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People, Experiments, Conclusions

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Democritus – Ancient Greeks (5th century B.C.) believed in a small, indivisible particle – “atomos” Dalton – Modern theory (1803)– indivisible atom J.J. Thomson – Plum Pudding model (1897); Used Cathode Ray tube to experiment; Atoms have identical, removable (- ) charged particles. hill.com/olcweb/cgi/pluginpop.cgi?it=swf::100%25::100%25 ::/sites/dl/free/ /117354/01_Cathode_Ray_Tube.s wf::Cathode%20Ray%20Tube

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Millikan (1909) – Oil Drop experiment to determine the charge on electron and mass of electron. Ernest Rutherford (1911) – Gold Foil Experiment – Discovered nucleus as dense (+) charged center. Niels Bohr (1913) – Electrons in orbit around the nucleus – based on work with Hydrogen. html

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Erwin Schrodinger (1926) – Quantum theory – treats electrons as waves. Quantized energy levels. No exact “path” for electrons – just a probability of location for the electrons in an area. Electron Cloud – picture of probability Surface of cloud gives approximate shape and 90% chance of e - location (Fig 13.2 in text)

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Quantum “numbers” - describe electrons Principal Quantum Numbers: n = 1, 2, 3, 4, 5 n = principal energy level Angular momentum Quantum Number In “English” – sublevel within each energy level. Each sublevel has a different shape. Current sublevels: s, p, d, f orbitals no g orbitals no g orbitals with g orbitals with g

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Magnetic Quantum Number Deals with distribution inside the sublevel. Orientation along the axes. Spin Quantum Number - +1/2 or -1/2

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Electron distribution Max # of e- in any given energy level is = 2n 2. (n=energy level) Level Max e-2818 Sublevels sss Availablepp d

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Orbitals – areas within the sublevel where electrons can “hang out” Each orbital can hold 2 electrons Sublevel# of orbitals# e- s1 2 p3 6 d5 10 f7 14

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Electron Configurations Sample: N = Say: one “s” two, two “s” two, two “p” three

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More Examples - C – F – Li – Mg P

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