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History of the Atomic Model and Scientists' Contributions

Explore the progression of the atomic model throughout history and the significant contributions made by scientists such as Aristotle, Democritus, John Dalton, J.J. Thomson, Robert Millikan, Ernest Rutherford, Niels Bohr, Erwin Schrodinger, and Louis De Broglie. Learn about the development of theories and experiments that shaped our understanding of atoms.

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History of the Atomic Model and Scientists' Contributions

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  1. HISTORY OF THE ATOMIC MODEL Chapters 4, 5.1,+ 25

  2. SCIENTISTSMODEL • Aristole: The Greek Model (400 BC – 320 BC) • 330 BC: Systematic concept of logic (Scientific Method) • Atoms did not exist • All matter was made up of 4 elements: Earth, Water, Fire, and Air • Democritus: The Greek Model (460 BC – 370 BC) • 400 BC: Matter can’t be divided forever; there must be a smallest piece (atomos) • Atoms are indestructible, indivisible, & the fundamental units of matter • Atom: smallest particle of an element that retains the properties of that element. • - no electric charge, electrically neutral • No experiments to test his theories

  3. SCIENTISTMODEL John Dalton: Dalton’s ModelEnglish (1766 – 1844) Dalton’s Atomic Theory (1803): • All elements are composed of atoms that are submicroscopic indivisible particles. • Atoms of the same elements are identical & atoms of different elements are different. • Atoms of different elements can physically mix together or chemically combine w/one another to form simple whole-number ratios to form compounds. • Chemical reactions occur when atoms are separated, rearranged or joined. Atoms of one element can never be changed into atoms of another element.

  4. SCIENTISTMODEL J.J. Thomson: Thomson’s Model English (1856-1940) • 1897: Used cathode ray tube to discover electrons • Cathode ray: glowing beam which travels from the cathode(-) to the anode(+). - are composed of electrons - are attracted to positive metal plate • Atoms had negatively charged particles • ELECTRON: negatively charged subatomic particle • not the original name (corpuscle) • “Plum Pudding” Model • (chocolate chip cookie) (watermelon) • - a ball of positive charge containing electrons

  5. POSITIVE CHARGE ELECTRONS EMBEDDED WITHIN Thomson’s ATOMIC Model

  6. Cathode Ray Tube: http://www.chem.uiuc.edu/demos/cathode.html

  7. Robert Millikan (1868-1953) Oil Drop Experiment (1909) • American • Determined the charge and mass of an electron • The mass is 1/1840 of the mass of a hydrogen atom (unit)?

  8. Empty Space ++++++ + Nucleus SCIENTISTMODEL Ernest Rutherford: Rutherford’s Model New Zealand (1871-1937) • Gold Foil Experiment (1911) • Discovered that most of atom’s mass is located in the positively charged nucleus NUCLEUS: center of the atom composed of PROTONS & NEUTRONS • is 99.9% of the atom’s mass • a marble in a football stadium http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/ruther14.swf

  9. Rutherford’s Gold Foil Experiment: (1911) http://micro.magnet.fsu.edu/electromag/java/rutherford/

  10. Gold Foil Experiment: Rutherford

  11. PROTON: positively charged subatomic particle discovered by Eugen Goldstein (1850-1930) • 1886: put holes in cathode and saw rays traveling in the opposite direction (canal rays) (1 amu) NEUTRON: subatomicparticle with no charge discovered by English scientist Sir James Chadwick (1891-1974) • 1932: mass is nearly equal to proton (1 amu) Thomson & Rutherford proved Dalton’s Theory incorrect: ATOMS ARE DIVISIBLE http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/ruther14.swf

  12. Electrons Energy Levels ++++++ SCIENTISTMODEL Niels Bohr: The Bohr Model (1885-1962) (Danish) • Electrons move in definite orbits around the nucleus (planets around the sun) • 1913: PLANETARY MODEL • Electrons are a part of energy levels located certain distances from the nucleus

  13. Energy Levels: region around the nucleus where the electron is likely to be moving. • a ladder that isn’t equally spaced • further the distance, closer the spacing • the higher the energy level the farther it is from the nucleus Electrons can jump from 1 energy level to another. • Quantum Energy: amount required to move an electron from its present energy level to the next higher one.

  14. SCIENTISTMODEL Erwin Schrodinger Quantum Mechanic Model (1887-1961) • Austrian(1926): Wave mechanics-mathematical • Probable location of electron • Cloud Shaped • Pattern similar to a Propeller blade Louis De Broglie Quantum Mechanic Model (1892-1987) • French (1929): Wave nature of electrons discovery • Electrons can act like waves (Quantum Physics) • Wave particle duality/High energy emission in elements—Synthetic elements

  15. Subatomic particles: Electrons, Protons, & Neutrons • Atomic Number: Number of Protons in the nucleus • Whole number written above chemical symbol Ex: Hydrogen=1(P) Oxygen=8(P) • Atomic Mass #: Sum of Protons + Neutrons Ex: Carbon Mass #12 = 6(P) + 6(N) Oxygen Mass #16 = 8(P) + 8(N) A.Mass # (#P + #N) - Atomic # (#P) =#Neutrons

  16. 12 6 Atomic Number (P) 6 C Carbon 12 Element Symbol LETS HAVE SOME PRACTICE Element Name Mass Number (P+N) Mass Number (P+N) C Atomic Number (P)

  17. WHAT GIVES AN ATOM ITS IDENTITY? • Isotope: Same # of Protons, different # of Neutrons • Different Mass Number • Same Atomic Number • Chemically alike Ex: Carbon-12 Mass #12 = 6(P) + 6(N) Carbon-13 Mass #13 = 6(P) + 7(N) Atomic Mass for isotopes of Carbon = 12.01 amu

  18. SO, WHAT GIVES AN ATOM ITS IDENTITY? • # of protons gives the atom its identity • # of electrons determines the chemistry of the atom • # of neutrons only changes the mass of the atom

  19. DO NOW There are 3 isotopes for Oxygen: O-16 O-17 O-18 (SYMBOL-MASS NUMBER) • Write the shorthand chemical symbol for all three isotopes C 12 6

  20. 16 8 O-16 O O-17 O O-18 O 17 8 18 8

  21. Average Atomic Mass 80% tests ---50 20% homework---100 What is your average? (50+100)/2=75, not the case, tests are weighted more .80x50= 40 .20x100=20 60 is your grade

  22. Average Atomic Mass Two isotopes of carbon are C-12 the abundance is 98.89% C-13 the abundance is 1.11% What is the average atomic mass (12+13)/2=12.5 C-12: (98.89%/100) x 12=11.87 C-13: (1.11%/100) x 13=0.14 11.87+0.14=12.01amu

  23. Average Atomic Mass • Do questions #23 & #24 on page 117

  24. DO NOW Determine the # of protons, neutrons, & electrons for the 4 isotopes of zinc: 64 30 66 30 67 30 70 30

  25. DO NOW Element X has two natural isotopes with mass 10.012 amu and a relative abundance of 19.91%. The isotope with mass 11.009 amu has a relative abundance of 80.09%. 1. Calculate the atomic mass of this element (show all work) and then name this element.

  26. Nuclear Chemistry • The study of changes in matter that originate in atomic nuclei • What makes a nucleus unstable? • Too many or Too few neutrons relative to the # of protons • The nuclei of unstable isotopes gain stability by undergoing changes

  27. Changes that Radioactive Isotopes Undergo • Alpha α particle • Release of helium nuclei • Rutherford’s Gold Foil Exp. • Beta β particle • Release of an electronfrom the breaking apart of a neutron in an atom • Gamma γ ray • Release of photons (light energy)

  28. What can they penetrate? • Alpha α =almost nothing • Inhalation (radon) • Open wounds • Can’t go through skin, paper, wood, plastic, lead, concrete • Beta β = somethings • Skin & paper • Can’t go through wood, plastic, lead, concrete • Gamma γ = a lot of stuff • Skin, paper, & wood • Can’t go through lead or concrete http://www.furryelephant.com/player.php?subject=physics&jumpTo=re/2Ms4

  29. Half-Life • Unstable isotopes have a rate of decay, known as half-life

  30. Uses of Radioactive Isotopes • Carbon dating (pg. 806, 814,815) (pg 883) • Geiger counter (pg 817) (pg 895) • Film Badge (pg 817) (pg 895) • Agriculture tracers (pg 818) (pg 896) • Treating Cancer (pg 819) (pg 897) • New textbook pages

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