# What is matter and measurement? how are they related? Contributed by: Chemistry is the study of matter, its properties, composition, and structure, and the changes it undergoes.
1. Chapter 1
Introduction: Matter and Measurement
Learning outcomes:
 Distinguish between the macroscopic world and the submicroscopic realm of atoms
and molecules.
 Categorize matter into its different states.
 Differentiate among pure substances, homogeneous mixtures, and heterogeneous
mixtures.
 Categorize the properties of pure substances.
 Contrast the different techniques used to separate mixtures of substances.
 Differentiate the different forms of energy.
 Use the metric system to express scientific quantities quantitatively.
 Differentiate between metric base units and derived units, such as volume, density,
and energy.
 Determine the uncertainty of a given quantity in scientific calculations.
 Analyze calculations to determine the correct number of significant figures for the
 Apply dimensional analysis to track the conversion from given units to desired units.
• Chemistry is the study of matter, its properties,
composition, and structure and the changes it undergoes.
• It is central to our fundamental understanding of many
science-related fields.
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2. Atomic and Molecular Perspective
Matter – Anything that has
mass and occupies space.
Atom – The smallest stable
building block of matter.
Molecule – Groups of atoms
held together with a specific
connectivity and shape.
Composition - the types of atoms that are present in a
compound and the ratio of these atoms (for example H2O, C2H6O).
Structure - how atoms are connected (bonded) to each other,
how far apart they are, and the shape of the molecule.
Methods of Classification of Matter
State of Matter - physical state is gas, liquid, or solid.
Composition of Matter - element, compound, or mixture
States of Matter
1) Gas (vapor) – has no fixed
volume or shape, uniformly
expands to fill its container,
diffusion occurs rapidly.
2) Liquid - has a distinct
volume independent of its
container, assumes the shape
of the portion of the
container it occupies, not
significantly compressible,
diffusion occurs but slower
than a gas.
3) Solid - has both a definite
shape and definite volume,
not significantly
compressible, diffusion
occurs extremely slowly.
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3. Elements, Compounds & Mixtures
Pure Substance Matter that has a fixed composition and distinct properties. All
pure substances are either elements or compounds.
Elements All atoms are the same kind, elements have only one type of atom. e.g.
oxygen (O2), gold (Au), silicon (Si) and diamond (C).
Compounds Contains more than one type of atom, but all molecules (or repeat
units) are the same, e.g. water (H2O), ethanol (C2H6O), quartz (SiO2), sodium
chloride (NaCl).
Mixture Have variable composition and can be separated into component parts by
physical methods. Mixtures contain more than one kind of molecule, and their
properties depend on the relative amount of each component present in the mixture.
Homogeneous & Heterogeneous Mixtures
The composition is variable for both
heterogeneous and homogeneous mixtures.
Heterogeneous Mixture - non-uniform.
Chocolate Chip Cookie – Chocolate, dough, etc.
Concrete – Cement, rocks, etc.
Nachos – Chips, cheese, jalapeños, salsa, etc.
Homogeneous Mixture – uniform throughout,
also called a solution.
Air – principle components include O2, N2 & CO2
Vodka – principle components are ethanol and water
Brass – solid solution of Cu and Zn
Ruby – solid solution of Al2O3 and Cr2O3
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4. Periodic Table Relative abundances of
elements in the Earth’s
crust and human body.
Elements are represented as symbols with one
or two letters; the first is always capitalized.
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5. Compound Elements
Elements can interact with other elements to form compounds,
and compounds can be decomposed into elements.
Law of Constant Composition
The elemental composition of a compound is always
the same, which is known as the Law of Constant
Composition (or Law of Definite Proportions).
Example: According to the law of definite proportions,
if a sample of compound A contains 12 g of sulfur and
6 g of oxygen, another sample of A that contains 50
grams of sulfur must contain ________ g oxygen.
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6. Chemical and Physical Properties
Physical Properties Some properties can be readily measured
with our senses, e.g. odor and color, instruments are needed to
measure other properties, such as electrical resistivity, hardness,
melting point, boiling point, density, mass, volume, etc.
Chemical Properties Describe the reactivity of a substance
toward other substances. Examples include:
Ethanol burns in air (reacts with oxygen)
Sodium reacts vigorously with water,
Corrosion of metal parts (rust),
Trinitrotoluene (TNT) is explosive.
Physical changes are changes in matter that
do not change the composition of a substance.
– Examples include changes of state,
temperature, and volume.
Chemical changes result in new substances.
– Examples include reactions: combustion,
oxidation, tarnishing, and decomposition.
Properties of Matter
• Intensive Properties:
□ Independent of the amount of the
substance that is present.
• Density, temperature, melting point,
boiling point, hardness, color, etc.
• Extensive Properties:
□ Dependent upon the amount of the
substance present.
• Mass, volume, energy, etc.
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7. Separation of Mixtures
Filtration Distillation
Chromatography
Energy
• Energy is the capacity to
do work or transfer heat.
• Work is the energy
transferred when a force
exerted on an object
causes a displacement of
that object.
• Heat is the energy used to
cause the temperature of
an object to increase.
• Force is any push or pull
on an object.
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8. Fundamental Forms of Energy
Kinetic energy is
the energy of motion.
– Its magnitude
depends on the
object’s mass
and its velocity:
KE = ½mv2
• Potential energy of an object depends on its relative position
compared to other objects.
• Potential energy also refers to the composition of an object,
including the chemical energy stored in chemical bonds.
One of the goals in chemistry is to related the energy changes in
the macroscopic world to the kinetic or potential energy of
substances at the molecular level.
Numbers and Units in Chemistry
Major role in quantifying:
– Units of measurement
– Quantities that are measured and calculated
– Uncertainty in measurement
– Significant figures
– Dimensional analysis
(e.g. 1 inch = 2.54 cm)
Energy is a derived unit, mass times velocity squared
Joule (J) = kg·m2/s2
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9. Metric System Prefixes
Example
Example 1: How many grams are in 10 kg?
10 kg 1×103 g
× =10×103 g
1 kg
In scientific notation, 1.0x104 g.
Example 2: Use appropriate metric prefixes to write
the following measurement without use of exponents,
7.6×10−10 g.
Approach: Use the closest prefix to the exponent (-10),
choose 1 ng = 1×10−9 g.
7.6×10−10 g 1 ng
× = 0.76 ng
1×10−9 g
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10. Volume
Volume is not a base unit for
SI; it is a derived unit from
length (m × m × m = m3).
The most commonly used
metric units for volume are the
liter (L) and the milliliter (mL).
– A liter is a cube 1
decimeter (dm) long on
each side.
– A milliliter (mL) is a cube
1 centimeter (cm) long on
each side, also called 1
cubic centimeter
(cm × cm × cm = cm3).
Glassware for Measuring Volume
Uncertainty in Measurements - Different measuring devices
have different uses and different degrees of precision.
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11. Temperature Scales
Temperature – the “hotness and coldness” of an object.
Heat flows spontaneously from an object with a higher
temperature to an object with a lower temperature.
Temperature
• The Kelvin is the SI
unit of temperature.
• It is based on the
properties of gases.
• There are no
negative Kelvin
temperatures.
• K = C + 273.15
Example: Express 25.00 °C in K.
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12. Temperature
•The Fahrenheit scale is not
used in scientific
measurements.
F = 9/5(C) + 32
C = 5/9(F − 32)
The ‘9/5’, ‘5/9’, and ’32’ are
exact numbers and do not
influence significant figures.
Examples:
Density
mass m Example: A piece of unknown metal with a
Density = d = =
volume V right rectangular prism shape has a width
of 3.2 cm, a length of 17.1 cm, and height
of 4.0 cm. Its mass is 1.5 kg. Calculate the
density of the metal in g/cm3.
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13. Example
The world’s largest gold bar in the Toi Gold Museum in
Japan can be seen and touched by visitors. It has an
irregular shape with dimensions about 17.9 in. by 8.9 in.
by 6.7 inches, with a volume of 12.94 L. Determine the
mass and weight (1 kg = 2.205 lb).
Numbers in Chemistry
• Exact numbers are counted or given by definition. For
example, there are 12 eggs in 1 dozen and 3 feet in 1 yard.
• Inexact (or measured) numbers depend on how they were
determined. Scientific instruments have limitations
(equipment errors) and individuals can read some
instrumentation differently (human errors).
The last digit measured is considered reliable, but not exact.
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14. Precision and Accuracy
• Precision is a measure of how closely individual
measurements to agree with one another.
• Accuracy refers to how closely individual
measurements agree with the correct “true” value.
Significant Figures
• The term significant figures refers to digits
that were measured.
• When rounding calculated numbers, we pay
attention to significant figures so we do not
overstate the accuracy of our answers.
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15. Significant Figures
1. Zeros between non-zero numbers are always
significant.
2. Zeros at the beginning of a number are never
significant, merely indicate the position of the
decimal point.
3. Zeros at the end of the number after a decimal
place are significant if the number contains a
decimal point.
4. Zeros at the end of a number before a decimal place
are ambiguous (e.g. 23,800 g), unless a decimal
point is written at the end (i.e. 23,800. g). Assume
the zeros are insignificant, unless there is a decimal
point. Avoid ambiguity by using scientific notation.
Significant Figures
How many significant figures are present in
each of the measured quantities?
0.0012
108
900.0
3.0012
0.002070
4.80  10-3
4.800  10-3
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16. Rounding*
After determining the appropriate number of
1. If the first digit you drop is greater than 5, add 1
to the last digit you keep. You are rounding up.
2. If the first digit you drop is less than 5, do
nothing to the digits you keep. You are rounding
down.
3. If the digit you drop is 5, and there are no
following digits, round down. If there are digits
following the 5, round up.
Significant Figures & Calculations
Line up the numbers at the decimal point and the
answer cannot have more decimal places than the
measurement with the fewest number of decimal
places.
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• The absolute uncertainty can be no
smaller than the least accurate number.
• 12.04
- 10.4
1.64 1.6
• The answer should have no more decimal
places than the least accurate number.
Multiplication and Division
The answer cannot have more significant
figures than the measurement with the
fewest number of significant figures.
3121  12 = 37452 = 3.7  104
# sig. digits 4 2 2
Know the number of appropriate digits throughout, round at the end.
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18. Mixed Operations
Determine accuracy in the same order as the
mathematical operations, # of significant digits in blue
•Retain at least one additional digit past the significant
figures in combined operations, so rounding doesn’t
affect results…
-keep track of the proper significant figures to use at
the end.
3 3
m 2.79 g 2.79 g
d= = =
v 8.34 mL - 7.58 mL 0.76 mL
3 3 2
d = 3.7 g/mL
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Evaluate the expression to the correct number of
significant figures. How many sig. figs. in the answer?
4.184 × 100.620 × (25.27 - 24.16) = _____
Using order of operations, evaluate subtraction first.
Then multiply each number and evaluate.
4.184 × 100.620 × (1.11) = 467.3034288
(4 s.f.) (6 s.f.) (3 s.f.) 467
Multiplication uses fewest number of sig. figures.
The answer should have three significant figures.
Retain at least one additional digit past the significant figures in
combined operations, so rounding doesn’t affect result…
-keep track of the proper significant figures for the final answer.
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19. Evaluate the expression to the correct number of
significant figures. How many sig. figs. in the answer?
9.6 x 100.65
+ 4.026 =
8.321
Evaluate the expression to the correct number of
significant figures. How many sig. figs. in the answer?
320.75 - (6102.1 / 3.1) =
Evaluate the expression to the correct number of
significant figures. How many sig. figs. in the answer?
(14.20000  0.7288) + (12.00536  0.0201) =
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20. Dimensional Analysis
Units are multiplied together or divided into
each other along with the numerical values.
• Keep track of both numerical values and units.
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Conversions: Two or More Factors
What is the mass in g, of 1.00 gal of H2O? The density
of water is 1.00 g/mL. 1 L = 1.057 qt, 1 gal = 4 qt
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21. Conversions Involving Volume
Express a volume of 1.250 L in mL and cm3
1 mL
(1.250 L)× =1,250. mL
1 × 10 L
1000 mL
(1.250 L)× =1,250. mL
1L
1000 cm3
(1.250 L)× =1,250. cm3
1L
Express a volume of 1,250. cm3 in m3.
The prefix centi is 10-2, 1 cm = 110-2 m for length.
Volume involves cubed units, create a conversion:
110−2 m 110−2 m 110−2 m 110−6 m3
=
1 cm 1 cm 1 cm 1 cm3
Use the conversion to express the volume in m3:
110−6 m3
(1250. cm3 ) × = 1.250 × 10−3 m3
1 cm3
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