2. Energy and Electrons Rev 11/05/08 Part One Pisgah High School M. Jones

3. Goals for this unit – Know about: 1.the lines in the hydrogen spectrum, and Bohr’s atomic theory, 2.the arrangement of electrons in atoms, and the shape of the periodic table, 3.energy diagrams and electron configurations, 4.valence electrons and dot diagrams.

4. Light and Energy

5. Some background terms and concepts

6. Frequency Symbol: f or f (or  – Greek: nu) Units: Hertz (Hz) or 1/sec Pronounced “reciprocal seconds” Historically, frequency was in units of “cycles per second”, but this made too much sense.

7. Wavelength Symbol:  Greek: lambda) Units: meters (m)  of radio waves in meters to micrometers.  of light in nanometers (10 -9 m). Wavelengths of light can be measured with a spectroscope.

8. The sine curve is often used to represent wave motion. Look at the following graphs. What is the relationship between frequency and wavelength?

9. O n e w a v e l e n g t h

12. One wave- length

13. What is the relationship between frequency and wavelength? As the frequency increases, the wavelength decreases.

14. The relationship between frequency and wavelength can be represented by: f = frequency (lambda) = wavelength

15. Frequency is inversely proportional to wavelength The relationship between frequency and wavelength can be represented by:

16. Frequency times wavelength equals a constant k = f

17. For electromagnetic energy, the equation is: c is the speed of light c = 3.00 x 10 8 m/sec c = f

18. The equation can also be written as: c = speed of light  = wavelength  = frequency c =

19. Light is part of the electromagnetic spectrum Refer to your reference tables.

20. Electromagnetic Spectrum Longer wavelength Lower frequency Lower energy Shorter wavelength Higher frequency Higher energy

21. Electromagnetic Spectrum Longer wavelength Lower frequency Lower energy Shorter wavelength Higher frequency Higher energy 400 nm 700 nm

22. The visible spectrum was discovered by … Dr. Roy G. Biv Red Orange Yellow Green Blue Indigo Violet 700 nm Lower energy 400 nm Higher energy

23. Electromagnetic waves carry energy.

24. …inversely proportional to the wavelength. …directly proportional to the frequency, The energy in an electromagnetic wave is …

25. Energy in an EM wave E =E = hc E =E = hc Small wavelength – large energy Large wavelength – small energy

26. Next, a demonstration … Look at sunlight and fluorescent light through the spectroscope. Use the spectroscope to look at the light coming from various gas discharge tubes. Record the colors and their order in the spectrum of hydrogen.

27. Stop Complete the observation of atomic spectra, then continue.

28. What did you see? Hydrogen Spectrum

29. And now something completely different, some history.

30. Thomson suggested the plum pudding model with many, many electrons throughout the atom. J. J. Thomson discovered the electron in 1897.

31. Thomson suggested the plum pudding model with many, many electrons throughout the atom. J. J. Thomson discovered the electron in 1897.

32. Rutherford suggested in 1911 that electrons might exist outside the nucleus in a “planetary” arrangement. Ernest Rutherford explained that atoms had a small, dense, positively nucleus. Hantaro Nagaoka Japan, 1904

33. Enter Niels Bohr Bohr also knew about Rutherford’s “planetary” model of the electrons. Bohr was there right after the “gold foil” explanation was published. Niels Bohr, with a brand new PhD in physics from U. Copenhagen, went first to Thompson, then to Rutherford to study in 1912.

34. Niels Bohr In 1913 he proposed a structure of the atom with Rutherford’s nucleus and electrons in discrete energy levels. While at Manchester, Niels Bohr studied the spectra of elements, and how these might relate to the internal structure of atoms.

35. Niels Bohr He used this new “quantum theory” to help explain how energy could be absorbed and emitted. Niels Bohr knew of the work done by Max Planck and incorporated it into his atomic theory.

36. Max Planck In 1900 Max Planck introduced an unusual idea. Energy exists in “packets”, or quanta. The beginnings of Quantum Theory. A quantum of energy is the smallest amount of energy possible. Energy exists only in multiples of these quanta.

37. Cadmium Sulfide Albert Einstein In 1905 Einstein used Max Planck’s idea of quanta to explain the photoelectric effect. Photon of light Electron knocked loose Gave credibility to quantum theory.

38. Albert Einstein A photon of just the right energy can knock an electron out of an atom. A photon is a “packet of light”, or quantum of energy. Einstein explained with the Quantum Theory what could not be explained with classical physics.

39. Niels Bohr … that electrons can only change energy levels when they absorb or give off a certain amount of energy. (1913) Bohr said that electrons could exist only in certain discrete energy levels, and …

40. Hydrogen atom nucleus Discrete energy levels for electrons electron

41. Hydrogen atom Electrons can exist at this level,

42. Hydrogen atom …or in this level,

43. Hydrogen atom …or in this level,

44. Hydrogen atom …but not in between the levels.

45. Hydrogen atom Unless the electron is absorbing energy, or … Giving off energy

46. Niels Bohr Why ??? When observed through a diffraction grating, specific lines of color are observed. When high voltage is connected to the hydrogen discharge tube, a bluish light is given off.

47. Niels Bohr … the electron moves up to a higher energy level. The electron in a hydrogen atom gains energy from the electricity passing thru the tube and …

48. Niels Bohr The electron in the excited state is “unstable”. … gives off light of a certain energy and wavelength. The electron drops to a lower energy level, and …

49. Don’t forget: E = hc = hc E and The wavelength is inversely proportional to the energy

50. Hydrogen atom

51. energy The electron absorbs energy and …

52. Hydrogen atom …the electron is elevated to the next energy level

53. The electron is unstable and “wants” to return to a lower energy level Hydrogen atom

54. Light of a particular wavelength is given off

55. Each line has a wavelength and color that corresponds to the difference in energy between the two levels. A line in the hydrogen spectrum is produced when an electron moves from higher energy level to a lower one.

56. Don’t forget: E = hc = hc E and The wavelength is inversely proportional to the energy

57. Energy and E1E1 E2E2 E1E1 E2E2  E is large  E is smallLonger  Shorter 

58. 5432154321 Hydrogen atom The energy levels are numbered …

59. 5432154321 Suppose an electron is in level 5 and drops to level 2. Hydrogen atom The energy levels are numbered …

60. 5432154321 Then purple light with a wavelength of 434 nm will be emitted. Hydrogen atom The energy levels are numbered …

61. 5432154321 Hydrogen atom Suppose an electron is in level 4 and drops to level 2. The energy levels are numbered …

62. 5432154321 Hydrogen atom Then you get blue-green light with a wavelength of 486 nm The energy levels are numbered …

63. 5432154321 Hydrogen atom Suppose an electron is in level 3 and drops to level 2. The energy levels are numbered …

64. 5432154321 Hydrogen atom Then red light with a wavelength of 656 nm will be emitted. The energy levels are numbered …

65. 5432154321 The colors of the visible lines come from the energy given off by electrons moving from higher energy levels down to level 2. Hydrogen atom The energy levels are numbered …

66. Don’t forget: E = hc = hc E and The wavelength is inversely proportional to the energy

67. The Hydrogen Spectrum Each color represents the transitions of gazillions of electrons in gazillions of H atoms going from higher energy levels to the second energy level.

68. Why can’t we see them? There can be transitions among seven energy levels. There must be a lot more lines in the spectrum of hydrogen.

69. Transitions in the H-spectrum 76543217654321 Transitions to the first energy level produce ultraviolet lines.

70. Transitions in the H-spectrum 76543217654321 Transitions to the second energy level produce visible lines.

71. Transitions in the H-spectrum 76543217654321 Transitions to the third and higher energy levels produce infrared lines.

72. Bohr successfully calculated the wavelengths of all the transitions in the hydrogen spectrum.

73. But only the hydrogen spectrum. The spectra of elements with more than one electron could not be accurately predicted.

74. “Regardless of it’s shortcomings and the modifications that were later applied, Bohr’s model of the atom was the first successful attempt to make the internal structure of the atom agree with spectroscopic data.” Asimov, 1964

75. The Arrangement of Electrons in the Quantum Mechanical Model of the Atom

76. The Modern View of the Atom: 1.A small, dense positively charged nucleus which contains protons and neutrons.

77. The Modern View of the Atom: 2.Electrons which exist outside of the nucleus at … 1.various distances from the nucleus, and at … 2. various energy levels.

78. The Electrons 3.The electrons can have both a mass, as does matter, and a wavelength, as does light energy.

79. The Electrons 4.The electrons themselves are not little solid spheres in orbit around the nucleus, but exist as a “fog” of half-energy, half- matter. The electrons can behave as either matter or energy, depending on the experiment.

80. Energy Levels 5.Based on the ideas of Bohr, the electrons are located … a)… in major energy levels, b)… in energy sublevels within major energy levels, c)… in orbitals within each sublevel.

81. The energy levels are like an organizational chart for a business:

82. Quantum Numbers 1.Each electron in an atom has a set of four (4) quantum numbers. 2.The quantum numbers are like an address: name, street, city, state. 3.Pauli exclusion principle: no two electrons in the same atom can have the same set of quantum numbers.

83. Quantum Numbers AddressQuntm. #Sym.What it tells StatePrincipal nMajor energy level CityAzmuthal LSublevel StreetMagnetic M L Orbital NameSpin M S Which e - in orbital Luckily, we will only deal with the Principal Quantum Number

84. What’s coming next? 1.2n 2, and the shape of the periodic table 2.Energy levels and sublevels 3.s, p, d, f and the periodic table 4.Orbitals, spin & energy diagrams 5.e - config., valence e -, dot diagrams Click here to go to Part Two.