constructing ideas in physical science joan abdallah, aaas darcy hampton, dcps davina pruitt-mentle,...

Constructing Ideas in Physical Science

Joan Abdallah , AAAS

Darcy Hampton, DCPS

Davina Pruitt-Mentle, University of Maryland

CIPS Institute for Middle School CIPS Institute for Middle School Science TeachersScience Teachers

AAAS/DCPS CIPS Workshop 8/2-8/13 2

Session 9 DebriefingSession 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


Deeper QuestionsDeeper Questions

• What deeper questions could you envision students asking?

• What misconceptions or misinterpretations can you foresee?

• How or what would you say?


CIPSCIPS

• Unit 4– Cycle 1– Activity 3


Energy & HeatEnergy & Heat

• Physical and chemical changes are always accomplished by energy transfer

• The most common form of energy transform or change is heat– Heat is a form of energy that

flows between a system and its surroundings

– Heat flows from a warmer object to a cooler one

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).


Energy (continued)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?


Dietary CaloriesDietary 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


Energy TransferEnergy 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


MatterMatter

Mixture• Most Natural Samples• Physical combination of

2 or more substances• Variable composition • Properties vary as

composition varies• Can separate by

physical means

Pure Substance• Few naturally pure gold

& diamond• Only 1 substance• Definite and constant

composition• Properties under a

given set of conditions


MixtureMixture

Heterogeneous• Visible difference in

parts and phases– Oil and vinegar– Cookie– Pizza– Dirt– Marble– Raw Milk

Homogeneous• Only 1 visible phase

– Homogenized milk– Air (pure)– Metal Alloy (14K

gold)– Sugar and Water– Gasoline


Pure SubstancePure Substance

Compound

aspirin, H2O, CO2

• Can be broken down into 2 or more simpler substances by chemical means

• Over six million known chemical combinations of 2 or more elements

• 7000 more discovered per week with chemical abstracts service

• Definite-constant element

composition

Element

Au, Ag, Cu, H+

• Pure and cannot be divided into simpler substances by physical or chemical means

• 90 naturally occurring• 22 synthetic

CompoundElement

Simpler Compound

Element

Element

MatterMatter

Heterogeneous materialsHeterogeneous materials Homogeneous materialsHomogeneous materials

SolutionsSolutions Pure substancesPure substances

MixturesMixtures CompoundsCompounds ElementsElements


CIPSCIPS

Unit 5

Cycle 1 & 2 Selected Examples


Subatomic ParticlesSubatomic ParticlesBuilding Blocks of AtomsBuilding Blocks of Atoms

• 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)– 1.675 x 10-24 g– Discovered by James

Chadwick (1932)– Inside nucleus

• Electron: (-)– Outside ‘e’ cloud– 9.109 x 10-28 g (1/1839 of a

proton)– Discovered by Joseph John

Thomson (1897)• It’s charge to mass ration

(e/m) = 1.758819 x 108 c/g

– c = charge of electron in Coulombs

– Millikan determined mass itself


AtomsAtoms

• 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 Matter1. All matter is made of atoms

2. All atoms of a given type are similar to one another and different from all other types

3. The relative number and arrangement of different types of atoms contained in a pure substance determines its identity (Law of Multiple Proportions)

4. Chemical change = a union, separation , or rearrangement of atoms to give a new substance

5. 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


Formula MassFormula Mass

H = 1

O = 16

H2O

2 x 1 = 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


Abundance of Elements in Abundance of Elements in MatterMatter

Universe• H 75-91%• He 9%

Earth

• O2 49.3%

• Fe 16.5%• Si 14.5%• Mg 14.2%

Atmosphere

• N2 78.3%

• O2 21%

Human Body

• H2 63%

• O2 25.5%

• C 9.5%

• N2 1.4%

Earth’s Crust

• O2 60%

• Si 20%• Al 6%

• H2 3%

• Ca 2.5%• Mg 2.4%• Fe 2.2%• Na 2.1%


Element Names – based onElement Names – based on

• Geographical Names– Germanium

(German)– Francium (France)– Polonium (Poland)

• Planets– Mercury– Uranium– Neptunium– Plutonium

• Gods– He (helios – sun’s

corona)

• Properties (color)– Chlorine - chloros –

greenish/yellow– Iridium –iris – various

colors


Chemical SymbolsChemical Symbols

• 1814 – Swedish – Jons Jakob Berzelius– Symbols = shorthand for name

• N = nitrogen• Ca = Calcium

– Latin or other name

– Latin

Iron Fe Ferrum

Gold Au Aurum

Antimony Sb Stibium

Copper Cu Cuprum

Lead Pb Plumbrum

Mercury Hg Hydrargyrum

Potassium KKalium

Silver Ag Argentum

Sodium Na Natrium

Tin Sn Stannum

– German

Tungsten W Wolfram


Generic Nomenclature: Generic Nomenclature: Provisional NamesProvisional 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 Unq• 105 un nil pentium Unp• 106 un nil hex ium Unh• 110 un un nil ium Uun

– Most nave been given names anyway


Atom InformationAtom 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


IsotopesIsotopes

• 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


More Atomic InfoMore 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)


Average Atomic MassAverage 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


ExamplesExamples

• 2 isotopes of Cl– Cl-35 34.9689 76.90%– Cl-37 36.9659 23.1%

= 35.453

• Mg– Mg-24 23.985 78.70%– Mg-25 24.986 10.13%– Mg-26 25.983 11.17%

• Ir– Ir-191 191 37.58%– Ir-193 193 62.42%

Notes SummaryNotes Summary


Quantitative vs. Qualitative Quantitative vs. Qualitative DataData

• Quantitative = numerical value

• Qualitative = descriptive explanation– 20 ml of a red thick liquid

• 20 ml = quantitative• Red, thick, liquid = qualitative


PropertiesProperties

• 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)


Physical vs. Chemical ChangePhysical 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


MatterMatter

• 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


SolutionsSolutions

• Solution = Solute + Solvent

• Solvent usually in larger quantity

Gas• Gas dissolved in gas

(air)• Liquid dissolved in a

gas (humidity)• Solid dissolved in a

gas (moth balls)

Liquid• Gas dissolved in a

liquid (soda)• Liquid dissolved in a

liquid (vinegar)• Solid dissolved in a

liquid (salt water)

Solid• Gas dissolved in a

solid (platinum)• Liquid dissolved in a

solid (dental filling)• Solid dissolved in a

solid (sterling Ag)

constructing ideas in physical science joan abdallah, aaas darcy hampton, dcps davina pruitt-mentle,...

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