The Wave Nature of Light: Electromagnetic Radiation

1. 7.1  The Wave Nature of Light • Electromagnetic Radiation 1. Wavelength λ 2. Frequency ν a. Hertz = cycles per second Hz b. f or ν 3. c = (wavelength sign)λ x f 4. c = 3.00 x 108 m/s a. Alias: Speed of light 5. 1m = 109 nm

2. Electromagnetic Spectrum

3. Bees see in UV Silverweed (Potentilla anserina): It is hard to imagine that these yellow flowers are actually hiding a two-tone pattern, as revealed in the ultraviolet image

8. 2 slit diffraction pattern

9. 7.2  Quantized NRG and Photons • Quantum = smallest quantity of NRG that can be emitted or absorbed • Planck’s constant = 6.63 x 10 -34 J-s = h 1. E = hν NRG is always emitted or absorbed in mult. of hv, = 2hv, 3hv, 2452345hv…

10. 6.2… What? Continued? YEAH. • In 1905 Albert Einstein decided to shine something on a metal surface, since everyone gets that urge at some point in their lives, and it emitted photons! Whoooo, photons! • Photon- tiny particle of light. Just like this letter!

11. 6.3  Bohr’s Model In 1913 Niels Bohr made a planetary model. But more importantly, he was DANISH.

12. 6.3 For cereal, this time. Line Spectra • Monochromatic- one color • Spectrum • Continuous spectrum – Rainbow! Which, incidentally, are always 60 degrees at the base. • Line spectra- a spectrum containing only specific wavelengths • Balmer, 1885, Swiss, Not nearly as cool as a Dane

13. ν = C{ (1/ ni2) – (1/nf2) } C= 3.29 x 1015 1/s or Hz ∆E = -2.178 x 10 -18J (1/ ni2) – (1/nf2)

14. And now back to your regularly scheduled programming F. Bohr Model p. 306 1. En = (-Rh)( 1/n2) n=1, 2, 3 … 2. Rh = Ryberg constant = 2.18 x 10 –18 J G. Ground state – lowest NRG state H. Excited state – anything higher I. ν = E/h = Rh/h { (1/ni2) – (1/nf2) }

15. 7.4 Wave Behavior of Matter • DeBroglie Equation 1. Wavelength = h/mv B. Mass x velocity = momentum, p C. Matter waves- wave characteristics of material particles D. Heisenberg’s Uncertainty Principle- impossible to know both location and momentum (velocity) of electron at the same time

16. 7.5 • Quantum Mechanics – both particle and wave properties of electrons • Schrodinger 1926 • Wave functions, psi ( trident thinger…) • (trident thinger)2 probability electron location = probability density • Electron density- regions where there is a high probability of finding an electron

17. II. Quantum Numbers  n, l, m, s n = principle quantum number (shell) n = 1, 2, 3, 4… 7 l = azimuthal quantum number (subshell) s = 0, p = 1, d = 2, f = 3 m = magnetic quantum number (orbital) aka me m= -1 … 0 … +1 s = spin (ms) s= + ½ , - ½

18. 7.7 Orbitals • s ARE ROUND 315 1. Nodes – area where trident2 approaches zero B. p are 2 lobes p315 C. d orbitals on page 316 D. f are on p317

20. s p d f

21. 6.7 Orbitals in many electron atoms • Effective nuclear charge – net positive charge attracting the electron • Screening effect – inner electrons are shielded • degenerate – orbitals with same NRG II. Pauli’s Exclusion Principle – no 2 electrons in an atom can have the same set of quantum numbers. Everyone is individual! III. Hund’s Rule – electrons fill orbitals one at a time

22. 7.8 Electronic Configuration • Exceptions Cr, Mo, Cu, Ag, and Au (pg. 327) • Valence electrons • Core electrons II. Orbital Filling diagrams III. Sections on the chart s p d f

23. The End! (this powerpoint was made by Tessa but edited by the infinitely smarter Lilly, because there were I’s in the equations. TAKE THAT, TESSA!) Enjoy the rest of your day.

24. Walter White form “Breaking Bad”