1st_Semester_Notebook

salvador77710

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

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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

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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

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Lorenzo Walker Technical High School

MUSTANG LABORATORIES

Chemistry Safety

Safety in the MUSTANG LABORATORIES - Chemistry Laboratory

Working in the chemistry laboratory is an interesting and rewarding experience. During your labs, you will be actively

involved from beginning to end—from setting some change in motion to drawing some conclusion. In the laboratory, you

will be working with equipment and materials that can cause injury if they are not handled properly.

However, the laboratory is a safe place to work if you are careful. Accidents do not just happen; they are caused—by

carelessness, haste, and disregard of safety rules and practices. Safety rules to be followed in the laboratory are listed

below. Before beginning any lab work, read these rules, learn them, and follow them carefully.

General

1. Be prepared to work when you arrive at the lab. Familiarize yourself with the lab procedures before beginning the lab.

2. Perform only those lab activities assigned by your teacher. Never do anything in the laboratory that is not called for in

the laboratory procedure or by your teacher. Never work alone in the lab. Do not engage in any horseplay.

3. Work areas should be kept clean and tidy at all times. Only lab manuals and notebooks should be brought to the work

area. Other books, purses, brief cases, etc. should be left at your desk or placed in a designated storage area.

4. Clothing should be appropriate for working in the lab. Jackets, ties, and other loose garments should be removed. Open

shoes should not be worn.

5. Long hair should be tied back or covered, especially in the vicinity of open flame.

6. Jewelry that might present a safety hazard, such as dangling necklaces, chains, medallions, or bracelets should not be

worn in the lab.

7. Follow all instructions, both written and oral, carefully.

8. Safety goggles and lab aprons should be worn at all times.

9. Set up apparatus as described in the lab manual or by your teacher. Never use makeshift arrangements.

10. Always use the prescribed instrument (tongs, test tube holder, forceps, etc.) for handling apparatus or equipment.

11. Keep all combustible materials away from open flames.

12. Never touch any substance in the lab unless specifically instructed to do so by your teacher.

13. Never put your face near the mouth of a container that is holding chemicals.

14. Never smell any chemicals unless instructed to do so by your teacher. When testing for odors, use a wafting motion to

direct the odors to your nose.

15. Any activity involving poisonous vapors should be conducted in the fume hood.

16. Dispose of waste materials as instructed by your teacher.

17. Clean up all spills immediately.

18. Clean and wipe dry all work surfaces at the end of class. Wash your hands thoroughly.

19. Know the location of emergency equipment (first aid kit, fire extinguisher, fire shower, fire blanket, etc.) and how to use them.

20. Report all accidents to the teacher immediately.

Handling Chemicals

21. Read and double check labels on reagent bottles before removing any reagent. Take only as much reagent as you

need.

22. Do not return unused reagent to stock bottles.

23. When transferring chemical reagents from one container to another, hold the containers out away from your body.

24. When mixing an acid and water, always add the acid to the water.

25. Avoid touching chemicals with your hands. If chemicals do come in contact with your hands, wash them immediately.

26. Notify your teacher if you have any medical problems that might relate to lab work, such as allergies or asthma.

27. If you will be working with chemicals in the lab, avoid wearing contact lenses. Change to glasses, if possible, or notify

the teacher.

Handling Glassware

28. Glass tubing, especially long pieces, should be carried in a vertical position to minimize the likelihood of breakage and

to avoid stabbing anyone.

29. Never handle broken glass with your bare hands. Use a brush and dustpan to clean up broken glass. Dispose of the

glass as directed by your teacher.

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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.

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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.

Salvador Gaspar

NAME __________________________

2

PERIOD ________________________

Felicia Rodas

PARENT NAME ____________________________

PARENT NUMBER _________________________

239-687-0029

SIGNATURE ____________________________

8-25-16

DATE ____________________________________

6


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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).


Science

Observation

Students will identify a phenomenon as science or not science.

Inference

Hypothesis

Identify which questions can be answered through science and which

questions are outside the boundaries of scientific investigation, such as

questions addressed by other ways of knowing, such as art, philosophy, and

religion.



Students will differentiate between problems and/or phenomenon that can and

those that cannot be explained or answered by science.

Students will differentiate between problems and/or phenomenon that can and

those that cannot be explained or answered by science.

Theory

Controlled experiment

Observation

Inference

Hypothesis

Describe how scientific inferences are drawn from scientific observations

and provide examples from the content being studied.





inferences.

Inference

Observation

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

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:

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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.

mixture

intensive property

solution

element

compound



mass

physical property

volume

chemical property

vapor

extensive property

Chapter 3

Scientific Measurements

The students will be able to solve conversion problems using

measurements.

Determine appropriate and consistent standards of measurement for the

data to be collected in a survey or experiment.


Students will participate in activities to collect data using standardized

measurement.

Students will be able to manipulate/convert data collected and apply the data


to scientific situations.

Scientific notation

International System of Units (SI)

Significant figures

Accepted value

Experimental value

Percent error

Dimensional analysis

9


cm x cm x cm = cm3

l*w*h

V

kelvin is plus or minus 273 to get to Celsius

ex=10C=283K

kelvin = 1Celcius

6.5*6.5*6.5=274.625

7.02*7.02*7.02=

(King)Kilo- (Henry)Hecto- (Died)Deka- (By)Base unit (Drinking)Deci- (Chocolate)Centi- (Milk)Milli-

1000 100 10 0 .1 .01 .001

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To use the Stair-Step method, find the prefix the original measurement starts with. (ex. milligram)

If there is no prefix, then you are starting with a base unit.

Find the step which you wish to make the conversion to. (ex. decigram)

Count the number of steps you moved, and determine in which direction you moved (left or right).

The decimal in your original measurement moves the same number of places as steps you moved and in the

same direction. (ex. milligram to decigram is 2 steps to the left, so 40 milligrams = .40 decigrams)

If the number of steps you move is larger than the number you have, you will have to add zeros to hold the

places. (ex. kilometers to meters is three steps to the right, so 10 kilometers would be equal to 10,000 m)

That’s all there is to it! You need to be able to count to 6, and know your left from your right!

1) Write the equivalent

a) 5 dm =_______m 0.5 b) 4 mL = ______L 0.004 c) 8 g = _______mg 8000

d) 9 mg =_______g 0.009 e) 2 mL = ______L 0.002 f) 6 kg = _____g 6000

g) 4 cm =_______m 0.04 h) 12 mg = ______ 0.0012 g i) 6.5 cm 3 = _______L 2.74625

0.0065L

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 = 0.382 _____kg

0.0382

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2. One cereal bar has a mass of 37 g. What is the mass of 6 cereal bars? Is that more than or less

than 1 kg? Explain your answer.

The mass of the 6 cereal bars is 222 g. It is less than 1 kg because 1 kg is equal to 1000 g.

Therefore, we need more cereal bars to have 1 kg in cereal bars.

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.

f

she has to make 110 trips in total because 110kg divide 10hg= 110hg

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.

jvj

250g because 1kg= 1000g 750g+250g= 1000g

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

1,400,000 divide 1000=1,400 km

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__________

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.

No because the tablecloth is 1.55m and the table is 1.6m long.

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 31.20m long.

10. The ceiling in Jan’s living room is 2.5 m high. She has a hanging lamp that hangs down 41 cm.

Her husband is exactly 2 m tall. Will he hit his head on the hanging lamp? Why or why not?

He will not hit the hanging lamp with his head because he is 2 meters tall and the lamp is 2.09 from the floor.

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Using SI Units

Match the terms in Column II with the descriptions in Column I. Write the letters of the correct term in

the blank on the left.

Column I Column II

_____ k 1. distance between two points

a. time

_____ e 2. SI unit of length

_____ m 3. tool used to measure length

_____ g 4. units obtained by combining other units

_____ b 5. amount of space occupied by an object

_____ h 6. unit used to express volume

_____ f 7. SI unit of mass

c

_____ 8. amount of matter in an object

_____ d 9. mass per unit of volume

o

_____ 10. temperature scale of most laboratory thermometers

_____ l 11. instrument used to measure mass

_____ a 12. interval between two events

_____ j 13. SI unit of temperature

_____ i 14. SI unit of time

_____ n 15. instrument used to measure temperature

b. volume

c. mass

d. density

e. meter

f. kilogram

g. derived

h. liter

i. second

j. Kelvin

k. length

1. balance

m. meterstick

n. thermometer

o. Celsius

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

16. Celsius degree, mass, Kelvin _____________________________________________________

They relate because they both belong to temperatures

________________________________________________________________________________

17. balance, second, mass __________________________________________________________

They realte because they belong to density and mass.

________________________________________________________________________________

18. kilogram, liter, cubic centimeter __________________________________________________

They relate because they belong to measuring volume.

________________________________________________________________________________

19. time, second, distance __________________________________________________________

They both relate because they measure time.

________________________________________________________________________________

20. decimeter, kilometer, Kelvin _____________________________________________________

They relate because they measure distance.

________________________________________________________________________________

14

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14

15

1. How many meters are in one kilometer? __________

1000

2. What part of a liter is one milliliter? __________ 1/1000 0.001

3. How many grams are in two dekagrams? __________ 20

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

kilograms?__________ 1000g= 1kg

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

1/10 0.1

In the blank at the left, write the term that correctly completes each statement. Choose from the terms

listed below.

Metric SI standard ten

prefixes ten tenth

6. An exact quantity that people agree to use for comparison is a ______________ standards .

7. The system of measurement used worldwide in science is _______________ si

.

8. SI is based on units of _______________ tens .

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

10. In SI, _______________ prefix 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.

Prefix

S.I prefixes and their meanings

Meaning

meli

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 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

_____ 6 e. hectometer

_____ 2 b. centimeter

_____ 4 c. meter

_____ 1 f. millimeter

_____ 3 g. decimeter

_____ 5 d. dekameter

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

provided.

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

think make up the decathlon?_____________________________________________________

10

8. How many years make up a decade? _______________________________________________

10

9. How many years make up a century? ______________________________________________

100

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

1/1000

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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

For Example:

11.) 0.0006.633

The decimal point moves from the first point to the second in 4 moves

therefore, making 6.633x10^-4

Since its making the number bigger, than the exponent is going to be a negative

12.) 4.69.4

The decimal point moves from the second one to the first decimal point in 2 moves

therefore, making 4.694x10^2

Since its making the number smaller, than the exponent is going to be positive

The decimal number has to be between 1 and 9, so this means that if you find a

scientific notation that isn't between those numbers, you have to change it to meet

the requirement of being considered a scientific notation.

When changing, the exponent changes as well because you have to move the

decimal point.

1. 7,485 6. 1.683

2. 884.2 7. 3.622

3. 0.00002887 8. 0.00001735

4. 0.05893 9. 0.9736

5. 0.006162 10. 0.08558

11. 6.633 X 10−⁴ 16. 1.937 X 10⁴

12. 4.445 X 10−⁴ 17. 3.457 X 10⁴

13. 2.182 X 10−³ 18. 3.948 X 10−⁵

14. 4.695 X 10² 19. 8.945 X 10⁵

15. 7.274 X 10⁵ 20. 6.783 X 10²

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SCIENTIFIC NOTATION RULES

How to Write Numbers in Scientific Notation

Scientific notation is a standard way of writing very large and very small numbers so that they're

easier to both compare and use in computations. To write in scientific notation, follow the form

N X 10 ᴬ

where N is a number between 1 and 10, but not 10 itself, and A is an integer (positive or negative

number).

RULE #1: Standard Scientific Notation is a number from 1 to 9 followed by a decimal and the

remaining significant figures and an exponent of 10 to hold place value.

Example:

5.43 x 10 2 = 5.43 x 100 = 543

8.65 x 10 – 3 = 8.65 x .001 = 0.00865

****54.3 x 10 1 is not Standard Scientific Notation!!!

RULE #2: When the decimal is moved to the Left the exponent gets Larger, but the value of the

number stays the same. Each place the decimal moves Changes the exponent by one (1). If you

move the decimal to the Right it makes the exponent smaller by one (1) for each place it is moved.

Example:

6000. x 10 0 = 600.0 x 10 1 = 60.00 x 10 2 = 6.000 x 10 3 = 6000

(Note: 10 0 = 1)

All the previous numbers are equal, but only 6.000 x 10 3 is in proper Scientific Notation.

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RULE #3: To add/subtract in scientific notation, the exponents must first be the same.

Example:

(3.0 x 10 2 ) + (6.4 x 10 3 ); since 6.4 x 10 3 is equal to 64. x 10 2 . Now add.

(3.0 x 10 2 )

+ (64. x 10 2 )

67.0 x 10 2 = 6.70 x 10 3 = 6.7 x 10 3

67.0 x 10 2 is mathematically correct, but a number in standard scientific notation can only

have one number to the left of the decimal, so the decimal is moved to the left one place and

one is added to the exponent.

Following the rules for significant figures, the answer becomes 6.7 x 10 3 .

RULE #4: To multiply, find the product of the numbers, then add the exponents.

Example:

(2.4 x 10 2 ) (5.5 x 10 –4 ) = ? [2.4 x 5.5 = 13.2]; [2 + -4 = -2], so

(2.4 x 10 2 ) (5.5 x 10 –4 ) = 13.2 x 10 –2 = 1.3 x 10 – 1

RULE #5: To divide, find the quotient of the number and subtract the exponents.

Example:

(3.3 x 10 – 6 ) / (9.1 x 10 – 8 ) = ? [3.3 / 9.1 = .36]; [-6 – (-8) = 2], so

(3.3 x 10 – 6 ) / (9.1 x 10 – 8 ) = .36 x 10 2 = 3.6 x 10 1

19


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

0.00002887

0.00001.735

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

6.633x10^-4

12. 0.0004445 17. 34,570

4.445x10^-4

1.937x10^6

3.457x10^4

13. 0.002182 18. 0.00003948

2.182x10^-3

3.948x10^-5

14. 469.5 19. 894,500

4.694x10^2

8.945x10^5

15. 727,400 20. 678.3

7.274x10^5

6.783x10^2

20

19


The Learning Goal for this assignment is:

Determine appropriate and consistent standards of measurement

for the data to be collected in a survey or experiment.

Notes Section:

There are three rules a numbers needs to have,

to be considered of having significant figures in them.

1.) Non-zero digits are always significant.

Ex. 46.78 and 3.94 these are all significant figures.

2.) Any zeros between two significant digits are significant.

Ex. 960 and 70050 have zeros that are all significant figures.

3.) A final zero or trailing zeros in the decimal portion only are significant.

Ex. 0.07030 and 0.00800 have zeros at the end that significant figures.

So the zeros in the beginning and in the middle of 7 and and 3 are NOT

significant figures.

Or in the second example, the zeros in the beginning 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

20

21


Significant Figures Rules

There are three rules on determining how many significant figures are in a

number:

1. Non-zero digits are always significant.

2. Any zeros between two significant digits are significant.

3. A final zero or trailing zeros in the DECIMAL PORTION ONLY are

significant.

Please remember that, in science, all numbers are based upon measurements (except for a very few

that are defined). Since all measurements are uncertain, we must only use those numbers that are

meaningful.

Not all of the digits have meaning (significance) and, therefore, should not be written down. In

science, only the numbers that have significance (derived from measurement) are written.

Rule 1: Non-zero digits are always significant.

If you measure something and the device you use (ruler, thermometer, triple-beam, balance, etc.)

returns a number to you, then you have made a measurement decision and that ACT of measuring

gives significance to that particular numeral (or digit) in the overall value you obtain.

Hence a number like 46.78 would have four significant figures and 3.94 would have three.

Rule 2: Any zeros between two significant digits are significant.

Suppose you had a number like 409. By the first rule, the 4 and the 9 are significant. However, to

make a measurement decision on the 4 (in the hundred's place) and the 9 (in the one's place), you

HAD to have made a decision on the ten's place. The measurement scale for this number would have

hundreds, tens, and ones marked.

Like the following example:

These are sometimes called "captured zeros."

If a number has a decimal at the end (after the one’s place) then all digits (numbers) are significant

and will be counted.

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

960.

70050.

21

22


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

22

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?

1) 2359 = ______ 4

2) 2.445 x 10−⁵= ______ 4

3) 2.93 x 10⁴= ______ 3

4) 1.30 x 10−⁷= ______ 3

5) 2604 = ______ 4

6) 9160 = ______ 3

7) 0.0800 = ______ 3

8) 0.84 = ______ 2

9) 0.0080 = ______ 2

10) 0.00040 = ______ 2

11) 0.0520 = ______ 3

12) 0.060 = ______ 2

13) 6.90 x 10−¹= ______ 3

14) 7.200 x 10⁵= ______ 4

15) 5.566 x 10−²= ______ 4

16) 3.88 x 10⁸= ______ 3

17) 3004 = ______ 4

18) 0.021 = ______ 2

19) 240 = ______ 2

20) 500 = ______ 1

23


24

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 = _______

0.8

4) 5.868 - 5.1 = _______

5) 59.9233 + 86.21 + 99.396 = _______

34.5

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 = _______

7.7

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

74.0

53.287

40.19

67.170

81.0

73.290

29.789

39.53

86.6

64.990

115.4

245.53

54.876

70.58

36.0

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 = __________ 20,000

2) 720 ÷ 7.7 = __________ 94

3) 929 x 0.3 = __________ 300

4) 300 ÷ 44.31 = __________ 7

5) 608 ÷ 9.8 = __________ 62

6) 0.06 x 0.079 = __________ 0.005

7) 0.008 x 72.91 x 7000 = __________ 4,000

8) 73.94 x 67 x 780 = __________ 3,900,000

9) 0.62 x 0.097 x 40 = __________ 2

10) 600 x 10 x 5030 = __________ 30,000,000

11) 5200 ÷ 4.46 = __________ 1,200

12) 0.0052 x 0.4 x 107 = __________ 0.2

13) 0.099 x 8.8 = __________ 0.87

14) 0.0095 x 5.2 = __________ 0.049

15) 8000 ÷ 4.62 = __________ 2,000

16) 0.6 x 0.8 = __________ 0.5

17) 2.84 x 0.66 = __________ 1.9

18) 0.5 x 0.09 = __________ 0.05

19) 8100 ÷ 34.84 = __________ 230

20) 8.24 x 6.9 x 8100 = __________ 460,000

25


Dimensional Analysis

This is a way to convert from one unit of a given substance to

another unit using ratios or conversion units. What this video

www.youtube.com/watch?v=aZ3J60GYo6U

Let’ look at a couple of examples:

1. Convert 2.6 qt to mL.

First we need a ratio or conversion unit so that we can go from quarts to milliliters. 1.00 qt = 946 mL

Next write down what you are starting with

2.6 qt

Then make you conversion tree

2.6 qt

Then fill in the units in your ratio so that you can cancel out the original unit and will be left with the

unit you need for the answer. Cross out units, one at a time that are paired, and one on top one on

the bottom.

2.6 qt mL

qt

Now fill in the values from the ratio.

2.6 qt 946 mL

1.00 qt

Now multiply all numbers on the top and multiply all numbers on the bottom and write them as a

fraction.

2.6 qt 946 mL = 2,459.6 mL

1.00 qt 1.00

Now divide the top number by the bottom number and write that number with the unit that was not

crossed out.

26


1qt=32 oz 1gal = 4qts 1.00 qt = 946 mL 1L = 1000mL

2. Convert 8135.6 mL to quarts

8135.6 mL 1 qts

8.6

946 mL

=

8.6 mL

3. Convert 115.2 oz to mL

115.2 oz 1 qts 946mL

32 oz

1 gts

=

3405.6

3406mL

4. Convert 2.3 g to Liters

2.3 g

4 qts

1 g

946 mL

1 qts

1 8.7

1000 mL

=

8.7 L

5. Convert 8.42 L to oz

8.42 L

1000 mL 1 qts

1 L

946 mL

32 oz

1 qts

=

284.820296

285 oz

Go to http://science.widener.edu/svb/tutorial/ chose #7 “Converting Volume” and do 5 more in the

space provided.

1. Convert _________ 8.2 L to _________ Gallons

8.2 L 1000 mL 1 qts

1 L

946 mL

4 qts

=

1 g 2.16

2.16 g

2. Convert _________ 2648.8 mL to _________ Gallons

2648.8 mL 1 qts

1 g

946 mL 4 qts

=

0.7

0.7000 g

3. Convert _________ 7095 mL to _________ Quarts

7095 mL

1 qts 7.5

946 mL

4. Convert _________ 4625.4 mLto _________ Gallons

4625.4 mL 1 gts

946 mL

1 g 1.22

4 qts

=

=

7.5 qts

1.22 g

5. Convert _________ 5.2 L to _________ Ounces

5.2 L 1000 mL 1 gts

1 L

946 mL

=

32 oz 175.89

1 qts

175.89 oz

27


The Learning Goal for this assignment is:

The students will learn what makes up atoms and how are atoms of one element

different from atoms of another element.

Notes Section

Atoms are the basic of the universe and matter is composed of atoms.

The number of protons determined what kind of element the atom is

The nuetrons determine if the atom is an isotope

The electrons tell you if the atom is an ion or not.

The atomic number of an element (proton number), tells you the number protons or positive

particles in an atom.

Electrovalence- something that has given up or taken electrons and becomes an ion.

Valence is a measure of how much an atom wanta to bond with other atoms.

2 types of Bonds: Covalent and electrovalent.

Ionic bonds are electrovalent bonds (groups of charged Ions held together by electric forces).

Isotopes are an atom is missing a nuetron or has an extra nuetron.

Electron "-"

Protons "+"

Nuetral "0"

Ion more "+" than "-" electrons in an atoms

Protons and nuetrons are found in the center of the atom(nucleus), hence the fact that they

make up the atom, and the electrons are found in the layers called shells or orbitals.

http://www.learner.org/interactives/periodic/basics_interactive.html

28


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.

29


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.

30


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.

31


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.

32


Summary

The article describes how atoms are made up of three parts: which are protons, electrons,

and nuetron. Protons have a positive charge, nuetrons have nuetral charge and electrons

have a negative charge. The article goes on to explain how the type of charge has to do with

whether or not an atom is an Ions have two types of form, which are cation(positive) and

anion(negative). Also, ions can either or electrovalent. Overall, elements don't have a precise

mass due to the effects positive and negative charge have on atoms.

33


34


Electron Configuration

Color the sublevel:

s = Red

d = Green

p = Blue

f = Orange

S

D

P

F

Write in sublevels

Write period, sublevel and super scripts.

Ctrl Shift =

gives you super scripts

35


The Learning Goal for this assignment is:

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

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

there are 4 diferent types of sublevel

The number of energy levels is how many subleval it has.

1s 2s 3s 4s 5s 6s 7s etc..

There are 4 different sublevels:

S-Sharp-2 Electron

P-Principle-6 Electrons

D-Diffuse-10 Electrons

F-Fundamental-14 Electrons

There are 3 rules of Clectron Configuration

1. Aufbau- NO^e ( n is the energy level , O is the orbital type or sublevel, e is the number electrons in that

orbital shell).

2. Hund rule- when electrons occupy orbitals of equal energy, one electron enters each orbital until all orbitals

cotain one electron with the same spin.

3. Puali exclusion principle- An orbital contains a maximum of 2 elelctrons and paired electrons will have

opposite spin.

36


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

37


In the space below, write the unabbreviated electron configurations of the following elements:

1) sodium ________________________________________________

1s2 2s2 2p6 3s1

2) iron ________________________________________________

1s2 2s2 2p6 3s2 3p6 4s2 3d6

3) bromine ________________________________________________

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5

4) barium ________________________________________________

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2

5) neptunium ________________________________________________

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f5

In the space below, write the abbreviated electron configurations of the following elements:

6) cobalt ________________________________________________

[Ar] 4s2 3d7

7) silver ________________________________________________

[Kr] 5s2 4d9

8) tellurium ________________________________________________

[Kr] 5s2 4d10 5p4

9) radium ________________________________________________

[Rn] 7s2

10) lawrencium ________________________________________________

[Rn] 7s2 5f14 6d1

Determine what elements are denoted by the following electron configurations:

11) 1s²s²2p⁶3s²3p⁴ ____________________

sulfer

12) 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s¹ ____________________

rubidium

13) [Kr] 5s²4d¹⁰5p³ ____________________

antimony

14) [Xe] 6s²4f¹⁴5d⁶ ____________________

osminium

15) [Rn] 7s²5f¹¹ ____________________

einsteinium

Identify the element or determine that it is not a valid electron configuration:

16) 1s²2s²2p⁶3s²3p⁶4s²4d¹⁰4p⁵ ____________________

not valid

17) 1s²2s²2p⁶3s³3d⁵ ____________________

not valid

18) [Ra] 7s²5f⁸ ____________________

not valid

19) [Kr] 5s²4d¹⁰5p⁵ ____________________

valid

20) [Xe] ____________________

not valid

1)sodium 1s 2 2s 2 2p 6 3s 1 2)iron 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6

3)bromine 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 4)barium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2

5)neptunium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 5 6)cobalt [Ar] 4s 2 3d 7

7)silver [Kr] 5s 2 4d 9 8)tellurium[Kr] 5s 2 4d 10 5p 4

9)radium [Rn] 7s 2 10)lawrencium[Rn] 7s 2 5f 14 6d 1

1s 2 2s 2 2p 6 3s 2 3p 4 sulfur 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)

38


calcium

nickel

carbon

xenon

sulfer

protactinium

one of the boxes in the 3rd section is not completely full.

the arrows in the 2s are not in the correct directions,

one has to go up and the other down. X= arrow going up

Y= arrow going down

1xy

1s:xy 2s:xy 2p:x

1s:xy 2s:xy 2p:xy xy xy 3s: x

[Ar] 4s:xy 3d:xy xy xy xy xy 4p: xy xy xy

[Ar] 4s:xy 310:x x x x

[Ne] 3s:xy 3p:x x x

1s:xy 2s:xy 2p:x x

[Ar] 4s:xy 3d:xy xy x x x

[Xe] 6s:xy 4f:xy xy xy xy xy xy xy 5d:xy xy xy x x

[Rn] 7s:xy 5f:x x x x x x

1s:xy 2s:xy 2p: x x x x

[Ar] 4s: x

39


40

Create groups for these Scientist and explain your groupings

(use the information you got from your research)


Research the Scientist and summarize their contributions to the Atomic Theory

Antoine Henri Becquerel

He discovered radioactivity which was an early contribution to atomic theory. He discovered this phenomenon while experimenting with uranium and a

photographic plate.

Niels Bohr

Bohr applied quantum theory to Rutherford’s atomic structure involving orbiting electrons. Bohr concluded that electrons traveled in stationary orbits, but

this also led to the discovery of energy levels and that there is a limited number of electron energies allowed.

Louis de Barogilie

Louis de Broglie was the scientist to introduce the theory of wave/particle duality, suggesting that particles act like waves and that waves act like particles.

This was described by the equation λ=h/p, where λ is wavelength, h is Planck’s constant, and p is momentum.

Glenn Seaborg

Glenn Seaborg discovered the element plutonium in late 1940. He went on to identify several more of the radioactive transuranium elements. He is also

responsible for discovering a wide variety of other elements in the periodic table.

Hantaro Nagaoka

Nagaoka proposed an alternative planetary model of the atom in which a positively charged center is surrounded by a number of revolving electrons. His

predictions were a very massive atomic center and how electrons revolve around the nucleus, bound by electrostatic forces.

Democritus

Democritus was the first scientist to suggest that all matter was composed of small, indivisible particles and that the properties of matter was determined

by the properties of these pieces of matter. Alot of his work is relfected in what the atom theory is now-a-days.

Marie and Pierre Curie

Pierre and Marie Curie are best known for their pioneering work in the study of radioactivity, which led to their discovery in 1898 of the elements radium

and polonium. Marie discovered that the amount of radiation depended upon the amount of element present in the compound.

Eugene Goldstein

Because of his Perforated Cathode Ray experiment, Goldstein concluded that atoms had a positively charged particle, the proton, because they flew

through the holes in the negative cathode. He is also credited with the discovery of canal rays.

Dmitri Mendeleev

Dmitri Mendeleev published his periodic table of elements in 1869. His table arranged the known elements according to their chemical properties and in

order of their relative atomic mass.

J.J. Thomson

Thomson discovered the electron through a series of experiments. He concluded that electrons were much smaller than the actual atom and the charge

to mass ratio was very large. Thomson also did experiments with cathode rays.

James Chadwick

James Chadwick discovered the neutron, a neutrally charged particle in the nucleus. His discovery lead to the fission of uranium-235 and the making of

the atomic bomb.

Erwin Shrodinger

Schrodinger, known for his quantum mechanical model, took the theories and ideas of other scientists before him and put them together to come up with

his own equation. This equation proved that energy was quantized and that orbitals were essential to electron location. This equation explained chemical

properties and reactivity of elements.

John Dalton

Dalton was the first scientist to theorize that atoms of different elements had different weights and proposed a number of ideas about the atom that

remains true today.

Lothar Meyer

He was one of the pioneers in developing the first periodic table of chemical elements. He worked both with Mendelee and Robert Bunsen. Meyer is best

known for his part in the periodic classification of the elements.

Robert Millikan

Millikan used his Oil Drop Experiment to prove the charge and mass of an electron. He also concluded that changes in energy occurred in tiny

increments, proving the Quantum Theory.

J.W. Dobereiner

Is best known for work that foreshadowed the periodic law for the chemical elements. Dobereiner also is known for his discovery of furfural and the

invention of the Dobereiner's triads.

Ernest Rutherford

Rutherford theorized that an atom had a very dense, positively charged core, due to particles being deflected by a sheet of gold foil. he also theorized that

negatively charged electrons orbited the nucleus like planets around the sun.

41


The Learning Goal for this Assignment is

The student will learn what information the periodic table provides and how periodic trends

can be explained.

Alkali Metals

Any of the elements lithium, sodium, potassium, rubidium, cesium, and francium, occupying Group IA (1) of the periodic table.

They are very reactive, electropositive, monovalent metals forming strongly alkaline hydroxides.

Alkali Earth Metals

The alkaline earth metals are six chemical elements in column (group) 2 of the Periodic table. They are beryllium (Be),

magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra).

Transitional Metals

Any of the set of metallic elements occupying a central block (Groups IVB–VIII, IB, and IIB, or 4–12) in the periodic table,

e.g., iron, manganese, chromium, and copper. Chemically they show variable valence and a strong tendency to form

coordination compounds, and many of their compounds are colored.

Inter Transitional Metals

An inner transition metal is one of a group of chemical elements on the periodic table. They are normally shown in two rows

below all of the other elements. They include elements 57-71 (lanthanides) and 89-103 (actinides). ... They have three

incomplete outermost electron shells and are all metals.

Metals

A solid material that is typically hard, shiny, malleable, fusible, and ductile, with good electrical and thermal conductivity

(e.g., iron, gold, silver, copper, and aluminum, and alloys such as brass and steel).

Metalloids

An element (e.g., germanium or silicon) whose properties are intermediate between those of metals and solid nonmetals.

They are electrical semiconductors.

Non Metals

An element or substance that is not a metal.

Noble Gases

Any of the gaseous elements helium, neon, argon, krypton, xenon, and radon, occupying Group 0 (18) of the periodic table.

They were long believed to be totally unreactive but compounds of xenon, krypton, and radon are now known.

42


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

43


Define Atomic Size:

Atomic Size

The size is determined by the amount of electrons or energy levels it has in the nucleus.

Explanation:

When you're looking at a group, the size is going to increase from

top to bottom becuase of the added energy levels.

44


Ionization Energy

Define Ionization Energy:

Ionization energy is the energy required to remove an electron from an atom.

Explanation:

It is easier to remove an electron if the size of the atom is smaller than having an atom with a

larger size, meaning that the size of the atom influences the amount of energy required to remove

an electron.

45


Define Electronegativity:

Electronegativity

Is the ability of an atom of an element to attract electrons when the atom is in a compound.

Explanation:

The bigger the mass, the more gravitational pull it has on pulling an electron. The smaller the mass, then the less

force it has to remove and electron because of how close it is to the nucleus

46


Ion Size

Define Ion Size:

Ion size is the size of an ion when refering to the elements in the periodic tables.

Explanation:

Ions may be larger or smaller than the neutral atom, depending on the ion's electric charge.

When an atom loses an electron to form a cation, the other electrons are more strongly

attracted to the nucleus, and the radius of the atom gets smaller. Similarly, when an electron is

added to an atom, forming an anion, the added electron increases the size of the electron cloud.

47


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

50


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

51


52

Salvador Gaspar

Name ____________________

Go to the web site www.darvill.clara.net/emag

1. Click on “How the waves fit into the spectrum” and fill in this table:

>: look out for the

RED words on the web site!

Frequency

Low __________, Long wavelength

Wavelength

High frequency, Short ______________

Radio Waves

Microwaves Infra-red Visible light Ultraviolet X-rays

Gamma rays

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

Communication

3. Click on “Microwaves”. They are used for cooking, mobile _________, Phones _______ Speed cameras and _________. Radar

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

cameras in police ____________ Helicopters , and alarm systems.

5. Click on “Visible Light”. This is used in 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 ________________ Radioactivity substances. We can use them to kill ________ Cancer cells,

which is called R_______________ adiotherapy .

9. My Quiz score is ____%. 100


10. Name ________________________________

Go to the web site www.darvill.clara.net/emag

Name How they’re made Uses Dangers

Gamma rays

X-rays

Ultraviolet

Visible light

Infra-red

Microwaves

Radiowaves

Given off by stars and some

radioactivity substances.

Given off by stars and strongly

by some types of nebula.

Made by special lamps and is

given off by the sun in large

quantities.

Given off by anything that's hot

enough to glow.

Given off by stars, lamps, and

flames.

Given off by transmitters in

phones and magnetrons in

microwaves.

Made by various types of

transmitters such as stars,

sparks, and lightning.

Kill living things

Radiotherapy

Radioactiv substances

See inside humans

Airport security

Kill microbes

Sun tan

CDs and DVDs

Laser printers

Remote Control

Heal sport injuries

Cook

Mobile

Wifi

Communications

Cause cancer

Cell damage

Mutations

Cell damages

Cancer

Damages retina

Cause sunburn

Skin cancer

Damgae eye retina

Overheat

Damages cataracts

Cancer

Leukemia

_____ Frequency _____ frequency,

Short wavelength ______ Wavelength

High

Long

Low

53


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:

Each type of radiation take different amounts of time to decay

Alpha radiation gives off an alpha particle:

2 protons, and 2 nuetrons

go down by 2

Beta radiation: 2 types of particles goes down by one

Negative

e-

Positive

e+

Gamma energy is a type energy given off by gamma radiation

Always associated with alpha and beta particles

very high is frequency

Half life- Time required for th quantity of a radioactive material to be reduced to one-life its original value

Fission- Breaking up large things to make them smaller

Fusion- Combining small things to make them into a bigger thing

Bohr's model- Consists of a central nucleus composed of protons and nuetrons which is surrounded by

electrons which orbit around the nucleus

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. After an atom ejects an alpha particle, a new parent atom is

formed which has two less neutrons and two less protons.

Beta decay is a radioactive process in which an electron is emitted from the nulceus of a radioactive

atom. The elctron released is a beta particle and when this happens, a nuetron becomes a proton.

Ex: rhenium has -187 decays(75z), by beta decay it becomes osmium -187 is created(76z).

B- decay: occurs when a nuetron turns into a proton and the nucleus emits an electron and

an antineutrino.

B+ decay: occurs when a proton inside a radionuclide nucleus is converted into a nuetron while

releasing a positron and an electron nuetrino.

54


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.

55


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.

56


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.

57


The Learning Goal for this assignment is:

Distinguish between bonding force 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:

58


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.

59


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.

60

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:

The number of valence electrons is the amount of dots on the element in a dot

diagram.

Ionic bonds are when metals and nonmetals are put together.

The cation gives its elemental name, and the anion has -ide in the en for the

name of the new substances.

Doesn't always happen with this with names.

The cation goes first and then the anion.

Ions are charged particles.

Ionic means molecules that consist of charge ions with opossing charges(Solid with highmelting

point and conduct electric currents.)

Ending of name with -ide means it is a negative ion.

Cations are positive and anions are negative and will usually always form an ionic bond.

its Cation if there is more positve ions and it is anion if there is more negative ions.

Electrostatic model- an application of the charge principles that opposite charges attract and

similar charges repel.

Electron affinity- the energy given off by an atom when it gain electrons.

Octet rule- the statement of that when atoms combine to form molecules, they generally each

lose, gain or share valence electrons until they meet have or share eight.

Lattice energy- energy released in the formation of an ionic compound.

Ionization potential- the energy required to remove an electron from a nuetral atom.

61


The Learning Goal for this assignment is:

Interpret formula representations of molecules and compounds in terms of composition and structure.

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.

62


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 other atoms to it. Subtract the amount of electrons in

bonds from total amount of electrons useful. Add the lone pairs to the terminal atom as well

as the central atom. Add double or triple bonds.

3: Find the formal charges. Try to get the charges to be as close to zero as possible by

moving the electrons and bonds. Add lone pairs to the ones that need it.

Two types of Covalent bonding: Non Polar and Polar bonding

Non Polar is when electrons are shared equally

Polar is when the electrons are shared unequally, but still have to meet the octet rule.

All atoms are required to meet the octet rule or else the drawing is not right.

Whe finding the middle element, you look for the least electronegative of the represented elements.

63


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.

64


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.

Add any left over bonds to make double or triple

bonds.

Then, if needed, use any lone pairs to make

double or triple bonds so that all atoms meet

the Octet Rule.

See Step 4 for total number of bonds.

Step 9

Find the formal charges for the atoms in the compound.

Arrange atoms so that all formal charges

are as close to 0 as possible.

Some central atoms do not meet the octet rule.

Boron can sometimes have only 6 electrons and

some elements in Periods 3—7 may exceed the

octet rule.

65


Linear

Molecular Geometry

Orbital Equation Lone Pairs Angle

sp AX 2 0 180

BeCl 2

Cl

Be

Cl

Beryllium D-Chloride

element bond lone pair

C

66


Trigonal planar

Molecular Geometry

Orbital Equation Lone Pairs Angle

sp 2 AX 3 0 120

BF 3

F

B

F

F

Boron Tri-Fluoride

element bond lone pair

C

67


Bent

Molecular Geometry

Orbital Equation Lone Pairs Angle

sp 2 AX 3 E 1 114

O 3

O

O

O

Tri-Oxide

element bond lone pair

C

68


Tetrahedral

Molecular Geometry

Orbital Equation Lone Pairs Angle

sp 3 AX 4 0 109.5

Phosphate

PO 4

3-

O

3-

O

P

O

O

element bond lone pair

C

69


Trigonal Pyramidal

Molecular Geometry

Orbital Equation Lone Pairs Angle

sp 3 AX 3 E 1 107

PH 3

Phosphorus Tri-Hydride

H

P

H

H

element bond lone pair

C

70


Bent

Molecular Geometry

Orbital Equation Lone Pairs Angle

sp 3 AX 2 E 2 2 104.5

H 2 O

H

O

H

Dihydrogen Oxide

element bond lone pair

C

71


Trigonal Bipyramidal

Molecular Geometry

Orbital Equation Lone Pairs Angle

sp 3 d AX 5 0 120.9

PCl 5

Cl

Cl

P

Cl

Cl

Cl

Phosphorus Pentachloride

element bond lone pair

C

72


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 Trifluoride

element bond lone pair

C

73


Octahedral

Molecular Geometry

Orbital Equation Lone Pairs Angle

sp 3 d 2 AX 6 0 90

SF 6

F

F

F

S

F

F

F

Sulfur Hexafluoride

element bond lone pair

C

74


Square Planar

Molecular Geometry

Orbital Equation Lone Pairs Angle

sp 3 d 2 AX 4 E 2 2 90

ICl 4

-

Cl

-1

Cl

Iodine Tetrachloride

Ion

I

Cl

Cl

element bond lone pair

C

75


Orbitals Equation

Lone

Pairs

Angle

Name

sp AX2 0 180

Linear

sp 2 AX3 0 120

Trigonal Planar

sp 2 AX3E 1 114

Bent

sp 3 AX4 0 109.5

Tetrahedral

sp 3 AX3E 1 107

Trigonal Pyramidal

sp 3 AX2E2 2 104.5

Bent

sp 3 AX5 0 120/90

Trig. Bipyramidal

sp 3 d AX3E2 2 90

T-Shaped

sp 3 d 2 AX6 0 90

Octahedral

sp 3 d 2 AX4E2 2 90

Square Planar

76


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-

77


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


http://www.bbc.co.uk/education/guides/zm9hvcw/revision

LG: The student will learn how Hydrocarbons are named and the

general properties of Hydrocarbons.

A homologous series is a family of hydrocarbons with similar chemical

properties who share the same general formula.

Alkanes- names end in -ane; are saturated; General formula- C n H 2n+2 ; all

contain only single bonds. Naming rules: longest unbranched chain

containing functional group with a number, and names the branches as well

as indicate the number

Alkenes- names end in -ene; are unsaturated; General Formula- C n H2 n ;

contain a carbon to carbon double bond. Naming rules: Same as Alkenes, but

the position of the carbon must be identified.

Cycloalkanes- names end in -ane and begin with cyclo-; General Formula

C n H 2n ; are saturated and contain only one bond.

All hydrocarbons can undergo combustion reactions with oxygen

to give the same product.

Also hydrocarbons burn when they react with oxygen in the air.

Combustion equation

Hydrogen+Oxygen -) Carbon dioxide+Water

Alkenes are more reactive than alkanes and cycloalkanes becuase

they have double bond.

Addition Reactions: double bonds breaks when the reaction molecule

attacks and adds on across it. Only Alkenes can participate

in addition reactions.

Bromine water can be used as a test for unsaturation.

The addition of bromine to an alkene is called bromination.

The addition of hydrogen to an alkene is called hydrogenation.

The addition of water to an alkene is called hydration.

ISOMERS- Compouds with the same molecular formula but diff.

chemical structure. Same number of type of atom, but may have

diff. physical and chemical property.

Molecular Formula- C4H10

80


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