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Unit 2:

Unit 2:. Atomic Theory. Video links. overview of atomic history http://www.youtube.com/watch?feature=player_detailpage&v=k1RHY8QcN1s. I. Atomic History. A. The Greeks Democritus [Philosopher] All matter is made of tiny, indivisible parts called ‘atoms’

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Unit 2:

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  1. Unit 2: Atomic Theory

  2. Video links • overview of atomic history • http://www.youtube.com/watch?feature=player_detailpage&v=k1RHY8QcN1s

  3. I. Atomic History A. The Greeks Democritus [Philosopher] • All matter is made of tiny, indivisible parts called ‘atoms’ • Developed word ‘atomos’ meaning not divisible

  4. John Dalton (1803-1808) • Used experiments with gases to develop the “Atomic Theory” • Determined atoms looked like ‘cannonballs’ or solid masses

  5. Dalton’s Atomic Theory • 1) All elements are made of atoms • 2) Atoms of each element are all the same, or have the same masses • 3) Atoms of different elements are different, or have different masses • 4) Atoms cannot be created or destroyed • 5) Atoms combine in small, whole number ratios

  6. J.J. Thomson (1897) • Developed ‘Cathode Ray’ experiment • Said atoms consisted of particles smaller than an entire atom • Discovered that the smaller particles within an atom had a negative charge • Discovered 1st subatomic particle: Electron • Founded “Plum Pudding Model”: Electrons were embedded within a positively charged mass

  7. Cathode Ray Tube Experiment Thomson manipulated cathode rays with a magnet to discover that subatomic particles existed and that they had negative charges

  8. Discovered alpha and beta radiation emitted from certain radioactive substances Developed and used Gold Foil Experiment First to separate the smaller parts of the atom Discovered the nucleus Placed electrons outside the nucleus Stated that atoms are composed of lots of empty space Ernest Rutherford (1898)

  9. Rutherford’s Gold Foil Experiment micro.magnet.fsu.edu/electromag/java/rutherford/

  10. Niels Bohr (1922) • Bohr analyzed work of others and studied atomic spectra, or light, given off by the elements • Described the “Atomic Spectra” of elements • Developed ‘Solar System’ model • Moved electrons from single, giant pathway into discrete energy levels around the nucleus • Each energy level contained 2, 8, 18, 32, etc. electrons total

  11. Bohr Model of the Atom • Stated that electrons moved around the nucleus in ‘orbits’ or energy levels • As electrons gain energy, they jump up energy levels, then release this energy to generate spectra

  12. Bohr’s model and the atomic spectrumhttp://jersey.uoregon.edu/vlab/elements/Elements.html • The spectral lines in the visible region of the atomic emission spectrum of barium are shown below. • Spectral lines exist in series in the different regions (infra-red, visible and ultra-violet) of the spectrum of electromagnetic radiation. • The spectral lines in a series get closer together with increasing frequency. • Each element has its own unique atomic emission spectrum.

  13. Erwin Schrodinger (1930) • Developed mathematical equations representing electrons • Electrons had wave and particle behaviors • Created “Wave-Mechanical” or “Modern” model • Most scientists use this model today • Placed electrons in orbitals

  14. “Electron Cloud” Model • Created paths for electrons within Bohr’s energy levels • Only 2 electrons per path • Electron paths, or ORBITALS, are mathematical equations describing probability densities for electrons • Developed sublevels with discrete paths within each energy level

  15. II. Subatomic Particles • Particles 1) Protons • found in the nucleus of an atom • charge of +1, mass of 1.0073a.m.u. 2) Neutrons • found in the nucleus of an atom • no charge, mass of 1.0087a.m.u.

  16. A. Subatomic Particles 3) Electrons • Found outside the nucleus in regions of probability [orbitals] • Charge of –1, mass of 5.46 x 10-4 a.m.u., or 1/1836 a.m.u. • Have particle and wave properties

  17. Atomic number = the number of protons in the nucleus All atoms of the same element have the same atomic number Atoms arranged on PT by increasing atomic numbers In neutral atoms: Atomic number equals number of electrons B. Atomic Number

  18. C. Isotopes • Isotopes = atoms of the same element that have differing numbers of neutrons in their nucleus, different mass number, but same atomic number • Same number of protons!!! • Changing number of neutrons affects properties [radioactivity…]

  19. Atomic Mass Number = number of protons plus the number of neutrons in the nucleus Whole number!! Mass Number changes when using different isotopes Written in isotopic notations, just subtract the top from bottom values: D. Atomic Mass

  20. E. Ions • Ions = atoms of the same element that have lost or gained electrons • Have overall (+) or (-) charge • Same numbers of protons, number of neutrons irrelevant • Positive ions: have LOST electrons • Negative ions: have GAINED electrons

  21. F. Atomic Mass (average) • Atomic Mass = weighted average of the natural isotopes times their percent abundance • Decimal value on PT • Accounts for the natural existence of various isotopes • Ex] calculate the atomic mass of carbon given that 98.92% is carbon-12 and 1.108% is carbon-13

  22. Virtual textbook http://www.chem1.com/acad/webtext/intro/int-1.html#SEC1

  23. A. EMS [Electromagnetic Spectrum] III. Electronic Structure

  24. A. EMS [Electromagnetic Spectrum] • EMS = continuous series of various types of energy, separated by their wavelengths and frequencies • Visible light = small portion; only part we can see without instruments • Continuous spectrum = picture of all colors of visible light as they pass through a prism

  25. Wavelength = distance between 2 peaks or troughs of 2 consecutive waves Symbol = λ [Greek letter “lambda”] Units are usually in ‘m’ or ‘nm’ Frequency = the number of peaks or troughs that pass a single point in one second Symbol = ʋ [Greek letter “nu”] Units are usually in ‘1/s’ or ‘s-1’ or ‘Hz’ EMS continued

  26. c = λν C = speed of light C = 3.0 x 10+8 m/s E = hν E = energy of photon h = Planck’s constant h = 6.63 x 10-34 Js All electromagnetic radiation travels at the speed of light Can calculate the energy of the radiation/electron given the wavelength Calculations using lambda and nu

  27. Planck’s Constant • Planck observed hot, glowing matter • Concluded: different substances glow different colors at different temperatures • Determined: matter releases energy in tiny, discrete packets called ‘quanta’ • Developed constant to relate energy and temperature, Planck’s constant, “h” h = 6.63 x 10-34 J*s

  28. Light traveling as waves All colors of light energy travel at the same speed, just different wavelengths!

  29. Particle vs. Wave Behavior of Light

  30. Wave behavior of light

  31. B. Photoelectric Effect • Einstein used Planck’s idea of quanta and photons to describe the photoelectric effect • Light of a certain wavelength shines on clean metal, causing the metal to eject electrons

  32. Bohr’s Model [conclusions made] • Bohr used the idea of ‘quanta’ to explain the bright-line emission spectra • Stated that each element’s atomic spectrum is unique • Electrons exist in ground state energy levels, as listed via the periodic table

  33. Bohr Model of the Atom • Stated that electrons moved around the nucleus in energy levels • Electrons will gain and lose energy at will • This generated the element’s atomic spectrum

  34. Bohr’s model

  35. Useful Websites and References //www.avogadro.co.uk/light/bohr/spectra.htm • shows formation of spectral lines for hydrogen • idea of ground vs. excited state //jersey.uoregon.edu/vlab/elements/Elements.html • Periodic table showing the absorption and emission spectra for each element Also check out Wikipedia under Bohr atom and Atomic spectra!

  36. Creation of an emission spectrum • If electrons absorb packets of energy, quanta, they temporarily move to into a higher energy level, called the excited state • The electrons then release this quanta of energy and fall back down to ground state • The release of energy generates the bright-line emission spectrum

  37. Examples of Bohr Diagrams

  38. IV. Electron Configurations • Energy Levels • These are areas with a high possibility of finding electrons with similar potential energies • 7 energy levels total

  39. Bohr Diagrams show the numbers of protons and neutrons in the nucleus Shows electrons in their respective energy levels Energy levels hold: 1st holds 2 electrons 2nd holds 8 electrons 3rd holds 18 electrons 4th holds 32 electrons Etc….. Bohr Diagrams and Energy Levels

  40. Sublevels are divisions within each energy level Represent the shapes and orientation in 3D space Too many electrons within the energy levels & they lose momentum and will crash into the nucleus--- not good! 1st energy level has 1 sublevel: “s” 2nd has 2 sublevels: “s” and “p” 3rd has 3 sublevels: “s, p, and d” 4th has 4 sublevels: “s, p, d, and f” B. Sublevels

  41. Sublevels and Shapes • “s” is spherical and has a max of 2 electrons • “p” is dumbbell shaped and has a max of 6 electrons • “d” is cloverleaf shaped and holds up to 10 electrons • “f” is a split cloverleaf with a max of 14 electrons • http://micro.magnet.fsu.edu/electromag/java/atomicorbitals/index.html

  42. Order of Sublevel Filling It does not go in order… 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s2 5p6 5d10 5f14 6S2 6P6 6d10 7s2 7p6

  43. Each sublevel consists of 1 to 7 orbitals [areas of probability for finding an electron] Each path or orbital only holds 2 electrons The 2 electrons within in each orbital each have a different spin This allows the electrons to exist in the same area without conflicting Orbitals within Sublevels

  44. Extended and Abbreviated Configurations • Electron Configurations = way to describe how the electrons are distributed around an atom and within the energy levels and sublevels • Ground state configurations are same order as electrons on PT • Excited state configurations have one electron shifted to a higher energy level

  45. Writing Electron Configurations • Electrons add in the same order as the atomic numbers of the PT • Aufbau Principle = adding electrons in the exact order of the PT

  46. Writing Configurations • Add in order of arrows for Neutral, Ground State atoms • Examples:

  47. Abbreviated configurations show only the placement of electrons added after the last ‘noble gas’ Ex] Bracket the configuration of the last noble gas [group 18] and add remaining electrons Ex] Abbreviated Configurations

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