Transcript EDUC 498Z Designing Learning with Technology

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 9 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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CIPS
• Unit 4
– Cycle 1
– Activity 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
Cycle 1 & 2 Selected Examples
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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%
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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
Gold
Antimony
Copper
Lead
Mercury
Potassium
Silver
Sodium
Tin
Fe
Au
Sb
Cu
Pb
Hg
K
Ag
Na
Sn
Ferrum
Aurum
Stibium
Cuprum
Plumbrum
Hydrargyrum
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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Atom Information
• Atomic Number = # of p, or # of e
• Mass number = # of p + # of n (nucleons)
• Number of n = mass # - atomic #
8 # of p and e
O element symbol
16 # of p+n
• ( ) on chart indicates unstable/synthetic … to
indicate uncalculated
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Isotopes
• Same atomic number, different mass
– Different number of neutrons
– Most elements in nature have isotopes
– Element with the most # of isotopes
• Xe – 36
– Cs – 1 stable/35 radioactive
– C – 13 isotopes
– U – 19 isotopes
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More Atomic Info
• Isobars – same mass but different atomic number
• Isotopes – same atomic number different mass
• Atomic Mass (or atomic weight) – Average relative
mass
– Scale of 12/6 C (12.0000 AMU’s standard)
– Takes into account isotopes and % abundance
as found in nature
– 1 amu = ½ the mass of 1 atom of C and =
1.6605x10-24g
– This is just an arbitrary standard (it used to be
oxygen -16)
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Average Atomic Mass
• Based on Carbon 12 standard
• One C-12 atom = mass of 12 amu
– e=9.10953x10-24g = 0.000549
– p=1.67265x10-24g = 1.0073 
– N=1.67495x10-24g = 1.0087 
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Examples
• 2 isotopes of Cl
– Cl-35
– Cl-37
34.9689 
36.9659 
76.90%
23.1%
= 35.453 
• Mg
– Mg-24
– Mg-25
– Mg-26
23.985 
24.986 
25.983 
78.70%
10.13%
11.17%
191 
193 
37.58%
62.42%
• Ir
– Ir-191
– Ir-193
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Notes Summary
Quantitative vs. Qualitative
Data
• Quantitative = numerical value
• Qualitative = descriptive explanation
– 20 ml of a red thick liquid
• 20 ml = quantitative
• Red, thick, liquid = qualitative
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Properties
• Physical
– Can be observed or measured without altering the identity of
the material
• Chemical
– Refers to the ability of a substance to undergo a change that
alters its identity
• Extensive physical
– Depend on the amount of the material present (ex. mass,
length, & volume)
• Intensive physical
– Does not depend on the amount of material present (ex.
density, boiling point, ductility, malleability, color)
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Physical vs. Chemical Change
• Physical
– Any change in a property of matter that does not result in a change
identity
• Ex. Changes of state – changes between the gaseous, liquid, and solid
state do not alter the identity of the substance
• Chemical
– Any change in which one or more substances are converted into
different substances with different characteristics
• Indications of a chemical change
– Heat/and or light produced
– Production of a gas
– Formation of a precipitate
• Chemical and Physical changes are accompanied by energy
changes: released (exothermic) or absorbed (endothermic)
• Examples
– Rusting, Burning – Chemical
– Tearing, Melting - Physical
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Matter
• Mixtures vs. Pure Substances
– Mixtures can be separated
• Homogeneous – the same composition
throughout – air/water
• Heterogeneous – different layers or parts –
pizza/blood/oil & vinegar
– Pure substances – cannot be separated
• Compounds can be further subdivided
chemically (water/carbon dioxide
• Elements – cannot be subdivided
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Solutions
• Solution = Solute + Solvent
• Solvent usually in larger quantity
Gas
Liquid
• Gas dissolved in gas
(air)
• Liquid dissolved in a
gas (humidity)
• Solid dissolved in a
gas (moth balls)
• Gas dissolved in a
• Gas dissolved in a
liquid (soda)
solid (platinum)
• Liquid dissolved in a
• Liquid dissolved in a
liquid (vinegar)
solid (dental filling)
• Solid dissolved in a
• Solid dissolved in a
liquid (salt water)
solid (sterling Ag)
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