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Unit Vocabulary:
1.
2.
3.
4.
5.
6.
7.
8.
9.
S.I. unit
Meter
Liter
Gram
Mass
Weight
Volume
Density
Intensive
10.
11.
12.
13.
14.
15.
16.
17.
18.
Extensive
Significant Figures
Precision
Accuracy
Matter
Element
Compound
Mixture
Heterogeneous Mixture
19.
20.
21.
22.
23.
24.
25.
Homogeneous Mixture
Pure Substance
Particle Diagram
Chromatography
Filtration
Distillation
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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Matter and Measurement Warm-up Questions
Question
Answer
Date: _________
Date: _________
Date: _________
Date: _________
Date: _________
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Matter
Can NOT be separated
by physical means
Can NOT be separated
by chemical means
Particle
Diagram
CAN be Separated by
PHYSICAL means
Same
composition
throughout
Separated by chemical
means, only
Particle
Diagram
Particle
Diagram
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Different
composition
throughout
Particle
Diagram
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 of
one compound and one element?
4. Which particle diagram represents a diatomic element?
4
Properties of Matter:
 Physical properties are the constants about a substance; can use
our senses to observe them; do not require chemical analysis
Example:
o Extensive Property: a property that depends on how
much material you are dealing with
Ex:
o Intensive Property: a property that does not depend on
how much material you are dealing with (help identify
matter; 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;
_________________________ ; looks different but __________________ to
original state; _______________________ as a product
Example:
Chemical Change – a reaction in which the composition of a substance is
changed (ex: ____________); properties ____________________________
matter; ____________________ (in the form of light, fire, heat etc)
Example:
Example: firewood burning
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
1. Circle ( ) all the elements and underline the compounds below.
2. On the line provided, record the number of different symbols within the species to
the left.
CO
___
C2H5OH
Mg
___
H2SO4 ___
O2
___
C
___
___
Al(CN)3 ___
He
___
___
Cl2
___
NI3
H2O ___
Cu
Co
___
___
NaCl
___
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? __
Element =
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 subscript
applies 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 this
number by the number outside of the parentheses.
Example:
# iron atoms __
# sulfur atoms __
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# oxygen atoms ___
* The Common Elements *
Rules for writing element symbols:
* Symbol *
Ag
Al
Ar
As
Au
B
Ba
Be
Br
C
Ca
Cl
Co
Cr
Cs
Cu
F
Fe
Fr
H
He
Hg
1)
2)
* Name *
silver
aluminum
argon
arsenic
gold
boron
barium
beryllium
bromine
carbon
calcium
chlorine
cobalt
chromium
cesium
copper
fluorine
iron
francium
hydrogen
helium
mercury
* Symbol *
I
K
Kr
Li
Mg
Mn
N
Na
Ne
Ni
O
P
Pb
Ra
Rb
Rn
S
Si
Sn
Sr
U
Xe
Zn
* Name *
iodine
potassium
krypton
lithium
magnesium
manganese
nitrogen
sodium
neon
nickel
oxygen
phosphorus
lead
radium
rubidium
radon
sulfur
silicon
tin
strontium
uranium
xenon
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)
Filtration
Watch Glass Evaporation
Crucible Evaporation
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Description of
Technique
What types of
matter will it
separate?
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 or
small 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 number
out 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 would
get 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 =
If you need to plug these values into your calculator at any time, follow
these steps using this value  2.3 x 10-5
1.
2.
3.
4.
5.
Type “2”
Type the decimal point
Type “3”
Then press the “ee” “EXP” or “
” key(s)
Press the “+/-“ key (NOT the “—“ or “subtract” key)
6. Type “5”
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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
tera
T
giga
G
mega
M
kilo
k
hecto
h
deca
da
no prefix:
deci
d
centi
c
milli
m
micro

nano
n
pico
p
femto
f
atto
a
Numerical (Multiply Root Word
by)*
1,000,000,000,000
1,000,000,000
1,000,000
1,000
100
10
1
0.1
0.01
0.001
0.000001
0.000000001
0.000000000001
0.000000000000001
0.000000000000000001
Exponential
1012
109
106
103
102
101
100
10¯1
10¯2
10¯3
10¯6
10¯9
10¯12
10¯15
10¯18
*Example: In the word kilometer, the root word (base unit) is “meter” and
the 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 that
measure the same type of quantity
Examples:
-
*Rest Assured! For the Regents, the most you will have to convert will be between the
milli-/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.
9. 12 mL = ______ L
4. 3 g = _______ kg
5. 1 km = ______ m
Compare by placing a <, >, or = on
the line provided:
6. 1 kg = _______ g
10.
56 cm
11.
7g
7. 1 L = ________ mL
8. 7 m = _______ mm
__ 6 m
__ 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.
Example
1. Write the term to be
converted (include both the
number and the unit)
Convert 6.0 cm to km
2. Write the conversion
formula (see Ref Tables)
3. Make a fraction of the
conversion formula such that
the denominator units are the
same as the units from step 1
and the numerator contains the
units you want to convert to.
4. Multiply the term is step 1
by the fraction in step 3.
5. Cancel out “like” units
6. Solve (everything on top of
fraction is multiplied and
divided by everything on
bottom)
Now you try one…How many minutes are there in the month of
October?
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ACCURACY VS. PRECISION
Accuracy –
Precision –
Cheryl, Cynthia, Carmen, and Casey shot the targets above at camp. Match each
target with the proper description (assume bulls eye is the desired result)
(a) Accurate and precise
(b) Accurate but not precise
(c) Precise but not accurate
(d) Neither precise nor accurate
The following data was collected during a lab experiment. The density of the cube
should be 10.8 g/mL. Would you say that this data is accurate, precise, neither, or
both relative to this accepted value? ___________________________
Trial Number
1
2
3
Density of Cube
6.2 g/mL
6.3 g/mL
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
abilities to measure ____________; significant figures are the _________
__________________ to the measurement
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)
2)
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 stands
for the Pacific coast and A stands for the Atlantic Coast.
Imagine the number you are looking at is a map of the USA. Begin counting
from the correct side of the number (Atlantic/right side or Pacific/left
side) based on what you determined in step 1. Consider the first nonzero
number you land on the start of your count. Consider each digit from here
on out significant as well until you reach the other end of the number.
Pacific Coast
3.
Atlantic Coast
Decimal is
Present
Decimal is
Absent
1. Start @ 1st
NONZERO
1. Start @ 1st
NONZERO
2. Count all
the way to the
Atlantic—NO
EXCEPTIONS
2. Count all
the way to the
Pacific—NO
EXCEPTIONS
Determine the number of significant numbers in each of the following:
1) 357
_______
5) 0.0357
2) 3570
_______
6) 3.570 x 103
_______
3) 3570.
_______
7) 0.3570
_______
4) 0.357
_______
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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 123456789 m
significant
1.23 x 102
2. Zeros located between nonzero numbers
are significant
40.7 L
87,009 km
3. For numbers less than one, all zeros to
the left of the 1st nonzero number are
NOT significant
0.009587 m
4. Zeros at the end of a number and to the
right of a decimal point are significant
85.00 g
5. Zeros at the end of a whole number may
be 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
6. Exact numbers have an infinite
number of significant figures
2000 m
PRACTICE:
Measurement
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
0.0009 kg
9.070000000 L
2000. m
1 ft = 12 inch
Number of Significant Figures
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Rule(s) Applied
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
Multiplication/Division
Perform operation as
normal & express
answer in least # sig
figs that were given to
you
Addition/Subtraction
Examples:
Examples
Line decimal points up;
round final answer to
lowest decimal place
(least accurate) value
given
12.257 x 1.162 =
+
3.95
2.879
213.6____
5.1456 – 2.31 = _______
69.25/45.8 = _________
Rules for Calculations with Numbers in Scientific Notation:
Rule
Example
Addition/Subtraction  All values must 4.5 x 106 - 2.3 x 105
have the same exponent. Result is the
sum or difference of the mantissas,
multiplied by the same exponent of 10
Multiplication  mantissas are
multiplied and exponents of 10 are
(3.1 x 103) (5.01 x 104)
added
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
MASS
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.
____________________________________________________
____________________________________________________
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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 of
48.0 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 is
its volume?
4) What is the mass of an object that has a volume of 34 cm3 and a density of
6.0 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 the
formula!
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.
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