CIPS Institute for Middle School
Science Teachers
Constructing Ideas in Physical Science
Joan Abdallah , AAAS
Darcy Hampton, DCPS
Davina Pruitt-Mentle, University of Maryland
Session 8 Debriefing
• What do you remember from yesterday’s
session (no peeking at text or notes)
• What were the “essential questions” being
asked/explored
• What conclusions did “we” decide
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Deeper Questions
• What deeper questions could you
envision students asking?
• What misconceptions or
misinterpretations can you foresee?
• How or what would you say?
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Deeper Questions or Possible
Misinterpretations
“What makes light”?
“What makes energy”?
“What makes different
colors”?
What would you say?
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Electromagnetic Spectrum
How Roy G. BV Lost a Vowel
This was originally "ROY G. BIV", because it used to be common to call the region between blue
and violet "indigo". In modern usage, indigo is not usually distinguished as a separate color in the
visible spectrum; thus Roy no longer has any vowels in his last name.
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Radiant Energy
See Handout:
Continuous and Line Spectra
Read Aloud
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So….
 Things are made up of atoms
 Atoms
 Protons
 Electrons
 Electrons do a lot of spinning and hopping around
(that’s what causes things to have certain shapes
and textures)
 When electrons get excited, they jump from lower
ground state to excited state and then back to rest
again
 This jumping back and forth = radiant energy
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How Do you Measure Radiant Energy?
• In order to understand behavior of e-, you need to know their:
– Velocity
– Location
• Werner Heisenberg (highs-en-berg), German, showed that it is
impossible to know both the exact position and the exact
momentum of an object (e-) at the same time (Heisenberg’s
Uncertainity Principle)
– Can not measure where the e- is since the “nature” of measuring is
to “move” something
– To know location you would have to “measure” it –but when you
measure it would effect (change) the velocity
– Smaller something is the more uncertain the position will be after
measuring it
– (Δx) (Δmv)  h/4
Need more coffee?
• (Δx) = change in position
• (Δmv) = momentum = mass x velocity (related to KE)
– h/4 = some constant (Planck’s
constant/4)
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The de Broglie Hypothesis
• This enabled de
Broglie to predict the
wavelength of a
particle when given
mass (m) and
velocity.
• General Trend
• 1923, de Broglie
(French)
• Used Planck’s/(and
Einstein) idea…that
radiation is made up of
packets of energy (this
gave waves properties
of particles)
• He wanted to prove that
particles could have
properties like waves
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– as mass (e-)
increases,  decreases
– e- mass ↑  ↓
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The de Broglie Hypothesis cont.
• Using E = mc2 (Einstein) and E = h  (Planck)
• Derived: mc2 = h 
• Substituted v (general velocity) for c
• Substituted v/  for , because the frequency of a
wave is equal to its velocity/by its wavelength
• mv2 = h/
• or  = h /mv2 = h/mv
See: http://cougar.slvhs.slv.k12.ca.us/~pboomer/chemtextbook/cch9.html
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The de Broglie Hypothesis cont.
• From this, shown that e- stream acts in the same way
as a ray of light
• Given credit for indicating how to predict the
wavelength of particular electrons
• Also showed that e- have properties of both waves
and particles = wave-particle duality of nature
• This is why you can not measure the velocity &
location of e- at the same time
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Two Formulas
•  (m) X  (s-1) = c (where c =speed of light or 3.00 x 108 m/s)
• Louis de Broglie suggested that the e- in its circular path about
the nucleus has associated with it particular wavelengths, and
also that the wavelength of the e- depends on its mass and
velocity
• He called this matter waves, and used it to describe the wave
characteristics of material particles
•  = h/mv
– mv also called momentum
– h= Planck’s constant 6.63 x 10-34 J.s (1 J = 1 kg m2/s2)
Show Subscripts and Symbols
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See handout (5.2)
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Making the Connection
• De Broglie = e- act as
waves (properties of 
and  )
• Schrodinger = e- act
as particles -different  (energy
property) &  (mass)
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• The link between
these two concepts
= h = Planck’s
constant
• Wave-particle
duality of nature
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Wave-particle Duality of Nature
• Shows or proves that a
beam of e- will produce
a diffraction pattern like
light patterns
• Bohr and Schodinger
called this : wave or
quantum mechanics
• i.e., where are we more
likely to “find” e- at a
given moment in time
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1s _
2s_
2p _ _ _
•
•
•
•
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1 & 2 = shells
S,p,d,f = subshells
_ = orbitals
Distance between the rings =
“nodes”, places where you
will not find e-
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Summary
• Planck’s hypothesis stated that energy
is given off on quanta (photon)
continuously
• Bohr showed that absorption of light at
set  correspond to definite changes in
energy of the e• Reasoned that orbits (rings) around
nucleus must have a definite diameter
and that e- could occupy only certain
orbits
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Summary cont.
• The energy absorbed when the atom was excited =
the energy difference between orbits
• Because these orbits represent definite energy
levels, a definite amount of energy is radiated
• The size of the smallest orbit an e- can occupy (one
closest to nucleus), the ground state, can be
calculated
• Energy is determined by the movement of e- between
energy levels that are specific for each element
• The same set of energy levels will always produce
the same spectrum
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To Learn More
• CEA Light Tour
[Local]
See Handout
From:
http://cse.ssl.berkeley.edu/light/light_tour.html
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Other Resources
• Waves –Virtual Lab [Local]
• Exploring Earth [Local] Observe the change in a
star's spectrum as its motion changes
• Electromagnetic spectrum - Wikipedia
• Discovery-The Color Spectrum How does it work?
[Local]
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But I can not see e-, so how do we know?
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States of Matter
• Matter- has mass, occupies space
• Physical states of matter
– Solid
– Liquid
– Gas
– Plasma
One of the four states of matter. (The other three are solid, liquid and gas.) Consists of
a gas of positively charged and negatively charged particles with approximately equal
concentrations of both so that the total gas is approximately charge neutral. A plasma
can be produced from a gas if enough energy is added to cause the electrically neutral
atoms of the gas to split into positively and negatively charged atoms and electrons.
See also: The Plasma State of Matter.
www.spacescience.org/ExploringSpace/Glossary/1.html
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Kinetic Molecular Theory of Matter
Solid
• States:
Liquid
Gas
Own definite
– All matter is in constant shape
motion
Own definite
volume
– An increase in
Indefinite
shape
Indefinite
shape
Definite
volume
Indefinite
volume
temperature increases Independent
motion and decreases of container
attraction forces holding shape
the matter together
– SLGP
Takes shape
of container
(until it fills it)
Takes
shapes of
container
and fills it
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Physical vs. Chemical Properties
Physical Properties
Chemical Properties
• Can be observed without
changing form
–
–
–
–
–
–
–
–
–
–
–
–
–
• Undergoes changes in
chemical composition
– Flammability or not
flammable
– Reacts/failure to react
with another
– Decomposes
– Rusting
– Combustion
Color
Odor
Taste
Size
BPo
MPo
Density
Specific heat (Cp)
Hardness
Solubility
Mass
Temperature
Heat capacity
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Properties can also be classified as:
Extensive
• Depend on sample
size
Intensive
• Values do not
depend on size of
portion
–
–
–
–
– Mass
– Volume
– Length
Temperature
MPo
BPo
FPo
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Properties
• Molecules vibrate faster when they are
stirred-therefore, this helps them dissolve
faster
• When heated dissolves faster
• At certain temperature (w/ a solid) when heat
added, the heat breaks the bonds. Solid
matter changes to liquid (Melting Point MPo)
• With a solid when freezes, attractive forces
cause molecules to lock together into solid
state (Freezing Point FPo)
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Liquid Changes
•
•
•
•
Change of liquid into vapor  evaporation
Change of vapor into a liquid  condensation
Opposite of condensation  evaporation
Opposite of evaporation  condensation
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Changes cont.
• As temperature falls, and the gaseous molecules slow down,
their weak attractive forces get an opportunity to bind the
molecules together and change the gas (vapor) into a liquid.
When water vapor touches cool dust particles in the air,
condensation takes place. The droplets of water, suspended in
the air, form clouds and rain
– Gas  Condensation  Liquid
• The changing of a solid into a gas without becoming liquid 
sublimation. A lot of heat is added to the solid. This added heat
causes the molecular vibrations to become so violent that the
molecules of the solid completely break away from each other
and enter into a gaseous state
– Solid  Sublimation  Gas
•
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Ex. Mothballs, vaporization (nuclear fallout)
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Changes cont.
• We know that water vapor will condense on a cool
speck of dust. If the water vapor touches a very cold
speck of dust in the air, the gaseous water may
crystallize without condensing first. The ice crystals,
suspended in the air, form clouds. If conditions are
right, these crystals may fall to the ground as snow.
– The changing of a gas into a solid = sublimation
• By definition, sublimation can indicate going from gas
to solid or from solid to gas…although in “chemistry”
usually implies going from the solid state to a gas
state.
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Putting It All Together
GAS
SOLID
LIQUID
Melts
Freezes
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Property Changes
Chemical Change
• Involves a change in
basic nature (chemical
composition)
• Change in at least one
new substance
– Sulfur & iron heated
– Burning paper
– Digesting
– Sour milk
– Detonation
Physical Change
• No new substance is
ever formed
– Tearing paper
– Sulfur & iron
– Sharpening
– Bite
– Chew
– Breaking glass
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Quiz
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
Rust
Melts
Sharpening
Digesting
Biting
Burning
Slicing
Detonation
Souring
Breaking
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C
P
P
C
P
C
P
C
C
P
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Break?
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CIPS
• Unit 4
– Cycle 1
– Activity 1,2 & 3
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Energy & Heat
• Physical and chemical
changes are always
accomplished by energy
transfer
• The most common form of
energy transform or
change is heat
Ex.
Object A = 25°C
Object B = 20°C
What happens when they are
mixed?
Energy will continue to transfer
until the temperature of the
objects are equal.
The energy transfer as a result
of a temperature difference
is called heat and is
represented by the letter (q).
– Heat is a form of energy that
flows between a system and
its surroundings
– Heat flows from a warmer
object to a cooler one
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Energy (continued)
• If energy is absorbed = endothermic reaction
• If energy is given off = exothermic reaction
– Match = exothermic
– Cold pack = endothermic
• Both forms require a certain amount of energy to get started – activation
energy
• Quantitative measurements of energy changes are expressed in joules
(J). This is a derived SI unit
–
–
–
–
Older unit = calorie
One calorie (c) = 4.184 J
(C) dietary unit  calorie (c)
The heat needed to raise 1 g of a substance by 1°C is called specific heat
(Cp) of the substance
Examples: Sand and water – different Cp values
Which gets hotter at the beach?
Which cools down faster?
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Dietary Calories
• The heat required to increase the temperature of 1g of water 1°C =
4.184J
• Dietary Calories (C) are 1000 times as large as a calorie (c)
• Caloric values are the amount of energy the human body can
obtain by chemically breaking down food
• The Law of Conservation of Energy shows that in an insulated
system, any heat loss by 1 quantity of matter must be gained by
another. The transfer of energy takes place between 2 quantities of
matter that are at different temperatures until they both reach an
equal temperature
Example: An average size backed potato (200g) has an energy value of 686,000 J. How many
calories is this?
4.184J = 1 c, 1000 c = 1 C
686000J/4.184 J = 164,000 c
164,000 c/ 1000 C=164C
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Energy Transfer
• The amount of heat energy transferred can be
calculated by:
– (heat gained) = (mass in grams)(change in T)(specific heat)
– q = (m)(T)(Cp)
– T = Tf - Ti
Example: How much heat is lost when a solid aluminum block with a mass
of 4100g cools from 660.0°C to 25°C? (Cp = 0.902 J/g°C)
q = (m)(T)(Cp)
T = 660.0°C - 25°C = 635°C
therefore: q = (4110g)(635°C)(0.902 J/g. °C) = 2,350,000 J
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Matter
Mixture
Pure Substance
• Most Natural Samples
• Physical combination of
2 or more substances
• Variable composition
• Properties vary as
composition varies
• Can separate by
physical means
• Few naturally pure gold
& diamond
• Only 1 substance
• Definite and constant
composition
• Properties under a
given set of conditions
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Mixture
Heterogeneous
• Visible difference in
parts and phases
–
–
–
–
–
–
Homogeneous
• Only 1 visible phase
– Homogenized milk
– Air (pure)
– Metal Alloy (14K
gold)
– Sugar and Water
– Gasoline
Oil and vinegar
Cookie
Pizza
Dirt
Marble
Raw Milk
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Pure Substance
Compound
aspirin, H2O, CO2
Element
Au, Ag, Cu, H+
• Can be broken down into 2 or
• Pure and cannot be divided into
more simpler substances by
simpler substances by physical
chemical means
or chemical means
• Over six million known chemical
• 90 naturally occurring
combinations of 2 or more
• 22 synthetic
elements
• 7000 more discovered per week
with chemical abstracts service
• Definite-constant element
Element
Simpler Compound
composition
Compound
Element
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Element
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Matter
Heterogeneous materials
Homogeneous materials
Solutions
Mixtures
Compounds
Pure substances
Elements
CIPS
Unit 5
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Subatomic Particles
Building Blocks of Atoms
• Electron: (-)
• Proton: (+)
– 1.673 x 10-28 g
– Discovered by Goldstein
(1886)
– Inside the nucleus
(credit given to Rutherford – beam of alpha particles
on thin metal foil experiment. Explained nucleus
in core, made up of neutrons and protons)
• Neutron: (no charge)
• It’s charge to mass ration
(e/m) = 1.758819 x 108 c/g
– c = charge of electron in
Coulombs
– Millikan determined mass
itself
– 1.675 x
g
– Discovered by James
Chadwick (1932)
– Inside nucleus
10-24
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– Outside ‘e’ cloud
– 9.109 x 10-28 g (1/1839 of a
proton)
– Discovered by Joseph John
Thomson (1897)
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Atoms
•
•
•
Atom – smallest particle of an element that can exist and still hold properties
“Atomos” – Greek – uncut/indivisible. Democritus proposed that elements are
composed of tiny particles
John Dalton (1808) published The Atomic Theory of Matter
1.
2.
3.
4.
5.
•
All matter is made of atoms
All atoms of a given type are similar to one another and different from all other types
The relative number and arrangement of different types of atoms contained in a pure
substance determines its identity (Law of Multiple Proportions)
Chemical change = a union, separation , or rearrangement of atoms to give a new
substance
Only whole atoms can participate in or result from any chemical change, since atoms
are considered indestructible during such changes (Law of Conservation of Mass)
Antonine Lavoier demonstrated via careful measurements that when combustion
is carried out in a closed container – the mass of the products = the mass of the
reactants
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Formula Mass
H=1
O = 16
H 2O
2x1=2
1 x 16 = 16
Total = 18 
Billy = 150
Susie = 100
Billy4Susie = 800
H2SO4
H = 2x1 = 2
S = 1 x 32 = 32
O = 4 x 16 = 64
Total 98
2CaCl2
Ca = 2x40 = 80
S = 4 x 36 = 144
Total 224
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Abundance of Elements in
Matter
Universe
• H 75-91%
• He 9%
Atmosphere
• N2 78.3%
• O2 21%
Earth
• O2 49.3%
• Fe 16.5%
• Si 14.5%
• Mg 14.2%
Human Body
• H2 63%
• O2 25.5%
• C 9.5%
• N2 1.4%
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Earth’s Crust
• O2 60%
• Si 20%
• Al 6%
• H2 3%
• Ca 2.5%
• Mg 2.4%
• Fe 2.2%
• Na 2.1%
45
Element Names – based on
• Geographical
Names
– Germanium
(German)
– Francium (France)
– Polonium (Poland)
• Gods
– He (helios – sun’s
corona)
• Properties (color)
– Chlorine - chloros –
greenish/yellow
– Iridium –iris – various
colors
• Planets
–
–
–
–
Mercury
Uranium
Neptunium
Plutonium
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Chemical Symbols
• 1814 – Swedish – Jons Jakob Berzelius
– Symbols = shorthand for name
• N = nitrogen
• Ca = Calcium
– Latin or other name
– German
Tungsten
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W
Wolfram
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– Latin
Iron
Fe
Gold
Au
Antimony Sb
Copper
Cu
Lead
Pb
Mercury
Hg
Hydrargyrum
Potassium K
Silver
Ag
Sodium
Na
Tin
Sn
Ferrum
Aurum
Stibium
Cuprum
Plumbrum
Kalium
Argentum
Natrium
Stannum
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Generic Nomenclature:
Provisional Names
• International Union of Pure and Applied Chemistry
(IUPAC)
• Latin – Greek Names
– 0 =nil, 1=un, 2=bi, 3=tri, 4=quad, 5=pent, 6=hex, 7=sept,
8=oct, 9=enn
– + ium
– i.e.
•
•
•
•
104 un nil quad ium
105 un nil pentium
106 un nil hex ium
110 un un nil ium
Unq
Unp
Unh
Uun
– Most nave been given names anyway
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