Chem 105 Fri 22 Oct 2010

Transcription

1 Chem 105 Fri 22 Oct ) Chapter 6 - Atomic Structure Hour Exam 2 will be returned Monday 10/22/2010 1

2 Mark your calendar: Hour Exam 3 (Ch 6, 7, 8) is Friday, Nov /22/2010 2

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4 Chap. 6. Atomic structure and Spectroscopy 10/22/2010 4

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6 fluorescence Energy O + e - O* + e - O* O + photon 10/22/2010 6

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8 Evidence for structure of atoms comes from their interaction with light. Spectroscopy = interaction of electromagnetic radiation with atoms and molecules. wavelength (mass of photon = 0. That s why it can go at the speed of light!) 10/22/2010 8

9 c = 3.0 x 10 8 m/sec = 186,000 miles/sec Wavelength ( lambda) units = m Frequency ( nu ) units = sec -1 1 sec -1 = 1 Hertz (Hz) 10/22/2010 9

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11 Max Planck first proposed in 1900 that light energy is quantized. That is, it exists as packets (photons) whose energy is proportional to frequency. Or E = h This conclusion first became obvious from studies on black-body radiation. ( But not by Planck himself: the history of this equation is complicated. Google ultraviolet catastrophe wiki 10/22/ Heating an iron bar to white-hot, and beyond.

12 Black-body radiation spectra show the amount of photons with different frequencies. As the bar gets hotter, the distribution goes toward the blue end of the scale. However, even at very high temps, the number of UV photons with λ < 200 nm is nil. 10/22/

13 the sharp drop in number of photons in the UV frequency range requires that there be a maximum level beyond which no excitation can occur. Heat energy promotes metal to higher quantized states. Where does the glowing light come from? As the metal atoms fall back to lower energy levels, they emit photons having frequencies α energy of quantized states 10/22/

14 The classic (incorrect) view of black body radiation assumed a continuum of energy states... which implied that as the temperature of the bar increased, the energy and frequency of the emitted photons would continue to increase indefinitely. UV 10/22/

15 E = h Frequency of photon (s -1 ) Planck s Constant = x Joule-sec Energy of a single photon (joules) Board: E, h, c,, E h hc High energy short wavelength high frequency Low energy long wavelength low frequency 10/22/

16 fluorescence E 10/22/

17 Photons from red aurora have a longer wavelength than green aurora. Where is this excited state on the energy scale? fluorescence 3 Above green Same level Below green E 1. Above green? (on the energy scale) 2. Same level? 3. Below green? 10/22/

18 What is he difference between these two processes, and why does red light originate at higher altitudes? Both these processes require the excited state of the atom to last undisturbed for seconds or even minutes, until finally the photon is released. Atoms at low altitudes suffer millions of collisions per second which deactivates excited states and prevents fluorescence. In the near-vacuum at the top of the atmosphere however, collisions are rare and fluorescence can occur. The lifetime of the lower energy excited state is about 5 min (!), and so it requires the extremely rare gas at the very top of the atmosphere to survive long enough to fluoresce. On the other hand, the lifetime of the higher energy state is about 1 sec. fluorescence E 1. Above green (on the energy scale) 2. Same level 3. Below green 10/22/

19 Tools of spectroscopy 1) Monochromator: physically separates beams of light with different wavelengths 2) Detector 10/22/

20 Electrically heated gases emit light at only a few defined wavelengths. E.g. Hydrogen atom spectrum is shown. 10/22/

21 Next week s experiment deals with various kinds of spectra Continuous Spectrum Hot Gas Emission Spectrum Cold Gas Absorption Spectrum 10/22/

22 Another early clue to quantized states of atoms were bright lines emission spectra 10/22/

23 Week-after-next experiment uses atomic absorption ( aa ) spectroscopy 10/22/

24 Flame tests NaCl KCN CuSO 4 10/22/

25 A major push in development of quantum theory came from attempts to explain emission spectra. In the Hydrogen spectrum, one set of lines is called the Balmer Series 411 nm 434 nm 486 nm 656 nm J. Rydberg noticed that this series of wavelengths could be fit approximately to a series 1 R λ n where n 2 R= Rydberg Constant = x 10 7 m -1 10/22/

26 Niels Bohr then suggested a simple Solar System Model for the H atom, which fit the Rydberg equation nicely (1913). The Bohr Model of H atom. 1) The electron is restricted to certain orbits. Therefore it cannot plunge into the nucleus its motions are quantized. 2) Assign a principal quantum number n = 1,2,3...to each orbit. 3) The higher values of n correspond to orbits with greater radii and higher potential energy. 4) Transitions between orbits are accompanied by absorption or emission of photons with defined energy. ΔE=h Based on this model, Bohr derived an equation for the energy of electron in the n th quantum level: E n Rhc 2 n 10/22/

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28 k = Rhc = 2.18 x J/atom = 1 Rydberg /22/ photon at 656 nm photon at 486 nm photon at 434 nm

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30 OWL means the absorption line in this graph, not in the emission spectrum. 10/22/

31 The End 10/22/