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Honors Chemistry

GUIDED NOTES Unit 1: Matter & Measurement

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Unit Vocabulary:

1. S.I. unit2. Meter3. Liter4. Gram5. Mass6. Weight7. Volume8. Density9. Intensive

10. Extensive11. Significant Figures12. Precision13. Accuracy14. Matter15. Element16. Compound17. Mixture18. Heterogeneous Mixture

19. Homogeneous Mixture20. Pure Substance21. Particle Diagram22. Chromatography23. Filtration24. Distillation25. Scientific Notation

Unit Objectives: When you complete this unit you will be able to do the following…

1) Classify types of matter

2) Draw particle diagrams to represent different types of matter

3) Recognize various techniques that can be used to separate matter

4) Convert between units of measurements

5) Differentiate between accuracy and precision

6) Write numbers in scientific notation

7) State rules to determine significant figures

8) Count significant figures

9) Understand the importance of significant figures

10) Calculate the volume and density of an object

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Can NOT be separated by physical means

CAN be Separated by PHYSICAL means

Can NOT be separated by chemical means

Separated by chemical means, only

Same composition throughout

Different composition throughout

Particle Diagram

Particle Diagram

Particle Diagram

Particle Diagram

Matter

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Practice Problems:

1. Which particle diagram(s) represent a mixture?

2. Which particle diagram(s) represent a pure substance?

3. Which of the following particle diagrams represents a mixture ofone compound and one element?

4. Which particle diagram represents a diatomic element?

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Properties of Matter:

Physical properties are the constants about a substance; can useour senses to observe them; do not require chemical analysis

Example:

o Extensive Property: a property that depends on howmuch material you are dealing with

Ex:

o Intensive Property: a property that does not depend onhow much material you are dealing with (help identifymatter; a constant about that particular type of matter)

Ex:

Chemical properties include behaviors substances adhere to when they__________ with other substances

Examples:

Guided Practice: Identify the following as being intensive, extensive, or chemical properties.

____________ 1. The mass of copper wire is 255 g.

____________ 2. The boiling point of ethyl alcohol is 77°C.

___________ 3. Baking soda reacts with vinegar to make carbon dioxide gas.

____________ 4. The density of mercury is 13.6g/mL.

____________ 5. The solubility of sodium chloride in water is 40g/100mL of water.

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Physical vs. Chemical Changes

Matter is always changing. Ice in your drink melts. Wood in your fire burns.

Physical Change – a change that does NOT alter the chemical properties of a substance (example: ______________, ______________); change in size or shape; ____________________________

*PHYSICAL processes can be reversed ________________________

Example: ice melting to become liquid (its still water!)

Chemical Change – a reaction in which the composition of a substance is changed (example: __________); properties _______________________

1. Signs of a chemical rxn:

Example: firewood burning

1. 2. 3.

Change of Matter Physical or Chemical?

Burning toast

Making ice cubes

Lighting a candle

Spoiling milk

Making kool-aid

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Elements vs. Compounds Element = _____________________________________________________ Compound = ____________________________________________________

1. Circle ( ) all the elements and underline the compounds below.2. On the line provided, record the number of different symbols within the species.

CO ___ Mg ___ Co ___

C2H5OH ___ Al(CN)3 ___ Cl2 ___

H2SO4 ___ He ___ NI3 ___

O2 ___ H2O ___ NaCl ___

C ___ Cu ___ I ___

Questions: 1) Does each compound have the same number of symbols? ____2) For each ELEMENT above, how many total symbols are listed? __3) What is the minimum number of symbols that must be present in

order for a species to be considered a compound? __

Understanding Compound Formulas: Within a compound, you may see subscripts. These subscripts tell you the

number of each type of atom that is present.

Example:

# carbon atoms __ # oxygen atoms __

If there are parentheses present around two or more atoms, the subscriptapplies to all atoms within the parentheses.

Example:

# aluminum atoms __ # carbon atoms __ # nitrogen atoms __

If one of the atoms within the parentheses has a subscript, you multiply thisnumber by the number outside of the parentheses.

Example:

# iron atoms __ # sulfur atoms __ # oxygen atoms ___

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* The Common Elements *Rules for writing element symbols: 1)

2)

* Symbol * * Name * * Symbol * * Name *Ag silver I iodine Al aluminum K potassium Ar argon Kr krypton As arsenic Li lithium Au gold Mg magnesium B boron Mn manganese Ba barium N nitrogen Be beryllium Na sodium Br bromine Ne neon C carbon Ni nickel Ca calcium O oxygen Cl chlorine P phosphorus Co cobalt Pb lead Cr chromium Ra radium Cs cesium Rb rubidium Cu copper Rn radon F fluorine S sulfur Fe iron Si silicon Fr francium Sn tin H hydrogen Sr strontium He helium U uranium Hg mercury Xe xenon

Zn zinc

MEMORIZE both directions (symbol to name, name to symbol) for Quiz on _____________

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Separation of Matter

Separation Apparatus Type of Separation (Physical or Chemical)

Description of Technique

What types of matter will it separate?

Filtration

Watch Glass Evaporation

Crucible Evaporation

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Separation of Matter (continued)

Distillation

Chromatography

On the other hand _____________________ requires reacting a sample with something else in order to turn it into a completely different compound

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SCIENTIFIC NOTATION – method for expressing very large orsmall numbers easily (Example: ___________________)

For example, the number 1,000,000 is in standard formation format. The scientific notation of this number is 1.0 x 106

We always move the decimal place to make the ____________(the numberout in front) between _______________

We then arrange the ___________ (the number up to the right of the ten) Now, _______________ if you were to take the 1.0 and move the

decimal place 6 places to the right (since it is a positive number), you wouldget the original number (1,000,000)

Example: 123000000000

Guided Practice – Write the following numbers in scientific notation (remember the mantissa rule!)

1. 34000000 =

2. 0.0000067 =

3. 25,864 =

Now, write the following scientific notations in standard (normal) notation form:

4. 5.7 x 108 =

5. 6.34 x 10-11 =

Calculator Practice: First, let’s enter the number 2.3 x 10-5 in scientific notation:

1. Type “2”2. Type the decimal point3. Type “3”4. Then press the “ee” “EXP” or “ ” key(s)5. Press the “+/-“ key (NOT the “—“ or “subtract” key)6. Type “5”

Next, let’s enter that number by 1 mole, or 6.02 x 1023. What do you get for

your answer? _________________

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Measurements and the Metric System In chemistry we measure matter using ____ units. This is an abbreviation for _________________________________.

SI BASE UNITS (AKA Base Units): **If you forget, use Table D in your Reference Tables!

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SI Metric Prefixes

Prefix Symbol Numerical (Multiply Root Wordby)*

Exponential

tera T 1,000,000,000,000 1012

giga G 1,000,000,000 109

mega M 1,000,000 106

kilo k 1,000 103

hecto h 100 102

deca da 10 101

no prefix: 1 100

deci d 0.1 10¯1

centi c 0.01 10¯2

milli m 0.001 10¯3

micro 0.000001 10¯6

nano n 0.000000001 10¯9

pico p 0.000000000001 10¯12

femto f 0.000000000000001 10¯15

atto a 0.000000000000000001 10¯18

*Example: In the word kilometer, the root word (base unit) is “meter” andthe prefix is “kilo.” Kilo means multiply the root word by 1000. Therefore, one kilometer is 1000 meters (1 km = 1000 m).

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Conversion Factors – a mathematical expression that relates two units thatmeasure the same type of quantity

Examples: -

*Rest Assured! For the Regents, the most you will have to convert will be between themilli-/kilo-/base unit (g, L, etc.). This is always a matter of ___________________. You must also make sure you move the decimal the ___________________ (right or left, which depends on whether you are converting from small to big or vice versa).

TRICK: kilo hecto deca base unit deci centi milli

k h d base unit d c m

Let’s practice! 1. A car travels 845 km. How many meters is this?

2. Convert 0.0290 L to milliliters.

3. Convert 2500mL to liters.

4. 3 g = _______ kg

5. 1 km = ______ m

6. 1 kg = _______ g

7. 1 L = ________ mL

8. 7 m = _______ mm

9. 12 mL = ______ L

Compare by placing a <, >, or = on the line provided:

10. 56 cm __ 6 m

11. 7 g __ 698 mg

Once you get your answer, check it! Does it make sense?

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Dimensional Analysis

Often you will be required to solve a problem with mixed units, or to convert from one set of units to another. Dimensional analysis is a simple method to accomplish this task.

Ex: How many minutes are there in 15 days?

Solution A:

STEP 1: Figure out the units that you have and the steps to get to the units that you need.

HAVE NEED

Days (d) Hours (h) Minutes (min)

STEP 2: Make a “grid” and plug in the numbers to make your first conversion. The number/units you HAVE goes in the top left, the number/units you NEED go in the top right, and the conversion factor goes in the bottom right.

Since 24 hours and 1 day are equivalent, you are actually multiplying 15 days by a factor of 1. This means that the magnitude of your number stays the same and only the units change. In other words, 15 days = 360 hours

STEP 3: Do the same for your second conversion. Now, we will use the conversion factor 60 minutes = 1 hour

STEP 4: Do the same for your second conversion.

Now you try one…How many minutes are there in the month of October?

21,600 min 360 h 60 min 1 h

=

15 d 24 h 1 d

= 360 h Have

Need

Conversion Factor

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ACCURACY VS. PRECISION

Accuracy – ____________________________________________ Ex: Hitting bulls eye when you are aiming for it

*For most experiments, ______________ means _____ fromthe expected value

Precision – _____________________________________________ _____________________________________________________

*For an experiment with +/- 5% as the margin for accuracy,that means the difference between the highest and lowestpercent error cannot exceed a ____________________

Ex: If the highest percent error for an experiment is +7.6%, and the lowest is -5.4% that range is 13.0%, which means that experiment was not precise

Practice: Cheryl, Cynthia, Carmen, and Casey take target practice in PE. Assuming that they were all aiming at the bulls eye, match each target with the proper description.

(a) Accurate and precise (b) Not precise, but one piece of data accurate (c) Precise but not accurate (d) Neither precise nor accurate

Practice: The following data was collected during a lab experiment. The density of the cube should be 10.8 g/mL. Is this data is accurate, precise, both, or neither? Justify your answer. ______________________________________________ ______________________________________________________________

Trial Number Density of Cube 1 6.2 g/mL2 6.3 g/mL3 6.5 g/mL

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SIGNIFICANT FIGURES - also known as Sig Figs (SF) A method for handling ____________________ in all measurements This arises due to the fact that we have different equipment with different

degrees of ___________________ Significant figures are associated with ____________________________ _______________________ do _______________ when determining sig figs

o Ex: Atomic masses on periodic tableConversions (1in = 2.54 cm)

Examples: 1. Reading a ruler

We know for sure that the object is more than _____, but less than _____

We know for sure that the object is more than _____, but less than _____

This ruler allows us to estimate the length to ________

2. Reading a graduated cylinder:

► Measurements are read from the bottom of the _________

►Which gives a volume reading of _______

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The Atlantic/Pacific Method - another way to determine the # sig figs in a number

1) Determine if a decimal point is present. If a decimal is present, think of “P”for “present.” If there is no decimal, think of “A” for “Absent.” P standsfor the Pacific coast and A stands for the Atlantic Coast.

2) Imagine the number you are looking at is a map of the USA. Begin countingfrom the correct side of the number (Atlantic/right side or Pacific/leftside) based on what you determined in step 1. Consider the first nonzeronumber you land on the start of your count. Consider each digit from hereon out significant as well until you reach the other end of the number.

3.

Determine the number of significant numbers in each of the following:

1) 357 _______

2) 3570 _______

3) 3570. _______

4) 0.357 _______

5) 0.0357 _______

6) 3.570 x 103 _______

7) 0.3570 _______

Pacific Coast

Decimal is Present

1. Start @ 1st

NONZERO

2. Count allthe way to the Atlantic—NO EXCEPTIONS

Atlantic Coast

Decimal is Absent

1. Start @ 1st

NONZERO

2. Count allthe way to the Pacific—NO EXCEPTIONS

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Rules for Determining Number of Significant Figures in a Given Number

Rule Example 1. All nonzero numbers (ex: 1 – 9) are always

significant 123456789 m

1.23 x 102

2. Zeros located between nonzero numbersare significant

40.7 L

87,009 km

3. For numbers less than one, all zeros tothe left of the 1st nonzero number are NOT significant

0.009587 m

0.0009 kg

4. Zeros at the end of a number and to theright of a decimal point are significant

85.00 g

9.070000000 L

5. Zeros at the end of a whole number maybe significant or not. If there is a decimal after the last zero, they are significant. If there is not a decimal point after the end zeros, they are NOT significant

2000 m

2000. m

6. Exact numbers have an infinitenumber of significant figures 1 ft = 12 inch

PRACTICE: Measurement Number of Significant Figures Rule(s) Applied

1020 mL 1200 m 1200. L 1200.00 mm 0.001 km 10.00 L 12000 m 00.100 cL 22.101 mm 101,000 km

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Rules for Using Sig Figs in Calculations General Rule Final answer must be expressed in the lowest amount of significant figures

that were originally given to you (you can’t create accuracy when you didn’t have it to start with!)

Operation Rule Examples

Multiplication/Division Perform operation as normal & express

answer in least # sig figs that were given to

you

12.257 x 1.162 =

Addition/Subtraction Line decimal points up; round final answer to lowest decimal place (least accurate) value

given

3.95 2.879

+ 213.6____

Examples: 5.1456 – 2.31 = _______

69.25/45.8 = _________

Rules for Calculations with Numbers in Scientific Notation: Rule Example Addition/Subtraction All values must have the same exponent. Result is the sum or difference of the mantissas, multiplied by the same exponent of 10

4.5 x 106 - 2.3 x 105

Multiplication mantissas are multiplied and exponents of 10 are added

(3.1 x 103) (5.01 x 104)

Division mantissas are divided and exponents are subtracted 7.63 x 103 / 8.6203 x 104

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MEASURING MATTER

1. Mass vs. Weight

*We really only work with ________ in chemistry class!** We have the same _________ whether we are on earth or on the moon. The different forces of gravity on each cause us to weigh more on earth than on the moon though (this is why we float on the moon!)

2. Volume - amount of _____________ an object takes up Techniques:

Liquids

Regular Solids

Irregular Solids

3. Density: amount of mass in a given space; _________ of mass to volume

Formula (Table T):

BOX A BOX B

Which box has a higher density? Explain your answer. ________________________________________________________________________________________________________

MASS WEIGHT

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Density Problems – Show all work!

*Note: the density of water is ______________

1) What is the density of an object with a mass of 60 g and a volume of 2 cm3?

2) If you have a gold brick that is 2.0 cm x 3.0 cm x 4.0 cm and has a mass of48.0 g, what is its density?

3) If a block of wood has a density of 0.6 g/ cm3 and a mass of 120 g, what isits volume?

4) What is the mass of an object that has a volume of 34 cm3 and a density of6.0 g/cm3?

5) Which is heavier, a ton of feathers or a ton of bowling balls?

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Percent Error

Measurement of the % that the measured value is “off” from accepted value Measured value = Accepted value =

Formula is found in Table T (back page 12) of your Reference Tables:

If negative, your measured value is ________________ the accepted value If positive, your measured value is ________________ the accepted value

*It is very important that you put the given values into the proper place in theformula!

Sample Problem: In a lab experiment, you are told by your teacher that the actual (or accepted) amount of sugar in a can of Coke is 39 g. You experimentally determine it to be 40 g based on your own data and calculations. What is your percent error? Express answer in the proper amount of significant figures.