1st Semester Notebook Salazar

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

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

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

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.

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__________________________
Arnaldo Salazar-Gomez
PERIOD ________________________
2
PARENT NAME ____________________________
Cornelia Gomez-Hernandez
PARENT NUMBER _________________________
(239) 384-8095
SIGNATURE ____________________________
DATE ____________________________________
August 25, 2016

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.
Inference
Hypothesis
Science
Observation
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
Observation
Hypothesis
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:

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
physical property
volume
chemical property
vapor
extensive property
Chapter 3
mixture
intensive property
solution
element
compound
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
International System of Units (SI)
Significant figures
Accepted value
Experimental value
Percent error
Dimensional analysis

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 80
d) 9 mg =_______g 0.9 e) 2 mL = ______L 0.002 f) 6 kg = _____g 6000
g) 4 cm =_______m 0.4 h) 12 mg = ______ 0.0012 g i) 6.5 cm 3 = _______L 0.0065
j) 7.02 mL =_____cm 7.02 3 k) .03 hg = _______ 30 dg l) 6035 mm _____cm 603.5
m) .32 m = _______cm 32 n) 38.2 g = 0.0382 _____kg

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.
The mass of 6 cereal bars is 0.222 kilograms. This is less than 1 kilogram because 1 is not equal
to .222 kilograms. 1 kilograms is 0.78 kilograms greater than 0.222 kilograms.
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.
Wanda must make 100 trips because 1 kilogram is equal to 10 hectagrams. Since she can only carry 1
kilograms per trip, it would be 110 trips.
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.
Dr. O needs 250 more grams because 1 kilogram equals 1000 grams. 1000-750 equals 250 grams.
5. Weather satellites orbit Earth at an altitude of 1,400,000 meters. What is this altitude in kilometers?
The altitude in kilometers is 1,400.00
6. Which unit would you use to measure the capacity? Write milliliter or liter.
a) a bucket __________
liter
b) a thimble __________
milliliter
c) a water storage tank__________
liter
d) a carton of juice__________
milliliter
7. Circle the more reasonable measure:
a) length of an ant 5mm or 5cm
b) length of an automobile 5 m or 50 m
c) distance from NY to LA 450 km or 4,500
km
d) height of a dining table 75 mm or 75 cm
8. Will a tablecloth that is 155 cm long cover a table that is 1.6 m long? Explain your answer.
No. The table cloth will only be 1.55 m long, which is 0.05 less than 1.6m
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?
The line would be 31.2 meters ling because each bill is 15.6, and 15.6*200 would equal 31.2 meters
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?
Her husband will NOT hit his head on the hanging lamp because he is only 2 meters tall, with .09 of a
meter still left before he hits his head.

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
b. volume
_____ 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
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 _____________________________________________________
Both are units to measure temperature
________________________________________________________________________________
17. balance, second, mass __________________________________________________________
Both are units to measure weight
________________________________________________________________________________
18. kilogram, liter, cubic centimeter __________________________________________________
Both are units of volume
________________________________________________________________________________
19. time, second, distance __________________________________________________________
Both are units to measure time
________________________________________________________________________________
20. decimeter, kilometer, Kelvin _____________________________________________________
Both are units to measure distance
________________________________________________________________________________

1. How many meters are in one kilometer? __________
1000 meters
2. What part of a liter is one milliliter? __________
0.001 L
3. How many grams are in two dekagrams? __________
20 grams
4. If one gram of water has a volume of one milliliter, what would the mass of one liter of water be in
kilograms?_1_________
5. What part of a meter is a decimeter? __________
0.1 decimeters
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 ten
.
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
.

Standards of Measurement
Fill in the missing information in the table below.
Prefix
S.I prefixes and their meanings
Meaning
milli
0.001
centi
0.01
deci- 0.1
deka
10
hecto- 100
kilo
1000
Circle the larger unit in each pair of units.
1. millimeter, kilometer
2. decimeter, dekameter
4. centimeter, millimeter
5. hectogram, kilogram
3. hectogram, decigram
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
_____ 2 b. centimeter
_____ 4 c. meter
_____ 6 e. hectometer
_____ 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? __________________________________
0.001

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
N X 10 A
ANY NON-ZERO IS ALWAYS SIGNIFICANT
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

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.

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

Convert each number from Scientific Notation to real numbers:
1. 7.485 X 10³ 6. 1.683 X 10⁰
7,485 1.683
2. 8.842 X 10² 7. 3.622 10⁰
884.2 3.662
3. 2.887 X 10−⁵ 8. 1.735 X 10−⁵
0.00002887 0.00001735
4. 5.893 X 10−² 9. 9.736 X 10−¹
0.05893 0.9736
5. 6.162 X 10−³ 10. 8.558 X 10−²
0.006162 0.08558
Convert each number from a real number to Scientific Notation:
11. 0.0006633 16. 1,937,000
6.633 X 10 -4
12. 0.0004445 17. 34,570
4.445 X 10 -4 1.937 X 10 6
15. 727,400 20. 678.3
7.274 X 10 5 6.783 X 10 2
3.457 X 10 4
13. 0.002182 18. 0.00003948
2.182 X 10 -3
3.948 X 10 -5
14. 469.5 19. 894,500
4.695 X 10 2
8.945 X 10 5

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 in scientific notation, just look at the
numbers, not the 10 a because they are
not significant.
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

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.

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.

How Many Significant Digits for Each Number?
4
1) 2359 = ______
2) 2.445 x 10−⁵= ______
3) 2.93 x 10⁴= ______
4) 1.30 x 10−⁷= ______
4
5) 2604 = ______
3
3
6) 9160 = ______
7) 0.0800 = ______
2
2
8) 0.84 = ______
9) 0.0080 = ______
10) 0.00040 = ______
11) 0.0520 = ______
12) 0.060 = ______
13) 6.90 x 10−¹= ______
14) 7.200 x 10⁵= ______
15) 5.566 x 10−²= ______
16) 3.88 x 10⁸= ______
17) 3004 = ______
18) 0.021 = ______
19) 240 = ______
20) 500 = ______
3
2
3
2
4
2
2
1
3
4
3
4
4
3
2.445 x 10 -5 =0.00002445

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 = _______
84.2
2) 87.54 - 3.3 = _______
3) 99.1498 + 6.5397 + 9.7 = _______
4) 5.868 - 5.1 = _______
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 = _______
55.36
115.4
20) 6.5238 + 1.7 + 27.79 = _______

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...
200
1) 0.6 x 65.0 x 602 = __________
2) 720 ÷ 7.7 = __________
3) 929 x 0.3 = __________
4) 300 ÷ 44.31 = __________
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.

1qt=32 oz 1gal = 4qts 1.00 qt = 946 mL 1L = 1000mL
2. Convert 8135.6 mL to quarts
8135.6 mL 1 qts
946 mL
=
8135.6
946
8.6 qts
3. Convert 115.2 oz to mL
115.2 oz 1 qt
32 oz
946 mL
1 qt
=
108979.2
32 3405.6 mL
4. Convert 2.3 g to Liters
2.3 gal 4 qts
1 gal
946 mL
1.00 qt
1 L
= 8703.2
1000 mL
1000
8.7032 L
5. Convert 8.42 L to oz
8.42 L
1000 mL
1 L
1 qt 32 oz
946 mL
=
1 qt 946
269440
284.820296
Go to http://science.widener.edu/svb/tutorial/ chose #7 “Converting Volume” and do 5 more in the
space provided.
1. Convert _________ 166.4 oz to _________ L
166.4 oz 1 qt 946 mL
32 oz 1 qt
=
1 L 157414.4
1000 mL
32000
4.9192 L
2. Convert _________ 8324.8 mL to _________ qt
8324.8 mL 1.00 qts = 8324.8
946 mL
946
8.8 qts
3. Convert _________ 6.34 L to _________ qts
=
4. Convert _________ to _________
=
mL
5. Convert _________ to _________
=

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
28

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
29

The Learning Goal for this assignment is:
Notes Section
http://www.learner.org/interactives/periodic/basics_interactive.html
30

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.
31

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.
32

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.
33

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.
34

Summary
35

36

Electron Configuration
Color the sublevel:
s = Red
p = Blue
d = Green
f = Orange
S
S
P
D
F
Write in sublevels
Write period, sublevel and super scripts.
Ctrl Shift =
gives you super scripts
37

The Learning Goal for this assignment is:
Relate properties of atoms and their position in the periodic table
to the arrangement of their electrons.
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
38

Electron Configuration
In order to write the electron configuration for an atom you must know the 3 rules of
electron configurations.
1. Aufbau
Notation
nO e
where
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
39

In the space below, write the unabbreviated electron configurations of the following elements:
1) sodium ________________________________________________
2) iron ________________________________________________
3) bromine ________________________________________________
4) barium ________________________________________________
5) neptunium ________________________________________________
In the space below, write the abbreviated electron configurations of the following elements:
6) cobalt ________________________________________________
7) silver ________________________________________________
8) tellurium ________________________________________________
9) radium ________________________________________________
10) lawrencium ________________________________________________
Determine what elements are denoted by the following electron configurations:
11) 1s²s²2p⁶3s²3p⁴ ____________________
12) 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s¹ ____________________
13) [Kr] 5s²4d¹⁰5p³ ____________________
14) [Xe] 6s²4f¹⁴5d⁶ ____________________
15) [Rn] 7s²5f¹¹ ____________________
Identify the element or determine that it is not a valid electron configuration:
16) 1s²2s²2p⁶3s²3p⁶4s²4d¹⁰4p⁵ ____________________
17) 1s²2s²2p⁶3s³3d⁵ ____________________
18) [Ra] 7s²5f⁸ ____________________
19) [Kr] 5s²4d¹⁰5p⁵ ____________________
20) [Xe] ____________________
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 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 rubidium
[Kr] 5s 2 4d 10 5p 3 antimony [Xe] 6s 2 4f 14 5d 6 osmium
[Rn] 7s 2 5f 11 einsteinium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 4d 10 4p 5 not valid (take a look at “4d”)
1s 2 2s 2 2p 6 3s 3 3d 5 not valid (3p comes after 3s) [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)
40

41

Create groups for these Scientist and explain your groupings
(use the information you got from your research)
Discovered Radioactivity in Elements:
Antoine Henri Becquerel
Marie and Pierre Curie
Knowledge of the atom and its parts (electrons, neutrons, protons, mass):
Neils Bohr
Louis de Broglie
Hantaro Nagaoka
J.J. Thomson
Erwin Shrodinger
Robert Millikan
Eugene Goldstein
James Chadwick
Ernest Rutherford
Formation of the Periodic Table:
Glenn Seaborg
Dmitri Mendeleev
Lothar Meyer
J.W. Dobereiner
Contributed to the atomic theory:
Democritus (Founder)
John Dalton
42

Research the Scientist and summarize their contributions to the Atomic Theory
Antoine Henri Becquerel
Niels Bohr He discovered that the electron is in a fixed position, and has shells of electrons on
the outer part of the electron. He also contributed that the energy is transferred only in
restricted amounts.
Louis de BroglieThe French physicist predicted the electron's wave nature. He also created the waveparticle
duality theorem, which states that any atom acts like a wave and a particle.
Glenn Seaborg Glenn Seaborg discovered many elements that were described as "transuranium"
elements. He discovered the transitional metals, and added to the periodic table
of elements..
Hantaro Nagaoka Hantaro Nagaoka proposed a model of the atom in which the nucleus was positively
charged, and surrounded by electrons. His prediction of the electrons in rings around
the nucleus was proven wrong by Ernest Rutherford.
Democritus Democritus was the student of the original founder of the atomic theory, and he stated that
atoms were small quantities of matter. He altered the original theory to prove it more
systematic.
Marie and Pierre Curie Marie and Pierre Curie added more to Becquerel's contribution to the atomic theory,
in terms of the presence of radiation in the elements, such as uranium. Marie Curie
also discovered two new radioactive elements (polonium and radium).
Eugene Goldstein Eugene Goldstein was behind the discovery of protons. He also discovered that the
electrons repel from the protons because they travel in opposite directions.
Dmitri Mendeleev Dmitri Mendeleev was the first person to publish the first Period Table of Elements. He
also created the Periodic Law, which stated that similar elements would appear in
certain spaces or intervals.
J.J. Thomson J.J. Thomson contributed to the atomic theory his discovery of the electron. He was able to
discover that electrons are negatively charged, and discovered that protons are positively
charged.
James ChadwickJames Chadwick discovered the neutron. He was able to do so by attacking
beryllium with alpha particles, and watched how they were repelled by the nucleus
because of the neutral electrical charge.
Erwin Shrodinger
Erwin Shrodinger modified Niels Bohr's model of the atom. He used mathematical
equations to describe how likely it is that an electron would be in a certain position.
John Dalton
theory.
Lothar Meyer
Ernest Rutherford
The French scientist discovered the concept of radioactivity in elements. He
was able to contribute this by using uranium and a photographic plate.
John Dalton showed that common elements disintegrated into common elements. He also
discovered that all things are made up of atoms, which is a fundamental concept of the atomic
Lothar Meyer pioneered the development of the periodic table of elements. He worked with
Rober Bunsen and Dmitri Mendeleev, and was one of the contributers to the development
of the periods in the table.
Robert Millikan Robert Millikan was the scientist who was behind the famous "Oil-Drop" experiment in
which he was able to conclude what charge electrons have. He was able to determine
that the electron has a negative charge.
J.W. Dobereiner
J.W. Dobereiner discovered families of elements with similar properties. There always
seemed to be three elements in these families, which led him to call them triads.
Ernest Rutherford decided to shoot positively-charged atoms at gold foil, and
discovered that the particles passed through. Most of the particle has empty
space, while the rest of the mass was in the heavy nucleus.
43

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
Alkali Earth Metals
Transitional Metals
Inter Transitional Metals
Metals
Metalloids
Non Metals
Noble Gases
44

Using Wikipedia, define the 8 categories of elements on the
left page.
Color your periodic table similar to the one on
pages 168—169 of your book.
alkali metals
alkaline metals
other metals
transitional metals
lanthanoids
metalloids
non metals
halogens
noble gases
unknown elements
actinoids
45

Define Atomic Size:
Atomic Size
Explanation:
46

Ionization Energy
Define Ionization Energy:
Explanation:
47

Define Electronegativity:
Electronegativity
Explanation:
48

Ion Size
Define Ion Size:
Explanation:
49

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

Name ____________________
Arnaldo Salazar Gomez
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 __________, frequency Long wavelength
High frequency, Short ______________
wavelength
Radio Waves
Microwaves Infra-red Visible Light Ultra-violet X-rays
Gamma rays
2. Click on “Radio waves”. They are used for _______________________
mainly communications
3. Click on “Microwaves”. They are used for cooking, mobile _________, wifi _______ 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 DVD ___ 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_______________ Radiotherapy .
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
X-Rays
Ultra Violet
Visible Light
Infra Red
Micro Waves
Radio Waves
Given off by stars, and by some
radioactive substances
These are made by stars and some X-ray
machines.
It's made by special lamps, and given off by
the Sun in large quantities. UV attracts
insects, and are electrocuted by UV.
They're made by anything that is hot enough
to glow. White light is made of all the colors
mixed together
These are given off by hot objects.
These are also able to be given off
by stars, lamps, flames, and anything
else warm.
These are made by various types of
transmitters. In a phone, they're made by a
transmitter chip and an antenna, and are
also given off by stars, only measuring a
couple of centimeters
They're made by various types of
transmitters, depending on the
wavelength. They're given off by
stars, sparks, and lightning.
These can be used to kill cancer cells
through radiotherapy, targeted radiotherapy,
tracers, to sterilize food to keep it fresh
longer, and to sterilize medical equipment.
These are used to see inside of people, are
used at airport security, and astronomers
use them because many objects in the
universe give off X-rays. These can also be
used to scan the brain.
It can be used for sun tans, detecting
forged bank notes, making clothes glow in
clubs, kill microbes, sterilize food and drug
products, and produce a suitable amount of
vitamin D.
It allows people to see things, used by
compact discs and DVD players, laser
printers, and aircraft weapon aiming
systems.
Remote controls in TVs use them, video
recorders, physiotherapists use them to
heal sports injuries. These are also used for
night sight, alarm systems, police
helicopters, weather forecasts.
It can be used to make substances hot,
and is used to cook many types of food.
Mobile phones generate these with a small
antenna. These are also used by fixed
traffic speed cameras, and radar.
Radio waves are used mainly for
communications, and have varying
wavelengths. These can be used for FM,
Police, Military Aircraft radios, and TV
transmissions.
These rays can cause cell
damage, cause mutations in
growing tissues, making
unborn babies very vulnerable
to them.
These can cause cell
damage and cancers, which
is why doctors stand behind
a shield to protect
themselves.
These rays can damage the
retina, cause sunburns, and
even skin cancer.
Too much light can damage the
retina, which can also happen when
you look at something very bright.
The damage can heal, and if it's too
bad it might be permanent.
The danger is overheating.
Prolonged exposure to these is
known to cause cataracts. Modern
military planes can "cook" people
because of powerful radar units. It
can also cause brain damage
Some dangers of radio waves
are the large doses are
believed to cause cancer,
leukemia, and other disorders.
_____ Low Frequency High _____ frequency,
Short wavelength ______ Wavelength
Short

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:
Protons (+) have a mass of 1 Atomic Mass Unit
Neutrons (/) have a mass of 1 Atomic Mass Unit. Neutrons have both positive and negative charge in
order to be able to have a neutral charge (get it? because positive+negative=neutral).
Electrons (-) have no mass (well, it actually does, but it's so small.
Protons and neutrons have the same mass, but different charges.
An ISOTOPE is an atom with an abnormal number of protons (whether too many, or too few). When talking
about isotopes, it is always referred to as the element-mass (i.e. Carbon-14). When there is a greater
difference of protons to neutrons, it could technically cause radiation. The end goal for all atoms is to stabilize,
but isotopes are not stable. EVERYTHING THAT IS LIVING has a specific amount of Carbon-14 in it. This
only applies to living things, not non-living things. Carbon dating is essentially taking a sample of Carbon-14
and figuring out how much the organism has/had when they were first born. It is then necessary to be able to
determine how many half-lifes the Carbon-14 has gone through. Once the amount of Carbon-14 is given,
then the half-life lengths will indicate how many is left to determine the true age of the organism. Alpha
particles are basically a helium atom (4/2He: Mass of 4 AMU and 2 Protons). A Helium is given off during
Alpha radiation.
EXAMPLE:
Uranium-238. To go through Alpha (a) radiation, subtract 4 from mass and 2 from protons to produce
Thorium-234 and Helium. Every Uranium-238 is ALWAYS going through an Alpha radiation. Radium-226 also
goes through an Alpha particle radiation, Radon-226,
Radon-224. There is truly no way to find out what every atom goes through.
When an atom goes through an Alpha radiation, the atom ALWAYS goes down 2 elements. Example: When
polonium undergoes Alpha radiation, it goes down to lead.
Beta particles: There are negative AND positive beta particles.
Example: Carbon-14 has 8 protons, and 6 neutrons. To make this even, a negative charge would have to be
given off, but that charge would have to give it off from the nucleus. It is necessary to convert a neutron to a
proton, but what would be left would be a negative charge.
Every carbon-14 that undergoes a negative beta radiation, it would become Nitrogen. When going through a
negative Beta radiation, it always goes up by one. Example: Carbon-14 becomes Nitrogen. Alpha particles
have very minimal damage, but Beta particles have an increased amount of energy, and could go through
paper, but not plastic. Positive Beta particles are called (Positrons). In positive beta radiatons, the atom
changes a proton to a neutron, which gives off a positron. These ALWAYS go down by one. Example:
Boron-8 becomes Berylium. Gamma radiation gives off nothing but energy. This HAS TO be associated with
Alpha and Beta; it never goes individually as a form of nuclear radiation. This is ONLY given off as a
byproduct. An Alpha particle is a Helium particle. This could be written as 4/2 He, where 4 is the atomic mass,
and 2 is the number of protons. The mass is always the larger number, and the bottom number is always
smaller because there are less protons than the atomic mass. In negative beta radiation, a neutron is
converted into a proton, and negative energy is released in the form of a neutron. In positive beta radiation, a
proton is converted into a neutron, and positive energy is released in the form of a positron (positive electron)
Reference:
Alpha Particles radiation ALWAYS decrease by 4 in mass, and 2 in protons. (Down by 2) (i.e. Sodium becomes Flourine)
Negative Beta radiation ALWAYS increase by 1 in protons. (Up by 1) (i.e. Hydrogen becomes Helium)
Positive Beta radiation ALWAYS decreases by 1 in protons (Down by 1) (i.e. Helium becomes Hydrogen).
Gamma radiation is a byproduct of any of the three radiation listed above.
Half-Life is the time it takes for a substance to lose half of the original mass.

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:
NH 4+ =Ammonium because it has 1 Nitrogen, and 4 Hydrogens. The purpose of all bonds is to
become stabilized. Ammonium has 1 less electron.
NH 3 =Ammonia because it is a compound.
NaCl is the common formula for Table Salt. This forms an ionic bond because it is a bond between a
metal and a nonmetal (Sodium is the metal, Chlorine is the nonmetal.
Using the Lewis Dot Diagram, a visual representation of any element can be formed. The cation
(sodium) is always given its elemental name, but the anion (chloride) is always given the suffix -ide.
An ionic compound consists of charged atoms (of both positive and negative charges). These
compounds are generally solids, have high melting points and conduct electrical current. These
compounds are generally formed from metal and a non-metal bonds. Ions are completely different in
physical and chemical properties from the neutral atoms of the elements. In other words, the atoms
look vastly different than the actual ions for that given element. A key difference between stable
elements and their ionic counterparts is that elements are stable (meaning hey have a balance
between the number of electrons:protons), whereas the ions have a charge, whether that is positive
or negative (an ion is positive if there are more protons than electrons (also called a cation), and is
called negative if here are more electrons than protons (also called an anion)). For any element that
has less than 4 valence electrons, these are bound to give up their electrons to meet the octet rule.
EXAMPLE: Sodium has eleven (11) electrons, with only one (1) electron on its outer shell (commonly
referred to as a valence electron). For sodium to meet the octet rule, it can either give up one electron
or gain seven (although this is highly unlikely). When an electron is given up, sodium reaches its octet
rule and has the same Lewis Dot Structure as Neon, the element with ten (10) electrons. For sodium
to meet the octet rule, it had to achieve a full shell of eight (8) electrons, which can be realistically
done by giving up an electron, giving sodium a positive one (1) charge because there is one more
proton than electrons. This sodium that has met the octet rule can be written as Na^+.
For any element that has more than four valence electrons, these are bound to gain electrons to meet
the octet rule.
EXAMPLE: Fluorine has nine (9) electrons, with seven (7) electrons on its outer shell. For fluorine to
meet the octet rule, it can either gain one electron or give up seven (although this is highly unlikely).
When an electron is gained, fluorine reaches its octet rule and has the same Lewis Dot Structure as
Neon, the element with ten (10) electrons. For fluorine to meet the octet rule, it had to achieve a full
shell of eight (8) electrons, which can be realistically done by gaining an electron, giving sodium a
negative one (1) charge because there is one more electron than protons. This fluorine that has met
the octet rule can be written as F^-. This new atom is also called Fluoride because it is a simple
negative ion.
Ionic compounds are a result of an attraction between a positive and a negative ion. In crystalline
compound, this net balance of forces is called the LATTICE ENERGY. The energy required to remove
an electron from a neutral atom is called the IONIZATION POTENTIAL.

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 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 electrons.
2. Find the central atom and bond the other atoms to it. Subtract the number of electrons in the
bonds from the total. Add lone pairs to the terminal atoms. Add Lone Pairs to the central atom or
double or triple bonds.
3. Find the Formal Charges. Try to get the charges as close to the zero as possible by moving
electrons and bonds.
Equation:
A=Central Atom, X=Things attached to the central atom, E=Lone Pairs
Covalent bonds typically occur between two non-metals, and there is no loss or gain of
electrons among the atoms. Non-polar binding is when there is an even share of electrons,
while polar bonding is when there is an uneven number of electrons being shared. Hydrogen is
the simplest non-polar molecule, as it only needs two (2) electrons to meet its octet rule since
the nearest noble gas is helium, which only has two (2) electrons in total. When two hydrogens
bonds, they share electrons equally. Oxygen needs two more electrons to complete the octet
rule, and when they bond with one another, they need a double bond to meet their octet rule.
These equally share the four electrons that are present in the double bonds. Hydrogen chloride
is a polar covalent molecule because chlorine gets an unequal share of the two electrons, as
the electrons spend more time with the chlorine, rather than the hydrogen. Water too is a
covalent compound.

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
Beryllium dichloride
Cl
Be
Cl
element bond lone pair
C

Trigonal Planar
Molecular Geometry
Orbital Equation Lone Pairs Angle
sp 2 AX 3 None 120
BF 3
Boron triflouride
F
B
F
F
element bond lone pair
C

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

Tetrahedral
Molecular Geometry
Orbital Equation Lone Pairs Angle
sp 3 AX 4 None 109.5
Phosphate
PO 4
3-
O
3-
O
P
O
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 3 AX 2 E 2 2 104.5
H 2 O
Dihydrogen oxide
H
O
H
element bond lone pair
C

Trigonal Bi Pyramidal
Molecular Geometry
Orbital Equation Lone Pairs Angle
sp 3 d AX 5 None 120/90
PCl 5
Cl
Cl
P
Cl
Cl Cl
Phosphorus pentachloride
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
F Cl
F
F
Chlorine triflouride
element bond lone
C

Octahedral
Molecular Geometry
Orbital Equation Lone Pairs Angle
sp 3 d 2 AX 6 None 90
SF 6
F
F
F
ulfur hexaflouride
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
-
Ioion
dine tetrachloride
Cl
Cl I
Cl
Cl
element bond lone pair
C

Orbitals Equation Lone Pairs
Angle
sp AX2 None 180
sp 2 AX3 None 120
sp 2 AX2E 1 116
sp 3 AX4 None 109.5
Linear
Name
Trigonal Planar
Bent
Tetrahedral
sp 3 AX3E 1 107
Trigonal Pyramidal
sp 3 AX2E2 2 104.5
Bent
sp 3 d AX5 None 120/90
Trigonal Bipyramidal
sp 3 d AX3E2 2 90
T-Shaped
sp 3 d 2 AX6 None 90
Octahedral
sp 3 d 2 AX4E2 2 90
Square Planar

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
78

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
79

80
http://www.bbc.co.uk/education/guides/zm9hvcw/revision
Learning Goal: 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.
Notes: Homologous series contain hydrocarbons that have
similar chemical properties and the same general formula.
Alkanes are the first homologous. These all end in -ane.
Examples of this are methane (which is a natural gas used
for cooking and heating), propane (which is used in gas
cylinders for BBQ, etc.), and octane (which is used in petrol
for cars). The general formula for alkanes is CnH2n+2. The
"n" in the formula represents the number of carbons in said
molecule. Example: Methane has the formula CH4. When
naming branched chain alkanes, there are several rules in
place. The longest unbranched chain that contains the
functional group is the parent molecule, or simply the
longest unbranched chain for alkanes. However, these
alkanes are bound to bending. The branches must also be
names, as well as the number of them. The alkanes don't
have a functional group and the branches are numbered
from the end that gives the lowest set of numbers. Alkenes
have the equation CnH2n. The position of the carbon to
carbon double bond must be identified. Alkanes are
SATURATED!!!!. Any that have double bonds are
unsaturated. Cycloalkanes have the same properties as the
alkenes.

81