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 )
3.
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
5.
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
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
10.
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
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
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
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.
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.
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
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
78. Use sum of first two quantum numbers to determine which orbital fills first
79.
80.
81.
82.
83.
84.
85.
86.
87.
88.
89.
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
