Quantum Mechanics


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Quantum Mechanics

  1. 1. Unit 3, Part 2 Quantum Mechanics Chemistry Notes
  2. 2. Light & Wave Motion <ul><li>Waves carry energy from 1 place to another </li></ul><ul><li>Light = Electromagnetic Radiation </li></ul><ul><ul><li>carries energy through space </li></ul></ul><ul><ul><li>Behaves as a wave or as a stream of particles of energy </li></ul></ul><ul><ul><li>( wave-particle nature of light ) </li></ul></ul>
  3. 3. A Typical Wave http://www.pas.rochester.edu/~afrank/A105/LectureV/FG03_003_PCT.gif
  4. 4. Wave Descriptions <ul><li>Wavelength : distance from crest to crest (or trough to trough); unit = m </li></ul><ul><li>Wave speed : how fast the wave travels over time; unit = m/s </li></ul><ul><li>Frequency : the # of waves that pass an area in a given amount of time; unit = 1/s = Hz </li></ul>
  5. 5. The Nature of Light <ul><li>The Photon – a bundle of a specific amount of electromagnetic energy </li></ul><ul><ul><li>Generally considered particle-like but also demonstrates wave-like properties </li></ul></ul>
  6. 6. The Electromagnetic Spectrum <ul><li>Gamma (shortest) </li></ul><ul><li>X-ray </li></ul><ul><li>Ultraviolet/UV </li></ul><ul><li>Visible </li></ul><ul><li>Infrared/IR </li></ul><ul><li>Microwave </li></ul><ul><li>Radio (longest) </li></ul>
  7. 7. The Electromagnetic Spectrum http://www.yorku.ca/eye/spectrum.gif
  8. 8. Looking At Wavelength <ul><li>Frequency </li></ul><ul><li>Short Wavelength </li></ul><ul><li>High Frequency </li></ul><ul><li>Energy </li></ul><ul><li>Short Wavelength </li></ul><ul><li>High Energy </li></ul><ul><li>Or </li></ul><ul><li>High Amplitude </li></ul><ul><li>High Energy </li></ul>http://www2.glenbrook.k12.il.us/gbssci/phys/Class/sound/u11l2a2.gif
  9. 9. Visible Light Spectrum <ul><li>Red </li></ul><ul><li>Orange </li></ul><ul><li>Yellow </li></ul><ul><li>Green </li></ul><ul><li>Blue </li></ul><ul><li>Indigo </li></ul><ul><li>Violet http://cache.eb.com/eb/image?id=73584&rendTypeId=35 </li></ul>Long Short
  10. 10. Absorption & Emission Spectra <ul><li>Absorption </li></ul><ul><li>something is being “taken in” </li></ul><ul><li>Emission </li></ul><ul><li>something is being “given off” </li></ul><ul><li>--Atoms will release energy as photons of light; the energy of the photons corresponds to the energy changes experienced by the atom </li></ul>
  11. 11. Absorption & Emission Spectra <ul><li>The “fingerprint” of an element; unique </li></ul><ul><li>Absorption – when an atom absorbs energy; electrons jump to higher energy levels; dark line spectra </li></ul><ul><li>Emission – when an atom releases photons of energy; electrons return to lower energy levels; bright line spectra </li></ul>
  12. 12. Absorption & Emission Spectra http://www.astro.columbia.edu/~archung/labs/fall2001/images/spectrum_line.gif
  13. 13. Hydrogen Atom Example Notice the spectral lines in the same positions? These represent specific energy level changes of the electrons in the atom. http://www.solarobserving.com/pics/hydrogen-spectra.jpg Dark Line Bright Line
  14. 14. The Bohr Model http://library.thinkquest.org/19662/images/eng/pages/model-bohr-2.jpg Light Energy Released Light Energy Taken In
  15. 15. Bohr Model of Atom <ul><li>Small nucleus with e - orbiting like planets around the sun </li></ul><ul><li>Electrons occupy specific energy levels (orbits) </li></ul><ul><li>Currently, atomic models suggest that e - occupy 3D regions called orbitals within an electron cloud </li></ul>
  16. 16. Ground State vs. Excited State http://imagine.gsfc.nasa.gov/docs/teachers/lessons/xray_spectra/images/emit.gif http://content.answers.com/main/content/wp/en/thumb/c/c4/180px-Energylevels.png Ground State = lowest energy level Excited State = a higher energy level Higher energy levels are closer together…
  17. 17. Energy Levels Are Quantized <ul><li>Atom absorbs energy, electrons jump from ground state to higher energy level </li></ul><ul><li>Electrons return to ground state, atom gives off photon of energy </li></ul><ul><li>The photon is quantized = carries a specific amount of energy that depends on the size of the “jump” from 1 energy level to the next </li></ul>
  18. 18. Chemical Behavior of Atom <ul><li>Determined by electron structure </li></ul><ul><li>How the electrons are arranged determines how they will behave in a reaction </li></ul>
  19. 19. Locating Electrons in an Atom <ul><li>Impossible to predict the exact location </li></ul><ul><li>The most probable location is described as the orbital within electron cloud </li></ul><ul><li>Describe the locations of electrons (or its address) as the electron configuration </li></ul>
  20. 20. The Electron Configuration <ul><li>Pauli Exclusion Principle </li></ul><ul><ul><li>No 2 electrons can have the same electron figuration </li></ul></ul><ul><ul><li>No 2 electrons can occupy the same location at the same time </li></ul></ul><ul><ul><li>Remember - electrons are always moving! </li></ul></ul>
  21. 21. Electrons & Energy Levels <ul><li>Energy levels are quantized = specific amounts of energy </li></ul><ul><li>Energy levels are called the principle energy levels </li></ul><ul><li>Each principle energy level can be divided into specific sublevels </li></ul>
  22. 22. The Hydrogen Example <ul><li>Atomic Number = 1 </li></ul><ul><ul><li>1 proton </li></ul></ul><ul><ul><li>1 electron </li></ul></ul><ul><li>Hydrogen has only 1 electron but more than 1 orbital. The electron can only occupy one orbital at a time. </li></ul>
  23. 23. Quantum Numbers <ul><li>4 quantum numbers make up the electron configuration </li></ul><ul><ul><li>(n) = principle energy level </li></ul></ul><ul><ul><li>(l) = energy sublevel </li></ul></ul><ul><ul><li>(m) = actual orbital </li></ul></ul><ul><ul><li>(s) = spin </li></ul></ul>
  24. 24. Quantum Number Analogy <ul><li>Hospital Complex Analogy </li></ul><ul><ul><li>(n) = building </li></ul></ul><ul><ul><li>(l) = floor </li></ul></ul><ul><ul><li>(m) = room </li></ul></ul><ul><ul><li>(s) = bed </li></ul></ul><ul><li>(assuming 2 beds per room) </li></ul>
  25. 25. Electron Configurations Written <ul><li>Example = 2p 3 </li></ul><ul><li>2 = principle energy level </li></ul><ul><li>p = sublevel </li></ul><ul><li>3 = # of electrons in sublevel </li></ul>
  26. 26. Quantum Number (n) <ul><li>Principle energy level </li></ul><ul><li>n = 1, 2, 3, 4, 5, 6, or 7 </li></ul><ul><li>n = corresponds to the row (period) on the periodic table </li></ul><ul><li>As n increases, energy increases </li></ul><ul><li>n refers to orbital size; as n increases, the orbital size increases </li></ul><ul><li>1 = close to nucleus, 7 = far away </li></ul>
  27. 27. Quantum Number ( l ) <ul><li>Energy sublevel </li></ul><ul><li>Not all e - in an energy level have the exact energy </li></ul><ul><li>l = s, p, d, or f </li></ul><ul><li>Sublevels may overlap </li></ul><ul><li>A filled sublevel is spherical in shape </li></ul><ul><li>Orbital shapes can be complex </li></ul>
  28. 28. Quantum Number (m) <ul><li>The actual orbital of the electron </li></ul><ul><li>Each orbital holds 2 electrons </li></ul><ul><li>s sublevel – 1 orbital, spherical shape </li></ul><ul><li>p sublevel – 3 orbitals, dumbell shape </li></ul><ul><li>d sublevel – 5 orbitals, weird shape </li></ul><ul><li>f sublevel – 7 orbitals, complex shape </li></ul>
  29. 29. Orbital Shapes http://lincoln.pps.k12.or.us/lscheffler/OrbitalShapesOverheads_files/image002.jpg
  30. 30. Quantum Number (s) <ul><li>Spin of the electron </li></ul><ul><li>Clockwise (+½) or Counterclockwise (-½) </li></ul><ul><li>Electrons in the same orbital have opposite spins </li></ul>
  31. 31. 3d ___ ___ ___ ___ ___ 4s ___ 3p ___ ___ ___ Principle Energy Level (n) Energy Sublevel (l) Orbital (m) holds 2 electrons total
  32. 32. Example Orbital Diagrams 1s 1s2s 1s2s2p http://www.uwgb.edu/dutchs/PETROLGY/WhatElmsLookLike.HTM
  33. 33. Electron Configuration Guidelines <ul><li>Elec. configuration lists ALL e - in the atom, giving the principle energy level, sublevel, and e - per sublevel </li></ul><ul><li>An atom is neutral: the # of protons equals the # of electrons , described by the atomic number </li></ul><ul><li>The period (row) describes the number of energy levels present </li></ul>
  34. 34. Electron Configuration Guidelines <ul><li>2n 2 = the maximum number of electrons in an energy level </li></ul><ul><ul><li>n = the energy level </li></ul></ul><ul><ul><li>n = 1, then 2 electrons </li></ul></ul><ul><ul><li>n = 2, then 8 electrons </li></ul></ul><ul><ul><li>n = 3, then 18 electrons </li></ul></ul>
  35. 35. Electron Configuration Guidelines <ul><li>The nth energy level has n sublevels present </li></ul><ul><ul><li>n = 1, then 1 sublevel </li></ul></ul><ul><ul><li>n – 2, then 2 sublevels </li></ul></ul><ul><li>Any element is “put into” the energy level (period) of its highest energy level electrons, even if these electrons are not filled in last </li></ul>
  36. 36. Electron Configuration Guidelines <ul><li>Lower energy levels fill in first; then higher energy electrons are filled in </li></ul><ul><li>Hund’s Rule : in any sublevel, one electron is placed in each orbital before pairing begins </li></ul><ul><ul><li>Some exceptions (discussed later) </li></ul></ul><ul><li>Degenerate : orbitals have the same energy & appear at same “level” on an orbital filling diagram </li></ul>
  37. 37. Electron Configuration Guidelines Place one electron in each orbital before pairing! Degenerate Orbitals: All 3 have the same energy and appear at the same level Hund’s Rule
  38. 38. Electron Configuration Guidelines <ul><li>Order of Orbital Filling </li></ul><ul><li>You will use the Periodic Table to help remember this! </li></ul><ul><li>1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d… </li></ul>
  39. 39. Orbital Filling Clues http://pegasus.cc.ucf.edu/~jparadis/chem2045/chapter06.html
  40. 40. Electron Configuration Guidelines <ul><li>Reminder! </li></ul><ul><li>Each Orbital Holds 2 Electrons </li></ul><ul><li>s 1 orbital 2 electrons </li></ul><ul><li>p 3 orbitals 6 electrons </li></ul><ul><li> d 5 orbitals 10 electrons </li></ul><ul><li> f 7 orbitals 14 electrons </li></ul>
  41. 41. Orbital Filling Diagrams <ul><li>See notes packet for diagram template </li></ul><ul><li>Aufbau Principle (to build up) </li></ul><ul><ul><li>Electrons fill lowest energy states first </li></ul></ul><ul><ul><li>If multiple orbitals, electrons fill all unoccupied orbitals first before pairing electrons with opposite spins ( Hund’s Rule ) </li></ul></ul>
  42. 42. Exceptions You Need to Know <ul><li>Copper (Cu) & Chromium (Cr) </li></ul><ul><ul><li>Electrons are found in unexpected places; there is extra stability in having half-filled d orbitals </li></ul></ul><ul><ul><li>The completion of a d sublevel is desired, so an electron will fill the d sublevel before the next s </li></ul></ul><ul><ul><li>1 electron in each d 5 </li></ul></ul>