Lecture Presentation
Chapter 1
Introduction:
Matter and
Measurement
Based on Power Point Presentation by
James F. Kirby
Quinnipiac University
Hamden, CT
© 2015 Pearson Education, Inc.
What is Chemistry?
•  the study of the properties and behavior of matter.
•  Central to our fundamental understanding of many
science-related fields.
Matter
And
Measurement
© 2015 Pearson Education, Inc.
What is Matter?
•  In general, anything that has mass and takes up space.
•  Chemical matter is composed of atoms and molecules.
Matter
And
Measurement
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Atoms: Building Blocks of Chemical Matter
•  Atoms are incredible tiny
(~10–10 m in diameter)
•  Element: simplest form
of chemical matter. Each
element corresponds to a
unique type of atom.
Note: Balls of different colors
represent atoms of different
elements. Individual atoms can
bond together to form larger
structures called molecules.
•  Compound: substance
composed of two or more
different of elements (i.e.
two or more atom types).
Matter
And
Measurement
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Classification of Matter According to State
Three states of chemical
matter:
Ø  Solid (e.g. ice)
Ø  Liquid (e.g. water in
its standard state)
Ø  Gas (e.g. water
vapor)
Click here for a YouTube tutorial
video on the 3 states of matter.
Be able to relate the observed
properties of each state to its
respective molecular-level model.
© 2015 Pearson Education, Inc.
Matter
And
Measurement
Classifying Matter Based on Composition
Follow this flow
chart to classify
any sample of
matter as one of
the following:
Ø  Element
Ø  Compound
Ø  Homogeneous
mixture
Ø  Heterogeneous
mixture
Matter
And
Measurement
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Classification of Matter—Pure Substances
Pure Substance: has distinct properties, and its
composition does not vary from sample to
sample. The two types of pure substances are
elements and compounds.
Ø Element: can’t be decomposed to simpler
substances.
Ø Compound: can be decomposed to simpler
substances. (Composed of two or more
elements.)
Matter
And
Measurement
© 2015 Pearson Education, Inc.
Compounds and Composition
•  Law of Constant Composition (or Law of Definite Proportions): A
given compound is always composed of the same two or more
elements, which are always present in the same ratio by mass. (e.g.
water always contains 8 grams of oxygen per gram of hydrogen.)
•  According to atomic theory, this is because the relative number of
atoms of each element that makes up the compound is the same in any
sample of the compound. (e.g. water is always composed of two H
atoms for every one O atom, and an O atom is 16 times more massive
than an H atom.)
Matter
And
Measurement
© 2015 Pearson Education, Inc.
Mixture
•  Physical blend of two or more pure
substances; components can usually be
separated by physical means.
•  Exhibits the properties of the substances that
make it up.
•  Heterogeneous Mixture: varies in
composition throughout a sample
•  Homogeneous Mixture (or Solution): has
the same composition throughout; thoroughly
mixed down to the atomic/molecular level.
Matter
And
Measurement
© 2015 Pearson Education, Inc.
Types of Properties
•  Physical Property: observed without changing a
substance into another substance.
◦  Examples: boiling point, density, mass, and volume.
•  Chemical Properties: observed as a consequence
of one substance changing into another substance.
◦  Examples: flammability, corrosiveness, and
reactivity with acid.
Matter
And
Measurement
© 2015 Pearson Education, Inc.
Types of Properties:
Another Way to Classify Them
Extensive Properties: depend on the amount of
the substance present.
Ø Examples: mass, volume, or energy.
Intensive Properties: independent of the amount
of the substance that is present.
◦  Examples: density, boiling point, and color.
◦  Note: any property that is a ratio of two extensive
properties (e.g. density = mass/volume) is an
intensive property.
Matter
And
Measurement
© 2015 Pearson Education, Inc.
Types of Changes
•  Physical Changes: changes in matter that do
not change the composition of a substance.
◦  Examples: changes of state, temperature, and
volume.
•  Chemical Changes: result in new substances.
◦  Examples: combustion, oxidation, and
decomposition.
Matter
And
Measurement
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Physical Changes in State of Matter
Ø Alter the physical form
but not the type of
matter present
Ø Examples include
changes of state
(e.g. when ice melts or
water evaporates,
there are still 2 H
atoms and 1 O atom in
each molecule.)
Matter
And
Measurement
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Chemical Reactions (Chemical Changes)
Ø  The starting substances
(reactants) are converted to
new substances (products)
with different properties.
Ø  Here, the elements
hydrogen (H2) and oxygen
(O2) become the compound
water (H2O).
Click here to view a video that demonstrates the
differences between chemical and physical
properties and changes. Also covered are the
topics of elements, compounds, and mixtures.
Matter
And
Measurement
© 2015 Pearson Education, Inc.
Separating Mixtures
Mixtures can be separated based on
physical properties of the components of the
mixture. Some methods used to separate
the components of mixtures are
Ø Filtration
Ø Distillation
Ø Chromatography
Matter
And
Measurement
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Filtration
•  Solid substances
are separated
from liquids and
solutions
•  Liquid portion
passes through
pores in the filter
paper; while solid
particles cannot.
Matter
And
Measurement
© 2015 Pearson Education, Inc.
Distillation
•  Uses differences in the
boiling points of
substances to separate
a homogeneous mixture
into its components.
•  Relative to the boiling
liquid mixture, the vapor
is more concentrated in
the more volatile
component
•  The vapor is then
condensed and
collected in a “receiving
flask.”
Matter
And
Measurement
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Chromatography
•  A mixture is dissolved in
a solvent (mobile phase)
that carries it through the
chromatography column.
•  Substances separate
based on differences in
their abilities to adhere to
the solid surface
(stationary phase)
•  In this case, “substance
a” adheres most strongly
to the stationary phase
and, thus, comes out
last. “Substance c”
adheres least strongly
and comes out first. Matter
And
Measurement
© 2015 Pearson Education, Inc.
Numbers and Chemistry
•  Chemistry is very quantitative, as it often
involves things with numerical values (i.e.
experimental measurements).
•  Important Numerical Concepts in Science
Ø Scientific Notation
Ø Units of measurement
Ø Quantities that are measured and calculated
Ø Uncertainty in measurement
Ø Significant figures
Ø Dimensional analysis
Matter
And
Measurement
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Units of Measurement: SI Units
•  Système International d’Unités (International System of Units)
•  Each fundamental quantity has its own base (see above table).
•  Derived SI units are constructed from the appropriate base units
(e.g. the SI unit for velocity is meters per second, m/s).
Matter
And
Measurement
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Scientific Notation
Convenient way of expressing both the very large and very
small quantities often encountered in science
o  Example: There are 6 x 1023 atoms in 1 gram of hydrogen.
o  Example: The mass of a proton is 1.67 x 10–27 kg.
Click here for a YouTube video
tutorial on Scientific Notation.
This 14 minute video gives lots
of examples.
Matter
And
Measurement
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Metric Prefixes: convert base units into
units that are more common and/or convenient
Matter
And
Measurement
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Mass and Length
•  Two of the most basic units commonly
measured in science.
•  Mass: measure of the amount of matter in
an object. SI unit is the kilogram (kg).
•  Length: measure of distance. SI is the
meter (m).
Matter
And
Measurement
© 2015 Pearson Education, Inc.
Volume: Example of a Derived Unit
•  Derived from a base/fundamental unit:
length
•  The SI unit of volume is the cubic meter:
m × m × m = m 3.
•  More commonly used metric units of
volume: liter (L) and milliliter (mL).
Ø A liter is a cube 1 decimeter (dm) long
on each side. (1 L = 1 dm3)
Ø A milliliter is a cube 1 centimeter (cm)
long on each side, also called 1 cubic
centimeter: cm × cm × cm = cm3.
(1 mL = 1 cm3)
Matter
And
Measurement
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Temperature
v  Generally related the “hotness and coldness” of an object.
v  Determines the direction of heat flow. Heat always flows
spontaneously from a higher temperature to a lower temperature.
v  Higher temperatures correspond to more thermal energy (random
atomic/molecular motion) in a substance.
v  In scientific measurements, the Celsius and Kelvin scales are used
most often used.
Matter
And
Measurement
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Temperature Units Used in Science
Celsius scale: based on the properties of water.
Ø  0 ˚C is the freezing point of water.
Ø  100 ˚C is the boiling point of water.
Kelvin scale: the “absolute” temperature
Ø  The SI unit of temperature.
Ø  The Kelvin unit has the same size as the Celsius unit, but the
two scales have different zero points
Ø  0 K is the lowest possible temperature (“absolute zero”). This
is the point at which all thermal energy would have been
drained from an object; so the temperature cannot get any
lower.
Ø  0 K = –273.15 ˚C
Conversion Between Scales (given on AP Exam)
Matter
K = ˚C + 273.15
And
Measurement
© 2015 Pearson Education, Inc.
Density
o  The mass to volume ratio of a substance:
D = m/V
o  a physical property, with units derived from units for
mass and volume.
•  The most common density unit for liquids and solids is
g/mL (Note: 1 g/ml = 1 g/cm3) For example, the
density of aluminum at 20 ˚C is 2.7 g/cm3.
•  A common density unit for gases is g/L. For example,
the density of air at sea level and 15 ˚C is 1.2 g/L.
•  The SI unit of density is kg/m3, but it is not commonly
used. For example, the density of aluminum at 20 ˚C
is 2.7 x 103 kg/m3.
Matter
And
Measurement
© 2015 Pearson Education, Inc.
Types Numbers Encountered in Science
Exact Numbers
Ø  Either counted explicitly (e.g. there are 16 students in a
classroom) or given by definition (e.g. there are 12 eggs in
1 dozen)
Ø  There is no uncertainty in their values.
Inexact (Measured) Numbers:
Ø  Involve a certain level of estimation, which is dependent
upon the circumstances of each measurement
Ø  Uncertainty is inherent in any measurement, since both
scientists and scientific instruments have limitations. (For
example, some scientists have better eyesight than
others. Also, some balances measure to ±0.01 g; others
measure to ±0.0001g.)
Matter
And
Measurement
© 2015 Pearson Education, Inc.
Accuracy versus Precision
Accuracy
²  refers to the closeness of a measurement
to the true value of a quantity.
²  Based on comparing the measured value
to the true/accepted value for the quantity.
Precision
²  refers to the proximity of several
measurements to each other.
²  Based on comparing the measured values
to each other. (The true value is not taken
into account when determining precision.)
Click here to view a tutorial
video on the difference between
accuracy and precision.
© 2015 Pearson Education, Inc.
Matter
And
Measurement
Significant Figures
•  Each measured or calculated value has a certain
number of significant figures. These are the
digits in which there is a good level of confidence.
•  Measuring the smallest significant digit involves
a certain level of estimation. This determines the
precision (i.e. level of reproducibility) of the
measured value.
Matter
And
Measurement
© 2015 Pearson Education, Inc.
Uncertainty in Measurements
•  Each measurement involves a certain degree of inherent random error,
which limits the precision (reproducibility) of the measured value.
•  The amount of precision (i.e. number of sig figs) in a measurement is
largely determined by the instrument used. For instance, the graduated
cylinder shown below measures volume to a precision of ± 1 mL, while
the burette is precise to ± 0.1 mL.
Click here to view a
video tutorial on
determining the
numbers of
significant figures in
measured values.
Matter
And
Measurement
© 2015 Pearson Education, Inc.
Rules for Determining the Number of
Significant Figures in a Numerical Value
1.  All nonzero digits are significant.
2.  Zeroes between two significant figures are
themselves significant.
3.  Zeroes at the beginning of a number are never
significant.
4.  Zeroes at the end of a number are significant if
a decimal point is written in the number.
Matter
And
Measurement
© 2015 Pearson Education, Inc.
Rules for Determining the Number of
Significant Figures in a Calculated Value
Note: the precision of a calculated value is limited by the precision of
the numbers used in the calculation. We must round the calculated
value to the correct number of significant figures so as to not
overstate the precision of the answer.
•  For addition or subtraction, answers are rounded to the least
significant decimal place.
•  For multiplication or division, answers are rounded to the number
of digits that corresponds to the least number of significant
figures in any of the numbers used in the calculation.
Click here to view a video with
a comprehensive review of how
to determine the numbers of sig
figs in both given numbers and
calculated values.
© 2015 Pearson Education, Inc.
Matter
And
Measurement
Dimensional Analysis
•  Used to convert one quantity to another.
•  Often utilizes conversion factors between equivalent quantities
(e.g., 1 in. = 2.54 cm exactly, by definition).
•  For instance, the equivalence between inches and centimeters can
be used to convert from cm to in. (1 in/2.54 cm) or from inches to
cm (2.54 cm/1 in.) We use the ratio which allows us to change
units (puts the units we have in the denominator to cancel).
Click here to view a video tutorial
involving a simple one-step unit
conversion of a measured value.
© 2015 Pearson Education, Inc.
Matter
And
Measurement
Additional Video Tutorials from YouTube
Click here to view a tutorial video that
includes two multi-step conversion problems.
Click here to view a video covering a
numerical word problem involving a density
calculation.
Click here to view a video covering a
numerical word problem to determine the
diameter of a red blood cell.
Matter
And
Measurement
© 2015 Pearson Education, Inc.