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.
How many joules are in 9.3 kJ?
_____2. I can convert between
different metric units.
How many grams are in 39,983 mg?
How many milliliters are in 452 L?
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.
_____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.
What is the density of a 27.8 g substance with a volume of 2.3 mL?
How many significant figures are there in 30.50 cm?
How many significant figures are there in 400 sec?
To the correct number of significant figures, what is the answer to
5.93 mL + 4.6 mL?
To the correct number of significant figures, what is the answer to
5.93 cm * 4.6 cm?
_____6. I can read the meniscus
on a graduated cylinder to the
correct number of significant
figures.
The volume is_______________ mL.
_____7. I can convert numbers
into scientific notation from
standard notation.
Convert 87,394,000,000,000 to scientific notation.
_____8. I can convert numbers
into standard notation from
scientific notation.
Convert 5.8 x 10 to standard notation.
Convert 0.0000040934 to scientific notation.
9
-5
Convert 4.3 x 10 to standard notation.
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:
Element:
_____1. I can define the terms
atom, element, compound, and
mixture.
Compound:
Mixture:
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.
Examples: C12H22O11, NaCl, Fe, salt water, air, CO2, H2, Ar, soda
Element
Compound
Mixture
The seven diatomic elements are:
Which of the following is a binary compound?
A) NaOH
B) H3PO4
C) H2O
D) MgSO4
Definitions:
Homogeneous Mixture:
_____5. I can define
homogeneous mixture and
heterogeneous mixture.
Heterogeneous Mixture:
Two examples of Homogeneous Mixtures:
A)
_____6. I can give examples of
homogeneous and
heterogeneous mixtures.
B)
Two examples of Heterogeneous Mixtures:
A)
B)
_____7. I can determine how
matter will be separated using
filtration.
When a mixture of sand, salt, sugar, and water is filtered, what passes
through the filter?
_____8. I can describe how
matter can be separated using
distillation.
Which physical property makes it possible to separate the components of
crude oil by means of distillation?
To separate a mixture of salt and water, use __________________________.
_____9. I can state which
separation process is best for a
given situation.
To separate a mixture of ethanol and water, use _______________________.
To separate a mixture of sand and water, use__________________________.
Monoatomic Element
Diatomic Element
____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
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.
Write “P” for physical or “C” for chemical on the line provided.
_____ Copper (II) sulfate is blue.
_____ Copper reacts with oxygen.
_____12. I can classify a
_____ Copper can be made into wire.
property as physical or chemical.
_____ Copper has a density of 8.96 g/cm3.
_____ Copper melts at 1358 K.
_____ Copper doesn’t dissolve in water.
Write “P” for physical or “C” for chemical on the line provided.
_____ Copper (II) sulfate dissolves in water.
_____ Copper reacts with oxygen to form solid copper (I) oxide.
_____13. I can classify a change
as physical or chemical.
_____ Solid copper is melted.
_____ A chunk of copper is pounded flat.
_____ Copper and zinc are mixed to form brass.
_____ A large piece of copper is chopped in half.
_____ 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
____15. I can use particle
diagrams to show the
arrangement and spacing of
atoms/molecules in different
phases.
Liquid
Solid
Shape
_____16. I can compare solids,
liquids, and gases in terms of
their shape, volume, particle
distance, and attraction between
particles.
Volume
Particle
Distance
Attraction
Between
Particles
Gas
Liquid
Gas
Fusion: change from _________________ to _________________.
Solidification: change from _________________ to _________________.
Condensation: change from _________________ to _________________.
Vaporization: change from _________________ to _________________.
_____17. I can state the change
of phase occurring during each
phase change.
Melting: change from _________________ to _________________.
Boiling: change from _________________ to _________________.
Sublimation: change from _________________ to _________________.
Deposition: change from _________________ to _________________.
Freezing: change from _________________ to _________________.
_____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:
Potential Energy:
_____19. I can define kinetic
energy, potential energy,
temperature, heat, endothermic,
and exothermic.
Temperature:
Heat:
Endothermic:
Exothermic:
Indicate whether the change is exothermic or endothermic:
fusion/melting _______________________
solidification/freezing _______________________
_____20. I can indicate if a
phase change is exothermic or
endothermic.
condensation _______________________
vaporization/boiling _______________________
sublimation _______________________
deposition _______________________
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 of heat flow.
A) Heat flows from object A to object B
B) Heat flows from object B to object A
C) Heat flows in both directions between the objects
D) No heat flow occurs between the objects
What Kelvin temperature is equal to 200˚C?
_____23. I can convert between
Celsius and Kelvin.
What Celsius temperature is equal to 200 K?
_____24. Given a
heating/cooling curve, I can
identify the process occurring
during each segment.
AB = _____________________________________________________
BC = _____________________________________________________
CD = _____________________________________________________
DE = _____________________________________________________
EF = _____________________________________________________
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?
What is the boiling point of this substance?
_____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?
What is the melting point of water in ˚C and K?
What is the boiling point of water in ˚C and K?
What is the condensing point of water in ˚C and K?
Energy is measured in ____________________.
Definitions:
Specific Heat Capacity:
_____34. I can define specific
heat capacity, heat of fusion,
and heat of vaporization.
Heat of Fusion:
Heat of Vaporization:
How much heat is needed to vaporize 10 g of water at 100˚C?
How much heat is needed to change the temperature of 100.0 g of water
from 20˚C to 35˚C?
_____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?
How many grams of water can be heated by 15˚C using 13,500 J of heat?
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?
The five parts of the Kinetic Molecular Theory are:
A.
B.
_____36. I can state the 5 parts
of the Kinetic Molecular
Theory.
C.
D.
E.
_____37. I can define an ideal
gas.
Definition:
Ideal Gas:
_____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 _____________ &
________________________ and some _________________________.
_______________.
A real gas will act most ideally under the conditions of __________ pressure
and ____________ temperature. (PLIGHT)
STP stands for ___________________________________________________.
_____41. I can state the
meaning of “STP” and the
Table on which it can be found.
The values can be found on Table______________.
Standard Pressure = _______________________________________.
Standard Temperature = ____________________________________.
Avogadro’s Hypothesis says_______________________________________
________________________________________________________________
_____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
____________________.
____________________.
increases, the pressure____________________.
Draw a curve for each of the graphs below:
_____46. I can recognize the
relationships stated above in
graph form.
V
P
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?
_____47. I can solve problems
using the combined gas law.
Definitions:
Vapor Pressure:
_____48. I can define vapor
pressure, boiling point, and
volatile.
Boiling Point:
Volatile:
The normal boiling point of a substance occurs at a pressure of
_____49. I can state the
conditions of temperature and
pressure that are used for normal
boiling points.
___________ atm or _______________kPa. This is known as
____________________________________and can be found on Table _____.
_____50. I can state the
relationship between
temperature and vapor pressure.
As the temperature of a substance increases, its vapor pressure
_____51. I can state the
relationship between
atmospheric pressure and
boiling point.
As the atmospheric pressure increases, the boiling point
_____52. I can state the
relationship between molecular
attraction and boiling point.
As the attraction between molecules increases, the 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?
_____54. I can determine the
vapor pressure of ethanol,
ethanoic acid, propanone, or
water at a given temperature.
_____55. I can determine the
temperature of ethanol, ethanoic
acid, propanone, or water at a
given vapor pressure.
What is the vapor pressure of ethanol at 56˚C?
____________________________.
____________________________.
____________________________.
When the vapor pressure of water is 30 kPa, what is the temperature of the
water?
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.
Dalton’s Model:
_____1. I can describe John
Dalton’s contribution to our
understanding of the atom.
Diagram:
Thomson’s Model:
_____2. I can describe JJ
Thomson’s contribution to our
understanding of the atom.
Diagram:
Rutherford’s Experiment:
Experiment Results:
_____3. I can describe Ernest
Rutherford’s contribution to our
understanding of the atom.
Experiment Conclusions:
Diagram:
Bohr’s Model:
_____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?
_____5. I can describe the
modern model (wave mechanical
model) of the atom.
Diagram:
Particle #1
Particle #2
Particle #3
Name
_____6. I can state the three
subatomic particles, their
location in an atom, and the
magnitude of their charges and
their masses (in amu).
Charge
Mass
Location in
Atom
_____7. I can identify the nuclear
charge of an atom.
_____8. I can explain why atoms
are electrically neutral.
What is the charge of an oxygen-18 nucleus?
Atoms are electrically neutral because the number of ________________ is
equal to the number of _____________________.
Definitions:
Atomic Number:
_____9. I can define atomic
number and mass number.
Mass Number:
_____10. I can use the Periodic
Table to determine the atomic
number of an element.
How many protons are in an atom of selenium?
_____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?
_____12. I can represent an atom
using different notations.
_____13. Given the mass
number, I can determine the
number of protons, neutrons,
and electrons in an atom.
How many protons are in an atom of silicon?
What are three other notations that can be used to represent 80Br?
In an atom of 212Po, how many protons are present?
In an atom of 212Po, how many electrons are present?
In an atom of 212Po, how many neutrons are present?
Definitions:
Ion:
_____14. I can define ion,
cation, and anion.
Cation:
Anion:
How many protons are in 18O-2?
How many neutrons are in 18O-2?
_____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?
What is this ion’s electron configuration?
How many electrons are in 27Al+3?
What is this ion’s electron configuration?
_____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:
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?
Definitions:
Principal Energy Level:
Orbital:
_____19. I can define principal
energy level, orbital, ground
state, excited state, electron
configuration, and bright line
spectrum.
Ground State:
Excited State:
Electron Configuration:
Bright Line Spectrum:
The 1st shell holds a maximum of ___________ 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 ___________ electrons.
The 3rd shell holds a maximum of ___________ electrons.
The 4th shell holds a maximum of ___________ electrons.
_____21. I can define valence
electrons.
_____22. I can define a stable
octet.
Definition:
Valence Electrons:
Definition:
Stable Octet:
How many valence electrons does rubidium have in the ground state?
_____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?
How many electrons does an aluminum atom have in its 2nd shell?
_____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.
_____26. I can identify an
electron configuration that
shows an atom in the excited
state.
As the distance between the nucleus and the electron increases, the energy
_____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
of the electron ____________________________.
As the number of the principal energy level increases, the distance to the
nucleus ____________________________.
Which electron configuration represents an atom of potassium in the
excited state?
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?
_____29. I can draw Lewis
electron dot diagrams for a given
element.
Draw the Lewis electron dot diagram for the following atoms:
Li
Be
B
C
N
O
F
Ne
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 ______________ to
_________________ ratio.
Type
Symbol
Notation
Charge
Mass
Penetrating
Power
alpha
_____2. I can compare types of
radiation.
beta
positron
gamma
The difference between natural transmutation and artificial transmutation is
that in natural transmutation an_____________
_____3. I can explain the
difference between natural
transmutation and artificial
transmutation.
__________breaks apart
on its own becoming a different element and in artificial transmutation a
_____________ ___________ is made ________________ 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.
_____4. I can state two synonyms Two synonyms for spontaneous decay are:___________________________
for spontaneous decay.
and _________________________________.
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.
Definition:
Half-life:
_____8. I can define half-life.
Based on Reference Table N, what fraction of a radioactive sample of Au198 will remain unchanged after 10.78 days?
_____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? (The half-life of K-37 is 1.23 seconds)
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?
_____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.
_____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?
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
Definitions:
Fission:
_____14. I can define fission and
fusion.
Fusion:
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:
_______________ temperature and __________________ 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.
B.
_____19. I can state 5 beneficial
uses for radioactive isotopes.
C.
D.
E.
Tc-99 and Co-60 are used for_______________________________________
_____20. I can state the scientific
use of specific radioactive
isotopes.
I-131 is used for_________________________________________________
C-14 is used for _________________________________________________
U-238 is used for ________________________________________________
Three risks associated with radioactivity and radioactive isotopes are:
A.
_____21. I can state three risks
associated with radioactivity and
radioactive isotopes.
B.
C.
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
___________________ ____________________.
Group 1 is called the _________________________________________.
Group 2 is called the _________________________________________.
_____2. I can state the group
names for elements in groups 1,
2, 3-12, 17, and 18.
Groups 3-12 are called the ____________________________________.
Group 17 is called the ________________________________________.
Group 18 is called the ________________________________________.
_____3. I can explain why
elements in the same group have
similar chemical properties.
Elements in the same group have similar chemical properties because
_____4. I can explain what
elements in the same period have
in common.
Elements in the same period have the same number of
Definitions:
Electronegativity:
_____5. I can define
electronegativity, ionization
energy, atomic radius, and ionic
radius.
Ionization Energy:
Atomic Radius:
Ionic Radius:
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.
_______B
________K
________Li
________C
_______Ar
_______Sb
________H
________Fe
________Au
_______S
_______F
________Si
________Fr
________He
_______Rn
_______Ge
________Al
________As
________Bi
_______I
_____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:
_____________________________ and ___________________________.
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.
_____________________________, _______________________________,
_____________________________, _______________________________,
_____________________________, _______________________________,
and _____________________________.
The properties of metals are:
_____9. I can list properties of
metals.
The properties of nonmetals are:
_____10. I can list properties of
nonmetals.
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 ____________________.
The most reactive metal is _________________________________.
As one reads down a group from top to bottom, the activity/reactivity of
NONMETALS __________________.
The most reactive nonmetal is ______________________________.
Metals lose electrons and become positive ions that are ________________
_____13. I can explain how
losing or gaining electrons
affects the radius of an ion.
than the original atom because ____________________________________.
Nonmetals gain electrons and become negative ions that are ____________
than the original atom because ____________________________________.
_____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
Definition:
Allotrope:
_____15. I can define allotrope
and give examples.
As one reads down a group from top to bottom, electronegativity
______________________ because __________________________________
_____16. I can state the periodic
trend for electronegativity and
explain why it occurs.
______________________________________________________________.
As one reads across a period from left to right, electronegativity
______________________ because __________________________________
______________________________________________________________.
As one reads down a group from top to bottom, ionization energy
______________________ because __________________________________
_____17. I can state the periodic
trend for ionization energy and
explain why it occurs.
______________________________________________________________.
As one reads across a period from left to right, , first ionization energy
______________________ because __________________________________
______________________________________________________________.
As one reads down a group from top to bottom, atomic radius
______________________ because __________________________________
_____18. I can state the periodic
trend for atomic radius and
explain why it occurs.
______________________________________________________________.
As one reads across a period from left to right, atomic radius
______________________ because __________________________________
______________________________________________________________.
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:
_____1. I can define
intramolecular forces and
intermolecular forces and give
examples of each.
Examples:
Intermolecular Forces:
Examples:
BARF stands for
__________________________________________________
This means that when a bond is FORMED, energy is __________________
and when a bond is BROKEN, energy is _________________________.
_____ 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 __________ valence electrons.
The three types of chemical bonds are ___________________________,
___________________________, and______________________________.
Definitions:
Ionic Bond:
_____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:
Metallic Bond:
In an ionic bond, the valence electrons of the ____________________are
_______________________ to the ___________________ so that each atom
attains a stable octet (like noble gases). In these compounds, the
_____6. I can define ionic,
covalent, and metallic bonds
based on what happens to the
valence electrons.
electronegativity difference is _____________________________________.
In a covalent bond, the valence electrons of the two ____________________
are _______________________ so that each atom attains a stable octet. In
these compounds, the electronegativity difference is ___________________.
In a metallic bond, positive ions are immersed in ____________________
______________________________________________________________.
_____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:
_____9. I can state the properties
of ionic substances.
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
_____ 11. I can state the number
of electrons that are shared in
single and multiple covalent
bonds.
B) C12H22O11
C)NaCl
D) C5H12
In a single covalent bond, ________ electrons or ________ pair is shared.
In a double covalent bond, ________ electrons or ________ pairs are shared.
In a triple covalent bond, ________ electrons or ________ pairs are shared.
Draw Lewis dot diagrams for the following molecular substances and
identify their shapes.
_____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.
H2O
CO2
I2
CH4
NH3
_____13. I can state the
properties of molecular
substances.
_____14. I can identify a
substance as “molecular” based
on its properties.
Physical properties of molecular substances are:
CHCl3
_____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
______________________________________________________ bonding.
State the type(s) of bonding in the following compounds:
NaCl __________________________ CO___________________________
Hg ________________________ Na3PO4 _____________________________
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.
Explain, in terms of valence electrons, why the bonding in HCl is different
than the bonding in NaCl.
Polar covalent bonds are formed when _____________________________
nonmetals share electrons ____________________________. The
_____18. I can explain the
difference between a polar
covalent bond and a nonpolar
covalent bond.
electronegativity difference is _____________________________________.
Nonpolar covalent bonds are formed when ___________________________
nonmetals share electrons ____________________________. The
electronegativity difference is ____________________________.
_____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.
_____21. I can determine which
end of a polar covalent bond is
slightly negative and which end
is slightly positive and explain
why.
The degree of polarity of a covalent bond is determined by the
_____________________________________________ 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.
Draw a water molecule and include slight charges. Explain why you
assigned each atom its charge.
Definitions:
Symmetrical:
_____22. I can define
symmetrical and asymmetrical.
Asymmetrical:
SNAP means____________________________________________________.
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.
Determine which molecules are polar and which are nonpolar.
Justify your answer.
_____24. I can determine if a
molecular is polar or nonpolar.
H2O
CO2
I2
CH4
NH3
CHCl3
Definitions:
Dipole-Dipole:
Dispersion:
_____25. I can define the
different types of intermolecular
forces.
Hydrogen Bonding:
Molecule-Ion Attraction:
_____26. I can state the
relationship between polarity
and intermolecular force
strength.
As the polarity of the molecule __________________________, the strength
_____27. I can state the
relationship between size of the
molecule and intermolecular
force strength.
As the size of the molecule___________________________, the strength
_____28. I can state the
relationship between the strength
of the intermolecular forces and
vapor pressure.
As the strength of the forces ____________________________, vapor
_____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.
of the forces ______________________________.
of the forces ______________________________.
pressure __________________________.
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.
At STP, CF4 boils at -127.8˚C and NH3 boils at -33.3˚C. Which substance
has stronger intermolecular forces? Justify your answer.
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__________________________________________
_____1. Given the chemical
formula, I can write the name for
binary compounds.
MgI2_________________________________________
Ni2O3_________________________________________
FeCl2_________________________________________
SnS___________________________________________
Write the chemical formula for the following compounds:
sodium bromide__________________
_____2. Given the name, I can
write the chemical formula for
binary compounds.
lithium iodide___________________
iron (III) fluoride________________
vanadium (V) oxide__________________
Write the name for the following compounds:
_____3. Given the chemical
formula, I can write the name for
ternary compounds.
CuSO4 __________________________________________
(NH4)3PO4________________________________________
Fe3(PO4)2 _________________________________________
_____4. Given the name, I can
write the chemical formula for
ternary compounds.
Write the chemical formula for the following compounds:
calcium oxalate_________________________________
nickel (II) thiosulfate_____________________________
Write the name for the following compounds:
_____5. Given the chemical
formula, I can write the name for
covalent molecules.
N2O __________________________________________
SF6___________________________________________
N2O3 _________________________________________
Write the chemical formula for the following compounds:
_____6. Given the name, I can
write the chemical formula for
covalent molecules.
diphosphorus pentasulfide_____________________
carbon monoxide_____________________________
dinitrogen tetroxide ___________________________
How many atoms are in N2?
_____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?
How many atoms of carbon are in Pb(C2H3O2)2?
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 + _____Ca(OH)2  _____Al(OH)3 + _____CaSO4
2 NH4Cl + CaO  2 NH3 + ______________ + CaCl2
Classify the following reactions as synthesis, decomposition, single
replacement, or double replacement.
Which of the following elements is most likely to react?
A) Cu
B) Al
C) Li
D) Mg
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 +
AgNO3 
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).
_____2. I can define a mole as it
pertains to chemistry.
The gram formula mass of a substance is the mass of ___________________.
Definition:
Mole:
What is the gfm for Ni2O3?
What is the gfm for Pb(C2H3O2)2?
_____3. I can determine the
gram-formula mass for any
compound or hydrate.
What is the gfm for CuSO4 • 5 H2O?
What is the percent by mass of Mg in Mg(NO3)2?
_____4. I can determine the
percent composition of an
element in a compound.
Definition:
Hydrate:
_____5. I can define hydrate.
What is the percent by mass of H2O in Na2SO4 • 10 H2O?
_____6. I can determine the
percent composition of water in
a hydrate.
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?
Definitions:
Empirical Formula:
_____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.
_____11. I can find the number of
moles of a substance if I am
given the mass and formula for
the substance.
Molecular Formula:
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?
What is the molecular formula of a compound that has the empirical
formula of CH2O and a molar mass of 180 g/mol?
How many moles of NaCl are equivalent to 94.3 g?
How many grams of LiBr would 3.5 moles of LiBr be?
_____12. I can find the mass of a
substance if I am given the moles
and formula for the substance.
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.
How many molecules are in 0.25 moles of H2?
How many liters does 4.6 moles of N2 occupy?
_____14. I can convert between
moles and volume.
_____15. I can answer questions
using Avogadro’s Hypothesis.
_____16. I can define
stoichiometry.
How many moles of CO2 are present in an 11.2 L sample?
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:
Given the following balanced equation, state the mole ratios between the
following substances.
_____17. Given a balanced
equation, I can state the mole
ratios between any of the
reactants and/or products.
C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O ()
The mole ratio between C3H8 and O2 is _______ to ________.
The mole ratio between C3H8 and CO2 is _______ to ________.
The mole ratio between C3H8 and H2O is _______ to ________.
The mole ratio between CO2 and O2 is _______ to ________.
The mole ratio between H2O and CO2 is _______ to ________.
Using the equation from question #17, determine how many moles of O2
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.
Using the equation from question #17, determine how many moles of CO2
are produced when 8.0 moles of H2O completely react.
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:
Solvent:
_____1. I can define solute,
solvent, solution, and solubility.
Solution:
Solubility:
“Like dissolves like” means
_____2. I can explain and apply
the expression “like dissolves
like.”
_____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.
Explain, in terms of molecular polarity, why ammonia is more soluble than
methane in water at 20˚C at standard pressure.
As the temperature increases, the solubility of solids and liquids
______________________________.
As the temperature increases, the solubility of a gas ___________________.
As the pressure increases, the solubility of a gas ______________________.
Gases are most soluble in ____________________ temperatures and
_____________________ 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.
_____potassium chlorate
_____silver bromide
_____lithium carbonate
_____calcium carbonate
How many grams of KClO3 must be dissolved in 100 grams of water at
20˚C to make a saturated solution?
_____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?
Definitions:
Concentration:
_____10. I can define
concentration, dilute, and
concentrated.
Dilute:
Concentrated:
What is the molarity of 3.5 moles of NaBr dissolved in 500 mL of water?
_____11. I can calculate the
concentration of a solution in
molarity.
What is the molarity of 1.5 L of a solution containing 52 g of LiF?
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.
_____14. I can explain how
adding a solute to a pure solvent
affects 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.
____16. I can determine which
solution will have the lowest
freezing point/highest boiling
point.
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?
When a solute is added to a solvent, the freezing point of the solvent
_________________________. The more solute that is added, the
____________________ the freezing point gets.
When a solute is added to a solvent, the boiling point of the solvent
__________________________. The more solute that is added, the
____________________ the boiling point gets.
Which solution has the highest boiling point?
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.
Arrhenius Theory
_____1. I can use two different
systems to define acids and
bases.
Alternative Theory
(Bronsted-Lowry)
Acid
Base
_____2. I can determine which
substance is an acid or base
according to the alternative
theory given a reaction.
Given the reaction: HI + H2O  H3O+ + I-, identify the acid and base.
Acid: __________________________ Base: __________________________
List the chemical names of three common acids and three common bases.
Acids
Bases
_____3. I can give examples of
the chemical names of common
acids and bases.
List the chemical formulas of three common acids and three common bases.
Acids
Bases
Acids
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.
Definitions:
Electrolyte:
Nonelectrolyte:
_____6. I can define electrolyte,
nonelectrolyte, hydronium ion,
and hydroxide ion.
Hydronium Ion:
Hydroxide Ion:
_____7. I can state another name
for the hydronium ion.
The hydronium ion is also known as the _____________________________.
_____8. I can state the two types
of substances that are
nonelectrolytes.
__________________________ and ______________________ are two
_____9. I can state the three
types of substances that are
electrolytes.
__________________, _________________, and _________________ are
_____10. I can identify
nonelectrolytes.
_____11. I can identify
electrolytes.
classes of compounds that are nonelectrolytes.
three classes of compounds that are electrolytes.
Circle all of the substances that are nonelectrolytes.
(1) C2H5OH
(2) C6H12O6
(3) C12H22O11
(4) CH3COOH
Which substance, when dissolved in water, forms a solution that conducts
an electric current?
(1) C2H5OH
(2) C6H12O6
(3) C12H22O11
(4) CH3COOH
Predict the products in the following reactions, then balance.
Zn + HCl 
_____12. I can complete
reactions between metals and
acids.
Mg + H3PO4 
Cu + H2SO4 
_____13. I can define
neutralization.
Definition:
Neutralization:
Which of the following equations is a neutralization reaction?
_____14. I can identify a
neutralization reaction from a
list of reactions.
(5) LiOH
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:
_____15. I can define pH, pOH,
and [ ].
pOH:
[
_____16. I can describe the
relationship between pH and
pOH and [H+] and [OH-].
]:
pH + pOH =
[H+] and [OH-] are ___________________________ related.
If the pH of a solution is 4.5, the solution is ____________________.
_____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 ____________________.
If the pH of a solution is 11, the solution is ____________________.
If the pH of a solution is 5.7, the solution is ____________________.
_____18. I can identify the pH
values of strong acids and strong
bases.
Strong acids have ___________________ pH values and strong bases have
_________________________ pH values.
In an acidic solution the [H+] is ____________________________ the [OH-].
_____19. I can describe acidic,
basic, or neutral solutions in
terms of ion concentrations.
In a basic solution the [H+] is ______________________________ the [OH-].
In a neutral solution the [H+] is ____________________________ the [OH-].
As the H+ concentration decreases, the pH __________________________.
_____20. I can state the
relationship between H+
concentration and pH.
As the H+ concentration increases, the pH __________________________.
A greater number of hydrogen ions results in a ___________________ pH,
indicating a ______________________ acid.
If the [H3O+] is 1 x 10-8, the pH of the solution will be________.
_____21. Given the hydronium
ion concentration, I can
determine the pH.
If the [H3O+] is 1 x 10-1, the pH of the solution will be________.
If the [H3O+] is 1 x 10-14, the pH of the solution will be________.
If the [H3O+] is 1 x 10-7, the pH of the solution will be________.
If the pH is 10, the [H+] is _________________________.
_____22. Given the pH, I can
determine the [H+] or [OH-].
If the pH is 3, the [H+] is _________________________.
If the pH is 6, the [OH-] is _________________________.
If the pH is 9, the [OH-] is _________________________.
If the H+ concentration is increased by a factor of 10,
the pH will _______________________ by _________.
_____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 _______________________ by _________.
If the H+ concentration is decreased by a factor of 1,000,
the pH will ________________________ by _________.
_____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?
What color is thymol blue in a pH of 11?
_____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?
_____26. I can state the purpose
of titration.
Why do scientists do titrations?
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?
_____27. I can solve for any
variable in the titration equation.
A student uses a 1.4 M HBr solution in a titration to determine the
concentration of a KOH solution. The data is below:
What is the molarity of the KOH solution?
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:
Ionic substances react ________________________ than covalent molecules
because ________________________________________________________.
As the concentration __________________________, the reaction rate
_____3. I can state and apply the
relationship between
concentration and reaction rate
in terms of the collision theory.
______________________________ because there are _________ 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 ____________________________, the reaction rate
_____________________________ because there are ___________ effective
_____4. I can state and apply the
relationship between surface area collisions between particles.
and reaction rate in terms of the
collision theory.
At STP, which 4.0 g sample of Zn(s) will react most quickly with dilute
hydrochloric acid?
A) lump
B) bar
C) powdered
D) sheet metal
As the temperature _____________________________, the reaction rate
_____5. I can state and apply the
relationship between
temperature and reaction rate in
terms of the collision theory.
____________________________ because there are ___________ 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.
Given the following balanced equation: 2 C + 3 H2  C2H6, what is the
heat of reaction if 3 moles of carbon react?
Definitions:
Potential Energy Diagram:
Heat of Reaction (∆H):
_____9. I can define potential
energy diagram, heat of reaction
(∆H), activation energy, and
catalyst.
Activation Energy:
Catalyst:
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 = _____________
PE of products = _____________
PE of activated complex = _____________
∆H= _____________
Activation Energy of Forward Reaction = _____________
Activation Energy of Reverse Reaction = _____________
_____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:
Will Not Change:
Give the potential energy diagram below, determine if the reaction is
exothermic or endothermic. Justify your answer.
_____12. Given a potential
energy diagram, I can determine
if the reaction is exothermic or
endothermic.
In a system at equilibrium the ________________ of the forward and reverse
_____13. I can state two
conditions that apply to all
systems at equilibrium.
reaction must be __________________ and the ________________________
of the reactants and products must be ___________________________.
Two types of physical equilibrium are ________________________, in
_____14. I can describe examples
of physical equilibrium.
which __________________________________ occur at the same rate, and
_________________________ equilibrium. In terms of saturation, a
solution that is at equilibrium must be ____________________________.
Definitions:
Forward Reaction:
_____15. I can define forward
reaction, reverse reaction,
reversible reaction, and closed
system.
Reverse Reaction:
Reversible Reaction:
Closed System:
_____16. I can state
LeChatelier’s Principle.
LeChatelier’s Principle states
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.
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.
Change
Direction of
Shift
Increased
Concentration
of
Decreased
Concentration
of
Increase [SO2]
Increase [SO3]
Decrease [O2]
Increase
temperature
Decrease
temperature
Increase pressure
Decrease pressure
Add a catalyst
_____18. I can define entropy.
_____19. I can rank the three
phases of matter from least
entropy to most entropy.
Entropy is
Least entropy
Most entropy
______________________<______________________<__________________
_____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)
_______________________ in entropy and a(n) _____________________
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 __________________ Li _________________ Na+ __________________
Assign oxidation numbers to each element in the compounds below.
_____2. I can assign oxidation
numbers to the elements in a
compound.
MnCl3: Mn ____________________ Cl _______________________
H2SO4: H __________________ S ________________ O ________________
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 _____________________
ClO3-1: Cl _____________________ O _________________________
The Law of Conservation of Charge states
Definitions:
Oxidation:
Reduction:
_____5. I can define oxidation,
reduction, redox reaction,
oxidizing agent, and reducing
agent.
Redox Reaction:
Oxidizing Agent:
Reducing Agent:
_____6. I can identify a redox
reaction from a list of chemical
reactions.
O ________________________
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:
Reduction Half-Reaction:
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.
_____11. I can state the two
types of electrochemical cells.
A spontaneous redox reaction will occur if the metal being oxidized is
__________________________ on Table J, or the nonmetal reduced is
__________________________ on Table J.
The two types of electrochemical cells are:
_____________________________ and _____________________________
Voltaic
Components
Oxidation Site
Reduction Site
Anode Charge
Cathode Charge
_____12. I can compare the two
types of electrochemical cells.
Mass at Anode
Mass at Cathode
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
Electrolytic
_____14. I can answer questions
about a voltaic cell.
What is the anode?
What is the cathode?
Write the oxidation half-reaction:
Write the reduction half-reaction:
In what direction do electrons flow?
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.
Explain, in terms of atoms and ions, why the mass of the anode decreases
during the operation of an electrochemical cell.
The diagram below represents an electrochemical cell:
_____16. I can answer questions
about an electrolytic cell.
What is the anode?
What is the cathode?
Write the oxidation half-reaction:
Write the reduction half-reaction:
In what direction do electrons flow?
What is the purpose of the battery?
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
______________________. Its symbol is_______________.
Definitions:
Organic Compound:
Saturated Hydrocarbon:
_____2. I can define organic
compound, saturated
hydrocarbon, unsaturated
hydrocarbon, isomers, and alkyl
group.
Unsaturated Hydrocarbon:
Isomers:
Alkyl Group:
_____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: _______________________ 2-Decene: _______________________
C4H8: _____________________
C7H12: _______________________
_____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: _______________________ 2-Decene: _______________________
CH3CH2CH2CH3: _______________ CH3CHCHCH3: __________________
Determine the homologous series of hydrocarbons to which each of the
following belongs:
belongs to the _______________________ series.
_____8. Given the structural
formula, I can determine which
homologous series a hydrocarbon
belongs to.
belongs to the _______________________ series.
belongs to the ________________________ 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: ________________________ Ethene: _________________________
3-Nonene: ______________________ 1-Pentyne: _______________________
Propene: ______________________ Methane: ________________________
_____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:
_____12. I can use Tables P & Q
to name hydrocarbons.
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.
Draw the following organic compounds:
_____16. I can use Tables P & R
to draw simple compounds in
any of the classes of organic
compounds.
2,2-dichloropentane
2-butanol
dimethyl ether
ethanal
2-pentanone
propanoic acid
methyl butanoate
ethanamine
pentanamide
ethyl propanoate
Substitution (halogenation) is the replacement of a __________________ in
a ____________________________ hydrocarbon with a halogen. It starts
with an _____________________ and ends with ____________ product(s).
Addition is the addition of _____________________________________ or
_____17. I can describe the 7
types of organic reactions.
_____________________________ at a double bond in an
________________________________ hydrocarbon. It starts with an
__________________________ and ends with _______________ product(s).
Esterification is when a(n) ______________________________ and a(n)
________________________ react to form ______________________ and
___________________________.
Saponification is the production of ___________________.
Fermentation is the breakdown of ___________________________ into
___________________________ and ___________________________.
Combustion is when a hydrocarbon reacts with ____________________ to
produce _________________________ and __________________________.
Polymerization is the formation of a large molecule called a _____________
from small molecules called ______________________. When monomers
are joined together by breaking double or triple bonds it is called
________________________________ polymerization. When monomers
are joined by the removal of water it is called _________________________
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.