Semester 1 Notebook-Martin


Evys Martin’s Chemistry


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


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


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.


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


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


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


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




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


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 __________________________

Evys Martin

PERIOD ________________________


PARENT NAME ____________________________

Ismary Reyes

PARENT NUMBER _________________________


SIGNATURE ____________________________

DATE ____________________________________


Chapter 1

Unit 1

Introduction to Chemistry

The students will learn why and how to solve problems using


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.





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


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.





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





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


Peer Review



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


Students will be able to identify indicators of physical and chemical changes.

Students will be able to calculate density.


physical property


chemical property


extensive property

Chapter 3


intensive property




Scientific Measurements

The students will be able to solve conversion problems using


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


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

Determine appropriate and consistant standards of measurements for the data to be collected in a survey

or equipment.

K-ing H-enry D-ied B-ecause D-rank C-hocolate M-ilk








one centemiter squared equals i milliliter.

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 .5

b) 4 mL = ______L .004 c) 8 g = _______mg 8000

d) 9 mg =_______g 0.009

e) 2 mL = ______L .oo2 f) 6 kg = _____g 0006

g) 4 cm =_______m 0.04 h) 12 mg = ______ 0.012g i) 6.5 cm 3 = _______L 0006.5

j) 7.02 mL =_____cm 702.0

3 k) .03 hg = _______ 0.3 dg l) 6035 mm _____cm 603.5

m) .32 m = _______cm .032

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 the bars will stay the same but the weight will change

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.

she will need to go 110 times, one kilogram aquals 1 hg

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.

she needs 250 grams

5. Weather satellites orbit Earth at an altitude of 1,400,000 meters. What is this altitude in kilometers?

1400 kilometers.

6. Which unit would you use to measure the capacity? Write milliliter or liter.

a) a bucket __________ liter

b) a thimble __________


c) a water storage tank__________ liter

d) a carton of juice__________ liter

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.


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?


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?

he wont hit the lamp, 1m=100cm so 41cm is 9cm away from

being half a meter.

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

_____ e 1. distance between two points

a. time

_____ 2. SI unit of length

m_____ 3. tool used to measure length

_____ g 4. units obtained by combining other units

_____ b 5. amount of space occupied by an object


_____ 6. unit used to express volume

_____ j 7. SI unit of mass

_____ d 8. amount of matter in an object

_____ j 9. mass per unit of volume

_____ o 10. temperature scale of most laboratory thermometers

_____ 11. instrument used to measure mass

_____ a 12. interval between two events

_____ j 13. SI unit of temperature

i_____ 14. SI unit of time

_____ n 15. instrument used to measure temperature

b. volume

c. mass

d. density

e. meter

f. kilogram

g. derived

h. liter

i. second

j. Kelvin

k. length

1. balance

m. meterstick

n. thermometer

o. Celsius

Circle the two terms in each group that are related. Explain how the terms are related.

16. Celsius degree, mass, Kelvin _____________________________________________________

these all measure the temp and movement of the


atoms in something. the mass will impact the temp.

17. balance, second, mass __________________________________________________________


18. kilogram, liter, cubic centimeter __________________________________________________

these can be used to measure the interior of



19. time, second, distance __________________________________________________________

these are all related, the distance to travel on


such time equals speed.

20. decimeter, kilometer, Kelvin _____________________________________________________

kalvin dont belong, the others are



1. How many meters are in one kilometer? __________ 100

2. What part of a liter is one milliliter? __________ a 4th

3. How many grams are in two dekagrams? __________


4. If one gram of water has a volume of one milliliter, what would the mass of one liter of water be in

kilograms?__________ .001

5. What part of a meter is a decimeter? __________


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 _______________ metric .

8. SI is based on units of _______________ ten


9. The system of measurement that was based on units of ten was the tenth _______________ 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 _______________ ten


Standards of Measurement

Fill in the missing information in the table below.



S.I prefixes and their meanings





deci- 0.1



hecto- 100



Circle the larger unit in each pair of units.

1. millimeter, kilometer 4. centimeter, millimeter

2. decimeter, dekameter 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

_____ 3 b. centimeter

_____ 5 c. meter


_____ d. dekameter


_____ e. hectometer


_____ f. millimeter



g. decimeter

Use your knowledge of the prefixes used in SI to answer the following questions in the spaces


7. One part of the Olympic games involves an activity called the decathlon. How many events do you

think make up the decathlon?_____________________________________________________


10 years

8. How many years make up a decade? _______________________________________________

100 years

9. How many years make up a century? ______________________________________________


10. What part of a second do you think a millisecond is? __________________________________

The Learning Goal for this assignment is:

determine appropriate and consistent standards of measurement of the

data to be collected in a survey.

Notes Section

to make the exponent smaller change the decimal

on the first number.

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


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


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.


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.


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.


(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.


(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.


(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

7485. 1.683

2. 8.842 X 10² 7. 3.622 10⁰

884.2 3.622

3. 2.887 X 10−⁵ 8. 1.735 X 10−⁵

.00002887 .00001735

4. 5.893 X 10−² 9. 9.736 X 10−¹



5. 6.162 X 10−³ 10. 8.558 X 10−²



Convert each number from a real number to Scientific Notation:


11. 0.0006633 16. 1,937,000

6.633 -4

7.274x10 2 6.783x10 2

12. 0.0004445 17. 34,570

4.445x10 -4


13. 0.002182 18. 0.00003948

2.182x10 -3

3.948 -5

14. 469.5 19. 894,500

4.695x10 2

8.945x10 2

15. 727,400 20. 678.3

The Learning Goal for this assignment is:

determine appropriate and consistant standards of measurement

for the data to be collected in survey or experiment.

Notes Section:

non-zeros are always significant bro.

any zero between significant digits, than those are significant.

after a decimal, if its before significant its not valuable, if its after then it is.

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


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


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


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.



Rule 3: A final zero or trailing zeros in the decimal portion ONLY are


This rule causes the most confusion among students.

In the following example the zeros are significant digits and highlighted in blue.



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.



These zeros serve only as space holders. They are there to put the decimal point in its correct


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.



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


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?

1) 2359 = ______ 4

2) 2.445 x 10−⁵= ______ 4

3) 2.93 x 10⁴= ______ 3

4) 1.30 x 10−⁷= ______ 2


5) 2604 = ______

6) 9160 = ______ 4

7) 0.0800 = ______ 2

8) 0.84 = ______ 2

9) 0.0080 = ______ 2

10) 0.00040 = ______ 2

11) 0.0520 = ______


12) 0.060 = ______ 2

13) 6.90 x 10−¹= ______ 2


14) 7.200 x 10⁵= ______


15) 5.566 x 10−²= ______

16) 3.88 x 10⁸= ______


17) 3004 = ______ 4

18) 0.021 = ______


19) 240 = ______ 3


20) 500 = ______

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


Solve the Problems and Round Accordingly...

1) 43.287 + 5.79 + 6.284 = ____ 55.36


2) 87.54 - 3.3 = _______ 84.2



3) 99.1498 + 6.5397 + 9.7 = _______


4) 5.868 - 5.1 = _______ 0.7


5) 59.9233 + 86.21 + 99.396 = _______ 245.52


6) 7.7 + 26.756 = _______ 34.4


7) 66.8 + 2.3 + 4.8516 = _______ 73.9


8) 9.7419 + 43.545 = _______

9) 4.8976 + 48.4644 + 1.514 = _______

10) 4.335 + 35.85 = _______

11) 9.448 - 1.7 = _______

12) 75.826 - 8.6555 = _______

13) 57.2 + 23.814 = _______

14) 77.684 - 4.394 = _______

15) 26.4496 + 3.339 = _______

16) 9.6848 + 29.85 = _______

17) 63.11 + 2.5412 + 4.93 = _______

18) 11.2471 + 75.4 = _______

19) 73.745 - 8.755 = _______

20) 6.5238 + 1.7 + 27.79 = _______

The following rule applies for multiplication and division:

The LEAST number of significant figures in any number of the problem determines the number of

significant figures in the answer.

This means you MUST know how to recognize significant figures in order to use this rule.

Solve the Problems and Round Accordingly...

1) 0.6 x 65.0 x 602 = __________

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 = __________

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.





electromagnetic radiation


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


Students will participate in activities to calculate half-life.


nuclear radiation

alpha particle

beta particle

gamma ray


½ life




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.





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


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


Students will be able to predict shapes of various compounds.


empirical formula





Evys Martin

Name ____________________

Go to the web site

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 __________, frequencies Long wavelength

High frequency, Short ______________


Radio Waves

microwaves infrared visisble light ultra violet x-rays

Gamma rays

2. Click on “Radio waves”. They are used for _______________________


3. Click on “Microwaves”. They are used for cooking, mobile _________, phones _______


cameras and _________. radars

4. Click on “Infra-red”. These waves are given off by _____ hot _________. objects They are used for remote controls,

cameras in police ____________ cars , and alarm systems.

5. Click on “Visible Light”. This is used in ___ dvdplayers and _______ color 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 ________



which is called R_______________ adiotherapy .

9. My Quiz score is ____%. 71

10. Name ________________________________

Go to the web site

Name How they’re made Uses Dangers

Gamma rays


ultra violet

visible light

infra red

micro waves

radio waves

given off by radioactive

substances and stars

stars, is fired a beem

of electrons forms


given off by sun , found

in sun beds

Given off my anything

that glows., given off by


hot objects, given off by stars,

lamps, and flames.

cellphones, and microwaves

Stars, Sparks, and even



see inside things,

including humans

tanning, fly killer,

bank notes

light bulbs, printers,

dvd players

remote controls,

night sights, healing



cellphones, wi-fi



can harm the


can cause cell

damage and

even cancer

damage to

retina, sunburn,

skin cancer

damege retina




large doses

can cause


_____ Frequency _____ frequency,

Short wavelength ______ Wavelength

Learning Goal for this section:

Explain and copare nuclear reactions, radioactive decay, fission, and the

energy, changes associated with them and their associated safety issues.

Notes Section:





carbon 14 could be radioactive because it

had a radioactive isotope.

carbon makes up everything, things have

specific amounts of carbon 14.

Beta particles have the mass of an electron, theres positive and negative beta

particles. Neutrons are both positive and negative, you take away the neutron

it becomes positive

Positive beta-Positron, has

practically no charge.

Can take away positive

charge from neutron and

make a negative, vise


If you do that is goes down by one EX: poloniom

would become bismouth.

Gamma(small amount)- ENRGY. -Not a huge amount but can be


Half life-loses protons, loses power, becomes weaker.

Iodine131 20mg has radioactivity- half life 8 days

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


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


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




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


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


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


* 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:

Non metal with metal bond is ionic, non metal with non metal is covalent bond.

gives off an electroneasier

to loose electrons than to gain.

Ionic compound diagram




could also add more of the elements, Be P Be P Be and this would be eaven.

The Learning Goal for this assignment is:

Distinguish between bonding forces holding compounds togather 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.


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


Instead the two hydrogen atoms SHARE both electrons equally.


Molecules of oxygen, present in about 20% concentration in air are

also covalent molecules. See the graphic on the left of the Lewis Dot


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, 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:

when it all equals zero than its good and you gucci.... add the elctrons to the terminal


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


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.


Molecular Geometry

Orbital Equation Lone Pairs Angle

sp AX 2 0 180

BeCl 2




Beryllium dichloride

element bond lone pair


Trigonal Planar

Molecular Geometry

Orbital Equation Lone Pairs Angle

sp 2 Ax 3 None 120

BF 3


Boron trifluoride




element bond lone pair



Molecular Geometry

Orbital Equation Lone Pairs Angle

SP 2 Ax2E 1 116

O 3





element bond lone pair



Molecular Geometry

Orbital Equation Lone Pairs Angle

SP 3 AX 4 0 109.5


PO 4







element bond lone pair


Trigonal Pyramidal

Molecular Geometry

Orbital Equation Lone Pairs Angle

Sp 3 AX 3 E 1 107

PH 3





Phosphorus trihybrid

element bond lone pair



Molecular Geometry

Orbital Equation Lone Pairs Angle

SP 3 AX 2 E 2 2 104.5

H 2 O




Dihydrogen oxide

element bond lone pair


Trigonal Bipyramid

Molecular Geometry

Orbital Equation Lone Pairs Angle

SP 3 d AX 5 none 120/90

PCl 5




Phosphorus pentacloride




element bond lone pair



Molecular Geometry

Orbital Equation Lone Pairs Angle

sp 3 d AX 3 E 2 2 90

ClF 3

Chlorine trifloride





element bond lone pair



Molecular Geometry

Orbital Equation Lone Pairs Angle

sp 3 d 2 AX 6 none 90

Sulfur hexafluoride

SF 6








element bond lone pair


Square Planar

Molecular Geometry

Orbital Equation Lone Pairs Angle

Sp 3 d 2 Ax 4 E 2 2 90

Iodine tetrachloride

ICl 4







element bond lon


Orbitals Equation Lone Pairs Angle


sp AX 2 0 180 Beryllium dichloride

sp2 AX3 None 120 Boron trifluoride

SP2 Ax2E 1 116 Trioxide


Ax 4 0 109.5 Phosphate

Sp3 Ax3E 1 107 Phosphorus trihybrid

SP3 Ax2E 2 2 104.5 Dihydrogen oxide

SP3d Ax5 none 120/90 Phosphorus pentacloride

sp3d Ax3E2 2 90 Phosphorus pentacloride

sp3d2 Ax6 none 90 Sulfur hexafluoride

Sp3d2 Ax4E2 2 90 Iodine tetrachloride

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-

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