Light, Energy, And More


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  • Light and its nature have caused a lot of ink to flow during these last decades. Its dual behavior is partly explained by (1)Double-slit experiment of Thomas Young - who represents the photon’s motion as a wave - and also by (2)the Photoelectric effect in which the photon is considered as a particle. A Revolution: SALEH THEORY solves this ambiguity and this difficulty presenting a three-dimensional trajectory for the photon's motion and a new formula to calculate its energy. More information on
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Light, Energy, And More

  1. 1. Light, Energy, and More October 23, 2007 Chemistry
  2. 2. Recap… <ul><li>Electromagnetic Spectrum </li></ul><ul><li>High Energy </li></ul><ul><li>Low Energy </li></ul><ul><li>Wave Nature of Light </li></ul>
  3. 4. What’s Going On Here?
  4. 5. When we heat metal what happens?
  5. 6. <ul><li>Does the wave model of light explain these changes? </li></ul><ul><ul><li>Does not explain different wavelengths and frequencies at different temperatures </li></ul></ul><ul><li>What is light? </li></ul><ul><ul><li>Radiation….what is radiation? </li></ul></ul><ul><ul><ul><li>Particles or rays of energy </li></ul></ul></ul><ul><li>What is temperature anyways? </li></ul><ul><ul><li>The measure of the average kinetic energy of the particles in an object </li></ul></ul><ul><ul><li>Kinetic Energy vs. Potential Energy </li></ul></ul><ul><li>Too many questions…. </li></ul>
  6. 7. Max Planck (1900) <ul><li>German Physicist </li></ul><ul><li>Began to look for answers </li></ul><ul><li>Matter can only gain or lose energy in small quantized amounts </li></ul><ul><li>What’s quantized? </li></ul>
  7. 8. Vocab Word!!! <ul><li>QUANTUM </li></ul><ul><ul><li>Minimum amount of energy that can be gained or lost by an atom </li></ul></ul><ul><ul><li>The emitted light from a glowing metal is a ENERGY…this energy is quantized </li></ul></ul>
  8. 9. If energy is now quantized…how can we determine the amount of energy of a quantum? <ul><li>What is energy measured in? </li></ul><ul><li>What are we observing? </li></ul><ul><li>What happens to the color when we increase the temperature (energy)? </li></ul><ul><ul><li>Proportional or inversely proportional? </li></ul></ul><ul><li>Now we need a constant… </li></ul><ul><ul><li>Planck’s constant, h=6.626 x 10^-34 J*s </li></ul></ul>
  9. 11. Time to put these words into action! <ul><li>What is the frequency and wavelength electromagnetic radiation that emits 1.68 x 10^-17 J of energy? What type of electromagnetic radiation is this? </li></ul><ul><li>Wavelength= 1.18 x 10^-8 m </li></ul><ul><li>Ultraviolet radiation </li></ul>
  10. 12. Some questions to answer… <ul><li>What is the color we see? </li></ul><ul><li>What happens to the energy of the radiation when we increase the frequency, v, of the radiation emitted? </li></ul><ul><li>Iron at room temp…color and E? </li></ul><ul><li>Iron with a little heat…color and E? </li></ul><ul><li>Iron with lots of heat…color and E? </li></ul>
  11. 13. According to Planck’s Theory… <ul><li>If we have a given v, matter can emit or absorb E only in whole number multiples of hv (1hv, 2hv, 3 hv…) </li></ul><ul><li>Matter can ONLY have specific amounts of energy </li></ul><ul><li>Wall of kids building blocks </li></ul><ul><ul><li>We can only add or take away in increments of whole blocks…we cannot remove half a block </li></ul></ul>
  12. 14. The Big Mystery of the 1900’s… <ul><li>The Photoelectric Effect… </li></ul><ul><ul><li>What caused these color changes in metals??? </li></ul></ul>
  13. 15. Photoelectric Effect <ul><li>Electrons (photoelectrons) are emitted from a metal’s surface when a light of a certain frequency shines on the surface </li></ul><ul><li>Certain specific amounts of energy (what’s this called???) needed to knock out electrons from metal atoms. </li></ul>
  14. 17. Albert Einstein (1905) <ul><li>Added onto Planck’s Theory… </li></ul><ul><li>Called the electron’s emitted, PHOTONS (the little energy packets Planck called quantums) </li></ul><ul><li>Now… E photon = hv </li></ul>
  15. 18. <ul><li>Planck paved the way for the explanation behind the mystery </li></ul><ul><li>But some one else came into the picture… </li></ul>
  16. 19. Now light is not just a wave… <ul><li>Einstein’s Dual Nature of Light </li></ul><ul><ul><li>Particle and wave characteristics </li></ul></ul><ul><ul><li>Light is a beam of tiny particles, called photons, acting like a wave </li></ul></ul>
  17. 20. NEW WORD!!! <ul><li>Photon </li></ul><ul><li>A particle of electromagnetic radiation with no mass that carries a quantum of energy </li></ul>
  18. 22. What Einstein added… <ul><li>Energy of a photon has a minimum or threshold value to eject photoelectrons </li></ul><ul><li>What must happen for the photoelectric effect to occur? </li></ul><ul><ul><li>Energy of a photon (particle of EM radiation) must have the minimum energy requirement to free the electron from the atom of metal </li></ul></ul>
  19. 23. Mystery Solved! <ul><li>No matter how long a light of a certain frequency is shone on metal (intensity), electrons will not be ejected unless the minimum amount of energy is shone. </li></ul><ul><li>Silver metal </li></ul><ul><ul><li>Photoelectrons ejected when a light with a frequency of at least 1.14 x 10^15 Hz or greater is used </li></ul></ul><ul><li>Sodium metal </li></ul><ul><ul><li>Red light </li></ul></ul><ul><ul><li>Violet light </li></ul></ul>
  20. 24. Revised Planck’s Work… <ul><li>Einstein piggy-backed off of Planck’s Theory and we now have….. </li></ul>Photon
  21. 25. Time to do a little work…. <ul><li>Tiny water drops in the air disperse the white light of the sun into a rainbow. What is the Energy oa a photon from the violet portion of the rainbow if it has a frequency of 7.23x10^14 Hz? </li></ul><ul><li>E=4.79 x 10^-19 J </li></ul><ul><ul><li>Energy in a photon of violet light </li></ul></ul>
  22. 26. A couple more…  <ul><li>A photon has an energy of 2.93 x 10^-25 J. What is its frequency? What type of electromagnetic radiation is the photon? </li></ul><ul><li>V=4.42 x 10^8 Hz </li></ul><ul><li>TV or FM waves </li></ul>
  23. 27. Practice makes perfect…  <ul><li>What is the energy of each photon in the following types of radiation? </li></ul><ul><ul><li>6.32 x 10^20 Hz </li></ul></ul><ul><ul><li>9.50 x 10^13 Hz </li></ul></ul><ul><ul><li>1.05 x 10^16 Hz </li></ul></ul><ul><li>What types of radiation are each? </li></ul><ul><li>4.19 x 10^-13 J gamma or x-ray </li></ul><ul><li>6.29 x 10^20 J infrared </li></ul><ul><li>6.96 x 10^-18 J ultraviolet </li></ul>
  24. 28. How does this work? (Neon Signs)
  25. 29. What do we know about neon signs? <ul><li>Electricity is passed through tube full of neon gas </li></ul><ul><li>Neon atoms in tube absorb this energy </li></ul><ul><ul><li>What happens when something absorbs energy? </li></ul></ul><ul><li>Neon atoms in tube become excited </li></ul><ul><ul><li>Stable of Unstable? </li></ul></ul><ul><ul><li>What happens when something is unstable? </li></ul></ul><ul><ul><li>What do we see released energy as? </li></ul></ul><ul><ul><ul><li>Electromagnetic radiation…visible light!!! </li></ul></ul></ul>
  26. 30. EM spectrum <ul><li>What happens when we pass sunlight through a prism? </li></ul><ul><ul><li>Continuous spectrum of colors </li></ul></ul><ul><ul><li>ROYGBIV </li></ul></ul>
  27. 31. What happens when we pass light from neon gas or hydrogen gas through prism? <ul><li>Separation of colors </li></ul><ul><li>Discontinuous spectrum </li></ul><ul><li>This is called… </li></ul>
  28. 32. ATOMIC EMISSION SPECTRUM (AES) <ul><li>AES of an element is the set of frequencies of the electromagnetic radiation emitted by the atoms of that element </li></ul><ul><li>Individual lines of color </li></ul><ul><li>Only certain lines of color appear for certain elements… </li></ul><ul><ul><li>What does this mean…???? </li></ul></ul><ul><li>Every element has a unique AES </li></ul><ul><li>Why is this important? </li></ul>
  29. 35. Hydrogen Atom <ul><li>Why did scientists want to use hydrogen? </li></ul><ul><ul><li>How many protons? </li></ul></ul><ul><ul><li>How many electrons? </li></ul></ul><ul><ul><li>Do you think it is easy to use? </li></ul></ul><ul><ul><li>Check out the AES of hydrogen gas… </li></ul></ul>
  30. 37. Neils Bohr (1913) <ul><li>Danish Physicist </li></ul><ul><li>Worked with Rutherford </li></ul><ul><li>Quantum Model of Hydrogen atom </li></ul><ul><ul><li>Predicted lines of Hydrogen AES </li></ul></ul>
  31. 38. <ul><li>Hydrogen has only one electron but why do we get different colored lines on AES??? </li></ul><ul><ul><li>We get hydrogen atoms excited… </li></ul></ul><ul><ul><li>Electrons move to excited levels </li></ul></ul><ul><li>H has certain allowable energy states…. </li></ul><ul><ul><li>The lowest energy state is called the GROUND STATE </li></ul></ul>
  32. 39. Bohr’s Hydrogen Orbits… <ul><li>He related H’s energy states to the motion of an electron in an atom </li></ul><ul><li>Single electron in moves around nucleus in circular orbits </li></ul><ul><li>Smaller orbit, smaller radius, closer to nucleus means…? </li></ul><ul><ul><li>Lower energy level </li></ul></ul><ul><li>Larger orbit, larger radius, farther from the nucleus means…? </li></ul><ul><ul><li>Higher energy level </li></ul></ul>
  33. 41. Bohr’s Quantum Model <ul><li>Assigned quantum numbers, n, to each orbit </li></ul><ul><li>Calculated orbits radius </li></ul><ul><ul><li>Chart on page 127 </li></ul></ul><ul><li>1 st orbit  n=1 (first energy level) </li></ul><ul><li>2 nd orbit  n=2 (second energy level) </li></ul><ul><li>3 rd orbit  n=3 (third energy level) </li></ul>
  34. 44. When we add energy, what happens to electron? <ul><li>Electron excited </li></ul><ul><li>Moves to next energy level </li></ul><ul><li>Excited=? </li></ul><ul><ul><li>unstable </li></ul></ul><ul><li>What happens when something is unstable? </li></ul><ul><ul><li>Wants to get back to being stable </li></ul></ul><ul><ul><li>Releases energy </li></ul></ul><ul><ul><li>Goes back down to lower energy level </li></ul></ul><ul><li>Photon is emitted corresponding to the 2 different energy levels associated with the 2 orbits </li></ul>
  35. 46. NEW EQUATION <ul><li>/_ E= E higher e- orbit - E lower e- orbit =E photon =hv </li></ul><ul><li>Only certain energies are possible so only certain frequencies, v, of EM radiation are emitted </li></ul><ul><li>Lets look at the AES of Hydrogen… </li></ul>
  36. 47. <ul><li>How many lines are there? </li></ul><ul><li>So how many different types of radiations are we seeing? </li></ul><ul><li>There are 4 electron transitions account for lines in the hydrogen spectrum </li></ul><ul><li>Going from 3 rd orbital to 2 nd orbital… </li></ul><ul><li>Going from 4 th orbital to 2 nd orbital… </li></ul><ul><li>Going from 5 th orbital to 2 nd orbital… </li></ul><ul><li>Going from 6 th orbital to 2 nd orbital… </li></ul>
  37. 48. Names for these lines… <ul><li>Balmer Series </li></ul><ul><ul><li>The 4 visible color lines </li></ul></ul><ul><ul><li>Electrons that drop into n=2 </li></ul></ul><ul><li>Other electrons transitions not visible </li></ul><ul><ul><li>Lyman series </li></ul></ul><ul><ul><ul><li>Ultraviolet light </li></ul></ul></ul><ul><ul><ul><li>Electrons drop into n=1 </li></ul></ul></ul><ul><ul><li>Paschen series </li></ul></ul><ul><ul><ul><li>Infrared </li></ul></ul></ul><ul><ul><ul><li>Electrons drop into n=3 </li></ul></ul></ul>
  38. 50. Problems with Bohr’s Model <ul><li>Predicted AES lines of H but not any other elements </li></ul><ul><li>Did not account for all chemical behavior </li></ul><ul><li>Big problem… </li></ul><ul><ul><li>Electrons don’t move in circular orbits </li></ul></ul><ul><ul><li>Time for a new model… </li></ul></ul>
  39. 52. Louis De Broglie (1924) <ul><li>French physics graduate student </li></ul><ul><li>Proposed idea that accounted for the fixed energy levels in Bohr’s model </li></ul>
  40. 53. If waves can have particle like characteristics, then can particles, such as electrons, have wave like characteristics???
  41. 54. What he knew… <ul><li>Electrons have wavelike motion (because it’s a particle) </li></ul><ul><li>An electron had restricted orbits </li></ul><ul><li>Each orbit had a fixed radius from the nucleus </li></ul><ul><li>Are a wide variety of wavelengths, frequencies, and energies possible? </li></ul>
  42. 56. <ul><li>No…there could only be allowed certain possible frequencies, wavelengths, and energies in an atom </li></ul><ul><li>De Broglie came up with an equation for the wavelength of a particle of mass (m) moving at velocity (v). </li></ul>
  43. 57. De Broglie’s Equation
  44. 58. What does this equation do? <ul><li>What are we using? </li></ul><ul><ul><li>Wavelength </li></ul></ul><ul><ul><li>Planck’s constant </li></ul></ul><ul><ul><li>Mass of the particle </li></ul></ul><ul><ul><li>Velocity </li></ul></ul><ul><li>Tells us that all moving particles have wave-like characteristics </li></ul>
  45. 59. Food for thought… <ul><li>Cars? </li></ul><ul><li>Baseball? </li></ul><ul><li>Do these have wavelike characteristics? Why or why not? </li></ul>
  46. 60. <ul><li>Yes…let’s look at the equation… </li></ul><ul><li>λ = h </li></ul><ul><li>mv </li></ul><ul><li>The car and the baseball do have a velocity and a mass… </li></ul><ul><li>Using De Broglie’s equation we do get a wavelength for the movement of a baseball and a car… </li></ul><ul><li>Let’s try the calculation… </li></ul>
  47. 61. Problem time… <ul><li>Mass of car= 910 kg </li></ul><ul><li>Velocity of car= 25m/s </li></ul><ul><li>What is the wavelength of the moving car? </li></ul><ul><ul><li>2.9 x 10^-38 m </li></ul></ul><ul><li>How big is this? </li></ul><ul><li>Can we see or measure this wavelength? </li></ul><ul><ul><li>No, much to small to be detected, even with the most sophisticated equipment </li></ul></ul>
  48. 62. Another one… <ul><li>Electron speed= 25 m/s </li></ul><ul><li>Electron mass= 9.11 x 10^-28 g </li></ul><ul><li>What is the wavelength of the moving electron? </li></ul><ul><ul><li>2.9 x 10^-5 m </li></ul></ul><ul><li>Do you think we can measure this wavelength and see it? </li></ul><ul><ul><li>Yes, with the right equipment </li></ul></ul>
  49. 63. Practice makes perfect  <ul><li>What is the wavelength of an electron of mass 9.11 x 10-28 kg traveling at a velocity of 2.00 x 108 m/s? (Planck's constant = 6.63 x 10-34 J/Hz. </li></ul><ul><li>3.64 x 10-15m. </li></ul>
  50. 65. Werner Heisenberg (1901-1976) <ul><li>German theoretical physicist </li></ul><ul><li>Drew conclusion from Rutherford, Bohr, and De Broglie’s models </li></ul>
  51. 66. Problem with finding the position of an electron <ul><li>Helium balloon in a dark room </li></ul><ul><li>How would you determine the location of this balloon? </li></ul><ul><li>Is the balloon going to stay in the same position? </li></ul><ul><li>Energy transfer </li></ul><ul><li>What if I gave you a flashlight? </li></ul><ul><ul><li>What happens when we shine a beam of light on the balloon? </li></ul></ul>
  52. 67. <ul><li>Photons from light that reflect off of the balloon reach our eyes and tell us where the balloon is </li></ul><ul><li>Is there a transfer of energy? </li></ul><ul><ul><li>How big is the balloon compared to the photons? </li></ul></ul><ul><li>Can we do the same thing with finding the location of an electron in an atom? </li></ul><ul><li>Heisenberg focused on the interactions between photons and electrons… </li></ul>
  53. 68. Heisenberg Uncertainty Principle <ul><li>It is fundamentally impossible to know precisely both the velocity and position of a particle at the same time </li></ul>
  54. 69. Erwin Schrodinger (1926) <ul><li>Austrian physicist </li></ul><ul><li>Furthered De Broglie’s wave-particle theory </li></ul><ul><li>Derived equation that treated hydrogen’s electron as a wave </li></ul><ul><li>Unlike Bohr’s, his fit well with atoms of different elements </li></ul>
  55. 70. Quantum Mechanical Model of the atom
  56. 71. The Quantum Mechanical Model <ul><li>Similar to Bohr’s… </li></ul><ul><ul><li>Limits an electron’s energy to certain values </li></ul></ul><ul><li>Unlike Bohr’s… </li></ul><ul><ul><li>What did Bohr say about the orbit of an electron around the nucleus? </li></ul></ul><ul><ul><li>The Quantum Mechanic Model makes no attempt to describe the electron’s path </li></ul></ul>
  57. 72. <ul><li>Schrodinger’s wave equation </li></ul><ul><ul><li>Solutions to equation called wave function </li></ul></ul><ul><ul><ul><li>Don’t worry about the equation its self…just know the basics…. </li></ul></ul></ul><ul><ul><li>Wave function  probability of finding the electron within a particular volume of space around the nucleus </li></ul></ul><ul><ul><li>High probability  more likely to occur </li></ul></ul><ul><ul><li>Low probability  less likely to occur </li></ul></ul>
  58. 73. What the wave function tells us <ul><li>The atomic orbital of the electron </li></ul><ul><ul><li>Atomic orbital  3-D region around nucleus </li></ul></ul><ul><li>Fuzzy Cloud </li></ul><ul><li>Density of the cloud at a given point is proportional to the probability of finding the electron at that point </li></ul>
  59. 74. New Word <ul><li>Orbital  region of space where there is a 90% probability of finding an electron of a given energy </li></ul><ul><li>“electron cloud” </li></ul>Orbital
  60. 75. What did Bohr assign to electron orbitals? <ul><li>Quantum numbers </li></ul><ul><li>Quantum Mechanical Model does the same… </li></ul>
  61. 76. <ul><li>Four Quantum Numbers: </li></ul><ul><ul><li>Specify the “address” (zip code) of each electron in an atom </li></ul></ul>
  62. 77. <ul><li>First number…Principal Quantum Number ( n) </li></ul><ul><li>Energy level (associated with the electron) </li></ul><ul><li>Size if orbital </li></ul><ul><ul><li>Lowest energy level is assigned principle quantum number of 1 (n=1) </li></ul></ul><ul><ul><ul><li>Ground state </li></ul></ul></ul><ul><ul><li>What do you think happens as we increase n? </li></ul></ul><ul><ul><ul><li>Orbital becomes larger </li></ul></ul></ul><ul><ul><ul><li>Electron spends more time farther away from the nucleus  atom’s energy increases </li></ul></ul></ul>
  63. 78. Principle energy levels contain… Energy Sublevels
  64. 79. <ul><li>Principle energy level 1  single sublevel </li></ul><ul><li>Principle energy level 2  two sublevels </li></ul><ul><li>Principle energy level 3  three sublevels </li></ul><ul><li>What pattern do you see in the number of sublevels as we move further away from the nucleus? </li></ul><ul><ul><li>They increase as n increases (the further we get from the nucleus) </li></ul></ul>UPPER LEVEL
  65. 80. <ul><li>Electron’s are labeled according to n value </li></ul><ul><li>In atom’s with more than one electron, two or more electron’s may have the same n value </li></ul><ul><ul><li>They are in the same “electron shell” </li></ul></ul>
  66. 81. Second quantum number Angular Momentum Quantum Number (l)
  67. 82. <ul><li>Each value of l corresponds to a different type of orbital with a different shape </li></ul><ul><li>Value of n controls l (subshells possible) </li></ul><ul><li>Angular momentum numbers can equal 0, 1, 2, 3… </li></ul><ul><li>l=n-1 </li></ul><ul><ul><li>When n=1, l=0  only one possible subshell </li></ul></ul><ul><ul><li>When n=2, l=0,1  two possible subshells </li></ul></ul>
  68. 83. What the number of l means… <ul><li>Corresponds to the name of the subshell </li></ul><ul><ul><li>L=0  subshell s </li></ul></ul><ul><ul><li>L=1  subshell p </li></ul></ul><ul><ul><li>L=2  subshell d </li></ul></ul><ul><ul><li>L=3  subshell f </li></ul></ul>
  69. 84. S P D F: THE SUBLEVELS <ul><li>Each of these 4 sublevels has a unique shape </li></ul><ul><li>Each orbital may contain at most, 2 electrons </li></ul><ul><li>LETTERS ORIGINATED FROM DESCRIPTIONS OF THEIR SPECTRAL LINES </li></ul><ul><ul><li>S  sharp…spherical </li></ul></ul><ul><ul><li>P  principal…dumbbell shaped </li></ul></ul><ul><ul><li>D  diffuse…not all the same shape </li></ul></ul><ul><ul><li>F  fundamental…not all the same shape </li></ul></ul>
  70. 87. <ul><li>When principle energy level n=1, then l=0, which means there is only a single sublevel (one orbital) which is the small, spherical 1s </li></ul><ul><li>When principle energy level n=2, then l can equal 0 or 1, which means that there are two sublevels (orbitals) 2s and 2p </li></ul><ul><ul><li>2s sublevel  bigger than 1s, still sphere </li></ul></ul><ul><ul><li>2p sublevel  three dumbbell shaped p orbitals of equal energy called 2px, 2py, and 2pz </li></ul></ul><ul><ul><ul><li>The letters are just there to tell you what axis the electrons go on: x,y, or z axis </li></ul></ul></ul><ul><li>When the principle energy level n=3, then l can equal 0,1, or 2, which means that there are 3 possible sublevels: </li></ul><ul><ul><li>3s, sphere, bigger than 1s and 2s </li></ul></ul><ul><ul><li>3p, dumbbells </li></ul></ul><ul><ul><li>3d </li></ul></ul><ul><ul><ul><li>Each d sublevel consists 5 orbitals of equal energy </li></ul></ul></ul><ul><ul><ul><li>Four d orbitals have same shape but different orientations </li></ul></ul></ul><ul><ul><ul><li>Fifth d orbital, 3d z2 is shaped and oriented different from the other four </li></ul></ul></ul><ul><li>When the principle energy level n=4, then 1 can equal 0,1,2, or 3 which means l=n-1=4 possible sublevels: </li></ul><ul><ul><li>Seven f orbitals of equal energy ( 2 electrons in each orbital) </li></ul></ul><ul><ul><li>4s, sphere </li></ul></ul><ul><ul><li>4p, dumbbells </li></ul></ul><ul><ul><li>4d, </li></ul></ul><ul><ul><li>4f </li></ul></ul>
  71. 88. n = # of sublevels per level n 2 = # of orbitals per level Sublevel sets: 1 s, 3 p, 5 d, 7 f
  72. 91. Orbitals combine to form a spherical shape. 2s 2p z 2p y 2p x
  73. 92. Remember… <ul><ul><li>1. Principal #  energy level </li></ul></ul><ul><ul><li>2. Ang. Mom. #  sublevel (s,p,d,f) </li></ul></ul>There are two more quantum numbers (3 and 4) we will discuss next class
  74. 93. Third Quantum Number <ul><li>M l  specifies the orientation of the orbital in space containing the electron </li></ul><ul><li>Tells us whether the orbital is on the x, y, or z axis </li></ul>
  75. 94. Fourth Quantum Number <ul><li>M s  related to the direction of the electron spin </li></ul><ul><li>Tells us if electron has a clockwise spin or counter clockwise spin </li></ul><ul><li>Specifies orientation of electrons spin axis </li></ul>
  76. 95. Recap… <ul><li>Bohr? </li></ul><ul><ul><li>Orbits explained hydrogen’s quantized energy states </li></ul></ul><ul><li>De Broglie? </li></ul><ul><ul><li>Dual particle and wave nature of electrons </li></ul></ul><ul><li>Schrodinger? </li></ul><ul><ul><li>Wave equation predicted existence of atomic orbitals containing electrons </li></ul></ul>
  77. 96. Electron Configuration <ul><li>Definition: arrangement of electrons in an atom </li></ul><ul><li>Basic rules for filling up orbital's with electrons </li></ul><ul><li>Which is more stable, low energy or high energy? </li></ul><ul><ul><li>So which orbitals are going to be filled up first? </li></ul></ul><ul><ul><li>We are going to want an arrangement that gives us the lowest possible energy </li></ul></ul>
  78. 97. Ground state electron configuration <ul><li>The most stable, lowest energy electron arrangement of an atom </li></ul><ul><li>Each element has a ground-state electron configuration </li></ul>
  79. 98. Three Rules for Electron Arrangement <ul><li>Aufbau Principle </li></ul><ul><li>Pauli Exclusion Principle </li></ul><ul><li>Hund’s Rule </li></ul>
  80. 99. Aufbau Principle <ul><li>Each electron occupies the lowest energy orbital available </li></ul><ul><li>In order to do this, you must learn the sequence of atomic orbitals from lowest to highest energy </li></ul><ul><li>Aufbau Diagram </li></ul><ul><ul><li>Each box represents an orbital </li></ul></ul><ul><ul><li>Each arrow represents an electron </li></ul></ul><ul><ul><li>Only two arrows per box… </li></ul></ul><ul><ul><ul><li>Only two electrons per orbital </li></ul></ul></ul>
  81. 102. Some important things to remember about Aufbau… <ul><li>All orbitals related to an energy sublevel are of equal energy </li></ul><ul><ul><li>All three 2p orbitals have the same energy </li></ul></ul><ul><li>In a multi-electron atom, the energy sublevels within a principle energy level have different energies </li></ul><ul><ul><li>All three 2p orbitals are of higher energy than the one 2s orbital </li></ul></ul>
  82. 104. <ul><li>In order of increasing energies, the sequence of energy sublevels within a principle energy level is s, p, d, f </li></ul><ul><li>Orbitals related to energy sublevels within one principle energy level can overlap orbitals related to energy sublevels within another principle level </li></ul><ul><ul><li>Ex. An orbital related to the atoms 4s sublevel has a lower energy than the five orbitals related to 3d sublevel. </li></ul></ul>
  83. 106. Pauli Exclusion Principle <ul><li>States that a maximum on 2 electrons can occupy a single atomic orbital but only if the electrons have opposite spins </li></ul><ul><li>Wolfgang Pauli </li></ul><ul><li>Austrian Physicist </li></ul><ul><li>Observed atoms in excited states </li></ul>
  84. 107. <ul><li>Each electron has a spin </li></ul><ul><li>Kinda like a spinning top </li></ul><ul><li>It can only spin in one of 2 directions </li></ul><ul><li>In order for electrons to be together in an orbital, they must have opposite spins </li></ul>
  85. 108. Hund’s Rule <ul><li>What kind of charge do electrons have? </li></ul><ul><li>Do they attract or repel each other? </li></ul><ul><li>So…….. </li></ul><ul><li>Hund’s Rule states that single electrons with the same spin must occupy all each energy equal orbital before additional electrons with opposite spins can occupy the same orbital </li></ul>
  86. 109. 2p orbitals
  87. 110. Read section 5-3!