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HONORS CHEMISTRY

MS. SONDERLEITER

Source: http://www.siraze.net/chemistry/sezennur/subjects/comics/comics22.htm

UNIT 1 – BASIC SKILLS (Measurement & Units, Lab Equipment & Skills, Significant Figures, Density,

Scientific Notation, Matter, Energy)

Name:__________________________________________

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Lesson 0 – Background Skills

Students MUST be able to

• Define the measured quantities of mass, length, area, perimeter, volume, time,

temperature, and energy.

• Choose the appropriate units for each measurement.

• Explain the basis of different measuring systems and discuss the benefits of the

metric system in science.

• Memorize the metric prefixes, their symbols, and their corresponding meanings.

• Use dimensional analysis to covert between units.

• Perform conversions within the metric system using memorized conversion

factors and between English and metric units using provided conversion factors.

• Write numbers in scientific notation.

• Convert numbers from scientific notation to standard notation.

Category Definition Measurement Tool Units

Mass

Length

Height

Width

Perimeter

Area

Volume

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Time

Temperature

Energy

Units in Everyday Life: We use units every day, often without even realizing it. In the

statements that follow, you will find a wide variety of interesting facts, but each is

missing a crucial piece of information – the dimensions (units)! All the statements are

meaningless until you supply the appropriate units. On the basis of your experiences, try

to match the appropriate units for the list provided.

carats

cm

degrees Celsius

degrees

Fahrenheit

feet

grams/mL

inches

kcal (Cal)

kilograms

kilometers

liters

miles

miles per hour

milligrams

pounds

stories

tons

yards

1) America’s tallest building (Sears Tower in Chicago) is 110 __________high.

2) The Empire State Building in New York is 1250 _______________high.

3) The Nile is the world’s longest river. It is 4180 ______________long.

4) The Amazon River in South America is 6296 _______________long.

5) The coldest temperature ever recorded was -128.6 ___________ in Vostok,

Antarctica, in 1983.

6) The highest recorded temperature in the United States was in Death Valley,

California, when the mercury reached 57____________.

7) The world record rainfall occurred in Cherrapunji, India, where 1042

____________ of rain fell in one year.

8) The largest recorded hailstone to ever fall landed in Coffeyville, Kansas, in 1979.

It had a diameter of 44.5 ________!

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9) The longest punt in NFL history was by Steve O’Neal of the New York Jets. He

kicked the football 98 ____________.

10) The largest seed in the world is that of the coc-de-mer coconut tree, which may

weigh as much as 40 _____________!

11) The world’s largest meteorite is located in Southwest Africa. It weighs 650

_________.

12) The most popular soft drink in the world is currently Coca Cola. More than 210

million ____________were consumed each day in 1990.

13) The largest diamond in the world was mined from South Africa in 1905 and

weighs 3106 ____________.

14) Earth is the densest of the nine planets, with an average density of 5.515

___________.

15) The world’s fastest aircraft is the Lockheed SR-71 Blackbird, clocking a record

speed of 2,193.67 ________________.

16) The largest gold nugget ever found has a mass of 100 _____________!

17) One large chicken egg contains an average of 274 _____________ of cholesterol.

18) A 16-year old male requires an average of 2800 ____________ of energy per day

while and average 16-year old female requires only 2100 ______________.

Concluding Questions/Main Points:

1) Why is it essential that appropriate units accompany all measurements?

2) Individuals who travel to regions of the world with poor sanitation are warned to

filter or boil their water before drinking it to remove deadly water-born pathogens

that cause diseases such as cholera and typhoid. If you were traveling in a region

known to have a polluted water supply, would you drink water that your host said

had been heated to 100 degrees for five minutes? Explain?

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English System Metric System

Description

Pros

Cons

1) Which system of units will we use in chemistry? Why?

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Metric Prefixes and Powers

Prefix Symbol Power

of 10

Meaning Decimal Notation Example

(mass)

109

106

103

10-2

10-3

10-6

10-9

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Perform the following unit conversions.

1) 11.7 mm to m

2) 0.9 cg to g

3) 44 nm to m

4) 4.53 kg to g

5) 9,350 µs to s

6) 0.0014 Gm to m

7) 76 mL to µL

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8) 0.450 Mm to cm

9) 0.089 kg to mg

10) 150.0 lbs to kg

11) 12,850 ft to miles

12) 48,987 min to hr

13) 18 in to cm

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14) 6.0 gal to mL

15) 56,098 min to days

16) 3.2 tons to grams

Additional Practice:

• http://chemistry.about.com/od/convertcalculate/a/conversions.htm

• http://proton.csudh.edu/homeworkcs/hwconvertvolcsn7.html

• http://proton.csudh.edu/homeworkcs/hwconvertmasscsn7.html

• http://proton.csudh.edu/homeworkcs/hwconvertlengthcsn7.html

Standard Notation Scientific Notation

8,700,000

3.4 x 10-3

0.00098

93

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Lesson 1 – Introduction to the Laboratory

Students will be able to

• Name and locate standard lab equipment and describe its use.

• Explain and identify the standard lab safety rules and procedures.

Equipment Name What do you think it is used for?

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Picture reference: Laboratory Equipment. Created by the Chemistry Faculty and the Media Center at Santa Monica College. 8/30/2005 <http://homepage.smc.edu/chem10/Equipment.html>

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Archer School for Girls

Science Safety Agreement

You are enrolled in a laboratory science course. Experiments are part of the foundation

of scientific knowledge. A laboratory should be a safe place to work and learn. The

experiments you will be conducting throughout the year have been selected, in part,

because of their ability to be done safely. Your teacher will go over the following rules

and guidelines for the lab. You will be given the opportunity to ask questions about these

rules. After your teacher has discussed these rules, you and your parents will sign a copy

of this agreement to be kept by your teacher.

1. Perform all experiments as directed. Never make substitutions or changes to a

procedure without first consulting your teacher.

2. Be properly prepared for each experiment. Be sure to read and complete the pre-

lab and know the hazards before you do the experiment.

3. Wear the appropriate protective equipment. Always wear safety goggles in the

lab when chemicals are being used – no exceptions!

4. Know the locations and uses of the safety devices in the lab. This includes the

shower, eye wash, fire extinguisher, sinks, and first aid kit.

5. Act in a responsible manner at all times. No running or horseplay in the lab.

6. Wear closed-toed shoes in the lab.

7. Tie back your hair if it is long, and refrain from wearing large jewelry.

8. Never taste a chemical, even if it is a substance found in households.

9. No food or drink is permitted in the lab at any time.

10. Turn off a Bunsen burner or heat source whenever you are not using it. Never

operate a heating source unattended.

11. Read chemical labels carefully. Know what you are using!

12. Never touch chemicals unless directed.

13. When checking for chemical odor – waft, don’t sniff the chemical.

14. Report all accidents, inures, or near misses to your teacher immediately, no matter

how small. (This includes cuts, burns, acid spills, etc.)

15. Dispose of chemicals properly. Your teacher will inform you of proper waste

disposal at the beginning of each experiment.

16. Do not use more material in a lab than is required, and never return unused

reagent back to the reagent bottle.

17. Clean up spills immediately. (This includes water.) Keep your lab area neat and

free of clutter.

18. Never take chemicals out of the laboratory.

19. Wash your hands with soap and water at the end of each experiment.

20. Allow plenty of time for hot glass to cool. Remember: hot glass looks like cool

glass!

21. Always use common sense and good judgment in the lab.

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* It is NOT recommended that contact lenses be worn in the lab. The permeability of

the lenses allows chemicals to enter the eye and can cause irritation. If you choose to

wear contact lenses, please indicate this at the bottom of the page.

- - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - -

I, ________________________, have read, understand, and agree to follow these science

safety rules and procedures. I agree to abide by any additional instructions, written or

verbal, provided by my science teacher during an experiment.

Student Signature ___________________________________ Date __________

Parent Signature ____________________________________ Date __________

I wear contact lenses: Yes _______ No_______

Please list any allergies below:

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Lesson 2 – Measurement & Uncertainty


• Explain why scientists must have a common system for performing

measurements.

• Report measurements with the appropriate number of digits using the rules of

significant figures.

Measure the length of the leaf below.

Source: http://cyberbridge.mcb.harvard.edu/math_2.html

Significant Figures Rule for Measurement in Science:

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1) Read the material provided at each station.

2) Get teacher signature for proper equipment use.

3) Fill out worksheet.

Station 1: Triple Beam Balance

Sample Mass of Beaker Mass of Beaker and

Sample

Mass of Sample

Itself

Sample #1

Sample #2

What is the smallest graduation (place value) on the triple beam balance?

What digit/place value do you have to estimate?

Station 2: Electronic Balance

Sample Mass of Sample Itself

Sample #1

Sample #2

Sample #3

Sample #4

What makes a top-loading balance easier to use than a triple beam balance?

What is the smallest graduation (place value) on the top-loading balance?

What digit (place value) do you have to estimate?

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Station 3: Beaker

Sample Color Volume

Beaker #1

Beaker #2

Beaker #3

Beaker #4

What is the smallest graduation (place value) on each of the 4 different beakers that you

used?

Beaker #1 ___________ Beaker #3_______________

Beaker #2 ___________ Beaker #4_______________

What digits (place value) do you have to estimate on each beaker?

Beaker #1 ___________ Beaker #3_______________

Beaker #2 ___________ Beaker #4_______________

Which beaker gives the most accurate measurement? Why?

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Station 4: Graduated Cylinder

Sample Color Volume

Cylinder #1

Cylinder #2

Cylinder #3

Cylinder #4

What is the smallest graduation (place value) of each of the cylinders used?

Cylinder #1 ___________ Cylinder #3_______________


What digit (place value) do you have to estimate on each of the cylinders?



Conclusion Questions/Main Points:

1) What determines the accuracy of a measuring device?

2) How do scientists reflect the accuracy of their equipment when reporting a

measurement?

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Lesson 3 – Significant Figures


• Explain the importance of significant figures.

• Count significant figures.

• Perform calculations and write answers with the proper significant figures.

What are significant figures?

Rules for Counting Significant Figures

How many significant figures are in each of the following numbers?

1. 3.000010

2. 98,001,000

3. 0.00107

4. 85,000.

5. 709

6. 0.007

7. 600

8. 9.000

9. 6.530

10. 9.007


• http://proton.csudh.edu/homeworkcs/hwsigfigurescsn7.html

• http://science.widener.edu/svb/tutorial/sigfigures.html

• http://www.lon-capa.org/~mmp/applist/sigfig/sig.htm

• http://www.dallassd.com/our%20schools/high%20School/Chemsite/hotpot/sf.htm

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Lesson 4 – Significant Figures & Calculations


• Count significant figures.

• Perform calculations and write answers with the proper significant figures.

Why?

Addition & Subtraction Rules

Practice:

1) 149.7 + 23.55 + 2000.34 =

2) 1.0322 x 103 + 4.34 x 10

3 =

3) 95.3 – 12.678 =

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Multiplication & Division Rules

Practice:

1) (0.0432)(2.909) (4.43 x 108) =

2) (8.507) / (0.0004) =

Mixed Math (Ahhhhhhhhhh!)

Practice:

1) 7.33(45.6 – 5.09) =

2) (0.8922) / [(0.00932)(4.03 x 102) =


• http://www.teacherbridge.org/public/bhs/teachers/Dana/SigFigOperations.html

• http://www.aaaknow.com/g71f_nx1.htm

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Significant? . . . . I Don’t Think So

By Michael Offutt, 1997, Chemistry Songbag II

You know, I really love my calculator,

When I multiply, divide, subtract or add.

I get so many digits in my answer.

Lots and lots of digits can’t be bad . . . right?

Well, my science teacher told me to be careful,

Some of those digits have to go,

She said I have to learn to round my answer.

There are rules that science students have to know.

This calculator is my special friend,

Its little brain never makes mistakes,

Every number on its screen is mathematically correct,

But significant? . . . I don’t think so.

Every scientific measurement’s uncertain,

Uncertain to varying degrees.

So when one or more are used in calculations,

The answer also has uncertainty.

When you multiply or divide uncertain values,

Your answer should be rounded when you’re through.

To the least number of sig figs in any of those values.

It’s really, really not that hard to do.

When you add or subtract uncertain values,

Decimal places are the key.

Count the fewest number of decimal places in those values.

And then you should round accordingly.

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Lesson 5 – Matter


• Define matter.

• Differentiate (verbally and with molecular level drawings) between the different

forms of matter (atoms, molecules, compounds, elements, pure substances,

mixtures, homogeneous mixtures, heterogeneous mixtures).

What is matter?

Sort the following into three categories: matter, not matter, or not sure.

peanut butter

water

fish

light

garbage

time

motion

the human brain

carbon dioxide

air

yourself

energy

an idea

heat

tree

energy

Pieces of Matter

Atoms Molecules

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BUT atoms and molecules are VERY small. We can’t see just 1!

Source: http://wps.prenhall.com/wps/media/objects/165/169061/GIFS/AAAUASO0.JPG

Molecular Level Drawings . . .

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Procedure: Examine the contents of the nine petri dishes that are set up in the lab.

Petri

Dish

#

Contents are either (choose one)

a pure substance or a mixture

Contents contain: element,

compound, or mixture

Atoms,

molecules, or both are present

1

2

3

4

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Petri

Dish

#

Contents are either (choose

one) a pure substance or a

mixture

Contents contain: element,

compound, or mixture

Atoms,

molecules, or both are present

5

6

7

8

9

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Lesson 6 – States of Matter & Drawings


• Describe the different states of matter in relation to molecular motion and

position.

• Identify a type of matter based on a molecular level drawing.

• Create a molecular level drawing to represent a specific type of matter.

Source: http://www.suntrek.org/images/states.gif

ANIMATION

Solid Liquid Gas Plasma

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Phase Changes

Source: http://www.gcsescience.com/Interconverting-Solid-Liquid-Gas.gif

Molecular Level Drawings

1) Draw molecular level diagrams to show the difference between a mixture and a

pure substance.

2) Draw molecular level diagrams to show the difference between a compound and a

mixture.

3) Draw molecular level diagrams to show the difference between a homogeneous

mixture and a heterogeneous mixture.

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Source: http://intro.chem.okstate.edu/HTML/SCFIMG/SCH111.gif

1) Which are solids? Liquids? Gases?

2) Which are mixtures and which are pure substances?

3) Which are atoms? Molecules?

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Lesson 7 – Physical Properties of Matter


• Define and give examples of physical properties of matter.

What is a physical property?

Physical Properties:

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Lesson 8 – Density


• Define and calculate the density of solids, liquids, and gases.

• Describe how to determine the density of various substances experimentally.

• Determine whether a substance will sink or float in water.

• Explain why the density of a substance is independent of the amount of the

substance present.

What is density? What does it mean when you say something is dense?

Finding Density in the Lab

Measuring Mass Measuring Volume

Solid

Triple Beam Balance Gizmo

Geometric Solid

Irregular Solid

Liquid

Gas

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Calculations

1) Given the following mass, volume, and density information, calculate the missing

quantity.

a. mass = ? ; volume = 124.1 mL; density = 0.821 g/mL

b. mass = 0.721 lbs; volume = 241 cm3; density = ?

2) A cube of metal weighs 1.45 kg and displaces 542 mL of water when immersed.

Calculate the density of the metal.

3) If 5.67 g of silver, which has a density of 10.5 g/cm3, is dropped into a graduated

cylinder containing 34.5 mL of water, to what volume will the water level rise?

4) You find a gold nugget that has a mass of 253.36g and a density of 19.3 g/mL.

a. What volume of gold do you have?

b. If it takes 1.32 mL of gold to make 1 ring, how many rings can you make?

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5) You find a mystery tube of gas that has a volume of 3.54 mL and a mass of

0.0063g. What is the density and identity of the gas?

Gas Density (g/mL)

Ar 0.001784

He 0.0001785

N2 0.001250

O2 0.001429

6) For a material to float on the surface of water, the material must have a density

less than that of water (1.0 g/mL) and must not react with the water or dissolve in

it. A spherical ball has a radius of 0.50 cm and weighs 2.0 g. Will this ball float or

sink when placed in water? (Hint: the volume of a sphere = (4/3)πr3)

7) The density of air at ordinary atmospheric pressure and 25°C is 1.19 x 10-3

g/mL.

What is the mass, in kilograms, of the air in a room that measures 12.5 x 15.5 x 20

ft?

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Lesson 9 – Separating Mixtures


• Describe how to separate simple mixtures.

• Define and describe the scientific reasoning behind the processes of filtration and

distillation.

Heterogeneous Mixtures

Mixture Composition How to separate . . .

Rocks and Water

Sand and Water

Homogeneous Mixtures

Mixture Composition How to separate . . .

Salt and Water

Rubbing Alcohol and Water

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Filtration

Distillation

Extraction

Coming Later On (Chromatography)

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Lesson 10 – Energy


• Define energy.

• Differentiate between energy and temperature.

• Explain and give examples of the law of conservation of energy.

• Convert between Celsius, Fahrenheit, and Kelvin temperature scales.

What is energy?

What are some of the different types of energy?

Key Terms:

• Energy Transfer –

• Law of Conservation of Energy –

• Heat –

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• Endothermic –

• Exothermic –

• Temperature –

Temperature Conversions:

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Lesson 11 – Energy Transfer & Specific Heat


• Define specific heat.

• Calculate energy changes, final and initial temperatures, or specific heat using the

equation q = mc∆T.

Heat Transfer:

Specific Heat:

Source:

http://www.explorelearning.com/ELContent/gizmos/ELScience_Deliverable/ExplorationGuides/images/EL_MSPS_Calorie2.gif

Relationship Between Energy Transfer, Temperature, and Specific Heat:

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Practice:

1. Calculate the specific heat of a substance if a 35 g sample absorbs 48 J as the

temperature is raised from 293 K to 313 K.

2. The temperature of a piece of metal with a mass of 95.4 g increases from 298.0 K

to 321.1 K when the metal absorbs 849 J of energy as heat. What is the specific

heat of the metal? Using the chart above, determine the identity of the metal.

3. If 980 kJ of energy is transferred to 6.2 L of water at 291 K, what will the final

temperature of the water be?

4. How much energy as heat must be transferred to raise the temperature of a 55 g

sample of aluminum from 22.4 °C to 94.6 °C? The specific heat of aluminum is

0.897 J/g°C.

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