Unit 1.
Matter and
Change
Do Now:
 What
are the parts of the scientific method and
explain each part?
Objective
 Identify
the common steps of scientific
methods
 Identify types of variables
 Describe the difference between a theory
and a scientific law
 Identify the characteristics of a substance
 Distinguish between physical and
chemical properties
 Differentiate among the physical states of
matter
Objectives Continued


Define physical and chemical changes and
list common changes
Apply the law of conservation of mass to
chemical reactions
 Activities:



Q&A
Solve problems
Exit Ticket
Scientific Method (p12)
 Systematic
approach used in scientific
study
 Method for scientists to verify the work of
others
Steps of Scientific Method
 Observation
 Hypothesis
 Experiments
 Conclusion
 Theory/
Scientific Law
Observation
 Act


of gathering information
Qualitative Data- color, shape, odor other
physical characteristics
Quantitative Data – some type of
measurement. It is numerical. Ex. Height,
weight, how fast, how slow etc.
Hypothesis
 Tentative
explanation for what has been
observed.
Experiment
 Set
of controlled observations used to test
the hypothesis
 Must carefully plan and set up one or
more laboratory experiments in order to
change and test one variable at a time.


Independent Variable – Variable that you
plan to change. ( what you can control)
Dependent Variable – variable that
changes based on the independent
variable
Conclusion
 Judgment
obtained
based on the information
Ch. 3 - Matter
I.
Kinetic Molecular Theory
States of Matter
A. Kinetic Molecular Theory
 KMT
 Particles
 The
of matter are always in motion.
kinetic energy (speed) of these
particles increases as temperature
increases.
Three States of Matter
 Solids
 very
low KE - particles vibrate
but can’t move around
 fixed
shape
 fixed
volume
Three States of Matter
 Liquids
 low
KE - particles can move
around but are still close
together
 variable
 fixed
shape
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.
Matter and its Properties
 It’s
Classified!
Physical vs. Chemical
 Physical

Property
can be observed without changing the
identity of the substance
 Extensive
 Chemical

or intensive properties
Property
describes the ability of a substance to
undergo changes in identity
Extensive vs. Intensive
 Extensive
Property

depends on the amount of matter present

ex,.-
 Intensive
Property

depends on the identity of substance, not
the amount

Ex.-
Extensive vs. Intensive



Extensive Property

depends on the amount of matter present

ex,.- Volume, mass, Energy
Intensive Property

depends on the identity of substance, not the
amount

Ex.- melting point, boiling point, conduct
electricity or heat
WHAT ABOUT DENSITY??
A. Extensive vs. Intensive
 Examples:

boiling point

volume

mass

density

conductivity
A. Extensive vs. Intensive
 Examples:

boiling point…… intensive

Volume …. extensive

Mass ….. extensive

Density ….. intensive

Conductivity ….. intensive
Physical vs. Chemical
 Physical
Change

changes the form of a substance without
changing its identity

properties remain the same
 Chemical
Change

changes the identity of a substance

products have different properties
Signs of a Chemical
 change
in color or odor
 formation
of a gas
 formation
of a precipitate (solid)
 change
in light or heat
B. Physical vs. Chemical
Examples:
melting
point
physical
flammable
chemical
density
physical
magnetic
physical
tarnishes
chemical
in air
B. Physical vs. Chemical
Examples:
melting
point
physical
flammable
chemical
density
physical
magnetic
physical
tarnishes
chemical
in air
B. Physical vs. Chemical
 Examples:
rusting
iron
dissolving
in water
burning
a log
melting
ice
grinding
spices
B. Physical vs. Chemical
 Examples:
rusting
iron
dissolving
in water
chemical
physical
burning
a log
chemical
melting
ice
physical
grinding
spices
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
Law of Conservation of
Mass
Pg. 78
#7 A student carefully placed 15.6 g of
sodium in a reactor supplied with an exess
quantity of chlorine gas. When the reaction
was complete, the student obtained 39.7f
of sodium chloride. Calculate how many
grams of chlorine gas reacted. How many
grams of sodium reacted?
Law of Conservation of
Mass
Pg. 78
#7 A student carefully placed 15.6 g of
sodium in a reactor supplied with an exess
quantity of chlorine gas. When the reaction
was complete, the student obtained 39.7f
of sodium chloride. Calculate how many
grams of chlorine gas reacted. How many
grams of sodium reacted?
24.1 g of chlorine gas, 15.6 of sodium
Law of Conservation of
Mass
Pg. 78
#7 A student carefully placed 15.6 g of
sodium in a reactor supplied with an exess
quantity of chlorine gas. When the reaction
was complete, the student obtained 39.7f
of sodium chloride. Calculate how many
grams of chlorine gas reacted. How many
grams of sodium reacted?
24.1 g of chlorine gas, 15.6 of sodium
HW
 Pg
79 (10 and 13 only)
Do Now:
 DO



NOW:
What is a solution?
Describe the difference between a
heterogeneous and homogenous mixture
Calculate % by Mass
Objective:
 SWBAT:




Contrast Mixtures and substances
Classify mixtures as homogeneous or
heterogeneous
List and describe several techniques used
to separate mixtures.
Distinguish between elements and
compounds
 Activities



Q&A
Solve problems
Exit Ticket
Objective:
 Activities:
A. 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
A. Matter Flowchart
 Examples:
graphite
pepper
sugar
paint
soda
(sucrose)
A. Matter Flowchart
 Examples:
graphite
element
pepper
hetero. mixture
sugar
compound
(sucrose)
paint
hetero. mixture
soda
solution
Pure Substances
 Element


composed of identical atoms
EX: copper wire, aluminum foil
Pure Substances
 Compound

composed of 2 or more elements in
a fixed ratio

properties differ from those of
individual elements

EX: table salt (NaCl)
Pure Substances
 Law

A given compound always contains the same, fixed
ratio of elements.
 Law

of Definite Composition
of Multiple Proportions
Elements can combine in different ratios to form
different compounds.
Pure Substances
 For
example…
Two different compounds,
each has a definite composition.
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
Diatomic Elements
 Hydrogen
 Nitrogen
 Oxygen
 Fluorine
 Chlorine
 Bromine
 Iodine
 There
are 7
diatomic elements
 These
atoms are
never alone, if they
are the pair up with
the same atom
C. Mixtures
 Variable
combination of 2 or more pure substances.
Heterogeneous
Homogeneous
C. Mixtures
 Solution



homogeneous
very small particles
no Tyndall effect
Tyndall Effect
C. Mixtures
 Colloid





heterogeneous
medium-sized particles
Tyndall effect
particles don’t settle
EX: milk
C. Mixtures
 Suspension





heterogeneous
large particles
Tyndall effect
particles settle
EX:
fresh-squeezed
lemonade
Mixtures
(variable composition)
 Homogeneous
Solutions

–
evenly distributed
 Heterogeneous

not evenly distributed
Tea – Homogeneous Mixture
Air – Homogeneous Mixture
Alloys – Homogeneous
Mixtures
Cereal – Heterogeneous
Mixture
Sand – Heterogeneous Mixture
C. Mixtures
 Examples:
mayonnaise
muddy
water
fog
saltwater
Italian
salad
dressing
C. Mixtures
 Examples:
mayonnaise
colloid
muddy
suspension
water
fog
colloid
saltwater
solution
Italian
suspension
salad
dressing
% by Mass = Percent by Mass
 AKA
% composition
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?
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 = (12.4/78.0) * 100 = 15.9%
Extra problems:
 Pg
88 (# 22-23)
HW
 Pg
90 (28)
 Pg 94 (32, 37, 38, 40, 42, 43,44,50,52,
57,58,60, 62, 64, 76, 92)
Do Now:
 Look
at your periodic table.
 What important information can you get
from the PTOE?
Objectives



Find patterns in the Periodic Table
Classify elements as metals, non-metals or
metalloids
Distinguish between metals, non-metals or
metalloids
 Activities


PPT
Group work
The Periodic
Table
A. Mendeleev
 Dmitri



Mendeleev (1869, Russian)
Organized elements
by increasing
atomic mass.
Elements with
similar properties
were grouped
together.
There were some
discrepancies.
C. Johannesson
A. Mendeleev
 Dmitri

Mendeleev (1869, Russian)
Predicted properties of undiscovered elements.
C. Johannesson
B. Moseley
 Henry
Moseley (1913, British)

Organized elements by increasing atomic
number.

Resolved discrepancies in Mendeleev’s
arrangement.
C. Johannesson
How PTOE is organized
 Metals
 Nonmetals
 Metalloids
C. Johannesson
B. Blocks
 Main Group Elements
 Transition Metals
 Inner Transition Metals
C. Johannesson
Periods and Families
 Periods:


horizontal rows on the periodic table
physical and chemical properties change
somewhat regularly across a row.
Elements closer to each other in the same
period tend to be similar than those that are
farther apart.
 Families:
groups

vertical rows of elements, aka
Each group contains similar chemical properties
Types of Elements

METALS:

Shiny
Conductors of heat and electricity


Most metals are malleable (can be pounded into thin sheets; a sugar
cube sized chunk of gold can be pounded into a thin sheet which will
cover a football field),

Most metals are ductile (can be drawn out into a thin wire).

All are solids at room temp (except Mercury, which is a liquid)
Metals tend to have low ionization energies, and typically lose
electrons (i.e. are oxidized) when they undergo chemical reactions
 Alkali metals are always 1+ (lose the electron in s subshell)
 Alkaline earth metals are always 2+ (lose both electrons in s
subshell)
Compounds of metals with non-metals tend to be ionic in nature.


Types of Elements

NON- METALS:
Vary greatly in appearance
 Non-lustrous
 Poor conductors of heat and electricity
 The melting points of non-metals are generally lower than
metals
 Seven non-metals exist under standard conditions as
diatomic molecules:
H2(g) N2(g) O2(g) F2(g) Cl2(g) Br2(l) I2(l) (volatile liquid evaporates readily)
 Nonmetals, when reacting with metals, tend to gain
electrons (typically attaining noble gas electron
configuration) and become anions: Nonmetal + Metal ->
Salt
 Compounds composed entirely of nonmetals are molecular
substances (not ionic)

Types of Elements
 Metalloids:
 Elements
may share properties of metals and
non-metals.
Exit Ticket:
 Classify
the following as either a
METAL, NON-METAL or METALLOID:
a. Au
b. Si
c. Br
d. An
element that is brittle and conducts
electricity
e. An element that is malleable
f. An element that has tendency to become an
anion
Review of Unit
 List
topics we have covered.
 Study
for test!