Matter – Properties and Changes
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Section 3.1 Properties of Matter
Chemistry: Study of matter and the changes it undergoes
Matter: anything that has mass and takes up space

Substance: matter with a uniform and unchanging composition; a pure substance
Ex. Water – always H2O
Nonexample: sea water – will have a different composition based upon where it
was collected
Extensive properties: depends upon
the amount of the substance
present
Physical Properties of Matter –
characteristics that can be observed or
measured without changing the
sample’s composition
Ex: mass, volume length
Examples:
Density
Color
Taste
Hardness
Boiling point
Intensive properties: independent of
the amount of the substance
present
Odor
Melting point
Ex: density
Chemical Properties of Matter –
Observing properties:
ability of a substance to combine with or
change into 1 or more other substances
The presence of certain properties
may vary depending on the
conditions of the environment

Also, if a substance does not react
with some other substances is also
a chemical property
States of matter:
physical state of a substance is a
physical property of that substance

State specific conditions in which
observations are made
Ex. Both chemical and physical
properties depend on
temperature and pressure
(water vs. steam)
Matter – Properties and Changes
The fact that a substance
can change form is an
important concept in
chemistry and is included
in a list of its physical
properties
2
States of Matter
Solid
Liquid
Definite shape and volume
Particles are tightly packed
When heated, particles
expand slightly
Indefinite shape, definite
volume
Indefinite shape, indefinite
volume
Takes shape of container
Conforms to the shape of the
container
Particles are less tightly
packed than solid; they are
able to slide past each other
(flow)
May not conform to the
shape of a container
Incompressible
Gas
Expands when heated
Virtually incompressible
Fills the entire volume of
the container
Particles are far apart
Easily compressed
gas
Gas
Gas: a
substance
that is
naturally in
the gaseous
state at room
temperature
freezing
solid
melting
liquid
Temperature goes up
= energy absorbed
Temperature goes down
= energy released
Vapor
Vapor:
gaseous state
of a
substance
that is a solid
or liquid at
room
temperature
Matter – Properties and Changes
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Section 3.2 Changes in Matter
Physical changes:

Chemical changes:
Any change to a substance without a change
in its composition
o Drawing copper into a wire
Crumpling up aluminum foil
Cutting paper
Breaking a crystal

State of matter depends on:
temperature, pressure of surroundings

As temperature and pressure changes, most
substances undergo a phase change

Phase change  physical change

The temperature and pressure at which a
substance undergoes a phase change is a
physical property of that substance, just like
its melting point or boiling point.

Ability of a substance to combine with
or change into 1 or more substances
AKA: chemical reaction –
o new substance formed in the
reaction that has different
composition and properties from
the original
o Crushing grapes  physical
change
Fermenting grape juice and
sugars into wine  chemical
change
o Iron rusting (forms iron oxide) 
chemical change
A+BC
reactants (A + B) yield product (C)
Verbs for chemical change:
Verbs for physical change:
boil, freeze, condense,
vaporize, melt
explode, rust, oxidize, corrode, tarnish,
ferment, burn, rot

Late 18th century
scientists used
quantitative tools to
monitor chemical change.
Color change generally means a
Conservation
of mass
Measure mass carefully before and after the chemical reaction:
New tools: analytical
balance for measuring
small changes in mass
chemical change has occurred
total mass remained constant
o Law of Conservation of Mass  mass is neither created
nor destroyed during a chemical reaction; it is
conserved
Mass reactants = Mass products


Matter – Properties and Changes
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Indicators of a chemical reaction:
Sample problem:
color change and production of gas
In an experiment, 10.00 g of red mercury (II) oxide
powder is placed in an open flask and heated until
it is converted to liquid mercury and oxygen gas.
The liquid mercury has a mass of 9.26 g. What is
the mass of oxygen formed in the reaction?
When reaction occurs in a closed
container, the gas cannot escape.
Section 3.3 Mixtures of Matter
Matter
Mixture: combination of 2+ pure
substances in which each substance
retains its individual properties

Most everyday matter occurs
as mixtures

sand and water
salt and water
Heterogeneous mixture
(sand/water)
Substances (pure)  form
of matter with a uniform
and unchanging
composition
Ways to separate
mixtures:
Homogeneous mixture (salt/water)
Constant composition throughout
Does not blend smoothly
throughout


Mixture is not
uniform
2+ substances
remain distinct


AKA: solution
Can be solid, liquid, gas
Examples:
solid – steel
liquid – vinegar (acetic acid &
water)
gas – air we breathe (nitrogen,
oxygen, and others)
1. Filtration – porous
barrier to separate
solid from liquid
2. Distillation –
separating
substances based
on boiling point
3. Crystallization –
(rock candy)
water evaporates
from sugar-water,
leaving sugar
crystals
4. Chromatography –
separate
substances by
their traveling
across material
Matter – Properties and Changes
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Section 3.4 Elements and Compounds
Matter
Mixtures
Pure substances
Compounds
Elements

pure substances that cannot be
separated into simpler substances by
physical or chemical means.
o

undergo chemical
changes
Elements 1-92 are naturally
occurring (exceptions: 43, 61 are
synthetic)
Unique chemical name and symbol
o Symbol consists of 1, 2, or 3 letters
(first letter is CAPITALIZED, rest
are lowercase)

a combination of 2+ different elements that
are combined chemically
o Ex. water, table salt, sugar

Chemical formulae are written from chemical
symbols

Can be broken down into simpler substances
by chemical means
o separating a compound into its
elements often requires external
energy (heat, electricity)

Properties of a compound are different from
those of its elements
Periodic Table of Elements
Law of Definite Proportions
1869: Russian chemist Dmitri Mendeleev
organized all of the elements known at
that time into rows and columns based on
similarities and their masses.
Regardless of amount, a compound is always
composed of the same elements in the same
ratio.





o
1st version of Periodic Table
Rows = periods
Columns = groups or families
“periods”  pattern of similar
properties repeats from period to
period
Mendeleev’s table left blank spots for
currently unknown elements
Made predictions of elements not
yet discovered by analyzing
similarities among elements and
patterns of repetition; predictions
were accurate.


H2O, NaCl, C12H22O11
Sum of the individual masses of the
elements that make up the compound =
mass of the compound
Percent Mass
Ratio of the mass of each element in a compound
to the compound as a whole.
Law of Multiple Proportions
Different compounds may be formed from the
same elements combined in different ratios.
Matter – Properties and Changes
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Sample problem:
Sample problem 2:
A compound is analyzed and found to be
50.0% sulfur and 50.0% oxygen. If the total
amount of the sulfur oxide compound is
12.5 g, how many grams of sulfur are
there?
Two unknown compounds are analyzed.
Compound 1 contains 5.63 g of tin and 3.37
g of chlorine, while compound II contains
2.5 g of tin and 2.98 g of chlorine. Are the
compounds the same?
Matter – Properties and Changes
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Matter – Properties and Changes
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