnCl2
_____12. I can complete
reactions between metals and
acids.
3 Mg + 2 H3PO4  3 H2 + Mg3(PO4)2
Cu + H2SO4  No Reaction
_____13. I can define
neutralization.
Definition:
Neutralization: a double replacement reaction that occurs between an
acid and a base producing water and a salt.
Which of the following equations is a neutralization reaction?
_____14. I can identify a
neutralization reaction from a
list of reactions.
A) 6 Na + B2O3  3 Na2O + 2 B
B) Mg(OH)2 + 2 HBr  MgBr2 + 2 H2O
C) 2 H2 + O2  2 H2O
D) 2 KClO3  2 KCl + 3 O2
Definitions:
pH: the concentration of hydrogen ions in solution; measures how acidic
or basic a solution is.
_____15. I can define pH, pOH,
and [ ].
pOH: the concentration of hydroxide ions in solution.
[
_____16. I can describe the
relationship between pH and
pOH and [H+] and [OH-].
]: concentration
pH + pOH = 14
[H+] and [OH-] are indirectly related.
If the pH of a solution is 4.5, the solution is acidic.
_____17. Based on pH, I can
determine if a solution is acidic,
basic, or neutral.
If the pH of a solution is 7.0, the solution is neutral.
If the pH of a solution is 11, the solution is basic.
If the pH of a solution is 5.7, the solution is acidic.
_____18. I can identify the pH
values of strong acids and strong
bases.
Strong acids have low pH values and strong bases have high pH values.
In an acidic solution the [H+] is greater than the [OH-].
_____19. I can describe acidic,
basic, or neutral solutions in
terms of ion concentrations.
In a basic solution the [H+] is less than the [OH-].
In a neutral solution the [H+] is equal to the [OH-].
As the H+ concentration decreases, the pH increases.
_____20. I can state the
relationship between H+
concentration and pH.
As the H+ concentration increases, the pH decreases.
A greater number of hydrogen ions results in a lower pH, indicating a
strong acid.
If the [H3O+] is 1 x 10-8, the pH of the solution will be 8.
If the [H3O+] is 1 x 10-1, the pH of the solution will be 1.
_____21. Given the hydronium
ion concentration, I can
determine the pH.
If the [H3O+] is 1 x 10-14, the pH of the solution will be 14.
If the [H3O+] is 1 x 10-7, the pH of the solution will be 7.
If the pH is 10, the [H+] is 1 x 10-10 M.
_____22. Given the pH, I can
determine the [H+] or [OH-].
If the pH is 3, the [H+] is 1 x 10-3 M.
If the pH is 6, the [OH-] is 1 x 10-8 M.
If the pH is 9, the [OH-] is 1 x 10-5 M.
If the H+ concentration is increased by a factor of 10,
the pH will decrease by 1.
_____23. I can determine the
change in pH when the [H+] of a
solution is changed.
If the H+ concentration is increased by a factor of 100,
the pH will decrease by 2.
If the H+ concentration is decreased by a factor of 1,000,
the pH will increase by 3.
_____24. I can answer indicator
questions using Table M.
Phenolphthalein tests colorless and bromthymol blue tests blue in samples
of fish tank water. What is the pH range for the water in this fish tank?
7.6 – 8.0
What color is thymol blue in a pH of 11? Blue
_____25. I can state the name of
the laboratory equipment that is
used to carry out a titration.
Which piece of laboratory equipment is used to carry out a titration?
Burette
_____26. I can state the purpose
of titration.
Why do scientists do titrations? To determine the concentration (molarity)
of an unknown acid or base.
In a titration, 36 mL of a 0.16 M NaOH solution exactly neutralizes 12 mL
of an H2SO4 solution. What is the concentration of the H2SO4 solution?
(H+)MAVA = MBVB(OH-)
(2)(MA)(12) = (0.16)(36)(1)
MA = 0.24 M
A student uses a 1.4 M HBr solution in a titration to determine the
concentration of a KOH solution. The data is below:
_____27. I can solve for any
variable in the titration equation.
What is the molarity of the KOH solution?
(H+)MAVA = MBVB(OH-)
(1)(1.4)(15.4) = (MB)(22.1)(1)
MB = 0.98 M
(subtract the volumes to get the volume used – plug those numbers in)
Unit 10: Kinetics & Equilibrium
Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence.
_____1. I can define an effective
collision according to the
collision theory.
_____2. I can determine the type
of compound that has a faster
reaction rate.
Definition:
Effective Collision: a collision between reactants that have enough energy
and proper angles; results in a reaction.
Ionic substances react faster than covalent molecules
because their reactions involve less bond rearrangement.
As the concentration increases, the reaction rate increases because there are
_____3. I can state and apply the
relationship between
concentration and reaction rate
in terms of the collision theory.
more effective collisions between particles.
At 20˚C, a reaction between powdered Zn(s) and hydrochloric acid will
occur most quickly if the concentration of the HCl is
A) 1.0 M
B) 1.5 M
C) 2.5 M
D) 2.8 M
As the surface area increases, the reaction rate increases because there are
_____4. I can state and apply the more effective collisions between particles.
relationship between surface area
and reaction rate in terms of the
At STP, which 4.0 g sample of Zn(s) will react most quickly with dilute
collision theory.
hydrochloric acid?
A) lump
B) bar
C) powdered
D) sheet metal
As the temperature increases, the reaction rate increases because there are
_____5. I can state and apply the
relationship between
temperature and reaction rate in
terms of the collision theory.
more effective collisions between particles.
Given the reaction: 2 Mg (s) + O2 (g)  2 MgO (s)
At which temperature would the reaction occur at the greatest rate?
A) 0˚C
_____6. I can use Table I to
determine if a reaction is
exothermic or endothermic.
B) 15˚C
C) 95˚C
D) 273 K
_____7. Based on the location of
the energy term, I can determine
if the reaction is exothermic or
endothermic.
_____8. I can determine the new
heat of reaction if the number of
moles in a reaction changes.
Given the following balanced equation: I + I  I2 + 146.3 kJ
Is this reaction exothermic or endothermic? Justify your answer.
Exothermic – energy is released (it is on the right of the arrow – product).
Given the following balanced equation: 2 C + 3 H2  C2H6, what is the
heat of reaction if 3 moles of carbon react?
2C = 3
-84
x
x = -126 kJ
Definitions:
Potential Energy Diagram: a graph that shows the changes in potential
energy over the course of a chemical reaction.
_____9. I can define potential
energy diagram, heat of reaction
(∆H), activation energy, and
catalyst.
Heat of Reaction (∆H): the difference between the potential energy of the
reactants and the potential energy of the products; Hproducts – Hreactants.
Activation Energy: the minimum amount of energy needed for a chemical
reaction to occur.
Catalyst: a substance that makes a reaction occur faster by lowering the
activation energy and providing an alternate pathway for the reaction to
occur.
Given the potential energy diagram below, determine the following:
_____10. Given a potential
energy diagram, I can determine
the PE of reactants, PE of
products, and PE of the activated
complex, ∆H, and activation
energy of forward and reverse
reactions.
PE of reactants = 40 kJ
PE of products = 15 kJ
PE of activated complex = 50 kJ
∆H= -25 kJ
Activation Energy of Forward Reaction = 10 kJ
Activation Energy of Reverse Reaction = 35 kJ
_____11. Given a potential
energy diagram for an
uncatalyzed reaction, I can
determine how the diagram will
change when a catalyst is added.
Draw a dotted line on the potential energy diagram below to indicate how it
will change if a catalyst is added. List 3 values that will change and 3 values
that will not change when a catalyst is added.
Will Change: PE of activated
complex, activation energy of
forward reaction, activation
energy of reverse reaction
Will Not Change: PE of
reactants, PE of products, heat
of reaction (∆H)
Give the potential energy diagram below, determine if the reaction is
exothermic or endothermic. Justify your answer.
Endothermic – energy of the products
is greater which shows that energy
was absorbed during the reaction.
The difference in PE of the reactants
and the PE of the products is positive.
_____12. Given a potential
energy diagram, I can determine
if the reaction is exothermic or
endothermic.
In a system at equilibrium the rates of the forward and reverse reaction
_____13. I can state two
conditions that apply to all
systems at equilibrium.
must be equal and the concentrations of the reactants and products must be
constant.
Two types of physical equilibrium are phase equilibrium, in which phase
_____14. I can describe examples
of physical equilibrium.
changes occur at the same rate, and solution equilibrium. In terms of
saturation, a solution that is at equilibrium must be saturated.
Definitions:
Forward Reaction: the chemical reaction read from left to right.
_____15. I can define forward
reaction, reverse reaction,
reversible reaction, and closed
system.
Reverse Reaction: the chemical reaction read from right to left.
Reversible Reaction: a chemical reaction that can occur in both directions
(forward and reverse).
Closed System: a system in which reactants and products may not enter
or leave.
_____16. I can state
LeChatelier’s Principle.
LeChatelier’s Principle states when a reaction at equilibrium is subjected
to a stress, it will shift (favor the forward or reverse reaction) to relieve
the stress.
Given the reaction at equilibrium:
2 SO2 (g) + O2 (g)
_____17. Given a balanced
equation at equilibrium, I can
predict the direction of shift in
the equilibrium when the
concentration, temperature, or
pressure is changed or if a
catalyst is added. I can predict
which substances will increase in
concentration and which
substances will decrease in
concentration due to the shift.
_____18. I can define entropy.
_____19. I can rank the three
phases of matter from least
entropy to most entropy.
2 SO3 (g) +
392 kJ
Predict the direction of shift in the equilibrium (right, left, none) when the
following changes are made to the system. Then predict which substances
from the above reaction will increase and decrease due to the shift.
Increase [SO2]
Right
Increased
Concentration
of
SO3
Increase [SO3]
Left
SO2, O2
SO3
Decrease [O2]
Increase
temperature
Decrease
temperature
Increase pressure
Left
SO2, O2
SO3
Left
SO2, O2
SO3
Right
SO3
SO2, O2
Right
SO3
SO2, O2
Decrease pressure
Left
SO2, O2
SO3
Add a catalyst
None
None
None
Change
Direction of
Shift
Decreased
Concentration
of
SO2, O2
Entropy is the amount of randomness or disorder.
Least entropy
Most entropy
Solid < Liquid < Gas
_____20. Given a balanced
equation, I can determine if the
reaction results in an overall
increase or decrease in entropy.
_____21. I can state the trends in
nature for entropy and energy.
Most systems in nature tend to undergo reactions that have a(n)
increase in entropy and a(n) decrease in energy.
Unit 11: Redox
Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence.
_____1. I can assign oxidation
numbers to any element or ion.
Assign oxidation number to each of the elements below.
O2 0
Li 0
Na+ +1
Assign oxidation numbers to each element in the compounds below.
_____2. I can assign oxidation
numbers to the elements in a
compound.
MnCl3: Mn +3
Cl -1
H2SO4: H +1
S +6
O -2
Assign oxidation numbers to each element in the polyatomic ions below.
_____3. I can assign oxidation
numbers to the elements in a
polyatomic ion.
_____4. I can state the Law of
Conservation of Charge.
PO4-3:
P +5
O -2
ClO3-1:
Cl +5
O -2
The Law of Conservation of Charge states in any reaction, the total charge
of the reactants must be the same as the total charge of the products.
Definitions:
Oxidation: the loss of electrons; the oxidation number will increase.
Reduction: the gain of electrons; the oxidation number will decrease.
_____5. I can define oxidation,
reduction, redox reaction,
oxidizing agent, and reducing
agent.
Redox Reaction: a reaction in which there is a transfer of electrons; one
substance will lose electrons and another substance will gain electrons.
Both oxidation and reduction occur simultaneously.
Oxidizing Agent: the substance that gets reduced.
Reducing Agent: the substance that gets oxidized.
_____6. I can identify a redox
reaction from a list of chemical
reactions.
Which half-reaction equation represents the reduction of a potassium ion?
A) K+ + 1 e-  K0
C) K+  K0 + 1 eB) K0 + 1 e-  K+
D) K0  K+ + 1 e-
_____7. I can distinguish
between an oxidation halfreaction and a reduction halfreaction.
The two half-reactions that come from the following equation are:
Li0 (s) + Ag+ (aq)  Li+ (aq) + Ag0 (s)
_____8. I can break a redox
reaction into its two halfreactions.
Oxidation Half-Reaction: Li0 (s)  Li+ (aq) + 1eReduction Half-Reaction: Ag+ (aq) + 1e-  Ag0 (s)
Given the reaction: Mg0 (s) + Fe+3 (aq)  Fe0 (s) + Mg+2 (aq)
When the equation is correctly balanced using smallest whole numbers, the
coefficient of Fe+3 will be
A) 1
_____9. I can balance a redox
reaction.
B) 2
C) 3
D) 6
_____10. I can describe the
conditions needed for a
spontaneous redox reaction to
occur.
A spontaneous redox reaction will occur if the metal being oxidized is
_____11. I can state the two
types of electrochemical cells.
The two types of electrochemical cells are: voltaic and electrolytic.
higher on Table J, or the nonmetal reduced is higher on Table J.
Voltaic
Electrolytic
One half-cell with the
anode, one half-cell
with the cathode, salt
bridge
One container with the
anode and cathode,
battery
Oxidation Site
Anode
Anode
Reduction Site
Cathode
Cathode
Anode Charge
Negative
Positive
Cathode Charge
Positive
Negative
Mass at Anode
Decreases
Decreases
Mass at Cathode
Increases
Increases
Anode  Cathode
Anode  Cathode
Chemical  Electrical
Electrical  Chemical
Spontaneous
Not Spontaneous
Batteries
Electroplating
Components
_____12. I can compare the two
types of electrochemical cells.
Direction of
Electron Flow
Energy Conversion
Nature of Process
Use
_____13. I can state the purpose
of the salt bridge in a voltaic cell.
The purpose of the salt bridge is to allow for the migration of ions; it
balances the charges and maintains neutrality in the solutions.
_____14. I can answer questions
about a voltaic cell.
What is the anode? Zn0
What is the cathode? Cu0
Write the oxidation half-reaction: Zn 0  Zn+2 + 2 eWrite the reduction half-reaction: Cu+2 + 2 e-  Cu 0
In what direction do electrons flow? Anode  Cathode (Zn0 to Cu 0)
Explain, in terms of atoms and ions, why the mass of the cathode increases
during the operation of an electrochemical cell.
_____15. I can explain, in terms
of atoms and ions, the changes in
mass that take place at the anode
and cathode of an
electrochemical cell.
The ions from the solution are reduced and become atoms that become
part of the cathode.
Explain, in terms of atoms and ions, why the mass of the anode decreases
during the operation of an electrochemical cell.
The atoms from the anode are oxidized producing ions that move into the
solution.
The diagram below represents an electrochemical cell:
_____16. I can answer questions
about an electrolytic cell.
What is the anode? Ag0
What is the cathode? Spoon
Write the oxidation half-reaction: Ag0  Ag+ + 1 eWrite the reduction half-reaction: Ag+ + 1 e-  Ag0
In what direction do electrons flow? Anode  Cathode (Ag to spoon)
What is the purpose of the battery? Provide energy for the reaction.
Unit 12: Organic Chemistry
Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence.
_____1. I can state the name and
symbol of the element that is
capable of forming rings, chains,
and networks.
The element that is capable of forming rings, chains, and networks is
carbon. Its symbol is C.
Definitions:
Organic Compound: a compound containing a carbon-hydrogen covalent
bond.
_____2. I can define organic
compound, saturated
hydrocarbon, unsaturated
hydrocarbon, isomers, and alkyl
group.
Saturated Hydrocarbon: a compound containing only carbon and
hydrogen in which all carbon atoms are held together by single bonds.
Unsaturated Hydrocarbon: a compound containing only carbon and
hydrogen in which there is at least one double or triple bond (multiple
bond) between the carbon atoms.
Isomers: molecules with the same molecular formula, but they have
different structures, properties, and names.
Alkyl Group: a hydrocarbon branch off of the main chain.
_____3. Given the formula, I can
determine if a compound is a
hydrocarbon.
Draw the complete structural formula for CH3CH2CH2CH2CH3.
_____4. I can expand a
condensed structural formula to
show the structural formula of an
organic compound.
_____5. Given the formula, I can
determine which compound has
the highest boiling point.
_____6. Given the name or
formula, I can use Table Q to
determine which class of
hydrocarbons a compound
belongs.
Draw the complete structural formula for CH3CHCHCH3.
Which of the following hydrocarbons has the highest boiling point?
A) C3H8
B) C4H10
C) CH4
D) C2H6
Determine the homologous series of hydrocarbons to which each of the
following belongs:
Decane: Alkane
2-Decene: Alkene
C4H8: Alkene
C7H12: Alkyne
_____7. Given the name or
formula, I can determine if the
hydrocarbon is saturated or
unsaturated.
Determine if each of the following is a saturated or unsaturated
hydrocarbon.
Decane: saturated
2-Decene: unsaturated
CH3CH2CH2CH3: saturated
CH3CHCHCH3: unsaturated
Determine the homologous series of hydrocarbons to which each of the
following belongs:
belongs to the alkane series.
_____8. Given the structural
formula, I can determine which
homologous series a hydrocarbon
belongs to.
belongs to the alkyne series.
belongs to the alkene series.
Determine how many carbon atoms are in each of the following
compounds:
_____9. Given the name, I can
use Table P to determine how
many carbons atoms are in a
compound.
Decane: 10
Ethene: 2
3-Nonene: 9
1-Pentyne: 5
Propene: 3
Methane: 1
_____10. Given a list of
compounds, I can determine
which ones are isomers.
Draw an isomer of the following:
_____11. I can draw an isomer of
a given compound.
Name the following hydrocarbons:
2-pentene
3-heptyne
nonane
_____12. I can use Tables P & Q
to name hydrocarbons.
2-hexene
2-methyl hexane
2,2,3-trimethyl butane
Draw the following hydrocarbons:
propyne
_____13. I can use Tables P & Q
to draw hydrocarbons.
3-ethyl hexane
3-hexene
2,3-dimethyl butane
_____14. Given a structural
formula, I can identify which
class of organic compounds a
substance belongs to.
Name the following organic compounds:
_____15. I can use Tables P & R
to name simple compounds in
any of the classes of organic
compounds.
butane
butanone
butanoic acid
2-butene
1-butanol
1-butanamine
2-butyne
butanal
butanamide
1-bromobutane
methyl propanaote
propyl methyl ether
Draw the following organic compounds:
2,2-dichloropentane
dimethyl ether
_____16. I can use Tables P & R
to draw simple compounds in
any of the classes of organic
compounds.
2-pentanone
methyl butanoate
pentanamide
2-butanol
ethanal
propanoic acid
ethanamine
ethyl propanoate
Substitution (halogenation) is the replacement of a hydrogen in a saturated
hydrocarbon with a halogen. It starts with an alkane and ends with two
product(s).
_____17. I can describe the 7
types of organic reactions.
Addition is the addition of hydrogens or halogens at a double bond in an
unsaturated hydrocarbon. It starts with an alkene and ends with one
product(s).
Esterification is when a(n) organic acid and a(n) alcohol react to form an
ester and water.
Saponification is the production of soap.
Fermentation is the breakdown of sugar into alcohol and carbon dioxide.
Combustion is when a hydrocarbon reacts with oxygen to produce water
and carbon dioxide.
Polymerization is the formation of a large molecule called a polymer
from small molecules called monomers. When monomers are joined
together by breaking double or triple bonds it is called addition
polymerization. When monomers are joined by the removal of water it is
called condensation polymerization.
_____18. Given an equation, I
can identify the type of organic
reaction that is occurring.
Complete the following reactions:
_____19. I can complete an
organic reaction by predicting
the missing products.