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Constructing Ideas in Physical Science

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)

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Constructing Ideas in Physical Science

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  1. CIPS Institute for Middle School Science Teachers Constructing Ideas in Physical Science Joan Abdallah, AAAS Darcy Hampton, DCPS Davina Pruitt-Mentle, University of Maryland

  2. 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 8/2-8/13

  3. Deeper Questions • What deeper questions could you envision students asking? • What misconceptions or misinterpretations can you foresee? • How or what would you say? 8/2-8/13

  4. CIPS • Unit 4 • Cycle 1 • Activity 3 8/2-8/13

  5. 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 calledheat and is represented by the letter (q). Energy & Heat 8/2-8/13

  6. 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? 8/2-8/13

  7. 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 8/2-8/13

  8. 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 8/2-8/13

  9. 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 Matter 8/2-8/13

  10. 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 Mixture 8/2-8/13

  11. 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 Pure Substance Element Simpler Compound Compound Element Element 8/2-8/13

  12. Heterogeneous materials Homogeneous materials Solutions Pure substances Mixtures Compounds Elements Matter

  13. CIPS Unit 5 Cycle 1 & 2 Selected Examples 8/2-8/13

  14. 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 Subatomic ParticlesBuilding Blocks of Atoms 8/2-8/13

  15. 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 • 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 8/2-8/13

  16. H = 1 O = 16 H2O 2 x 1 = 2 1 x 16 = 16 Total = 18  Billy = 150 Susie = 100 Billy4Susie = 800 Formula Mass 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 8/2-8/13

  17. 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% Abundance of Elements in Matter Earth’s Crust • O2 60% • Si 20% • Al 6% • H2 3% • Ca 2.5% • Mg 2.4% • Fe 2.2% • Na 2.1% 8/2-8/13

  18. 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 Element Names – based on 8/2-8/13

  19. Chemical 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 K Kalium Silver Ag Argentum Sodium Na Natrium Tin Sn Stannum • German Tungsten W Wolfram 8/2-8/13

  20. 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 Unq • 105 un nil pentium Unp • 106 un nil hex ium Unh • 110 un un nil ium Uun • Most nave been given names anyway 8/2-8/13

  21. 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 8/2-8/13

  22. 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 8/2-8/13

  23. 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) 8/2-8/13

  24. 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  8/2-8/13

  25. Examples • 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% 8/2-8/13

  26. Notes Summary

  27. Quantitative vs. Qualitative Data • Quantitative = numerical value • Qualitative = descriptive explanation • 20 ml of a red thick liquid • 20 ml = quantitative • Red, thick, liquid = qualitative 8/2-8/13

  28. 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) 8/2-8/13

  29. 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 8/2-8/13

  30. 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 8/2-8/13

  31. Solutions • 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) 8/2-8/13

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