Atomic Structure


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Atomic Structure

  1. 1. The Wave Nature of Light<br />All forms of NRG/Light have characteristic wavelengths (λ) and frequency (υ).<br />Inversely related <br />λ υ = c (the speed of light)<br />Light visible to the naked eye exists as a tiny portion of the electromagnetic spectrum<br />
  2. 2. Electromagnetic Spectrum<br />
  3. 3. Max Planck<br />Transfer of energy was not continuous<br />Only came in certain values (quantized)<br />ΔE = hν<br />h = Planck’s constant = 6.626 x 10-34 Js<br />Packets of energy (quantum<br />
  4. 4. Albert Einstein<br />Proposed that electromagnetic radiation was quantized and made up of a stream of particles<br />Photons<br />The dual nature of light<br />λ = h/mv (deBroglie equation)<br />
  5. 5. Electrons as Waves<br />Louis de Broglie<br />~1924<br />Louis de Broglie (1924)<br />Applied wave-particle theory to electrons<br />electrons exhibit wave properties<br />QUANTIZED WAVELENGTHS<br />Standing Wave<br />Second Harmonic or First Overtone<br />Fundamental mode<br /> 200<br /> 150<br /> 100<br /> 50<br /> 0<br />- 50<br />-100<br />-150<br />-200<br /> 200<br /> 150<br /> 100<br /> 50<br /> 0<br />- 50<br />-100<br />-150<br />-200<br /> 200<br /> 150<br /> 100<br /> 50<br /> 0<br />- 50<br />-100<br />-150<br />-200<br />0 50 100 150 200<br />0 50 100 150 200<br />0 50 100 150 200<br />Adapted from work by Christy Johannesson<br />
  6. 6. Electrons as Waves<br />QUANTIZED WAVELENGTHS<br />Courtesy Christy Johannesson<br />
  7. 7. Electrons as Waves<br />ELECTRONS<br />VISIBLE LIGHT<br />Evidence: DIFFRACTION PATTERNS<br />Courtesy Christy Johannesson<br />Davis, Frey, Sarquis, Sarquis, Modern Chemistry2006, page 105<br />
  8. 8. Spectrum <br />Light is diffracted (bent) through different objects<br />Hydrogen Spectra gave interesting results<br />
  9. 9. Niels Bohr <br />In the Bohr Model (1913) the neutrons and protons occupy a dense central region called the nucleus, and the electrons orbit the nucleus much like planets orbiting the Sun.<br />They are not confined to a planar orbit like the planets are. <br />
  10. 10. Bohr Model<br />Planetary<br /> model<br />After Rutherford’s discovery, Bohr proposed that electrons travel in definite orbits around the nucleus.<br />
  11. 11. Bohr Cont.<br />When the electrons are in their lowest possible NRG level, they are in their ground state. <br />Electrons absorb NRG and go to a higher NRG level(Excited State)<br />NRG (Light) is released when the electron jumps from a higher NRG level to a lower NRG level(Fluorescence)<br />Constantly happening<br />
  12. 12. An unsatisfactory model for the hydrogen atom<br />According to classical physics, light<br />should be emitted as the electron <br />circles the nucleus. A loss of energy<br />would cause the electron to be drawn<br />closer to the nucleus and eventually<br />spiral into it.<br />Hill, Petrucci, General Chemistry An Integrated Approach 2nd Edition, page 294<br />
  13. 13. Quantum Mechanical Model<br />Niels Bohr &<br />Albert Einstein<br /> Modern atomic theory describes the electronic structure of the atom as the probability of finding electrons within certain regions of space (orbitals).<br />
  14. 14. Modern View<br />The atom is mostly empty space<br />Two regions<br />Nucleus <br />protons and neutrons<br />Electron cloud<br />region where you might find an electron<br />
  15. 15. Quantum Mechanics<br />Werner Heisenberg<br />~1926<br />Heisenberg Uncertainty Principle<br />Impossible to know both the velocity and position of an electron at the same time<br />g<br />Microscope<br />Electron<br />
  16. 16. Quantum Mechanics<br />SchrödingerWave Equation(1926)<br />finite # of solutions quantized energy levels<br />defines probability of finding an electron<br />Erwin Schrodinger<br />~1926<br />Courtesy Christy Johannesson<br />
  17. 17. Quantum Mechanics<br />Orbital<br />Orbital (“electron cloud”)<br />Region in space where there is 90% probability of finding an electron<br />90% probability of<br />finding the electron<br />Electron Probability vs. Distance<br />40<br />30<br />20<br />Electron Probability (%)<br />10<br />0<br />100<br />150<br />200<br />250<br />50<br />0<br />Distance from the Nucleus (pm)<br />Courtesy Christy Johannesson<br />
  18. 18. Quantum Numbers<br />UPPER LEVEL<br />Four Quantum Numbers:<br />Specify the “address” of each electron in an atom<br />Courtesy Christy Johannesson<br />
  19. 19. Quantum Numbers<br />Principal Quantum Number( n)<br />Angular Momentum Quantum #( l)<br />Magnetic Quantum Number( ml)<br />Spin Quantum Number( ms)<br />
  20. 20. Quantum Numbers<br />1. Principal Quantum Number( n)<br />Energy level<br />Size of the orbital<br />Integral values <br />1s<br />2s<br />3s<br />Courtesy Christy Johannesson<br />
  21. 21. Quantum Numbers<br />s<br />p<br />d<br />f<br />2. Angular Momentum Quantum #( l)<br />Energy sublevel<br />Shape of the orbital<br />Courtesy Christy Johannesson<br />
  22. 22. Angular Momentum Quantum #( l )<br />Has integral values from 0 to n-1<br />Related to the shape of the orbital<br />l = 0 is called s<br />l = 1 is called p<br />l = 2 is called d<br />l = 3 is called f<br />
  23. 23. Shapes of s, p, and d-Orbitals<br />
  24. 24. Shape of f-orbitals<br />
  25. 25. Quantum Numbers<br />3. Magnetic Quantum Number( ml)<br />Orientation of orbital<br />Specifies the exact orbital within each sublevel<br />Has values between l and -l<br />Courtesy Christy Johannesson<br />
  26. 26. Quantum Numbers<br />4. Spin Quantum Number( ms)<br />Electron spin  +½ or -½<br />An orbital can hold 2 electrons that spin in opposite directions.<br />Courtesy Christy Johannesson<br />