Scientific Measurement curve, I can determine which sections of the curve show changes in kinetic energy. On the graph, circle the sections that show a change in kinetic energy. _____27

The volume is_______________ mL.

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.

_____2. I can convert between

different metric units.

How many joules are in 9.3 kJ?

How many grams are in 39,983 mg?

How many milliliters are in 452 L?

_____3. I can solve for any

variable of the density equation.

What is the volume of 54.6 g of beryllium (density = 1.85 g/cm3)?

What is the density of a 27.8 g substance with a volume of 2.3 mL?

_____4. I can determine the

number of significant figures in

a measurement.

How many significant figures are there in 30.50 cm?

How many significant figures are there in 400 sec?


answer to a math problem to the

correct number of significant

figures.

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.

_____7. I can convert numbers

into scientific notation from

standard notation.

Convert 87,394,000,000,000 to scientific notation.

Convert 0.0000040934 to scientific notation.

_____8. I can convert numbers

into standard notation from

scientific notation.

Convert 5.8 x 109 to standard notation.

Convert 4.3 x 10-5 to standard notation.

Unit 1: Matter & Energy


_____1. I can define the terms

atom, element, compound, and

mixture.

Definitions: Atom:

Element:

Compound:

Mixture:

_____2. I can classify substances

as a pure substance (element or

compound) or a mixture.

Put each of the following examples into the correct column.

Examples: C12H22O11, NaCl, Fe, salt water, air, CO2, H2, Ar, soda

Element

Compound Mixture

_____3. I can list the 7 diatomic

elements.

The seven diatomic elements are:

_____4. I can identify a binary

compound.

Which of the following is a binary compound?

A) NaOH B) H3PO4 C) H2O D) MgSO4

_____5. I can define

homogeneous mixture and

heterogeneous mixture.

Definitions: Homogeneous Mixture:

Heterogeneous Mixture:

_____6. I can give examples of

homogeneous and

heterogeneous mixtures.

Two examples of Homogeneous Mixtures:

A)

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?

_____9. I can state which

separation process is best for a

given situation.

To separate a mixture of salt and water, use __________________________.

To separate a mixture of ethanol and water, use _______________________.

To separate a mixture of sand and water, use__________________________.

____10. I can draw particle

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.

Monoatomic Element

Diatomic Element

Compound

Mixture of 2 compounds

Mixture of an element and compound

_____11. I can answer questions

about the Law of Conservation

of Matter (Mass).

All chemical reactions have conservation of

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.

_____12. I can classify a

property as physical or chemical.

Write “P” for physical or “C” for chemical on the line provided.

_____ Copper (II) sulfate is blue.

_____ Copper reacts with oxygen. _____ Copper can be made into wire.

_____ Copper has a density of 8.96 g/cm3.

_____ Copper melts at 1358 K.

_____ Copper doesn’t dissolve in water.

_____13. I can classify a change

as physical or chemical.

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.

_____ 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.

_____14. In a particle diagram, I

can distinguish between a

physical change and a chemical

change.

Circle the particle diagram that best represents Substance A after a physical change has occurred.

Substance A

____15. I can use particle

diagrams to show the

arrangement and spacing of

atoms/molecules in different

phases.

Draw a particle diagram to represent atoms of Li in each phase.

Solid Liquid Gas

_____16. I can compare solids,

liquids, and gases in terms of

their shape, volume, particle

distance, and attraction between

particles.

Solid Liquid Gas

Shape

Volume

Particle

Distance

Attraction

Between

Particles

_____17. I can state the change

of phase occurring during each

phase change.

Fusion: change from _________________ to _________________.

Solidification: change from _________________ to _________________.

Condensation: change from _________________ to _________________.

Vaporization: change from _________________ to _________________.

Melting: change from _________________ to _________________.

Boiling: change from _________________ to _________________.

Sublimation: change from _________________ to _________________.

Deposition: change from _________________ to _________________.

Freezing: change from _________________ to _________________.


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

_____19. I can define kinetic

energy, potential energy,

temperature, heat, endothermic,

and exothermic.

Definitions: Kinetic Energy:

Potential Energy:

Temperature:

Heat:

Endothermic:

Exothermic:

_____20. I can indicate if a

phase change is exothermic or

endothermic.

Indicate whether the change is exothermic or endothermic:

fusion/melting _______________________

solidification/freezing _______________________

condensation _______________________

vaporization/boiling _______________________ sublimation _______________________

deposition _______________________


about average kinetic energy.

Which sample of water contains particles having the highest 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

_____22. I can state the

direction of heat flow.

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?

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


Celsius and Kelvin.

What Kelvin temperature is equal to 200˚C?

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 = _____________________________________________________

_____25. Given a


determine which sections of the

curve show changes in potential

energy.

On the graph, circle the sections that show a change in potential energy.

_____26. Given a



curve show changes in kinetic

energy.

On the graph, circle the sections that show a change in kinetic energy.

_____27. Given a



curve show phase changes.

On the graph, circle the sections that show a phase change.

_____28. Given a


determine the temperature at

which a substance melts/freezes

or vaporizes/condenses.

What is the freezing point of this substance?

What is the boiling point of this substance?


temperature at which water

freezes in ˚C and K.

What is the freezing point of water in ˚C and K?



melts in ˚C and K.

What is the melting point of water in ˚C and K?



boils in ˚C and K.

What is the boiling point of water in ˚C and K?



condenses in ˚C and K.

What is the condensing point of water in ˚C and K?

_____33. I can state the unit

used to measure energy.

Energy is measured in ____________________.

53oC

113oC

_____34. I can define specific

heat capacity, heat of fusion,

and heat of vaporization.

Definitions: Specific Heat Capacity:

Heat of Fusion:

Heat of Vaporization:

_____35. I can use the heat

equations to solve for any

variable.

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?

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?

_____36. I can state the 5 parts

of the Kinetic Molecular

Theory.

The five parts of the Kinetic Molecular Theory are:

A.

B.

C.

D.

E.

_____37. I can define an ideal

gas.

Definition: Ideal Gas:

V

P

T

V

T

P

_____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

________________________ and some _________________________.

_____39. I can state the two

elements that act ideally most of

the time.

The two elements that act ideally most of the time are _____________ &

_______________.


conditions under which real

gases behave most like ideal

gases.

A real gas will act most ideally under the conditions of __________ pressure

and ____________ temperature. (PLIGHT)


meaning of “STP” and the

Table on which it can be found.

STP stands for ___________________________________________________.

The values can be found on Table______________.

Standard Pressure = _______________________________________.

Standard Temperature = ____________________________________.

_____42. I can state Avogadro’s

Hypothesis (Law) and answer

questions based on this concept.

Avogadro’s Hypothesis says_______________________________________

________________________________________________________________

_______________________________________________________________.

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)


relationship between pressure

and volume for gases.

At constant temperature, as the pressure on a gas increases, the volume

____________________.


relationship between

temperature and volume for

gases.

At constant pressure, as the temperature of a gas increases, the volume

____________________.



temperature and pressure for

gases.

In a fixed container (has constant volume), as the temperature of a gas

increases, the pressure____________________.

_____46. I can recognize the

relationships stated above in

graph form.

Draw a curve for each of the graphs below:

_____47. I can solve problems

using the combined gas law.

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?

_____48. I can define vapor

pressure, boiling point, and

volatile.

Definitions: Vapor Pressure:

Boiling Point:

Volatile:


conditions of temperature and

pressure that are used for normal

boiling points.

The normal boiling point of a substance occurs at a pressure of

___________ atm or _______________kPa. This is known as

____________________________________and can be found on Table _____.



temperature and vapor pressure.

As the temperature of a substance increases, its vapor pressure

____________________________.



atmospheric pressure and

boiling point.

As the atmospheric pressure increases, the boiling point ____________________________.


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?


vapor pressure of ethanol,

ethanoic acid, propanone, or

water at a given temperature.

What is the vapor pressure of ethanol at 56˚C?


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?

Unit 2: The Atom


_____1. I can describe John

Dalton’s contribution to our

understanding of the atom.

Dalton’s Model:

Diagram:

_____2. I can describe JJ

Thomson’s contribution to our


Thomson’s Model:

Diagram:

_____3. I can describe Ernest

Rutherford’s contribution to our


Rutherford’s Experiment:

Experiment Results:

Experiment Conclusions:

Diagram:

_____4. I can describe Niels

Bohr’s contribution to our


Bohr’s Model:

Diagram:

_____5. I can describe the

modern model (wave mechanical

model) of the atom.

Where are electrons likely to be found according to the modern model?

Diagram:

_____6. I can state the three

subatomic particles, their

location in an atom, and the

magnitude of their charges and

their masses (in amu).

Particle #1 Particle #2 Particle #3

Name

Charge

Mass

Location in

Atom

_____7. I can identify the nuclear

charge of an atom.

What is the charge of an oxygen-18 nucleus?

_____8. I can explain why atoms

are electrically neutral.

Atoms are electrically neutral because the number of ________________ is

equal to the number of _____________________.

_____9. I can define atomic

number and mass number.

Definitions: Atomic 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?

How many protons are in an atom of silicon?


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.

What are three other notations that can be used to represent 80Br?

_____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?

In an atom of 212Po, how many electrons are present?

In an atom of 212Po, how many neutrons are present?

_____14. I can define ion,

cation, and anion.

Definitions: Ion:

Cation:

Anion:

_____15. Given the mass number

and the charge, I can determine

the number of protons, neutrons,

and electrons in an ion.

How many protons are in 18O-2?

How many neutrons are in 18O-2?

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.

Definition: Isotope:

_____17. I can identify isotopes

based on their symbols.

Which symbols represent atoms that are isotopes?

A) 14C and 14N B) 16O and 18O C) 131I and 131I D) 222Rn and 222Ra

_____18. I can calculate average

atomic mass given the masses of

the naturally occurring isotopes

and the percent abundances.

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?

_____19. I can define principal

energy level, orbital, ground

state, excited state, electron

configuration, and bright line

spectrum.

Definitions: Principal Energy Level:

Orbital:

Ground State:

Excited State:

Electron Configuration:

Bright Line Spectrum:


maximum number of electrons

that will fit into each of the first

four shells.

The 1st shell holds a maximum of ___________ electrons.

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.

Definition: Valence Electrons:

_____22. I can define a stable

octet.

Definition: Stable Octet:

_____23. I can locate and

interpret an element’s electron

configuration on the Periodic

Table.

How many valence electrons does rubidium have in the ground state?

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?


relationship between distance

from the nucleus and energy of

an electron.

As the distance between the nucleus and the electron increases, the energy

of the electron ____________________________.


relationship between the number

of the principal energy level and

the distance to the atom’s

nucleus.

As the number of the principal energy level increases, the distance to the

nucleus ____________________________.

_____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?

A) 2-8-7-1 B) 2-8-8-1 C) 2-8-7-2 D) 2-8-8-2

_____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

_____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


_____1. I can explain what

makes a nucleus stable or

unstable.

The stability of the nucleus is dependent on the ______________ to

_________________ ratio.

_____2. I can compare types of

radiation.

Type Symbol Notation Charge Mass Penetrating

Power

alpha

beta

positron

gamma

_____3. I can explain the

difference between natural

transmutation and artificial

transmutation.

The difference between natural transmutation and artificial transmutation is

that in natural transmutation an_____________ __________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

for spontaneous decay.

Two synonyms for spontaneous decay are:___________________________

and _________________________________.

_____5. I can identify a natural

decay reaction from a list of

reactions.

Which equation represents a natural decay?

_____6. I can identify an

artificial transmutation reaction

from a list of reactions.

Which equation represents artificial transmutation?

_____7. I can show how mass

number and electrical charge are

conserved in any nuclear

reaction.

Complete the following nuclear equations:

________

_____8. I can define half-life.

Definition: Half-life:

_____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.

Based on Reference Table N, what fraction of a radioactive sample of Au-198 will remain unchanged after 10.78 days?

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)

_____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?

_____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?

_____12. Using Table N, I can

determine the length of half-life

and/or decay mode for a specific

radioactive isotope.

Compared to K-37, the isotope K-42 has

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

_____14. I can define fission and

fusion.

Definitions: Fission:

Fusion:

_____15. I can identify a fission

reaction from a list of reactions.

Which equation represents fission?

_____16. I can identify a fusion

reaction from a list of reactions.

Which equation represents fusion?


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.

_____18. Given a list of

reactions, I can differentiate a

nuclear reaction from a chemical

reaction.

Which of the following equations represent NUCLEAR reactions?

_____19. I can state 5 beneficial

uses for radioactive isotopes.

Five beneficial uses for radioactive isotopes are:

A.

B.

C.

D.

E.

_____20. I can state the scientific

use of specific radioactive

isotopes.

Tc-99 and Co-60 are used for_______________________________________

I-131 is used for_________________________________________________

C-14 is used for _________________________________________________

U-238 is used for ________________________________________________

_____21. I can state three risks

associated with radioactivity and

radioactive isotopes.

Three risks associated with radioactivity and radioactive isotopes are:

A.

B.

C.

Unit 4: Periodic Table


_____1. I can state how the

elements on the Periodic Table

are arranged.

The elements on the Periodic Table are arranged by increasing

___________________ ____________________.

_____2. I can state the group

names for elements in groups 1,

2, 3-12, 17, and 18.

Group 1 is called the _________________________________________.

Group 2 is called the _________________________________________.

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


electronegativity, ionization

energy, atomic radius, and ionic

radius.

Definitions: Electronegativity:

Ionization Energy:

Atomic Radius:

Ionic Radius:

_____6. I can classify elements as

metals, nonmetals, or metalloids

based on their placement on the

Periodic Table.

Classify each of the following elements as metals (M), nonmetals (NM), or

metalloids (MTLD) (have properties of both metals and nonmetals).

_______B ________K ________Li ________C _______Ar

_______Sb ________H ________Fe ________Au _______S

_______F ________Si ________Fr ________He _______Rn

_______Ge ________Al ________As ________Bi _______I


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 ___________________________.


names/symbols of the 11

elements that are gases at STP.

The 11 elements that are gases at STP are:

_____________________________, _______________________________,

_____________________________, _______________________________,

_____________________________, _______________________________,

_____________________________, _______________________________,

_____________________________, _______________________________,

and _____________________________.

_____9. I can list properties of

metals.

The properties of metals are:


nonmetals.

The properties of nonmetals are:

_____11. I can state the trend for

activity/reactivity for METALS

as one reads down a group.

As one reads down a group from top to bottom, the activity/reactivity of

METALS ____________________.

The most reactive metal is _________________________________.

_____12. I can state the trend for

activity/reactivity for

NONMETALS as one reads

down a group.

As one reads down a group from top to bottom, the activity/reactivity of

NONMETALS __________________.

The most reactive nonmetal is ______________________________.

_____13. I can explain how

losing or gaining electrons

affects the radius of an ion.

Metals lose electrons and become positive ions that are ________________

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

_____15. I can define allotrope

and give examples.

Definition: Allotrope:

_____16. I can state the periodic

trend for electronegativity and

explain why it occurs.

As one reads down a group from top to bottom, electronegativity

______________________ because __________________________________

______________________________________________________________.

As one reads across a period from left to right, electronegativity

______________________ because __________________________________

______________________________________________________________.


trend for ionization energy and


As one reads down a group from top to bottom, ionization energy

______________________ because __________________________________

______________________________________________________________.

As one reads across a period from left to right, , first ionization energy

______________________ because __________________________________

______________________________________________________________.


trend for atomic radius and


As one reads down a group from top to bottom, atomic radius

______________________ because __________________________________

______________________________________________________________.

As one reads across a period from left to right, atomic radius

______________________ because __________________________________

______________________________________________________________.

Unit 5: Bonding



intramolecular forces and

intermolecular forces and give

examples of each.

Definitions: Intramolecular Forces:

Examples:

Intermolecular Forces:

Examples:

_____ 2. I can explain and apply

the meaning of BARF as it

applies to chemical bonding.

BARF stands for

__________________________________________________

This means that when a bond is FORMED, energy is __________________

and when a bond is BROKEN, energy is _________________________.

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.

Atoms are most stable when they have __________ valence electrons.


types of chemical bonds.

The three types of chemical bonds are ___________________________,

___________________________, and______________________________.

_____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.

Definitions: Ionic Bond:

Covalent Bond:

Metallic Bond:

_____6. I can define ionic,

covalent, and metallic bonds

based on what happens to the

valence electrons.

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

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 ____________________

______________________________________________________________.


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

_____8. I can draw a Lewis dot

diagram to represent an ionic

compound.

Draw Lewis dot diagrams for the following ionic compounds:

NaCl CaCl2

_____9. I can state the properties

of ionic substances.

Physical properties of ionic substances are:

_____10. I can identify a

substance as ionic based on its

properties.

A solid substance was tested in the lab. The results are shown below.

*dissolves in water *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

_____ 11. I can state the number

of electrons that are shared in

single and multiple covalent

bonds.

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.

_____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.

Draw Lewis dot diagrams for the following molecular substances and

identify their shapes.

H2O CO2

I2 CH4

NH3 CHCl3


properties of molecular

substances.

Physical properties of molecular substances are:


substance as “molecular” based

on its properties.

_____15. I can state the type of

bonding that occurs in

compounds containing

polyatomic ions.

Compounds containing polyatomic ions have

______________________________________________________ bonding.

_____16. Given the chemical

formula for a compound, I can

determine the type(s) of bonding

in the compound.

State the type(s) of bonding in the following compounds:

NaCl __________________________ CO___________________________

Hg ________________________ Na3PO4 _____________________________

_____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 methane (CH4) is similar to the bonding in water (H2O).

Explain, in terms of valence electrons, why the bonding in HCl is different than the bonding in NaCl.

_____18. I can explain the

difference between a polar

covalent bond and a nonpolar

covalent bond.

Polar covalent bonds are formed when _____________________________

nonmetals share electrons ____________________________. The

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.

The degree of polarity of a covalent bond is determined by the

_____________________________________________ between the elements.

_____20. I can explain why one

covalent bond is more or less

polar than another covalent

bond.

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.

_____21. I can determine which

end of a polar covalent bond is

slightly negative and which end

is slightly positive and explain

why.

Draw a water molecule and include slight charges. Explain why you assigned each atom its charge.


symmetrical and asymmetrical.

Definitions: Symmetrical:

Asymmetrical:

_____23. I can explain and apply

the meaning of SNAP as it

relates to determining molecule

polarity.

SNAP means____________________________________________________.

Why is a molecule of CH4 nonpolar even though the bonds between the

carbon and hydrogen are polar?

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.

_____24. I can determine if a

molecular is polar or nonpolar.

Determine which molecules are polar and which are nonpolar.

Justify your answer.

H2O CO2

I2 CH4

NH3 CHCl3

_____25. I can define the

different types of intermolecular

forces.

Definitions: Dipole-Dipole:

Dispersion:

Hydrogen Bonding:

Molecule-Ion Attraction:


relationship between polarity

and intermolecular force

strength.

As the polarity of the molecule __________________________, the strength

of the forces ______________________________.


relationship between size of the

molecule and intermolecular

force strength.

As the size of the molecule___________________________, the strength

of the forces ______________________________.


relationship between the strength

of the intermolecular forces and

vapor pressure.

As the strength of the forces ____________________________, vapor

pressure __________________________.

_____29. Given the physical state

of some substances, I can

compare the relative strength of

the intermolecular forces.

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.

_____30. Given the boiling

points (or freezing points) of

some substances, I can compare

the relative strength of the

intermolecular forces.

At STP, CF4 boils at -127.8˚C and NH3 boils at -33.3˚C. Which substance

has stronger intermolecular forces? Justify your answer.


about hydrogen bonding.

Which compound has hydrogen bonding between its molecules?

A) CH4 B) CaH2 C) KNO3 D) H2O

The relatively high boiling point of water is due to water having

A) hydrogen bonding B) metallic bonding

C) nonpolar covalent bonding D) strong ionic bonding


about molecule-ion attraction.

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?

Unit 6: Chemical Reactions



formula, I can write the name for

binary compounds.

Write the name for the following compounds:

CrO__________________________________________

MgI2_________________________________________

Ni2O3_________________________________________

FeCl2_________________________________________

SnS___________________________________________

_____2. Given the name, I can

write the chemical formula for

binary compounds.

Write the chemical formula for the following compounds:

sodium bromide__________________

lithium iodide___________________

iron (III) fluoride________________

vanadium (V) oxide__________________



ternary compounds.


CuSO4 __________________________________________

(NH4)3PO4________________________________________

Fe3(PO4)2 _________________________________________



ternary compounds.


calcium oxalate_________________________________

nickel (II) thiosulfate_____________________________



covalent molecules.


N2O __________________________________________

SF6___________________________________________

N2O3 _________________________________________



covalent molecules.


diphosphorus pentasulfide_____________________

carbon monoxide_____________________________

dinitrogen tetroxide ___________________________

_____7.Given the chemical

symbol/formula, I can determine

how many atoms are present.

How many atoms are in N2?

What is the total number of atoms in Pb(C2H3O2)2?

How many atoms of carbon are in Pb(C2H3O2)2?

_____8. I can use particle

diagrams to show conservation

of mass in a chemical equation.

Using the symbols shown below, complete the equation below to illustrate

conservation of mass.

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.

_____ Al2(SO4)3 + _____Ca(OH)2 _____Al(OH)3 + _____CaSO4

_____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.

2 NH4Cl + CaO 2 NH3 + ______________ + CaCl2

_____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.

Classify the following reactions as synthesis, decomposition, single replacement, or double replacement.

_____12. From a list of given list

of elements, I can determine

which element is most active.

Which of the following elements is most likely to react?

A) Cu B) Al C) Li D) Mg

_____13. I can predict if a single

replacement reaction will occur.

Which reaction occurs spontaneously?

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

_____14. I can predict the

products in a double replacement

reaction.

Complete and balance the following reaction:

MgCl2 + AgNO3

= Al

= Br

Unit 7: Moles


_____1. I can define gram

formula mass (molar mass).

The gram formula mass of a substance is the mass of ___________________.

_____2. I can define a mole as it

pertains to chemistry.

Definition: Mole:


gram-formula mass for any

compound or hydrate.

What is the gfm for Ni2O3?

What is the gfm for Pb(C2H3O2)2?

What is the gfm for CuSO4 • 5 H2O?


percent composition of an

element in a compound.

What is the percent by mass of Mg in Mg(NO3)2?

_____5. I can define hydrate.

Definition: Hydrate:


percent composition of water in

a hydrate.

What is the percent by mass of H2O in Na2SO4 • 10 H2O?

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?

_____7. I can define empirical

formula and molecular formula.

Definitions: Empirical Formula:

Molecular Formula:

_____8. Given the molecular

formula, I can determine the

empirical formula of a

compound.

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

_____9. Given the percent

composition, I can determine the

empirical formula of a

compound.

A compound consists of 25.9% nitrogen and 74.1% oxygen by mass. What

is the empirical formula of the compound?

_____10. Given the empirical

formula and the molar mass, I

can determine the molecular

formula of a compound.

What is the molecular formula of a compound that has the empirical

formula of CH2O and a molar mass of 180 g/mol?

_____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?

_____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?


moles and numbers of particles

using Avogadro’s number.

How many moles of O2 contain 3.01 x 1023 molecules?

How many molecules are in 0.25 moles of H2?


moles and volume.

How many liters does 4.6 moles of N2 occupy?

How many moles of CO2 are present in an 11.2 L sample?


using Avogadro’s Hypothesis.

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)


stoichiometry.

Definition: Stoichiometry:

_____17. Given a balanced

equation, I can state the mole

ratios between any of the

reactants and/or products.

Given the following balanced equation, state the mole ratios between the following substances.

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 ________.

_____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 O2

are needed to completely react with 7.0 moles of C3H8.

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


_____1. I can define solute,

solvent, solution, and solubility.

Definitions: Solute:

Solvent:

Solution:

Solubility:

_____2. I can explain and apply

the expression “like dissolves

like.”

“Like dissolves like” means

Explain, in terms of molecular polarity, why ammonia is more soluble than

methane in water at 20˚C at standard pressure.

_____3. I can describe the trend

in solubility for solids and liquids

as the temperature changes.

As the temperature increases, the solubility of solids and liquids

______________________________.


in solubility for gases as the

temperature changes.

As the temperature increases, the solubility of a gas ___________________.


in solubility for gases as the

pressure changes.

As the pressure increases, the solubility of a gas ______________________.


conditions in which gases are

most soluble.

Gases are most soluble in ____________________ temperatures and

_____________________ pressures.

_____7. I can use Table F to

determine if a substance will be

soluble in water.

Write “S” for soluble and “NS” for not soluble.

_____potassium chlorate _____silver bromide

_____lithium carbonate _____calcium carbonate

_____8. I can use Table G to

determine how much solute to

add at a given temperature to

make a saturated solution.

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?


concentration, dilute, and

concentrated.

Definitions: Concentration:

Dilute:

Concentrated:

_____11. I can calculate the

concentration of a solution in

molarity.

What is the molarity of 3.5 moles of NaBr dissolved in 500 mL of water?

What is the molarity of 1.5 L of a solution containing 52 g of LiF?


solution is the most concentrated

or the most dilute.

Which solution is most concentrated?

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

_____13. I can calculate the

concentration of a solution in

ppm.

What is the concentration, in ppm, of a 2,600 g solution containing 0.015 g of CO2?

_____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

_________________________. The more solute that is added, the

____________________ the freezing point gets.

____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

__________________________. The more solute that is added, the

____________________ the boiling point gets.

____16. I can determine which

solution will have the lowest

freezing point/highest boiling

point.

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


_____1. I can use two different

systems to define acids and

bases.

Arrhenius Theory

Alternative Theory

(Bronsted-Lowry)

Acid

Base


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: __________________________


the chemical names of common

acids and bases.

List the chemical names of three common acids and three common bases.

Acids Bases


the chemical formulas of

common acids and bases.

List the chemical formulas of three common acids and three common bases.

Acids Bases


common acids and bases.

Acids Bases

_____6. I can define electrolyte,

nonelectrolyte, hydronium ion,

and hydroxide ion.

Definitions: Electrolyte:

Nonelectrolyte:

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

classes of compounds that are nonelectrolytes.


types of substances that are

electrolytes.

__________________, _________________, and _________________ are

three classes of compounds that are electrolytes.

_____10. I can identify

nonelectrolytes.

Circle all of the substances that are nonelectrolytes.

(1) C2H5OH (2) C6H12O6 (3) C12H22O11 (4) CH3COOH (5) LiOH

_____11. I can identify

electrolytes.

Which substance, when dissolved in water, forms a solution that conducts an electric current?

(1) C2H5OH (2) C6H12O6 (3) C12H22O11 (4) CH3COOH

_____12. I can complete

reactions between metals and

acids.

Predict the products in the following reactions, then balance.

Zn + HCl

Mg + H3PO4

Cu + H2SO4


neutralization.

Definition: Neutralization:


neutralization reaction from a

list of reactions.

Which of the following equations is a neutralization reaction?

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

_____15. I can define pH, pOH,

and [ ].

Definitions: pH:

pOH:

[ ]:


relationship between pH and

pOH and [H+] and [OH-].

pH + pOH =

[H+] and [OH-] are ___________________________ related.

_____17. Based on pH, I can

determine if a solution is acidic,

basic, or neutral.

If the pH of a solution is 4.5, the solution is ____________________.


If the pH of a solution is 11, 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.

_____19. I can describe acidic,

basic, or neutral solutions in

terms of ion concentrations.

In an acidic solution the [H+] is ____________________________ the [OH-].

In a basic solution the [H+] is ______________________________ the [OH-].

In a neutral solution the [H+] is ____________________________ the [OH-].


relationship between H+

concentration and pH.

As the H+ concentration decreases, the pH __________________________.

As the H+ concentration increases, the pH __________________________. A greater number of hydrogen ions results in a ___________________ pH,

indicating a ______________________ acid.

_____21. Given the hydronium

ion concentration, I can

determine the pH.

If the [H3O

+] is 1 x 10-8, the pH of the solution will be________.

If the [H3O


If the [H3O


If the [H3O


_____22. Given the pH, I can

determine the [H+] or [OH-].

If the pH is 10, the [H+] is _________________________.

If the pH is 3, the [H+] is _________________________.

If the pH is 6, the [OH-] is _________________________.

If the pH is 9, the [OH-] is _________________________.


change in pH when the [H+] of a

solution is changed.

If the H+ concentration is increased by a factor of 10,

the pH will _______________________ by _________.

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?

_____27. I can solve for any

variable in the titration equation.

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?

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


_____1. I can define an effective

collision according to the

collision theory.

Definition: Effective Collision:

_____2. I can determine the type

of compound that has a faster

reaction rate.

Ionic substances react ________________________ than covalent molecules

because ________________________________________________________.

_____3. I can state and apply the


concentration and reaction rate

in terms of the collision theory.

As the concentration __________________________, the reaction rate

______________________________ 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


relationship between surface area

and reaction rate in terms of the

collision theory.

As the surface area ____________________________, the reaction rate

_____________________________ because there are ___________ effective


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



temperature and reaction rate in

terms of the collision theory.

As the temperature _____________________________, the reaction rate

____________________________ because there are ___________ effective


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 B) 15˚C C) 95˚C D) 273 K

_____6. I can use Table I to

determine if a reaction is

exothermic or endothermic.

_____7. Based on the location of

the energy term, I can determine

if the reaction is exothermic or

endothermic.

Given the following balanced equation: I + I I2 + 146.3 kJ

Is this reaction exothermic or endothermic? Justify your answer.

_____8. I can determine the new

heat of reaction if the number of

moles in a reaction changes.

Given the following balanced equation: 2 C + 3 H2 C2H6, what is the

heat of reaction if 3 moles of carbon react?

_____9. I can define potential

energy diagram, heat of reaction

(∆H), activation energy, and

catalyst.

Definitions: Potential Energy Diagram:

Heat of Reaction (∆H):

Activation Energy:

Catalyst:

_____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.

Given the potential energy diagram below, determine the following:

PE of reactants = _____________ PE of products = _____________

PE of activated complex = _____________ ∆H= _____________

Activation Energy of Forward Reaction = _____________

Activation Energy of Reverse Reaction = _____________


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:


energy diagram, I can determine

if the reaction is exothermic or

endothermic.

Give the potential energy diagram below, determine if the reaction is

exothermic or endothermic. Justify your answer.

_____13. I can state two

conditions that apply to all

systems at equilibrium.

In a system at equilibrium the ________________ of the forward and reverse

reaction must be __________________ and the ________________________

of the reactants and products must be ___________________________.

_____14. I can describe examples

of physical equilibrium.

Two types of physical equilibrium are ________________________, in

which __________________________________ occur at the same rate, and

_________________________ equilibrium. In terms of saturation, a

solution that is at equilibrium must be ____________________________.

_____15. I can define forward

reaction, reverse reaction,

reversible reaction, and closed

system.

Definitions: Forward Reaction:

Reverse Reaction:

Reversible Reaction:

Closed System:

_____16. I can state

LeChatelier’s Principle.

LeChatelier’s Principle states


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.

Given the reaction at equilibrium:

2 SO2 (g) + O2 (g) 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.

Entropy is

_____19. I can rank the three

phases of matter from least

entropy to most entropy.

Least entropy Most entropy

______________________<______________________<__________________


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


_____1. I can assign oxidation

numbers to any element or ion.

Assign oxidation number to each of the elements below.

O2 __________________ Li _________________ Na+ __________________


numbers to the elements in a

compound.

Assign oxidation numbers to each element in the compounds below.

MnCl3: Mn ____________________ Cl _______________________

H2SO4: H __________________ S ________________ O ________________


numbers to the elements in a

polyatomic ion.

Assign oxidation numbers to each element in the polyatomic ions below.

PO4

-3: P _____________________ O ________________________

ClO3

-1: Cl _____________________ O _________________________

_____4. I can state the Law of

Conservation of Charge.

The Law of Conservation of Charge states

_____5. I can define oxidation,

reduction, redox reaction,

oxidizing agent, and reducing

agent.

Definitions: Oxidation:

Reduction:

Redox Reaction:

Oxidizing Agent:

Reducing Agent:

_____6. I can identify a redox

reaction from a list of chemical

reactions.

_____7. I can distinguish

between an oxidation half-

reaction and a reduction half-

reaction.

Which half-reaction equation represents the reduction of a potassium ion?

A) K+ + 1 e- K0 C) K+ K0 + 1 e-

B) K0 + 1 e- K+ D) K0 K+ + 1 e-

_____8. I can break a redox

reaction into its two half-

reactions.

The two half-reactions that come from the following equation are:

Li0 (s) + Ag+ (aq) Li+ (aq) + Ag0 (s)

Oxidation Half-Reaction:

Reduction Half-Reaction:

_____9. I can balance a redox

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 B) 2 C) 3 D) 6


conditions needed for a

spontaneous redox reaction to

occur.

A spontaneous redox reaction will occur if the metal being oxidized is

__________________________ on Table J, or the nonmetal reduced is

__________________________ on Table J.

_____11. I can state the two

types of electrochemical cells.

The two types of electrochemical cells are:

_____________________________ and _____________________________

_____12. I can compare the two

types of electrochemical cells.

Voltaic Electrolytic

Components

Oxidation Site

Reduction Site

Anode Charge

Cathode Charge

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


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?

_____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 cathode increases during the operation 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.


about an electrolytic cell.

The diagram below represents an electrochemical 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


_____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_______________.

_____2. I can define organic

compound, saturated

hydrocarbon, unsaturated

hydrocarbon, isomers, and alkyl

group.

Definitions: Organic Compound:

Saturated Hydrocarbon:

Unsaturated Hydrocarbon:

Isomers:

Alkyl Group:

_____3. Given the formula, I can

determine if a compound is a

hydrocarbon.

_____4. I can expand a

condensed structural formula to

show the structural formula of an

organic compound.

Draw the complete structural formula for CH3CH2CH2CH2CH3.

Draw the complete structural formula for CH3CHCHCH3.

_____5. Given the formula, I can

determine which compound has

the highest boiling point.

Which of the following hydrocarbons has the highest boiling point?

A) C3H8 B) C4H10 C) CH4 D) C2H6

_____6. Given the name or

formula, I can use Table Q to

determine which class of

hydrocarbons a compound

belongs.

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: __________________

_____8. Given the structural

formula, I can determine which

homologous series a hydrocarbon

belongs to.

Determine the homologous series of hydrocarbons to which each of the following belongs:

belongs to the _______________________ series.

belongs to the _______________________ series.

belongs to the ________________________ series.


use Table P to determine how

many carbons atoms are in a

compound.

Determine how many carbon atoms are in each of the following

compounds:

Decane: ________________________ Ethene: _________________________

3-Nonene: ______________________ 1-Pentyne: _______________________

Propene: ______________________ Methane: ________________________

_____10. Given a list of

compounds, I can determine

which ones are isomers.

_____11. I can draw an isomer of

a given compound.

Draw an isomer of the following:

_____12. I can use Tables P & Q

to name hydrocarbons.

Name the following hydrocarbons:

_____13. I can use Tables P & Q

to draw hydrocarbons.

Draw the following hydrocarbons:

propyne 3-hexene

3-ethyl hexane 2,3-dimethyl butane

_____14. Given a structural

formula, I can identify which

class of organic compounds a

substance belongs to.

_____15. I can use Tables P & R

to name simple compounds in

any of the classes of organic

compounds.

Name the following organic compounds:

_____16. I can use Tables P & R

to draw simple compounds in

any of the classes of organic

compounds.

Draw the following organic compounds:

2,2-dichloropentane 2-butanol

dimethyl ether ethanal

2-pentanone propanoic acid

methyl butanoate ethanamine

pentanamide ethyl propanoate

_____17. I can describe the 7

types of organic reactions.

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 _____________________________ 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.

_____19. I can complete an

organic reaction by predicting

the missing products.

Complete the following reactions:

Documents

Scientific Measurement curve, I can determine which sections of the curve show changes in kinetic energy. On the graph, circle the sections that show a change in kinetic energy. _____27