CHAPTER 3: MATTER
PROPERTIES AND CHANGES
Mrs. Faria
Chapter BIG idea
Everything is
made up of
matter!!!
Do Now
How did you study for your test?
Be specific.
How can you improve your study
methods….
Do Now: Answer the following questions in your notebook
 What are the parts of a pencil?
 What is each part of the pencil made of?
 Where does each part originate from?
 Do any of the parts have anything in common?
SECTION 1:
PROPERTIES OF
MATTER
Section 1: Essential Questions & Vocabulary
 What characteristics identify a substance?
 What distinguishes physical properties from chemical properties?
 How do the properties of the physical states of matter differ?
Vocabulary
 Density
 Vapor
 States of matter
 Physical Property
 Solid
 Extensive Property
 Liquid
 Intensive Property
 Gas
 Chemical Property
Properties of Matter: Main Idea
 Most common substances exist as solids, liquids, and gases,
which have diverse physical and chemical properties.
 What observations can you make about the glass of ices
water?
 What will happen if you leave the glass of ice water out on the
counter for a long time?
 Does the composition of the water change as the ice melts?
Matter & Substances
 Matter – anything that has a mass and takes up
space.
 Substance – matter with a uniform and unchanging
composition
 Examples: Table Salt & Water
Kinetic Molecular Theory (KMT)
 Particles of matter are always in motion.
 The kinetic energy (speed) of these
particles increases as temperature
increases.
States of Matter
 All matter that exits naturally on Earth can be classified as one of the
physical form called states of matter.
 The three common states of matter can be distinguished by the way they fill a container.
 SOLIDS: matter that has its own definite shape and volume.
 LIQUIDS: matter that flows, has constant volume, and takes the shape of its
container.
 GASES: matter that flows to conform to the shape of its container and fills
the entire volume of its container.
Three States of Matter: SOLIDS
 Very low KE - particles vibrate but can’t move
around
 Definite (fixed)shape
 Definite (fixed) volume
Three States of Matter: LIQUIDS
 Low KE - particles can move around but are still
close together
 Variable shape
 Fixed volume
Three States of Matter: GASES
 High KE - particles can separate and move
throughout container
 Variable shape
 Variable volume
 Vapor- Gaseous state of a substance that is a
solid or liquid at room temperature.
States of Matter: Practice
SOLIDS
Kinetic Energy
Shape
Volume
LIQUIDS
GASES
States of Matter: Practice Solutions
SOLIDS
LIQUIDS
GASES
Kinetic Energy
Very Low
Low
High
Shape
Definite
Definite
Not Definite
Volume
Definite
Definite
Not Definite
Physical and Chemical Properties
Physical Property
Chemical Property
• can be observed
• Describes the ability of
without changing the
a substance to
identity of the
undergo changes in
substance
identity
• Extensive or
intensive properties
Physical vs. Chemical Properties: Practice
Property
Melting point
Flammable
Density
Magnetic
Tarnished in air
Chemical or Physical
Physical vs. Chemical Properties: Practice Solutions
Property
Chemical or Physical
Melting point
Physical
Flammable
Chemical
Density
Physical
Magnetic
Physical
Tarnished in air
Chemical
Extensive vs. Intensive Properties
Extensive Property
Depends on the amount of
matter present
Intensive Property
Depends on the identity of
substance, not the amount
Extensive vs. Intensive Properties: Practice
Physical Property
Boiling Point
Volume
Mass
Density
Conductivity
Extensive or Intensive
Extensive vs. Intensive Properties: Practice
Solutions
Physical Property
Extensive or Intensive
Boiling Point
Intensive
Volume
Extensive
Mass
Extensive
Density
Intensive
Conductivity
Intensive
SECTION 2: CHANGES
IN MATTER
Essential Questions & Vocabulary
 What is a physical change?
 What are common examples of physical changes?
 What defines a chemical change?
 What evidence can be cited to show that a chemical changes has
taken place?
 How does the law of conservation of mass apply to chemical
reactions?
Vocabulary
 Physical Change
 Chemical Change
 Phase Change
 Law of conservation of mass
Physical vs. Chemical Changes
Physical Change
Chemical Change
Changes the form of a
substance without changing
its identity (composition)
Changes the identity of a
substance
Properties remain the same
Products have different
properties
Signs of a Chemical Change
 Change in color or odor
 Formation of a gas (bubbling)
 Formation of a precipitate (solid)
 Change in light or heat (flammability)
 Temperature change
Phase change
 Transition of matter from one state to another
 Phase change can occur with temperature change and/or pressure
change
Physical vs. Chemical Changes: Practice
Change
Rusting Iron
Dissolving in water
Burning a log
Grinding spices
Melting Ice
Chemical or Physical Change
Physical vs. Chemical Changes: Practice
Solutions
Change
Chemical or Physical Change
Rusting Iron
Chemical
Dissolving in water
Physical
Burning a log
Chemical
Grinding spices
Physical
Melting Ice
Physical
Law of Conservation of Mass
 Mass is neither created nor destroyed during a
chemical reaction. It is conserved
 In a chemical reaction, the mass of the reactants
must equal the mass of the products
 Reactants: Starting substances that will change into new
substances.
 Products: Ending substances that will have changed into the new
substance.
 Create a drawing/ cartoon to represent this
SECTION 3: MIXTURES
OF MATTER
Essential Questions & Vocabulary
 How do mixtures and substances differ?
 Why are some mixtures classified as homogeneous, while others are
classified as heterogeneous?
 What are several techniques used to separate mixtures?
Vocabulary
•
•
•
•
•
Mixture
Heterogeneous mixture
Homogeneous mixture
Solution
Filtration
•
•
•
•
Distillation
Crystallization
Sublimation
Chromatography
Mixtures
 Variable combination of 2 or more pure substances.
Heterogeneous
Homogeneous
Mixtures
Heterogeneous Mixtures
• A mixture that does NOT blend
smoothly throughout
• Individual substances remain
distinct
Homogeneous Mixtures
• A mixture that has constant
composition throughout
• Single phase
Mixtures- variable composition
 Homogeneous – Solutions
 evenly distributed
 Heterogeneous
 not evenly distributed
Mixtures
Alloy
 Homogeneous mixture of metals
 Example: Brass – Mixture of Copper and Zinc
 Homogeneous mixture of metals and nonmetals
(metal is the major component)
 Example: Steel – Mixture of Iron and Carbon
Tyndal Effect
• Light scattering by particles in a
colloid or particles in a fine
suspension.
Mixtures
Solution
 homogeneous
 very small particles
 no Tyndall effect
Tyndall Effect
Mixtures
Colloid
 heterogeneous
 medium-sized particles
 Tyndall effect
 particles don’t settle
 EX: milk
Mixtures
Suspension
 heterogeneous
 large particles
 Tyndall effect
 particles settle
 EX: fresh-squeezed lemonade
Mixtures – Practice
Tea
Brass
Sand
Alloy Wheel
Cheerios w/
Strawberries
Air
Mixtures – Practice Solutions
Tea
Homogeneous Mixture
Brass
Homogeneous Mixture
Sand
Heterogeneous Mixture
Cheerios w/
Strawberries
Heterogeneous Mixture
Alloy Wheel
Homogeneous Mixture
Air
Homogeneous Mixture
Mixtures – Colloids, Solutions, Suspensions
Practice
Mixture
Saltwater
Italian Salad Dressing
Mayonnaise
Muddy Water
Fog
Colloid, Solution, Suspension
Mixtures – Colloids, Solutions, Suspensions
Practice Solutions
Mixture
Colloid, Solution, Suspension
Saltwater
Solution
Italian Salad Dressing
Suspension
Mayonnaise
Colloid
Muddy Water
Suspension
Fog
Colloid
Matter Flowchart
Separating Mixtures
Filtration
 Uses a porous barrier to separate a
solid from a liquid.
 Ideal for separating
heterogeneous mixtures.
Separating Mixtures
Distillation
 Physical separation technique that
is based on differences in the
boiling points of the substances
involved.
 Ideal for homogenous mixtures -
solutions
Separating Mixtures
Crystallization
 Separation technique that results in
the formation of pure solid particles
of a substance from a solution
containing the dissolved substance.
 Example: Rock Candy
Separating Mixtures
 Sublimation
 Process during which a solid changes
to vapor without melting, that is,
without going through the liquid state.
 Can be used to separate two solids in
a mixture when one of the solids
sublimates but the other does not.
Separating Mixtures
Chromatography
 Technique that separates the
components of a mixture
dissolved in either a gas or a
liquid based on the ability of
each component to travel or
be drawn across the surface of
a fixed substrate.
SECTION 4:
ELEMENTS AND
COMPOUNDS
Essential Questions & Vocabulary
 What distinguishes elements from compounds?
 How is the periodic table organized?
 What are the laws of definite and multiple proportions and why are
they important?
Vocabulary
• Element
• Law of definite proportions
• Periodic table
• Percent by mass
• Compound
• Law of multiple proportions
Matter Flowchart
MATTER
yes
Can it be physically
separated?
MIXTURE
yes
Is the composition
uniform?
Homogeneous
Mixture
(solution)
PURE SUBSTANCE
no
Heterogeneous
Mixture
Colloids
no
yes
Can it be chemically
decomposed?
Compound
Suspensions
no
Element
Periodic Table
 Dimitri Mendeleev (1834-1907) – first organized the all known
elements according to their mass.
 Organizes the elements into a grid of horizontal rows called periods
and vertical columns called groups or families.
 Elements in the same group (column) have similar physical and
chemical properties.
 Pattern of similar properties repeats from period to period.
Element Information - Periodic Table
Pure Substances - constant composition
 Elements
 Listed on the Periodic Table
 Cannot be broken down
into unique components
 Na, Cl, Al, O2, S8
 Compounds
 Made of elements
that are chemically
joined
 Can be broken down
 NaCl, H2O, AlCl3,
H2SO4
Pure Substances
Element
 composed of identical atoms
 EX: copper wire, aluminum foil
Diatomic Elements
 These atoms are never alone. If they are found alone, then it is always
paired up with the same atom.
 There are 7 diatomic elements
 Hydrogen
 Chlorine
 Nitrogen
 Bromine
 Oxygen
 Iodine
 Fluorine
Pure Substances
Compound
 composed of 2 or more elements in a fixed ratio
 properties differ from those of individual elements
 EX: table salt (NaCl)
Element Symbols – Practice
 Element Symbols – either 1 or 2 letters
 First letter is always Capital
 2nd letter is always lower case
Element Symbol
H
He
Cl
Fe
O
Element Name
Element Symbols – Practice Solutions
 Element Symbols – either 1 or 2 letters
 First letter is always Capital
 2nd letter is always lower case
Element Symbol
H
Element Name
Hydrogen
He
Cl
Helium
Chlorine
Fe
O
Iron
Oxygen
Compounds (Formulas)
 Compound Formulas – Write the element symbols
 No subscript – only one atom represented
 Subscript – the number of atoms of that element in compound
 Parentheses with subscript – the units in the parentheses are
repeated
NaCl
• 1 Na (Sodium)
• 1 Cl (Chlorine)
H 2O
• 2 H (Hydrogen)
• 1 O (Oxygen)
𝐶𝑢(𝑁𝑂3 )2
• 1 Cu (copper)
• 2 N (Nitrogen)
• 6 O (Oxygen)
Pure Substances
 Law of Definite Composition
 A given compound always contains the same, fixed ratio of
elements.
 Law of Multiple Proportions
 Elements can combine in different ratios to form different
compounds.
Chemistry Joke
 Two scientists walk into a restaurant. The first one says “I’ll
have some H2O.” The second one says, “I’ll have some H2O
too.” Then he dies.
Hydrogen Peroxide
Law of Definite Proportions
 A compound is always composed of the same elements in
the same proportion by mass.
 The mass of the compound is equal to the sum of the
masses of the elements that make up the compound.
% by Mass = Percent by Mass
 Also known as percent composition
𝒎𝒂𝒔𝒔 𝒐𝒇 𝒆𝒍𝒆𝒎𝒆𝒏𝒕
𝑷𝒆𝒓𝒄𝒆𝒏𝒕 𝒃𝒚 𝒎𝒂𝒔𝒔 =
× 𝟏𝟎𝟎
𝒎𝒂𝒔𝒔 𝒐𝒇 𝒄𝒐𝒎𝒑𝒐𝒖𝒏𝒅
Pure Substances
 For example…
Two different compounds,
each has a definite composition.
% Mass – Carbon Monoxide & Carbon Dioxide Examples
 Carbon Monoxide – CO
 Carbon Dioxide – CO2
 Carbon = 12.0 g
 Carbon = 12.0 g
 Oxygen = 16.0 g
 Oxygen = 32.0 g
 % of carbon =
12 𝑔
28.0 𝑔
× 100 - 42.9%
 % of Oxygen =
16 𝑔
28.0 𝑔
× 100
 57.1%
 % of carbon =
12 𝑔
44.0 𝑔
× 100 - 27.3%
 % of Oxygen =
32 𝑔
44.0 𝑔
× 100
 72.7%
Example
 P 88 # 19
 A 78.0 g sample of an unknown compounds contains 12.4 g
of hydrogen. What is the present by mass of hydrogen in
the compound?
% mass =
𝟏𝟐.𝟒 𝒈 𝑯𝒚𝒅𝒓𝒐𝒈𝒆𝒏
𝟕𝟖 𝒈 𝑺𝒂𝒎𝒑𝒍𝒆
×100 = 15.9%
Law of Multiple Proportions
 When different compounds are formed by a combination of the same
elements, different masses of one element combine with the same
fixed mass of the other element in a ratio of small whole numbers.
Water versus Hydrogen Peroxide
• Water - H2O
• Hydrogen Peroxide – H2O2
• Hydrogen – 2.0 g
• Hydrogen – 2.0 g
• Oxygen – 16.0 g
• Oxygen – 32.0 g
• 2 g Hydrogen : 16 g of Oxygen
• 2 g Hydrogen : 32 g of Oxygen
• Ratio 1 to 8
• Ratio 1 to 16
Hydrogen Ratio: 1 to 1
Oxygen Ratio: 8 to 16
1 to 2
Law of Multiple Proportions – Practice
Compound
NaCl
CuO
H2O
H2O2
Simple Whole Number Ratio
(Ex. 1:1, 2:2, etc.)
Law of Multiple Proportions – Practice
Solutions
Compound
Simple Whole Number Ratio
(Ex. 1:1, 2:2, etc.)
NaCl
1:1
CuO
1:1
H2O
2:1
H2O2
2:2