1 / 31

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)

chen
Télécharger la présentation

Constructing Ideas in Physical Science

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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

More Related