1. Lab 9 Atomic Structure Emission Spectrum Electron Configuration

2. HISTORY OF THE ATOM 460 BC Democritus develops the idea of atoms he pounded up materials in his mortar and pestle until he had reduced them to smaller and smaller particles which he called ATOMS ( greek for indivisible )

4. HISTORY OF THE ATOM 1808 John Dalton suggested that all matter was made up of tiny spheres that were able to bounce around with perfect elasticity and called them ATOMS

6. HISTORY OF THE ATOM 1898 Joseph John Thompson found that atoms could sometimes eject a far smaller negative particle which he called an ELECTRON

7. Thomson and the Discovery of Electrons

9. J. J. Thomson’s Experiment Devised an experiment to find the ratio of the cathode ray particle’s mass ( m e ) to the charge ( e ) m e / e = –5.686 x 10 –12 kg C –1

11. HISTORY OF THE ATOM Thomson develops the idea that an atom was made up of electrons scattered unevenly within an elastic sphere surrounded by a soup of positive charge to balance the electron's charge 1904 PLUM PUDDING MODEL

12. Millikan’s Oil Drop Experiment Measuring the Charge on an electron

14. Radioactivity

15. Goldstein’s Discovery of Protons

16. Mass Spectrometer- Determining the Percent Abundance of Different Isotopes of same element

17. Mass Spectrometer If a stream of positive ions having equal velocities is brought into a magnetic field, the lightest ions are deflected the most, making a tighter circle

18. Mass Spectrometry EOS A record of the separation of ions is called a mass spectrum

21. HISTORY OF THE ATOM 1910 Ernest Rutherford oversaw Geiger and Marsden carrying out his famous experiment. They fired Helium nuclei at a piece of gold foil which was only a few atoms thick. They found that although most of them passed through. About 1 in 10,000 hit

22. Rutherford’s Gold Foil Experiment

23. Rutherford’s Gold Foil Experiment gold foil helium nuclei They found that while most of the helium nuclei passed through the foil, a small number were deflected and, to their surprise, some helium nuclei bounced straight back.

25. Visualizing the Pathway of Alpha Particles through a Gold Atom

26. ELECTROMAGNETIC RADIATION

27. Electromagnetic radiation.

29. Electromagnetic Radiation

30. The Electromagnetic Spectrum

33. Electromagnetic Radiation wavelength Visible light wavelength Ultaviolet radiation Amplitude Node

34. Electro magnetic Spectrum In increasing energy, R O Y G B I V

36. Sunlight viewed through a spectroscope

38. Photoelectric Effect

41. Light Spectrum Lab! Slit that allows light inside Line up the slit so that it is parallel with the spectrum tube (light bulb)

42. The Emission Spectrum of Hydrogen- Discrete Bands of Colored Light

43. Excited Gases & Atomic Structure

44. Emission Spectra of Different Atoms: A Fingerprint to Identify

49. Bohr’s Calculations of the Energy Δ E = -2.18 x 10 -18 J (1/n f 2 – 1/n i 2 ) n = the energy level Δ E = positive when electron climbs up levels absorbing energy increasing PE Δ E = negative when e- falls down levels releasing energy decreasing PE

50. Niels Bohr (1885-1962) Δ E = -2.18 x 10 -18 J (1/n f 2 – 1/n i 2 ) Calculate the energy as an electron drops from level 6 down to level 2. Calculate the frequency and wavelength of this photon.

51. ultraviolet infrared

52. Visualizing the Movement of the Electron

54. Line Spectra of Other Elements Oops. Bohr’s equation does NOT predict these wavelengths.

56. Visualizing the “falling” e- Where does the electron have more potential energy?

59. E. Schrodinger 1887-1961 W. Heisenberg 1901-1976 Wave Functions: Calculating the Probability of locating an electron in a region of space The Uncertainty Principle: Cannot both determine location and energy of electron

60. The Wave Function and Orbitals

61. The region near the nucleus is separated from the outer region by a spherical node - a spherical shell in which the electron probability is zero EOS

64. Increasing Radius of s-orbital with higher values of n

65. s orbital p orbital d orbital

68. f Orbitals

72. The s-orbital Spherical shaped orbital ℓ = 0 m ℓ = 0 Only one s-orbital in any energy level

73. The p-orbital Double-lobe shaped orbital ℓ = 1 mℓ = -1 or 0 or +1 Only three p-orbitals in any energy level- except for level one

74. Planes of zero probability

75. Model of d-orbital

76. Only electrons with opposite spins can be in the same orbital

78. Use sum of first two quantum numbers to determine which orbital fills first

90. Nickel Electron Configuration and quantum numbers 1s 2s 2p 3s 3p 1 0 0 ½ 2 0 0 ½ 2 1 -1 ½ 2 1 0 ½ 3 0 0 ½ 3 1 0 ½ 2 1 0 ½ 3 1 -1 ½ 3 1 0 ½ 4s 3d 4 01 0 ½ 3 2 -2 ½ 3 2 -1 ½ 3 2 0 ½ 3 2 1 ½ 3 2 1 ½ 1 st # indicates energy level n = 1 1 st level n = 2 2 nd level 2 nd # type of orbital l = 0 is s-orbital l = 1 is p-orbital l = 2 is d-orbital