1. Chapter 1:
Matter – Its Properties and
Measurement

2. The Scientific Method
Observation Hypothesis Experiment Established
Theory
modify modify
Experiment Theory
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3. Properties of Matter
• Matter: occupies space and displays
mass and inertia
• Composition: relative proportions of the
components of a sample of matter
ex. water is 11.19% H and 88.81% O by mass
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4. Properties of Matter
• Physical property:
– a property that can be measured or observed without
changing the matter’s composition
• Chemical property:
– a property that comes with observing a change in
chemical composition
• Extensive property: depends on the quantity of
matter present
• Intensive property: does NOT depend on the
quantity of matter present
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5. Elements and Compounds
• Element:
– cannot be decomposed into a simpler substance
through chemical processes; distinguished by the
unit of the atom
• Compound:
– a substance made from the atoms of two or more
elements bonded chemically in defined proportions
• Compounds can only be decomposed into
their respective elements via chemical
processes
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6. Pure Substances and Mixtures
• A pure substance
– A substance with a fixed and uniform composition and distinct
properties (ex: pure water)
• A mixture:
– A combination of two or more pure substances which can
vary in composition and properties
a) homogeneous: ex: salt water
b) heterogeneous: ex: oil and water
• It is possible to separate mixtures through physical
porcesses
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7. Pure Substances and Mixtures
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8. Measuring Matter
an observed measurement not followed by a unit is meaningless!
The seven base SI units are:
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9. Who cares about units anyway?
• Mars Climate Orbiter
• probe sent by NASA to Mars to
study its weather
• the $168 million probe was
destroyed in 1999 after entering
the Martian atmosphere
• desired altitude: 140-150 km
• altitude attained: 57 km
• investigation revealed that the
on board computer used SI
units, while the computers on
Earth were using BE units
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10. SI Prefixes
Value
Prefix
Symbol
1012
tera-
T
109
giga-
G
106
mega-
M
103
kilo-
k
102
hecto-
h
101
deca-
da
10-1
deci-
d
10-2
centi-
c
10-3
milli-
m
10-6
micro-
µ
10-9
nano-
n
10-12
pico-
p
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11. Mass versus Weight
• mass:
– measures the quantity of matter in an object
• weight:
– the force of gravity on an object
The kilogram (kg) is the official SI unit, but we will most often
use the gram (g):
1 kg = 1000 g
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12. Volume
• volume: the size of a
cube (i.e., m3)
• we will most often use
the litre (L) for
measuring volumes
1000 mL = 1 L
1000 L = 1 m3
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13. Temperature
• the SI unit is the kelvin (K)
• absolute zero temperature is 0 K or -273.15oC
• the freezing point of water is 273.15 K or 0oC
• the boiling point of water is 373.15 K or 100oC
always use the temperature in K in your calculations!
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14. Accuracy and Precision
• Accuracy:
– indicates how close a measured value is to
the actual (or accepted) value
• Precision:
– indicates the degree of reproducibility of a
measured quantity
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15. Accuracy and Precision
accurate not accurate,
and precise but precise
not precise, neither accurate
but accurate nor precise
• accurate measurements are usually precise, but a systematic
error will produce values which are precise but not accurate
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16. Scientific measurements
• Scientific notation: N x 10n
6.022 045 x 1023 instead of 602 204 500 000 000 000 000 000
N=6.022 045 and n=23
• Significant figures
– digits considered to be significant in the calculation or
measurement of a quantity
this balance is precise to ±0.01 kg
an object that has a mass of 6.732 kg
will give a measurement of 6.73 ± 0.01 kg
__________
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17. Rules for sig figs…
• all non zero digits are significant
4
6.732 kg has __ significant figures
• zeros between two sig figs are also significant
5
6.0061 kg has __ significant figures
• zeros to the left of a sig fig are not significant
3
0.0502 kg has __ significant figures
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18. Rules for sig figs…
• if the value is greater than 1, all zeros to the right of the
decimal point are significant
4
6.000 kg has __ significant figures
• when converting to scientific notation, it may sometimes
be ambiguous whether hanging zeros are significant or
not
4500 kg could be 4.5 x 103, 4.50 x 103, or 4.500 x 103 kg
therefore 4500 kg could have 2, 3, or 4 sig figs!
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19. Rules for sig figs
• a whole number with perfect precision has an infinite
number of significant figures
if we determine the average of 3 trials, we can assume it s
3.000 000 000 … trials
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20. Rules for sig figs…
• addition/subtraction:
– the answer must have the same number of sig figs after
the decimal as the element of the calculation with the
least number of sig figs after the decimal point
+ 0.2225
+ 2.73 + 2.06
+ 0.321 ! 1.1
3.27
rounded to ______ 1.0
rounded to ______
+ 3.2735 + 0.96
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21. Rules for sig figs…
• multiplication/division:
– the answer must have the same number of sig figs
as the element of the calculation with the least
number of sig figs
2.2 x 3.7845 = 8.32590 8.3
rounded to ______
3.76 / 4.236 = 0.8876298 0.888
rounded to ______
(2.27 x 7.324) / 3.3 = 5.0380 5.0
rounded to ______
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22. Rules for sig figs…
• Logarithms
– the answer must have the same number of sig figs as the log
element
log(957) = 2.980911... = 2.98 ??
= log(9.57 x 102)
= log(9.57) + log(102)
= 0.980911... + 2.00000... = 2.981
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23. In summary…
• on tests and the final exam,
• on homework assignments,
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24. Conversion Factors
• to convert a quantity from one unit to another,
we need to use a conversion factor
Dimensional Analysis
Quantity with Quantity with Conversion
desired unit = given unit X factor
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25. Example 1: Conversion factors
Convert 345.3 cm into metres.
Solution
100 cm = 1 m
1m
? m = 345.3 cm x = 3.453 m
100 cm
• N.B. the number of sig figs in the conversion
factor is infinite
!
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26. Example 2: Conversion factors
The density of the lightest metal, lithium (Li) is 5.34 x 102 kg/m3.
Convert this value to g/cm3.
Solution
1000 g = 1 kg 100 cm = 1 m
3 2 kg 1000 g 1 m %3
? g/cm = 5.34 x 10 3
• •$ ' =
m 1 kg # 100 cm
2kg 1000 g 1 m3
5.34 x 10 • • 6 = 0.534 g/cm3
m3 1 kg 10 cm3
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27. Density
mass
• density = ρ = volume
• density is a intensive property and is a very useful
conversion factor
!
• the SI unit is kg/m3, but we will most often use g/cm3
for solids and liquids and g/L for gases
1 g/cm3 = 1 g/mL = 1000 kg/m3
1 g/L = 0.001 g/mL
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28. Example : Using density
A piece of platinum has a density of 21.5 g/cm3 and a
volume of 4.49 cm3. What is its mass?
Solution
m
= ! m = •V
V
3 21.5 g Pt
? g Pt = 4.49 cm Pt • 3
= 96.5 g Pt
cm Pt
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29. Percent Composition
• number of parts of a component in 100
parts of the whole
– ex. 10% means 10 parts x per 100 parts of
the whole
• IMPORTANT: must be defined by a unit!
– ex. a rock contains 3.5% gold by mass
means 3.5 g of gold per 100 g of rock
– ex. a bottle of wine contains 10.7% alcohol
by volume means 10.7 mL of alcohol per
100 mL of wine
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30. Percent Composition
• when expressed as a conversion factor,
the numerator and denominator must
have the SAME UNITS
– ex. a rock contains 3.5% gold by mass
3.5 g gold 3.5 kg gold 3.5 oz gold
= =
100 g rock 100 kg rock 100 oz rock
– ex. a bottle of wine contains 10.7% alcohol
by volume
!
10.7 mL EtOH 10.7 L EtOH 10.7 tbsp EtOH
= =
100 mL wine 100 L wine 100 tbsp wine
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31. Example : Using percent composition
A solution of sucrose in water is 28.0% sucrose by mass
and has a density of 1.118 g/mL. What mass of sucrose
(in grams) is in 3.50 L of this solution?
Solution
1000 mL 1.118 g solution 28.0 g sucrose
? g sucrose = 3.50 L solution • • •
1L mL solution 100 g solution
= 1.10x10 3 g sucrose
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32. Chapter 1: Key Concepts
1. the forms and properties of matter
2. SI units and prefixes
3. accuracy vs. precision
4. significant figures
5. scientific notation
6. conversion factors
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33. Chapter 1: Suggested Problems
13, 14, 19, 20, 23,
27, 31, 45, 47, 51,
63, 65, 80, 81, 89
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