Atomic structure – part ii

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Atomic structure – part ii

  1. 1. Atomic Structure – Part II<br />Signor Rinno D. Montales<br />
  2. 2. Ernest Rutherford (1871-1937)<br />British physicist, believed in the plum pudding model of the atom.<br />With his associate Hans Geiger(1882-1945) a German physicist, studied the alpha particles emitted by radium which was isolated by Marie and Pierre Curie.<br />Alpha particles are found to be helium atoms with their electrons removed, positively charged and mass of 2500 times that of the electron.<br />
  3. 3. Gold foil<br />a–Particle<br />emitter<br />Slit<br />Detecting screen<br />Rutherford’s Experimental Design<br />(a)<br />
  4. 4. Gold foil<br />a–Particle<br />emitter<br />Slit<br />Detecting screen<br />(a)<br />
  5. 5. Gold foil<br />a–Particle<br />emitter<br />Slit<br />Detecting screen<br />(a)<br />
  6. 6. Florescent <br />Screen<br />Lead block<br />Uranium<br />Gold Foil<br />
  7. 7. Rutherford’s Hypotheses<br />The alpha particles to pass through without changing direction very much<br />Because most of the mass of the atom (positive charges) were spread. Alone they were not enough to stop the alpha particles<br />If the Thomson model were correct, all the alpha particles, travelling at high speeds and massive, would have passed through the metal foil undeflected or only slightly deflected<br />
  8. 8. What he expected<br />
  9. 9. Because<br />
  10. 10. Because, he thought the mass was evenly distributed in the atom<br />
  11. 11. Because, he thought the mass was evenly distributed in the atom<br />
  12. 12. What he got<br />They observed that although majority of the alpha particles passed through undeflected, some were only slightly deflected, some were scattered by more than 90 degrees and a few by nearly 180 degrees or almost completely turned back <br />
  13. 13. +<br />How he explained it<br />Atom consists of a very small nucleus <br />surrounded by electrons. Rutherford<br />estimated the radius at 10-12 to <br />10-13 cm compared to radius of <br />the atom of about 10-8 cm<br />The nucleus contains most of <br />the mass of the atom and all of <br />its positive charge.<br />Alpha particles are deflected by<br />nucleus it if they get close enough at each other<br />
  14. 14. +<br />
  15. 15. Ernest Rutherford’s Model<br />(Nuclear Model of an Atom)<br />
  16. 16. Bohr’s Model<br />Electrons move in circular orbits around the nucleus<br />Adopted Planck’s idea that energies are quantized<br />
  17. 17. Three Phenomena<br />Black-body radiation<br />Photoelectric effect<br />Emission spectra <br />
  18. 18. ELECTROMAGNETIC SPECTRUM<br />The waves in the spectrum all travel at same speed through a vacuum but differ in the frequency and, therefore, wavelength. <br />
  19. 19.
  20. 20. LINE SPECTRUM<br />Spectrum containing radiation of specific wavelengths <br />
  21. 21.
  22. 22. Energy states of a Hydrogen Atom<br />When a sample of gaseous H atoms is excited, different atoms absorb different quantities of energy<br />Each atom has one electron, but so many atoms are present that all the energy levels (orbit) are populated by electrons<br />Ground state – lowest energy level (n = 1)<br />Excited state - higher energy level (n= 2…) <br />
  23. 23. When dropped from n = 3 orbit (second excited state) = infrared series or lines were emitted by photons –PASCHEN SERIES<br />When dropped from n = 2 orbit (first excited state) = visible series or lines – BALMER SERIES<br />When dropped from n = 1 orbit (ground state) = ultraviolet series or lines – LYMAN SERIES<br />
  24. 24.
  25. 25. An electron could jump from one allowed energy state to another by emitting or absorbing photons whose energy corresponds exactly to the energy difference between the two states. <br />ΔE = Ef– Ei<br />Line spectra are produced because these energy changes correspond to photons of specific wavelengths<br />
  26. 26.
  27. 27. The PARTICLE Nature of Light <br />Blackbody Radiation<br />Light given off by hot objects<br /> Wavelength distribution of the radiation depends on temperature<br />“red-hot” object being cooler than a “white-hot” one<br />
  28. 28.
  29. 29. MAX PLANCK ( 1858-1947)<br />Energy can be released or absorbed by atoms only in discrete “chunks” of some minimum size.<br />Quantum – “fixed amount”, smallest amount of energy that can be mitted or absorbed as electromagnetic radiation.<br />
  30. 30. Hot glowing object could emit (or absorb) only certain quantities of energy<br />E = hv<br />E = energy of radiation<br />v= frequency<br />h= Planck’s constant ( 6.63 x 10-34 joule-seconds) <br />
  31. 31. Hot object’s radiation is emitted by the atoms contained within it.<br />The atom itself can have only certain quantities of energy.<br />The energy is quantized- values are restricted only in certain quantities <br />
  32. 32.
  33. 33.
  34. 34.
  35. 35.
  36. 36. CONTINUOUS<br />QUANTIZED<br />
  37. 37. Photoelectric Effect<br />Emission of electrons from metal surfaces on which light shines<br />
  38. 38. ALBERT EINSTEIN (1905)<br />Used Planck’s quantum theory to explain the photoelectric effect<br />
  39. 39. Radiant energy striking the metal surface is a stream of energy packets<br />PHOTON<br />Behaves like a particle<br />Has an energy proportional to <br />Energy of photon= E = hv<br />
  40. 40. A photon transfers its energy to an electron in the metal.<br />A certain amount of energy is required to overcome the attractive forces that hold it within the metal. <br />
  41. 41.
  42. 42. Three postulates of Bohr<br />Only orbits of certain radii, corresponding to certain energies, are permitted for electrons in an atom called STATIONARY STATES. – fixed circular orbit<br />An electron in a permitted orbit has a specific energy and is in an “allowed” energy state. Electron will not radiate energy while moving within an orbit.<br />Energy is only emitted or absorbed by an electron as it changes from one allowed energy state to another. This energy is emitted or absorbed as a photon.<br />Ephoton = Estate A – Estate B = hv<br />

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