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3-Atomic Structure

3-Atomic Structure. Overview Characteristics of Atoms Interaction b/tw matter and light Photoelectric Effect Absorption and Emission Spectra Electron behavior Quantum numbers. Atomic Structure . Atomic orbitals Orbital energies Electron configuration and the periodic table

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3-Atomic Structure

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  1. 3-Atomic Structure Overview • Characteristics of Atoms • Interaction b/tw matter and light • Photoelectric Effect • Absorption and Emission Spectra • Electron behavior • Quantum numbers

  2. Atomic Structure • Atomic orbitals • Orbital energies • Electron configuration and the periodic table • Periodic table • Periodic properties • Energy

  3. Characteristics of Atoms • Atoms possess mass • Atoms contain positive nuclei • Atoms contain electrons • Atoms occupy volume • Atoms have various properties • Atoms attract one another • Atoms can combine with one another to form molecules

  4. Atomic Structure • Atomic structure studied through atomic interaction with light • Light: electromagnetic radiation • carries energy through space • moves at 3.00 x 108 m/s in vacuum • wavelike characteristics

  5. Electromagnetic Spectrum

  6. Visible Spectrum

  7.  Wavelength () & Frequency () amplitude  = number of complete cycles to pass given point in 1 second

  8. Energy c =  x = 3.00 x 108 m/s long wavelength  low frequency Low Energy High Energy short wavelength  high frequency

  9. Energy Mathematical relationship: E = h E = energy h = Planck’s constant: 6.63 x 10–34J s  = frequency in s–1

  10. E = Energy Mathematical relationship: E = h c =  x Energy: directly proportional to frequency inversely proportional to wavelength

  11. Problems 3-1, 2, & 3 • a) Calculate the wavelength of light with a frequency  = 5.77 x 1014 s–1 b) What is the energy of this light? 2. Which is higher in energy, light of wave-length of 250 nm or light of 5.4 x 10–7 m? 3. a) What is the frequency of light with an energy of 3.4 x 10–19 J? b) What is the wavelength of light with an energy of 1.4 x 10–20 J?

  12. Photoelectric Effect • Light on metal surface • Electrons emitted • Threshold frequency, o If  < o, no photoelectric effect If  > o, photoelectric effect As , kinetic energy of electrons 

  13. Photoelectric Effect Einstein: energy  frequency If  < o electron doesn’t have enough energy to leave the atom If  > o electron does have enough energy to leave the atom Energy is transferred from light to electron, extra is kinetic energy of electron Ephoton = hphoton = ho + KEelectron KEelectron = hphoton – ho Animation

  14. Problem 3-4 A given metal has a photoelectric threshold frequency of o = 1.3 x 1014 s1. If light of  = 455 nm is used to produce the photoelectric effect, determine the kinetic energy of the electrons that are produced.

  15. Bohr Model Line spectra Light through a prism  continuous spectrum: Ordinary white light

  16. Bohr Model Line spectra Light from gas-discharge tube through a prism  line spectrum: H2 discharge tube

  17. Line Spectra (emission) White light H He Ne

  18. Line Spectra (absorption) Gas-filled tube Light source

  19. Bohr Model For hydrogen: C = 3.29 x 1015 s–1 Niels Bohr: Electron energy in the atom is quantized. n = 1, 2, 3,…. RH = 2.18 x 10–18 J

  20. Bohr Model Eatom = Eelectron = h E = Ef– Ei = h Minus sign: free electron has zero energy Line spectrum Photoelectric effect:

  21. Bohr Energy Levels

  22. Electrons • All electrons have same charge and mass • Electrons have properties of waves and particles (De Broglie)

  23. Heisenberg Uncertainty Principle Cannot simultaneously know the position and momentum of electron x = h Recognition that classical mechanics don’t work at atomic level.

  24. Schrödinger Equation Erwin Schrödinger 1926 Wave functions with discrete energies Less empirical, more theoretical n En n wave functions or orbitals n2 probability density functions

  25. Quantum Numbers Each orbital defined by 3 quantum numbers Quantum number: number that labels state of electron and specifies the value of a property

  26. Quantum Numbers Principal quantum number, n (shell) Specifies energy of electron (analogous to Bohr’s n) Average distance from nucleus n = 1, 2, 3, 4…..

  27. Quantum Numbers Azimuthal quantum number,  (subshell) • = 0, 1, 2… n–1 n = 1,  = 0 n = 2,  = 0 or 1 n = 3,  = 0, 1, or 2 Etc.

  28. Quantum Numbers Magnetic quantum number, m Describes the orientation of orbital in space m = –….+  If  = 2, m = –2, –1, 0, +1, +2

  29. Problem 3-5 Fill in the quantum numbers in the table below.

  30. Schrödinger Equation Wave equations:  Each electron has  & E associated w/ it Probability Density Functions: 2 -graphical depiction of high probability of finding electron

  31. Probability Density Functions Link to Ron Rinehart’s page  energy 2 probability density function s, p, d, f, g 1s 3s 2s Node: area of 0 electron density

  32. 3p Probability Density Functions 2p Node: area of 0 electron density nodes Link to Ron Rinehart’s page

  33. Electrons and Orbitals Pauli Exclusion Principle: no two electrons in the same atom may have the same quantum numbers Electron spin quantum number ms = ½ Electrons are spin paired within a given orbital

  34. Electrons and Orbitals n = 1  = 0, m = 0, ms = ½ 2 electrons possible: 1,0,0,+½ and 1,0,0,–½ 2 electrons per orbital 1s1 H 1s2 He

  35. Electrons and Orbitals n = 2  = 0, m = 0, ms = ½ 2,0,0, ½ 2 electrons possible n = 2  = 1, m = –1,0,+1, ms = ½ 2,1,–1, ½ 2,1,0, ½ 2,1,+1, ½ 6 electrons possible

  36. Electron Configurations n = 1 1s 2 electrons possible H 1e– 1s1  He 2e– 1s2 

  37. Electron Configurations n = 2 2s 2 electrons possible Li 3e– 1s2 2s1  2s  1s Be 4e– 1s2 2s2  2s  1s

  38. 2p 2s 1s Electron Configurations n = 2 2p  = 1, m = –1, 0, +1 3 x 2p orbitals (px, py, pz): 6 electrons possible B 5e– 1s2 2s2 2p1

  39. Electron Configurations n = 2 2p  = 1, m = –1, 0, +1 3 x 2p orbitals (px, py, pz): 6 electrons possible 2p  B 5e– 1s2 2s2 2p1  2s 1s 

  40. 2p    2s 1s  Electron Configurations n = 2 2p  = 1, m = –1, 0, +1 C 6e– 1s2 2s2 2p2 Hund’s Rule: for degenerate orbitals, the lowest energy is attained when electrons w/ same spin is maximized

  41. Problem 3-6 Write electron configurations and depict the electrons for N, O, F, and Ne.

  42. 3p  3s 2p     2s 1s  Electron Configurations n = 3 3s, 3p, 3d Na 11e– 1s2 2s2 2p63s1

  43. 3p  3s 2p     2s 1s  Electron Configurations n = 3 3s, 3p, 3d Mg 12e– 1s2 2s2 2p63s2

  44. 3p   3s 2p     2s 1s  Electron Configurations n = 3 3s, 3p, 3d Al 13e– 1s2 2s2 2p63s23p1

  45. 3p    3s 2p     2s 1s  Electron Configurations n = 3 3s, 3p, 3d Si 14e– 1s2 2s2 2p63s23p2

  46. 3p     3s 2p     2s 1s  Electron Configurations n = 3 3s, 3p, 3d P 15e– 1s2 2s2 2p63s23p3

  47. 3p     3s 2p     2s 1s  Electron Configurations n = 3 3s, 3p, 3d S 16e– 1s2 2s2 2p63s23p4

  48. 3p     3s 2p     2s 1s  Electron Configurations n = 3 3s, 3p, 3d Cl 17e– 1s2 2s2 2p63s23p5

  49. 3p     3s 2p     2s 1s  Electron Configurations n = 3 3s, 3p, 3d Ar 18e– 1s2 2s2 2p63s23p6

  50. Electron Configurations 3d vs. 4s Filling order 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 5g 6s 6p 6d 7s 7p

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