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

2. AAAS/DCPS CIPS Workshop8/2-8/13 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

3. AAAS/DCPS CIPS Workshop8/2-8/13 3 Deeper Questions What deeper questions could you envision students asking? What misconceptions or misinterpretations can you foresee? How or what would you say?

4. AAAS/DCPS CIPS Workshop8/2-8/13 4 CIPS Unit 4 –Cycle 1 –Activity 3

5. AAAS/DCPS CIPS Workshop8/2-8/13 5 Energy & 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).

6. AAAS/DCPS CIPS Workshop8/2-8/13 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?

7. AAAS/DCPS CIPS Workshop8/2-8/13 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. AAAS/DCPS CIPS Workshop8/2-8/13 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 = T f - T i 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

9. AAAS/DCPS CIPS Workshop8/2-8/13 9 Matter 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

10. AAAS/DCPS CIPS Workshop8/2-8/13 10 Mixture 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

11. AAAS/DCPS CIPS Workshop8/2-8/13 11 Pure Substance Compound aspirin, H 2 O, CO 2 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 Compound Element Simpler Compound Element

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

13. AAAS/DCPS CIPS Workshop8/2-8/13 13 CIPS Unit 5 Cycle 1 & 2 Selected Examples

14. AAAS/DCPS CIPS Workshop8/2-8/13 14 Subatomic Particles Building 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 10 8 c/g –c = charge of electron in Coulombs –Millikan determined mass itself

15. AAAS/DCPS CIPS Workshop8/2-8/13 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 1.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

16. AAAS/DCPS CIPS Workshop8/2-8/13 16 Formula Mass H = 1 O = 16 H 2 O 2 x 1 = 2 1 x 16 = 16 Total = 18  Billy = 150 Susie = 100 Billy 4 Susie = 800 H 2 SO 4 H = 2x1 = 2 S = 1 x 32 = 32 O = 4 x 16 = 64 Total 98  2CaCl 2 Ca = 2x40 = 80 S = 4 x 36 = 144 Total 224 

17. AAAS/DCPS CIPS Workshop8/2-8/13 17 Abundance of Elements in Matter Universe H 75-91% He 9% Earth O 2 49.3% Fe 16.5% Si 14.5% Mg 14.2% Atmosphere N 2 78.3% O 2 21% Human Body H 2 63% O 2 25.5% C 9.5% N 2 1.4% Earth’s Crust O 2 60% Si 20% Al 6% H 2 3% Ca 2.5% Mg 2.4% Fe 2.2% Na 2.1%

18. AAAS/DCPS CIPS Workshop8/2-8/13 18 Element 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

19. AAAS/DCPS CIPS Workshop8/2-8/13 19 Chemical Symbols 1814 – Swedish – Jons Jakob Berzelius –Symbols = shorthand for name N = nitrogen Ca = Calcium –Latin or other name –Latin IronFeFerrum GoldAuAurum AntimonySbStibium CopperCuCuprum LeadPbPlumbrum MercuryHg Hydrargyrum PotassiumKKalium SilverAgArgentum SodiumNaNatrium TinSnStannum –German TungstenWWolfram

20. AAAS/DCPS CIPS Workshop8/2-8/13 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 iumUnq 105 un nil pentiumUnp 106 un nil hex iumUnh 110 un un nil iumUun –Most nave been given names anyway

21. AAAS/DCPS CIPS Workshop8/2-8/13 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

22. AAAS/DCPS CIPS Workshop8/2-8/13 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

23. AAAS/DCPS CIPS Workshop8/2-8/13 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 -24 g –This is just an arbitrary standard (it used to be oxygen -16)

24. AAAS/DCPS CIPS Workshop8/2-8/13 24 Average Atomic Mass Based on Carbon 12 standard One C-12 atom = mass of 12 amu –e=9.10953x10 -24 g = 0.000549  –p=1.67265x10 -24 g = 1.0073  –N=1.67495x10 -24 g = 1.0087 

25. AAAS/DCPS CIPS Workshop8/2-8/13 25 Examples 2 isotopes of Cl –Cl-35 34.9689  76.90% –Cl-3736.9659  23.1% = 35.453  Mg –Mg-2423.985  78.70% –Mg-2524.986  10.13% –Mg-2625.983  11.17% Ir –Ir-191191  37.58% –Ir-193193  62.42%

26. Notes Summary

27. AAAS/DCPS CIPS Workshop8/2-8/13 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

28. AAAS/DCPS CIPS Workshop8/2-8/13 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)

29. AAAS/DCPS CIPS Workshop8/2-8/13 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

30. AAAS/DCPS CIPS Workshop8/2-8/13 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

31. AAAS/DCPS CIPS Workshop8/2-8/13 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)