Chapter 5 Electrons in Atoms

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Chapter 5 Electrons in Atoms

  1. 1. Chapter 5 Electrons in Atoms
  2. 2. Light and Quantized Energy (5.1) <ul><li>The study of light led to the development of the quantum mechanical model. </li></ul><ul><li>Light is a kind of electromagnetic radiation EM). </li></ul><ul><li>All move at 3.00 x 10 8 m/s (c) Speed of light. </li></ul>
  3. 3. Parts of a wave Origin Wavelength Amplitude Crest Trough
  4. 4. Parts of Wave <ul><li>Crest - high point on a wave </li></ul><ul><li>Trough - Low point on a wave </li></ul><ul><li>Amplitude - distance from origin to crest </li></ul><ul><li>Wavelength (  ) - distance from crest to crest. To calculate use:  =c/v. </li></ul><ul><ul><ul><li>c = speed of light ( 3.00 x 10 8 m/s). </li></ul></ul></ul><ul><ul><ul><li>V = frequency (HZ) </li></ul></ul></ul>
  5. 5. Frequency <ul><li>Frequency (v) is the number of waves that pass a given point per second. Units are cycles/sec or hertz ( Hz ). To calculate use: </li></ul><ul><ul><ul><li>v = c/     </li></ul></ul></ul>
  6. 6. Frequency and wavelength <ul><li>Are inversely related (v = c/   </li></ul><ul><li>As one goes up the other goes down. </li></ul><ul><li>Different frequencies of light show as different colors of light. </li></ul><ul><li>The whole range is called the electromagnetic (EM) spectrum </li></ul>
  7. 7. Spectrum Radio waves Microwaves Infrared . Ultra-violet X-Rays Gamma Rays Long Wavelength Short Wavelength Visible Light Low energy High energy Low Frequency High Frequency
  8. 8. Light is a Particle <ul><li>Light is energy, Energy is quantized, therefore, Light must be quantized. </li></ul><ul><li>These quantized pieces of light are called photons . </li></ul><ul><li>Energy and frequency of the photons are directly related. E = h x  </li></ul><ul><li>(i.e.. High frequency = high energy) </li></ul>
  9. 9. Energy and frequency <ul><li>A photon is a particle of EM radiation with no mass that carries a quantum of energy. To calculate its energy use: </li></ul><ul><ul><li>E Photon = h x  </li></ul></ul><ul><ul><ul><li>E is the energy of the photon </li></ul></ul></ul><ul><ul><ul><li>  is the frequency </li></ul></ul></ul><ul><ul><ul><li>h is Planck’s constant (6.626 x 10 -34 Joules sec). </li></ul></ul></ul>
  10. 10. Photoelectric Effect <ul><li>In the photoelectric effect , electrons, called photoelectrons, are emitted from a metals surface when light of a certain frequency shines on it. (solar calculator) </li></ul><ul><li>Can be used to identify the type of metal. </li></ul>
  11. 11. Examples <ul><li>What is the frequency of red light with a wavelength of 4.2 x 10 -5 cm? </li></ul><ul><li>What is the wavelength of KFI, which broadcasts at with a frequency of 640 kHz? </li></ul><ul><li>What is the energy of a photon of each of the above? </li></ul>
  12. 12. Atomic Emission Spectrum <ul><li>How color tells us about atoms? </li></ul><ul><li>The atomic emission spectrum of an element is the set of frequencies of the EM waves emitted by atoms of the element. </li></ul><ul><li>Each is unique to the individual element giving a pattern of visible colors when viewed through a prism. </li></ul>
  13. 13. Prism <ul><li>White light is made up of all the colors of the visible spectrum. </li></ul><ul><li>Passing it through a prism separates it into colors. </li></ul>
  14. 14. If the light is not white <ul><li>By heating a gas or with electricity we can get it to give off colors. </li></ul><ul><li>Passing this light through a prism shows a unique color pattern </li></ul>
  15. 15. Atomic Emission Spectrum <ul><li>Each element gives off its own characteristic colors. </li></ul><ul><li>Can be used to identify the atom. </li></ul><ul><li>This is how we know what stars are made of. </li></ul>
  16. 16. <ul><li>These are called line spectra </li></ul><ul><li>unique to each element. </li></ul><ul><li>These are emission spectra </li></ul><ul><li>Mirror images are absorption spectra </li></ul><ul><li>Light with black missing </li></ul>
  17. 17. An explanation of the Atomic Emission Spectra
  18. 18. Where the electron starts <ul><li>When we write electron configurations we are starting at the writing the lowest energy level. </li></ul><ul><li>The energy level an electron starts from is called its ground state . </li></ul>
  19. 19. Changing the energy <ul><li>Let’s look at a hydrogen atom </li></ul>
  20. 20. <ul><li>Heat or electricity or light can move the electron up energy levels </li></ul>Changing the energy
  21. 21. <ul><li>As the electron falls back to ground state it gives the energy back as light </li></ul>Changing the energy
  22. 22. <ul><li>May fall down in steps </li></ul><ul><li>Each with a different energy </li></ul>Changing the energy
  23. 23. The Bohr Ring Atom n = 3 n = 4 n = 2 n = 1
  24. 24. { { {
  25. 25. <ul><li>The Further the electrons fall, the more the energy and the higher the frequency. </li></ul>Ultraviolet Visible Infrared
  26. 26. Light is also a wave <ul><li>Light is a particle - it comes in chunks. </li></ul><ul><li>Light is also a wave- we can measure its wave length and it behaves as a wave </li></ul><ul><li>The wavelength of a particle is calculated using  = h/mv . (de Broglie equation) </li></ul>
  27. 27. Diffraction <ul><li>When light passes through, or reflects off, a series of thinly spaced lines, it creates a rainbow effect because the waves interfere with each other. </li></ul>
  28. 28. A wave moves toward a slit.
  29. 29. A wave moves toward a slit.
  30. 30. A wave moves toward a slit.
  31. 31. A wave moves toward a slit.
  32. 32. A wave moves toward a slit.
  33. 35. Comes out as a curve
  34. 36. Comes out as a curve
  35. 37. Comes out as a curve
  36. 38. with two holes
  37. 39. with two holes
  38. 40. with two holes
  39. 41. with two holes
  40. 42. with two holes
  41. 43. with two holes Two Curves
  42. 44. with two holes Two Curves
  43. 45. Two Curves with two holes Interfere with each other
  44. 46. Two Curves with two holes Interfere with each other crests add up
  45. 47. Several waves
  46. 48. Several waves
  47. 49. Several waves
  48. 50. Several waves
  49. 51. Several waves
  50. 52. Several waves
  51. 53. Several waves
  52. 54. Several waves
  53. 55. Several waves
  54. 56. Several waves
  55. 57. Several waves Several Curves
  56. 58. Several waves Several Curves
  57. 59. Several waves Several Curves
  58. 60. Several waves Several Curves
  59. 61. Several waves Several waves Interference Pattern Several Curves
  60. 62. Diffraction <ul><li>Light shows interference patterns </li></ul><ul><li>What will an electron do when going through two slits? </li></ul><ul><ul><li>If it goes through one slit or the other, it will make two spots. </li></ul></ul><ul><ul><li>If it goes through both slits, then it makes an interference pattern. </li></ul></ul>
  61. 63. Electron “gun” Electron as Particle
  62. 64. Electron “gun” Electron as wave
  63. 65. Heisenberg Uncertainty Principle <ul><li>It is impossible to know exactly the speed and position of a particle. </li></ul>
  64. 66. Quantum Theory and the Atom (5.2) <ul><li>Rutherford’s model Discovered the nucleus </li></ul><ul><ul><li>small dense and positive </li></ul></ul><ul><ul><li>Electrons moved around in Electron cloud </li></ul></ul>
  65. 67. Bohr’s Model <ul><li>Why don’t the electrons fall into the nucleus? </li></ul><ul><li>Electrons move like planets around the sun. </li></ul><ul><ul><li>In circular orbits at different levels. </li></ul></ul><ul><ul><li>Energy separates one level from another. </li></ul></ul>
  66. 68. Bohr’s Model Nucleus Electron Orbit Energy Levels
  67. 69. Bohr’s Model Nucleus Electron Orbit Energy Levels
  68. 70. Bohr’s Model <ul><li>Further away from the nucleus means more energy. </li></ul><ul><li>There is no “in between” energy levels </li></ul>Increasing energy Nucleus First Second Third Fourth Fifth }
  69. 71. The Quantum Mechanical Model <ul><li>Energy is quantized. It comes in chunks. </li></ul><ul><li>Quanta - the amount of energy needed to move from one energy level to another. </li></ul><ul><li>Quantum is the leap in energy. </li></ul><ul><li>Schrödinger derived an equation that described the energy and position of the electrons in an atom </li></ul><ul><li>Treated electrons as waves. De Broglie equation predicts wave characteristics of moving particles. (  = h/mv) </li></ul>
  70. 72. <ul><li>Does have energy levels for electrons. </li></ul><ul><li>Orbits are not circular. </li></ul><ul><li>It can only tell us the probability of finding an electron a certain distance from the nucleus. </li></ul>The Quantum Mechanical Model
  71. 73. <ul><li>The electron is found inside a blurry “electron cloud” </li></ul><ul><li>An area where there is a chance of finding an electron. </li></ul><ul><li>Draw a line at 90 % probability. </li></ul>The Quantum Mechanical Model
  72. 74. Atomic Orbitals <ul><li>Principal Quantum Number (n) = the energy level of the electron ( 1,2,3,4,5 ). </li></ul><ul><li>Within each energy level, there are sublevels that have specific shapes (s, p, d, f) </li></ul><ul><li>Sublevels have atomic orbitals. These are regions where there is a high probability of finding an electron. (s=1,p=3,d=5,f=7) </li></ul><ul><li>Each orbital can hold up to 2 electrons. Electrons held: s=2, p=6, d=10, f=14 </li></ul>
  73. 75. <ul><li>An atomic orbital is a three-dimensional region around the nucleus that describes the electrons probable location. </li></ul><ul><li>There is one “s” </li></ul><ul><li>orbital for every energy </li></ul><ul><li>level (1s,2s,3s,4s,5s). </li></ul><ul><li>*It is Spherical shaped and can hold 2 electrons each. </li></ul>“S” orbitals
  74. 76. “P” orbitals <ul><li>Starts at the second energy level (2p,3p,4p,5p) </li></ul><ul><li>Dumbbell shaped (3 types) </li></ul><ul><li>Each can hold 2 electrons (6-total) </li></ul>
  75. 77. “ P” Orbitals (aligned on the x,y,z axis)
  76. 78. “D” orbitals <ul><li>Start at the third energy level (3d,4d,5d) </li></ul><ul><li>5 different shapes </li></ul><ul><li>Each can hold 2 electrons (10-total) </li></ul>
  77. 79. “F” orbitals <ul><li>Start at the fourth energy level (4f,5f) </li></ul><ul><li>Have seven different shapes </li></ul><ul><li>2 electrons per shape (14-total) </li></ul>
  78. 80. “F” orbitals
  79. 81. Summary 1 2 3 4 s 1 2 S P d S P D f Energy Level (n) S P Sublevels (S, p, d, f) Number of orbitals (Odd 1,3,5,7) 1 3 2 6 1 3 5 2 6 10 1 3 5 7 2 6 10 14 Maximum Number of Electrons (orbital x 2)
  80. 82. By Energy Level <ul><li>First Energy Level </li></ul><ul><li>only s orbital </li></ul><ul><li>only 2 electrons total </li></ul><ul><li>Written as 1s 2 </li></ul><ul><li>Second Energy Level </li></ul><ul><li>s and p orbitals are available </li></ul><ul><li>2 in s, 6 in p </li></ul><ul><li>Written as 2s 2 2p 6 </li></ul><ul><li>8 total electrons total </li></ul>
  81. 83. Filling order <ul><li>Lowest energy level fills first. </li></ul><ul><li>Each box gets 1 electron before anyone gets 2. </li></ul><ul><li>Orbitals can overlap </li></ul><ul><li>Counting system </li></ul><ul><ul><li>Each box is an orbital shape </li></ul></ul><ul><ul><li>Has Room for two electrons </li></ul></ul>
  82. 84. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 7p 3d 4d 5d 6d 4f 5f
  83. 85. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f
  84. 86. Electron Configurations (5.3) <ul><li>Shows the way electrons are arranged in atoms. </li></ul><ul><li>Aufbau principle - electrons enter the lowest energy first. </li></ul><ul><li>This causes difficulties because of the overlap of orbitals of different energies. </li></ul><ul><li>Pauli Exclusion Principle - at most 2 electrons per orbital - opposite spins </li></ul>
  85. 87. Electron Configuration <ul><li>Hund’s Rule - When electrons occupy orbitals of equal energy they don’t pair up until they have to . </li></ul><ul><li>Let’s determine the electron configuration for Phosphorus </li></ul><ul><li>Need to account for 15 electrons </li></ul>
  86. 88. <ul><li>The first to electrons go into the 1s orbital </li></ul><ul><li>Notice the opposite spins </li></ul><ul><li>only 13 more </li></ul>Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f
  87. 89. <ul><li>The next electrons go into the 2s orbital </li></ul><ul><li>only 11 more </li></ul>Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f
  88. 90. <ul><li>The next electrons go into the 2p orbital </li></ul><ul><li>only 5 more </li></ul>Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f
  89. 91. <ul><li>The next electrons go into the 3s orbital </li></ul><ul><li>only 3 more </li></ul>Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f
  90. 92. <ul><li>The last three electrons go into the 3p orbitals. </li></ul><ul><li>They each go into separate shapes </li></ul><ul><li>3 unpaired electrons </li></ul><ul><li>1s 2 2s 2 2p 6 3s 2 3p 3 </li></ul>Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f
  91. 93. The easy way to remember <ul><li>1s 2 </li></ul><ul><li>2 electrons </li></ul>1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f
  92. 94. Fill from the bottom up following the arrows <ul><li>1s 2 2s 2 </li></ul><ul><li>4 electrons </li></ul>1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f
  93. 95. Fill from the bottom up following the arrows <ul><li>1s 2 2s 2 2p 6 3s 2 </li></ul><ul><li>12 electrons </li></ul>1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f
  94. 96. Fill from the bottom up following the arrows <ul><li>1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 </li></ul><ul><li>20 electrons </li></ul>1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f
  95. 97. Fill from the bottom up following the arrows <ul><li>1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 </li></ul><ul><li>38 electrons </li></ul>1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f
  96. 98. Fill from the bottom up following the arrows <ul><li>1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 </li></ul><ul><li>56 electrons </li></ul>1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f
  97. 99. Fill from the bottom up following the arrows <ul><li>1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 </li></ul><ul><li>88 electrons </li></ul>1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f
  98. 100. Fill from the bottom up following the arrows <ul><li>1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 14 6d 10 7p 6 </li></ul><ul><li>118 electrons </li></ul>1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f
  99. 101. Rewrite when done <ul><li>Group the energy levels together </li></ul><ul><li>1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 7s 2 7p 6 </li></ul><ul><li>1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 14 6d 10 7p 6 </li></ul>
  100. 102. Exceptions to Electron Configuration (optional)
  101. 103. Orbitals fill in order <ul><li>Lowest energy to higher energy. </li></ul><ul><li>Adding electrons can change the energy of the orbital. </li></ul><ul><li>Filled and half-filled orbitals have a lower energy. </li></ul><ul><li>Makes them more stable. </li></ul><ul><li>Changes the filling order of d orbitals </li></ul>
  102. 104. Write these electron configurations <ul><li>Titanium - 22 electrons </li></ul><ul><li>1s 2 2s 2 2p 6 3s 2 3p 6 3d 2 4s 2 </li></ul><ul><li>Vanadium - 23 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 3d 3 4s 2 </li></ul><ul><li>Chromium - 24 electrons </li></ul><ul><li>1s 2 2s 2 2p 6 3s 2 3p 6 3d 4 4s 2 is expected </li></ul><ul><li>But this is wrong!! </li></ul>
  103. 105. Chromium is actually <ul><li>1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1 </li></ul><ul><li>Why? </li></ul><ul><li>This gives us two half filled orbitals. </li></ul>
  104. 106. Chromium is actually <ul><li>1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1 </li></ul><ul><li>Why? </li></ul><ul><li>This gives us two half filled orbitals. </li></ul>
  105. 107. Chromium is actually <ul><li>1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1 </li></ul><ul><li>Why? </li></ul><ul><li>This gives us two half filled orbitals. </li></ul><ul><li>Slightly lower in energy. </li></ul><ul><li>The same principle applies to copper. </li></ul>
  106. 108. Copper’s electron configuration <ul><li>Copper has 29 electrons so we expect </li></ul><ul><li>1s 2 2s 2 2p 6 3s 2 3p 6 3d 9 4s 2 </li></ul><ul><li>But the actual configuration is </li></ul><ul><li>1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 </li></ul><ul><li>This gives one filled orbital and one half filled orbital. </li></ul><ul><li>Remember these exceptions </li></ul><ul><ul><li>d 4 s 2  d 5 s 1 </li></ul></ul><ul><ul><li>d 9 s 2  d 10 s 1 </li></ul></ul>
  107. 109. In each energy level <ul><li>The number of electrons that can fit in each energy level is calculated with </li></ul><ul><li>Max e - = 2n 2 where n is the energy level </li></ul><ul><li>1 st </li></ul><ul><li>2 nd </li></ul><ul><li>3 rd </li></ul>

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