Vergarachemistrynotebook

Jamyled Vergara’s Chemistry Notebook

Honors Chemistry
Class Policies and Grading
The students will receive a Unit Outline at the beginning of each Unit. It will
have information about the assignments that they will do, what it’s grade
classification will be, what action they will need to do to complete the
assignment and when it is due.
The students will receive a Weekly Memo of the activities they will be
responsible for that week. It will serve to inform the students of the learning
goal for the week. It will also give the students any special information
about that week.
The students will also receive daily lectures and assignments that are
designed to teach and re-enforce information related to the learning goal.
This will be time in which new material will be taught and reviewed and will
give the students the opportunity to ask questions regarding the concepts
being taught.
The students will work with a Lab partner and also be in a Lab group, but it
will be up to the individual student to do his or her part of all assignments
and the individual student will ultimately be responsible for all information
presented in the class.
The students will be required to follow all District and School Policies and to
follow all Lab Safety Procedures, which they will be given and will sign,
while performing labs. Students should come to class on time and with the
supplies needed for that class.
The following grading policy will be used.
Percent of Final Grade
Notebook 40%
Test/Projects 30%
Labs/Quizzes 20%
Work 10%
The students will be given a teacher generated Mid Term and a District
Final.

Unit 1
Measurement Lab
Separation of Mixtures Lab with Lab Write Up
Unit 2
Flame Test Lab
Nuclear Decay Lab
Element Marketing Project
Unit 3
Golden Penny Lab with Lab Write Up
Molecular Geometry
Research Presentation on a Chemical
Mid Term
Unit 4
Double Displacement Lab
Stoichiometry Lab with Lab Write Up
Mole Educational Demonstration Project
Unit 5
Gas Laws Lab with Lab Write Up
States of Matter Lab
Teach a Gas Law Project
Unit 6
Dilutions Lab
Titration Lab
District Final
1

Unit 1 (22 days)
Chapter 1 Introduction to Chemistry
Honors Chemistry
2016/2017 Syllabus
3 days
1.1 The Scope of Chemistry 1.3 Thinking Like a Scientist
1.2 Chemistry and You 1.4 Problem Solving in Chemistry
Chapter 2 Matter and Change
2.1 Properties of Matter 2.3 Elements and Compounds
2.2 Mixtures 2.4 Chemical Reactions
Chapter 3 Scientific Measurement
9 days
10 days
3.1 Using and Expressing Measurements 3.3 Solving Conversion Problems
3.2 Units of Measurement
Unit 2 (15 days)
Chapter 4 Atomic Structure
5 days
4.1 Defining the Atom 4.3 Distinguishing Among Atoms
4.2 Structure of the Nuclear Atom
Chapter 5 Electrons in Atoms
5 days
5.1 Revising the Atomic Model 5.2 Electron Arrangement in Atoms
5.3 Atomic Emission Spectrum and the Quantum Mechanical Model
Chapter 6 The Periodic Table
6.1 Organizing the Elements 6.3 Periodic Trends
6.2 Classifying Elements
Unit 3 (22 days)
Chapter 25 Nuclear Chemistry
25.1 Nuclear Radiation 25.3 Fission and Fusion
25.2 Nuclear Transformations 25.4 Radiation in Your Life
Chapter 7 Ionic and Metallic Bonding
7.1 Ions 7.3 Bonding in Metals
7.2 Ionic Bonds and Ionic Compounds
Chapter 8 Covalent Bonding
5 days
6 days
8 days
8 days
8.1 Molecular Compounds 8.3 Bonding Theories
8.2 The Nature of Covalent Bonding 8.4 Polar Bonds and Molecules
Unit 4 (14 days)
Chapter 9 Chemical Names and Formulas
6 days
9.1 Naming Ions 9.3 Naming & Writing Formulas Molecular Compounds
9.2 Naming and Writing Formulas for Ionic Compounds 9.4 Names for Acids and Bases
Chapter 22 Hydrocarbons Compounds
22.1 Hydrocarbons 22.4 Hydrocarbon Rings
Chapter 23 Functional Groups
4 days
4 days
23.1 Introduction to Functional Groups 23.4 Alcohols, Ethers, and Amines
2

Unit 5 (28 days)
Chapter 10 Chemical Quantities 8 days
10.1 The Mole: A Measurement of Matter 10.3 % Composition & Chem. Formulas
10.2 Mole-Mass and Mole-Volume Relationships
Chapter 11 Chemical Reactions 8 days
11.1 Describing Chemical Reactions 11.3 Reactions in Aqueous Solutions
11.2 Types of Chemical Reactions
Chapter 12 Stoichiometry 12 days
12.1 The Arithmetic of Equations 12.3 Limiting Reagent and % Yield
12.2 Chemical Calculations
Unit 6 (22 days)
Chapter 13 States of Matter 6 days
13.1 The Nature of Gases 13.3 The Nature of Solids
13.2 The Nature of Liquids 13.4 Changes in State
Chapter 14 The Behavior of Gases 10 days
14.1 Properties of Gases 14.3 Ideal Gases
14.2 The Gas Laws 14.4 Gases: Mixtures and Movement
Chapter 15 Water and Aqueous Systems 6 days
15.1 Water and its Properties 15.3 Heterogeneous Aqueous Systems
15.2 Homogeneous Aqueous Systems
Unit 7 (18 days)
Chapter 16 Solutions 8 days
16.1 Properties of Solutions 16.3 Colligative Properties of Solutions
16.2 Concentrations of Solutions 16.4 Calc. Involving Colligative Property
Chapter 17 Thermochemistry 5 days
17.1 The Flow of Energy 17.3 Heat in Changes of State
17.2 Measuring and Expressing Enthalpy Change 17.4 Calculating Heats in Reactions
Chapter 18 Reaction Rates and Equilibrium 5 days
18.1 Rates of Reactions 18.3 Reversible Reaction & Equilibrium
18.2 The Progress of Chemical Reactions 18.5 Free Energy and Entropy
Unit 8 (14 days)
Chapter 19 Acid and Bases 10 days
19.1 Acid-Base Theories 19.4 Neutralization Reactions
19.2 Hydrogen Ions and Acidity 19.5 Salts in Solutions
19.3 Strengths of Acids and Bases
Chapter 20 Oxidation-Reduction Reactions 4 days
20.1 The Meaning of Oxidation and Reduction 20.3 Describing Redox Equations
20.2 Oxidation Numbers
3

Lorenzo Walker Technical High School
MUSTANG LABORATORIES
Chemistry Safety
Safety in the MUSTANG LABORATORIES - Chemistry Laboratory
Working in the chemistry laboratory is an interesting and rewarding experience. During your labs, you will be actively
involved from beginning to end—from setting some change in motion to drawing some conclusion. In the laboratory, you
will be working with equipment and materials that can cause injury if they are not handled properly.
However, the laboratory is a safe place to work if you are careful. Accidents do not just happen; they are caused—by
carelessness, haste, and disregard of safety rules and practices. Safety rules to be followed in the laboratory are listed
below. Before beginning any lab work, read these rules, learn them, and follow them carefully.
General
1. Be prepared to work when you arrive at the lab. Familiarize yourself with the lab procedures before beginning the lab.
2. Perform only those lab activities assigned by your teacher. Never do anything in the laboratory that is not called for in
the laboratory procedure or by your teacher. Never work alone in the lab. Do not engage in any horseplay.
3. Work areas should be kept clean and tidy at all times. Only lab manuals and notebooks should be brought to the work
area. Other books, purses, brief cases, etc. should be left at your desk or placed in a designated storage area.
4. Clothing should be appropriate for working in the lab. Jackets, ties, and other loose garments should be removed. Open
shoes should not be worn.
5. Long hair should be tied back or covered, especially in the vicinity of open flame.
6. Jewelry that might present a safety hazard, such as dangling necklaces, chains, medallions, or bracelets should not be
worn in the lab.
7. Follow all instructions, both written and oral, carefully.
8. Safety goggles and lab aprons should be worn at all times.
9. Set up apparatus as described in the lab manual or by your teacher. Never use makeshift arrangements.
10. Always use the prescribed instrument (tongs, test tube holder, forceps, etc.) for handling apparatus or equipment.
11. Keep all combustible materials away from open flames.
12. Never touch any substance in the lab unless specifically instructed to do so by your teacher.
13. Never put your face near the mouth of a container that is holding chemicals.
14. Never smell any chemicals unless instructed to do so by your teacher. When testing for odors, use a wafting motion to
direct the odors to your nose.
15. Any activity involving poisonous vapors should be conducted in the fume hood.
16. Dispose of waste materials as instructed by your teacher.
17. Clean up all spills immediately.
18. Clean and wipe dry all work surfaces at the end of class. Wash your hands thoroughly.
19. Know the location of emergency equipment (first aid kit, fire extinguisher, fire shower, fire blanket, etc.) and how to use them.
20. Report all accidents to the teacher immediately.
Handling Chemicals
21. Read and double check labels on reagent bottles before removing any reagent. Take only as much reagent as you
need.
22. Do not return unused reagent to stock bottles.
23. When transferring chemical reagents from one container to another, hold the containers out away from your body.
24. When mixing an acid and water, always add the acid to the water.
25. Avoid touching chemicals with your hands. If chemicals do come in contact with your hands, wash them immediately.
26. Notify your teacher if you have any medical problems that might relate to lab work, such as allergies or asthma.
27. If you will be working with chemicals in the lab, avoid wearing contact lenses. Change to glasses, if possible, or notify
the teacher.
Handling Glassware
28. Glass tubing, especially long pieces, should be carried in a vertical position to minimize the likelihood of breakage and
to avoid stabbing anyone.
29. Never handle broken glass with your bare hands. Use a brush and dustpan to clean up broken glass. Dispose of the
glass as directed by your teacher.
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5
30. Always lubricate glassware (tubing, thistle tubes, thermometers, etc.) with water or glycerin before attempting to insert
it into a rubber stopper.
31. Never apply force when inserting or removing glassware from a stopper. Use a twisting motion. If a piece of glassware
becomes "frozen" in a stopper, take it to your teacher.
32. Do not place hot glassware directly on the lab table. Always use an insulating pad of some sort.
33. Allow plenty of time for hot glass to cool before touching it. Hot glass can cause painful burns. (Hot glass looks cool.)
Heating Substances
34. Exercise extreme caution when using a gas burner. Keep your head and clothing away from the flame.
35. Always turn the burner off when it is not in use.
36. Do not bring any substance into contact with a flame unless instructed to do so.
37. Never heat anything without being instructed to do so.
38. Never look into a container that is being heated.
39. When heating a substance in a test tube, make sure that the mouth of the tube is not pointed at yourself or anyone
else.
40. Never leave unattended anything that is being heated or is visibly reacting.
First Aid in the MUSTANG LABORATORIES - Chemistry Laboratory
Accidents do not often happen in well-equipped chemistry laboratories if students understand safe laboratory procedures
and are careful in following them. When an occasional accident does occur, it is likely to be a minor one.
The instructor will assist in treating injuries such as minor cuts and burns. However, for some types of injuries, you must
take action immediately. The following information will be helpful to you if an accident occurs.
1. Shock. People who are suffering from any severe injury (for example, a bad burn or major loss of blood) may be in a
state of shock. A person in shock is usually pale and faint. The person may be sweating, with cold, moist skin and a weak,
rapid pulse. Shock is a serious medical condition. Do not allow a person in shock to walk anywhere—even to the campus
security office. While emergency help is being summoned, place the victim face up in a horizontal position, with the feet
raised about 30 centimeters. Loosen any tightly fitting clothing and keep him or her warm.
2. Chemicals in the Eyes. Getting any kind of a chemical into the eyes is undesirable, but certain chemicals are
especially harmful. They can destroy eyesight in a matter of seconds. Because you will be wearing safety goggles at all
times in the lab, the likelihood of this kind of accident is remote. However, if it does happen, flush your eyes with water
immediately. Do NOT attempt to go to the campus office before flushing your eyes. It is important that flushing with water
be continued for a prolonged time—about 15 minutes.
3. Clothing or Hair on Fire. A person whose clothing or hair catches on fire will often run around hysterically in an
unsuccessful effort to get away from the fire. This only provides the fire with more oxygen and makes it burn faster. For
clothing fires, throw yourself to the ground and roll around to extinguish the flames. For hair fires, use a fire blanket to
smother the flames. Notify campus security immediately.
4. Bleeding from a Cut. Most cuts that occur in the chemistry laboratory are minor. For minor cuts, apply pressure to the
wound with a sterile gauze. Notify campus security of all injuries in the lab. If the victim is bleeding badly, raise the
bleeding part, if possible, and apply pressure to the wound with a piece of sterile gauze. While first aid is being given,
someone else should notify the campus security officer.
5. Chemicals in the Mouth. Many chemicals are poisonous to varying degrees. Any chemical taken into the mouth
should be spat out and the mouth rinsed thoroughly with water. Note the name of the chemical and notify the campus
office immediately. If the victim swallows a chemical, note the name of the chemical and notify campus security
immediately.
If necessary, the campus security officer or administrator will contact the Poison Control Center, a hospital emergency
room, or a physician for instructions.
6. Acid or Base Spilled on the Skin.
Flush the skin with water for about 15 minutes. Take the victim to the campus office to report the injury.
7. Breathing Smoke or Chemical Fumes.
All experiments that give off smoke or noxious gases should be conducted in a well-ventilated fume hood. This will make
an accident of this kind unlikely. If smoke or chemical fumes are present in the laboratory, all persons—even those who
do not feel ill—should leave the laboratory immediately. Make certain that all doors to the laboratory are closed after the
last person has left. Since smoke rises, stay low while evacuating a smoke-filled room. Notify campus security
immediately.

MUSTANG LABORATORIES
COMMITMENT TO SAFETY IN THE LABORATORY
As a student enrolled in Chemistry at Lorenzo Walker Technical High
School, I agree to use good laboratory safety practices at all times. I
also agree that I will:
1. Conduct myself in a professional manner, respecting both my personal safety and the safety of
others in the laboratory.
2. Wear proper and approved safety glasses or goggles in the laboratory at all times.
3. Wear sensible clothing and tie back long hair in the laboratory. Understand that open-toed shoes
pose a hazard during laboratory classes and that contact lenses are an added safety risk.
4. Keep my lab area free of clutter during an experiment.
5. Never bring food or drink into the laboratory, nor apply makeup within the laboratory.
6. Be aware of the location of safety equipment such as the fire extinguisher, eye wash station, fire
blanket, first aid kit. Know the location of the nearest telephone and exits.
7. Read the assigned lab prior to coming to the laboratory.
8. Carefully read all labels on all chemical containers before using their contents, remove a small
amount of reagent properly if needed, do not pour back the unused chemicals into the original
container.
9. Dispose of chemicals as directed by the instructor only. At no time will I pour anything down the
sink without prior instruction.
10. Never inhale fumes emitted during an experiment. Use the fume hood when instructed to do so.
11. Report any accident immediately to the instructor, including chemical spills.
12. Dispose of broken glass and sharps only in the designated containers.
13. Clean my work area and all glassware before leaving the laboratory.
14. Wash my hands before leaving the laboratory.
NAME __________________________
Jamyled Vergara
PARENT NAME ____________________________
Mayte Alonso
PERIOD ________________________
1 PARENT NUMBER _________________________
239-537-9085
SIGNATURE ____________________________
DATE ____________________________________
08/25/16
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7

Chapter 1
Unit 1
Introduction to Chemistry
The students will learn why and how to solve problems using
chemistry.
Identify what is science, what clearly is not science, and what superficially
resembles science (but fails to meet the criteria for science).
Students will identify a phenomenon as science or not science.
Science
Inference
Observation
Hypothesis
Identify which questions can be answered through science and which
questions are outside the boundaries of scientific investigation, such as
questions addressed by other ways of knowing, such as art, philosophy, and
religion.
Students will differentiate between problems and/or phenomenon that can and
those that cannot be explained or answered by science.
Students will differentiate between problems and/or phenomenon that can and
those that cannot be explained or answered by science.
Observation
Inference
Hypothesis
Theory
Controlled experiment
Describe how scientific inferences are drawn from scientific observations
and provide examples from the content being studied.
Students will conduct and record observations.
Students will make inferences.
Students will identify a statement as being either an observation or inference.
Students will pose scientific questions and make predictions based on
inferences.
Inference
Hypothesis
Observation
Controlled experiment
Identify sources of information and assess their reliability according to the
strict standards of scientific investigation.
Students will compare and assess the validity of known scientific information
from a variety of sources:
8

Print vs. print
Online vs. online
Print vs. online
Students will conduct an experiment using the scientific method and compare
with other groups.
Controlled experiment
Investigation
Peer Review
Accuracy
Precision
Percentage Error
Chapter 2
Matter and Change
The students will learn what properties are used to describe
matter and how matter can change its form.
Differentiate between physical and chemical properties and physical and
chemical changes of matter.
Students will be able to identify physical and chemical properties of various
substances.
Students will be able to identify indicators of physical and chemical changes.
Students will be able to calculate density.
mass
mixture
physical property
intensive property
volume
solution
chemical property
element
vapor
compound
extensive property
Chapter 3
Scientific Measurements
The students will be able to solve conversion problems using
measurements.
Determine appropriate and consistent standards of measurement for the
data to be collected in a survey or experiment.
Students will participate in activities to collect data using standardized
measurement.
Students will be able to manipulate/convert data collected and apply the data
to scientific situations.
Scientific notation
Experimental value
International System of Units (SI)
Percent error
Significant figures
Dimensional analysis
Accepted value
9

Benchmark
Determine appropriate and consistent standards of measurement for the data to be collected in a survey or
experiment.
K H D B D C M
1000 100 10 base 1/10 1/100 1/1000
*King Henry Died By Drinking chocolate milk
cm x cm x cm = cm
10

To use the Stair-Step method, find the prefix the original measurement starts with. (ex. milligram)
If there is no prefix, then you are starting with a base unit.
Find the step which you wish to make the conversion to. (ex. decigram)
Count the number of steps you moved, and determine in which direction you moved (left or right).
The decimal in your original measurement moves the same number of places as steps you moved and in the
same direction. (ex. milligram to decigram is 2 steps to the left, so 40 milligrams = .40 decigrams)
If the number of steps you move is larger than the number you have, you will have to add zeros to hold the
places. (ex. kilometers to meters is three steps to the right, so 10 kilometers would be equal to 10,000 m)
That’s all there is to it! You need to be able to count to 6, and know your left from your right!
1) Write the equivalent
a) 5 dm =_______m 0.5 b) 4 mL = ______L 0.004 c) 8 g = _______mg 8000
d) 9 mg =_______g .009 e) 2 mL = ______L .002 f) 6 kg = _____g 6000
g) 4 cm =_______m 400
h) 12 mg = .012 ______ g i) 6.5 cm 3 = _______L .0065
j) 7.02 mL =_____cm 7.02 3 k) .03 hg = _______ .3 dg l) 6035 mm 603500 _____cm
m) .32 m = _______cm 32
n) 38.2 g = _____kg .0382
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2. One cereal bar has a mass of 37 g. What is the mass of 6 cereal bars? Is that more than or less
than 1 kg? Explain your answer.
222g 1k =l000g
less because 1k= 1000g
3. Wanda needs to move 110 kg of rocks. She can carry l0 hg each trip. How many trips must she
make? Explain your answer.
110 10hg = 1kg
4. Dr. O is playing in her garden again She needs 1 kg of potting soil for her plants. She has 750 g.
How much more does she need? Explain your answer.
5. Weather satellites orbit Earth at an altitude of 1,400,000 meters. What is this altitude in kilometers?
6. Which unit would you use to measure the capacity? Write milliliter or liter.
liter
a) a bucket __________
b) a thimble __________
milliliter
c) a water storage tank__________ liter
d) a carton of juice__________
liter
250g
1400 km
7. Circle the more reasonable measure:
a) length of an ant 5mm or 5cm
5mm
b) length of an automobile 5 m or 50 m
5m
c) distance from NY to LA 450 km or 4,500 km
4500km
d) height of a dining table 75 mm or 75 cm
75cm
8. Will a tablecloth that is 155 cm long cover a table that is 1.6 m long? Explain your answer.
no, because its 0.05 meters too short
9. A dollar bill is 15.6 cm long. If 200 dollar bills were laid end to end, how many meters long would
the line be?
36.2 meters long
10. The ceiling in Jan’s living room is 2.5 m high. She has a hanging lamp that hangs down 41 cm.
Her husband is exactly 2 m tall. Will he hit his head on the hanging lamp? Why or why not?
no, because he is 9cm shorter than the light.
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Using SI Units
Match the terms in Column II with the descriptions in Column I. Write the letters of the correct term in
the blank on the left.
Column I Column II
_____ k 1. distance between two points
a. time
_____ e 2. SI unit of length
_____ m 3. tool used to measure length
_____ g 4. units obtained by combining other units
_____ b 5. amount of space occupied by an object
_____ h 6. unit used to express volume
_____ f 7. SI unit of mass
_____ c 8. amount of matter in an object
_____ d 9. mass per unit of volume
_____ o 10. temperature scale of most laboratory thermometers
_____ l 11. instrument used to measure mass
_____ a 12. interval between two events
_____ j 13. SI unit of temperature
_____ i 14. SI unit of time
_____ n 15. instrument used to measure temperature
b. volume
c. mass
d. density
e. meter
f. kilogram
g. derived
h. liter
i. second
j. Kelvin
k. length
1. balance
m. meterstick
n. thermometer
o. Celsius
Circle the two terms in each group that are related. Explain how the terms are related.
16. Celsius degree, mass, Kelvin _____________________________________________________
Celsius degree and kelvin are both SI units of temperature
________________________________________________________________________________
17. balance, second, mass __________________________________________________________
balance and mass both are dealing with mass
________________________________________________________________________________
18. kilogram, liter, cubic centimeter __________________________________________________
liter and cubic centimeter are both units of volume
________________________________________________________________________________
19. time, second, distance __________________________________________________________
time and second are both units of time
________________________________________________________________________________
20. decimeter, kilometer, Kelvin _____________________________________________________
decimeter and kilometer are both distance
________________________________________________________________________________
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1. How many meters are in one kilometer? __________ 1000
2. What part of a liter is one milliliter? __________
0.001
3. How many grams are in two dekagrams? __________ 20
4. If one gram of water has a volume of one milliliter, what would the mass of one liter of water be in
kilograms?__________ 1000
5. What part of a meter is a decimeter? __________ 0.1
In the blank at the left, write the term that correctly completes each statement. Choose from the terms
listed below.
Metric SI standard ten
prefixes ten tenth
6. An exact quantity that people agree to use for comparison is a ______________ standard .
7. The system of measurement used worldwide in science is _______________ SI
.
8. SI is based on units of _______________ ten
.
9. The system of measurement that was based on units of ten was the _______________ metric
system.
10. In SI, _______________ prefixes are used with the names of the base unit to indicate the multiple of ten
that is being used with the base unit.
11. The prefix deci- means _______________ tenth
.
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Standards of Measurement
Fill in the missing information in the table below.
Prefix
mili
S.I prefixes and their meanings
Meaning
0.001
0.01
centi
deci- 0.1
deka
10
hecto- 100
kilo
1000
Circle the larger unit in each pair of units.
1. millimeter, kilometer 4. centimeter, millimeter
kilometer
cenitmeter
2. decimeter, dekameter 5. hectogram, kilogram
dekameter
3. hectogram, decigram
kilogram
hectogram
6. In SI, the base unit of length is the meter. Use this information to arrange the following units of
measurement in the correct order from smallest to largest.
Write the number 1 (smallest) through 7 - (largest) in the spaces provided.
_____ 7 a. kilometer
6
_____ e. hectometer
_____ 2 b. centimeter
_____
4
c. meter
_____ 1 f. millimeter
_____ 3 g. decimeter
_____ 5 d. dekameter
Use your knowledge of the prefixes used in SI to answer the following questions in the spaces
provided.
7. One part of the Olympic games involves an activity called the decathlon. How many events do you
think make up the decathlon?_____________________________________________________
10
8. How many years make up a decade? _______________________________________________
10
9. How many years make up a century? ______________________________________________
100
10. What part of a second do you think a millisecond is? __________________________________
1000
15

The Learning Goal for this assignment is:
determine appropriate and consistent standards of measurement for the data to be
collected in a survey of experiment.
Notes Section
if the exponent is a negative number the decimal moves to the left
if the exponent is a positive number the decimal moves to the right
if the decimal moves to left the exponent gets larger
if the decimal moves to the right the exponent gets smaller
1. 7,485 6. 1.683
2. 884.2 7. 3.622
3. 0.00002887 8. 0.00001735
4. 0.05893 9. 0.9736
5. 0.006162 10. 0.08558
11. 6.633 X 10−⁴ 16. 1.937 X 10⁴
12. 4.445 X 10−⁴ 17. 3.457 X 10⁴
13. 2.182 X 10−³ 18. 3.948 X 10−⁵
14. 4.695 X 10² 19. 8.945 X 10⁵
15. 7.274 X 10⁵ 20. 6.783 X 10²
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SCIENTIFIC NOTATION RULES
How to Write Numbers in Scientific Notation
Scientific notation is a standard way of writing very large and very small numbers so that they're
easier to both compare and use in computations. To write in scientific notation, follow the form
N X 10 ᴬ
where N is a number between 1 and 10, but not 10 itself, and A is an integer (positive or negative
number).
RULE #1: Standard Scientific Notation is a number from 1 to 9 followed by a decimal and the
remaining significant figures and an exponent of 10 to hold place value.
Example:
5.43 x 10 2 = 5.43 x 100 = 543
8.65 x 10 – 3 = 8.65 x .001 = 0.00865
****54.3 x 10 1 is not Standard Scientific Notation!!!
RULE #2: When the decimal is moved to the Left the exponent gets Larger, but the value of the
number stays the same. Each place the decimal moves Changes the exponent by one (1). If you
move the decimal to the Right it makes the exponent smaller by one (1) for each place it is moved.
Example:
6000. x 10 0 = 600.0 x 10 1 = 60.00 x 10 2 = 6.000 x 10 3 = 6000
(Note: 10 0 = 1)
All the previous numbers are equal, but only 6.000 x 10 3 is in proper Scientific Notation.
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RULE #3: To add/subtract in scientific notation, the exponents must first be the same.
Example:
(3.0 x 10 2 ) + (6.4 x 10 3 ); since 6.4 x 10 3 is equal to 64. x 10 2 . Now add.
(3.0 x 10 2 )
+ (64. x 10 2 )
67.0 x 10 2 = 6.70 x 10 3 = 6.7 x 10 3
67.0 x 10 2 is mathematically correct, but a number in standard scientific notation can only
have one number to the left of the decimal, so the decimal is moved to the left one place and
one is added to the exponent.
Following the rules for significant figures, the answer becomes 6.7 x 10 3 .
RULE #4: To multiply, find the product of the numbers, then add the exponents.
Example:
(2.4 x 10 2 ) (5.5 x 10 –4 ) = ? [2.4 x 5.5 = 13.2]; [2 + -4 = -2], so
(2.4 x 10 2 ) (5.5 x 10 –4 ) = 13.2 x 10 –2 = 1.3 x 10 – 1
RULE #5: To divide, find the quotient of the number and subtract the exponents.
Example:
(3.3 x 10 – 6 ) / (9.1 x 10 – 8 ) = ? [3.3 / 9.1 = .36]; [-6 – (-8) = 2], so
(3.3 x 10 – 6 ) / (9.1 x 10 – 8 ) = .36 x 10 2 = 3.6 x 10 1
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Convert each number from Scientific Notation to real numbers:
1. 7.485 X 10³ 6. 1.683 X 10⁰
2. 8.842 X 10² 7. 3.622 10⁰
3. 2.887 X 10−⁵ 8. 1.735 X 10−⁵
4. 5.893 X 10−² 9. 9.736 X 10−¹
5. 6.162 X 10−³ 10. 8.558 X 10−²
Convert each number from a real number to Scientific Notation:
11. 0.0006633 16. 1,937,000
12. 0.0004445 17. 34,570
13. 0.002182 18. 0.00003948
14. 469.5 19. 894,500
15. 727,400 20. 678.3
19

The Learning Goal for this assignment is:
determine appropriate and consistent standards of measurement for the data to be
collected in a survey or experiment
Notes Section:
when a number is multiplied by an exponent its either gets bigger or smaller depending
whether the number is a negative or a
postive
Question Sig Figs Question Add & Subtract Question Multiple & Divide
1 4 1 55.36 1 20,000
2 4 2 84.2 2 94
3 3 3 115.4 3 300
4 3 4 0.8 4 7
5 4 5 245.53 5 62
6 3 6 34.5 6 0.005
7 3 7 74.0 7 4,000
8 2 8 53.287 8 3,900,000
9 2 9 54.876 9 2
10 2 10 40.19 10 30,000,000
11 3 11 7.7 11 1,200
12 2 12 67.170 12 0.2
13 3 13 81.0 13 0.87
14 4 14 73.290 14 0.049
15 4 15 29.789 15 2,000
16 3 16 39.53 16 0.5
17 4 17 70.58 17 1.9
18 2 18 86.6 18 0.05
19 2 19 64.990 19 230
20 1 20 36.0 20 460,000
20

Significant Figures Rules
There are three rules on determining how many significant figures are in a
number:
1. Non-zero digits are always significant.
2. Any zeros between two significant digits are significant.
3. A final zero or trailing zeros in the DECIMAL PORTION ONLY are
significant.
Please remember that, in science, all numbers are based upon measurements (except for a very few
that are defined). Since all measurements are uncertain, we must only use those numbers that are
meaningful.
Not all of the digits have meaning (significance) and, therefore, should not be written down. In
science, only the numbers that have significance (derived from measurement) are written.
Rule 1: Non-zero digits are always significant.
If you measure something and the device you use (ruler, thermometer, triple-beam, balance, etc.)
returns a number to you, then you have made a measurement decision and that ACT of measuring
gives significance to that particular numeral (or digit) in the overall value you obtain.
Hence a number like 46.78 would have four significant figures and 3.94 would have three.
Rule 2: Any zeros between two significant digits are significant.
Suppose you had a number like 409. By the first rule, the 4 and the 9 are significant. However, to
make a measurement decision on the 4 (in the hundred's place) and the 9 (in the one's place), you
HAD to have made a decision on the ten's place. The measurement scale for this number would have
hundreds, tens, and ones marked.
Like the following example:
These are sometimes called "captured zeros."
If a number has a decimal at the end (after the one’s place) then all digits (numbers) are significant
and will be counted.
In the following example the zeros are significant digits and highlighted in blue.
960.
70050.
21

Rule 3: A final zero or trailing zeros in the decimal portion ONLY are
significant.
This rule causes the most confusion among students.
In the following example the zeros are significant digits and highlighted in blue.
0.07030
0.00800
Here are two more examples where the significant zeros are highlighted in blue.
When Zeros are Not Significant Digits
4.7 0 x 10−³
6.5 0 0 x 10⁴
Zero Type # 1 : Space holding zeros in numbers less than one.
In the following example the zeros are NOT significant digits and highlighted in red.
0.09060
0.00400
These zeros serve only as space holders. They are there to put the decimal point in its correct
location.
They DO NOT involve measurement decisions.
Zero Type # 2 : Trailing zeros in a whole number.
In the following example the zeros are NOT significant digits and highlighted in red.
200
25000
For addition and subtraction, look at the decimal portion (i.e., to the right of the decimal point)
of the numbers ONLY. Here is what to do:
1) Count the number of significant figures in the decimal portion of each number in the problem. (The
digits to the left of the decimal place are not used to determine the number of decimal places in the
final answer.)
2) Add or subtract in the normal fashion.
3) Round the answer to the LEAST number of places in the decimal portion of any number in the
problem
The following rule applies for multiplication and division:
The LEAST number of significant figures in any number of the problem determines the number of
significant figures in the answer.
This means you MUST know how to recognize significant figures in order to use this rule.
22

How Many Significant Digits for Each Number?
1) 2359 = ______ 4
2) 2.445 x 10−⁵= ______ 4
3) 2.93 x 10⁴= ______ 3
4) 1.30 x 10−⁷= ______ 2
5) 2604 = ______ 4
6) 9160 = ______ 3
7) 0.0800 = ______ 3
8) 0.84 = ______ 2
9) 0.0080 = ______ 2
10) 0.00040 = ______ 2
11) 0.0520 = ______ 3
12) 0.060 = ______ 2
13) 6.90 x 10−¹= ______ 2
14) 7.200 x 10⁵= ______ 2
15) 5.566 x 10−²= ______ 4
16) 3.88 x 10⁸= ______ 3
17) 3004 = ______ 4
18) 0.021 = ______ 2
19) 240 = ______ 2
20) 500 = ______ 1
23

For addition and subtraction, look at the decimal portion (i.e., to the right of the decimal point) of the
numbers ONLY. Here is what to do:
1) Count the number of significant figures in the decimal portion of each number in the problem. (The
digits to the left of the decimal place are not used to determine the number of decimal places in the
final answer.)
2) Add or subtract in the normal fashion.
3) Round the answer to the LEAST number of places in the decimal portion of any number in the
problem.
Solve the Problems and Round Accordingly...
1) 43.287 + 5.79 + 6.284 = _______ 55.36
2) 87.54 - 3.3 = _______ 84.2
3) 99.1498 + 6.5397 + 9.7 = _______ 115.4
4) 5.868 - 5.1 = _______ 0.8
5) 59.9233 + 86.21 + 99.396 = _______
6) 7.7 + 26.756 = _______
7) 66.8 + 2.3 + 4.8516 = _______
8) 9.7419 + 43.545 = _______
9) 4.8976 + 48.4644 + 1.514 = _______
10) 4.335 + 35.85 = _______
11) 9.448 - 1.7 = _______
12) 75.826 - 8.6555 = _______
13) 57.2 + 23.814 = _______
14) 77.684 - 4.394 = _______
15) 26.4496 + 3.339 = _______
16) 9.6848 + 29.85 = _______
17) 63.11 + 2.5412 + 4.93 = _______
18) 11.2471 + 75.4 = _______
19) 73.745 - 8.755 = _______
20) 6.5238 + 1.7 + 27.79 = _______
24

25
The following rule applies for multiplication and division:
The LEAST number of significant figures in any number of the problem determines the number of
significant figures in the answer.
This means you MUST know how to recognize significant figures in order to use this rule.
Solve the Problems and Round Accordingly...
1) 0.6 x 65.0 x 602 = __________ 20000
2) 720 ÷ 7.7 = __________
94
3) 929 x 0.3 = __________ 300
4) 300 ÷ 44.31 = __________
7
5) 608 ÷ 9.8 = __________
6) 0.06 x 0.079 = __________
7) 0.008 x 72.91 x 7000 = __________
8) 73.94 x 67 x 780 = __________
9) 0.62 x 0.097 x 40 = __________
10) 600 x 10 x 5030 = __________
11) 5200 ÷ 4.46 = __________
12) 0.0052 x 0.4 x 107 = __________
13) 0.099 x 8.8 = __________
14) 0.0095 x 5.2 = __________
15) 8000 ÷ 4.62 = __________
16) 0.6 x 0.8 = __________
17) 2.84 x 0.66 = __________
18) 0.5 x 0.09 = __________
19) 8100 ÷ 34.84 = __________
20) 8.24 x 6.9 x 8100 = __________

Dimensional Analysis
This is a way to convert from one unit of a given substance to
another unit using ratios or conversion units. What this video
www.youtube.com/watch?v=aZ3J60GYo6U
Let’ look at a couple of examples:
1. Convert 2.6 qt to mL.
First we need a ratio or conversion unit so that we can go from quarts to milliliters. 1.00 qt = 946 mL
Next write down what you are starting with
2.6 qt
Then make you conversion tree
2.6 qt
Then fill in the units in your ratio so that you can cancel out the original unit and will be left with the
unit you need for the answer. Cross out units, one at a time that are paired, and one on top one on
the bottom.
2.6 qt mL
qt
Now fill in the values from the ratio.
2.6 qt 946 mL
1.00 qt
Now multiply all numbers on the top and multiply all numbers on the bottom and write them as a
fraction.
2.6 qt 946 mL = 2,459.6 mL
1.00 qt 1.00
Now divide the top number by the bottom number and write that number with the unit that was not
crossed out.
26

1qt=32 oz 1gal = 4qts 1.00 qt = 946 mL 1L = 1000mL
2. Convert 8135.6 mL to quarts
8135.6 1
946
=
8135.6
946
8.5qt
3. Convert 115.2 oz to mL
115.2 1
32
946
= 108979.2
32 3405.6ml
1
4. Convert 2.3 g to Liters
2.3
4
1
946
1
1
= 8703.2
1000
1000
8.7032L
5. Convert 8.42 L to oz
8.42
1000
1
1
946
32
1
=
269440
946
284.820296oz
Go to http://science.widener.edu/svb/tutorial/ chose #7 “Converting Volume” and do 5 more in the
space provided.
1. Convert _________
300.8oz
to _________ ml
300.8 1
946 284556.8
32
1
=
32
8892.4ml
0.48gal
2. Convert _________ to _________ l
=
3. Convert _________ to _________
=
4. Convert _________ to _________
=
5. Convert _________ to _________
=
27

Chapter 4
Unit 2
Atomic Structure
The students will learn what makes up atoms and how are atoms
of one element different from atoms of another element.
Explore the scientific theory of atoms (also known as atomic theory) by
describing changes in the atomic model over time and why those changes
were necessitated by experimental evidence.
• Students will be able to draw/identify each atomic model.
• Students will be able to compare/contrast the different atomic models.
• Students will be able to describe how results of experimental evidence caused
the atomic model to change.
proton
electron
neutron
nucleus
electron cloud
Explore the scientific theory of atoms (also known as atomic theory) by
describing the structure of atoms in terms of protons, neutrons and
electrons, and differentiate among these particles in terms of their mass,
electrical charges and locations within the atom.
• Students will compare/contrast the characteristics of subatomic particles.
atomic number
mass number
isotope
atomic mass unit (amu)
atomic mass

Chapter 5 Electrons in Atoms
The students will be able to describe the arrangement of
electrons in atoms and predict what will happen when
electrons in atoms absorb or release energy.
Describe the quantization of energy at the atomic level.
• Students will participate in activities to view emission spectrums using a
diffraction grating or a spectroscope.
• Students will be able to explain how the spectrum lines relate to electron motion.
energy level
atomic orbital
quantum mechanical model
Chapter 6 The Periodic Table
The student will learn what information the periodic table
provides and how periodic trends can be explained.
Relate properties of atoms and their position in the periodic table to the
arrangement of their electrons.
• Students will be able to compare and contrast metals, nonmetals, and
metalloids.
• Students will be able to describe the traits of various families on the periodic
table.
• Students will be able to explain periodicity.
• Students will write/represent electron configuration of various elements.
• Students will be able to use a periodic table to calculate the number of p + , e - ,
and n 0 .
• Students will be able to calculate the average weight of mass.
periodic law
halogen metals
noble gas
nonmetals
transition metal
metalloid atomic
radius alkali metal
ionization energy
alkaline earth

metal
electronegativity
The Learning Goal for this assignment is:
Explore the scientific theory of atoms (also known as atomic theory) by describing changes in the atomic
model overt time and why those changes were necessitated by experimental evidence.
Explore the scientific theory of atoms (also known as atomic theory) by describing the structure of atoms in
terms of protons, neutrons and electrons, and differentiate among these particles in terms of their mass,
electrical charges and locations within the atom.
Notes Section
Atoms are the basis of the universe and matter is composed of atoms.
- The number of protons determine what kind of element the atom
is -The neutrons determine if the atom is an isotope - Electrons tell
you if the atom is an ion or not.
-Nucleons and quarks are super-small particles moving around in atoms and can be found
inside the protons and neutrons.
-Protons and neutrons are always in the center of the atom (nucleus)
-Electrons are found circling around the center in areas called shells or orbitals.
"+", "-", or "0" refers to the charge of the particle.
Electron - "-"
Proton- "+"
Neutral- "0"
Ion- more "+" than "-" electrons in an atom
-The atomic number of an element (proton number), tells you the number of protons
or positive particles in an atom.
-Electrovalence- something that has given up or taken electrons and become an ion

-Valence is a measure of how much an atom wants to bond with other atoms.
-Two types of bonding: covalent and electrovalent
-Ionic bonds are electrovalent bonds (groups of charged ions held together by electric forces
)
-Isotope- an atom is missing a neutron or has an extra neutron
http://www.learner.org/interactives/periodic/basics_interactive.html

Atoms Are Building Blocks
Atoms are the basis of chemistry. They are the basis for everything in the Universe. You
should start by remembering that matter is composed of atoms. Atoms and the study of
atoms are a world unto themselves. We're going to cover basics like atomic structure
and bonding between atoms.
Smaller Than Atoms?
Are there pieces of matter that are smaller than atoms?
Sure there are. You'll soon be learning that atoms are
composed of pieces like electrons, protons, and neutrons.
But guess what? There are even smaller particles moving
around in atoms. These super-small particles can be
found inside the protons and neutrons. Scientists have
many names for those pieces, but you may have heard of
nucleons and quarks. Nuclear chemists and physicists
work together at particle accelerators to discover the
presence of these tiny, tiny, tiny pieces of matter.
Even though super-tiny atomic particles exist, you only need to remember the three
basic parts of an atom: electrons, protons, and neutrons. What are electrons, protons,
and neutrons? A picture works best to show off the idea. You have a basic atom. There
are three types of pieces in that atom: electrons, protons, and neutrons. That's all you
have to remember. Three things! As you know, there are
almost 120 known elements in the periodic table. Chemists and physicists haven't
stopped there. They are trying to make new ones in labs every day. The thing that
makes each of those elements different is the number of electrons, protons, and
neutrons. The protons and neutrons are always in the center of the atom. Scientists
call the center region of the atom the nucleus. The
nucleus in a cell is a thing. The nucleus in an atom is a
place where you find protons and neutrons. The electrons
are always found whizzing around the center in areas
called shells or orbitals.
You can also see that each piece has either a "+", "-", or
a "0." That symbol refers to the charge of the particle.
Have you ever heard about getting a shock from a
socket, static electricity, or lightning? Those are all different types of
electric charges. Those charges are also found in tiny particles of matter. The electron
always has a "-", or negative, charge. The proton always has a "+", or positive, charge.
If the charge of an entire atom is "0", or neutral, there are equal numbers of positive and

negative pieces. Neutral means there are equal numbers of electrons and protons. The
third particle is the neutron. It has a neutral charge, also known as a charge of zero. All
atoms have equal numbers of protons and electrons so that they are neutral. If there are
more positive protons or negative electrons in an atom, you have a special atom called
an ion.
Looking at Ions
We haven’t talked about ions before, so let’s get down to basics. The
atomic number of an element, also called a proton number, tells you
the number of protons or positive particles in an atom. A normal atom
has a neutral charge with equal numbers of positive and negative
particles.
That means an atom with a neutral charge is one where the number of
electrons is equal to the atomic number. Ions are atoms with extra
electrons or missing electrons. When you are missing an electron
or two, you have a positive charge. When you have an extra
electron or two, you have a negative charge.
What do you do if you are a sodium (Na) atom? You have eleven
electrons — one too many to have an entire shell filled. You need
to
find another element that will take that electron away from you. When you lose that
electron, you will you’ll have full shells. Whenever an atom has full shells, we say it is
"happy." Let's look at chlorine (Cl). Chlorine has seventeen electrons and only needs one
more to fill its third shell and be "happy." Chlorine will take your extra sodium electron and
leave you with 10 electrons inside of two filled shells. You are now a happy atom too. You
are also an ion and missing one electron. That missing electron gives you a positive
charge. You are still the element sodium, but you are now a sodium ion (Na + ). You have
one less electron than your atomic number.

Ion Characteristics
So now you've become a sodium ion. You have ten electrons.
That's the same number of electrons as neon (Ne). But you
aren't neon. Since you're missing an electron, you aren't
really a complete sodium atom either. As an ion you are now
something completely new. Your whole goal as an atom was
to become a "happy atom" with completely filled electron
shells. Now you have those filled shells. You have a lower
energy. You lost an electron and you are "happy." So what
makes you interesting to other atoms? Now that you have
given up the electron, you are quite electrically attractive.
Other electrically charged atoms (ions) of the opposite charge
(negative) are now looking at you and seeing a good partner to
bond with. That's where the chlorine comes in. It's not only chlorine. Almost any ion with
a negative charge will be interested in bonding with you.
Electrovalence
Don't get worried about the big word. Electrovalence is just another word for something
that has given up or taken electrons and become an ion. If you look at the periodic
table, you might notice that elements on the left side usually become positively charged
ions (cations) and elements on the right side get a negative charge (anions). That trend
means that the left side has a positive valence and the right side has a negative
valence. Valence is a measure of how much an atom wants to bond with other atoms. It
is also a measure of how many electrons are excited about bonding with other atoms.
There are two main types of bonding, covalent and electrovalent. You may have heard
of the term "ionic bonds." Ionic bonds are electrovalent bonds. They are just groups of
charged ions held together by electric forces. When in the presence of other ions, the
electrovalent bonds are weaker because of outside electrical forces and attractions.
Sodium and chlorine ions alone have a very strong bond, but as soon as you put those
ions in a solution with H + (Hydrogen ion), OH - (Hydroxide), F - (Fluorine ion) or Mg ++
(Magnesium ion), there are charged distractions that break the Na-Cl bond.

Look at sodium chloride (NaCl) one more time. Salt is a very strong bond when it is
sitting on your table. It would be nearly impossible to break those ionic/electrovalent
bonds. However, if you put that salt into some water (H2O), the bonds break very
quickly. It happens easily because of the electrical attraction of the water. Now you have
sodium (Na + ) and chlorine (Cl - ) ions floating around the solution. You should remember
that ionic bonds are normally strong, but they are very weak in water.

Neutron Madness
We have already learned that ions are atoms that are
either missing or have extra electrons. Let's say an atom
is missing a neutron or has an extra neutron. That type
of atom is called an isotope. An atom is still the same
element if it is missing an electron. The same goes for
isotopes. They are still the same element. They are just
a little different from every other atom of the same
element.
For example, there are a lot of carbon (C) atoms in the
Universe. The normal ones are carbon-12. Those atoms have 6 neutrons. There are a
few straggler atoms that don't have 6. Those odd ones may have 7 or even 8 neutrons.
As you learn more about chemistry, you will probably hear about carbon-14. Carbon-14
actually has 8 neutrons (2 extra). C-14 is considered an isotope of the element carbon.
Messing with the Mass
If you have looked at a periodic table, you may have noticed that the atomic mass of an
element is rarely an even number. That happens because of the isotopes. If you are an
atom with an extra electron, it's no big deal. Electrons don't have much of a mass when
compared to a neutron or proton.
Atomic masses are calculated by figuring out the
amounts of each type of atom and isotope there are in
the Universe. For carbon, there are a lot of C-12, a
couple of C-13, and a few C-14 atoms. When you
average out all of the masses, you get a number that is
a little bit higher than 12 (the weight of a C-12 atom).
The average atomic mass for the element is actually
12.011. Since you never really know which carbon atom
you are using in calculations, you should use the
average mass of an atom.
Bromine (Br), at atomic number 35, has a greater variety of isotopes. The atomic mass
of bromine (Br) is 79.90. There are two main isotopes at 79 and 81, which average out
to the 79.90amu value. The 79 has 44 neutrons and the 81 has 46 neutrons. While it
won't change the average atomic mass, scientists have made bromine isotopes with
masses from 68 to 97. It's all about the number of neutrons. As you move to higher
atomic numbers in the periodic table, you will probably find even more isotopes for
each element.

Summary :
Atoms are what make up everything in this world. Despite its size, they are also made
of three different charges known as protons, electrons, and neutrons. Protons are the
positive charge of an atom while electrons are the negative charge and neutrons have
a charge of 0. Protons and neutrons are located in the nucleus of an atom and
electrons are found circling around the center in areas called shells or orbitals. In an
element, the protons and neutrons can be found by the atomic number while the
electrons can be found by subtracting the atomic number by the atomic mass. They
may be what make up an atom but there are even smaller particles known as
nucleons and quarks.Nucleons and quarks are located inside the protons but they
arent as important. Protons, electrons, and neutrons are significant and so are ions.
Ions are atoms with extra electrons or missing electrons. So when you are missing an
electron two, you have a positive charge and when you have an extra electron or two,
you have a negative charge. Ions are formed when
something has given up or taken electron.This is called an electrovalence. Valence is
a measure of how much an atom wants to bond with other atoms. its also a measure of
how many electrons are "excited" about bonding with other atoms. There are two
different ways of bonding: covalent and electrovalent. Ionic bonds are electrovalent
bonds which are just groups held together by electrical forces. An atom that is missing
a neutron or has an extra neutron are Isotopes. As you move to higher atomic
numbers in the periodic table. you will probably find even more isotopes for each
element.

Electron Configuration
Color the sublevel:
s = Red
d = Green
p = Blue
f = Orange
Write in sublevels

Write period, sublevel and super scripts.
Ctrl Shift =
gives you super scripts
The Learning Goal for this assignment is:
Describe the quantization of energy at the atomic level.
www.youtube.com/watch?v=jtYzEzykFdg
www.youtube.com/watch?
annotation_id=annotation_2076&feature=iv&src_vid=jtYzEzykFdg&v=cOlac8ruD_0
www.youtube.com/watch?
annotation_id=annotation_570977&feature=iv&src_vid=cOlac8ruD_0&v=lR2vqHZWb5A
Notes Section
Principle Energy Levels- Bohr's Model
1,2,3,4...
Sublevels : s- sharp (holds 2 electrons)
p- principle (holds 6 electrons)
d- diffuse (holds 10 electrons)
f- fundamental (holds 14 electrons)
The number of energy levels is how many sublevels it has.

1s
2s, 2p
3s, 3p, 3d
4s, 4p, 4d, 4f
5s, 5p ,5d, 5f
6s, 6p, 6d, 6f
7s, 7p, 7d, 7f
Order of sublevel energies:
1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10,
5p6, 6s2,4f14, 5d10, 6p6, 7s2, 5f14, 6d10, 7p6
3 rules of electron configuration
Rule #1. Aufbau - nO^e ( n is the energy level, O is the orbital type or
sublevel, e is the number of electrons in that orbital shell)
Rule #2. Hund rule - when electrons occupy orbitals of equal energy, one electron
enters each orbital until all orbitals contain one electron with the
same spin.
Rule #3. Pauli exclusion principle - an orbital conatins a maximum of 2 electrons
and
paired electrons will have opposite spin.
Electron Configuration
In order to write the electron configuration for an atom you must know the 3 rules of
electron configurations.
1.
Notation e
where
nO
n is the energy level
O is the orbital type (s, p, d, or f)
e is the number of electrons in that orbital shell
Principle

electrons will first occupy orbitals of the lowest energy level
2. Hund rule
when electrons occupy orbitals of equal energy, one electron enters each orbital until
all the orbitals contain one electron with the same spin.
3. Pauli exclusion principle
an orbital contains a maximum of 2 electrons and
paired electrons will have opposite spin

In the space below, write the unabbreviated electron configurations of the following elements:
1) sodium ________________________________________________1s^2, 2s^2, 2p^6, 3s^1
2) iron ________________________________________________1s^2, 2s^2, 2p^6, 3s^2, 3p^6,
4s^2, 3d^6
3) bromine ________________________________________________1s^2, 2s^2, 2p^6, 3s^3, 3p^6,
4s^2, 3d^10, 4p^5
4) barium 1________________________________________________s^2, 2s^2, 2p^6, 3s^2, 3p^6, 4s^2,
3d^10, 4p^6, 5s^2, 4d^10, 5p^6, 6s^ 2
5) neptunium 1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 4s^2, 3d^10, 4p^6, 5s^2, 4d^10, 5p^6, 6s^2, 4f^14,
5d^10,________________________________________________
6p^6, 7s^2, 5f^5
In the space below, write the abbreviated electron configurations of the following elements:
6) cobalt [________________________________________________Ar] 4s^2, 3d^7
7) silver ________________________________________________[Kr] 5s^2, 4d^9
8) tellurium ________________________________________________[Kr] 5s^2, 4d^10, 5p^4
9) radium ________________________________________________[Rn] 7s^2
10) lawrencium ________________________________________________[Rn] 7s^2, 5f^14, 6d^1
Determine what elements are denoted by the following electron configurations:
11) 1s²s²2p⁶3s²3p⁴ ____________________Sulfur
12) 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s¹ Rubidium____________________
13) [Kr] 5s²4d¹⁰5p³ ____________________Antimony
14) [Xe] 6s²4f¹⁴5d⁶ Osmium____________________
15) [Rn] 7s²5f¹¹ Einsteinium____________________
Identify the element or determine that it is not a valid electron configuration:
16) 1s²2s²2p⁶3s²3p⁶4s²4d¹⁰4p⁵ Not valid____________________
17) 1s²2s²2p⁶3s³3d⁵ Not valid____________________
18) [Ra] 7s²5f⁸ ____________________Not valid

Calcium[Ca]
Nickel[Ni]
Carbon[C]
Xenon[Xe]
Sulfur[S]
Uranium[U]
All the boxes need to be complete before
moving on to another section
There are two arrows going up in 2s.
[ x ]
[ xo] [xo] [x ][ ][ ]
[ xo] [xo] [xo][xo][xo] [x ]
[ Ar] [xo] [xo][xo][xo][xo][xo] [xo][xo][xo ]
[ Ar] [xo] [x ][x ][x ][x ][ ]
[ He]
[ Ar] [xo]
[xo] [x ][x ][ ]
[ Xe] [xo] [xo][xo][xo][xo][xo][xo][xo] [xo][xo][x ][x
[ Rn] [xo] [x ][x ][x ][x ][x ][x ][ ]
[ He] [xo] [xo][x ][x ]
19) [Kr] 5s²4d¹⁰5p⁵ ____________________Iodine
20) [Xe] ____________________Not valid
1)sodium 1s 2 2s 2 2p 6 3s 1 2)iron 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6
3)bromine 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 4)barium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2
5)neptunium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 5 6)cobalt [Ar] 4s 2 3d 7
7)silver [Kr] 5s 2 4d 9 8)tellurium[Kr] 5s 2 4d 10 5p 4
9)radium [Rn] 7s 2 10)lawrencium[Rn] 7s 2 5f 14 6d 1
1s 2 2s 2 2p 6 3s 2 3p 4 sulfur
[Kr] 5s 2 4d 10 5p 3 antimony
[Rn] 7s 2 5f 11 einsteinium
1s 2 2s 2 2p 6 3s 3 3d 5 not valid (3p comes after 3s)
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 rubidium
[Xe] 6s 2 4f 14 5d 6 osmium
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 4d 10 4p 5 not valid (take a look at “4d”)
[Ra] 7s 2 5f 8 not valid (radium isn’t a noble gas)
[Kr] 5s 2 4d 10 5p 5 valid iodine
20)[Xe] not valid (an element can’t be its own electron configuration)
[Ne] [xo] [x ][x ][x ] [Ar] [x ]
Create groups for these Scientist and explain your groupings (use
the information you got from your research)
Antoine Henri Becquerel

Glenn Seaborg
Marie and Pierre Curie
Dmitri Mendeleev
Lothar Meyer
J.W Dobereiner
All scientists contributed in finding radioactive
elements by radioactivity.
Niels Bohr
Ernest Rutherford
Louis de Barogilie
James Chadwick
All scientists used the quantum theory to identify
their discoveries
J.J Thomson
Hantaro Nagoka
Eugene Goldstein
Erwin Schrodinger
Robert Millikan
All scientists contributed in the discovery of
electrons and cathode rays by using each other's
ideas.
Democritus
John Dalton
Both scientists believed atoms were indestructible
and everything is made up of atoms.
Mendeleev and Meyes worked on the periodic
table together while Dobereiner dicovered triads.
Research the Scientist and summarize their contributions to the Atomic Theory
Antoine Henri Becquerel
- Antoine Henri Becquerel was the first person to discover the existence of radioactivity through his investigations of
uranium and other substances.Becquerel also discoved one type of radioactivity beta particles which is due to high-speed
electrons leaving the nucleus of the atom. Niels Bohr - Niels Bohr suggested a theory for the hydrogen atom based on

quantum theory that energy is transferred only in certain well defined quantities. Bohr's theory was reasonable enough to
explain why atoms emitted light in fixed wavelengths. Louis de Barogilie
- Louis de Barogilie was best known for his discovery of the wave nature of electrons or quantum theory and suggestedthat
all matter has wave properties which is called the Barogilie hypothesis.
Glenn Seaborg
- Glenn Seaborg discovered the element plutonium. He then went on to identify several more of the
radioactive"transuranium" elements. Hantaro Nagaoka
- Hantaro Nagaoka was the first to disagree with J.J Thomson's model. He modified Thomson's model, stating that the
negatively charged electrons were actually located outside the atom, and that they orbit around it. He called this model
the "Saturn model" because it resembled the planet Saturn.
Democritus
-Democritus adopted the theory that Leucippus originally proposed and developed it further. Democritus' atomic theory states
that everything is made up of atoms, which are indestructible, eternal and invisible, small and unable to be diminished. Marie
and Pierre Curie
- Marie and Pierre Curie are best known for their study of radioactivity, which led to their discovery of the elements radium
and polonium.They concluded that radioactivity does not depend on how atoms are arranged into molecules, but rather that it
originates within the atoms themselves.
Eugene Goldstein
- Eugene Goldstein is known for his work on electrical phenomena in gases and on cathode rays. Goldstein called the other
ray that travels from the anode toward the cathode " Canal Rays" because the rays pass through the channels, in the cathode.
Dmitri Mendeleev
- Dmitri Mendeleev is known for his work on the Periodic Law and the create of the first Periodic table. Mendeleev predicted the existence
and properties of new chemical elements by using his periodic table. His dicovery helped many scientists with their discoveries.
J.J. Thomson
- J.J Thomson discoved the electron by experimenting with a cathode ray, tube. He demonstated that cathode rays were
negatively charged. He aslo studied positively charged particles in neon gas.
James Chadwick
- James Chadwick discovered the neutron. He didn't understand Joliot-Curies' experiment so he decided to do his own.
Chadwick drew the conclusion that the protons had actually been hit by the nuetron. Erwin Shrodinger
- Erwin Schrodinger used mathematical equations to describe the likelihood of finding an electron in a certain position. This
atomic model is known as the quantum mechanical model of the atom. John Dalton
- John Dalton's atomic theory said that elements of tiny particles are called atoms. He proposed that all matter was
composed of atoms, indivisible and indestructible building blocks. Lothar Meyer - Lothar Meyer developed a periodic
classification of the chemical elements. Meyer and Mandeleev worked together on the periodic table predicting new elements.
Meyer published a revised and expanded version of Mandeleev's periodic table. Robert Millikan
- Robert Millikan's oil experiment helped to quantify the charge of an electron, which contributed greatly to our understanding
of the structure of the atom and atomis theory. J.W. Dobereiner
- J.W Dobereiner discovered the existence of families of elements with similar chemical properties. He called these families
triads because there always seemed to be three elements. He also discovered the organic cmpound furfural and developed
the seperation of calcium and magnesium.
Ernest Rutherford
- Ernest Rutherford was the first to discover that atoms have a small charged nucleus surrounded by largely empty
space,and are circled by tiny electrons which became known as the Rutherford model of the atom. He is also credited with the
discovery of the proton and hypothesized the existence of the neutron.

The Learning Goal for this Assignment is
Relate properties of atoms and their position in the periodic
table to the arrangement of their electrons.
Alkali Metals
- A group in the s-block of the periodic table consisting of the chemical elements lithium (Li), sodium
(Na), potassium (K), rubidium (Rb),caesium (Cs), and francium (Fr).They are all shiny,soft,highly
reactive metals at a standard temp. and pressure and readily lose their outermost electron to form
cations with charge+1.
Alkali Earth Metals
-Six chemical elements in group 2 of the periodic table consisting of nerylium (Be), magnesium (Mg), calcium
(Ca), strontium (Sr), barium (Ba), and radium (Ra).They are all shiny,silvery-white, somewhat reactive metals
at standard temp. and pressure and readily use to form cations with charge
+2, and an oxidation state (oxidation number) of +2.
Transitional Metals
-Elements in the d-block of the periodic table, which includes groups
3 to 12. Atom has a partially filled d sub shell, or which can give rise to cations. They chemically show
variable valence and and a strong tendency to form coordination compounds and many of their
compounds are colored.
Inter Transitional Metals
-A group in the f-block consisting of chemical elements 57-71 lanthanides and 89-103 actinides on the
periodic table.They have three incomplete outermost electron shells and are all metals. The lanthanides
are very similar and the actinides are all radioactive.
Metals
- About 91 of the 118 elements in the periodic table are metals. It is an element that is
typically hard, opaque, shiny, and has good electrical and thermal conductivity. Examples
of metals are aluminium, copper, iron, tin, gold, lead, silver, titanium, uranium, and zinc. They are elements
that form cations when compounds of it are in solution and oxides of the element from hydroxides rather than
acids in water.
Metalloids
- A group of elements in the p-block that are shiny but brittle, might be dull but conduct electricity, they have
metal and non-metal properties. Consists of the elements boron (B), silicon (Si), germanium (Ge), arsenic (As),
antimony (Sb), tellurium (Te) and polonium (Po).
Non Metals
- Chemical elements that does not have the properties of a metal. Consists of hydrogen (H), carbon (C),
nitrogen(N), oxygen (O), phosphorus (P), sulfur (S) and selenium (Se).

Noble Gases
* A group of elements that are all gases, all are monoatomic which means that each molecule
is a single atom and are totally unreactive. They are in the p- block and consists of helium
(He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn).

Using Wikipedia, define the 8 categories of elements on the
left page.
Color your periodic table similar to the one on
alkali metals
alkaline metals
other metals
pages 168—169 of your book.
metalloids
non metals
halogens
transitional metals
lanthanoids
noble gases
unknown elements
actinoids

Atomic Size
Define Atomic Size: The size of an atom. The number of energy levels and size of the nucleus
determines the size of the atom.
Explanation: Atomic size increases from top to bottom within a group and decreases from left to
right across a period.
Define Ionization Energy:
The amount of energy needed to take an electron away
from a specefic atom.

Ionization Energy
Explanation:
First ionization energy tends to decrease from top to bottom within a group and
increase form left to right across a period.
Electronegativity
Define Electronegativity: The ability of an atom to take an electron away from another atom.

Explanation: Electronegativity values decrease from top to bottom within a group. For
representative elements, the values tend to increase from left to right
across a period.
Define Ion Size:
The size of the ion's atom in comparison to the size of the elemental atom.

Ion Size
Explanation: Ionic size tends to increase from top to bottom within a group. Generally, the size
cations and anions decrease from left to right across a period.

Unit 3
Chapter 25 Nuclear Chemistry
The students will learn what happens when an unstable
nucleus decays and how nuclear chemistry affects their lives.
Explore the theory of electromagnetism by comparing and contrasting the
different parts of the electromagnetic spectrum in terms of wavelength,
frequency, and energy, and relate them to phenomena and applications.
Students will be able to compare and contrast the different parts of the
electromagnetic spectrum.
Students will be able to apply knowledge of the EMS to real world phenomena.
Students will be able to quantitatively compare the relationship between energy,
wavelength, and frequency of the EMS.
amplitude
wavelength
frequency
hertz
electromagnetic radiation
photon
Planck’s constant
Explain and compare nuclear reactions (radioactive decay, fission and
fusion), the energy changes associated with them and their associated
safety issues.
Students will be able to compare and contrast fission and fusion reactions.
Students will be able to complete nuclear decay equations to identify the type of
decay.
Students will participate in activities to calculate half-life.
radioactivity
nuclear radiation
alpha particle
beta particle
gamma ray
positron
½ life
transmutation
fission
fusion

Chapter 7
Ionic and Metallic Bonding
The students will learn how ionic compounds form and how
metallic bounding affects the properties of metals.
Compare the magnitude and range of the four fundamental forces
(gravitational, electromagnetic, weak nuclear, strong nuclear).
Students will compare/contrast the characteristics of each fundamental force.
gravity
electromagnetic
strong
weak
Distinguish between bonding forces holding compounds together and other
attractive forces, including hydrogen bonding and van der Waals forces.
Students will be able to compare/contrast traits of ionic and covalent bonds.
Students will be able to compare/contrast basic attractive forces between
molecules.
Students will be able to predict the type of bond or attractive force between
atoms or molecules.
ionic bond
covalent bond
metallic bond
polar covalent bond
hydrogen bond
van der Waals forces
London dispersion forces
Chapter 8
Covalent Bonding
The students will learn how molecular bonding is different
than ionic bonding and electrons affect the shape of a
molecule and its properties.
Interpret formula representations of molecules and compounds in terms of
composition and structure.
Students will be able to interpret chemical formulas in terms of # of atoms.
Students will be able to differentiate between ionic and molecular compounds.
Students will be able to list various VSEPR shapes and identify examples of
each.
Students will be able to predict shapes of various compounds.
Molecule
empirical formula
Atom
Electron
Element
Compound

jamyled vergara
Name ____________________
Go to the web site www.darvill.clara.net/emag
1. Click on “How the waves fit into the spectrum” and fill in this table:
>: look out for the
RED words on the web site!
Low __________, frequencyLong wavelength
High frequency, Short ______________
wavelength
Radio Waves
microwaves infrared visible light ultra-violet x-rays
Gamma rays
2. Click on “Radio waves”. They are used for _______________________
communication
3. Click on “Microwaves”. They are used for cooking, mobile _________, phones _______ speed cameras and _________. radar
4. Click on “Infra-red”. These waves are given off by _____ hot _________. objects They are used for remote controls,
cameras in police ____________ helicopters , and alarm systems.
5. Click on “Visible Light”. This is used in ___ cd players and _______ laser printers, and for seeing where we’re going.
6. “UV” stands for “ ________ ultra ___________”. violet This can damage the _________ retina in your eyes, and cause
sunburn and even _______ skin cancer. Its uses include detecting forged ______ bank _______. notes
7. X-rays are used to see inside people, and for _________ airport security.
8. Gamma rays are given off by some ________________ radioactive substances. We can use them to kill ________ cancer cells,
which is called R_______________ adiotherapy .
9. My Quiz score is ____%. 100

10. Name ________________________________
Go to the web site www.darvill.clara.net/emag
Name How they’re made Uses Dangers
Gamma rays
radiowaves
they are made by various types
of transimiters
they are mainly used for
communication
they can cause
cancer, leukemia, and
other disorders
_____ Frequency _____ frequency,
Short wavelength ______ Wavelength

Learning Goal for this section:
Explain and compare nuclear reactions(radioactive decay,fission, and fusion),
the energy changes associated with them and their associated safety issues.
Notes Section:
*Alpha particles go down by two. is a helium. 4/2 He
*Beta particles look like an electron
(e-)
mass stays the same
converting a neutron to a proton.
up 1
positron goes down by one
*Gamma is all energy
gamma radiation is a high frequency
that goes through anything.
*Half life take substance you started
with and

The Nucleus
A typical model of the atom is called the Bohr Model, in
honor of Niels Bohr who proposed the structure in 1913. The Bohr atom consists of a central nucleus
composed of neutrons and protons, which is surrounded by electrons which “orbit” around the nucleus.
Protons carry a positive charge of one and have a mass of about 1 atomic mass unit or amu (1 amu =1.7x10-
27 kg, a very, very small number). Neutrons are electrically “neutral” and also have a mass of about 1 amu. In
contrast electron carry a negative charge and have mass of only 0.00055 amu. The number of protons in a
nucleus determines the element of the atom. For example, the number of protons in uranium is 92 and the
number in neon is 10. The proton number is often referred to as Z.
Atoms with different numbers of protons are called elements, and are arranged in the periodic table with
increasing Z.
Atoms in nature are electrically neutral so the number of electrons orbiting the nucleus equals the number of
protons in the nucleus.
Neutrons make up the remaining mass of the nucleus and provide a means to “glue” the protons in place.
Without neutrons, the nucleus would split apart because the positive protons would repel each other. Elements
can have nucleii with different numbers of neutrons in them. For example hydrogen, which normally only has
one proton in the nucleus, can have a neutron added to its nucleus to from deuterium, ir have two neutrons
added to create tritium, which is radioactive. Atoms of the same element which vary in neutron number are
called isotopes. Some elements have many stable isotopes (tin has 10) while others have only one or two. We
express isotopes with the nomenclature Neon-20 or 20 Ne 10, with twenty representing the total number of
neutrons and protons in the atom, often referred to as A, and 10 representing the number of protons (Z).
Alpha Particle
Decay
Alpha decay is a radioactive process in which a
particle with two neutrons and two protons is
ejected from the nucleus of a radioactive atom. The particle is identical to the nucleus of a helium atom.
Alpha decay only occurs in very heavy elements such as uranium, thorium and radium. The nuclei of these
atoms are very “neutron rich” (i.e. have a lot more neutrons in their nucleus than they do protons) which makes
emission of the alpha particle possible.
After an atom ejects an alpha particle, a new parent atom is formed which has two less neutrons and two less
protons. Thus, when uranium-238 (which has a Z of 92) decays by alpha emission, thorium-234 is created
(which has a Z of 90).
Because alpha particles contain two protons, they have a positive charge of two. Further, alpha particles are
very heavy and very energetic compared to other common types of radiation. These characteristics allow alpha
particles to interact readily with materials they encounter, including air, causing many ionizations in a very short
distance. Typical alpha particles will travel no more than a few centimeters in air and are stopped by a sheet of
paper.

Beta Particle Decay
Beta decay is a radioactive process in which an electron is emitted from the nucleus of a radioactive
atom Because this electron is from the nucleus of the atom, it is called a beta particle to distinguish it
from the electrons which orbit the atom.
Like alpha decay, beta decay occurs in isotopes which are “neutron rich” (i.e. have a lot more
neutrons in their nucleus than they do protons). Atoms which undergo beta decay are located below
the line of stable elements on the chart of the nuclides, and are typically produced in nuclear reactors.
When a nucleus ejects a beta particle, one of the neutrons in the nucleus is transformed into a proton.
Since the number of protons in the nucleus has changed, a new daughter atom is formed which has
one less neutron but one more proton than the parent. For example, when rhenium-187 decays
(which has a Z of 75) by beta decay, osmium-187 is created (which has a Z of 76). Beta particles
have a single negative charge and weigh only a small fraction of a neutron or proton. As a result, beta
particles interact less readily with material than alpha particles. Depending on the beta particles
energy (which depends on the radioactive atom), beta particles will travel up to several meters in air,
and are stopped by thin layers of metal or plastic.
Positron emission or beta plus decay (β+ decay) is a subtype of radioactive decay called beta decay,
in which a proton inside a radionuclide nucleus is converted into a neutron while releasing a positron
and an electron neutrino (νe). Positron emission is mediated by the weak force.
An example of positron emission (β+ decay) is shown with magnesium-23 decaying into sodium-23:
23 Mg12 → 23 Na11 + e +
Because positron emission decreases proton number relative to neutron number, positron decay
happens typically in large "proton-rich" radionuclides. Positron decay results in nuclear transmutation,
changing an atom of one chemical element into an atom of an element with an atomic number that is
less by one unit.
Positron emission should not be confused with electron emission or beta minus decay (β− decay),
which occurs when a neutron turns into a proton and the nucleus emits an electron and an
antineutrino.

Gamma
Radiation
After a decay reaction, the nucleus is often in an
“excited” state. This means that the decay has
resulted in producing a nucleus which still has
excess energy to get rid of. Rather than emitting another beta or alpha particle, this energy is lost by
emitting a pulse of electromagnetic radiation called a gamma ray. The gamma ray is identical in
nature to light or microwaves, but of very high energy.
Like all forms of electromagnetic radiation, the gamma ray has no mass and no charge. Gamma rays
interact with material by colliding with the electrons in the shells of atoms. They lose their energy
slowly in material, being able to travel significant distances before stopping. Depending on their initial
energy, gamma rays can travel from 1 to hundreds of meters in air and can easily go right through
people.
It is important to note that most alpha and beta emitters also emit gamma rays as part of their decay
process. However, their is no such thing as a “pure” gamma emitter. Important gamma emitters
including technetium-99m which is used in nuclear medicine, and cesium-137 which is used for
calibration of nuclear instruments.
Half Life
Half-life is the time required for the quantity of a
radioactive material to be reduced to one-half its
original value.
All radionuclides have a particular half-life, some
of which a very long, while other are extremely
short. For example, uranium-238 has such a
long half life, 4.5x109 years, that only a small fraction has decayed since the earth was formed. In
contrast, carbon-11 has a half-life of only 20 minutes. Since this nuclide has medical applications, it
has to be created where it is being used so that enough will be present to conduct medical studies.

The Learning Goal for this assignment is:
Distinguish between bonding forces holding
compounds together and other attractive
forces, including hydrogen bonding and van
der Waals forces.
Introduction to Ionic Compounds
Those molecules that consist of charged ions with opposite charges are called IONIC. These ionic
compounds are generally solids with high melting points and conduct electrical current. Ionic
compounds are generally formed from metal and a non-metal elements. See Ionic Bonding below.
Ionic Compound Example
For example, you are familiar with the fairly benign unspectacular behavior of common white
crystalline table salt (NaCl). Salt consists of positive sodium ions (Na + ) & negative chloride ions (Cl - ).
On the other hand the element sodium is a silvery gray metal composed of neutral atoms which react
vigorously with water or air. Chlorine as an element is a neutral greenish-yellow, poisonous, diatomic
gas (Cl2).
The main principle to remember is that ions are completely different in physical and chemical
properties from the neutral atoms of the elements.
The notation of the + and - charges on ions is very important as it conveys a definite meaning.
Whereas elements are neutral in charge, IONS have either a positive or negative charge depending
upon whether there is an excess of protons (positive ion) or excess of electrons (negative ion).
Formation of Positive Ions
Metals usually have 1-4 electrons in the outer energy level. The electron arrangement of a rare gas is
most easily achieved by losing the few electrons in the newly started energy level. The number of
electrons lost must bring the electron number "down to" that of a prior rare gas.
How will sodium complete its octet?
First examine the electron arrangement of the atom. The atomic number is eleven, therefore, there
are eleven electrons and eleven protons on the neutral sodium atom. Here is the Bohr diagram and
Lewis symbol for sodium:

This analysis shows that sodium has only one electron in its outer level. The nearest rare gas is neon
with 8 electron in the outer energy level. Therefore, this electron is lost so that there are now eight
electrons in the outer energy level, and the Bohr diagrams and Lewis symbols for sodium ion and
neon are identical. The octet rule is satisfied.
Ion Charge?
What is the charge on sodium ion as a result of losing one electron? A comparison of the atom and
the ion will yield this answer.
Sodium Atom
Sodium Ion
11 p+ to revert to 11 p + Protons are identical in
12 n an octet 12 n
the atom and ion.
Positive charge is
11 e- lose 1 electron 10 e-
caused by lack of
0 charge + 1 charge
electrons.
Formation of Negative Ions
How will fluorine complete its octet?
First examine the electron arrangement of the atom. The atomic number is nine, therefore, there are
nine electrons and nine protons on the neutral fluorine atom. Here is the Bohr diagram and Lewis
symbol for fluorine:
This analysis shows that fluorine already has seven electrons in its outer level. The nearest rare gas
is neon with 8 electron in the outer energy level. Therefore only one additional electron is needed to
complete the octet in the fluorine atom to make the fluoride ion. If the one electron is added, the Bohr
diagrams and Lewis symbols for fluorine and neon are identical. The octet rule is satisfied.

Ion Charge?
What is the charge on fluorine as a result of adding one electron? A comparison of the atom and the
ion will yield this answer.
Fluorine Atom Fluoride Ion *
9 p+ to complete 9 p + Protons are identical in
10 n octet 10 n
9 e- add 1 electron 10 e-
0 charge - 1 charge
the atom and ion.
Negative charge is
caused by excess
electrons
* The "ide" ending in the name signifies a simple negative ion.
Summary Principle of Ionic Compounds
An ionic compound is formed by the complete transfer of electrons from a metal to a nonmetal and
the resulting ions have achieved an octet. The protons do not change. Metal atoms in Groups 1-3
lose electrons to non-metal atoms with 5-7 electrons missing in the outer level. Non-metals gain 1-4
electrons to complete an octet.
Octet Rule
Elemental atoms generally lose, gain, or share electrons with other atoms in order to achieve the
same electron structure as the nearest rare gas with eight electrons in the outer level.
The proper application of the Octet Rule provides valuable assistance in predicting and explaining
various aspects of chemical formulas.
Introduction to Ionic Bonding
Ionic bonding is best treated using a simple
electrostatic model. The electrostatic model
is simply an application of the charge
principles that opposite charges attract and
similar charges repel. An ionic compound
results from the interaction of a positive and
negative ion, such as sodium and chloride in
common salt.
The IONIC BOND results as a balance
between the force of attraction between
opposite plus and minus charges of the ions
and the force of repulsion between similar
negative charges in the electron clouds. In
crystalline compounds this net balance of
forces is called the LATTICE ENERGY.
Lattice energy is the energy released in the
formation of an ionic compound.
DEFINITION: The formation of an IONIC
BOND is the result of the transfer of one or
more electrons from a metal onto a nonmetal.

Metals, with only a few electrons in the outer energy level, tend to lose electrons most readily. The
energy required to remove an electron from a neutral atom is called the IONIZATION POTENTIAL.
Energy + Metal Atom ---> Metal (+) ion + e-
Non-metals, which lack only one or two electrons in the outer energy level have little tendency to lose
electrons - the ionization potential would be very high. Instead non-metals have a tendency to gain
electrons. The ELECTRON AFFINITY is the energy given off by an atom when it gains electrons.
Non-metal Atom + e- --- Non-metal (-) ion + energy
The energy required to produce positive ions (ionization potential) is roughly balanced by the energy
given off to produce negative ions (electron affinity). The energy released by the net force of
attraction by the ions provides the overall stabilizing energy of the compound.
Notes Section:
An ion is a charged pariticle.
metals and non metals form an ionic bond3

The Learning Goal for this assignment is:
Introduction to Covalent Bonding:
Bonding between non-metals consists of two electrons shared between two atoms. Using the Wave
Theory, the covalent bond involves an overlap of the electron clouds from each atom. The electrons
are concentrated in the region between the two atoms. In covalent bonding, the two electrons shared
by the atoms are attracted to the nucleus of both atoms. Neither atom completely loses or gains
electrons as in ionic bonding.
There are two types of covalent bonding:
1. Non-polar bonding with an equal sharing of electrons.
2. Polar bonding with an unequal sharing of electrons. The number of shared electrons depends on
the number of electrons needed to complete the octet.
NON-POLAR BONDING results when two identical non-metals equally share electrons between
them. One well known exception to the identical atom rule is the combination of carbon and hydrogen
in all organic compounds.
Hydrogen
The simplest non-polar covalent molecule is hydrogen. Each hydrogen
atom has one electron and needs two to complete its first energy level.
Since both hydrogen atoms are identical, neither atom will be able to
dominate in the control of the electrons. The electrons are therefore
shared equally. The hydrogen covalent bond can be represented in a
variety of ways as shown here:
The "octet" for hydrogen is only 2 electrons since the nearest rare gas is
He. The diatomic molecule is formed because individual hydrogen atoms
containing only a single electron are unstable. Since both atoms are
identical a complete transfer of electrons as in ionic bonding is
impossible.
Instead the two hydrogen atoms SHARE both electrons equally.
Oxygen
Molecules of oxygen, present in about 20% concentration in air are
also covalent molecules. See the graphic on the left of the Lewis Dot
Structure.
There are 6 electrons in the outer shell, therefore, 2 electrons are
needed to complete the octet. The two oxygen atoms share a total of
four electrons in two separate bonds, called double bonds.
The two oxygen atoms equally share the four electrons.

POLAR BONDING results when two different non-metals unequally share electrons between them.
One well known exception to the identical atom rule is the combination of carbon and hydrogen in all
organic compounds.
The non-metal closer to fluorine in the Periodic Table has a greater tendency to keep its own electron
and also draw away the other atom's electron. It is NOT completely successful. As a result, only
partial charges are established. One atom becomes partially positive since it has lost control of its
electron some of the time. The other atom becomes partially negative since it gains electron some of
the time.
Hydrogen Chloride
Hydrogen Chloride forms a polar covalent molecule. The graphic
on the left shows that chlorine has 7 electrons in the outer shell.
Hydrogen has one electron in its outer energy shell. Since 8
electrons are needed for an octet, they share the electrons.
However, chlorine gets an unequal share of the two electrons,
although the electrons are still shared (not transferred as in ionic
bonding), the sharing is unequal. The electrons spends more of the
time closer to chlorine. As a result, the chlorine acquires a "partial"
negative charge. At the same time, since hydrogen loses the
electron most - but not all of the time, it acquires a "partial" charge.
The partial charge is denoted with a small Greek symbol for delta.
Water
Water, the most universal compound on all of the earth, has the property of
being a polar molecule. As a result of this property, the physical and
chemical properties of the compound are fairly unique.
Dihydrogen Oxide or water forms a polar covalent molecule. The graphic on
the left shows that oxygen has 6 electrons in the outer shell. Hydrogen has
one electron in its outer energy shell. Since 8 electrons are needed for an
octet, they share the electrons.
Notes Section:
1) count the valence e-
2)find central atom and bond other atoms to it. Submit bonds from total and put lone pairs and
add to the terminal atoms and double and triple bonds as needed.
3)find formal charges. Try to get it as close to zero as possible.

C 2 H 6 O Ethanol CH 3 CH 2 O
Step 1
Find valence e- for all atoms. Add them together.
C: 4 x 2 = 8
H: 1 x 6 = 6
O: 6
Total = 20
Step 2
Find octet e- for each atom and add them together.
C: 8 x 2 = 16
H: 2 x 6 = 12
O: 8
Total = 36
Step 3
Subtract Step 1 total from Step 2.
Gives you bonding e-.
36 – 20 = 16e-
Step 4
Find number of bonds by diving the number in step 3 by 2
(because each bond is made of 2 e-)
16e- / 2 = 8 bond pairs
These can be single, double or triple bonds.
Step 5
Determine which is the central atom
Find the one that is the least electronegative.
Use the periodic table and find the one farthest
away from Fluorine or
The one that only has 1 atom.

Step 6
Put the atoms in the structure that you think it will
have and bond them together.
Put Single bonds between atoms.
Step 7
Find the number of nonbonding (lone pairs) e-.
Subtract step 3 number from step 1.
20 – 16 = 4e- = 2 lone pairs
Step 8
Complete the Octet Rule by adding the lone
pairs.
Then, if needed, use any lone pairs to make
double and triple bonds so that all atoms meet
the Octet Rule.
See Step 4 for total number of bonds.

Linear
Molecular Geometry
Orbital Equation Lone Pairs Angle
sp Ax 2 none 180
BeCl 2
Cl
Be
Cl
Beryllium Di
chloride
element bond lone pair
C

Trigonal planar
Molecular Geometry
Orbital Equation Lone Pairs Angle
sp 2 Ax 3 none 120
BF 3
Boron trifluoride
F
B
F
F
element bond lone pair
C

Bent
Molecular Geometry
Orbital Equation Lone Pairs Angle
sp 3 AX 2 E 2 2 104.5
OF 2
O
F
F
trioxyde
element bond lone pair
C

tretrahedral
Molecular Geometry
Orbital Equation Lone Pairs Angle
sp 3 AX 4 none 109.5
PO 4
3-
O
O
p
O
phosphate
O
element bond lone pair
C

Trigonal pyramidal
Molecular Geometry
Orbital Equation Lone Pairs Angle
Sp 3 AX 3 E 1 107
PH 3
Phosphorus trihydride
H
P
H
H
element bond lone pair
C

Bent
Molecular Geometry
Orbital Equation Lone Pairs Angle
sp 2 AX 2 1 116
O 3
Oxygen trifluoride
O O
O
element bond lone pair
C

Trigonal Bi pyramidal
Molecular Geometry
Orbital Equation Lone Pairs Angle
SP AX 5 none 120/ 90
PCl 5
Phosphorus
Cl
Cl
pentachloride
Cl
P
cl
Cl
element bond lone pair
C

T-shaped
Molecular Geometry
Orbital Equation Lone Pairs Angle
sp 3 d AX 3 E 2 2 90
ClF 3
Chlorine Trifluoride
F
Cl
F
F
element bond lone pair
C

octahedral
Molecular Geometry
Orbital Equation Lone Pairs Angle
Sp 3 d 2 AX 6 none 90
Sulfur hexafluoride
SF 6
F
F
F
S
F
F
F
element bond lone pair
C

Square Planar
Molecular Geometry
Orbital Equation Lone Pairs Angle
Sp 3 d 2 AX 4 E 2 2 90
ICl 4
-
Iodine tetrachloride
Cl
Cl
I
Cl
Cl
element bond lone pair
C

Name Formula Charge
Dichromate Cr₂O₇ 2-
Sulfate SO₄ 2-
Hydrogen Carbonate HCO₃ 1-
Hypochlorite ClO 1-
Phosphate PO₄ 3-
Nitrite NO₂ 1-
Chlorite ClO₂ 1-
Dihydrogen phosphate H₂PO₄ 1-
Chromate CrO₄ 2-
Carbonate CO₃ 2-
Hydroxide OH 1-
Hydrogen phosphate HPO₄ 2-
Ammonium NH₄ 1+
Acetate C₂H₃O₂ 1-
Perchlorate ClO₄ 1-
Permanganate MnO₄ 1-
Chlorate ClO₃ 1-
Hydrogen Sulfate HSO₄ 1-
Phosphite PO₃ 3-
Sulfite SO₃ 2-
Silicate SiO₃ 2-
Nitrate NO₃ 1-
Hydrogen Sulfite HSO₃ 1-
Oxalate C₂O₄ 2-
Cyanide CN 1-
Hydronium H₃O 1+
Thiosulfate S₂O₃ 2-

Chapter 9
Unit 4
Chemical Names and Formulas
The students will learn how the periodic table helps them
determine the names and formulas of ions and compounds.
Chapter 22 Hydrocarbon Compounds
The student will learn how Hydrocarbons are named and the
general properties of Hydrocarbons.
Describe how different natural resources are produced and how their rates
of use and renewal limit availability.
Students will explore local, national, and global renewable and nonrenewable
resources.
Students will explain the environmental costs of the use of renewable and
nonrenewable resources.
Students will explain the benefits of renewable and nonrenewable resources.
Nuclear reactors
Natural gas
Petroleum
Refining
Coal

Chapter 23 Functional Groups
The student will learn what effects functional groups have on
organic compounds and how chemical reactions are used in
organic compounds.
Describe the properties of the carbon atom that make the diversity of carbon
compounds possible.
Identify selected functional groups and relate how they contribute to
properties of carbon compounds.
Students will identify examples of important carbon based molecules.
Students will create 2D or 3D models of carbon molecules and explain why this
molecule is important to life.
covalent bond
single bond
double bond
triple bond
monomer
polymer

LG: describe the properties of the carbon atom that make the diversity of carbon compounds possible.
identify selected functional groups and relate how they contribute to properties of carbon compounds.
http://www.bbc.co.uk/education/guides/zm9hvcw/revision
Notes
*homologous: a family of hydrocarbons with similar chemical properties who
share the same general formula
*First group is the alkanes
* the general formula of alkanes CnH2n+2
Naming branched chain alkanes
*the longest unbranched chain containing the functional group is the parent
molecule
*indicate the functional group with a number
*name the branches and indicate how many there are
*alkenes don't contain a functional group and so the branches are numbered
from the end that gives the lowest set of position numbers for the branches
Second group the Alkenes
*alkenes contain a carbon to carbon double bond which makes them more
reactive than alkanes
*The alkenes general formula is C n H 2n
*Can be used as fuels, solvents, and starting products
*same as naming alkanes but the double carbon must be identified
The last group is the Cycloalkanes
*same general formulas as alkenes but contain only single bonds
*combustion reaction: all hydrocarbons can undergo combustion reactions
with oxygen and give off the same two products.
Only alkenes can participate in addition reactions due to their double bond
*isomers are compounds with the same molecular of fomulae but not
structure

Unit 5
Chapter 10 Chemical Quantities
The student will learn why the mole is important and how the
molecular formula of a compound can be determined
experimentally.
Chapter 11 Chemical Reactions
The students will learn how chemical reactions obey the law of
conservation of mass and how they can predict the products
of a chemical reaction.
Characterize types of chemical reactions, for example: redox, acid-base,
synthesis, and single and double replacement reactions.
Students will be able to identify the type of chemical reaction that occurs.
Students will be able to compare/contrast reactants and products of various
types of chemical reactions.
Students will be able to predict the product of various reactants.
Students will be able to write balanced chemical equations for each type of
reaction.
Decomposition
Combustion
Redox
Acid-Base
Synthesis
single-replacement
double-replacement
Differentiate between chemical and nuclear reactions.
Students will compare/contrast chemical and nuclear reactions.
fission
fusion

Chapter 12 Stoichiometry
The students will learn how balanced chemical equations are
used in stoichiometric calculations and how to calculate
amounts of reactants and products in a chemical equation.
Apply the mole concept and the law of conservation of mass to calculate
quantities of chemicals participating in reactions.
Students will be able to use a balanced equation to determine mole ratios.
Students will be able to apply law of conservation of mass to chemical equations.
Students will be able to calculate empirical and molecular formulas.
Students will be able to calculate the % composition of a compound.
Students will be able to calculate theoretical yield.
Students will be able to calculate % error.
Students will be able to calculate molar mass.
Students will be able to perform stoichiometric calculations, including limiting
reagents.
mole
Avogadro’ s number
molar mass
gram formula mass

www.youtube.com/watch?v=AsqEkF7hcII
www.youtube.com/watch?v=tEn0N4R2dqA
www.youtube.com/watch?v=Pft2CASl0M0
www.youtube.com/watch?v=rwhJklbK8R0
The mole is Avogadro's number =6.02x10 23
convert 5.72g of c 6 h 12 o 6
180
The Mole

www.youtube.com/watch?v=BTRm8PwcZ3U
www.youtube.com/watch?v=F9NkYSKJifs
www.youtube.com/watch?v=xPdqEX_WMjo
Molar Mass

Categories of Reactions
All chemical reactions can be placed into one of six categories. Here they are, in no
particular order:
1) Synthesis: A synthesis reaction is when two or more simple compounds combine to form a
more complicated one. These reactions come in the general form of:
reactants--->products
A+B--->AB
One example of a synthesis reaction is the combination of iron and sulfur to form iron (II) sulfide:
8 Fe + S8 ---> 8 FeS
If two elements or very simple molecules combine with each other, it’s probably a synthesis reaction.
The products will probably be predictable using the octet rule to find charges.
2) Decomposition: A decomposition reaction is the opposite of a synthesis reaction - a
complex molecule breaks down to make simpler ones. These reactions come in the general form:
AB---> A+B
One example of a decomposition reaction is the electrolysis of water to make oxygen and hydrogen
gas:
2 H2O ---> 2 H2 + O2
If one compound has an arrow coming off of it, it’s probably a decomposition reaction. The products
will either be a couple of very simple molecules, or some elements, or both.
3) Single displacement: This is when one element trades places with another element in a
compound. These reactions come in the general form of:
A+BC---> B+AC
One example of a single displacement reaction is when magnesium replaces hydrogen in water to
make magnesium hydroxide and hydrogen gas:
Mg + 2 H2O ---> Mg(OH)2 + H2
If a pure element reacts with another compound (usually, but not always, ionic), it’s probably a single
displacement reaction. The products will be the compounds formed when the pure element switches
places with another element in the other compound.
Important note: these reactions will only occur if the pure element on the reactant side of the equation
is higher on the activity series than the element it replaces.

4) Double displacement: This is when the anions and cations of two different molecules
switch places, forming two entirely different compounds. These reactions are in the general form:
AB+CD--->AD+CB
One example of a double displacement reaction is the reaction of lead (II) nitrate with potassium
iodide to form lead (II) iodide and potassium nitrate:
Pb(NO3)2 + 2 KI ---> PbI2 + 2 KNO3
If two ionic compounds combine, it’s probably a double displacement reaction. Switch the cations
and balance out the charges to figure out what will be made.
Important note: These reactions will only occur if both reactants are soluble in water and only one
product is soluble in water.
5) Acid-base: This is a special kind of double displacement reaction that takes place when an
acid and base react with each other. The H + ion in the acid reacts with the OH - ion in the base,
causing the formation of water. Generally, the product of this reaction is some ionic salt and water:
HA+BOH--->AB+H 2 O
One example of an acid-base reaction is the reaction of hydrobromic acid (HBr) with sodium
hydroxide:
HBr + NaOH ---> NaBr + H2O
If an acid and a base combine, it’s an acid-base reaction. The products will be an ionic compound
and water.
6) Combustion: A combustion reaction is when oxygen combines with another compound to
form water and carbon dioxide. These reactions are exothermic, meaning they produce heat. An
example of this kind of reaction is the burning of napthalene: Organic Cbase+ O 2 --> H 2 O + CO
C10H8 + 12 O2 ---> 10 CO2 + 4 H2O
If something that has carbon and hydrogen reacts with oxygen, it’s probably a combustion reaction.
The products will be CO2 and H2O.
Follow this series of questions. When you can answer "yes" to a question, then
stop!
1) Does your reaction have two (or more) chemicals combining to form one chemical? If yes, then it's
a synthesis reaction
2) Does your reaction have one large molecule falling apart to make several small ones? If yes, then
it's a decomposition reaction
3) Does your reaction have any molecules that contain only one element? If yes, then it's a single
displacement reaction
4) Does your reaction have water as one of the products? If yes, then it's an acid-base reaction
5) Does your reaction have oxygen as one of it's reactants and carbon dioxide and water as
products? If yes, then it's a combustion reaction
6) If you haven't answered "yes" to any of the questions above, then you've got a double
displacement reaction.

List what type the following reactions are:
1) NaOH + KNO3 --> NaNO3 + KOH
Double Displacement
2) CH4 + 2 O2 --> CO2 + 2 H2O
Combustion
3) 2 Fe + 6 NaBr --> 2 FeBr3 + 6 Na
single displacement
4) CaSO4 + Mg(OH)2 --> Ca(OH)2 + MgSO4
Double Displacement
5) NH4OH + HBr --> H2O + NH4Br
Acid base
6) Pb + O2 --> PbO2
7) Na2CO3 --> Na2O + CO2
synthesis
Decomposition

Determine the Type of Reaction for each equation.
Then predict the products of each of the following chemical reactions. If a reaction will not occur,
explain why not.
Then Balance the equation.
1) __Ag2SO4 + 2 __NaNO3 →
2AgNo 3 + Na 2 SO 4
2)__NaI 2 + __CaSO4 →
Na 2 SO 4 +CaI 2
3) __HNO3 + __Ca(OH)2 →
Ca(NO 3 ) + 2 H 2 O
4) __CaCO3 →
CaO+ CO 2
Ca=1 Ca=1
C=1 C=1
O=3 O=3
5) __AlCl3 + __(NH4)PO4 →
Al=1
Cl=3
N=3
P=1
O=4
6) __Pb + __Fe(NO3)3 H=12 →
No reaction
AlPO 4 + 3NH 4 Cl
Al=1
CL=1X3
N=lx3
P=1
O=4
H=4 =12
7) __C3H6 + 4.5 __O2 →
c=3
H=6
O=2
3H 2 O+ 3CO 2
C=1=3
H=2=6
O=3=7=9
8) __Na + __CaSO4 →

How to Balance Chemical Equations
A chemical equation is a theoretical or written representation of what happens during a chemical
reaction. The law of conservation of mass states that no atoms can be created or destroyed in a
chemical reaction, so the number of atoms that are present in the reactants has to balance the
number of atoms that are present in the products. Follow this guide to learn how to balance chemical
equations.
Step 1
Write down your given equation. For this example, we will use:
C3H8 + O2 --> H2O + CO2
Step 2
Write down the number of atoms that you have on each side of the equation. Look at the subscripts
next to each atom to find the number of atoms in the equation.
Left side: 3 carbon, 8 hydrogen and 2 oxygen
Right side: 1 carbon, 2 hydrogen and 3 oxygen

Step 3
Always leave hydrogen and oxygen for last. This means that you will need to balance the carbon
atoms first.
Step 4
Add a coefficient to the single carbon atom on the right of the equation to balance it with the 3 carbon
atoms on the left of the equation.
C3H8 + O2 --> H2O + 3CO2
The coefficient 3 in front of carbon on the right side indicates 3 carbon atoms just as the subscript 3
on the left side indicates 3 carbon atoms.
In a chemical equation, you can change coefficients, but you should never alter the subscripts.

Step 5
Balance the hydrogen atoms next. You have 8 on the left side, so you'll need 8 on the right side.
C3H8 + O2 --> 4H2O + 3CO2
On the right side, we added a 4 as the coefficient because the subscript showed that we already
had 2 hydrogen atoms.
When you multiply the coefficient 4 times the subscript 2, you end up with 8.
Step 6
Finish by balancing the oxygen atoms.
Because we've added coefficients to the molecules on the right side of the equation, the number of
oxygen atoms has changed. We now have 4 oxygen atoms in the water molecule and 6 oxygen
atoms in the carbon dioxide molecule. That makes a total of 10 oxygen atoms.
Add a coefficient of 5 to the oxygen molecule on the left side of the equation. You now have 10
oxygen molecules on each side.
C3H8 + 5O2 --> 4H2O + 3CO2.
The carbon, hydrogen and oxygen atoms are balanced. Your equation is complete.

1) ___ 2 NaNO3 + ___ PbO ___ Pb(NO3)2 + ___ Na2O
left
Na=1x2 =2
N=1=2
O=4=7
Pb=1=1
right
Na=2
N=2
O=7
Pb=1
2) ___ 6 AgI + ___ Fe2(CO3)3 ___ 2 FeI3 + ___ Ag2CO3
Ag=1
I=1
Fe=2
C=1
O=9
Ag=2
I=3=6
Fe=1=2
C=1
O=3
3) ___ C2H4O2 + ___ O2 ___ 2 CO2 + ___ H2O
C=2
H=4
O=4
C=1x2=2
H=2x2=4
O=3+2=5
4) ___ ZnSO4 + ___ Li2CO3 ___ ZnCO3 + ___ Li2SO4
Zn=1
Zn=1
S=1
S=1
O=6
O=7
Li=2
Li=2
C=1
C=1
5)
5
___ V2O5 + ___ CaS ___
5
CaO + ___ V2S5
V=2
V=2
O=5
O=1x5
Ca=1x5=5
Ca=1x5=5
S=1x5
S=5

6) ___ Mn(NO2)2 + ___ BeCl2 ___ Be(NO2)2 + ___ MnCl2
7) ___ AgBr + ___ GaPO4 ___ Ag3PO4 + ___ GaBr3
8) ___ H2SO4 + ___ B(OH)3 __ B2(SO4)3 + ___ H2O
9) ___ S8 + ___ O2 ___ SO2
10) ___ Fe + ___ AgNO3 ___ Fe(NO3)2 + ___ Ag

1) 2 NaNO3 + PbO Pb(NO3)2 + Na2O
2) 6 AgI + Fe2(CO3)3 2 FeI3 + 3 Ag2CO3
3) C2H4O2 + 2 O2 2 CO2 + 2 H2O
4) ZnSO4 + Li2CO3 ZnCO3 + Li2SO4
5) V2O5 + 5 CaS 5 CaO + V2S5
6) Mn(NO2)2 + BeCl2 Be(NO2)2 + MnCl2
7) 3 AgBr + GaPO4 Ag3PO4 + GaBr3
8) 3 H2SO4 + 2 B(OH)3 B2(SO4)3 + 6 H2O
9) S8 + 8 O2 8 SO2
10) Fe + 2 AgNO3 Fe(NO3)2 + 2 Ag
Additional Notes:

The Learning Goal for this assignment is:
Stoichiometry and Balancing Reactions
Stoichiometry is a section of chemistry that involves using relationships between reactants and/or
products in a chemical reaction to determine desired quantitative data. In Greek, stoikhein means
element and metron means measure, so stoichiometry literally translated means the measure of
elements. In order to use stoichiometry to run calculations about chemical reactions, it is important to
first understand the relationships that exist between products and reactants and why they exist, which
require understanding how to balanced reactions.
Balancing
In chemistry, chemical reactions are frequently written as an equation, using chemical symbols. The
reactants are displayed on the left side of the equation and the products are shown on the right, with
the separation of either a single or double arrow that signifies the direction of the reaction. The
significance of single and double arrow is important when discussing solubility constants, but we will
not go into detail about it in this module. To balance an equation, it is necessary that there are the
same number of atoms on the left side of the equation as the right. One can do this by raising the
coefficients.
Reactants to Products
A chemical equation is like a recipe for a reaction so it displays all the ingredients or terms of a
chemical reaction. It includes the elements, molecules, or ions in the reactants and in the products as
well as their states, and the proportion for how much of each particle is create relative to one another,
through the stoichiometric coefficient. The following equation demonstrates the typical format of a
chemical equation:
2Na(s) + 2HCl(aq) → 2NaCl(aq) + H2(g)
In the above equation, the elements present in the reaction are represented by their chemical
symbols. Based on the Law of Conservation of Mass, which states that matter is neither created nor
destroyed in a chemical reaction, every chemical reaction has the same elements in its reactants and
products, though the elements they are paired up with often change in a reaction. In this reaction,
sodium (Na), hydrogen (H), and chloride (Cl) are the elements present in both reactants, so based on
the law of conservation of mass, they are also present on the product side of the equations.
Displaying each element is important when using the chemical equation to convert between
elements.
Stoichiometric Coefficients
In a balanced reaction, both sides of the equation have the same number of elements. The
stoichiometric coefficient is the number written in front of atoms, ion and molecules in a chemical
reaction to balance the number of each element on both the reactant and product sides of the
equation. These stoichiometric coefficients are useful since they establish the mole ratio between
reactants and products. In the balanced equation:
2Na(s)+2HCl(aq)→2NaCl(aq)+H2(g)

we can determine that 2 moles of HCl will react with 2 moles of Na(s) to form 2 moles of NaCl(aq) and 1
mole of H2(g). If we know how many moles of Na we start out with, we can use the ratio of 2 moles
of NaCl to 2 moles of Na to determine how many moles of NaCl were produced or we can use the
ration of 1 mole of H2 to 2 moles of Na to convert to NaCl. This is known as the coefficient factor. The
balanced equation makes it possible to convert information about one reactant or product to
quantitative data about another element. Understanding this is essential to solving stoichiometric
problems.
Example 1
Lead (IV) hydroxide and sulfuric acid react as shown below. Balance the reaction.
Solution
___Pb(OH)4 +___H2SO4→___Pb(SO4)2 +___H2O
Start by counting the number of atoms of each element.
Unbalanced
Pb 1 1 Pb
O 8 9 O
H 6 2 H
S 1 2 S
The reaction is not balanced; the reaction has 16 reactant atoms and only 14 product atoms and does
not obey the conservation of mass principle. Stoichiometric coefficients must be added to make the
equation balanced. In this example, there are only one sulfur atom present on the reactant side, so a
coefficient of 2 should be added in front of H2SO4to have an equal number of sulfur on both sides of
the equation. Since there are 12 oxygen on the reactant side and only 9 on the product side, a 4
coefficient should be added in front of H2O where there is a deficiency of oxygen. Count the number
of elements now present on either side of the equation. Since the numbers are the same, the
equation is now balanced.
Pb(OH)4 + 2H2SO4→ Pb(SO4)2 + 4H2O
Balanced
Pb 1 1 Pb
O 8 12 12 9 O
H 6 8 8 2 H
S 1 2 2 2 S
Balancing reactions involves finding least common multiples between numbers of elements present
on both sides of the equation. In general, when applying coefficients, add coefficients to the
molecules or unpaired elements last.
A balanced equation ultimately has to satisfy two conditions.
1. The numbers of each element on the left and right side of the equation must be equal.
2. The charge on both sides of the equation must be equal. It is especially important to pay
attention to charge when balancing redox reactions.

Stoichiometry and Balanced Equations
In stoichiometry, balanced equations make it possible to compare different elements through the
stoichiometric factor discussed earlier. This is the mole ratio between two factors in a chemical
reaction found through the ratio of stoichiometric coefficients. Here is a real world example to show
how stoichiometric factors are useful.
Example 2
There are 12 party invitations and 20 stamps. Each party invitation needs 2 stamps to be sent. How
many party invitations can be sent?
Solution
The equation for this can be written as
I+2S→IS2
where
I represent invitations,
S represents stamps, and
IS 2 represents the sent party invitations consisting of one invitation and two stamps.
Based on this, we have the ratio of 2 stamps for 1 sent invite, based on the balanced equation.
Invitations Stamps Party Invitations Sent
In this example are all the reactants (stamps and invitations) used up? No, and this is normally the
case with chemical reactions. There is often excess of one of the reactants. The limiting reagent, the
one that runs out first, prevents the reaction from continuing and determines the maximum amount of
product that can be formed.
Example 3
What is the limiting reagent in this example?
Solution
Stamps, because there was only enough to send out invitations, whereas there were enough
invitations for 12 complete party invitations. Aside from just looking at the problem, the problem can
be solved using stoichiometric factors.
12 I x 1IS2 = 12 IS2 possible
1I
20 S x 1IS2 = 10 IS2 possible
2S

When there is no limiting reagent because the ratio of all the reactants caused them to run out at the
same time, it is known as stoichiometric proportions.
Types of Reactions
There are 6 basic types of reactions.
Combustion: Combustion is the formation of CO2 and H2O from the reaction of a chemical
and O2
Combination (synthesis): Combination is the addition of 2 or more simple reactants to form a
complex product.
Decomposition: Decomposition is when complex reactants are broken down into simpler
products.
Single Displacement: Single displacement is when an element from on reactant switches with
an element of the other to form two new reactants.
Double Displacement: Double displacement is when two elements from on reactants
switched with two elements of the other to form two new reactants.
Acid-Base: Acid- base reactions are when two reactants form salts and water.
Molar Mass
Before applying stoichiometric factors to chemical equations, you need to understand molar mass.
Molar mass is a useful chemical ratio between mass and moles. The atomic mass of each individual
element as listed in the periodic table established this relationship for atoms or ions. For compounds
or molecules, you have to take the sum of the atomic mass times the number of each atom in order to
determine the molar mass.
Example 4
What is the molar mass of H2O?
Solution
Molar mass = 2 × (1.00g/mol) + 1×(16.0g/mol) = 18.0g/mol
Using molar mass and coefficient factors, it is possible to convert mass of reactants to mass of
products or vice versa.
Example 5: Combustion of Propane
Propane (C3H8) burns in this reaction:
C3H8 + 5O2 → 4H2O + 3CO2
If 200 g of propane is burned, how many g of H2Ois produced?
Solution
Steps to getting this answer: Since you cannot calculate from grams of reactant to grams of products
you must convert from grams of C3H8 to moles of C3H8 then from moles of C3H8 to moles of H2O.
Then convert from moles of H2O to grams of H2O.

Step 1: 200g C3H8 is equal to 4.54 mol C3H8.
Step 2: Since there is a ratio of 4:1 H2O to C3H8, for every 4.54 mol C3H8 there are 18.18 mol
H2O.
Step 3: Convert 18.18 mol H2O to g H2O 18.18 mol H2O is equal to 327.27 g H2O.
Variation in Stoichiometric Equations
Almost every quantitative relationship can be converted into a ratio that can be useful in data
analysis.
Density
Density (ρ) is calculated as mass/volume. This ratio can be useful in determining the volume of a
solution, given the mass or useful in finding the mass given the volume. In the latter case, the inverse
relationship would be used.
Volume x (Mass/Volume) = Mass
Mass x (Volume/Mass) = Volume
Percent Mass
Percent establishes a relationship as well. A percent mass states how many grams of a mixture are of
a certain element or molecule. The percent X% states that of every 100 grams of a mixture, X grams
are of the stated element or compound. This is useful in determining mass of a desired substance in
a molecule.
Example 6
A substance is 5% carbon by mass. If the total mass of the substance is 10.0 grams, what is the
mass of carbon in the sample? How many moles of carbon are there?
Solution
10 g sample x 5 g carbon = 0.5 g carbon
100 g sample
0.5g carbon x 1 mol carbon = 0.0416 mol carbon
12.0g carbon
Molarity
Molarity (moles/L) establishes a relationship between moles and liters. Given volume and molarity, it
is possible to calculate mole or use moles and molarity to calculate volume. This is useful in chemical
equations and dilutions.

Example 7
How much 5M stock solution is needed to prepare 100 mL of 2M solution?
Solution
100 mL of dilute solution (1 L/1000 mL) (2 mol/1L solution) (1 L stock solution/5 mol solution) (1000
ml stock solution/1L stock solution) = 40 mL stock solution.
These ratios of molarity, density, and mass percent are useful in complex examples ahead.
Determining Empirical Formulas
An empirical formula can be determined through chemical stoichiometry by determining which
elements are present in the molecule and in what ratio. The ratio of elements is determined by
comparing the number of moles of each element present.
Example 8
1. Find the molar mass of the empirical formula CH2O.
12.0g C + (1.00g H) * (2H) + 16.0g O = 30.0 g/mol CH2O
2. Determine the molecular mass experimentally. For our compound, it is 120.0 g/mol.
3. Divide the experimentally determined molecular mass by the mass of the empirical formula.
(120.0 g/mol) / (30.0 g/mol) = 3.9984
4. Since 3.9984 is very close to four, it is possible to safely round up and assume that there was a
slight error in the experimentally determined molecular mass. If the answer is not close to a whole
number, there was either an error in the calculation of the empirical formula or a large error in the
determination of the molecular mass.
5. Multiply the ratio from step 4 by the subscripts of the empirical formula to get the molecular
formula.
CH2O * 4 =?
C: 1 * 4 = 4
H: 2 * 4 = 8
O 1 * 4 = 4
CH2O * 4 = C4H8O4
6. Check your result by calculating the molar mass of the molecular formula and comparing it to the
experimentally determined mass.
molar mass of C4H8O4= 120.104 g/mol
experimentally determined mass = 120.056 g/mol
% error = | theoretical - experimental | / theoretical * 100%
% error = | 120.104 g/mol - 120.056 g/mol | / 120.104 g/mol * 100%
% error = 0.040 %

Stoichiometry and balanced equations make it possible to use one piece of information to calculate
another. There are countless ways stoichiometry can be used in chemistry and everyday life. Try and
see if you can use what you learned to solve the following problems.
Problem 1
Why are the following equations not considered balanced?
a. H2O(l)→H2(g)+O2(g)
b. Zn(s)+Au + (aq) →Zn 2+ (aq) +Ag(s)
Problem 2
Hydrochloric acid reacts with a solid chunk of aluminum to produce hydrogen gas and aluminum ions.
Write the balanced chemical equation for this reaction.
Problem 3
Given a 10.1M stock solution, how many mL must be added to water to produce 200 mL of 5M
solution?
Problem 4
If 0.502g of methane gas react with 0.27g of oxygen to produce carbon dioxide and water, what is the
limiting reagent and how many moles of water are produced? The unbalanced equation is provided
below.
CH4(g)+O2(g)→CO2(g)+H2O(l)

Theoretical and Actual Yields
Key Terms
(Excess reagent, limiting reagent)
Theoretical and actual yields
Percentage or actual yield
Skills to Develop
Use stoichiometric calculation to determine excess and limiting reagents in a chemical reaction
and explain why.
Calculate theoretical yields of products formed in reactions that involve limiting reagents.
Evaluate percentage or actual yields from known amounts of reactants
Theoretical and Actual Yields
Reactants not completely used up are called excess reagents, and the reactant that completely
reacts is called the limiting reagent. This concept has been illustrated for the reaction:
2Na+Cl2 →2NaCl
Amounts of products calculated from the complete reaction of the limiting reagent are called
theoretical yields, whereas the amount actually produced of a product is the actual yield. The ratio of
actual yield to theoretical yield expressed in percentage is called the percentage yield.
percent yield = actual yield / theoretical yield ×100
Chemical reaction equations give the ideal stoichiometric relationship among reactants and products.
Thus, the theoretical yield can be calculated from reaction stoichiometry. For many chemical
reactions, the actual yield is usually less than the theoretical yield, understandably due to loss in the
process or inefficiency of the chemical reaction.
Example 1
Methyl alcohol can be produced in a high-pressure reaction
CO(g) + 2H2(g) →CH3OH(l)
If 6.1 metric tons of methyl alcohol is obtained from 1.2 metric tons of hydrogen reacting with excess
amount of CO, estimate the theoretical and the percentage yield?
Hint:
To calculate the theoretical yield, consider the reaction
CO(g) + 2H2(g) → CH3OH(l)
28.0 + 4.0 = 32.0 (stoichiometric masses in, g, kg, or tons)
1.2 tons H2 × 32.0 CH3OH = 9.6 tons CH3OH
4.0 H2
Thus, the theoretical yield from 1.2 metric tons (1.2x10 6 g) of hydrogen gas is 9.6 tons. The actual
yield is stated in the problem, 6.1 metric tons. Thus, the percentage yield is

%yield = 6.1 tons × 100 = 64%
9.6tons
Discussion
Due to chemical equilibrium or the mass action law, the limiting reagent may not be completely
consumed. Thus, a lower yield is expected in some cases. Losses during the recovery process of the
product will cause an even lower actual yield.
Example 2
A solution containing silver ion, Ag + , has been treated with excess of chloride ions Cl − . When dried,
0.1234 g of AgCl was recovered. Assuming the percentage yield to be 98.7%, how many grams of
silver ions were present in the solution?
Hint:
The reaction and relative masses of reagents and product are:
The calculation,
Ag + (aq) + Cl − (aq) → AgCl(s)
107.868 + 35.453 = 143.321
0.1234 g AgCl ×107.868 g Ag + =0.09287 g Ag +
143.321g AgCl
shows that 0.1234 g dry AgCl comes from 0.09287g Ag + ions. Since the actual yield is only 98.7%,
the actual amount of Ag + ions present is therefore
0.09287 g Ag + = 0.09409 g Ag +
0.987
Discussion
One can also calculate the theoretical yield of AgCl from the percentage yield of 98.7% to be
0.1234 g AgCl =0.1250 g AgCl
0.987
From 0.1250 g AgCl, the amount of Ag + present is also 0.09409 g.
Stoichiometry - A Review
Skills Taught
evaluate molecular weight for a given formula
evaluate weight (mass) percentages of elements for a given formula
evaluate amounts (in mass and mole units) produced in a chemical reaction from given
conditions
classify reactions by types: combination, combustion, displacement, formation, etc
determine the chemical formula when weight percentages are given and then evaluate the
mole percentages of elements in the formula
determine the chemical formula when weight percentages are given and molecular weight is
known
determine the amount produced, the actual yield, and other stoichiometry quantities for a given
reaction

Review Purposes
To get an overall view of stoichiometry.
Apply skills learned to perform quantitative chemical analysis.
Apply theories and rules of chemistry to solve problems.
Assess areas of strength and weakness for review purposes.
Improve problem solving strategy and learning efficiency.
Stoichiometry
STOICHIOMETRY is the quantitative relationship of reactants and products. This unit has been
divided into the following objects. A brief review is given here so that you can get a birds'-eye or
overall view of stoichiometry.
1.Amounts of substances
Express amounts of substance in mass units of g, kg, tons, and convert them to moles,
kilomoles, or millimoles.
2.Chemical formulas
Represent a substance with a formula that reflects its chemical composition, structure, and
bonding; evaluate weight and mole percentages of elements in a substance; and determine
chemical formula by elemental analysis.
3.Reaction features
Define some common features of chemical reactions; classify chemical reactions by common
features such as combination, combustion, decomposition, displacement, and redox reactions.
4.Reaction equations
Express quantitative relationships using chemical reaction equations; evaluate quantities of
reactants and products in a chemical reaction; and solve reaction stoichiometry problems.
5.Excess and limiting reagents
Define excess and limiting reagents; determine excess and limiting reagents in a reaction
mixture; and determine quantities produced in a chemical reaction.
6.Yields
Define theoretical and actual yields due to limiting reagent; apply the concept of limiting
reagent to evaluate theoretical yield; convert actual yield to percentage yields.

Use the space provided to write out the steps you take to solve different types of problems.
Use any additional space for notes. These 2 pages should be full when you turn in your notebook.
Steps to solve problems
Mole to Mole:

Unit 6
Chapter 13 States of Matter
The students will learn what are the factors that determine and
characteristics that distinguish gases liquids and solids and
how substances change from one state to another.
Differentiate among the four states of matter.
Students will measure the physical characteristics of matter such as temperature
and density.
Students will compare and contrast the physical characteristics of the 4 states of
matter.
solid
liquid
gas
plasma
Relate temperature to the average molecular kinetic energy.
Students will be able to compare and contrast the motion of particles of a sample
at various temperatures.
Kinetic energy
Kinetic theory
Temperature
Describe phase transitions in terms of kinetic molecular theory.
Students will be able to identify and describe phase changes.
Students will be able to compare and contrast the change in particle motion
for phase changes.
Students will be able to interpret heating/cooling curves and phase diagrams.
melting point
freezing point
boiling point
condensation
sublimation
phase diagram
kinetic molecular theory

Chapter 14 The Behavior of Gases
The students will learn how gases respond to changes in
pressure, volume, and temperature and why the ideal gas law
is useful even though ideal gases do not exist.
Interpret the behavior of ideal gases in terms of kinetic molecular theory.
Students will be able to describe the behavior of an ideal gas.
Students will participate in activities to apply the Ideal Gas Law and its
component laws to predict gas behavior.
Students will be able to perform temperature/pressure conversions.
Compressibility
Boyle’s Law
Charles’s Law
Gay-Lussac’s Law
Combine Gas Law
Ideal Gas Law
Partial pressure
Dalton’s Law of partial pressure
Diffusion
Effusion
Graham’s Law of effusion
Chapter 15 Water and Aqueous Systems
The students will learn how the interactions between water
molecules account for the unique properties of water and how
aqueous solutions form.
Discuss the special properties of water that contribute to Earth's suitability
as an environment for life: cohesive behavior, ability to moderate
temperature, expansion upon freezing, and versatility as a solvent.
Students will be able to prepare a solution of known molarity
Students will participate in activities to calculate molarity
Surface tension
Surfactant
Aqueous solutionSolvent
Solute

The Learning Goal for this assignment is: Differentiate among the four
states of matter.
Take note over the following chapter. Use the Headings provided to organize your notes. Define and number all highlighted
vocabulary (total 23 ) as well as summarize the sections. You may add pictures where needed. The pictures should be an
appropriate size. Use Arial 12 for all text. This document should be 3 pages and should be saved as a pdf before you submit it into
Angel.
13.1 The Nature of Gases
Chapter 13 States of Matter
Pages 420 - 439
Kinetic Theory and a Model for Gases
Kinetic energy 1 is energy that is always in motion
What are the three assumptions of the kinetic theory 2 (All matter consists of tiny particles
that are in constant motion) as it applies to the gases?
The particles in a gas are considered to be small, hard spheres with an insignificant
volume. There is empty space between the particles.
The motion of the particles is rapid, constant, and random. They have no sense of
direction and an aimless path and collide with other particles.
All collisions between particles in a gas are perfectly elastic. An elastic collision is when
kinetic energy is transferred without losing particles.
Gas Pressure
How does kinetic theory explain gas pressure?
3
Gas pressure (results from the force exerted by a gas per unit surface area of an object)
the results of rapidly moving particles in a gas simultaneously colliding with an object.
Gas pressure caused by a mixture of gases.
What causes this force?
Moving bodies exert a force when they collide with other bodies.
If there are not particles colliding with one another or no particles at all and it's all empty
4
it creates a vacuum
The air exerts pressure on Earth because gravity holds the particles in the air within the
atmosphere. The collision of atoms and molecules in air with objects results in
5
atmospheric pressure.
To measure the atmospheric pressure, we use a device called the barometer
7
The SI unit of pressure is Pascal (Pa). It represents a small amount of pressure.
Standard atmosphere (atm):
8
1 atm= 760 mm Hg= 101.3 kPa
Kinetic Energy and Temperature
6

What is the relationship between the temperature in kelvins and the average kinetic
energy of particles? The kelvin temperature of a substance is directly proportional to the
average kinetic energy of the particles of the substance
When the heat increases the particles absorb energy. The increase in kinetic energy
results in the increase of temperature.
The average kinetic energy of the particles in a substance is directly related to the
substance temperature. At any given temperature the particles of a substance have the
same average kinetic energy.
13.2 The Nature of Liquids
A Model of Liquids
What factors determine the physical properties of a liquid?
The interplay between the disruptive motions in a particle in a liquid and attractions
among the particles determines the physical properties of liquids
According to the kinetic theory both the particles in gases and liquids have kinetic
energy.
Intermolecular attraction reduces the amount of space between the particles in a liquid.
Making it denser than a gas.
Evaporation
What is the relationship between evaporation and kinetic energy?
During evaporation
9(a process of when conversion occurs at the surface of a liquid that
is not boiling), only those molecules with a certain minimum kinetic energy can escape
from the surface of the liquid.
Just like kinetic energy and temperature evaporation occurs quickly because the liquid
increases the average kinetic energy of its particles.
*the conversion of a liquid to a gas or vapor is called vaporization.
Vapor Pressure
When can a dynamic equilibrium exist between a liquid and a vapor?
In a system at constant vapor pressure, a dynamic equilibrium exists between the vapor
and the liquid. The system is in equilibrium because the rate of evaporation of liquid
equals the rate of condensation of vapor.
To measure the vapor pressure 11of a liquid you can utilize a device called a manometer.
Boiling Point
Under what condition does boiling occur?
When a liquid is heated to a temperature at which particles throughout the liquid have
enough kinetic energy to vaporize, the liquid begins to boil.
The boiling point(bp)
12is the temperature at which the vapor pressure of a liquid is just as
equal to the external pressure on the liquid
Normal boiling
point 13 = 101kPa
The average kinetic energy of the particles increases and the temperature rises.
13.3 The Nature of Solids
A Model of Solids
10

How are the structures and properties of a solid related?
The general properties of a solid reflect in the orderly arrangement of their particles and
the fixed location of their particles.
The particles in a solid differ from the ones in a liquid since they do not have plenty of
motion.
As we could tell that the previous ones share the same property that when they are
heated the particles move more rapidly as their kinetic energy increases. When the
14
temperature increases the solid melts. The melting point is when a solid transforms into
a liquid. Changing the temperature of the states of matter changes the movement of its
15
particles and its physical property. The freezing point is the temperature where a liquid
converts into a solid.
Compounds have different melting points depending on how strong the force is that
holds them together. Some solids at times don’t melt.
Crystal Structure and Unit Cell
What determines the shape of a crystal?
The shape of a crystal reflects the arrangement of the particles within the solid.
16
A crystal is the most solid substance. Its a homogeneous structure, the same
composition throughout. It has symmetrically arranged plane faces.
The crystals are placed in seven different groups. It is classified by the angle of the
structure.
The smallest group of particles within a crystal is known as a unit cell. Each crystal
system can have up to four types of unit cells.
18
Allotrope is each of multiple different physical forms in an element that can exist. They
have different properties and structure.
19
All though some solids are crystalline not all solids are. An amorphous solid lacks an
ordered internal structure. They are not ordered symmetrically like crystal their order is
more random.
20
A type of amorphous solids is glass. Glass is a transparent fusion product of inorganic
substances that cooled and did not crystallize.
13.4 Changes of State
Sublimation
When can sublimation occur?
Sublimation occurs in solids with vapor pressure that exceed atmospheric pressure at or
near room temperature.
Sublimation also has vapor pressure like liquids.
Sublimation has useful applications to daily life, for example brewed coffee.
Phase Diagrams
How are the conditions at which phases are in equilibrium represented on a phase
diagram?
The conditions of pressure and temperature at which two phases exist in equilibrium are
indicated on a phase diagram by a line separating the two regions representing the
phases. sublimation 21 changes from a solid directly to a gas. Phase diagram 22 gives
conditions of temperature and pressure. Triple point 23 where three lines in a digram
meet at the same point.
17

Name: Milca Mazariegos
Name: Jamyled Vergara
Grade: 10
States of Matter Project
You and you lab partner are going to create a study aid in the form of a game for the
information in Chapter 13 States of Matter.
First, each of you, independently from each other, will summarize the chapter on 3
pages of a pdf which will be submitted in Angel by the end of class on Wednesday Feb
22.
Second, you and you lab partner will be given a game platform which you will use for
your questions and answers, either Jeopardy or Kahoot.
Third, you will fill in the information at the bottom of this page with your username,
passwords and/or websites so that you do not forget this and I have a copy in case
anything gets misplaced. This page will be submitted into Angel as a Word Document
on Wednesday February 22 during class.
Fourth, you will use your notes to generate the questions and answers.
Finally you will give me access to your game by putting the website or Game Number
on this page adding this page to your 3 pages of notes and resubmitting it in Angel as
a pdf by the end of the class on Friday Feb 24.
This page is due by the end of class on Wednesday February 22.
This Project is due by the end of class on Friday February 24.
Jeopardy (https://jeopardylabs.com)
Password: Chemistry23
Edit Link: https://jeopardylabs.com/edit/states-of-matter698
Play Link: https://jeopardylabs.com/play/states-of-matter698
Kahoot (https://getkahoot.com)
Username:
Email:
Password:
Game PIN:

Gas Laws
Pressure
Volume
Temperature
Average Kinetic energy
https://www.gre.nasa.gov/www /k-12/air
plane/ An imation/frg lab2. html
When holding one thing constant one or the other will be changed
keeping the pressure and volume constant you can predict the temperature is a effected

Notes
mass is the amount particle s there are in there.
if you decrease a volume of a balloon it increases the pressure and vice versa- inverse
relationship( when one thing goes up the other thing goes down
Multiply

Notes
keep the pressure constant
direct proportion= as one thing increase the other thing will increase.
how does the volume increase the temperature?
The molecules pick up more kinetic energy to maintain the pressure
on the outside which increases the temperature.
*K is the constant
must be equal at the end

notes
As we increase the temperature we increase the pressure
Directly proportionate= have to be division
keeping the container the same size not allowing it to get bigger.
108kPa= x
41c =22c
change the Celsius into kelvin
add 273 to the number
108kPa=x
314k =293k
cross multiply
314k x X= 31860
314k 314k
101.4649682 = 101 kPa

1) 1 atm= 760 torr
19980 torr- 7730 ml
2)
KHBDCM
5140ml/1000
P1 x 5.14L= 2070torr x 8.79L
P1 x 5140ml = 2070 torr x 8.79L
Pl x 5.14L= 18195l.3 torr x X
5.14L
5.14L
2.605atm * 7730ml= 2.95atm *ml
axb=c x X
20136.65= 2.95x X
2.95 2.95
68259.983051 ml
answer:6830ml
3539.941634 torr 1 atm
4.6578179atm 760 torr
P1= 4.66

To find K from C + 273
k---->
1) V1 = 3550 ml
85.4C 294.0 k
convert c to K 85.4 + 273 = 358.4
v1
358.4K = 3550ml
294.4 K
cross multiply
358.4 x 3550ml
294.0 K x X= 1272320 k x ml
294.0 k 294.0 k
4.327619044= 4.33l
2) 2920ml
T1
= 6660ml
853.0k
cross multiply
2920mlx 853.0 k
6660mlx X=24907660ml x k
6660ml 6660ml
373.987988k =374k

1)
2) 198kPa = 225kPa
27c x
= convert c to kelvin + 273 = 300 198kPa = 225kPa
273k x
cross multiply
198kPa x X= 67500
198kPa 198kPa
340.90= 341k
All three laws combine together
the Combined gas laws
P1 x V1 = P2 x V2
T1 T2
155kPa x 1.00L
25.0c
= 605kPa x X
125c
convert into kelvin
155kPax 1.00= 605 x X
298K 398k
Ideal gas Laws
P x V = n RT
155kPaxL= 605x X
298k 398k
61690=180290kpa x k x X
180290 180290
.342170947
v=.342L

The Combined Gas Laws
The combined gas laws make use of the relationships shared by pressure, volume and temperature: the
variables found in other gas laws such Boyle’s, Charles’, and Gay- Lussac’s law.
T1 T2
P1 x V1 = P2 x V2
450 mL of a gas occupies a container that has a temperature of 28°C and a pressure of 788 mmHg.
What is the temperature if the volume is reduced to 50 mL at 760 mmHg? As the unit of
measurement is in Celsius, remember to convert it to Kelvin

Ideal Gas Laws
The state variables, pressure, volume, and temperature are used to determine the ideal gas law
and the state of n moles.
P x V = n RT

The Learning Goal for this assignment is:
The students will learn how the interaction between water molecules account for the
unique properties of water and how aqueous solutions form
Take note over the following chapter. Use the Headings provided to organize your notes. Define and number all highlighted vocabulary (total 22 ) as well
as summarize and take notes over the sections. You may add pictures where needed. The pictures should be an appropriate size. Use Arial 12 for all
text. This document should be 2 pages and should be saved as a pdf before you submit it into Angel.
Chapter 15 Water and Aqueous Systems
Pages 488 - 507
15.1 Water and Its Properties
Water in the Liquid State
What factor causes the high surface tension, low vapor pressure, and high boiling point of water?
Many unique important properties of water including its high surface tension, low vapor pressure and
high boiling point- result from hydrogen bonding.
How do the polarities of the two O-H bonds affect the polarity of the molecule? The shape of the
molecule is the determining factor. The bond of the angle is estimated at 105 degrees, which gives
the molecule its bent shape.
Surface tension 1 is the inward force, or pull, that tends to reduce the surface area of a liquid. So like
the liquid a centimeter above the glass just as it is about to spill but doesn’t. There is a lot of tension
in water but there is a way to decrease it. Surfactant 2 is a substance that interferes with hydrogen
bonding minimizing the surface tension.
Water in the Solid State
How can you describe the structure of ice?
The structure of ice is a regular open framework of water molecules in a hexagonal arrangement.
Water molecules are denser than ice.
What is the melting point for ice? The melting point of ice is at 0 degrees Celsius
15.2 Homogeneous Aqueous Systems
Solutions
What types of substances dissolve most readily in water? Substances that dissolve most readily in
water include ionic compounds and polar covalent compounds.
Methane is a non-covalent compound.
All water contains minerals and gases. An aqueous solution 3 is water that contains dissolved
substances.
In a solution, the dissolving medium is the solvent 4 . The dissolved particles are the solute 5 .
*solutes and solvents may be gases, liquids, or solids.
Solvation 6 is the process by which positive and negative ions of an ionic solid become surrounded by
solvent molecules.
Electrolytes and Nonelectrolytes
An electrolyte 7 is a compound that conducts an electric current when it is an aqueous solution or in a
molten state. *requires mobile ions.
Why are all ionic compounds electrolytes?
All ionic compounds are electrolytes because they dissociate into ions.
Aqueous solution of hydrogen is an electrolyte= HCl(g)+H2O(l) H3O +(aq) + Cl - (aq)

A nonelectrolyte 8 does not conduct an electric current because they are not composed of ions.
In a strong electrolyte 9 almost all the solutes exist as ions. In a weak electrolyte 10 it conducts poorly
because nearly no solutes exist as ions.
Hydrates
The forces holding the water molecules in hydrates are not very strong, so the water is easily lost and
regained. Water of hydration is water contained in a crystal. A hydrate is a compound that contains
water of hydration. Effloresce is a suspension differs from a solution because the particles of a
suspension are much larger and do not stay suspended indefinitely. A suspension is mixture from
which particles settle out upon standing Suspension are heterogeneous because at least two
substances can be identified. Anhydrous: A substance that does not contain water. Hygroscopic are
hydrates or other compounds that remove moisture from air. Desiccant is substance sued to absorb
moisture from air and create dry atmosphere. Deliquescent remove sufficient water from the air to
dissolve completely and form solutions.
15.3 Heterogeneous Aqueous Systems
Suspensions
A suspension differs from a solution because the particles of a suspension are much larger and do
not stay suspended indefinitely. Suspension: A mixture from which particles settle out upon standing
Suspension are heterogeneous because at least two substances can be identified.
Colloids
Colloids have particles smaller than those in suspensions and larger than those in solutions. Colloids:
A heterogeneous mixture containing particles that range in size from 1 n to 1000 nm. Typical effect:
the scattering of visible light by colloidal particles. Brownian motion: Chaotic movement of colloidal
particles. Emulsion: A colloidal dispersion of a liquid in a liquid. Particles are spread or dispersed.
They’re cloudy so they can be distinguished by Tyndall effect and by the observation of Brownian
motion. They are also subject to coagulation or clumping together and they can be emulsified or
made stable. Coagulation are when they absorb ions from the from the dispersing medium onto the
surface.

Unit 7
Chapter 16 Solutions
The students will learn what properties are used to describe
the nature of solutions and how to quantify the concentration
of a solution.
Chapter 17 Thermochemistry
The student will learn how energy is converted in a chemical
or physical process and how to determine the amount of
energy is absorbed or released in that process.
Differentiate among the various forms of energy and recognize that they can
be transformed from one form to others.
Students will participate in activities to investigate and describe the
transformation of energy from one form to another (i.e. batteries, food, fuels,
etc.)
Explore the Law of Conservation of Energy by differentiating among open,
closed, and isolated systems and explain that the total energy in an isolated
system is a conserved quantity.
Students will be able to calculate various energy changes:
o q = mc∆t
o ∆Hfus
o ∆Hmelt
Thermochemistry
Heat
System
Surrounding
Law of conservation of energy
Bond Making is exothermic
Bond Breaking is endothermic
Heat capacity
Specific heat
Calorimetry
Enthalpy
Thermochemical equation
Molar heat of (fusion, solidification,
vaporization, condensation, solution)
Distinguish between endothermic and exothermic chemical processes.
Students will be able to recognize exothermic and endothermic reactions through
experimentation.
Students will participate in activities (Pasco) to create exothermic and
endothermic graphs.
Endothermic
Exothermic

Create and interpret potential energy diagrams, for example: chemical
reactions, orbits around a central body, motion of a pendulum
Students will participate in activities (Pasco) to create exothermic and
endothermic graphs.
Students will be able to interpret exothermic and endothermic reaction graphs.
Potential energy diagram
Thermochemical equations
Chapter 18 Reaction Rates and Equilibrium
The student will learn how the rate of a chemical reaction can
be controlled, what the role of energy is and why some
reactions occur naturally and others do not.
Explain how various factors, such as concentration, temperature, and
presence of a catalyst affect the rate of a chemical reaction.
Students will be able to describe how each factor may affect the rate of a
chemical reaction.
Students will be able to compare the relative effect of each factor on the rate of a
chemical reaction.
Rate
Collision theory
Activation energy
Catalyst
Activated complex
Inhibitor
Explain the concept of dynamic equilibrium in terms of reversible processes
occurring at the same rates.
Students will be able to describe a system in dynamic equilibrium.
Students will be able to describe how factors may affect the equilibrium of a
reaction.
Reversible reaction
Chemical equilibrium
Le Chatelier principle
Explain entropy’s role in determining the efficiency of processes that convert
energy to work.
Students will be able to describe the change in entropy of a reaction.
Students will be able to determine if a reaction is spontaneous
Entropy
Law of disorder
Spontaneous/nonspontaneous reaction

The Learning Goal for this assignment is:
Defining Concentration
Measures of Concentration
Concentration is defined as the amount of dissolved solute in a given amount of solvent or
solution. There are several terms that describe concentration. Some of these terms are relative;
that is, they can be used only to compare the concentration of one solution to another. Dilute and
concentrated are two such terms. A dilute solution contains less dissolved solute than a
concentrated solution (in equal volumes of solution).
The terms saturated, unsaturated, and supersaturated are terms that describe concentration more
precisely.
Saturated: The maximum amount of solute is dissolved in a given amount of solvent at a
particular temperature. Such solutions are stable.
Unsaturated: Less than the maximum amount of solute is dissolved in a given amount of
solvent at a particular temperature. Such solutions are stable.
Supersaturated: More than the maximum amount of solute is dissolved in a given amount of
solvent at a particular temperature. Such solutions are unstable.
Look at the solubility curve shown below:
The solubility of NaNO3 is 86.0 g/100 mL H2O (at 20 °C). If you
prepare a solution of 50.0 g NaNO3 dissolved in 100 mL H2O, an unsaturated solution results
(Point A on the graph).
continue adding NaNO to the solution until 86.0 g are dissolved, a saturated solution results
(Point B).
heat the solution to 50 °C, 113 grams NaNO3 can be dissolved (Point C). When the solution
cools back down to 20 °C, it will be supersaturated (Point D).
Quantifying Concentration
To describe the concentration of a solution even more precisely, various measures of concentration
can be used. Some of the ways concentration can be quantified include calculating the
Mass of solute per solution mass (expressed as a percent or parts per million)
Moles of solute per kilogram solvent (molality)
Mass of solute per liter of solution (grams/liter)
Moles of solute per liter of solution (molarity)

Part 1: Mass Percent
Mass percent (also called percent by mass, weight percent, or percent by weight) compares the
mass of the solute to the entire mass of the solution.
Notes:
Part 2: Parts per Million
Parts per million (ppm) is another measure of concentration. It is similar to mass percent. But
mass percent indicates the number of grams of solute per 100 g solution. Parts per million
indicates the number of grams of solute per 1,000,000 g solution. This measure of concentration is
often used to express the concentrations of very dilute solutions.
Notes:
Part 3: Molality
Molality (m) is the ratio of the moles of solute to the kilograms of solvent. Note: this is the first
measure of concentration that is concerned with the mass of the solvent, not the mass of the
solution as a whole.
Notes:
Part 4: Grams per Liter
To express the concentration of a solution in grams per liter, you must know the mass of the solute
and the volume of the solution, not just the volume of the solvent.
Notes:
Suppose you wanted to know what the concentration would be before making the solution. Could that
be done? In order to relate the volume or mass of solvent to the volume of solution, you would have
to know the density of the solution. You will see how solution density can be used to calculate
molarity in the next section.

Part 5: Molarity
Molarity (M) is the most common measure of concentration. The concentration of most solutions
you see in the lab are expressed in terms of molarity. Just like g/L, molarity calculations require
that you know either the volume or density of the resulting solution.
Notes:
Using Density to Calculate Molarity
If the volume of the resulting solution is not known, molarity is calculated as follows:
Convert the volume of solvent to grams. (The simulation does this step for you.)
Determine the total mass of the solution (mass of solute + mass of solvent).
Convert the solution mass to volume in milliliters, using its density (volume = mass / density).
Convert the solution volume to liters (divide by 1000).
Convert solute grams to moles.
Calculate the molarity (moles solute / L solution).
Reinforcing What You've Learned
1. 35.0 g potassium dichromate are dissolved in 354.0 g distilled water. What is the concentration
of the resulting solution, expressed as percent by mass? 9.00%
2. A chemist discovers that 2587.0 g distilled water are contaminated with 13.0 g NaNO3. What is
the concentration of NaNO3, expressed in ppm? 5.00 X 10 3
3. 175.00 g H2SO4 are added to 61.49 g distilled H2O. What is the molality of the resulting solution?
29.04m
4. 225 g NaBr are dissolved in 525 g distilled water. If the volume of the resulting solution is
584 mL, what is its concentration when expressed in g/L? 385g/L
5. 470.0 g Na2CO3 are dissolved in 4230.0 g distilled water. What is the molarity of the resulting
solution (D = 1.1029 g/mL)?
1.1561 M
Applying What You've Learned
6. Think about the equations for mass percent and parts/million. Consider how you might convert
between these two measures of concentration.
Mass percent = ppm ÷ 10,000 ppm = mass percent X .0001
7. Think about the equations for grams/liter and molarity. Consider how you might convert
between these two measures of concentration.
grams / L = molarity X molar mass
molarity = grams / L ÷ molar mass

Glossary
Concentration: Amount of dissolved solute in a given amount of solvent or solution.
Density: Amount of matter per unit volume. Density is calculated by dividing an object's mass by
its volume.
Mass: Amount of matter an object contains or, more scientifically, the measure of an object's
resistance to changes in motion. The SI (Systéme International) unit for mass is the kilogram. In
the lab, mass is often measured in grams.
Mass percent: Also referred to as percent by mass and, occasionally, weight percent or percent
by weight. Mathematically, mass % = (mass of solute / mass of entire solution) x 100.
Molar mass: Mass of a compound, calculated by adding up the individual masses for its component
atoms, which are obtained from the periodic table of the elements. Molar mass is expressed in
grams/mole and is sometimes referred to as molecular mass (for molecular compounds) or formula
mass (for ionic compounds).
Molality (m): Moles of solute per kilogram of solvent. Mathematically, m = moles of solute / kilogram of
solvent.
Molarity (M): Moles of solute per liter of solution. Mathematically, M = moles of solute / liter of
solution.
Mole (mol): Counting unit used to express the large numbers of particles, such as atoms or
molecules, that are involved in chemical processes. One mole of particles contains 6.02 x 10
particles. The mass of one mole of an element, in grams, is equivalent to the atomic mass for that
element, as indicated on the periodic table.
Parts per million (ppm): Measure of concentration often used for dilute solutions. Mathematically,
ppm = (mass of solute / mass of solution) x 10.
Solubility: Measure of the maximum amount of solute that can be dissolved in a given amount of
solvent at a given temperature, forming a stable solution.
Solute: Dissolved substance in a solution. The solute is generally the solution component present
in the lesser amount.
Solution: Homogeneous mixture in which one substance has been dissolved in another.
Solvent: Substance in which a solute is dissolved to form a solution. The solvent is generally the
solution component present in the greater amount.
Volume: Amount of space an object occupies. The SI (Systéme International) unit for volume is
the cubic meter. In the lab, volume is often measured in cubic centimeters, milliliters, or liters.
SAS Curriculum Pathways VLab #866

The System and the Surroundings in Chemistry
Thermochemistry
The system is the part of the universe we wish to focus our attention on. In the world of chemistry, the
system is the chemical reaction. For example:
2H2 + O2 ---> 2H2O
The system consists of those molecules which are reacting.
The surroundings are everything else; the rest of the universe. For example, say the above reaction is
happening in gas phase; then the walls of the container are part of the surroundings.
There are two important issues:
1. a great majority of our studies will focus on the change in the amount of energy, not the
absolute amount of energy in the system or the surroundings.
2. regarding the direction of energy flow, we have a "sign convention."
Two possibilities exist concerning the flow of energy between system and surroundings:
1. The system can have energy added to it, which increases its amount and lessens the energy
amount in the surroundings.
2. The system can have energy removed from it, thereby lowering its amount and increasing the
amount in the surroundings.
We will signify an increase in energy with a positive sign and a loss of energy with a negative sign.
Also, we will take the point-of-view from the system. Consequently:
1. When energy (heat or work) flow out of the system, the system decreases in its amount. This
is assigned a negative sign and is called exothermic.
2. When energy (heat or work) flows into the system, the system increases its energy amount.
This is assigned a positive sign and is called endothermic.
We do not discuss chemical reactions from the surrounding's point-of-view. Only from the system's.
Notes:

Specific Heat
Here is the definition of specific heat:
the amount of heat necessary for 1.00 gram of a substance to change 1.00 °C
Note the two important factors:
1. It's 1.00 gram of a substance
2. and it changes 1.00 °C
Keep in mind the fact that this is a very specific value. It is only for one gram going one degree. The
specific heat is an important part of energy calculations since it tells you how much energy is needed
to move each gram of the substance one degree.
Every substance has its own specific heat and each phase has its own distinct value. In fact, the
specific heat value of a substance changes from degree to degree, but we will ignore that.
The units are often Joules per gram-degree Celsius (J/g*°C). Sometimes the unit J/kg K is also used.
This last unit is technically the most correct unit to use, but since the first one is quite common, you
will need to know both.
I will ignore calorie-based units almost entirely.
Here are the specific heat values for water:
Phase J g¯1 °C¯1 J kg¯1
K¯1
Gas 2.02 2.02 x 10 3
Liquid 4.184 4.184 x 10 3
Solid 2.06 2.06 x 10 3
Notice that one set of values is simply 1000 times bigger than the other. That's to offset the influence
of going from grams to kilograms in the denominator of the unit.
Notice that the change from Celsius to Kelvin does not affect the value. That is because the specific
heat is measured on the basis of one degree. In both scales (Celsius and Kelvin) the jump from one
degree to the next are the same "distance." Sometimes a student will think that 273 must be involved
somewhere. Not in this case.
Specific heat values can be looked up in reference books. Typically, in the classroom, you will not be
asked to memorize any specific heat values. However, you may be asked to memorize the values for
the three phases of water.
As you go about the Internet, you will find different values cited for specific heats of a given
substance. For example, I have seen 4.186 and 4.187 used in place of 4.184 for liquid water. None of
the values are wrong, it's just that specific heat values literally change from degree to degree. What
happens is that an author will settle on one particular value and use it. Often, the one particular value
used is what the author used as a student.
Hence, 4.184.

The Time-Temperature Graph
We are going to heat a container that has 72.0 grams of ice (no liquid water yet!) in it. To make the
illustration simple, please consider that 100% of the heat applied goes into the water. There is no loss
of heat into heating the container and no heat is lost to the air.
Let us suppose the ice starts at -10.0 °C and that the pressure is always one atmosphere. We will
end the example with steam at 120.0 °C.
There are five major steps to discuss in turn before this problem is completely solved. Here they are:
1. the ice rises in temperature from -10.0 to 0.00 °C.
2. the ice melts at 0.00 °C.
3. the liquid water then rises in temperature from zero to 100.0 °C.
4. the liquid water then boils at 100.0 °C.
5. the steam then rises in temperature from 100.0 to 120.0 °C
Each one of these steps will have a calculation associated with it. WARNING: many homework and
test questions can be written which use less than the five steps. For example, suppose the water in
the problem above started at 10.0 °C. Then, only steps 3, 4, and 5 would be required for solution.
To the right is the type of graph which is typically used to
show this process over time.
You can figure out that the five numbered sections on the
graph relate to the five numbered parts of the list just above
the graph.
Also, note that numbers 2 and 4 are phases changes: solid
to liquid in #2 and liquid to gas in #4.
Q=mcΔT
where ΔT is (Tf – Ti)
Here are some symbols that will be used, A LOT!!
Δt = the change in temperature from start to finish in degrees Celsius (°C)
m = mass of substance in grams
c = the specific heat. Its unit is Joules per gram X degree Celsius (J / g °C is one way to write
the unit; J g¯1 °C¯1 is another)
q = the amount of heat involved, measured in Joules or kilojoules (symbols = J and kJ)
mol = moles of substance.
ΔH is the symbol for the molar heat of fusion and ΔH is the symbol for the molar heat of
vaporization.
We will also require the molar mass of the substance. In this example it is water, so the molar mass is
18.0 g/mol.
Notes:

Step One: solid ice rises in temperature
As we apply heat, the ice will rise in temperature until it
arrives at its normal melting point of zero Celsius.
Once it arrives at zero, the Δt equals 10.0 °C.
Here is an important point: THE ICE HAS NOT MELTED
YET.
At the end of this step we have SOLID ice at zero
degrees. It has not melted yet. That's an important point.
Each gram of water requires a constant amount of energy
to go up each degree Celsius. This amount of energy is
called specific heat and has the symbol c.
72.0 grams of ice (no liquid water yet!) has changed 10.0 °C. We need to calculate the energy
needed to do this.
This summarizes the information needed:
Δt = 10 °C
The mass = 72.0 g
c = 2.06 Joules per gram-degree Celsius
The calculation needed, using words & symbols is:
q = (mass) (Δt) (c)
Why is this equation the way it is?
Think about one gram going one degree. The ice needs 2.06 J for that. Now go the second degree.
Another 2.06 J. Go the third degree and use another 2.06 J. So one gram going 10 degrees needs
2.06 x 10 = 20.6 J. Now we have 72 grams, so gram #2 also needs 20.6, gram #3 needs 20.6 and so
on until 72 grams.
With the numbers in place, we have:
q = (72.0 g) (10 °C) (2.06 J/g °C)
So we calculate and get 1483.2 J. We won't bother to round off right now since there are four more
calculations to go. Maybe you can see that we will have to do five calculations and then sum them all
up.
One warning before going on: three of the calculations will yield J as the unit on the answer and two
will give kJ. When you add the five values together, you MUST have them all be the same unit.
In the context of this problem, kJ is the preferred unit. You might want to think about what 1483.2 J is
in kJ.
Notes:

Step Two: solid ice melts
Now, we continue to add energy and the ice begins to
melt.
However, the temperature DOES NOT CHANGE. It
remains at zero during the time the ice melts.
Each mole of water will require a constant amount of
energy to melt. That amount is named the molar heat of
fusion and its symbol is ΔHf. The molar heat of fusion is
the energy required to melt one mole of a substance at its
normal melting point. One mole of solid water, one mole
of solid benzene, one mole of solid lead. It does not
matter. Each substance has its own value.
During this time, the energy is being used to overcome water molecules' attraction for each other,
destroying the three-dimensional structure of the ice.
The unit for this is kJ/mol. Sometimes you see older references that use kcal/mol. The conversion
between calories and Joules is 4.184 J = 1.000 cal.
Sometimes you also see this number expressed "per gram" rather than "per mole." For example,
water's molar heat of fusion is 6.02 kJ/mol. Expressed per gram, it is 334.16 J/g.
Typically, the term "heat of fusion" is used with the "per gram" value.
72.0 grams of solid water is 0.0 °C. It is going to melt AND stay at zero degrees. This is an important
point. While the ice melts, its temperature will remain the same. We need to calculate the energy
needed to do this.
This summarizes the information needed:
ΔHf = 6.02 kJ/mol
The mass = 72.0 g
The molar mass of H2O = 18.0 gram/mol
The calculation needed, using words & symbols is:
q = (moles of water) (ΔHf)
We can rewrite the moles of water portion and make the equation like this:
q = (grams water / molar mass of water) (ΔHf)
Why is this equation the way it is?
Think about one mole of ice. That amount of ice (one mole or 18.0 grams) needs 6.02 kilojoules of
energy to melt. Each mole of ice needs 6.02 kilojoules. So the (grams water / molar mass of water) in
the above equation calculates the amount of moles.
With the numbers in place, we have:
q = (72.0 g / 18.0 g mol¯1 ) (6.02 kJ / mol)
So we calculate and get 24.08 kJ. We won't bother to round off right now since there are three more
calculations to go. We're doing the second step now. When all five are done, we'll sum them all up.

Step Three: liquid water rises in temperature
Once the ice is totally melted, the temperature can now
begin to rise again.
It continues to go up until it reaches its normal boiling
point of 100.0 °C.
Since the temperature went from zero to 100, the Δt is
100.
Here is an important point: THE LIQUID HAS NOT
BOILED YET.
At the end of this step we have liquid water at 100 degrees. It has not turned to steam yet.
Each gram of water requires a constant amount of energy to go up each degree Celsius. This amount
of energy is called specific heat and has the symbol c. There will be a different value needed,
depending on the substance being in the solid, liquid or gas phase.
72.0 grams of liquid water is 0.0 °C. It is going to warm up to 100.0 °C, but at that temperature, the
water WILL NOT BOIL. We need to calculate the energy needed to do this.
This summarizes the information needed:
Δt = 100.0 °C (100.0 °C – 0.0 °C)
The mass = 72.0 g
c = 4.184 Joules per gram-degree Celsius
The calculation needed, using words & symbols is:
q = (mass) (Δt) (c)
Why is this equation the way it is?
Think about one gram going one degree. The liquid water needs 4.184 J for that. Now go the second
degree. Another 4.184 J. Go the third degree and use another 4.184 J. So one gram going 100
degrees needs 4.184 x 100 = 418.4 J. Now we have 72 grams, so gram #2 also needs 418.4, gram
#3 needs 418.4 and so on until 72 grams.
With the numbers in place, we have:
q = (72.0 g) (100.0 °C) (4.184 J/g °C)
So we calculate and get 30124.8 J. We won't bother to round off right now since there are two more
calculations to go. We will have to do five calculations and then sum them all up.
Notes:

Step Four: liquid water boils
Now, we continue to add energy and the water begins to
boil.
However, the temperature DOES NOT CHANGE. It
remains at 100 during the time the water boils.
Each mole of water will require a constant amount of
energy to boil. That amount is named the molar heat of
vaporization and its symbol is ΔH. The molar heat of
vaporization is the energy required to boil one mole of a
substance at its normal boiling point. One mole of liquid water, one mole of liquid benzene, one mole
of liquid lead. It does not matter. Each substance has its own value.
During this time, the energy is being used to overcome water molecules' attraction for each other,
allowing them to move from close together (liquid) to quite far apart (the gas state).
The unit for this is kJ/mol. Sometimes you see older references that use kcal/mol. The conversion
between calories and Joules is 4.184 J = 1.000 cal.
Typically, the term "heat of vaporization" is used with the "per gram" value.
72.0 grams of liquid water is at 100.0 °C. It is going to boil AND stay at 100 degrees. This is an
important point. While the water boils, its temperature will remain the same. We need to calculate the
energy needed to do this.
This summarizes the information needed:
ΔH = 40.7 kJ/mol
The mass = 72.0 g
The molar mass of H2O = 18.0 gram/mol
The calculation needed, using words & symbols is:
q = (moles of water) (ΔH)
We can rewrite the moles of water portion and make the equation like this:
q = (grams water / molar mass of water) (ΔH)
Why is this equation the way it is?
Think about one mole of liquid water. That amount of water (one mole or 18.0 grams) needs 40.7
kilojoules of energy to boil. Each mole of liquid water needs 40.7 kilojoules to boil. So the (grams
water / molar mass of water) in the above equation calculates the amount of moles.
With the numbers in place, we have:
q = (72.0 g / 18.0 g mol¯1 ) (40.7 kJ / mol)
So we calculate and get 162.8 kJ. We won't bother to round off right now since there is one more
calculation to go. We're doing the fourth step now. When all five are done, we'll sum them all up.

Step Five: steam rises in temperature
Once the water is completely changed to steam, the
temperature can now begin to rise again.
It continues to go up until we stop adding energy. In this
case, let the temperature rise to 120 °C.
Since the temperature went from 100 °C to 120°C, the Δt
is 20°C.
Each gram of water requires a constant amount of energy
to go up each degree Celsius. This amount of energy is called specific heat and has the symbol c.
There will be a different value needed, depending on the substance being in the solid, liquid or gas
phase.
72.0 grams of steam is 100.0 °C. It is going to warm up to 120.0 °C. We need to calculate the energy
needed to do this.
This summarizes the information needed:
Δt = 20 °C
The mass = 72.0 g
c = 2.02 Joules per gram-degree Celsius
The calculation needed, using words & symbols is:
q = (mass) (Δt) (c)
Why is this equation the way it is?
Think about one gram going one degree. The liquid water needs 2.02 J for that. Now go the second
degree. Another 2.02 J. Go the third degree and use another 2.02 J. So one gram going 20 degress
needs 2.02 x 20 = 44 J. Now we have 72 grams, so gram #2 also needs 44, gram #3 needs 44 and
so on until 72 grams.
I hope that helped.
With the numbers in place, we have:
q = (72.0 g) (20 °C) (2.02 J/g °C)
So we calculate and get 2908.8 J. We won't bother to round off right now since we still need to sum
up all five values.
Notes:

The following table summarizes the five steps and their results. Each step number is a link back to
the explanation of the calculation.
Converting to kJ gives us this:
1.4832 kJ
24.08 kJ
30.1248 kJ
162.8 kJ
2.9088 kJ
Step q 72.0 g of H2O
1 1483.2 J Δt = 10 (solid)
2 24.08 kJ melting
3 30124.8 J Δt = 100 (liquid)
4 162.8 kJ boiling
5 2908.8 J Δt = 20 (gas)
Summing up gives 221.3968 kJ and proper significant digits gives us 221.4 kJ for the answer.
Notice how all units were converted to kJ before continuing on. Joules is a perfectly fine unit; it's just
that 221,396.8 J is an awkward number to work with. Usually Joules is used for values under 1000,
otherwise kJ is used.
By the way, on other sites you may see kj used for kilojoules. I've also seen Kj used. Both of these
are wrong symbols. kJ is the only correct symbol.
Enthalpy
When a process occurs at constant pressure, the heat evolved (either released or absorbed) is equal
to the change in enthalpy. Enthalpy (H) is the sum of the internal energy (U) and the product of
pressure and volume (PV) given by the equation:
H=U+PV
When a process occurs at constant pressure, the heat evolved (either released or absorbed) is equal
to the change in enthalpy. Enthalpy is a state function which depends entirely on the state
functions T, P and U. Enthalpy is usually expressed as the change in enthalpy (ΔH) for a process
between initial and final states:
ΔH=ΔU+ΔPVΔ
If temperature and pressure remain constant through the process and the work is limited to pressurevolume
work, then the enthalpy change is given by the equation:
ΔH=ΔU+PΔV
Also at constant pressure the heat flow (q) for the process is equal to the change in enthalpy defined
by the equation:
ΔH=q
By looking at whether q is exothermic or endothermic we can determine a relationship between ΔH
and q. If the reaction absorbs heat it is endothermic meaning the reaction consumes heat from the
surroundings so q>0 (positive). Therefore, at constant temperature and pressure, by the equation
above, if q is positive then ΔH is also positive. And the same goes for if the reaction releases heat,

then it is exothermic, meaning the system gives off heat to its surroundings, so q

The Learning Goal for this section is: Explain how various factors, such as concentration, temperature, and
presence of a catalyst affects the rate of a chemical reaction.
You are going to answer these 15 questions first in the order they are given to you. This will be a quiz grade. You are then going to explain the
chemical concepts that are used in each question. You can do this in the order given or group them by concepts. You can use your book or information
found on the internet but all information must be written in your own word. The font needs to be Arial 12. This is due on Monday April 24 by midnight in
the drop box. This document should be a 6-page pdf.
1.B 6.B 11.A
2.C 7.B 12.D
3.C 8.D 13.C
4.D 9.A 14.A
5.A 10.B 15.D
1 This graph represents the change in energy for two laboratory trials of the same reaction.
Which factor could explain the energy difference between the trials?
A Heat was added to trial #2.
B A catalyst was added to trial #2.
C Trial #1 was stirred.
D Trial #1 was cooled.
One of the four factors that can influence the rate of a chemical reaction is a catalyst.
Trial number two is an exothermic reaction releasing heat.
Catalysts speed up reactions, meaning the substance doesn’t need to use so much energy to carry
out the reaction it is making
2 Consider this balanced chemical equation:
Which will increase the rate of the reaction?
A increasing pressure on the reaction
2H2O2 (aq) → 2H2O(l) + O2 (g)

B decreasing concentration of the reactants
C adding a catalyst to the reaction
D decreasing the temperature of the reaction
A catalyst is not consumed during a reaction
What is a catalyst?
A catalyst is a substance which speeds up a reaction, but is chemically unchanged at the end of the
reaction. When the reaction has finished, it’s the same as it was when you started.
3 For the reaction
A + (aq) + B — (aq) → AB (s)
increasing the temperature increases the rate of the reaction. Which is the best explanation for this
happening?
A The pressure increases, which in turn increases the production of products.
B The concentration of reactants increases with an increase in temperature.
C The average kinetic energy increases, so the likelihood of more effective collisions between ions
increases.
D Systems are more stable at high temperatures.
The rate constants are the catalyst and the temperature
What is collision theory?
The Collison Theory is when atoms, ions, and molecules form products by colliding with one another.
It explains why some reactants are extremely slow at room temperature.
If having enough kinetic energy those who do not have enough energy will bounce apart, unchanged
What is activation energy?
Activation Energy – minimum energy that particles who collide must have to react
4 Which statement explains why the speed of some reactions is increased when the surface area of
one or all the reactants is increased?
A increasing surface area changes the electronegativity of the reactant particles
B increasing surface area changes the concentration of the reactant particles
C increasing surface area changes the conductivity of reactant particles
D increasing surface area enables more reactant particles to collide
Particles can only react when they collide. If you heat a substance, the particles move faster and so
collide more frequently. That will speed up the rate of reaction.
Collisions only result in a reaction if the particles collide with enough energy to get the reaction
started. This minimum energy required is called the activation energy for the reaction
catalyst
C6H6 + Br2 → C6H5Br + HBr
5 Which of the following changes will cause an increase in the rate of the above reaction?
A increasing the concentration of Br2
B decreasing the concentration of C6H6
C increasing the concentration of HBr
D decreasing the temperature
What is concentration?
Concentration is the amount of a substance that is defined by it SA
Trying to squeeze particles into a certain spot will increase concentration.
Increasing the concentration of a reactant will increase the rate of the reaction.
Increasing the concentration of a product will not increase the rate of the reaction.

2CO + O2 → 2CO2
6 If the above reaction takes place inside a sealed reaction chamber, then which of these procedures
will cause a decrease in the rate of reaction?
A raising the temperature of the reaction chamber
B increasing the volume inside the reaction chamber
C removing the CO2 as it is formed
D adding more CO to the reaction chamber
increasing the volume of the reaction chamber will cause more products to form
What is a reaction chamber?
It is a chamber that stores matter. The chamber measures how much matter there is before and after
the reaction.
By increasing the volume of the chamber the particles will be farther apart thus decreasing the rate of
reaction
7 A catalyst can speed up the rate of a given chemical reaction by
A increasing the equilibrium constant in favor of products.
B lowering the activation energy required for the reaction to occur.
C raising the temperature at which the reaction occurs.
D increasing the pressure of reactants, thus favoring products.
The catalyst does not change the total free energy from reactants to products
To increase the rate of a reaction you need to increase the number of successful collisions. One
possible way of doing this is to provide an alternative way for the reaction to happen which has a
lower activation energy.
Having a lower activation energy can speed up the reaction rate.
A catalyst provides an alternative route for the reaction. That alternative route has a lower activation
energy
8 Which reaction diagram shows the effect of using the appropriate catalyst in a chemical reaction?
A
C
B
D

Letters A and C don’t fluctuate more than once as if the energy lowered then peaked, then lowered again,
yet raised back up then down, this would only happen if there were two reactions happening.
As it demonstrates in the graph the activation energy without a catalyst is much higher than the activation
energy with a catalyst, with a catalyst it has lower activation energy.
9 H2O2, hydrogen peroxide, naturally breaks down into H2O and O2 over time. MnO2, manganese
dioxide, can be used to lower the energy of activation needed for this reaction to take place and, thus,
increase the rate of reaction. What type of substance is MnO2?
A a catalyst
B an enhancer
C an inhibitor
D a reactant
Why is manganese dioxide breaking down the hydrogen peroxide?
Ho22 hydrogen peroxide naturally breaks down into ho2 and o2 over time mno2 manganese dioxide
can be used to lower the energy of activation needed for this reaction to take place and thus increase
the rate of reaction what type of substance is mno 2
What is an inhibitor?
An inhibitor is a substance that comes between the action of a catalyst.
Disclosure: if particles collide with enough energy they can still react the same way as if the catalyst
wasn't there.
10 When a reaction is at equilibrium and more reactant is added, which of the following changes is
the immediate result?
A The reverse reaction rate remains the same.
B The forward reaction rate increases.
C The reverse reaction rate decreases.
D The forward reaction rate remains the same.
What is equilibrium?
Equilibrium is when concentrations of reactants and products are in an unchanging ratio,
It’s when forward reactions occur at equal rates.
What is the reverse reaction?
the forward and reverse reactions are in balance at equilibrium, so there is no net change in the
concentrations of the reactants or products
by adding more reactant, the forward reaction would be the one to decrease.
It would not have anything to do with a reverse reaction because there is nothing being taken out of
the reaction.
A reverse reaction is when “reactants form products that, in turn, react together to give the reactants
back”, and that is clearly not what is going on in this problem.
11 In which of the following reactions involving gases would the forward reaction be favored by an
increase in pressure?
A A + B ⇄AB
B A + B ⇄ C + D
C 2A + B ⇄ C + 2D
D AC ⇄ A + C
The answer is letter A because an increase in pressure would mean that the equation is just the
simplest form for reacting.
Letter C does not even have an equivalent equation.
Letter D would be a reverse reaction because the product is given before the reactants are given.

12 Which action will drive the reaction to the right?
A heating the equilibrium mixture
B adding water to the system
C decreasing the oxygen concentration
D increasing the system’s pressure
4HCl(g) + O2(g) ⇄ 2H2O(l) + 2Cl2(g) + 113 kJ
by increasing the system’s pressure, there is a change being made to the equation, which would
move it from one side of the equation to the other, such as the right side which is where this is
wanting to be moved to.
By heating the mixture, you would be increasing the temperature but that would not mean that the
reactants would form the product they do.
NO2(g) + CO(g) ⇄ NO(g) + CO2(g)
13 The reaction shown above occurs inside a closed flask. What action will shift the reaction to the
left?
A pumping CO gas into the closed flask
B raising the total pressure inside the flask
C increasing the NO concentration in the flask
D venting some CO2 gas from the flask
by increasing the concentration in the flask, that would make the reaction happen on the left side of
the equation
NH4Cl(s) + heat ⇄ NH3(g) + HCl(g)
14 What kind of change will shift the reaction above to the right to form more products?
A a decrease in total pressure
B an increase in the concentration of HCl
C an increase in the pressure of NH3
D a decrease in temperature
What is total pressure?
The total pressure is the force per unit area that is felt when a flowing fluid is brought to rest and is
usually measured with a pitot tube type instrument, the total pressure is the sum of the static pressure
and the dynamic pressure.
increasing the pressure or concentration, the reaction would be moved to the left, so if it wants to be
moved to the right you would have to decrease the pressure
Static pressure is felt when the fluid is at rest or when the measurement is taken when traveling along
with the fluid flow. The difference between the total and static pressure is the dynamic
pressure, which represents the kinetic energy of the flowing fluid. Dynamic pressure is a function of
the fluid velocity and its density and can be calculated from

15 In a sealed bottle that is half full of water, equilibrium will be attained when water molecules
A cease to evaporate.
B begin to condense.
C are equal in number for both the liquid and the gas phase.
D evaporate and condense at equal rates.
Equilibrium is when two competing forces are balanced
when water molecules are evaporating but they’re are being condensed at the same rate, it’s as if
nothing is happening to the water molecules.

Unit 8
Chapter 19 Acid and Bases
The student will learn what are the different ways chemists
define aids and bases, what the pH of a solution means and
how chemist use acid-base reactions.
Relate acidity and basicity to hydronium and hydroxyl ion concentration and
pH.
Students will be able to use a pH scale to identify substances as acids or bases.
Students will be able to use various equipment (probeware, universal pH, etc.) to
identify the pH of substances.
Students will be able to calculate H3O+ and OH- concentration of various
substances.
pH scale
Hydronium ion
Arrhenius acid/base
Lewis acid/base
Bronsted-Lowry acid/base
Strong acid/base
Weak acid/base
Neutralization reaction
Titration
Chapter 20 Oxidation-Reduction Reactions
The student will learn what happens during oxidation and
reduction and how to balance redox equations.
Describe oxidation-reduction reactions in living and non-living systems.
Students will be able to compare and contrast redox reactions.
Students will be able to assign oxidation numbers to redox reactions.
Students will be able to write half reactions
Oxidation

Reduction
Oxidation reduction reaction
Oxidation number
Half reaction
Electrochemical process
Battery
Cathode
Anode
Electrolysis

The Learning Goal for this section is:
Acids and Bases
The Observable Properties of Acids and Bases
The words acid and alkaline (an older word for base) are derived from direct sensory experience.
Acid Property #1:
The word acid comes from the Latin word acere, which means "sour." All acids taste sour. Well
known from ancient times were vinegar, sour milk and lemon juice. Aspirin (scientific name:
acetylsalicylic acid) tastes sour if you don't swallow it fast enough. Other languages derive their word
for acid from the meaning of sour. So, in France, we have acide. In Germany, we have säure from
saure and in Russia, kislota from kisly.
Base Property #1:
The word "base" has a more complex history (see below) and its name is not related to taste. All
bases taste bitter. For example, mustard is a base. It tastes bitter. Many medicines, because they are
bases, taste bitter. This is the reason cough syrups are advertised as having a "great grape taste."
The taste is added in order to cover the bitterness of the active ingredient in cough syrup.
Acid Property #2:
Acids make a blue vegetable dye called litmus turn red.
Base Property #2:
Bases are substances which will restore the original blue color of litmus after having been reddened
by an acid.
Acid Property #3:
Acids destroy the chemical properties of bases.
Base Property #3:
Bases destroy the chemical properties of acids.
Neutralization is the name for this type of reaction.
Acid Property #4:
Acids conduct an electric current.
Base Property #4:
Bases conduct an electric current.
This is a common property shared with salts. Acids, bases and salts are grouped together into a
category called electrolytes, meaning that a water solution of the given substance will conduct an
electric current.
Non-electrolyte solutions cannot conduct a current. The most common example of this is sugar
dissolved in water.

So far, the properties have an obvious relationship: taste, color change, mutual destruction, and
response to electric current. This last property is related, but in a less obvious way. The property
below identifies a unique chemical reaction that acids and bases engage in.
Acid Property #5:
Upon chemically reacting with an active metal, acids will evolve hydrogen gas (H2). The key word, of
course, is active. Some metals, like gold, silver or platinum, are rather unreactive and it takes rather
extreme conditions to get these "unreactive" metals to react. Not so with the metals in this property.
They include the alkali metals (Group I, Li to Rb), the alkaline earth metals (Group II, Be to Ra), as
well as zinc and aluminum. Just bring the acid and the metal together at anything close to room
temperature and you get a reaction. Here's a sample reaction:
Zn + 2 HCl(aq) ---> ZnCl2 + H2
Another common acid reaction some sources mention is that acids react with carbonates (and
bicarbonates) to give carbon dioxide gas:
HCl + NaCO3 ---> CO2 + H2O + NaCl
Base Property #5:
Bases feel slippery, sometimes people say soapy. This is because they dissolve the fatty acids and
oils from your skin and this cuts down on the friction between your fingers as you rub them together.
In essence, the base is making soap out of you. Yes, bases are involved in the production of soap! In
the early years of soap making, the soaps were very harsh on the skin and clothes due to the high
base content. Even today, people with very sensitive skin must sometimes use a non-soap-based
product for bathing.
It was not until more modern times that the chemical nature (as opposed to observable properties) of
acids and bases began to be explored. That leads to this property that is not directly observable by
the senses.
Acid Property #6:
Acids produce hydrogen ion (H + ) in solution. A more correct formula for what is produced is that of the
hydronium ion, H3O + . Both formulas are used interchangeably.
Acid base theories: Svante Arrhenius
I. Introduction
The basic idea is that certain substances remain ionized in solution all the time. Today, everyone
accepts this without question, but it was the subject of much dissention and disagreement in 1884,
when a twenty-five-year-old Arrhenius presented and defended his dissertation.
II. The Acid Base Theory
Acid - any substance which delivers hydrogen ion (H + ) to the solution.
Base - any substance which delivers hydroxide ion (OH¯) to the solution.
Here is a generic acid dissociating, according to Arrhenius:
HA ---> H + + A¯

This would be a generic base:
XOH ---> X + + OH¯
When acids and bases react according to this theory, they neutralize each other, forming water and a
salt:
HA + XOH ---> H2O + XA
Keeping in mind that the acid, the base and the salt all ionize, we can write this:
Finally, we can drop all spectator ions, to get this:
H + + A¯ + X + + OH¯ ---> H2O + X + + A¯
H + + OH¯ ---> H2O
These ideas covered all of the known acids at the time (the usual suspects like hydrochloric acid,
acetic acid, and so on) and most of the bases (sodium hydroxide, potassium hydroxide, calcium
hydroxide and so on). HOWEVER, and it is a big however, the theory did not explain why ammonia
(NH3) was a base. There are other problems with the theory also.
III. Problems with Arrhenius' Theory
1. The solvent has no role to play in Arrhenius' theory. An acid is expected to be an acid in any
solvent. This was found to not be the case. For example, HCl is an acid in water, behaving in
the manner Arrhenius expected. However, if HCl is dissolved in benzene, there is no
dissociation, the HCl remaining as un-dissociated molecules. The nature of the solvent plays a
critical role in acid-base properties of substances.
2. All salts in Arrhenius' theory should produce solutions that are neither acidic or basic. This is
not the case. If equal amounts of HCl and ammonia react, the solution is slightly acidic. If equal
amounts of acetic acid and sodium hydroxide are reacted, the resulting solution is basic.
Arrhenius had no explanation for this.
3. The need for hydroxide as the base led Arrhenius to propose the formula NH4OH as the
formula for ammonia in water. This led to the misconception that NH4OH is the actual base,
not NH3.
In fact, by 1896, several years before Arrhenius announced his theory, it had been recognized that
characteristic base properties where just as evident in such solvents as aniline, where no hydroxide
ions were possible.
4. H + , a bare proton, does not exist for very long in water. The proton affinity of H2O is about 799
kJ/mol. Consequently, this reaction:
H2O + H + ---> H3O +
happens to a very great degree. The "concentration" of free protons in water has been estimated to
be 10¯130 M. A rather preposterous value, indeed.
The Arrhenius theory of acids and bases will be fully supplanted by the theory proposed
independently by Johannes and Thomas Lowry in 1923.

The acid base theory of Brønsted and Lowry
I. Introduction
In 1923, within several months of each other, Johannes Nicolaus Brønsted (Denmark) and Thomas
Martin Lowry (England) published essentially the same theory about how acids and bases behave.
Since they came to their conclusions independently of each other, both names have been used for
the theory name.
II. The Acid Base Theory
Using the words of Brønsted:
". . . acids and bases are substances that are capable of splitting off or taking up hydrogen ions,
respectively."
Or an acid-base reaction consists of the transfer of a proton from an acid to a base. KEEP THIS
THOUGHT IN MIND!!
Here is a more recent way to say the same thing:
An acid is a substance from which a proton can be removed.
A base is a substance that can remove a proton from an acid.
Remember: proton, hydrogen ion and H + all mean the same thing
Very common in the chemistry world is this definition set:
An acid is a "proton donor."
A base is a "proton acceptor."
In an acid, the hydrogen ion is bonded to the rest of the molecule. It takes energy (sometimes a little,
sometimes a lot) to break that bond. So the acid molecule does not "give" or "donate" the proton, it
has it taken away. In the same sense, you do not donate your wallet to the pickpocket, you have it
removed from you.
The base is a molecule with a built-in "drive" to collect protons. As soon as the base approaches the
acid, it will (if it is strong enough) rip the proton off the acid molecule and add it to itself.
Now this is where all the fun stuff comes in that you get to learn. You see, some bases are stronger
than others, meaning some have a large "desire" for protons, while other bases have a weaker drive.
It's the same way with acids, some have very weak bonds and the proton is easy to pick off, while
other acids have stronger bonds, making it harder to "get the proton."
One important contribution coming from Lowry has to do with the state of the hydrogen ion in solution.
In Brønsted's announcement of the theory, he used H + . Lowry, in his paper (actually a long letter to
the editor) used the H3O + that is commonly used today.
III. Sample Equations written in the Brønsted-Lowry Style
A. Reactions that proceed to a large extent:

HCl + H2O ⇌ H3O + + Cl¯
HCl - this is an acid, because it has a proton available to be transferred.
H2O - this is a base, since it gets the proton that the acid lost.
Now, here comes an interesting idea:
H3O + - this is an acid, because it can give a proton.
Cl¯ - this is a base, since it has the capacity to receive a proton.
Notice that each pair (HCl and Cl¯ as well as H2O and H3O + differ by one proton (symbol = H + ). These
pairs are called conjugate pairs.
HNO3 + H2O ⇌ H3O + + NO3¯
The acids are HNO3 and H3O + and the bases are H2O and NO3¯.
Remember that an acid-base reaction is a competition between two bases (think about it!) for a
proton. If the stronger of the two acids and the stronger of the two bases are reactants (appear on the
left side of the equation), the reaction is said to proceed to a large extent.
Here are some more conjugate acid-base pairs to look for:
H2O and OH¯
HCO3¯ and CO3 2¯
H2PO4¯ and HPO4 2¯
HSO4¯ and SO4 2¯
NH4 + and NH3
CH3NH3 + and CH3NH2
HC2H3O2 and C2H3O2¯
B. Reactions that proceed to a small extent:
If the weaker of the two acids and the weaker of the two bases are reactants (appear on the left side
of the equation), the reaction is said to proceed to only a small extent:
HC2H3O2 + H2O ⇌ H3O + + C2H3O2¯
NH3 + H2O ⇌ NH4 + + OH¯
Identify the conjugate acid base pairs in each reaction.
HC 2H 3O 2 and C 2H 3O 2¯
is one conjugate pair.
H 2O and H 3O + is the other.
NH 3 and NH 4
+
is one pair.
H 2O and OH¯ is the other.
Notice that H 2O in the first equation is acting as a base and in the second equation is acting as an acid.

IV. Problems with the Theory
This theory works very nicely in all protic solvents (water, ammonia, acetic acid, etc.), but fails to
explain acid base behavior in aprotic solvents such as benzene and dioxane. That job will be left for a
more general theory, such as the Lewis Theory of Acids and Bases.
The Lewis theory of acids and bases
I. Introduction
Lewis gives his definition of an acid and a base:
"We are inclined to think of substances as possessing acid or basic properties, without having a
particular solvent in mind. It seems to me that with complete generality we may say that a basic
substance is one which has a lone pair of electrons which may be used to complete the stable group
of another atom, and that an acid is one which can employ a lone pair from another molecule in
completing the stable group of one of its own atoms."
"In other words, the basic substance furnishes a pair of electrons for a chemical bond, the acid
substance accepts such a pair."
It is important to make two points here:
1. NO hydrogen ion need be involved.
2. NO solvent need be involved.
The Lewis theory of acids and bases is more general than the "one sided" nature of the Bronsted-
Lowry theory. Keep in mind that Bronsted-Lowry, which defines an acid as a proton donor and a base
as a proton acceptor, REQUIRES the presence of a solvent, specifically a protic solvent, of which
water is the usual example. Since almost all chemistry is done in water, the fact that this limits the
Bronsted-Lowry definition is of little practical consequence.
The Lewis definitions of acid and base do not have the constraints that the Bronsted-Lowry theory
does and, as we shall see, many more reactions were seen to be acid base in nature using the Lewis
definition than when using the Bronsted-Lowry definitions.
II. The Acid Base Theory
The modern way to define a Lewis acid and base is a bit more concise than above:
Acid: an electron acceptor.
Base: an electron donor.
A "Lewis acid" is any atom, ion, or molecule which can accept electrons and a "Lewis base" is any
atom, ion, or molecule capable of donating electrons. However, a warning: many textbooks will say
"electron pair" where I have only written "electron." The truth is that it sometimes is an electron pair
and sometimes it is not.
It turns out that it may be more accurate to say that "Lewis acids" are substances which are electrondeficient
(or low electron density) and "Lewis bases" are substances which are electron-rich (or high
electron density).

Several categories of substances can be considered Lewis acids:
1. positive ions
2. having less than a full octet in the valence shell
3. polar double bonds (one end)
4. expandable valence shells
Several categories of substances can be considered Lewis bases:
1. negative ions
2. one of more unshared pairs in the valence shell
3. polar double bonds (the other end)
4. the presence of a double bond
Sören Sörenson and the pH scale
I. Short Historical Introduction
In the late 1880's, Svante Arrhenius proposed that acids were substances that delivered hydrogen ion
to the solution. He has also pointed out that the law of mass action could be applied to ionic
reactions, such as an acid dissociating into hydrogen ion and a negatively charged anion.
This idea was followed up by Wilhelm Ostwald, who calculated the dissociation constants (the
modern symbol is Ka) of many weak acids. Ostwald also showed that the size of the constant is a
measure of an acid's strength.
By 1894, the dissociation constant of water (today called Kw) was measured to the modern value of
1 x 10¯14 .
In 1904, H. Friedenthal recommended that the hydrogen ion concentration be used to characterize
solutions. He also pointed out that alkaline (modern word = basic) solutions could also be
characterized this way since the hydroxyl concentration was always 1 x 10¯14 ÷ the hydrogen ion
concentration. Many consider this to be the real introduction of the pH scale.
III. The Introduction of pH
Sörenson defined pH as the negative logarithm of the hydrogen ion concentration.
pH = - log [H + ]
Remember that sometimes H3O + is written, so
pH = - log [H3O + ]
means the same thing.
So let's try a simple problem: The [H + ] in a solution is measured to be 0.010 M. What is the pH?
The solution is pretty straightforward. Plug the [H + ] into the pH definition:
pH = - log 0.010
An alternate way to write this is:
pH = - log 10¯2
Since the log of 10¯2 is -2, we have:
pH = - (- 2)
Which, of course, is 2.

Let's discuss significant figures and pH.
Another sample problem: Calculate the pH of a solution in which the [H3O + ] is 1.20 x 10¯3 M.
For the solution, we have:
pH = - log 1.20 x 10¯3
This problem can be done very easily using your calculator. However, be warned about putting
numbers into the calculator.
So you enter (-), log, 1.20, X10 n , (-), 3, enter.
The answer, to the proper number of significant digits is: 2.921.
III. Significant Figures in pH
Here is the example problem: Calculate the pH of a solution where the [H + ] is 0.00100 M. (This could
also be a pOH problem. The point being made is the same.)
OK, you say, that's pretty easy, the answer is 3. After all, 0.00100 is 10¯3 and the negative log of 10¯3
is 3.
You would be graded wrong!! Why? Because the pH is not written to reflect the number of significant
figures in the concentration.
Notice that there are three sig figs in 0.00100. (Hopefully you remember significant figures, since you
probably studied them months ago before getting to acid base stuff. THEY ARE STILL IMPORTANT!)
So, our pH value should also reflect three significant figures.
However, there is a special rule to remember with pH (and pOH) values. The whole number portion
DOES NOT COUNT when figuring out how many digits to write down.
Let's phrase that another way: in a pH (and a pOH), the only place where significant figures are
contained is in the decimal portion.
So, the correct answer to the above problem is 3.000. Three sig figs and they are all in the decimal
portion, NOT (I repeat NOT) in the whole number portion.
Practice Problems
Convert each hydrogen ion concentration into a pH. Identify each as an acidic pH or a basic pH.
1. 0.0015
2. 5.0 x 10¯9
3. 1.0
4. 3.27 x 10¯4
5. 1.00 x 10¯12
6. 0.00010

1. 2.82
2. 8.30
3. 0.00
4. 3.485
5. 12.000
6. 4.00
Sörenson also just mentions the reverse direction. That is, suppose you know the pH and you want to
get to the hydrogen ion concentration ([H + ])?
Here is the equation for that:
[H + ] = 10¯pH
That's right, ten to the minus pH gets you back to the [H + ] (called the hydrogen ion concentration).
This is actually pretty easy to do with the calculator. Here's the sample problem: calculate the [H + ]
from a pH of 2.45.
This problem can be done very easily using your calculator. However, be warned about putting
numbers into the calculator.
So you enter 2nd, 10 x , (-), 2.45, enter.
The answer, to the proper number of significant digits is: .00355.
The pH of an acidic pond is 5. What is the hydrogen ion concentration (moles per liter)?
The answer is:
pH = -log (hydrogen ion concentration)
The answer was .00001. Thus, 5 = -log (.00001).
We'll take the formula that you started with (pH = -log([H+])) and work to the answer (solve for [H+]).
pH = - log ([H+]) Given.
pH = log ([H+] (-1) ) Since logarithms are like exponents, when you multiply a log by
something, you can just move it to the inside of log as an exponent.
10 pH = 10 log ([H+] (-1)) Take each side to tenth power.
10 pH = [H+] (-1) Since "log" is just another notation for "log base 10", when you
raise a log to the tenth power, the log cancels out.
[H+] = 10 (-pH)
Take the reciprocal of both sides.
That is the general form. To answer the specific question,
5 = - log ([H+])
5 = log ([H+] (-1) )
10 5 = [H+] (-1)
10 (-5) = [H+]
[H+]
= .00001 mol/L

On your calculator you would input 10, ^, (-), 5 and you would get 0.00001.
This is also the way to find the amount of OH + that are present in a base.
To find the pH: -log(concentration)
To find the concentration: 10 -pH
Define these terms:
pH scale
Hydronium ion
Arrhenius acid/base
Lewis acid/base
Bronsted-Lowry acid/base
Strong acid/base
Weak acid/base
Neutralization reaction
Titration

The Learning Goal for this assignment is:
1. Oxidation Numbers
Redox Reactions
Oxidation and reduction
Every atom, ion or polyatomic ion has a formal oxidation number associated with it. This value
compares the number of protons in an atom (positive charge) and the number of electrons assigned
to that atom (negative charge).
In many cases, the oxidation number reflects the actual charge on the atom, but there are many
cases where it does not. Think of oxidation numbers as a bookkeeping exercise simply to keep track
of where electrons go.
2. Reduction
Reduction means what it says: the oxidation number is reduced in reduction.
This is accomplished by adding electrons. The electrons, being negative, reduce the overall oxidation
number of the atom receiving the electrons.
3. Oxidation
Oxidation is the reverse process: the oxidation number of an atom is increased during oxidation.
This is done by removing electrons. The electrons, being negative, make the atom that lost them
more positive.
I use this mnemonic to help me remember which is which: LEO the lion says GER.
LEO = Loss of Electrons is Oxidation
GER = Gain of Electrons is Reduction
Another well-known mnemonic is this: OIL RIG
OIL = Oxidation Is Loss (of Electrons)
RIG = Reduction Is Gain (of Electrons)
Another way is to simply remember that reduction is to reduce the oxidation number. Therefore,
oxidation must increase the value.
4. Reduction-Oxidation Reactions
There are many chemical reactions in which one substance gets reduced in oxidation number
(reduction) while another participating substance gets increased in oxidation number (oxidation).
Such a reaction is called called a REDOX reaction. The RED, of course, comes from REDuction and
OX from OXidation. However, it is pronounced re-dox and not red-ox.
Here is a simple example of a redox reaction:

Ag+ + Cu ---> Ag + Cu2+
I have deliberately not balanced it and I have also written it in net ionic form. I have found that kids
studying redox get confused by net ionic form and how to change a full equation into net ionic form.
Redox equations need to be balanced but, except for the simplest ones, it cannot be done by
inspection (also called trial and error). I take that back, complex ones can be done by trial and error. It
typically takes quite a bit of work, especially when compared to how long it takes when the proper
technique is used.
There is a technique used to balance redox reactions. It is called "balancing by half-reactions." The
basic plan will be to split the full equation into two simpler parts (called half-reactions), balance them
following several standard steps, then recombine the balanced half-reactions into the final answer.
This is another technique called the "ion-electron method." I plan to ignore it.
Notes:
5. Some Definitions
Oxidizing Agent - that substance which oxidizes somebody else. It is reduced in the process.
Reducing Agent - that substance which reduces somebody else. It is oxidized in the process.
It helps me to remember these definitions by the opposite nature of what happens. By that, I mean
the oxidizing agent gets reduced and the reducing agent gets oxidized.
6. Rules for Assigning Oxidation Numbers
The Oxidation Number of an element corresponds to the number of electrons, e - , that an atom loses,
gains, or appears to use when joining with other atoms in compounds. In determining the Oxidation
Number of an atom, there are seven guidelines to follow:
1. The Oxidation Number of an individual atom is 0. This includes diatomic elements such as
O2 or others like P4 and S8
2. The total Oxidation Number of all atoms in: a neutral species is 0 and in an ion is equal to the
ion charge.
3. Group 1 metals have an Oxidation Number of +1 and Group 2 an Oxidation Number of +2
4. The Oxidation Number of fluorine is -1 in compounds
5. Hydrogen generally has an Oxidation Number of +1 in compounds, except hydrides
6. Oxygen generally has an Oxidation Number of -2 in compounds, except peroxides
7. In binary metal compounds, Group 17 elements have an Oxidation Number of -1, Group 16 of -
2, and Group 15 of -3.
Note: The sum of the Oxidation Number s is equal to zero for neutral compounds and equal to the
charge for polyatomic ion species.
Now, some examples:
1. What is the oxidation number of Cl in HCl?
Since H = +1, the Cl must be -1 (minus one).

2. What is the oxidation number of Na in Na2O?
Since O = -2, the two Na must each be +1.
3. What is the oxidation number of Cl in ClO¯?
The O is -2, but since a -1 must be left over, then the Cl is +1.
4. What is the oxidation number for each element in KMnO4?
K = +1 because KCl exists. We know the Cl = -1 because HCl exists.
O = -2 by definition
Mn = +7. There are 4 oxygens for a total of -8, K is +1, so Mn must be the rest.
5. What is the oxidation number of S in SO4 2¯
O = -2. There are four oxygens for -8 total. Since -2 must be left over, the S must = +6.
Please note that, if there is no charge indicated on a formula, the total charge is taken to be zero.
Practice Problems
Find oxidation numbers
1. N in NO3¯
2. C in CO3 2¯
3. Cr in CrO4 2¯
4. Cr in Cr2O7 2¯
5. Fe in Fe2O3
6. Pb in PbOH +
7. V in VO2 +
8. V in VO 2+
9. Mn in MnO4¯
10. Mn in MnO4 2¯
Notes:

7. Half Reactions
A half-reaction is simply one which shows either reduction OR oxidation, but not both. Here is the
example redox reaction used in a different file:
Ag + + Cu → Ag + Cu 2+
It has BOTH a reduction and an oxidation in it. That is why we call it a redox reaction, from REDuction
and OXidation.
What you must be able to do is look at a redox reaction and separate out the two half-reactions in it.
To do that, identify the atoms which get reduced and get oxidized. Here are the two half-reactions
from the above example:
Ag+ → Ag
Cu → Cu 2+
The silver is being reduced, its oxidation number going from +1 to zero. The copper's oxidation
number went from zero to +2, so it was oxidized in the reaction. In order to figure out the halfreactions,
you MUST be able to calculate the oxidation number of an atom.
Keep in mind that a half-reaction shows only one of the two behaviors we are studying. A single halfreaction
will show ONLY reduction or ONLY oxidation, never both in the same equation.
Also, notice that the reaction is read from left to right to determine if it is reduction or oxidation. If you
read the reaction in the opposite direction (from right to left) it then becomes the other of our two
choices (reduction or oxidation). For example, the silver half-reaction above is a reduction, but in the
reverse direction it is an oxidation, going from zero on the right to +1 on the left.
There will be times when you want to switch a half-reaction from one of the two types to the other. In
that case, rewrite the entire equation and swap sides for everything involved. If I needed the silver
half-reaction to be oxidation, I'd write Ag → Ag+ rather than just doing it mentally.
The next step is that both half-reactions must be balanced. However, there is a twist. When you
learned about balancing equation, you made equal the number of atoms of each element on each
side of the arrow. That still applies, but there is one more thing: the total amount of charge on each
side of the half-reaction MUST be the same.
When you look at the two half-reactions above, you will see they are already balanced for atoms with
one Ag on each side and one Cu on each side. So, all we need to do is balance the charge. To do
this you add electrons to the more positive side. You add enough to make the total charge on each
side become EQUAL.
To the silver half-reaction, we add one electron:
To the copper half-reaction, we add two electrons:
Ag+ + e¯ ---> Ag
Cu ---> Cu 2+ + 2e¯

One point of concern: notice that each half-reaction wound up with a total charge of zero on each
side. This is not always the case. You need to strive to get the total charge on each side EQUAL, not
zero.
One more point to make before wrapping this up. A half-reaction is a "fake" chemical reaction. It's just
a bookkeeping exercise. Half-reactions NEVER occur alone. If a reduction half-reaction is actually
happening (say in a beaker in front of you), then an oxidation reaction is also occurring. The two halfreactions
can be in separate containers, but they do have to have some type of "chemical
connection" between them.
Half-Reaction Practice Problems
Balance each half-reaction for atoms and charge:
1) Cl2 → Cl¯
2) Sn → Sn 2+
3) Fe 2+ → Fe 3+
4) I3¯ → I¯
5) ICl2¯ → I¯ (I'm being mean on this one. Hint: the iodine is the only thing reduced or oxidized.)
Separate each of these redox reactions into their two half-reactions (but do not balance):
6) Sn + NO3¯ →SnO2 + NO2
7) HClO + Co →Cl2 + Co 2+
8) NO2 →NO3¯ + NO
Here are the two half-reactions to be combined:
Here is the rule to follow:
Ag+ + e¯ → Ag
Cu → Cu 2+ + 2e¯
the total electrons MUST cancel when the two half-reactions are added.
Another way to say it:
the number of electrons in each half-reaction MUST be equal when the two half-reactions are
added.
What that means is that one (or both) equations must be multiplied through by a factor. The value of
the factor is selected so as to make the number of electrons equal.
In our example problem, the top reaction (the one with silver) must be multiplied by two, like this:
2Ag+ + 2e¯ → 2Ag

Notice that each separate substance is increased by the factor amount. Occasionally, a student will
multiply ONLY the electrons by the factor. That is incorrect.
When the two half-reactions are added, we get:
2Ag+ + 2e¯ + Cu → 2Ag + Cu 2+ + 2e¯
With two electrons on each side, they may be canceled, resulting in:
2Ag+ + Cu → 2Ag + Cu 2+
This is the correct answer. Notice that there are two silvers on each side and one copper. Notice also
that the total charge on each side is +2. It is balanced for both atoms and charge. Sometimes, I am
asked if the order matters, if the Cu could be first on the left-hand side. The answer is that the order
does not matter. There happen to be certain styles about where particular substances are put in the
final answer, but these are only styles. They do not affect the chemical correctness of the answer.
Notes:
Practice Problems
Separate into half-reactions, balance them and then recombine.
1) Sn + Cl2 ---> Sn 2+ + Cl¯
2) Fe 2+ + I3¯ ---> Fe 3+ + I¯
1. N in NO 3¯
The O is -2 and three of them makes -6. Since -1 is left over, the N must be +5
2. C in CO 3
2¯
The O is -2 and three of them makes -6. Since -2 is left over, the C must be +4
3. Cr in CrO 4
2¯
The O is -2 and four of them makes -8. Since -2 is left over, the Cr must be +6
4. Cr in Cr 2 O 7
2¯
The O is -2 and seven of them makes -14. Since -2 is left over, the two Cr must be +12
What follows is an important point. Each Cr is +6. Each one. We do not speak of Cr 2 being +12. Each Cr atom is considered individually.
5. Fe in Fe 2 O 3
The O is -2 and three of them makes -6. Each Fe must then be +3
6. +2 9. +7
7. +5 10. +6
8. +4
1) Cl 2 + 2e¯ →2Cl¯ 5) ICl 2¯
+ 2e¯ →I¯ + 2Cl¯
2) Sn →Sn 2+ + 2e¯ 6) Sn →SnO 2 and NO 3¯
→NO 2
3) Fe 2+ →Fe 3+ + e¯ 7) HClO →Cl 2 and Co ---> Co 2+
4) I 3¯ → 3I¯ + 2e¯ 8) NO 2 →NO 3¯
and NO 2 →NO
Note that in this last redox equation the NO 2 (actually just the N) is both being reduced AND being oxidized. In the first half-reaction, the N goes from +4
to +5 (oxidation) and in the second, the N goes from +4 to +2 (reduction).
By the way, we could flip the reaction so that NO 3¯
and NO are reacting together to produce only one product, the NO 2 . In that case, the two halfreactions
would be reversed.

Reference Tables for Physical Setting/CHEMISTRY
Table A
Standard Temperature and Pressure
Name Value Unit
Standard Pressure 101.3 kPa kilopascal
1 atm atmosphere
Standard Temperature 273 K kelvin
0°C degree Celsius
Table D
Selected Units
Symbol Name Quantity
m meter length
g gram mass
Pa pascal pressure
K kelvin temperature
Table B
Physical Constants for Water
mol
J
mole
joule
amount of
substance
energy, work,
quantity of heat
Heat of Fusion
Heat of Vaporization
Specific Heat Capacity of H 2
O()
Specific Heat Capacity of H 2
O(s)
Specific Heat Capacity of H 2
O(g)
Table C
Selected Prefixes
Factor Prefix Symbol
10 3 kilo- k
334 J/g
2260 J/g
4.18 J/g•K
2.10 J/g•K
2.01 J/g•K
s second time
min minute time
h hour time
d day time
y year time
L liter volume
ppm parts per million concentration
M
molarity
solution
concentration
u atomic mass unit atomic mass
10 –1 deci- d
10 –2 centi- c
10 –3 milli- m
10 –6 micro- μ
10 –9 nano- n
10 –12 pico- p
R1

Table E
Selected Polyatomic Ions
Formula Name Formula Name
H 3
O +
hydronium
CrO 4
2–
chromate
Hg 2
2+
mercury(I)
Cr 2
O 7
2–
dichromate
NH 4
+
C 2
H 3
O
–
2 –}
CH 3
COO
CN –
CO 3
2–
HCO
–
3
C 2
O
2–
4
ClO –
ammonium
acetate
cyanide
carbonate
hydrogen
carbonate
oxalate
hypochlorite
MnO 4
–
NO
–
2
NO
–
3
O
2–
2
OH –
PO 4
3–
SCN –
SO 3
2–
permanganate
nitrite
nitrate
peroxide
hydroxide
phosphate
thiocyanate
sulfite
ClO 2
–
chlorite
SO 4
2–
sulfate
ClO 3
–
chlorate
HSO 4
–
hydrogen sulfate
ClO 4
–
perchlorate
S 2
O 3
2–
thiosulfate
Table F
Solubility Guidelines for Aqueous Solutions
Ions That Form
Soluble Compounds
Group 1 ions
(Li + , Na + , etc.)
ammonium (NH 4 + )
nitrate (NO 3 – )
acetate (C 2
H 3
O 2 – or
CH 3
COO – )
hydrogen carbonate
(HCO 3 – )
chlorate (ClO 3 – )
halides (Cl – , Br – , I – )
Exceptions
when combined with
Ag + , Pb 2+ , or Hg 2
2+
sulfates (SO 4 2– ) when combined with Ag + ,
Ca 2+ , Sr 2+ , Ba 2+ , or Pb 2+
Ions That Form
Insoluble Compounds*
Exceptions
carbonate (CO 3 2– ) when combined with Group 1
ions or ammonium (NH 4 + )
chromate (CrO 4 2– ) when combined with Group 1
ions, Ca 2+ , Mg 2+ , or
ammonium (NH 4 + )
phosphate (PO 4 3– ) when combined with Group 1
ions or ammonium (NH 4 + )
sulfide (S 2– ) when combined with Group 1
ions or ammonium (NH 4 + )
hydroxide (OH – ) when combined with Group 1
ions, Ca 2+ , Ba 2+ , Sr 2+ , or
ammonium (NH 4 + )
*compounds having very low solubility in H 2 O
R2

150.
140.
Table G
Solubility Curves at Standard Pressure
KI
NaNO 3
130.
120.
KNO 3
110.
100.
Solubility (g solute/100. g H 2
O)
90.
80.
70.
60.
HCl
NH 4
Cl
KCl
50.
40.
30.
NaCl
KClO 3
NH 3
20.
10.
SO 2
0
0 10. 20. 30. 40. 50. 60. 70. 80. 90. 100.
Temperature (°C)
R3

Table H
Vapor Pressure of Four Liquids
200.
propanone
ethanol
150.
water
Vapor Pressure (kPa)
100.
101.3 kPa
ethanoic
acid
50.
0
0 25 50. 75 100. 125
R4

Table I
Heats of Reaction at 101.3 kPa and 298 K
Reaction
ΔH (kJ)*
CH 4
(g) + 2O 2
(g) CO 2
(g) + 2H 2
O() –890.4
C 3
H 8
(g) + 5O 2
(g) 3CO 2
(g) + 4H 2
O() –2219.2
2C 8
H 18
() + 25O 2
(g) 16CO 2
(g) + 18H 2
O() –10943
2CH 3
OH() + 3O 2
(g) 2CO 2
(g) + 4H 2
O() –1452
C 2
H 5
OH() + 3O 2
(g) 2CO 2
(g) + 3H 2
O() –1367
C 6
H 12
O 6
(s) + 6O 2
(g) 6CO 2
(g) + 6H 2
O() –2804
2CO(g) + O 2
(g) 2CO 2
(g) –566.0
C(s) + O 2
(g) CO 2
(g) –393.5
4Al(s) + 3O 2
(g) 2Al 2
O 3
(s) –3351
N 2
(g) + O 2
(g) 2NO(g) +182.6
N 2
(g) + 2O 2
(g) 2NO 2
(g) +66.4
2H 2
(g) + O 2
(g) 2H 2
O(g) –483.6
2H 2
(g) + O 2
(g) 2H 2
O() –571.6
N 2
(g) + 3H 2
(g) 2NH 3
(g) –91.8
2C(s) + 3H 2
(g) C 2
H 6
(g) –84.0
2C(s) + 2H 2
(g) C 2
H 4
(g) +52.4
2C(s) + H 2
(g) C 2
H 2
(g) +227.4
H 2
(g) + I 2
(g) 2HI(g) +53.0
KNO 3
(s) H 2 O K + (aq) + NO 3 – (aq) +34.89
NaOH(s) H 2 O Na + (aq) + OH – (aq) –44.51
NH 4
Cl(s) H 2 O NH 4 + (aq) + Cl – (aq) +14.78
NH 4
NO 3
(s) H 2 O NH 4 + (aq) + NO 3 – (aq) +25.69
NaCl(s) H 2 O Na + (aq) + Cl – (aq) +3.88
LiBr(s) H 2 O Li + (aq) + Br – (aq) –48.83
H + (aq) + OH – (aq) H 2
O() –55.8
*The ΔH values are based on molar quantities represented in the equations.
A minus sign indicates an exothermic reaction.
Most
Active
Least
Active
Table J
Activity Series**
Metals Nonmetals Most
Active
Li
Rb
K
Cs
Ba
Sr
Ca
Na
Mg
Al
Ti
Mn
Zn
Cr
Fe
Co
Ni
Sn
Pb
H 2
Cu
Ag
Au
F 2
Cl 2
Br 2
I 2
**Activity Series is based on the hydrogen
standard. H 2 is not a metal.
Least
Active
R5

Table K
Common Acids
Table N
Selected Radioisotopes
HCl(aq)
Formula
HNO 2
(aq)
HNO 3
(aq)
H 2
SO 3
(aq)
H 2
SO 4
(aq)
H 3
PO 4
(aq)
H 2
CO 3
(aq)
or
CO 2
(aq)
CH 3
COOH(aq)
or
HC 2
H 3
O 2
(aq)
Name
hydrochloric acid
nitrous acid
nitric acid
sulfurous acid
sulfuric acid
phosphoric acid
carbonic acid
ethanoic acid
(acetic acid)
Nuclide Half-Life Decay
Mode
Nuclide
Name
198 Au 2.695 d β – gold-198
14 C 5715 y β – carbon-14
37 Ca 182 ms β + calcium-37
60 Co 5.271 y β – cobalt-60
137 Cs 30.2 y β – cesium-137
53 Fe 8.51 min β + iron-53
220 Fr 27.4 s α francium-220
3 H 12.31 y β – hydrogen-3
131 I 8.021 d β – iodine-131
37 K 1.23 s β + potassium-37
42 K 12.36 h β – potassium-42
Table L
Common Bases
85 Kr 10.73 y β – krypton-85
16 N 7.13 s β – nitrogen-16
Formula
NaOH(aq)
KOH(aq)
Ca(OH) 2
(aq)
NH 3
(aq)
Name
sodium hydroxide
potassium hydroxide
calcium hydroxide
aqueous ammonia
19 Ne 17.22 s β + neon-19
32 P 14.28 d β – phosphorus-32
239 Pu 2.410 × 10 4 y α plutonium-239
226 Ra 1599 y α radium-226
222 Rn 3.823 d α radon-222
90 Sr 29.1 y β – strontium-90
Table M
Common Acid–Base Indicators
Approximate
Indicator pH Range Color
for Color Change
Change
methyl orange 3.1–4.4 red to yellow
bromthymol blue 6.0–7.6 yellow to blue
phenolphthalein 8–9 colorless to pink
litmus 4.5–8.3 red to blue
bromcresol green 3.8–5.4 yellow to blue
thymol blue 8.0–9.6 yellow to blue
99 Tc 2.13 × 10 5 y β – technetium-99
232 Th 1.40 × 10 10 y α thorium-232
233 U 1.592 × 10 5 y α uranium-233
235 U 7.04 × 10 8 y α uranium-235
238 U 4.47 × 10 9 y α uranium-238
Source: CRC Handbook of Chemistry and Physics, 91 st ed., 2010–2011,
CRC Press
Source: The Merck Index, 14 th ed., 2006, Merck Publishing Group
R6

Table O
Symbols Used in Nuclear Chemistry
Name Notation Symbol
alpha particle
4
2
He or 4 2 α α
beta particle
0
–1
e or 0
–1 β β–
gamma radiation
0
0
γ γ
neutron
1
0
n n
proton
1
1
H or 1 1 p p
positron
0
+1
e or 0
+1 β β+
Table P
Organic Prefixes
Prefix
meth- 1
eth- 2
prop- 3
but- 4
pent- 5
hex- 6
hept- 7
oct- 8
non- 9
dec- 10
Number of
Carbon Atoms
Table Q
Homologous Series of Hydrocarbons
Name General Examples
Formula Name Structural Formula
R7
alkanes C n
H 2n+2
ethane
alkenes C n
H 2n
ethene
alkynes C n
H 2n–2
ethyne
Note: n = number of carbon atoms
H H
H C C H
H H
H
H
C C
H
H
H C C H

Table R
Organic Functional Groups
Class of
Compound
Functional
Group
General
Formula
Example
halide
(halocarbon)
F (fluoro-)
Cl (chloro-)
Br (bromo-)
I (iodo-)
R X
(X represents
any halogen)
CH 3
CHClCH 3
2-chloropropane
alcohol
OH
R
OH
CH 3
CH 2
CH 2
OH
1-propanol
ether
O
R O R′
CH 3
OCH 2
CH 3
methyl ethyl ether
aldehyde
O
C H
R
O
C H
O
CH 3
CH 2
C H
propanal
ketone
O
C
O
R C R′
O
CH 3
CCH 2
CH 2
CH 3
2-pentanone
organic acid
O
C OH
R
O
C OH
O
CH 3
CH 2
C OH
propanoic acid
ester
O
C O
O
R C O R′
O
CH 3
CH 2
COCH 3
methyl propanoate
amine
N
R
R′
N R′′
CH 3
CH 2
CH 2
NH 2
1-propanamine
amide
O
C NH
R
O R′
C NH
O
CH 3
CH 2
C NH 2
propanamide
Note: R represents a bonded atom or group of atoms.
R8

0
6.941
+1
Li
3
2-1
Na
39.0983
K +1
19
2-8-8-1
85.4678 +1
Rb
Cs
(223)
Fr
87
-18-32-18-8-1
+1
Ra
88
-18-32-18-8-2
39
138.9055
La
57
2-8-18-18-9-2
+2 (227)
Ac
89
-18-32-18-9-2
47.867
Ti
22
2-8-10-2
91.224
Zr
40
2-8-18-10-2
+3 178.49
Hf
72
*18-32-10-2
+3 (261)
Rf
104
+2
+3
+4
+4
+4
50.9415
V
23
2-8-11-2
+2
+3
+4
+5
51.996
Cr
24
2-8-13-1
95.94
Mo
42
2-8-18-13-1
183.84
W
74
-18-32-12-2
+2
+3
+6
+6
+6
54.9380
Mn
25
2-8-13-2
+2
+3
+4
+7
55.845
Fe
26
2-8-14-2
+2
+3 58.9332
Co
27
2-8-15-2
+2
+3
58.693
Ni
28
2-8-16-2
+2
+3 63.546 Cu
2-8-18-1
107.868
Ag
47
2-8-18-18-1
79
+1
+2
+1
65.409
Zn
30
2-8-18-2
10.81
+3 12.011
B
5
2-3
26.98154
Al
13
2-8-3
+2 69.723
Ga
31
2-8-18-3
+3
+3
–4
+2
+4
C
6
2-4
28.0855
Si
14
2-8-4
72.64
Ge
32
2-8-18-4
Pb
–4
+2
+4
+2
+4
74.9216
As
33
2-8-18-5
Sb
–3
+3
15.9994 O
–2 18.9984
8
2-6 2-7
78.96
Se
34
2-8-18-6
127.60
Te
52
2-8-18-18-6
(209)
Po
84
-18-32-18-6
–2
+4
+6
–2
+4
+6
+2
+4
F
79.904
Br
35
2-8-18-7
126.904
l
53
2-8-18-18-7
(210)
At
85
-18-32-18-7
( ? )
Uus
117
4.00260 0
He
2
2
–1 20.180
Ne
10
2-8
0
22.98977
11
2-8-1
1
1.00794 +1
–1
H
1
1
1
37
2-8-18-8-1
–1
+1
+5
–1
+1
+5
+7
83.798
Kr
36
2-8-18-8
131.29
Xe
54
2-8-18-18-8
(222)
Rn
86
-18-32-18-8
0
+2
0
+2
+4
+6
0
132.905
55
2-8-18-18-8-1
Symbol
Relative atomic masses are based
Group on 12 C = 12 (exact)
Group
2
13 14 15 16 17 18
Atomic Number
+1
+1
9.01218 +2
Be
4
2-2
24.305
Mg
12
2-8-2
40.08
Ca
20
2-8-8-2
87.62
Sr
38
2-8-18-8-2
137.33
Ba
56
2-8-18-18-8-2
(226)
+2
+2
+2
+2
3
44.9559
Sc
21
2-8-9-2
88.9059
Y
2-8-18-9-2
+3
+3
4
KEY
92.9064
Nb +3
+5
41
2-8-18-12-1
180.948
Ta
73
-18-32-11-2
(262)
105
5
Periodic Table of the Elements
Atomic Mass
Electron Configuration
+4
Db
+5
6
(266)
Sg
106
12.011 2-4
–4
6
C
+2
+4
(98)
Tc
43
2-8-18-13-2
186.207
Re
75
-18-32-13-2
(272)
Bh
107
7
Group
+4
+6
+7
+4
+6
+7
8
101.07
Ru
44
2-8-18-15-1
190.23
Os
76
-18-32-14-2
(277)
Hs
108
+3
+3
+4
Selected Oxidation States
Note: Numbers in parentheses
are mass numbers of the most
stable or common isotope.
9
102.906
Rh
45
2-8-18-16-1
192.217
Ir
77
-18-32-15-2
(276)
Mt
109
+3
+3
+4
106.42
Pd
46
2-8-18-18
195.08
Pt
78
-18-32-17-1
+2
+4
+2
+4
196.967
Au
-18-32-18-1
(281)
Ds (280) Rg
110
10
29
111
11 12
+1
+3
112.41
Cd
48
2-8-18-18-2
200.59
Hg
80
-18-32-18-2
(285)
Cn
112
+2 114.818
In
+1
+2
49
2-8-18-18-3
204.383
Tl
81
-18-32-18-3
(284)
Uut
113**
+3
+1
+3
118.71
Sn
50
2-8-18-18-4
207.2
82
-18-32-18-4
(289)
Uuq
114
+2
+4
+2
+4
14.0067 –3
–2
N
–1
7
2-5
30.97376
P
15
2-8-5
121.760
51
2-8-18-18-5
208.980
Bi
83
-18-32-18-5
(288)
Uup
115
+1
+2
+3
+4
+5
–3
+3
+5
+5
–3
+3
+5
+3
+5
32.065
S
16
2-8-6
(292)
Uuh
116
–2
+4
+6
35.453
Cl
17
2-8-7
–1
+1
+5
+7
39.948
Ar
18
2-8-8
18
(294)
Uuo
118
140.116
Ce
58
232.038
Th
90
+3
+4
140.908
Pr +3
59
144.24
Nd
60
+4 231.036
Pa +4 238.029 +5
U +3
+4
+5
+6
91
92
+3
(145)
Pm
61
+3
150.36
Sm
62
+2
+3
151.964
Eu
63
+2
+3
157.25
Gd
64
+3
158.925
(237)Np (244) Pu (243) Am (247) Cm +3 (247) Bk +3
+3
+4
+5
+6
93 94
+3
+4
+5
+6
65
+3
+4
+5
+6
95 96 97
Tb
+3
+4
162.500
Dy
66
(251)
+3
164.930
Ho
67
+3
167.259
Er
68
Cf +3 (252) Es (257) Fm
100
98 99
+3
+3
+3
168.934
Tm +3
69
(258)
Md
101
+2
+3
173.04
Yb
70
(259)
No
102
+2
+3
+2
+3
174.9668
Lu
71
(262)
Lr
103
+3
+3
*denotes the presence of (2-8-) for elements 72 and above
**The systematic names and symbols for elements of atomic numbers 113 and above
will be used until the approval of trivial names by IUPAC.
Source: CRC Handbook of Chemistry and Physics, 91 st ed., 2010–2011, CRC Press
9
R9
Period
1
2
3
4
5
6
7

Table S
Properties of Selected Elements
First
Atomic Symbol Name Ionization
Electro- Melting Boiling* Density** Atomic
Number Energy negativity Point Point (g/cm 3 ) Radius
(kJ/mol) (K) (K) (pm)
1 H hydrogen 1312 2.2 14 20. 0.000082 32
2 He helium 2372 — — 4 0.000164 37
3 Li lithium 520. 1.0 454 1615 0.534 130.
4 Be beryllium 900. 1.6 1560. 2744 1.85 99
5 B boron 801 2.0 2348 4273 2.34 84
6 C carbon 1086 2.6 — — .— 75
7 N nitrogen 1402 3.0 63 77 0.001145 71
8 O oxygen 1314 3.4 54 90. 0.001308 64
9 F fluorine 1681 4.0 53 85 0.001553 60.
10 Ne neon 2081 — 24 27 0.000825 62
11 Na sodium 496 0.9 371 1156 0.97 160.
12 Mg magnesium 738 1.3 923 1363 1.74 140.
13 Al aluminum 578 1.6 933 2792 2.70 124
14 Si silicon 787 1.9 1687 3538 2.3296 114
15 P phosphorus (white) 1012 2.2 317 554 1.823 109
16 S sulfur (monoclinic) 1000. 2.6 388 718 2.00 104
17 Cl chlorine 1251 3.2 172 239 0.002898 100.
18 Ar argon 1521 — 84 87 0.001633 101
19 K potassium 419 0.8 337 1032 0.89 200.
20 Ca calcium 590. 1.0 1115 1757 1.54 174
21 Sc scandium 633 1.4 1814 3109 2.99 159
22 Ti titanium 659 1.5 1941 3560. 4.506 148
23 V vanadium 651 1.6 2183 3680. 6.0 144
24 Cr chromium 653 1.7 2180. 2944 7.15 130.
25 Mn manganese 717 1.6 1519 2334 7.3 129
26 Fe iron 762 1.8 1811 3134 7.87 124
27 Co cobalt 760. 1.9 1768 3200. 8.86 118
28 Ni nickel 737 1.9 1728 3186 8.90 117
29 Cu copper 745 1.9 1358 2835 8.96 122
30 Zn zinc 906 1.7 693 1180. 7.134 120.
31 Ga gallium 579 1.8 303 2477 5.91 123
32 Ge germanium 762 2.0 1211 3106 5.3234 120.
33 As arsenic (gray) 944 2.2 1090. — 5.75 120.
34 Se selenium (gray) 941 2.6 494 958 4.809 118
35 Br bromine 1140. 3.0 266 332 3.1028 117
36 Kr krypton 1351 — 116 120. 0.003425 116
37 Rb rubidium 403 0.8 312 961 1.53 215
38 Sr strontium 549 1.0 1050. 1655 2.64 190.
39 Y yttrium 600. 1.2 1795 3618 4.47 176
40 Zr zirconium 640. 1.3 2128 4682 6.52 164
R10

First
Atomic Symbol Name Ionization
Electro- Melting Boiling* Density** Atomic
Number Energy negativity Point Point (g/cm 3 ) Radius
(kJ/mol) (K) (K) (pm)
41 Nb niobium 652 1.6 2750. 5017 8.57 156
42 Mo molybdenum 684 2.2 2896 4912 10.2 146
43 Tc technetium 702 2.1 2430. 4538 11 138
44 Ru ruthenium 710. 2.2 2606 4423 12.1 136
45 Rh rhodium 720. 2.3 2237 3968 12.4 134
46 Pd palladium 804 2.2 1828 3236 12.0 130.
47 Ag silver 731 1.9 1235 2435 10.5 136
48 Cd cadmium 868 1.7 594 1040. 8.69 140.
49 In indium 558 1.8 430. 2345 7.31 142
50 Sn tin (white) 709 2.0 505 2875 7.287 140.
51 Sb antimony (gray) 831 2.1 904 1860. 6.68 140.
52 Te tellurium 869 2.1 723 1261 6.232 137
53 I iodine 1008 2.7 387 457 4.933 136
54 Xe xenon 1170. 2.6 161 165 0.005366 136
55 Cs cesium 376 0.8 302 944 1.873 238
56 Ba barium 503 0.9 1000. 2170. 3.62 206
57 La lanthanum 538 1.1 1193 3737 6.15 194
Elements 58–71 have been omitted.
72 Hf hafnium 659 1.3 2506 4876 13.3 164
73 Ta tantalum 728 1.5 3290. 5731 16.4 158
74 W tungsten 759 1.7 3695 5828 19.3 150.
75 Re rhenium 756 1.9 3458 5869 20.8 141
76 Os osmium 814 2.2 3306 5285 22.587 136
77 Ir iridium 865 2.2 2719 4701 22.562 132
78 Pt platinum 864 2.2 2041 4098 21.5 130.
79 Au gold 890. 2.4 1337 3129 19.3 130.
80 Hg mercury 1007 1.9 234 630. 13.5336 132
81 Tl thallium 589 1.8 577 1746 11.8 144
82 Pb lead 716 1.8 600. 2022 11.3 145
83 Bi bismuth 703 1.9 544 1837 9.79 150.
84 Po polonium 812 2.0 527 1235 9.20 142
85 At astatine — 2.2 575 — — 148
86 Rn radon 1037 — 202 211 0.009074 146
87 Fr francium 393 0.7 300. — — 242
88 Ra radium 509 0.9 969 — 5 211
89 Ac actinium 499 1.1 1323 3471 10. 201
Elements 90 and above have been omitted.
*boiling point at standard pressure
**density of solids and liquids at room temperature and density of gases at 298 K and 101.3 kPa
— no data available
Source: CRC Handbook for Chemistry and Physics, 91 st ed., 2010–2011, CRC Press
R11

Table T
Important Formulas and Equations
d = density
m
Density d = m = mass
V
V = volume
Mole Calculations number of moles =
given mass
gram-formula mass
measured value – accepted value
Percent Error % error = × 100
accepted value
mass of part
Percent Composition % composition by mass = × 100
mass of whole
mass of solute
parts per million = × 1000000
mass of solution
Concentration
molarity =
moles of solute
liter of solution
Combined Gas Law
P
P = pressure
1
V 1
P
= 2
V 2
V = volume
T 1
T 2 T = temperature
M A
= molarity of H + M B
= molarity of OH –
Titration M A
V A
= M B
V B
V A
= volume of acid V B
= volume of base
q = mCΔT q = heat H f
= heat of fusion
Heat q = mH f
m = mass H v
= heat of vaporization
q = mH v
C=specific heat capacity
ΔT = change in temperature
Temperature
K = °C + 273
K = kelvin
°C = degree Celsius
R12

Collier County CHEMISTRY EXAM FORMULA AND RESOURCE PACKET
GENERAL
D m V
[ ExperimentalValue AcceptedVa lue]
% error
x100
AcceptedVa lue
% yield
ExperimentalYield
TheoreticalYield
x100
CONCENTRATIONS
moles of solute
M = Molarity
liters of solution
KEY
P = pressure
V = volume
T = temperature
n = number of moles
m = mass
M = molar mass (grams/mole)
D = density
KE = kinetic energy
Avogadro’s Number = 6.02 x 10 23
GASES, LIQUIDS, SOLUTIONS
m = Molality
M1V1 M2V2
S1
P1
S 2
P 2
ACID/BASE
pH = - log[H + ]
[H + ]=10 -pH
moles of solute
kilograms of solvent
Gas constant
R 8.314 L kPa L atm L mmHg
0.0821 62.4
K mol K mol K mol
1 atm = 760 mmHg = 760 torr = 101.3 kPa
K = o C + 273
o C = K - 273
STP = Standard Temperature and Pressure = 0 o C
and 1 atm
P V
1 1
P2V
2
pOH = - log [OH - ]
[OH - ]= 10 -pOH
pH + pOH = 14
Kw = [H + ] x [OH - ] = 1.0 x 10 -14 M 2
V
T
1
1
P
T
1
1
V
T
P
T
2
2
2
2
Or V1T2 = V2T1
Or P1T2 = P2T1
THERMOCHEMISTRY
ΔH= mCΔT, where ΔT = T f - T
P1V
1
T
1
P2V
2
T
2
Or
P1V1T2=P2V2T1
q = mCΔT
PV
nRT
Specific Heat of Water = 4.18 J/g*˚C or 1.0 cal/g*˚C
Specific Heat of Ice = 2.1 J/g*˚C or 0.5 cal/g*˚C
Specific Heat of Steam = 2.0 J/g*˚C or 0.4 cal/g*˚C
P
Total
P
1
P
2
Rate A
Rate B
...
Molar MassB
Molar MassA
R13

CHEMISTRY EXAM FORMULA AND RESOURCE PACKET
Solubility of Compounds at 25 o C and 1 atm
acetate
bromide
carbonate
chlorate
chloride
hydroxide
iodide
nitrate
oxide
perchlorate
phosphate
sulfate
sulfide
aluminum S S - S S I S S I S I S d
ammonium S S S S S S S S - S S S S
barium S S I S S S S S sS S I I d
calcium S S I S S S S S sS S I sS I
copper(II) S S - S S I S S I S I S I
iron(II) S S I S S I S S I S I S I
iron(III) S S - S S I S S I S I sS d
lithium S S sS S S S S S S S sS S S
magnesium S S I S S I S S I S I S d
potassium S S S S S S S S S S S S S
silver sS I I S I - I S I S I sS I
sodium S S S S S S S S S S S S S
strontium S S I S S S S S S S I I I
zinc S S I S S I S S I S I S I
S=soluble
sS = slightly soluble in water
I = insoluble in water
d = decomposes in water
- = no such compound
R14

CHEMISTRY EXAM FORMULA AND RESOURCE PACKET
Common Polyatomic Ions
1- Charge 2- Charge 3- Charge
Formula Name Formula Name Formula Name
Dihydrogen
Phosphate
Hydrogen
phosphate
Phosphite
Acetate Oxalate Phosphate
Hydrogen
sulfite
Sulfite
Hydrogen
sulfate
Sulfate
Hydrogen
carbonate
Carbonate
Nitrite Chromate 1+ Charge
Nitrate Dichromate Formula Name
Cyanide Silicate Ammonium
Hydroxide
Permanganate
Hypochlorite
Chlorite
Chlorate
Perchlorate
R15

Activity Series of Metals
Name
Symbol
D
Lithium
Li
e
Potassium
K
c
r
Barium
Ba
e
Calcium
Ca
a
Sodium
Na
s
i
Magnesium
Mg
n
Aluminum
Al
g
Zinc
Zn
Iron
Fe
A
c
Nickel
Ni
t
Tin
Sn
i
v
Lead
Pb
i
(Hydrogen)
(H)*
t
Copper
Cu
y
Mercury
Hg
Silver
Ag
Gold
Au
*Metals from Li to Na will replace H from acids and water; from Mg to
Pb they will replace H from acids only.
Decreasing
Activity
Activity Series of Nonmetal (Halogens)
Name
Symbol
Fluorine F 2
Chlorine Cl 2
Bromine Br 2
Iodine I 2
R16