Chapter 1: Chemistry and
Measurement
Renee Y. Becker
Valencia Community College
CHM 1045
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Properties of Matter
• Chemistry: The study of composition, properties,
and transformations of matter
• Matter: Anything that has both mass & volume
• Hypothesis: Interpretation of results
• Theory: Consistent explanation of observations
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Conservation of Mass
• Law of Mass Conservation: Mass is
neither created nor destroyed in chemical
reactions.
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Example 1: Conservation of Mass
C(s) + O2(g)  CO2(g)
a) 12.3g C reacts with 32.8g O2, ?g CO2
b) 0.238g C reacts with ?g O2 to make .873g CO2
c) ?g C reacts with 1.63g O2 to make 2.24g CO2
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Dalton’s Atomic Theory
• Law of Definite Proportions: Different samples
of a pure chemical substance always contain the
same proportion of elements by mass.
– Any sample of H2O contains 2 hydrogen atoms for
every oxygen atom
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Matter
• Matter is any substance that has mass and
occupies volume.
• Matter exists in one of three physical states:
– solid
– liquid
– gas
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Gases
• In a gas, the particles of matter are far apart and
uniformly distributed throughout the container.
• Gases have an indefinite shape and assume the shape
of their container.
• Gases can be compressed and have an indefinite
volume.
• Gases have the most energy of the three states of
matter.
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Liquid
• In a liquid, the particles of matter are loosely packed
and are free to move past one another.
• Liquids have an indefinite shape and assume the
shape of their container.
• Liquids cannot be compressed and have a definite
volume.
• Liquids have less energy than gases but more energy
than solids.
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Solid
• In a solid, the particles of matter are tightly packed
together.
• Solids have a definite, fixed shape.
• Solids cannot be compressed and have a definite
volume.
• Solids have the least energy of the three states of
matter.
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Phases
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Changes in Physical State
• Most substances can exist as either a solid, liquid, or
gas.
• Water exists as a solid below 0 °C; as a liquid
between 0 °C and 100 °C; and as a gas above 100°C.
• A substance can change physical states as the
temperature changes.
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Solid  Liquid
• When a solid changes to a liquid, the phase change is
called melting.
– A substance melts as the temperature increases.
• When a liquid changes to a solid, the phase change is
called freezing.
– A substance freezes as the temperature decreases.
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Liquid  Gas
• When a liquid changes to a
gas, the phase change is called
vaporization.
• A substance vaporizes as the
temperature increases.
• When a gas changes to a
liquid, the phase change is
called condensation.
• A substance condenses as the
temperature decreases.
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Solid  Gas
When a solid changes directly to a gas,
the phase change is called sublimation.
A substance sublimes as the temperature
increases.
When a gas changes
directly to a solid, the phase
change is called deposition.
A substance undergoes
deposition as the
temperature decreases.
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15
Classifications of Matter
• Matter can be divided into two classes:
– mixtures
– pure substances
• Mixtures are composed of more than one substance
and can be physically separated into its component
substances.
• Pure substances are composed of only one substance
and cannot be physically separated.
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Mixtures
• There are two types of mixtures:
– homogeneous mixtures
– heterogeneous mixtures
• Homogeneous mixtures have uniform properties
throughout.
– Salt water is a homogeneous mixture.
• Heterogeneous mixtures do not have uniform
properties throughout.
– Sand and water is a heterogeneous mixture.
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Pure Substances
• There are two types of pure substances:
– Compounds
– Elements
• A compound is a substance composed of two or
more elements chemically combined
– Compounds can be chemically separated into individual
elements.
– Water is a compound that can be separated into hydrogen
and oxygen.
• An element cannot be broken down further by
chemical reactions.
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19
Properties of Matter
• Properties: describe or identify matter
• Intensive Properties do not depend on amount
– temperature, boiling point, melting point
• Extensive Properties do depend on amount.
– length and volume
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Properties of Matter
• Physical Properties can be determined without
changing the chemical makeup of the sample.
• Some typical physical properties are:
– Melting Point, Boiling Point, Density, Mass, Touch, Taste,
Temperature, Size, Color, Hardness, Conductivity.
• Some typical physical changes are:
– Melting, Freezing, Boiling, Condensation, Evaporation,
Dissolving, Stretching, Bending, Breaking.
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Properties of Matter
• Chemical Properties are those that do change
the chemical makeup of the sample.
• Some typical chemical properties are:
– Burning, Cooking, Rusting, Color change, Souring of milk,
Ripening of fruit, Browning of apples, Taking a photograph,
Digesting food.
• Note: Chemical properties are actually chemical changes
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Properties of Matter
PHYSICAL
CHANGE
PROPERTIES
New form of old
substance.
No new substances
formed.
CHEMICAL
Old substance
destroyed.
New substance
formed.
Description by senses – List of chemical
shape, color, odor, etc.
changes possible.
Measurable properties –
density, boiling point,
etc.
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Example 2: Matter
Which of the following represents a mixture?
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Accuracy, Precision, and Significant Figures in Measurement
• Accuracy is how close
to the true value a
given measurement is.
• Precision is how well a
number of independent
measurements agree
with one another.
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Scientific Notation
• Changing numbers into scientific notation
– Large # to small #
– Moving decimal place to left, positive exponent
123,987 = 1.23987 x 105
– Small # to large #
– Moving decimal place to right, negative
exponent
0.000239 = 2.39 x 10-4
How to put into calculator
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Example 3: Scientific Notation
Put into or take out of scientific notation
a) 87542
b) 2.1956 x 10-3
c) 0.784
d) 2.78 x 106
e) 92000
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Accuracy, Precision, and Significant Figures in Measurement
• Significant Figures are the total number of digits
in the measurement.
• The results of calculations are only as reliable as
the least precise measurement!!
• Rules exist to govern the use of significant figures
after the measurements have been made.
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Accuracy, Precision, and Significant Figures in Measurement
• Rules for Significant Figures:
– Zeros in the middle of a number are significant
– Zeros at the beginning of a number are not
significant
– Zeros at the end of a number and following a
period are significant
– Zeros at the end of a number and before a period
may or may not be significant.
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Example 4: Significant Figures
How many Sig. Figs ?
a) 0.000459
b) 12.36
c) 36,450
d) 8.005
e) 28.050
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Accuracy, Precision, and Significant Figures in Measurement
• Rules for Calculating Numbers:
– During multiplication or division, the answer
can’t have more sig figs than any of the original
numbers.
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Example 5: Significant Figures
a) 238.5 x 79 =
b) 12 / 0.1272 =
c) 0.2895 x 0.29 =
d) 32.567 / 22.98 =
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Accuracy, Precision, and Significant Figures in Measurement
-During addition or subtraction, the answer can’t
have more digits to the right of the decimal point than
any of the original numbers.
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Example 6: Significant Figures
a) 238.5 + 79 =
b) 12.3 - 0.1272 =
c) 0.2895 + 0.29 =
d) 32.567 - 22.98 =
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Accuracy, Precision, and Significant Figures in Measurement
• Rules for Rounding Numbers:
– If the first digit removed is less than 5
• round down (leave # same)
– If the first digit removed is 5 or greater
• round up
– Only final answers are rounded off, do not round
intermediate calculations
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Example 7: Rounding and Significant Figures
Round off each of the following measurements
(a) 3.774499 L to four significant figures
(b) 255.0974 K to three significant figures
(c) 55.265 kg to four significant figures
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Example 8: Accuracy & Precision
• Which of the following is precise but not
accurate?
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Measurement and Units
SI Units
Physical Quantity
Mass
Length
Temperature
Amount of substance
Time
Electric current
Luminous intensity
Name of Unit
kilogram
meter
kelvin
mole
second
ampere
candela
Abbreviation
kg
m
K
mol
s
A
cd
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Measurement and Units Some prefixes for multiples of SI units
*
*
*
*
*
*
*
Factor
1,000,000,000 = 109
1,000,000 = 106
1,000 = 103
100 = 102
10 = 101
0.1 = 10-1
0.01 = 10-2
0.001 = 10-3
0.000,001 = 10-6
0.000,000,001 = 10-9
0.000,000,000,001 = 10-12
* Important
Prefix
giga
mega
kilo
hecto
deka
deci
centi
milli
micro
nano
pico
Symbol
G
M
k
h
da
d
c
m
µ
n
p
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Measurement and Units
• Temperature
Conversions:
The Kelvin and Celsius
degree are essentially
the same because both
are one hundredth of the
interval between freezing
and boiling points of
water.
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Measurement and Units
• Temperature Conversions:
– Celsius (°C) — Kelvin temperature conversion:
Kelvin (K) = °C + 273.15
– Fahrenheit (°F) — Celsius (C) temperature
conversions:
C = 5/9 (F – 32)
F = (9/5 x C) + 32
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Example 9: Temp. Conversions
Carry out the indicated temperature conversions:
(a) –78°C = ? K
(b) 158°C = ? °F
(c) 375 K = ? °C
(d) 98.6°F = ? °C
(e) 98.6°F = ? K
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Measurement and Units
• Density: relates the mass of an object to its
volume.
Density = mass / Volume
D=m/V
V=m/D
m=VD
• Density decreases as a substance is heated because
the substance’s volume increases.
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Example 10: Density
What is the density of glass (in grams per cubic
centimeter) if a sample weighing 26.43 g has a
volume of 12.40 cm3?
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Example 11: Density
What is the volume of an unknown solution if the
mass is 12.567 g and the density is 14.621 g/mL ?
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Example 12: Density
What is the mass of an unknown solution if the mass
is 20.2 mL and the density is 2.613 g/mL ?
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Measurements and Units
• Dimensional-Analysis method uses a conversion
factor to express the relationship between units.
Original quantity x conversion factor = equivalent quantity
Example: express 2.50 kg  lb.
Conversion factor: 1.00 kg = 2.205 lb
2.50 kg x 2.205 lb = 6.00 lb
1.00 kg
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Measurements and Units
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Example 13: Conversions
a) 1.267 km  m  cm
b) .784 L  mL
c) 3.67 x 105 cm  in
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Example 14: Conversions
a) 79 oz  g
b) 9.63 x 10-3 yd  ft
c) 23.5 cm2  m2
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Example 15: Conversions
a) 1.34 x 1012 pm  m
b) 4.67 x 10-7 nm  pm
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