MR. PEPE’S CHEMISTRY CLASS RULES AND REGULATIONS ................................ 8
INTRO.................................................................................................................. 10
PRACTICING CONVERSIONS WORKSHEET ....................................................... 11
UNIT CONVERSIONS PRACTICE ........................................................................ 14
SCIENTIFIC NOTATION AND UNIT PREFIXES ...................................................... 16
SIGNIFICANT FIGURES WORKSHEET .................................................................. 18
CHAPTER 1 ......................................................................................................... 20
PHYSICAL BEHAVIOR OF MATTER ..................................................................... 20
HEAT AND SPECIFIC HEAT WORKSHEET ............................................................ 21
PHASE CHANGE WORKSHEET............................................................................ 23
CHAPTER 2 ......................................................................................................... 25
ATOMIC CONCEPTS .......................................................................................... 25
SUBATOMIC PARTICLES ..................................................................................... 26
LIGHT AND PHOTONS WORKSHEET ................................................................... 28
BRIGHT LINE SPECTRA AND PRINCIPAL ENERGY LEVELS.................................. 29
CHAPTER 3 ......................................................................................................... 30
NUCLEAR CHEMISTRY ........................................................................................ 30
NUCLEAR CHEMISTRY WORKSHEET ................................................................... 31
NUCLEAR CHEMISTRY WORKSHEET II ................................................................ 33
1
NUCLEAR APPLICATIONS .................................................................................. 34
ISLAND OF STABILITY ......................................................................................... 35
NUCLEAR REACTORS ......................................................................................... 36
CHAPTER 4 ......................................................................................................... 37
PERIODIC TABLE ................................................................................................. 37
PERIODIC TRENDS WORKSHEET ......................................................................... 38
PERIODIC TRENDS CLASS WORKSHEET ............................................................. 39
LEWIS DOT STRUCTURES ..................................................................................... 41
METALS, METALLOIDS, AND NONMETALS ......................................................... 42
CHAPTER 5 ......................................................................................................... 43
CHEMICAL BONDING........................................................................................ 43
INTRODUCTION TO BONDING .......................................................................... 44
CHEMICAL BONDING CLASS WORKSHEET ....................................................... 45
ELECTRON CONFIGURATIONS WORKSHEET ..................................................... 47
NAMING COMPOUNDS WORKSHEET ............................................................... 48
NAMING IONIC COMPOUNDS ......................................................................... 50
NAMING COVALENT COMPOUNDS WORKSHEET ............................................ 52
LEWIS DOT STRUCTURES OVERVIEW .................................................................. 53
LEWIS STRUCTURES ............................................................................................. 54
INTERMOLECULAR FORCES ............................................................................... 55
2
INTERMOLECULAR FORCES WORKSHEET .......................................................... 56
CHEMICAL FORMULA WRITING WORKSHEET.................................................... 57
LEWIS STRUCTURES, VSEPR, POLARITY, IM FORCES .......................................... 58
EARTH: THE WATER PLANET ................................................................................ 60
MOLE POEM ....................................................................................................... 61
CHAPTER 6 ......................................................................................................... 62
MOLES / STIOCHIOMETRY ................................................................................. 62
MOLAR CONVERSIONS WORKSHEET ................................................................ 62
MOLES WORKSHEET ........................................................................................... 63
GRAMS TO MOLES, MOLES TO GRAMS ............................................................ 64
MOLES, MOLECULES, AND GRAMS WORKSHEET .............................................. 66
PERCENT, ACTUAL, AND THEORETICAL YIELD .................................................. 68
BALANCING CHEMICAL EQUATIONS ............................................................... 70
SIX TYPES OF CHEMICAL REACTION WORKSHEET............................................ 71
WORD EQUATIONS WORKSHEET ....................................................................... 72
A VOYAGE THROUGH EQUATIONS .................................................................. 73
HYDRATES WORKSHEET ...................................................................................... 77
MOLECULAR FORMULA WORKSHEET ................................................................ 78
PERCENTAGE COMPOSITION WORKSHEET....................................................... 79
PERCENT COMPOSITION AND MOLECULAR FORMULA WORKSHEET .............. 81
3
STOICHIOMETRY: MOLE-MOLE CALCULATIONS ............................................. 82
STOICHIOMETRY MOLE-MASS CALCULATIONS .............................................. 84
LIMITING REAGENT WORKSHEET ....................................................................... 86
CHAPTER 7 ......................................................................................................... 87
GAS LAWS .......................................................................................................... 87
KINETIC THEORY OF GASES ............................................................................... 88
DESCRIPTION OF MEASUREMENTS FOR GASES ................................................ 89
GAS MEASUREMENTS ........................................................................................ 90
BOYLES’ LAW ...................................................................................................... 91
CHARLES’ LAW WORKSHEET .............................................................................. 93
GAY – LUSSAC WORKSHEET .............................................................................. 95
COMBINED GAS LAW PROBLEMS ..................................................................... 96
THE IDEAL AND COMBINED GAS LAWS ............................................................ 98
MOLAR VOLUME, DENSITY, AND VOLUME-VOLUME PROBLEMS .................... 99
CHAPTER 9 ....................................................................................................... 101
SOLUTIONS ....................................................................................................... 101
SOLUTIONS VOCABULARY .............................................................................. 102
SOLUBILITY CURVE TABLE G............................................................................. 104
TABLE F – SOLUBILITY GUIDELINES ................................................................... 106
MOLARITY PRACTICE PROBLEMS .................................................................... 107
4
DILUTION PROBLEMS........................................................................................ 109
M1V1 = M2V2 ..................................................................................................... 109
PARTS PER MILLION (PPM) ............................................................................... 110
CHAPTER 9 ....................................................................................................... 111
KINETICS/EQUILIBRIUM .................................................................................... 111
KINETICS ........................................................................................................... 112
POTENTIAL ENERGY DIAGRAM - 1 .................................................................. 116
POTENTIAL ENERGY DIAGRAM - 2 .................................................................. 117
CONSTRUCTING A POTENTIAL ENERGY DIAGRAM........................................ 118
ENDOTHERMIC OR EXOTHERMIC (TABLE H) ................................................... 119
REVIEW OF EQUILIBRIUM ................................................................................. 121
COLLISION THEORY ......................................................................................... 123
THERMODYNAMICS WORKSHEET .................................................................... 124
CHANGES OF STATE ........................................................................................ 125
CHAPTER 10 ..................................................................................................... 126
ACIDS/BASES ................................................................................................... 126
ACIDS AND BASES WORKSHEET - (PART I) ..................................................... 127
ACIDS AND BASES WORKSHEET - (PART II) .................................................... 128
ACID AND BASE WORKSHEET III...................................................................... 129
TABLE M ............................................................................................................ 131
5
NEUTRALIZATION.............................................................................................. 133
PH CALCULATIONS .......................................................................................... 135
CHAPTER 11 ..................................................................................................... 136
OXIDATION-REDUCTION ................................................................................. 136
HOW TO WRITE AND BALANCE A REDOX EQUATION.................................... 137
REDOX WORKSHEET ......................................................................................... 139
OXIDATION STATE ............................................................................................ 140
OXIDATION AND REDUCTION PRACTICE ....................................................... 141
ELECTROCHEMICAL CELL ................................................................................ 142
ELECTROLYTIC CELL ......................................................................................... 144
ELECTROPLATING ............................................................................................. 145
ELECTROCHEMISTRY ........................................................................................ 146
ELECTROCHEMICAL CELL ................................................................................ 148
CHAPTER 11 ..................................................................................................... 149
ORGANIC CHEMISTRY ..................................................................................... 149
ORGANIC FUNCTIONAL GROUPS................................................................... 150
ORGANIC CHEMISTRY ..................................................................................... 151
The Alkanes ..............................................................................................................................................151
The Alkenes ..............................................................................................................................................151
The Alkynes ..............................................................................................................................................151
The Benzene Series ................................................................................................................................152
6
The Haloalkanes .....................................................................................................................................152
The Alcohols.............................................................................................................................................153
The Aldehydes ........................................................................................................................................153
The Ketones ..............................................................................................................................................153
The Carboxylic acids ............................................................................................................................154
The Ethers ..................................................................................................................................................154
The Esters...................................................................................................................................................154
The Amines ...............................................................................................................................................155
The Amides ...............................................................................................................................................155
ORGANIC CHEMISTRY ..................................................................................... 156
VOCABULARY .................................................................................................. 158
7
Mr. Pepe’s Chemistry Class Rules and Regulations
Expectations:
Respect is a mutual relationship for both student and teacher.
o In order to maintain an orderly classroom, the following is a list
of acceptable behavior:
No talking when someone is speaking.
Come to class on time.
If you disagree with a teacher’s decision, it is your right
to discuss it with the teacher after class.
Come to class prepared.
Respect other students in the class, as this in an
environment in which everyone is here to learn.
Cheating and/or Plagiarism
Both are unacceptable. Any student caught copying, cheating or
plagiarizing or allowing other students to copy their schoolwork will
result in a zero for that assignment.
If you are caught cheating, it is an automatic zero and a
conference will be scheduled with your parents and school staff.
All local school rules are in effect
NO eating during class periods at all due to safety issues.
You should be here on time.
Remember, there is no directed study, so come to class every day
of the week.
I will post all materials on http://web.mac.com/empepe/
, so there is no excuse for missing anything.
Homework:
Should be completed on the date due.
If you are absent, I expect you to check with a classmate, me, to
find out what assignments were given.
For each day that you are absent you will get an extension of one
day to turn in the assignment.
You will be allowed 2 late homework’s per quarter. These
assignments can be handed in up to one week late without losing
credit.
If you are unable to complete a homework assignment because
you found it too difficult, you can still receive full credit if you write
out the question(s) and come see me within a couple of days of
when it was due.
There will be a separate contract for Lab Safety. The lab is a safe
place as along as the rules are observed.
8
If for any reason you are having a difficult time, please
come see me.
Grading Policy:
HOMEWORK: WILL CHECKED 20 TIMES PER QUARTER. EACH MISSED
ASSIGNMENT WIL RESULT IN THE LOSS OF ONE AVERAGE POINT.
Essentially, homework/class work could make the difference between
getting a C or an A in this class.
Things to bring to EACH AND EVERY class:
1. A positive attitude and the desire to succeed.
2. Pens
3. Pencils (2 or 3 colors works best)
4. Regents Reference Table (1st copy is free)
5. Calculator
6. Notebook for class
7. Composition Graph Notebook for lab (No Spirals)
8. Chemistry Lab Book on Lab days
Name (Print) ___________________________________
Signature ______________________________________ Date ___________
Parents Name __________________________________
Signature______________________________________ Date ___________
9
Intro
10
Practicing Conversions Worksheet
10 dimes = 1 dollar
20 nickels = 1 dollar
4 quarters = 1 dollar
These conversions are known and are easy because you are
familiar with them.
Practice these conversions:
1) How many quarters are in 6 dollars?
2) How many dimes are in 11 dollars?
3) How many nickels are in 3 dollars?
_____________________________________________________________
How did you get these answers? ....You simply multiplied The
conversion number by however many dollars you had.... Like
this:
3 dollars X 20 nickels = 60 nickels
1 dollar
_____________________________________________________________
Now try these:
4) How many dollars do you have if you had 40 quarters?
5) How many dollars do you have if you had 60 dimes?
6) How many dollars do you have if you had 160 nickels?
11
How did you get these answers? ....You simply divided The
conversion number into however many coins you had.... Like
this:
160 nickels X 1 dollar = 8 dollars
20 nickels
_____________________________________________________________
You are able to do this because you can (in your head)
visualize the conversions of dollars to coins and visa-versa. It’s
easy. It makes sense (and cents). Doing conversions in
chemistry is no different. The only problem is that you are not
used to the units. Everyone knows want a nickel is... but what is
a kilojoule?
First, accept that these “units” that you will learn are just simple
words to describe the size of a measurement. A unit describes
how big, small, heavy, long, or intense some measurement is. A
meter is a unit of length. It describes how long something is (just
like an inch or a foot or a mile are units of length). Just like there
are 12 inches in a foot, there are 1000 meters in a kilometer. In
fact, any unit with the prefix kilo- in front of it means 1000 times
that unit.
1 kilometer = 1000 meters 1 kilogram = 1000 grams
1 kilojoule = 1000 joules 1 kilocalorie = 1000 calories
How do you convert kilometers to meters? Try these:
1) How many meters are in 6 kilometers? ...Well if 1 kilometer =
1000 meters,
then, 6 kilometers must be 6 times greater.
6 kilometers X 1000 meters = 6000 meters
1 kilometer
12
2) How many kilometers are in 200 meters? ..Well if 1 kilometer
= 1000 meters, then, 200 meters must be divided into 1000.
200 meters X 1 kilometer = 0.2 kilometers
1000 meters
Other conversions work the same way. You either multiply or
divide by some number. Your units will always cancel, leaving
you with just the units you want left. Practice the rest of these
conversions on the following page:
Convert the following:
112 kilograms into grams:
32 grams into kilograms
0.5 kilometers into meters
12 joules into kilojoules
13
Unit Conversions Practice
There are 5280 feet in one mile
There are 0.034 ounces in one milliliter
There are 0.454 kg in one pound
There are 1.6 kilometers in one mile
There are 73 gallons in 2 barrels
There are 1.05 quarts in one liter
There are 4 quarts in one gallon
Do the following one-step unit conversions:
1) Convert 23 miles to feet.
2) Convert 120 lbs to kilograms.
3) Convert 451 mL to ounces.
4) Convert 6 feet to miles.
5) Convert 4 quarts to liters.
6) Convert 0.045 barrels to gallons.
14
Do the following multi-step unit conversions:
7) Convert 75 minutes to days.
8) Convert 46 inches to miles (there are 12 inches in one foot).
9) Convert 65 ounces to liters. (There are 1000 mL in one liter).
10) Convert one million seconds to years.
11) Convert 12 liters to barrels.
12) Find your age in seconds.
15
Scientific Notation and Unit Prefixes
Make the following conversions:
1) 3.4 liters to milliliters 6) 45 meters to centimeters
2) 876 millimeters to meters 7) 11.7 grams to kilograms
3) 78,999 milligrams to grams 8) 0.0009 kiloliters to liters
4) 0.9 centigrams to grams 9) 44 centimeters to meters
5) 112 meters to millimeters 10) 277 kilograms to grams
________________________________________________________________
Convert the following to scientific notation:
11) 45,700 ______________________________
12) 0.009 ______________________________
13) 23 ______________________________
14) 0.9 ______________________________
15) 24,212,000 ______________________________
16) 0.000665 ______________________________
16
Convert the following to scientific notation:
17) 21.9 ______________________________
18) 0.00332 ______________________________
19) 321 ______________________________
20) 0.119 ______________________________
21) 1492 ______________________________
22) 0.2713 ______________________________
23) 314159 ______________________________
24) 6022 ______________________________
25) 0.12011 ______________________________
Convert the following numbers in scientific notation to expanded form:
26) 3.825 x 103 ______________________________
27) 6.3 x 104 ______________________________
28) 2.3 x 10-2 ______________________________
29) 4.44 x 10-6 ______________________________
30) 7.121 x 109 ______________________________
31) 1.2 x 10-1 ______________________________
32) 1.8 x 102 ______________________________
33) 8.1 x 10-4 ______________________________
34) 6.7 x 105 ______________________________
35) 3.4 x 107 ______________________________
17
Significant Figures Worksheet
The number of digits in a number that tell you useful information. For
example, when you weigh yourself on a bathroom scale, it says something
like 150 pounds rather than 150.32843737 pounds. Why? Because the
thing can only weigh accurately to the nearest pound. Any other digits
that are on this number don't mean anything, because they're probably
wrong anyway.
Significant digits, which are also called significant figures, are very
important in Chemistry. Each recorded measurement has a certain
number of significant digits. Calculations done on these measurements
must follow the rules for significant digits. The significance of a digit has to
do with whether it represents a true measurement or not. Any digit that is
actually measured or estimated will be considered significant.
Placeholders, or digits that have not been measured or estimated, are not
considered significant. The rules for determining the significance of a digit
will follow.
Rules For Significant Digits
1. Digits from 1-9 are always significant.
2. Zeros between two other significant digits are always significant
3. One or more additional zeros to the right of both the decimal place
and another significant digit are significant.
4. Zeros used solely for spacing the decimal point (placeholders) are
not significant.
Recognizing significant digits will become much easier over time, as
you continue to practice the rules. Below are some examples, which
show the number of significant digits in a group of numbers, and an
explanation why the digits are significant.
Table 1.1 Examples of Significant Digits
EXAMPLES # OF SIG. DIG. COMMENT
453 kg 3 All non-zero digits are always significant.
5057 L 4 Zeros between 2 sig. dig. are significant.
5.00 3 Additional zeros to the right of decimal and a sig. dig. are
significant.
0.007 1 Placeholders are not sig.
18
How many significant figures are in each of the following numbers?
1) 5.40 ____ 6) 1.2 x 103 ____
2) 210 ____ 7) 0.00120 ____
3) 801.5 ____ 8) 0.0102 ____
4) 1,000 ____ 9) 9.010 x 10-6 ____
5) 101.0100 ____ 10) 2,370.0 ____
11) Why are significant figures important when taking data in the
laboratory?
12) Why are significant figures NOT important when solving problems in
your math class?
13) Using two different instruments, I measured the length of my foot to
be 27 centimeters and 27.00 centimeters. Explain the difference
between these two measurements.
14) I can lift a 20 kilogram weight over my head ten times before I get
tired. Write this measurement to the correct number of significant
figures.
19
Chapter 1
Physical Behavior of Matter
20
Heat and Specific Heat Worksheet
Heat = mass x specific heat x change in temperature
(oC)
q = m x C x ∆t
Example: If a 5.00 gram sample of a glass of water is heated
from 20.0 oC to 30.0 oC, How much heat was gained by the
water?
Given: The specific heat of water is 4.18 joules/gram.oC
q = m x C x ∆t
q = 5.00 grams x 4.18 J/g.oC x (30.0 oC - 20.0 oC)
q = 5.00 grams x 4.18 J/g.oC x 10.0 oC
q = 5.00 grams x 4.18 J x 10.0 oC = 209 joules
gram.oC
1. If 25.0 grams of water is heated from 25.0 oC to 30.0 oC,
how much heat was gained by the water?
2. If 15.0 grams of water is cooled from 20.0 oC to 15.0 oC,
how much heat was lost by the water?
21
3. If a sample of water that is heated from 25.0 oC to 35.0 oC
gains 250.8 joules of heat, what was the mass of the water
sample.
4. If 30.0 grams of water is heated and gains 376.2 joules of
heat, what was the change in temperature of the water
sample?
5. If 10.0 grams of water is heated from 25.0 oC and gains 334.4
joules of heat, what was the final temperature of the water
sample?
22
PHASE CHANGE WORKSHEET
Phase Change Diagram for 50.0 grams of Water
T 120.0
E
M
GAS
P
E 100.0 Heat of Vaporization
R
A
T LIQUID
U
R Heat of Fusion
E 0
(OC)
SOLID
-10.0
Heat Absorbed (joules) ---------------------------->
A 50.0 gram sample of water is heated from –10.0 OC to 120.0 OC (thus
going through all three phases of matter). How much heat is absorbed in
this process?
1. First find the heat absorbed from the heating of ice at –10.0 OC to ice
at 0.0 OC using the specific heat formula that you’ve been taught. (Note
that ice has a specific heat of 2.09 joules/g.OC)
23
2. Next find the heat of fusion for 50.0 grams of ice-water. (Use Reference
Table B)
3. Next find the heat absorbed from the heating of water from 0.0 OC to
100.0 OC using the specific heat formula that you’ve been taught. (note
that water has a specific heat of 4.18 joules/g.OC)
4. Next find the heat of vaporization for 50.0 grams of water-steam. (Use
Reference Table B)
5. Next find the heat absorbed from the heating of steam from 100.0 OC
to 120.0 OC using the specific heat formula that you’ve been taught.
(Note that steam has a specific heat of 2.09 joules/g.OC)
6. Now add up all the heats. What’s your total?
24
Chapter 2
Atomic Concepts
25
Subatomic Particles
The table below contains information about several elements. In each
case, enough information has been provided for you to fill in the blanks.
Assume all atoms are neutral.
# of # of
Isotope Nuclear Atomic Mass # of
Electron Neutron
Name Symbol Number Number Protons
s s
1. calcium-40
2. 12 24
3. 1 2
4.
5. 26 30
6. 201 80
7. 17 18
PART B – AVERAGE ATOMIC MASS
8. Calculate the average atomic mass for neon if its abundance in
nature is 90.5% neon-20, 0.3% neon-21, and 9.2% neon-22.
9. Calculate the average atomic mass of silver if 13 out of 25 atoms are
silver-107 and 12 out of 25 atoms are silver-109.
26
10. Distinguish between mass number, relative atomic mass, and average
atomic mass.
27
Light and Photons Worksheet
1. Define the following terms:
continuous spectrum
excited state
ground state
line-emission spectrum
orbital
photoelectric effect
photon
quanta of energy (quantum)
2. If the wavelength of radiation for a specific energy is doubled, what
happens to the frequency?
3. If the frequency of radiation for a specific energy is doubled, what
happens to the energy?
4. Draw a picture of the Bohr model of the hydrogen atom showing the
excited and ground states; show how a photon of light energy is released
in the picture.
What is so unique about the line emission spectra of an element?
28
Bright Line Spectra And Principal Energy Levels
1. The characteristic bright-line spectrum of an element is produced
when electrons
(1) absorb quanta and return to lower energy levels
(2) absorb quanta and move to higher energy levels
(3) release quanta and return to lower energy levels
(4) release quanta and move to higher energy levels
2. Answer the following questions:
3. The excited atom has an electron configuration of 1-6. What is the
atom? Write the ground state of the same atom.
4. The number of electrons that the first principal energy level can
hold is ___________.
5. The number of electrons that the second principal energy level can
hold is ___________.
6. The number of electrons that the third principal energy level can
hold is ___________.
7. The number of electrons that the fourth principal energy level can
hold is ___________.
8. The formula for determining the number of electrons that a
principal energy can hold is: ____________
29
Chapter 3
Nuclear Chemistry
30
NUCLEAR CHEMISTRY WORKSHEET
1. List the three types of natural transmutations (radiation).
A.
B.
C.
2. What is an alpha particle? What is a beta particle?
3. Define the term half-life?
4. In the following nuclear reactions, what is X?
226 222
a. Ra --> Rn + X
88 86
232 232
Th --> Pa + X
90 91
b.
232 4
Th --> X + He
90 2
c.
5. After 50 days, only 5 grams remained from a 10 gram sample of a
substance. What is the half-life of that substance?
6. After 50 days, only 25 grams remained from a 100 gram sample of a
substance. What is the half-life of that substance?
31
7. Carbon-14 has a half-life of 5730 years. What fraction of a 1 gram
sample would remain after 17,190 years?
8. Why is carbon dating? Why would you use carbon to date an ancient
piece of wood?
9. What is carbon dating not used in determining the age of dinosaur
bones?
32
Nuclear Chemistry Worksheet II
Using your knowledge of nuclear chemistry, write the equations for the
following processes:
1) The alpha decay of radon-198
2) The beta decay of uranium-237
3) Positron emission from silicon-26
4) Sodium-22 undergoes electron capture
5) What is the difference between nuclear fusion and nuclear fission?
6) What is a “mass defect” and why is it important?
7) Name three uses for nuclear reactions.
33
NUCLEAR APPLICATIONS
In order to learn more about how nuclear chemistry affects our lives, you
will research one of the topics below using your textbook and other
available materials. In teams, you will create a poster on your topic and
share the information you have learned with the class.
Your poster should:
include a diagram appropriate to your topic
be neat and attractive
You should:
have each member of your team participate in the preparation and
presentation
be able to explain the fundamentals of your topic – You’re not expected
to be an expert, but quality of information is better than mere quantity.
Topic Choices:
Medical applications – Explain how radioactive isotopes are used for
medical diagnosis and treatment.
Nuclear reactors – Explain how a nuclear fission reactor works. Include
a diagram of a reactor.
Nuclear power – Compare the use of fission and fusion as sources of
energy. Briefly describe how each method works (fuel, reactor type,
etc.) and discuss the pros and cons of each method.
Nuclear waste – Explain the problems associated with nuclear waste
and methods used for storage and disposal
Nuclear weapons – Explain how fission (atomic) and fusion (hydrogen)
bombs work. Include simple diagrams for each type of bomb.
Radiation exposure & detection – Explain sources of radiation, safe
exposure levels, and how it is detected.
Radioactive dating – Explain how carbon-14 and other isotopes are
used to determine the age of an object.
Synthetic elements – Explain how nuclear reactions are used to create
new elements
34
Island of Stability
35
Nuclear Reactors
36
Chapter 4
Periodic Table
37
Periodic Trends Worksheet
1) Rank the following elements by increasing atomic radius: carbon,
aluminum, oxygen, potassium.
2) Rank the following elements by increasing electronegativity: sulfur,
oxygen, neon, aluminum.
3) What is the difference between electron affinity and ionization
energy?
4) Why does fluorine have a higher ionization energy than iodine?
5) Why do elements in the same family generally have similar
properties?
38
Periodic Trends Class Worksheet
1. What are atomic radii?
2. What happens to the atomic radius of an atom as you go down a
group?
3. What happens to the atomic radius of an atom as you go across a
period?
4. What is Ionization Energy? Show an equation that demonstrates
Ionization Energy.
5. What happens to the Ionization Energy of an atom as you go down a
group?
6. What happens to the Ionization Energy of an atom as you go across a
period?
7. What is Electron Affinity?
39
8. What is Electronegativity?
9. What happens to the Electronegativity of an atom as you go down a
group?
10. What happens to the Electronegativity of an atom as you go across a
period?
11. Find the Ionization Energies of the following atoms from your reference
table:
a) lithium
b) sodium
c) magnesium
d) fluorine
e) chlorine
f) oxygen
12. Find the Electronegativities of the following atoms from your reference
table:
a) lithium
b) sodium
c) magnesium
d) fluorine
e) chlorine
f) oxygen
13. What atom is the most electronegative?
40
Lewis Dot Structures
Please draw the Lewis dot structures of the following elements on the
Periodic table:
1
2
3
41
Metals, Metalloids, and Nonmetals
1. What’s the difference between a chemical and a physical property?
Give two examples of each and explain how they are different.
2. Give four properties that are generally present in metals.
3. If steel (a metal) is hard and granite (a nonmetal) is hard, why don’t
we make automobile engines out of granite?
4. What are metalloids used for, and how does this affect modern
technology?
42
Chapter 5
Chemical Bonding
43
Introduction to Bonding
Define chemical bond –
Bonding Comparison Chart
IONIC COVALENT METALLIC
Types of
Atoms
Involved
Method of
Bond
Formation
Type of
Structure
Physical
State
Melting
Point
Solubility in
Water
Electrical
Conductivit
y
Other
Properties
Image
44
Chemical Bonding Class Worksheet
1. What are Valence Electrons?
2. What is a covalent bond?
3. a) How many electrons (normally) does each atom want to have in its
outer shell to be complete?
b) What word is used to describe this number of electrons?
4. What is a diatomic molecule? Show seven of them?
5. What is an Ionic bond?
6. How can you determine if a molecule is Ionic or covalent?
7. What is a polar covalent bond? Give an example.
8. What is a non-polar covalent bond? Give an example.
9. Show the Ionic bonding between Potassium (K) and Fluorine (F).
45
10. a) Show the covalent bonding between Hydrogen (H) and Bromine
(Br).
b) Is this bond polar or non-polar?
11. a) Show the covalent bonding between two Chlorine atoms.
b) Is this bond polar or non-polar?
12. a) Show the covalent bonding of a water molecule.
b) Is this a polar or non-polar molecule?
13. What is a double bond? What is a triple bond?
14. Show the bonding for CO2.
15. Show the bonding for N2.
46
Electron Configurations Worksheet
Write the complete ground state electron configurations for the following:
1) lithium ________________________________________
2) oxygen ________________________________________
3) calcium ________________________________________
4) titanium ________________________________________
5) rubidium ________________________________________
6) lead ________________________________________
7) erbium ________________________________________
Write the abbreviated ground state electron configurations for the
following:
8) helium ________________________________________
9) nitrogen ________________________________________
10) chlorine ________________________________________
11) iron ________________________________________
12) zinc ________________________________________
13) barium ________________________________________
14) polonium ________________________________________
47
Naming Compounds Worksheet
You will relearn how to name compounds by separating into
your groups and helping each other on the following
categories:
SECTION I - Binary Ionic compounds: Name the following:
NaCl Al2O3
MgCl2 Na2S
Li2O FeO
CaO Fe2O3
AlF3 FeCl3
SECTION II - Binary Ionic compounds: Find the formulas for the following:
potassium chloride iron (II) sulfide
calcium chloride iron (III) sulfide
sodium oxide copper(I) oxide
aluminum iodide copper (II) oxide
aluminum sulfide iron (III) bromide
SECTION III - Polyatomic compounds: Find the formulas for the
following:
aluminum hydroxide calcium phosphate
ammonium sulfide sodium sulfate
ammonium sulfate magnesium phosphate
sodium chlorate lithium carbonate
magnesium chlorate calcium carbonate
SECTION IV - Polyatomic compounds: Name the following:
NaOH AlPO4
Mg(OH)2 Na2SO4
Li2(CO3) (NH4)2O
Ca(SO4) (NH4)2SO4
Al(ClO4)3 Fe2(SO4)3
48
SECTION V - Molecular compounds
Name the following:
NO
CO2
N2O
CO
SO3
SECTION VI - Molecular compounds
Find the formulas for the following:
dinitrogen pentoxide
carbon tetrafluoride
dihydrogen monosulfide
carbon monosulfide
dinitrogen tetroxide
49
Naming Ionic Compounds
Give the name and molar mass of the following ionic compounds:
1) Na2CO3 ____________________________________________________
2) NaOH _____________________________________________________
3) MgBr2 _____________________________________________________
4) KCl _______________________________________________________
5) FeCl2 ______________________________________________________
6) FeCl3 ______________________________________________________
7) Zn(OH)2 ___________________________________________________
8) Be2SO4 ___________________________________________________
9) CrF2 ______________________________________________________
10) Al2S3 _____________________________________________________
11) PbO ______________________________________________________
12) Li3PO4 ____________________________________________________
13) TiI4 _______________________________________________________
14) Co3N2 ____________________________________________________
15) Mg3P2 ____________________________________________________
16) Ga(NO2)3 __________________________________________________
17) Ag2SO3 ____________________________________________________
18) NH4OH ____________________________________________________
19) Al(CN)3 ____________________________________________________
20) Be(CH3COO)2 ______________________________________________
50
For the following compounds, give the formulas and the molar masses:
22) sodium phosphide ___________________________________________
23) magnesium nitrate ___________________________________________
24) lead (II) sulfite ______________________________________________
25) calcium phosphate ___________________________________________
26) ammonium sulfate ___________________________________________
27) silver cyanide _______________________________________________
28) aluminum sulfide ____________________________________________
29) beryllium chloride ____________________________________________
30) copper (I) arsenide ___________________________________________
31) iron (III) oxide _______________________________________________
32) gallium nitride _______________________________________________
33) iron (II) bromide _____________________________________________
34) vanadium (V) phosphate ______________________________________
35) calcium oxide _______________________________________________
36) magnesium acetate __________________________________________
37) aluminum sulfate ____________________________________________
38) copper (I) carbonate __________________________________________
39) barium oxide ________________________________________________
40) ammonium sulfite ____________________________________________
41) silver bromide _______________________________________________
42) lead (IV) nitrite ______________________________________________
51
Naming Covalent Compounds Worksheet
Write the formulas for the following covalent compounds:
1) antimony tribromide __________________________________
2) hexaboron silicide __________________________________
3) chlorine dioxide __________________________________
4) hydrogen iodide __________________________________
5) iodine pentafluoride __________________________________
6) dinitrogen trioxide __________________________________
7) ammonia __________________________________
8) phosphorus triiodide __________________________________
Write the names for the following covalent compounds:
9) P4S5 __________________________________
10) O2 __________________________________
11) SeF6 __________________________________
12) Si2Br6 __________________________________
13) SCl4 __________________________________
14) CH4 __________________________________
15) B2Si __________________________________
16) NF3 __________________________________
52
Lewis Dot Structures Overview
List the Lewis dot structure for atoms in the following groups:
1 2 13 14 15 16 17 18
List the Lewis Dot structures for the following groups in IONIC bonding.
1 2 13 14 15 16 17 18
List the Lew Dot Structures for the following groups in covalent bonding.
1 2 13 14 15 16 17 18
53
Lewis Structures
1. Draw the Lewis structures of the first twenty atoms on the
periodic table.
1
2
3
2. Draw the Lewis dot structures of the following:
1. Hydrogen atom
2. Hydrogen cation
3. Hydrogen anion
4. H2
5. Cl2
6. O2
7. NaCl
8. MgCl2
9. PCl3
10. NH4+1
11. H3O+
54
Intermolecular Forces
For questions 1-5, identify the main type of intermolecular force in each
compound:
1) carbon disulfide
2) ammonia
3) oxygen
4) CH2F2
5) C2H6
6) What type of bond usually accounts for the unusually high boiling
point of water ?
7) Explain how dipole-dipole forces cause molecules to be attracted
to one another.
8) When you go down Group 17 or down Group 18, why does the
boiling point increase?
9) Which liquid in Table H has the weakest intermolecular forces?
Explain how you reached your answer.
10) Compare and contrast asymmetrical and symmetrical molecules.
55
Intermolecular Forces Worksheet
1) Using your knowledge of molecular structure, identify the main
intermolecular force in the following compounds. You may find it
useful to draw Lewis structures to find your answer.
a) PF3 _____________________________
b) H2CO _____________________________
c) HF _____________________________
2) Explain how dipole-dipole forces cause molecules to be attracted
to one another.
3) Rank the following compounds from lowest to highest boiling point:
calcium carbonate, methane, methanol (CH4O), di-methyl ether
(CH3OCH3).
4) Explain why nonpolar molecules usually have much lower surface
tension than polar ones.
56
Chemical Formula Writing Worksheet
Write chemical formulas for the compounds in each box. The names are
found by finding the intersection between the cations and anions.
Example: The first box is the intersection between the “zinc” cation and
the “chloride” anion, so you should write “ZnCl2”, as shown.
zinc iron (II) iron (III) gallium silver lead (IV)
chloride ZnCl2
acetate
nitrate
oxide
nitride
sulfate
Write the formulas for the following compounds:
1) copper (II) chloride ____________________________________
2) lithium acetate ____________________________________
3) vanadium (III) selenide ____________________________________
4) manganese (IV) nitride ____________________________________
5) beryllium oxide ____________________________________
6) sodium sulfate ____________________________________
7) aluminum arsenide ____________________________________
8) potassium permanganate ____________________________________
9) chromium (VI) cyanide ____________________________________
10) tin (II) sulfite ____________________________________
11) vanadium (V) fluoride ____________________________________
12) ammonium nitrate ____________________________________
57
Lewis Structures, VSEPR, Polarity, IM Forces
For each of the following molecules, draw the Lewis structure (with any
resonance structures, if applicable), indicate the molecular shapes and
bond angles, indicate the molecular polarity (if any), and identify the
major intermolecular force in each compound. Hint – in this worksheet, as
in all chemistry problems you’ll see, polyatomic ions aren’t drawn as big
lines of atoms.
1) carbon tetrafluoride
2) BF3
3) NF3
4) H2CS
58
5) carbonate ion
6) CH2F2
7) nitrate ion
8) O2
9) PF3
10) H2S
59
Earth: The Water Planet
The amount of water on Earth greatly exceeds that known on
or within any other planet in the solar system. Liquid water,
which is essential for life to survive, has unique and amazing
properties; it covers 70% of Earth’s surface. Where did all Earth’s
water come from?
If the Earth and solar system evolved from a swirling cloud of
dust and gas, almost no water would reside near Earth’s
present orbit. Any water (liquid or ice) that close to the Sun
would vaporize and be blown by solar wind to the outer
reaches of the solar system, a as we see happening with water
vapor in the tails of comets.
Did comets or meteorites deliver Earth’s water? Although
comets contain considerable water,b they could not have
brought much water to Earth, because comets contain too
much heavy hydrogen, relatively rare in Earth’s oceans.
Comets also contain too much argon. If comets were the
source of only 1% of Earth’s water, then, using evolutionists’
assumptions, our atmosphere would contain 400 times more
argon than it does.c The few types of meteorites that contain
considerable water also have too much heavy hydrogen.d
[Pages 268–318 explain why comets and some types of
meteorites contain so much water and heavy hydrogen.
Heavy hydrogen is described on page 276.]
These observations have caused some to conclude that water
was transported from the outer solar system to Earth by objects
that no longer exist. e If so, many of these “water tankers”
should have collided with the other inner planets (Mercury,
Venus, and Mars), producing water characteristics similar to
those of Earth. In fact, their water characteristics are not like
those of Earth. f Instead of imagining “water tankers” that
conveniently disappeared, perhaps we should ask if the Earth
was created with its water already present.
www.creationscience.com
60
Mole Poem
Carbon has six protons and six neutrons, it’s true.
Add them both together for the mass in a.m.u..
So if carbon-12 is what you have, 12 a.m.u. is carbon's mass.
But this is just one isotope, that is studied in this class.
And a mole's worth of carbon would then be 12 grams.
Remember this example and it will get you out of jams.
Atoms are quite tiny, too small to weigh just one.
Therefore, Avogadro sought to have a little fun.
He found the mole fits perfectly to tell the mass in grams.
You can find the mass of any substance if you add up its nucleons.
So if a carbon atom weighs 12 a.m.u., don't think it's lame.
A mole of carbon atoms weighs 12 grams, (Note the number is the same).
Now take, for an example, a mole of water clear.
Sum the mass of all atoms in the molecule you find there.
Then add them all together, and you will see its true,
That the mass of all the atoms is the mass in a.m.u..
That sum is 18, and its mass is in a.m.u..
And if you had a mole of water, the grams is 18 too.
Mr. Avogadro discovered what you need.
Made it more convenient to do measurements, indeed.
He found the number of particles and that number is the mole.
Use it in your calculations to help you reach your goal.
So memorize the number, and know it when its heard.
A mole is 6.02 times ten to the twenty-third.
The number of atoms found in one mole is also said to be
The number of molecules in a mole of compound of chemistry.
Don’t let it confuse you, for the number is the same.
Be it a number of atoms or a number of molecules, a mole is its name.
So memorize the number, and know it when its heard.
A mole is 6.02 times ten to the twenty-third.
61
Chapter 6
Moles / Stiochiometry
Molar Conversions Worksheet
1. Find the formulas of the following compounds, then find their Molar
Mass
(Note that some are ionic and some are molecular)
Formula Molar Mass
a. lithium phosphate _____________ ____________
b. magnesium phosphate _____________ ____________
c. sodium sulfide _____________ ____________
d. sodium sulfate _____________ ____________
e. dinitrogen tetraoxide _____________ ____________
f. water _____________ ____________
g. aluminum carbonate ____________ ____________
h. carbon dioxide _____________ ____________
62
Moles Worksheet
1) Define “mole”.
2) How many moles are present in 34 grams of Cu(OH)2?
3) How many moles are present in 2.45 x 1023 molecules of CH4?
4) How many grams are there in 3.4 x 1024 molecules of NH3?
5) How much does 4.2 moles of Ca(NO3)2 weigh?
6) What is the molar mass of MgO?
7) How are the terms “molar mass” and “atomic mass” different from
one another?
8) Which is a better unit for expressing molar mass, “amu” or
“grams/mole”?
63
GRAMS TO MOLES, MOLES TO GRAMS
Given the following, find the number of moles:
1) 30 grams of H3PO4
2) 25 grams of HF
3) 110 grams of NaHCO3
4) 1.1 grams of FeCl3
5) 987 grams of Ra(OH)2
6) 564 grams of copper
7) 12.3 grams of CO2
8) 89 grams of Pb(CH3COO)4
64
Given the following, find the number of grams:
9) 4 moles of Cu(CN)2
10) 5.6 moles of C6H6
11) 21.3 moles of BaCO3
12) 1.2 moles of (NH4)3PO3
13) 9.3 x 10-3 moles of SmO
14) 6.6 moles of ZnO
15) 5.4 moles of K2SO4
16) 88.4 moles of NI3
65
Moles, Molecules, and Grams Worksheet
1) How many molecules are there in 24 grams of FeF3?
2) How many molecules are there in 450 grams of Na2SO4?
3) How many grams are there in 2.3 x 1024 atoms of silver?
4) How many grams are there in 7.4 x 1023 molecules of AgNO3?
5) How many grams are there in 7.5 x 1023 molecules of H2SO4?
6) How many molecules are there in 122 grams of Cu(NO3)2?
7) How many grams are there in 9.4 x 1025 molecules of H2?
8) How many molecules are there in 230 grams of CoCl 2?
66
9) How many molecules are there in 2.3 grams of NH4SO2?
10) How many grams are there in 3.3 x 1023 molecules of N2I6?
11) How many molecules are there in 200 grams of CCl4?
12) How many grams are there in 1 x 1024 molecules of BCl3?
13) How many grams are there in 4.5 x 1022 molecules of Ba(NO2)2?
14) How many molecules are there in 9.34 grams of LiCl?
15) How many grams do 4.3 x 1021 molecules of UF6 weigh?
16) How many molecules are there in 230 grams of NH4OH?
67
Percent, Actual, and Theoretical Yield
1) LiOH + KCl LiCl + KOH
a) I began this reaction with 20 grams of lithium hydroxide. What is
my theoretical yield of lithium chloride?
b) I actually produced 6 grams of lithium chloride. What is my
percent yield?
2) C3H8 + 5 O2 3 CO2 + 4 H2O
a) If I start with 5 grams of C3H8, what is my theoretical yield of
water?
b) I got a percent yield of 75% How many grams of water did I
make?
3) Be + 2 HCl BeCl2 + H2
My theoretical yield of beryllium chloride was 10.7 grams. If my
actual yield was 4.5 grams, what was my percent yield?
4) 2 NaCl + CaO CaCl2 + Na2O
What is my theoretical yield of sodium oxide if I start with 20 grams
of calcium oxide?
68
5) FeBr2 + 2 KCl FeCl2 + 2 KBr
a) What is my theoretical yield of iron (II) chloride if I start with 34
grams of iron (II) bromide?
b) What is my percent yield of iron (II) chloride if my actual yield is 4
grams?
6) TiS + H2O H2S + TiO
What is my percent yield of titanium (II) oxide if I start with 20 grams
of titanium (II) sulfide and my actual yield of titanium (II) oxide is 22
grams?
7) U + 3 Br2 UBr6
What is my actual yield of uranium hexabromide if I start with 100
grams of uranium and get a percent yield of 83% ?
8) H2SO4 H2O + SO3
If I start with 89 grams of sulfuric acid and produce 7.1 grams of
water, what is my percent yield?
69
Balancing Chemical Equations
Balance the equations below:
1) ____ N2 + ____ H2 ____ NH3
2) ____ KClO3 ____ KCl + ____ O2
3) ____ NaCl + ____ F2 ____ NaF + ____ Cl2
4) ____ H2 + ____ O2 ____ H2O
5) ____ Pb(OH)2 + ____ HCl ____ H2O + ____ PbCl2
6) ____ AlBr3 + ____ K2SO4 ____ KBr + ____ Al2(SO4)3
7) ____ CH4 + ____ O2 ____ CO2 + ____ H2O
8) ____ C3H8 + ____ O2 ____ CO2 + ____ H2O
9) ____ C8H18 + ____ O2 ____ CO2 + ____ H2O
10) ____ FeCl3 + ____ NaOH ____ Fe(OH)3 + ____NaCl
11) ____ P + ____O2 ____P2O5
12) ____ Na + ____ H2O ____ NaOH + ____H2
13) ____ Ag2O ____ Ag + ____O2
14) ____ S8 + ____O2 ____ SO3
15) ____ CO2 + ____ H2O ____ C6H12O6 + ____O2
16) ____ K + ____ MgBr ____ KBr + ____ Mg
17) ____ HCl + ____ CaCO3 ____ CaCl2 + ____H2O + ____ CO2
70
Six Types of Chemical Reaction Worksheet
Balance the following reactions and indicate which of the six types of
chemical reaction are being represented:
1) ____ NaBr + ____ Ca(OH)2 ___ CaBr2 + ____ NaOH
Type of reaction: _____________________________
2) ____ NH3+ ____ H2SO4 ____ (NH4)2SO4
Type of reaction: _____________________________
3) ____ C5H9O + ____ O2 ____ CO2 + ____ H2O
Type of reaction: _____________________________
4) ____ Pb + ____ H3PO4 ____ H2 + ____ Pb3(PO4)2
Type of reaction: _____________________________
5) ____ Li3N + ____ NH4NO3 ___ LiNO3 + ___ (NH4)3N
Type of reaction: _____________________________
6) ____ HBr + ___ Al(OH)3 ___ H2O + ___ AlBr3
Type of reaction: _____________________________
7) What’s the main difference between a double displacement
reaction and an acid-base reaction?
8) Combustion reactions always result in the formation of water. What
other types of chemical reaction may result in the formation of
water? Write examples of these reactions on the opposite side of
this paper.
71
Word Equations Worksheet
Write the word equations for each of the following chemical reactions:
1) When dissolved beryllium chloride reacts with dissolved silver nitrate
in water, aqueous beryllium nitrate and silver chloride powder are
made.
2) When isopropanol (C3H8O) burns in oxygen, carbon dioxide, water,
and heat are produced.
3) When dissolved sodium hydroxide reacts with sulfuric acid (H2SO4),
aqueous sodium sulfate, water, and heat are formed.
4) When fluorine gas is put into contact with calcium metal at high
temperatures, calcium fluoride powder is created in an exothermic
reaction.
5) When sodium metal reacts with iron (II) chloride, iron metal and
sodium chloride are formed.
72
A Voyage through Equations
After working on this worksheet, you should be able to do the following:
1) Given an equation, you should be able to tell what kind of reaction
it is.
2) Predict the products of a reaction when given the reactants.
Section 1: Identify the type of reaction
For the following reactions, indicate whether the following are examples
of synthesis, decomposition, combustion, single displacement, double
displacement, or acid-base reactions:
1) Na3PO4 + 3 KOH 3 NaOH + K3PO4 _________________________
2) MgCl2 + Li2CO3 MgCO3 + 2 LiCl _________________________
3) C6H12 + 9 O2 6 CO2 + 6 H2O _________________________
4) Pb + FeSO4 PbSO4 + Fe _________________________
5) CaCO3 CaO + CO2 _________________________
6) P4 + 3 O2 2 P2O3 _________________________
7) 2 RbNO3 + BeF2 Be(NO3)2 + 2 RbF ________________________
8) 2 AgNO3 + Cu Cu(NO3)2 + 2 Ag ________________________
9) C3H6O + 4 O2 3 CO2 + 3 H2O _________________________
10) 2 C5H5 + Fe Fe(C5H5)2 _________________________
11) SeCl6 + O2 SeO2 + 3Cl2 _________________________
12) 2 MgI2 + Mn(SO3)2 2 MgSO3 + MnI4 _________________________
13) O3 O. + O2 _________________________
14) 2 NO2 2 O2 + N2_________________________
73
Section 2: Practicing equation balancing
Before you can write a balanced equation for a problem which asks you
to predict the products of a reaction, you need to know how to balance
an equation. Because some of you may not fully remember how to
balance an equation, here are some practice problems:
1) __ C6H6 + __ O2 __ H2O + __ CO2
2) __ NaI + __ Pb(SO4)2 __ PbI4 + __ Na2SO4
3) __ NH3 + __ O2 __ NO + __ H2O
4) __ Fe(OH)3 __ Fe2O3 + __ H2O
5) __ HNO3 + __ Mg(OH)2 __H2O + __ Mg(NO3)2
6) __ H3PO4 + __ NaBr __ HBr + __ Na3PO4
7) __ C + __ H2 __ C3H8
8) __ CaO + __ MnI4 __ MnO2 + __ CaI2
9) __ Fe2O3 + __ H2O __ Fe(OH)3
10) __ C2H2 + __ H2 __ C2H6
11) __ VF5 + __ HI __ V2I10 + __ HF
12) __ OsO4 + __ PtCl4 __ PtO2 + __ OsCl8
13) __ CF4 + __ Br2 __ CBr4 + __ F2
14) __ Hg2I2 + __ O2 __ Hg2O + __ I2
15) __ Y(NO3)2 + __ GaPO4 __ YPO4 + __ Ga(NO3)2
74
Section 3: Predicting the products of chemical reactions
Predict the products of the following reactions:
1) __ Ag + __CuSO4
Type:___________________________
2) __ NaI + __ CaCl2
Type:___________________________
3) __ O2 + __ H2
Type:___________________________
4) __ HNO3 + __ Mn(OH)2
Type:___________________________
5) __ AgNO2 + __ BaSO4
Type:___________________________
6) __ HCN + __ CuSO4
Type:___________________________
7) __ H2O + __ AgI
Type:___________________________
8) __ HNO3 + __Fe(OH)3
Type:___________________________
9) __ LiBr + __ Co(SO3)2
Type:___________________________
10) __ LiNO3 + __Ag
Type:___________________________
75
11) __ N2 + __ O2
Type:___________________________
12) __ H2CO3
Type:___________________________
13) __ AlCl3 + __ Cs
Type:___________________________
14) __ Al(NO3)3 + __ Ga
Type:___________________________
15) __ H2SO4 + __ NH4OH
Type:___________________________
16) __ CH3COOH + __ O2
Type:___________________________
17) __ C4H8 + __ O2
Type:___________________________
18) __ KCl + __ Mg(OH)2
Type:___________________________
19) __ Zn + __ Au(NO2)2
Type:___________________________
20) __ KOH + __ H2SO4
Type:___________________________
21) __ BaS + __ PtCl2
Type:___________________________
76
Hydrates Worksheet
1) How is a hydrate different from other chemical compounds?
2) Define the following terms:
anhydrate
dehydration
3) Name the following compounds:
a) FeCl3. 6 H2O ___________________________________________
b) CuSO4 . 5 H2O _________________________________________
4) Write the formulas for the following compounds:
a) barium chloride dihydrate _________________________________
b) magnesium sulfate heptahydrate ___________________________
5) What is the percent composition of water in the compound in
problem 4b?
6) If 125 grams of magnesium sulfate heptahydrate is completely
dehydrated, how many grams of anhydrous magnesium sulfate will
remain?
77
Molecular Formula Worksheet
Write the molecular formulas of the following compounds:
1) A compound with an empirical formula of C2OH4 and a
molar mass of 88 grams per mole.
2) A compound with an empirical formula of C4H4O and a
molar mass of 136 grams per mole.
3) A compound with an empirical formula of CFBrO and a
molar mass of 254.7 grams per mole.
4) A compound with an empirical formula of C2H8N and a
molar mass of 46 grams per mole.
5) Chemical analysis of a liquid shows that it is 60.0% C, 13.4%
H, 26.6% O by mass. Calculate the empirical formula of the
substance.
78
Percentage Composition Worksheet
Give the % composition of all elements in these compounds. Show all work!
________________________________________________________________
1) ammonium sulfite % N ______
% H ______
% S ______
% O ______
________________________________________________________________
2) aluminum acetate % Al ______
% C ______
% H ______
% O ______
________________________________________________________________
3) sodium bromide % Na ______
% Br ______
________________________________________________________________
4) copper (II) hydroxide % Cu ______
% O ______
% H ______
________________________________________________________________
5) magnesium carbonate % Mg ______
% C ______
% O ______
79
________________________________________________________________
6) iron (II) phosphate % Fe ______
% P ______
% O ______
________________________________________________________________
7) beryllium nitride % Be ______
% N ______
________________________________________________________________
8) potassium cyanide % K ______
% C ______
% N ______
_____________________________________________________________
9) manganese (III) nitrate % Mn ______
% N ______
% O ______
________________________________________________________________
10) lithium phosphide % Li ______
% P ______
________________________________________________________________
11) nickel (III) sulfate % Ni ______
% S ______
% O ______
80
Percent Composition and Molecular Formula Worksheet
1) What’s the empirical formula of a molecule containing
65.5% carbon, 5.5% hydrogen, and 29.0% oxygen?
2) If the molar mass of the compound in problem 1 is 110
grams/mole, what’s the molecular formula?
_____________________________________________________________
3) What’s the empirical formula of a molecule containing
18.7% lithium, 16.3% carbon, and 65.0% oxygen?
4) If the molar mass of the compound in problem 3 is 73.8
grams/mole, what’s the molecular formula?
81
Stoichiometry: Mole-Mole Calculations
You will be asked to calculate the moles of a product produced (or the
moles of a reactant needed), given a known quantity of moles to start
with. This is simply done by setting up a proportion and solving for X.
For example, given the equation:
4 NH3 + 3 O2 -----> 2 N2 + 6 H2O
4 moles of NH3 are combined with 3 moles of O2 to make
2 moles of N2 and 6 moles of H2O.
This means that the ratio can be rewritten as: 4:3:2:6.
Knowing this ratio, allows you to solve for any unknown as long as you are
given one quantity.
Example: Suppose I told you that you had exactly 2.500 moles of NH3.
How would you find the moles of oxygen needed and the moles of
nitrogen and water produced?
4 NH3 + 3 O2 -----> 2 N2 + 6 H2O
2.500 moles ____moles ____moles ____moles
Set up theses proportions to find the answers as follows:
4 moles of NH3 = 3 moles of O2 4 X = 7.5 X = 1.875
2.500 moles of NH3 X moles of O2
4 moles of NH3 = 2 moles of N2 4 X = 5.0 X = 1.250
2.500 moles of NH3 X moles of N2
4 moles of NH3 = 6 moles of H2O 4 X = 15 X = 3.750
2.500 moles of NH3 X moles of H2O
4 NH3 + 3 O2 -----> 2 N2 + 6 H2O
2.500 moles 1.875 moles 1.250 moles 3.750 moles
82
Given the following equations, find the moles of all the
unknown products and reactants, from the moles of one of the
compounds given.
1. 2 Fe + 3 S ----> Fe2S3 Given: 3 moles of Iron
2. 2 NaBr + Cl2 -------> 2 NaCl + Br2 Given: 0.5 moles of NaBr
3. Cu2O + H2 --------> 2 Cu + H2O Given: 2.75 moles of H2
4. Mg(OH)2 + 2HNO3 -------> Mg(NO3)2 + 2 H2O Given: 5.25 moles of
Mg(NO3)2
5. 2 C2H6 + 7 O2 ------> 4 CO2 + 6 H2O Given: 3 moles of H2O
6. 2 Sb +3 Cl2 -----> 2 SbCl3 Given: 4 moles of Sb
7. Na2SO3 + H2SO4 -----> Na2SO4 + H2SO3 Given: 5 moles of H2SO4
8. CaCl2 ----> Ca + Cl2 Given: 1 mole of Ca
9. 2 H2O ----> 2 H2 + O2 Given: 2 moles of H2
10. 2 C4H10 + 13 O2 ------> 8 CO2 +10 H2O Given: 0.75 moles of C4H10
83
Stoichiometry Mole-Mass Calculations
You will be asked to calculate the mass of a product produced (or the
mass of a reactant needed), given a known quantity of moles to start
with. This is done in two steps:
Step 1: Set up a proportion to find moles.
Moles of A -----(use proportion)-----> Moles of B
For example, given the equation:
4 NH3 + 3 O2 -----> 2 N2 + 6 H2O
If you were given 2.50 moles of NH3, how would you find the mass of
water produced?
4 NH3 + 3 O2 -----> 2 N2 + 6 H2O
2.50 moles ______moles
4 moles of NH3 = 6 moles of H2O 4 X = 15
2.50 moles of NH3 X X = 3.75 moles of H2O
Step 2: To convert Moles to Mass multiply by the molar mass.
Moles of B -----(x Molar Mass)-----> Mass of B
H2O has a molar mass of 18 (Oxygen = 16, two Hydrogens = 2)
3.75 moles of H2O x 18 grams = 67.5 grams
mole
84
Given the following equations, find the mass of the unknown
product or reactant from the moles of one of the compounds
given.
1. 2 Al + 3 S ----> Al2S3 Given: 3 moles of Al
Find: Mass of Al2S3
2. 2 NaI + Cl2 -------> 2 NaCl + I2 Given: 0.5 moles of NaI
Find: Mass of I2
3. Na2O + H2 --------> 2 Na + H2O Given: 2.75 moles of H2
Find: Mass of Na
5. 2 C2H6 + 7 O2 ------> 4 CO2 + 6 H2O Given: 3 moles of H2O
Find: Mass of O2
6. 2 Sb +3 Cl2 -----> 2 SbCl3 Given: 4 moles of Sb
Find: Mass of SbCl3
7. Na2SO3 + H2SO4 -----> Na2SO4 + H2SO3
Given: 5 moles of H2SO4
Find: Mass of Na2SO4
8. CaCl2 ----> Ca + Cl2 Given: 1 mole of Ca
Find: Mass of CaCl2
9. 2 H2O ----> 2 H2 + O2 Given: 2 moles of H2
Find: Mass of H2O
10. 2 C4H10 + 13 O2 ------> 8 CO2 +10 H2O Given: 0.75 moles of C4H10
Find: Mass of CO2
85
Limiting Reagent Worksheet
Using your knowledge of stoichiometry and limiting reagents, answer the following
questions:
1) Write the balanced equation for the reaction of lead (II) nitrate with sodium iodide
to form sodium nitrate and lead (II) iodide:
2) If I start with 25.0 grams of lead (II) nitrate and 15.0 grams of sodium iodide, how
many grams of sodium nitrate can be formed?
3) What is the limiting reagent in the reaction described in problem 2?
4) How much of the nonlimiting reagent will be left over from the reaction in
problem #2?
86
Chapter 7
Gas Laws
87
Kinetic Theory of Gases
On the notes you’ve taken, in your own words, answer the
following questions:
1. On what two factors is the kinetic theory based?
1.
2.
2. What is an “ideal gas”?
3. State the 5 assumptions of the kinetic theory of gases.
1.
2.
3.
4.
5.
4. What are the 5 physical properties of a gas?
1.
2.
3.
4.
5.
5. What is diffusion?
6. What are the 4 types of measurements that can be used to describe
the condition of a gas?
1.
2.
3.
4.
88
Description of Measurements for Gases
Temperature Units Conversion
oC (degrees Celsius) oC + 273 = K
K (Kelvin) K - 273 = oC
Volume Units Conversion
L (liters) 1 L = 1000 mL ( so x1000 to convert to mL)
mL (milliliters) 1000 mL = 1 L ( so /1000 to convert to L)
cm3 (cubic centimeters) 1 mL = 1 cm3
Pressure Units Conversion
mm Hg (millimeters of mercury) 760 mm Hg = 1 atm
( so /760 to convert to atm)
atm (atmospheres) 1 atm = 760 mm Hg
( so x760 to convert to mm Hg)
torr (same as mm Hg) 760 torr = 760 mm Hg
Pa (pascals) 1.013 x105 Pa = 1 atm
kPa (kilopascals) 101.3 kPa = 1 atm
Amount Units Conversion
mol (moles) 1 mole = 6.02 x 1023 particles
89
Gas Measurements
Convert the following measurements using your gas conversion
chart.
1. Convert the following temperatures from degrees Celsius to Kelvin:
a. 0 oC = _________
b. 27 oC = _________
c. -50 oC = _________
d. -273 oC = _________
2. Convert the following temperatures from Kelvin to degrees Celsius:
a. 273 K = _________
b. 350 K = _________
c. 100 K = _________
d. 0K = _________
3. Convert the following volumes from liters to milliliters:
a. 1 L = _________
b. 2.50 L = _________
c. 0.350 L = _________
d. 0.010 L = _________
5. Convert the following volumes from milliliters to liters:
a. 1000 mL = _________
b. 750 mL = _________
c. 45 mL = _________
d. 6850 mL = _________
90
Boyles’ Law
Use Boyles’ Law to answer the following questions:
1) 1.00 L of a gas at standard temperature and pressure is compressed
to 473 mL. What is the new pressure of the gas?
2) In a thermonuclear device, the pressure of 0.050 liters of gas within
the bomb casing reaches 4.0 x 106 atm. When the bomb casing is
destroyed by the explosion, the gas is released into the atmosphere
where it reaches a pressure of 1.00 atm. What is the volume of the
gas after the explosion?
3) Synthetic diamonds can be manufactured at pressures of 6.00 x 104
atm. If we took 2.00 liters of gas at 1.00 atm and compressed it to a
pressure of 6.00 x 104 atm, what would the volume of that gas be?
4) The highest pressure ever produced in a laboratory setting was
about 2.0 x 106 atm. If we have a 1.0 x 10-5 liter sample of a gas at
that pressure, then release the pressure until it is equal to 0.275 atm,
what would the new volume of that gas be?
91
5) Atmospheric pressure on the peak of Mt. Everest can be as low as
150 mm Hg, which is why climbers need to bring oxygen tanks for
the last part of the climb. If the climbers carry 10.0 liter tanks with an
internal gas pressure of 3.04 x 104 mm Hg, what will be the volume of
the gas when it is released from the tanks?
6) Part of the reason that conventional explosives cause so much
damage is that their detonation produces a strong shock wave that
can knock things down. While using explosives to knock down a
building, the shock wave can be so strong that 12 liters of gas will
reach a pressure of 3.8 x 104 mm Hg. When the shock wave passes
and the gas returns to a pressure of 760 mm Hg, what will the
volume of that gas be?
7) Submarines need to be extremely strong to withstand the extremely
high pressure of water pushing down on them. An experimental
research submarine with a volume of 15,000 liters has an internal
pressure of 1.2 atm. If the pressure of the ocean breaks the
submarine forming a bubble with a pressure of 250 atm pushing on
it, how big will that bubble be?
8) Divers get “the bends” if they come up too fast because gas in their
blood expands, forming bubbles in their blood. If a diver has 0.05 L
of gas in his blood under a pressure of 250 atm, then rises
instantaneously to a depth where his blood has a pressure of 50.0
atm, what will the volume of gas in his blood be? Do you think this
will harm the diver?
92
Charles’ Law Worksheet
1) The temperature inside my refrigerator is about 40 Celsius. If I place
a balloon in my fridge that initially has a temperature of 220 C and a
volume of 0.5 liters, what will be the volume of the balloon when it is
fully cooled by my refrigerator?
2) A man heats a balloon in the oven. If the balloon initially has a
volume of 0.4 liters and a temperature of 20 0C, what will the
volume of the balloon be after he heats it to a temperature of 250
0C?
3) On hot days, you may have noticed that potato chip bags seem to
“inflate”, even though they have not been opened. If I have a 250
mL bag at a temperature of 19 0C, and I leave it in my car which
has a temperature of 600 C, what will the new volume of the bag
be?
4) A soda bottle is flexible enough that the volume of the bottle can
change even without opening it. If you have an empty soda bottle
(volume of 2 L) at room temperature (25 0C), what will the new
volume be if you put it in your freezer (-4 0C)?
93
5) Some students believe that teachers are full of hot air. If I inhale 2.2
liters of gas at a temperature of 180 C and it heats to a temperature
of 380 C in my lungs, what is the new volume of the gas?
6) How hot will a 2.3 L balloon have to get to expand to a volume of
400 L? Assume that the initial temperature of the balloon is 25 0C.
7) I have made a thermometer which measures temperature by the
compressing and expanding of gas in a piston. I have measured
that at 1000 C the volume of the piston is 20 L. What is the
temperature outside if the piston has a volume of 15 L? What would
be appropriate clothing for the weather?
94
Gay – Lussac Worksheet
1. Convert 25.0 C at 63.0 atm to its new temperature at standard
pressure.
2. A gas has a pressure of 0.370 atm and 50.0C What is the pressure at
standard temperature?
3. A gas has a pressure of 699.0 mmHg at 40.0C. What is the temperature
at standard pressure?
4. If a gas is cooled from 323.0 K to 273.15 K and the volume is constant,
what final pressure would result if the original pressure was 750.0 mmHg
?
5. If a gas in a closed container, with an original temperature 0f 25.0C, is
pressureized from 15.0 atmospheres to 16.0 atmospheres, what would
the final temperature of the gas be?
95
Combined Gas Law Problems
Use the combined gas law to solve the following problems:
1) If I initially have a gas at a pressure of 12 atm, a volume of 23 liters,
and a temperature of 200 K, and then I raise the pressure to 14 atm
and increase the temperature to 300 K, what is the new volume of
the gas?
2) A gas takes up a volume of 17 liters, has a pressure of 2.3 atm, and
a temperature of 299 K. If I raise the temperature to 350 K and
lower the pressure to 1.5 atm, what is the new volume of the gas?
3) A gas that has a volume of 28 liters, a temperature of 45 0C, and an
unknown pressure has its volume increased to 34 liters and its
temperature decreased to 35 0C. If I measure the pressure after the
change to be 2.0 atm, what was the original pressure of the gas?
4) A gas has a temperature of 14 0C, and a volume of 4.5 liters. If the
temperature is raised to 29 0C and the pressure is not changed,
what is the new volume of the gas?
96
5) If I have 17 liters of gas at a temperature of 67 0C and a pressure of
88.89 atm, what will be the pressure of the gas if I raise the
temperature to 94 0C and decrease the volume to 12 liters?
6) I have an unknown volume of gas at a pressure of 0.5 atm and a
temperature of 325 K. If I raise the pressure to 1.2 atm, decrease the
temperature to 320 K, and measure the final volume to be 48 liters,
what was the initial volume of the gas?
7) If I have 21 liters of gas held at a pressure of 78 atm and a
temperature of 900 K, what will be the volume of the gas if I
decrease the pressure to 45 atm and decrease the temperature to
750 K?
8) If I have 2.9 L of gas at a pressure of 5 atm and a temperature of 50
0C, what will be the temperature of the gas if I decrease the volume
of the gas to 2.4 L and decrease the pressure to 3 atm?
9) I have an unknown volume of gas held at a temperature of 115 K in
a container with a pressure of 60 atm. If by increasing the
temperature to 225 K and decreasing the pressure to 30 atm causes
the volume of the gas to be 29 liters, how many liters of gas did I
start with?
97
The Ideal and Combined Gas Laws
Use your knowledge of the ideal and combined gas laws to
solve the following problems. Hint: Figuring out which equation
you need to use is the hard part!
1) If four moles of a gas at a pressure of 5.4 atmospheres
have a volume of 120 liters, what is the temperature?
2) If I initially have a gas with a pressure of 84 kPa and a
temperature of 350 C and I heat it an additional 230
degrees, what will the new pressure be? Assume the
volume of the container is constant.
3) My car has an internal volume of 2600 liters. If the sun
heats my car from a temperature of 200 C to a
temperature of 550 C, what will the pressure inside my car
be? Assume the pressure was initially 760 mm Hg.
4) How many moles of gas are in my car in problem #3?
5) A toy balloon filled with air has an internal pressure of 1.25
atm and a volume of 2.50 L. If I take the balloon to the
bottom of the ocean where the pressure is 95
atmospheres, what will the new volume of the balloon
be? How many moles of gas does the balloon hold?
(Assume T = 285 K)
98
Molar Volume, Density, and Volume-Volume Problems
1. The density of a gas is 2.0 grams/liter at STP. What is its molar mass?
(Remember your units!)
2. Given the reaction: 2 C2H6 + 7 O2 4 CO2 + 6 H2O, at STP, what is
the total volume of CO2 formed when 6.0 liters of C2H6 are completely
consumed?
(Remember your units!)
3. At STP, the volume occupied by 32 grams of a gas is 11.2 liters. What is
the molar mass of this gas? (Remember your units!)
4. How many liters of (dry) helium gas would there be in a filled
container at STP, if the container held exactly 2 moles of helium gas?
(Remember your units!)
5. What is the volume of 0.500 moles of neon at STP?(Remember your
units!)
99
6. What is the density of the F2 gas at STP? (Remember your units!)
7. How many moles of Argon gas would occupy 67.2 liters at STP?
(Remember your units!)
8. The gram molecular mass (molar mass) of a gas is 56 grams/mole.
What is its density at STP? (Remember your units!)
9. A 15 gram sample of a gas has a volume of 30. liters at STP. What is the
density of this gas at STP? (Remember your units!)
10. At STP, 32 grams of O2 would occupy the same volume as how many
grams of helium? (Remember your units!)
100
Chapter 9
Solutions
101
Solutions Vocabulary
1. Define these words
Soluble -
Solution -
Solvent -
Solute -
Suspension -
Colloid –
Tyndall Effect -
Electrolyte -
Non-electrolyte -
2. Define these words:
Solution equilibrium -
Saturated solution -
Unsaturated solution-
102
Supersaturated solution-
Solubility -
Miscible -
Immiscible -
Hydration -
Solvated –
Heat of solution -
3. What are the 3 factors that affect the RATE of dissolving?
4. What is Le Chateleir’s principle? – Explain
5. What is Henry’s law?
103
Solubility Curve Table G
1. List three compounds that are gases in table G.
a. __________________, _________________, _____________________
2. What do these three gaseous compounds have in common?
3. For solids listed on this table, what is the relationship between solubility
and temperature?
4. What substance on Table G is least soluble at 50C ?
5. What substance is most soluble in Table G ?
6. How much potassium nitrate will dissolve in 100 grams of water at 4C ?
7. How much ammonia gas will dissolve in 200 grams of water at 90C ?
8. How much sodium nitrate will dissolve in 50 of water at 45C ?
9. 100 grams of solution of ammonium chloride contains 60 grams of
solute at 70 C. Is the solution saturated, unsaturated, or
supersaturated?
10. 100 grams of solution of sodium nitrate contains 120 grams of solute at
50C. Is the solution saturated, unsaturated, or supersaturated?
11. 100 grams of solution of sodium chloride contains 40 grams of solute at
95C. Is the solution saturated, unsaturated, or supersaturated?
12. At what temperature is potassium chlorate and sodium chloride
equally soluble?
104
13. 100 grams of sodium nitrate solution at 30C contains 90 grams of
solute. What mass of solute must be added to the solution to saturate
the solution?
14. 100 grams of saturated solution ammonium chloride solution at 85C,
when the temperature drops to 45C, how much solute will leave to
precipitate?
15. 100 grams of solution potassium nitrate contains 105 grams of solute,
what temperature must be obtained to saturate the solution?
16. 100 grams of saturated solution potassium chloride solution at 145C,
when the temperature drops to 45C, how much solute will leave to
precipitate?
17. 100 grams of solution sodium nitrate contains 115 grams of solute, what
temperature must be obtained to saturate the solution?
18. How much solute must be added to 100 mL of water with 50 grams of
ammonium chloride at 80C to saturate the solution?
19. How much solute must be added to 200 mL of water with 100 grams of
sodium nitrate at 70C to saturate the solution?
20. If you were to see a closed system with 100 mL of clear colorless
aqueous solution of ammonium chloride with 20.0 g of solid sitting on
the bottom of the container at 60C.
a. How much ammonium chloride is in the solution?
b. What is the total mass of ammonium chloride in the container?
105
Table F – Solubility Guidelines
Answer the following questions by using Table F Solubility Curve on the
Reference Tables for the Physical Setting/Chemistry.
If the compounds formed from the ions on this table are soluble indicate
by placing an “S” in the appropriate square.
If the compounds formed from the ions on this table are insoluble indicate
by placing an “I” in the appropriate square.
PO4 -3 CO3 -2 SO4 -2 Cl -1 OH -1 NO3 -1
ANIONS
CATIONS
Li +1 Li3PO4 Li2CO3 Li2SO4 LiCl LiOH LiNO3
Ag +1
Ca +2 Ca3(PO4)2 CaCO3 CaSO4 CaCl2 Ca(OH)2 Ca(NO3)2
Pb +2
Sr +2
NH4 +1 (NH4)3PO4 (NH4)2SO4 NH4OH
106
Molarity Practice Problems
1) How many grams of potassium carbonate are needed to make 200
mL of a 2.5 M solution?
2) How many liters of 4 M solution can be made using 100 grams of
lithium bromide?
3) What is the concentration of a 450 mL solution that contains 200
grams of iron (II) chloride?
4) How many grams of ammonium sulfate are needed to make a 0.25
L solution at a concentration of 6 M?
5) What is the concentration of a solution that has a volume of 2.5 L
and contains 660 grams of calcium phosphate?
6) How many grams of copper (II) fluoride are needed to make 6.7
liters of a 1.2 M solution?
107
7) How many liters of 0.88 M solution can be made with 25.5 grams of
lithium fluoride?
8) What is the concentration of a solution that with a volume of 660
that contains 33.4 grams of aluminum acetate?
9) How many liters of 0.75 M solution can be made using 75 grams of
lead (II) oxide?
10) How many grams of manganese (IV) oxide are needed to make a
5.6 liters of a 2.1 M solution?
11) What is the concentration of a solution with a volume of 9 mL that
contains 2 grams of iron (III) hydroxide?
12) How many liters of 3.4 M solution can be made using 78 grams of
isopropanol (C3H8O)?
13) What is the concentration of a solution with a volume of 3.3 mL that
contains 12 grams of ammonium sulfite?
108
Dilution Problems
Molarity x Volume = Molarity x Volume
M1V1 = M2V2
1. How much water would you add to 200 mL 5.0 HCl solution to change
the concentration of a solution to 2.5 M?
2. How much water would you add to 600 mL 0.3 NaCl solution to
change the concentration of a solution to 0.1 M?
3. 400 mL of water are added to 200 mL 0.25M KCl solution. Calculate the
concentration of the new solution. (Hint: the volume of the new
solution will be 600 mL.)
4. 2.0 L of water are added to 6.0 L 0.01 M HI solution. Calculate the
concentration of the new solution.
5. 20 mL was added to 80 mL of 16M H2SO4. Calculate the concentration
of the new solution.
6. How much water must be added to 250 mL of a 2.0 M H3PO4 to
change the concentration to 0.5 M?
7. How much water must be added to 750 mL of a 2.5M H3PO4 to change
the concentration to 2.0 M?
109
Parts Per Million (PPM)
Parts per million is used to report the concentration of a dilute solution.
Parts per million is similar to percentage composition because it compares
masses. Parts per million (ppm) is a ratio between the mass of a solute and
the total mass of the solution.
1. Approximately 0.0043 grams of oxygen can be dissolved in 100 mL of
water at 20.0 C. Express this in terms of parts per million.
2. Approximately 5.0 grams of sulfur dioxide (SO2) at 25.0 C will dissolve in
100 mL of water. Express in terms of ppm.
3. An unknown gas was found to dissolve 0.00089 grams in 100 mL of
water. Calculate the concentration in parts per million.
4. A gas dissolved in water to create 1007.1 grams of solution with a
concentration of 35 ppm. How many grams of the gas were dissolved
in the aqueous solution.
110
Chapter 9
Kinetics/Equilibrium
111
Kinetics
1. Define the following words:
Exothermic -
Endothermic -
Kinetics -
Enthalpy-
Entropy -
Heat of Rxn -
Heat of Formation -
Activation Energy -
Catalyst -
Principles of Kinetic Molecular Theory
1.
2.
3.
Five Factors that Affect the Rate of Reaction
1.
2.
3.
4.
5.
112
Collision Theory Worksheet
1) Explain why all reactions have an activation energy, using your
knowledge of collision theory.
2) Describe how the activation energy of a reaction affects the overall
rate of the chemical reaction.
3) A rule of thumb used by organic chemists is that the rate of a
chemical reaction can be doubled by increasing the reaction
temperature by ten degrees Celsius. Explain this drastic increase in
reaction rate using your knowledge of collision theory.
4) It has been observed that more gas station fires occur on hot days
than on cold days. Explain this phenomenon using your knowledge
of collision theory. (Hint: It’s not just the temperature increase that
causes this!)
5) It has been observed with one variety of paint that the rate of paint
drying can be drastically increased by adding a small amount of
“accelerant”. Based on what you know of catalysts, is it reasonable
to think of this accelerant as being a catalyst? Explain.
113
Potential Energy Diagrams and Worksheet
An Exothermic Reaction
P
O
T
E - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - -- - - - - - - - - - - -
N
T #3
I
A
L #4
Reactants
E - -- - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - -
N
E #5
R Products
G - - - - - - - - - - - -- - - - - - - - - - - -- - - - - - - - - - - -- - - - -
Y
#1 #6 #2
Reaction ----------------------------------------->
Arrow #1 - is the potential energy of the reactants
Arrow #2 - is the potential energy of the products
Arrow #3 - is the forward activation energy
Arrow #4 - is the reverse activation energy
Arrow #5 - is the change in energy or enthalpy or Heat of reaction or rxn
(Note that rxn is always products – reactants)
Arrow #6 - is the activated complex
114
FOR EACH OF THE FOLLOWING, USE THE DIAGRAM ON THE PREVIOUS PAGE
TO SHOW HOW YOU CALCULATED YOUR ANSWERS.
1. If the energy of the reactants is 25 kcal/mole, and the energy of the
products is 10 kcal/mole, what is the heat of this reaction?
2. From the information in question #1, and given that the activation
energy of the forward reaction is 45 kcal/mole, what is the activation
energy of the reverse reaction?
3. From the information in question #1 and #2, what is the activated
complex for this reaction?
115
Potential Energy Diagram - 1
116
Potential Energy Diagram - 2
117
Constructing a Potential Energy Diagram
118
Endothermic or Exothermic (Table H)
Label the following reactions as Endothermic or Exothermic:
C (s) + O2 (g) ----> CO2 (g) + heat
heat + N2 (g)+ O2 (g) -----> 2NO (g)
4.3 kcal + I2 (g) + Cl2 (g) -----> 2 ICl (g)
2 K (s) + Cl2 (g) ----> 2 KCl (g) + 98.3 kcal
Each of the following reactions shows an increase in entropy, explain why.
CO2 (s) ----> CO2 (g)
NaCl (s) ------> Na+ (aq) + Cl- (aq)
2 KClO3 (s) ----> 2 KCl (s) + 3 O2 (g)
H2O (g) at 110 oC -------> H2O (g) at 140 oC
119
Le Châtlier’s Principle
Explain how the following changes in reaction conditions will affect the
position of the equilibrium below, and explain your reasoning.
A(g) + B(aq) C(s) ΔHrxn= -453 kJ/mol
1) The pressure of A in the reaction chamber is increased.
2) The temperature of the reaction is increased by 200 C.
3) A catalyst is added to the system.
4) As the reaction progresses, more of compound B is steadily added
to the reaction chamber.
5) An inhibitor is added to the reaction chamber.
6) Argon gas is added to the reaction chamber, doubling the
pressure.
120
Review of Equilibrium
1. Write out the equilibrium expression (Keq) for the following reaction:
N2 (g) + 3 H2 (g) 2 NH3 (g)
2. There are four ways to alter the equilibrium of a reaction (shift the
reaction to the right or the left). List the four ways and explain how they
shift the equilibrium.
3. What does it mean when Keq > 1 ? What does it mean when Keq H+12O-2 + K+1Cl-1 + Mn+2Cl-12 + Cl02
Step 2) Find out which element is being oxidized and which element is
being reduced; then start to write out the 1/2 rxns for each.
Oxidation: Cl-1 Cl02
Reduction: Mn+7 Mn+2
Step 3) Then Balance both the Number of atoms and the Charge...
Oxidation: 2 Cl-1 Cl02 + 2 e-1
Reduction: 5 e-1 + Mn+7 Mn+2
Step 4) Balance out the electrons so that the number of electrons gained
= the number lost.
Oxidation: (2 Cl-1 Cl02 + 2 e-1) x 5 = 10 Cl-1 ----> 5 Cl02 + 10 e-1
Reduction: (5 e-1 + Mn+7 Mn+2) x 2 = 10e-1 + 2 Mn+7 ---> 2 Mn+
137
Step 5) Now add the two reactions together like a math problem,
canceling out the electrons.
Oxidation: 10 Cl-1 5 Cl02 + 10 e-1
Reduction: 10e-1 + 2 Mn+7 2 Mn+2
-----------------------------------------------------------
10 Cl-1 + 2 Mn+7 2 Mn+2 + 5 Cl02
Step 6) Reinsert the coefficients into the original equation.
H+1Cl-1 + K+1Mn+7O-24 H+12O-2 + K+1Cl-1 + Mn+2Cl-12 + Cl02
becomes.....
10 HCl + 2 KMnO4 H2O + KCl + 2 MnCl2 + 5 Cl2
Step 7) Now check to make sure everything is balanced, and if not, the try
to rebalance it.
Since the K’s were not balanced, The KCl is multiplied by 2.
10 HCl + 2 KMnO4 H2O + 2 KCl + 2 MnCl2 + 5 Cl2
Now to balance the Cl’s (since there are 16 on the right, we need
16 on the left.)
16 HCl + 2 KMnO4 H2O + 2 KCl + 2 MnCl2 + 5 Cl2
Finally, to balnce the H’s (and also the O’s) 8 H2O’s are needed.
16 HCl + 2 KMnO4 8 H2O + 2 KCl + 2 MnCl2 + 5 Cl2
138
Redox Worksheet
1. What is oxidation? Give an example of this process.
2. What is reduction? Give an example of this process.
3. What is meant by a “redox reaction”?
4. In the following reactions, identify which elements are being oxidized,
and which are being reduced. Show the oxidation numbers of the
oxidized and reduced elements:
a. 2 KNO3 (s) ----> 2 KNO2 (s) + O2 (g)
b. H2 (g) + CuO (s) -----> Cu (s) + H2O (l)
5. For the equations above, write the electronic equations for the
elements oxidized and reduced in each.
139
Oxidation State
In each of the following chemicals, determine the oxidation states of
each element:
1) sodium nitrate ____________________________________
2) ammonia ____________________________________
3) zinc oxide ____________________________________
4) water ____________________________________
5) calcium hydride ____________________________________
6) carbon dioxide ____________________________________
7) nitrogen ____________________________________
8) sodium sulfate ____________________________________
9) aluminum hydroxide ____________________________________
10) magnesium phosphate ____________________________________
In each of the following reactions, determine what was oxidized and what
was reduced.
11) Ca + H2O CaO + H2
Element oxidized: ____________________________________
Element reduced: ____________________________________
12) 2 H2 + O2 2 H2O
Element oxidized: ____________________________________
Element reduced: ____________________________________
140
Oxidation and Reduction Practice
In each of the following equations, indicate the element that has been
oxidized and the one that has been reduced. You should also label the
oxidation state of each before and after the process:
1) 2 Na + FeCl2 2 NaCl + Fe
2) 2 C2H2 + 5 O2 4 CO2 + 2 H2O
3) 2 PbS + 3 O2 2 SO2 + 2 PbO
4) 2 H2 + O2 2 H2O
5) Cu + HNO3 CuNO3 + H2
6) AgNO3 + Cu CuNO3 + Ag
141
Balancing Redox Reactions (1/2 RXN method)
Balance the following equations:
1. Al + Fe+2 Al+3 + Fe
2. Cl2 + Br- Br2 + Cl-
3. Al + Sn+4 Al+3 + Sn
4. Na + H2O Na+ + OH- + H2
5. Fe + Cu2+ Fe2+ + Cu
6. Zn + 2HCl → ZnCl2 + H2
7. 2Al + 3Cu2+ → 2Al3+ + 3Cu
8. HCl + KMnO4 KCl + MnCl2 + H2O + Cl2
142
ELECTROCHEMICAL CELL
A device that produces usuable
electric energyfrom a spontaneous
chemical rxn; a battery.
Zn0(s) Zn2+ + 2e- Cu2+ + 2e- Cu0(s)
oxidation reduction
Anode Cathode
- is negative - is positive
- oxidation occurs - reduction occurs
- e- flow from anode - e- flow to the cathode
- e- generated by the - e- accumulated and used
oxidation 1/2 rxn to make reduction happen
143
ELECTROLYTIC CELL
A device that drives a non-
spontaneous chemical rxn by using
an external electrical source. (It uses
a battery)
2 Cl- Cl2(g) + 2e- 2 Na+ +2e- 2 Na0(s)
Anode Cathode
- is positive - is negative
- oxidation occurs at anode - reduction occurs at cathode
- neg. ions attracted to anode; - positive ions attracted to
- electrons produced here. cathode, battery supplies
electrons to cathode.
144
ELECTROPLATING
Electric current used to deposit a layer of of
metal, such as silver, on the object to be plated.
Same system as the electrolytic cell.
Ag Ag+ + 1e- Ag+ + 1e- Ag
Anode Cathode
- is positive - is negative
- oxidation occurs at anode - reduction occurs at cathode
- positive ions provided - positive ions attracted to
by anode cathode
- electroplating occurs at
cathode
145
Electrochemistry
1. Write the word anode or cathode next to each line.
a. The electrode at which oxidation occurs. _________________
b.The electrode at which reduction occurs. _________________
c.The negative electrode in an electrochemical cell. _________________
d.The negative electrode in an electrolytic cell. _________________
e.In an electrolytic cell, the electrode that attracts Cl - ions . _____________
f.The part of an electroplating system that is provided by a fork. _________
g.The part of an electroplating system that generates Ag+ ions. _________
2. Define the term reducing agent.
3. Define the term oxidizing agent.
4. The most active reducing agent among the elements is:
a. iodine b. cesium c. fluorine d. lithium
Explain why:
____________________________________________________________
5. The most active oxidizing agent among the elements is:
a. iodine b. cesium c. fluorine d. lithium
Explain why:
____________________________________________________________
6. How do anodes and cathodes differ in electrochemical cells vs.
electrolytic cells?
146
7. Write out the net ionic (or skeletal redox) equation for an
electrochemical cell that has
MgCl2 in one 1/2 cell, and SnCl2 in the other 1/2 cell.
8. Draw the full electrochemical cell diagram for question #7 on this
sheet.
9.
10.
147
Electrochemical Cell
You are given two beakers. In Beaker #1 is a solution of lead
nitrate ( Pb(NO3)2 ) and a piece of solid lead metal. In Beaker
#2 is a solution of nickel nitrate ( Ni(NO3)2 ) and a piece of solid
nickel metal. The two pieces of metal are connected by wires
to each other through a voltmeter which reads electrode
potential (Ecell) in volts. Finally, a salt bridge is placed between
the two beakers. The salt bridge contains a solution of
potassium nitrate.
1. Draw a diagram of the picture mentioned above.
2. Which beaker ( half-cell) is undergoing oxidation and which
is undergoing reduction?
3. Write the 1/2 reactions for oxidation and reduction of these
two half-cells.
4. What does the Ecell equal for this reaction?
5. Which electrode is the anode and which is the cathode?
6. Which electrode is positive and which is negative?
7. In which direction are the electrons flowing (from what to
what)?
8. Why will this reaction occur spontaneously?
9. What is the function of the salt bridge?
10. What will happen to the Nickel metal over time, and what
will happen to the Lead metal over time?
11. Which of the two solutions will have more metal ions in it
over time and which will have less over time?
148
Chapter 11
Organic Chemistry
149
Organic Functional Groups
Identify the functional groups in each of the following organic
compounds:
1)
2)
3)
4)
5)
150
Organic Chemistry
The Alkanes
These are Single-bonded hydrocarbons.
IUPAC name Molecular Formula Generic Formula
methane CH4
ethane C2H6
propane C3H8
butane C4H10
pentane C5H12
hexane C6H14 CnH2n+2
heptane C7H16
octane C8H18
nonane C9H20
decane C10H22
The Alkenes
These are Double-bonded hydrocarbons that have the same prefix as
the alkanes but have the suffix ending -ene.
The generic formula is CnH2n
Examples: ethene, propene, butene...
Propene: C3H6 CH2=CH-CH3
The Alkynes
These are Triple-bonded hydrocarbons that have the same prefix as the
alkanes but have the
suffix ending -yne. The generic formula is CnH2n-2
Examples: ethyne, propyne, butyne...
Propyne: C3H4 CH = C-CH3
151
The Benzene Series
These are Aromatic hydrocarbons that have rings and double bonds that
show resonance.
The generic formula is CnH2n-6
Examples: benzene, toluene...
Benzene: C6H6
Toluene: C7H8
CH3
The Haloalkanes
These are hydrocarbons with a halogen attached to the chain instead of
a hydrogen. The prefix begins with the halogen name.
Examples: iodomethane, bromoethane, 2-fluoropropane, 1-
chlorobutane
Bromoethane: C2H5Br CH3-CH2-Br
152
The Alcohols
These are hydrocarbons with an hydroxyl functional group (-OH) attached
to the chain instead of a hydrogen. The suffix ends with -ol.
Examples: methanol, ethanol. ethenol. 1,2-ethanediol
Ethanol: C2H5OH CH3-CH2-OH
The Aldehydes
These are hydrocarbons with a carbonyl functional group (=O) attached
to the end of a chain. The double-bonded oxygen takes the place of two
of the hydrogens. The suffix ends with -al.
Examples: methanal, ethanal. propanal
O
Ethanal: C2H4O CH3-CH
The Ketones
These are hydrocarbons with a carbonyl functional group (=O) attached
to the middle of a chain. The double-bonded oxygen takes the place of
two of the hydrogens. The suffix ends with -one.
Example: 2-pentanone, propanone (more commonly known as
acetone)
Acetone: C3H6O O
CH3-C-CH3
153
The Carboxylic acids
These are hydrocarbons with a carboxyl functional group (-OOH)
attached to the end of a chain. They have a double-bonded oxygen
(=O) and an hydroxyl (-OH) group attached to the same carbon, taking
the place of three hydrogens. The suffix ends with -ic acid.
Examples: methanoic acid, ethanoic acid (also known as acetic
acid or vinegar)
O
Acetic acid: CH3COOH CH3-C-OH
The Ethers
These are hydrocarbons with an oxygen bonded between two carbon
chains. It is simply called an ether functional group (R’-O-R); where the R’s
stand for any type of a carbon chain. The name of the compound ends
with the word ether.
Examples: dimethyl ether, ethylmethyl ether
Dimethyl ether: CH3OCH3 CH3-O-CH3
The Esters
These are hydrocarbons with a double bonded oxygen and an -OR group
attached to a carbon chain. (Remember that R stands for any type of a
carbon chain). The suffix ends with -oate. Though the common name
ending is usually acetate. They look like this: R’-C-OOR
Examples: ethyl methanoate, ethyl ethanoate, methyl
ethanoate (methyl acetate)
O
Methyl acetate : CH3COOCH3 CH3-C-O-CH3
154
The Amines
These are hydrocarbons with an amino group (-NH2) on a chain of
carbons.
The suffix ends with -amine..
Examples: methylamine, ethylamine
Methylamine: CH3NH2 CH3-NH2
The Amides
Hydrocarbons with an amino group (-NH2) and a carbonyl group (=O) on
the same carbon.
The suffix ends with -amide..
Examples: methylamide, ethylamide
O
Methylamide: CHONH2 H-C-NH2
155
Organic Chemistry
1. Name the following compounds:
a. CH3-CH-CH2-CH2-CH2-CH3
CH3
CH3
b. CH3-CH2-CH2-C-CH3
CH3
c. CH3-CH2-CH2-CH-CH3
CH2-CH3
CH2-CH3
d. CH3-CH2-CH-CH-CH2-CH2-CH3
CH2-CH3
2. Draw the following compounds (complete with hydrogens):
a. 4-propyl decane
b. 3,5-diethyl octane
c. 3,3-diethyl octane
156
3. Draw the following compounds (complete with hydrogens):
a. 3-ethyl,4-methyl,5-propyl nonane
b. 3,3,5 -trimethyl octane
4. How many bonds does carbon make?
5. What is the general formula of all alkanes?
6. What is the molecular formula of butane?
7. Draw the structural formula for butane?
8. Draw an isomer of butane and name it (other than iso-butane).
9. Pentane has a higher boiling point than butane which has a higher
boiling point then methane.
Can you come up with a simple explanation for why this is true?
10. Draw three isomers for C5H12
157
Vocabulary
absolute temperature: This is a temperature reading made relative to
absolute zero. We use the unit of Kelvins for these readings.
absolute zero: This is the lowest temperature possible. If you remember
that temperature is a measurement of how much atoms move around
in a solid, you can guess that they stop moving entirely at absolute
zero. In reality, bonds still vibrate a little bit, but for the most part you
don't see much happening.
accuracy: When you measure something, the accuracy is how close
your measured value is to the real value. For example, if you're
actually six feet tall and your brother measures your height as six feet,
one inch, he's pretty accurate. However, if your cousin measures your
height as twelve feet, 13 inches, he's not accurate at all.
acid: This is anything that gives off H+ ions in water. Acids have a pH
less than 7 and are good at dissolving metals. They turn litmus paper
red and phenolphthalein colorless.
acid anhydride: This is an oxide that forms an acid when you stick it in
water. An example is SO3 - when you add water it turns into sulfuric
acid, H2SO4.
acid dissociation constant (Ka): This is equal to the ratio of the
concentrations of an acid's conjugate base and the acid present
when a weak acid dissociates in water. That is, if you have a solution
of Acid X where the concentration of the conjugate base is 0.5 M and
the concentration of the acid is 10 M, the acid dissociation constant is
0.5/10 = 0.05.
activated complex: In a chemical reaction, the reagents have to join
together into a great big blob before they can fall back apart into the
products. This great big blob is called the activated complex (a.k.a.
transition state)
activation energy: The minimum amount of energy needed for a
chemical reaction to take place. For some reactions this is very small
(it only takes a spark to make gasoline burn). For others, it's very high
(when you burn magnesium, you need to hold it over a Bunsen burner
for a minute or so).
activity series: This is when you arrange elements in the order of how
much they tend to react with water and acids.
actual yield: When you do a chemical reaction, this is the amount of
chemical that you actually make (i.e. The amount of stuff you can
weigh).
addition reaction: A reaction where atoms add to a carbon-carbon
multiple bond.
158
adsorption: When one substance collects of the surface of another
one.
alcohol: An organic molecule containing an -OH group
aldehyde: An organic molecule containing a -COH group
alkali metals: Group I in the periodic table.
alkaline earth metals: Group II in the periodic table.
alkane: An organic molecule which contains only single carbon-
carbon bonds.
alkene: An organic molecule containing at least one C=C bond
alkyne: An organic molecule containing at least one C-C triple bond.
allotropes: When you have different forms of an element in the same
state. The relationship that white phosphorus and red phosphorus
have to each other is that they're allotropes.
alloy: A mixture of two metals. Usually, you add very small amounts of
a different element to make the metal stronger and harder.
alpha particle: A radioactive particle equivalent to a helium nucleus
(2 protons, 2 neutrons)
amine: An organic molecule which consists of an ammonia molecule
where one or more of the hydrogen atoms has been replaced by
organic groups.
amino acid: The basic building blocks of proteins. They're called
"amino acids" because they're both amines (they contain nitrogen)
and acids (carboxylic acids, to be precise)
amphiprotic: When something is both an acid and a base. Like amino
acids, for example.
amphoteric: When something is both an acid and a base. Sounds
familiar, huh?
anode: The electrode where oxidation occurs. In other words, this is
where electrons are lost by a substance.
aqueous: dissolved in water
atomic mass unit (a.m.u.): This is the smallest unit of mass we use in
chemistry, and is equivalent to 1/12 the mass of carbon-12. To all
intents and purposes, protons and neutrons weigh 1 a.m.u.
atomic radius: This is one half the distance between two bonded
nuclei. Why don't we just measure the distance from the nucleus to
the outside of the atom - after all, isn't that the same thing as a radius?
It is, but atoms are also (theoretically) infinitely large (due to quantum
mechanics), making this impossible to measure.
atomic solid: A solid where there's a bunch of atoms in the lattice. This
is different from an ionic solid, where ions are the things that are
sticking together.
Aufbau principle: When you add protons to the nucleus to build up
the elements, electrons are added into orbitals.
159
Avogadro's Law: If you've got two gases under the same conditions of
temperature, pressure, and volume, they've got the same number of
particles (atoms or molecules). This law only works for ideal gases,
none of which actually exist.
base anhydride: An oxide that forms a base when water is added.
CaO is an example, turning into calcium hydroxide in water.
base: A compound that gives off OH- ions in water. They are slippery
and bitter and have a pH greater than 7.
battery: This is when a bunch of voltaic cells are stuck together.
beta particle: A radioactive particle equivalent to an electron.
bidentate ligand: A ligand that can attach twice to a metal ion.
binary compound: A compound only having two elements
binding energy: The amount of energy that holds the neutrons and
protons together in the nucleus of an atom. It's a lot of energy, which is
why you don't see nuclei falling apart all over the place.
bond energy: The amount of energy it takes to break one mole of
bonds.
bond length: The average distance between the nuclei of two
bonded atoms.
Boyle's Law: The volume of a gas at constant temperature varies
inversely with pressure. In other words, if you put big pressure on
something, it gets small.
Bronsted-Lowry acid: Acids donate protons [H+ ions] and bases grab
them
buffer: A liquid that resists change in pH by the addition of acid or
base. It consists of a weak acid and it's conjugate base (acetic acid
and sodium acetate, for example).
calorimetry: The study of heat flow. Usually you'd do calorimetry to
find the heat of combustion of a compound or the heat of reaction of
two compounds.
carboxylic acid: An organic molecule with a -COOH group on it.
Acetic acid is the most famous one.
catalyst: A substance that speeds up a chemical reaction without
being used up by the reaction. Enzymes are catalysts because they
allow the reactions that take place in the body to occur fast enough
that we can live.
cathode: The electrode in which reduction occurs. Reduction is when
a compound gains electrons.
chain reaction: A reaction in which the products from one step
provide the reagents for the next one. This is frequently referred to in
nuclear fission (when large nuclei break apart to form smaller ones)
and in free-radical reactions.
160
Charles's Law: The volume of a gas at constant pressure is directly
proportional to the temperature. In other words, if you heat something
up, it gets big.
chemical equation: The recipe that describes what you need to do to
make a reaction take place.
chemical properties: Properties that can only be described by making
a chemical change (by making or breaking bonds). For example,
color isn't a chemical property because you don't need to change
something chemically to see what color it is. Flammability, on the other
hand, is a chemical property, because you can't tell if something burns
unless you actually try to burn it.
chirality: When a molecule has a nonsuperimposable mirror image. To
imagine this, put your hands together. Although they are mirror
images, you can't put them right on top of each other so they are
interchangable. Well, normal people can't, anyway.
chromatography: This is when you use a system containing a mobile
phase (usually a liquid in general chemistry classes) and a stationary
phase (something dissolved in the liquid) to separate different
compounds. This is usually done by exploiting the differing polarities of
solutes, though you can do it a whole slew o' ways.
circuit: The closed path in a circuit through which electrons flow.
coagulation: When you destroy a colloid by letting the particles settle
out.
colligative property: Any property of a solution that changes when the
concentration changes. Examples are color, flavor, boiling point,
melting point, and osmotic pressure.
colloid: It's a suspension.
combustion: When a compound combines with oxygen gas to form
water, heat, and carbon dioxide
common ion effect: When the equilibrium position of a process is
altered by adding another compound containing one of the same
ions that's in the equilibrium.
complex ion: An ion in which a central atom (usually a transition
metal) is surrounded by a bunch of molecules like water or ammonia
(called "ligands")
concentration: A measurement of the amount of stuff (solute)
dissolved in a liquid (solvent). The most common concentration unit is
molarity (M), which is equal to the number of moles of solute divided
by the number of liters of solution.
condensation: When a vapor reforms a liquid. This is what happens on
your bathroom mirror when you take a shower.
conductance: A measurement of how well electricity can flow
through an object.
161
conjugate acid: The compound formed when a base gains a proton
(hydrogen atom).
conjugate base: The compound formed when an acid loses a proton
(hydrogen atom).
continuous spectrum: A spectrum that gives off all the colors of light,
like a rainbow. This is caused by blackbody emission.
covalent bond: A chemical bond formed when two atoms share two
electrons.
critical mass: The minimum amount of radioactive material needed to
undergo a nuclear chain reaction.
critical point: The end point of the liquid-vapor line in a phase
diagram. Past the critical point, you get something called a
"supercritical liquid", which has weird properties.
crystal lattice: see "lattice"
crystal: A large chunk of an ionic solid.
Dalton's law of partial pressures: The total pressure in a mixture of
gases is equal to the sums of the partial pressures of all the gases put
together.
decomposition: When a big molecule falls apart to make two or more
little ones.
degenerate: Things (usually orbitals) are said to be degenerate if they
have the same energy. This term is used a whole lot in quantum
mechanics. Also when dealing with kids who steal cars.
delocalization: This is when electrons can move around all over a
molecule. This happens when you have double bonds on adjacent
atoms in a molecule (conjugated hydrocarbon)
denature: When the 3-D structure of a protein breaks down due to
heat (or pH, etc), it's said to be denatured. This means that it unravels
because the intermolecular forces between atoms in the chain aren't
strong enough to hold it together anymore.
diffusion: When particles move from areas of high concentration to
areas of low concentration. For example, if you open a bottle of
ammonia on one end of the room, the concentration of ammonia
molecules in the air is very high on that side of the room. As a result,
they tend to migrate across the room, which explains why you can
smell it after a little while. Be careful not to mix this up with effusion (see
definition)
dilution: When you add solvent to a solution to make it less
concentrated.
dipole moment: When a molecule has some charge separation
(usually because the molecule is polar), it's said to have a dipole
moment.
dipole-dipole force: When the positive end of a polar molecule
becomes attracted to the negative end of another polar molecule.
162
dissociation: When water dissolves a compound.
distillation: This is when you separate a mixture of liquids by heating it
up. The one with the lowest boiling point evaporates first, followed by
the one with the next lowest boiling point, etc.
double-displacement reaction (a.k.a. double replacement reaction):
When the cations of two ionic compounds switch places.
effusion: When a gas moves through an opening into a chamber that
contains no pressure. Effusion is much faster than diffusion because
there are no other gas molecules to get in the way.
electrolysis: When electricity is used to break apart a chemical
compound.
electrolyte: An ionic compound that dissolves in water to conduct
electricity. Strong electrolytes break apart completely in water; weak
electrolytes only fall apart a little bit.
electron affinity: The energy change that accompanies the addition
of an electron to an atom in the gas phase.
electronegativity: A measurement of how much an atom tends to
steal electrons from atoms that it's bonded to. Elements at the top
right of the periodic table (excluding the noble gases) are very
electronegative while atoms in the bottom left are not very
electronegative (a.k.a. "electropositive")
electropositive: When something is not at all electronegative. In fact,
it tends to lose electrons rather than to gain them. Elements that are
electropositive are generally to the left and bottom of the periodic
table.
empirical formula: A reduced molecular formula. If you have a
molecular formula and you can reduce all of the subscripts by some
constant number, the result is the empirical formula.
emulsion: When very small drops of a liquid are suspended in another.
An example of an emulsion is salad dressing after you've shaken it up.
enantiomers: molecules that are nonsuperimposable mirror images of
each other.
endothermic: When a process absorbs energy (gets cold).
endpoint: The point where you actually stop a titration, usually
because an indicator has changed color. This is different than the
"equivalence point" because the indicator might not change colors at
the exact instant that the solution is neutral.
energy level: A possible level of energy that an electron can have in
an atom.
enthalpy: A measurement of the energy content of a system.
entropy: A measurement of the randomness in a system.
enzyme: A biological molecule that catalyzes reactions in living
creatures.
163
equilibrium: When the forward rate of a chemical reaction is the same
as the reverse rate. This only takes place in reversible reactions
because these are the only type of reaction in which the forward and
backward reactions can both take place.
equivalence point: The point in a titration at which the solution is
completely neutral. This is different than the "endpoint" (see above).
ester: An organic molecule with R-CO-OR' functionality.
excess reagent: Sometimes when you do a chemical reaction, there's
some of one reagent left over. That's called the excess reagent.
excited state: A higher energy level that electrons can jump to when
energy is added.
exothermic: When a process gives off energy (gets hot).
family: The same thing as a "group" (see above)
first law of thermodynamics: The energy of the universe is constant. It's
the same thing as the Law of conservation of energy.
fission: A nuclear reaction where a big atom breaks up into little ones.
This is what happens in nuclear power plants.
free energy: also called "Gibbs free energy", it's the capacity of a
system to do work.
free radical: An atom or molecule with an unpaired electron. They're
way reactive.
functional group: A generic term for a group of atoms that cause a
molecule to react in a specific way. It's really common to talk about
this in organic chemistry, where you have "aldehydes, carboxylic acids,
amines" and so on.
gamma ray: High energy light given off during a nuclear process.
When a nucleus gives off this light, it goes to a lower energy state,
making it more stable.
geometrical isomer: isomerism where atoms or groups of atoms can
take up different positions around a double bond or a ring. This is also
called cis- trans- isomerism.
ground state: The lowest energy state possible for an electron.
group: A column (the things up and down) in the periodic table.
Elements in the same group tend to have the same properties. These
are also called "families".
half-life: The time required for half of the radioactive atoms in a
sample to decay. When talking about chemical reactions, it's the
amount of time required to make half the reagent react.
half-reaction: The oxidation or reduction part of a redox reaction.
halogen: The elements in group 17. They're really reactive.
heat of reaction: The amount of heat absorbed or released in a
reaction. Also called the "enthalpy of reaction"
heat: The kinetic energy of the particles in a system. The faster the
particles move, the higher the heat.
164
Hess's Law: The enthalpy change for a change is the same whether it
takes place in one big step or in many small ones.
heterogeneous mixture: A mixture where the substances aren't equally
distributed.
homogeneous mixture: A mixture that looks really "smooth" because
everything is mixed up really well.
Hund's rule: The most stable arrangement of electrons occurs when
they're all unpaired.
hybrid orbital: An orbital caused by the mixing of s, p, d, and f-orbitals.
hydration: When a molecule has water molecules attached to it.
hydrocarbon: A molecule containing carbon and hydrogen.
hydrogen bond: The tendency of the hydrogen atom stuck to an
electronegative atom to become attracted to the lone pair electrons
on another electronegative atom. It's a pretty strong intermolecular
force, which explains why water has such a high melting and boiling
point.
hydrogenation: When hydrogen is added to a carbon-carbon multiple
bond.
hydronium ion: The H+ ion, made famous by acids.
hydroxide ion: The OH- ion, made famous by bases.
ideal gas law: PV=nRT
ideal gas: A gas in which the particles are infinitely small, have a
kinetic energy directly proportional to the temperature, travel in
random straight lines, and don't attract or repel each other. Needless
to say, there's no such thing as an ideal gas in the real world. However,
we use ideal gases anyway because they make the math work out
well for equations that describe how gases behave.
ideal solution: A solution in which the vapor pressure is directly
proportional to the mole fraction of solvent present
immiscible: When two substances don't dissolve in each other. Think
of oil and water. They're immiscible. Organic compounds and water
are frequently immiscible.
indicator: A compound that turns different colors at different pH
values. We generally like to have the color change at a pH of around
seven because that's where the equivalence point of a titration is.
inhibitor: A substance that slows down a chemical reaction.
inorganic compound: Any compound that doesn't contain carbon
(except for carbon dioxide, carbon monoxide, and carbonates).
insoluble: When something doesn't dissolve.
intermediate: A molecule which exists for a short time in a chemical
reaction before turning into the product.
intermolecular force: A force that exists between two different
molecules. Examples are hydrogen bonding (which is strong), dipole-
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dipole forces (which are kind of weak), and London dispersion forces
(a.k.a. Van der Waal forces), which are very weak.
ionic bond: A bond formed when charge particles stick together.
ionization energy: The amount of energy required to pull an electron
off of a gaseous atom.
irreversible reaction: A chemical reaction in which the reagents make
products but the products can't reform reagents. Most chemical
reactions in basic chemistry classes are thought of as being irreversible.
isotonic solutions: Solutions containing the same osmotic pressure.
isotope: When an element has more than one possibility for the
number of neutrons, these are called isotopes. All known elements
posess isotopes. For the record, the word "isotope" doesn't imply that
something is radioactive. TV told you that, and TV is stupid.
Kelvin: A unit used to measure temperature. One Kelvin is equal in size
to one degree Celsius. To convert between degrees Celsius and
Kelvins, simply add 273.15 to the temperature in degrees Celsius to get
Kelvins.
ketone: A molecule containing a R-CO-R' functional group. Acetone
(dimethyl ketone) is a common one.
kinetic energy: The energy due to the movement of an object. The
more something moves, the more kinetic energy it has.
Lanthanide contraction: The tendency of the lanthanides to get small
when you go from left to right in the periodic table.
lattice energy: The energy released when one mole of a crystal is
formed from gaseous ions.
lattice: The three-dimensional arrangement of atoms or ions in a
crystal.
law of conservation of energy: The amount of energy in the universe
never changes, ever. It just changes form.
law of conservation of mass: The amount of stuff after a chemical
reaction takes place is the same as the amount of stuff you started
with.
Le Chatlier's Principle: When you disturb an equilibrium (by adding
more chemical, by heating it up, etc.), it will eventually go back into
equilibrium under a different set of conditions.
Lewis acid: An electron-pair acceptor (carbonyl groups are really
good ones)
Lewis base: An electron-pair donor. Things with lone pairs like water
and ammonia are really good ones.
Lewis structure: A structural formula that shows all of the atoms and
valence electrons in a molecule.
ligand: A molecule or ion that sticks to the central atom in a complex.
Common examples are ammonia, carbon monoxide, or water.
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limiting reagent: If you do a chemical reaction and one of the
chemicals gets used up before the other one, the one that got used
up is called the "limiting reagent" because it limited the amount of
product that could be formed. The other one is called the excess
reagent.
line spectrum: A spectrum showing only certain wavelengths.
London dispersion force: The forces between nonpolar atoms or
molecules which is caused by momentary induced dipoles. It's real
weak.
lone pair: two electrons that aren't involved in chemical bonding. Also
frequently referred to as an "unshared pair".
main-block elements: Groups 1,2, and 13-18 in the periodic table.
They're called main block elements because the outermost electron is
in the s- or p- orbitals. What that has to do with the term "main block" is
unclear to me, but hey, that's life.
mass defect: The difference between the mass of an atom and the
sum of the masses of its individual components. Atoms usually weigh a
little less than if you added up the weights of all the particles. This is
because that extra mass was converted into the energy which holds
the atom together (see "binding energy")
mass: The amount of matter in an object. The more mass, the more
stuff is present.
mechanism: A step-by-step sequence that shows how the products of
a reaction are made from the reagents. Mechanisms are very
frequently shown during organic chemistry.
molality: The number of moles of solute per kilogram of solvent in a
solution. This is a unit of concentration that's not anywhere near as
handy or common as molarity.
molar mass: The mass of one mole of particles.
molar volume: The volume of one mole of a substance at STP. If you
believe that everything is an ideal gas, this is always 22.4 liters.
Unfortunately, there's no such thing as an ideal gas.
molarity: A unit of concentration equal to moles of solute divided by
liters of solution.
mole fraction: The number of moles of stuff in a mixture that are due to
one of the compouds.
mole ratio: The ratio of moles of what you've been given in a reaction
to what you want to find. Handy in stoichiometry.
mole: 6.02 x 1023 things.
molecular compound: A compound held together by covalent
bonds.
molecular formula: A formula that shows the correct quantity of all of
the atoms in a molecule.
monatomic ion: An ion that has only one atom, like the chloride ion.
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neutralization reaction: The reaction of an acid with a base to form
water and a salt.
node: A location in an orbital where there's no probability of finding an
electron.
nonpolar covalent bond: A covalent bond where the electrons are
shared equally between the two atoms.
normal boiling point: The boiling point of a substance at 1.00 atm.
normal melting point: The melting point of a substance at 1.00 atm.
normality: The number of equivalents of a substance dissolved in a liter
of solution.
nuclar fusion: When many small atoms combine to form a large one.
This occurs during a thermonuclear reaction.
nuclear fission: This is when the nucleus of an atom breaks into many
parts.
nuclear reaction: Any reaction that involves a change in the nucleus
of an atom. Nuclear reactions take loads of energy, which is why you
don't see them much around the lab.
nucleon: A particle (such as proton or neutron) that's in the nucleus of
an atom.
octet rule: All atoms want to be like the nearest noble gas. (Well, they
all want to have the same number of valence electrons, anyway). To
do this, they either gain or lose electrons (to form ionic compounds) or
share electrons (to form covalent compounds).
optical isomerism: Isomerism in which the isomers cause plane
polarized light to rotate in different directions.
orbital: This is where the electrons in an atom live.
organic compound: A compound that contains carbon (except
carbon dioxide, carbon monoxide, and carbonates)
osmosis: The flow of a pure liquid into an area of high concentration
through a semi-permeable membrane.
oxidation number: The apparent charge on an atom.
oxidation: When a substance loses electrons.
partial pressure: The pressure of one gas in a mixture. For example, if
you had a 50:50 mix of helium and hydrogen gases and the total
pressure was 2 atm, the partial pressure of hydrogen would be 1 atm.
Pauli exclusion principle: No two electrons in an atom can have the
same quantum numbers.
percent yield: The actual yield divided by the theoretical yield, times
100.
period: A row (left to right) in the periodic table.
periodic law: The properties of elements change with increasing
atomic number in a periodic way. That's why you can stick the
elements into a big chart and have the elements line up in nice
families.
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pH: -log[H+]
phase diagram: A chart which shows how the phase depends on
various conditions of temperature and pressure.
phase: The state of a compound (solid, liquid, or gas)
physical property: A property which can be determined without
changing something chemically. If that doesn't make sense, see the
definition of "chemical change".
pi-bond: A double bond.
polar covalent bond: A covalent bond where one atom tries to grab
the electrons from the other one. This occurs because the
electronegativities of the two atoms aren't the same.
polyatomic: contains more than one atom.
polymer: A molecule containing many repeating units. Plastics are
polymers and are formed by free radical chain reactions.
polyprotic acid: An acid that can give up more than one hydronium
ion. Examples are sulfuric acid and phosphoric acid.
potential energy: The energy something has because of where it is.
Things that are way up high have more potential energy than things
that are way down low because they have farther to fall.
precision: A measurement of how repeatable a measurement is. The
more significant figures, the more precise the measurement.
pressure: Force/area
product: The thing you make in a chemical reaction.
quantum theory: The branch of physical chemistry that describes how
energy can only exist at certain levels and makes generalizations
about how atoms behave from this assumption.
radioactive: When a substance has an unstable nucleus that can fall
apart, it's referred to as radioactive.
Raoult's Law: The vapor pressure of a solution is directly proportional to
the mole fraction of the solvent.
rate determining step: The slowest step in a chemical reaction.
rate law: A mathematical expression for the speed of a reaction as a
function of concentration. A hint: It's usually true that things go faster if
you have more stuff in the first place.
redox reaction: A reaction that has both an oxidation and reduction.
resonance structure: When more than one valid Lewis structure can
be drawn for a molecule, these structures are said to be resonance
structures. Resonance structures arise from the fact that the electrons
are delocalized.
reversible reaction: A reaction in which the products can make
reagents, as well as the reagents making products.
root mean square velocity (RMS velocity): The square root of the
average of the squares of the individual velocities of the gas particles
in a mixture. To put it in a way that a normal human can understand,
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it's the average of how fast the particles in a gas are going (assuming
you ignore the direction they're traveling in).
salt: An ionic compound.
saturated: When the maximum amount of solute is dissolved in a liquid
Second law of thermodynamics: Whenever you do something, the
universe gets more random.
semiconductor: A substance that conducts electricity poorly at room
temperature, but has increasing conductivity at higher temperatures.
Metalloids are usually good semiconductors.
shielding effect: The outer electrons aren't pulled very tightly by the
nucleus because the inner electrons repel them. This repulsion is called
the shielding effect, and can be used to explain lots of neat-o stuff.
sigma bond: A real fancy way of saying "single bond"
significant figure: The number of digits in a number that tell you useful
information. For example, when you weigh yourself on a bathroom
scale, it says something like 150 pounds rather than 150.32843737
pounds. Why? Because the thing can only weigh accurately to the
nearest pound. Any other digits that are on this number don't mean
anything, because they're probably wrong anyway.
single-displacement reaction (a.k.a. single replacement reaction):
When one unbonded element replaces an element in a chemical
compound. These are frequently redox reactions.
solubility: A measurement of how much of a solute can dissolve in a
liquid.
solubility product constant: Abbreviated Ksp, this value indicates the
degree to which a compound dissociates in water. The higher the
solubility product constant, the more soluble the compound.
solute: The solid that gets dissolved in a solution.
solvent: The liquid that dissolves the solid in a solution.
specific heat capacity: The amount of heat required to increase the
temperature of one gram of a substance by one degree.
spectator ions: The ions in a reaction that don't react.
spontaneous change: A change that occurs by itself. All exothermic
reactions are spontaneous. However, this doesn't mean that all
exothermic reactions are fast. The combustion of gasoline is
spontaneous, but not very fast unless you add a little energy.
standard temperature and pressure: One atmosphere and 273 K.
steric hindrance: This is the idea that the functional groups on big
molecules get in the way of a chemical reaction, making it go slower.
Imagine a fat guy trying to get into a Honda Prelude - that's steric
hindrance.
stoichiometry: The art of figuring how much stuff you'll make in a
chemical reaction from the amount of each reagent you start with.
STP: See standard temperature and pressure.
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strong acid: An acid that fully dissociates in water
strong nuclear force: The force that holds the nucleus together. As the
name suggests, this force is strong.
structural formula: See Lewis structure.
sublimation: When a solid can change directly into a gas. Dry ice
does this.
supercooling: When you cool something below its normal freezing
point
supersaturated: When more solute is dissolved in a liquid than is
theoretically possible. This doesn't happen much, as you might
imagine.
surface tension: A measurement of how much the molecules on a
liquid tend to like to stick to each other. If something has a high
surface tension, it likes to bead up.
suspension: A mixture that looks homogeneous when you stir it, but
where the solids settle out when you stop. Mud is a very short-lived
suspension, while peanut butter is a very long-lived suspension.
synthesis: When you make a big molecule from two or more smaller
ones.
system: Everything you're talking about at the moment.
temperature: A measurement of the average kinetic energy of the
particles in a system.
theoretical yield: The amount of product which should be made in a
chemical reaction if everything goes perfectly.
thermodynamics: The study of energy
Third law o' thermodynamics: The randomness of a system at 0 K is
zero.
titration: When the concentration of an acid or base is determined by
neutralizing it.
transition state: See "activated complex"
triple point: The temperature and pressure at which all three states of a
substance can exist in equilibrium.
unit cell: The simplest part of a crystal that can be repeated over and
over to make the whole thing.
unsaturated: When you haven't yet dissolved all of the solute that's
possible to dissolve in a liquid.
unshared electron pair: two electrons that aren't involved in chemical
bonding. Also frequently referred to as a "lone pair".
valence electron: The outermost electrons in an atom.
vapor pressure: The pressure of a substance that's present above it's
liquid. For example, you can tell that ammonia has a high vapor
pressure because the smell of it is very strong above liquid ammonia.
vaporization: When you boil a liquid.
volatile: A substance with a high vapor pressure.
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