Estructura electronica atomica

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Estructura electronica atomica

  1. 1. Chapter 6 Electronic Structure of Atoms Prentice Hall © 2003 Chapter 6 CHEMISTRY The Central Science 9th Edition David P. White
  2. 2. Prentice Hall © 2003 Chapter 6 <ul><li>All waves have a characteristic wavelength,  , and amplitude, A . </li></ul><ul><li>The frequency,  , of a wave is the number of cycles which pass a point in one second. </li></ul><ul><li>The speed of a wave, v , is given by its frequency multiplied by its wavelength: </li></ul><ul><li>For light, speed = c . </li></ul>The Wave Nature of Light
  3. 3. Prentice Hall © 2003 Chapter 6 The Wave Nature of Light
  4. 4. The Wave Nature of Light
  5. 5. Prentice Hall © 2003 Chapter 6 <ul><li>Modern atomic theory arose out of studies of the interaction of radiation with matter. </li></ul><ul><li>Electromagnetic radiation moves through a vacuum with a speed of 2.99792458  10 -8 m/s. </li></ul><ul><li>Electromagnetic waves have characteristic wavelengths and frequencies. </li></ul><ul><li>Example: visible radiation has wavelengths between 400 nm (violet) and 750 nm (red). </li></ul>The Wave Nature of Light
  6. 6. The Wave Nature of Light
  7. 7. Prentice Hall © 2003 Chapter 6
  8. 8. Prentice Hall © 2003 Chapter 6 The Wave Nature of Light
  9. 9. Prentice Hall © 2003 Chapter 6
  10. 10. Prentice Hall © 2003 Chapter 6 <ul><li>Planck : energy can only be absorbed or released from atoms in certain amounts called quanta . </li></ul><ul><li>The relationship between energy and frequency is </li></ul><ul><li>where h is Planck’s constant (6.626  10 -34 J.s). </li></ul><ul><li>To understand quantization consider walking up a ramp versus walking up stairs: </li></ul><ul><ul><li>For the ramp, there is a continuous change in height whereas up stairs there is a quantized change in height. </li></ul></ul>Quantized Energy and Photons
  11. 11. Prentice Hall © 2003 Chapter 6 <ul><li>The Photoelectric Effect and Photons </li></ul><ul><li>The photoelectric effect provides evidence for the particle nature of light -- “quantization”. </li></ul><ul><li>If light shines on the surface of a metal, there is a point at which electrons are ejected from the metal. </li></ul><ul><li>The electrons will only be ejected once the threshold frequency is reached. </li></ul><ul><li>Below the threshold frequency, no electrons are ejected. </li></ul><ul><li>Above the threshold frequency, the number of electrons ejected depend on the intensity of the light. </li></ul>Quantized Energy and Photons
  12. 12. Prentice Hall © 2003 Chapter 6 <ul><li>The Photoelectric Effect and Photons </li></ul><ul><li>Einstein assumed that light traveled in energy packets called photons . </li></ul><ul><li>The energy of one photon: </li></ul>Quantized Energy and Photons
  13. 14. Prentice Hall © 2003 Chapter 6
  14. 15. Prentice Hall © 2003 Chapter 6
  15. 16. Prentice Hall © 2003 Chapter 6
  16. 17. Prentice Hall © 2003 Chapter 6 <ul><li>Radiation composed of only one wavelength is called monochromatic. </li></ul><ul><li>Radiation that spans a whole array of different wavelengths is called continuous. </li></ul><ul><li>White light can be separated into a continuous spectrum of colors. </li></ul><ul><li>Note that there are no dark spots on the continuous spectrum that would correspond to different lines. </li></ul>Line Spectra and the Bohr Model
  17. 18. Prentice Hall © 2003 Chapter 6
  18. 20. Prentice Hall © 2003 Chapter 6
  19. 21. Prentice Hall © 2003 Chapter 6 <ul><li>Line Spectra </li></ul><ul><li>Balmer: discovered that the lines in the visible line spectrum of hydrogen fit a simple equation. </li></ul><ul><li>Later Rydberg generalized Balmer’s equation to: </li></ul><ul><li>where R H is the Rydberg constant (1.096776  10 7 m -1 ), h is Planck’s constant (6.626  10 -34 J·s), n 1 and n 2 are integers ( n 2 > n 1 ). </li></ul>Line Spectra and the Bohr Model
  20. 22. Prentice Hall © 2003 Chapter 6 <ul><li>Bohr Model </li></ul><ul><li>Rutherford assumed the electrons orbited the nucleus analogous to planets around the sun. </li></ul><ul><li>However, a charged particle moving in a circular path should lose energy. </li></ul><ul><li>This means that the atom should be unstable according to Rutherford’s theory. </li></ul><ul><li>Bohr noted the line spectra of certain elements and assumed the electrons were confined to specific energy states. These were called orbits. </li></ul>Line Spectra and the Bohr Model
  21. 23. Prentice Hall © 2003 Chapter 6 <ul><li>Bohr Model </li></ul><ul><li>Colors from excited gases arise because electrons move between energy states in the atom. </li></ul>Line Spectra and the Bohr Model
  22. 24. Prentice Hall © 2003 Chapter 6 <ul><li>Bohr Model </li></ul><ul><li>Since the energy states are quantized, the light emitted from excited atoms must be quantized and appear as line spectra. </li></ul><ul><li>After lots of math, Bohr showed that </li></ul><ul><li>where n is the principal quantum number (i.e., n = 1, 2, 3, … and nothing else). </li></ul>Line Spectra and the Bohr Model
  23. 25. Prentice Hall © 2003 Chapter 6 <ul><li>Bohr Model </li></ul><ul><li>The first orbit in the Bohr model has n = 1, is closest to the nucleus, and has negative energy by convention. </li></ul><ul><li>The furthest orbit in the Bohr model has n close to infinity and corresponds to zero energy. </li></ul><ul><li>Electrons in the Bohr model can only move between orbits by absorbing and emitting energy in quanta ( h  ). </li></ul><ul><li>The amount of energy absorbed or emitted on movement between states is given by </li></ul>Line Spectra and the Bohr Model
  24. 26. Prentice Hall © 2003 Chapter 6
  25. 27. Prentice Hall © 2003 Chapter 6
  26. 28. Prentice Hall © 2003 Chapter 6 <ul><li>Bohr Model </li></ul><ul><li>We can show that </li></ul><ul><li>When n i > n f , energy is emitted. </li></ul><ul><li>When n f > n i , energy is absorbed </li></ul>Line Spectra and the Bohr Model
  27. 29. Bohr Model Line Spectra and the Bohr Model
  28. 30. Prentice Hall © 2003 Chapter 6
  29. 31. Prentice Hall © 2003 Chapter 6
  30. 32. Prentice Hall © 2003 Chapter 6 <ul><li>Limitations of the Bohr Model </li></ul><ul><li>Can only explain the line spectrum of hydrogen adequately. </li></ul><ul><li>Electrons are not completely described as small particles. </li></ul>Line Spectra and the Bohr Model
  31. 33. Prentice Hall © 2003 Chapter 6 <ul><li>Knowing that light has a particle nature, it seems reasonable to ask if matter has a wave nature. </li></ul><ul><li>Using Einstein’s and Planck’s equations, de Broglie showed: </li></ul><ul><li>The momentum, mv , is a particle property, whereas  is a wave property. </li></ul><ul><li>de Broglie summarized the concepts of waves and particles, with noticeable effects if the objects are small. </li></ul>The Wave Behavior of Matter
  32. 34. Prentice Hall © 2003 Chapter 6
  33. 35. Prentice Hall © 2003 Chapter 6
  34. 36. Prentice Hall © 2003 Chapter 6
  35. 37. Prentice Hall © 2003 Chapter 6 <ul><li>The Uncertainty Principle </li></ul><ul><li>Heisenberg’s Uncertainty Principle : on the mass scale of atomic particles, we cannot determine exactly the position, direction of motion, and speed simultaneously. </li></ul><ul><li>For electrons: we cannot determine their momentum and position simultaneously. </li></ul><ul><li>If  x is the uncertainty in position and  mv is the uncertainty in momentum, then </li></ul>The Wave Behavior of Matter
  36. 38. Prentice Hall © 2003 Chapter 6 H  = x 
  37. 39. Prentice Hall © 2003 Chapter 6 <ul><li>Schrödinger proposed an equation that contains both wave and particle terms. H  = x  </li></ul><ul><li>Solving the equation leads to wave functions. </li></ul><ul><li>The wave function gives the shape of the electronic orbital. </li></ul><ul><li>The square of the wave function, gives the probability of finding the electron, </li></ul><ul><li>that is, gives the electron density for the atom. </li></ul>Quantum Mechanics and Atomic Orbitals
  38. 40. Prentice Hall © 2003 Chapter 6 Quantum Mechanics and Atomic Orbitals
  39. 41. Prentice Hall © 2003 Chapter 6 <ul><li>Orbitals and Quantum Numbers </li></ul><ul><li>If we solve the Schrödinger equation, we get wave functions and energies for the wave functions. </li></ul><ul><li>We call wave functions orbitals . </li></ul><ul><li>Schrödinger’s equation requires 3 quantum numbers: </li></ul><ul><ul><li>Principal Quantum Number, n. This is the same as Bohr’s n . As n becomes larger, the atom becomes larger and the electron is further from the nucleus. </li></ul></ul>Quantum Mechanics and Atomic Orbitals
  40. 42. Prentice Hall © 2003 Chapter 6 <ul><li>Orbitals and Quantum Numbers </li></ul><ul><ul><li>Azimuthal Quantum Number, l. This quantum number depends on the value of n. The values of l begin at 0 and increase to ( n - 1). We usually use letters for l ( s , p , d and f for l = 0, 1, 2, and 3). Usually we refer to the s , p , d and f -orbitals. </li></ul></ul><ul><ul><li>Magnetic Quantum Number, m l . This quantum number depends on l . The magnetic quantum number has integral values between - l and + l . Magnetic quantum numbers give the 3D orientation of each orbital. </li></ul></ul>Quantum Mechanics and Atomic Orbitals
  41. 43. Prentice Hall © 2003 Chapter 6
  42. 44. Prentice Hall © 2003 Chapter 6
  43. 45. Prentice Hall © 2003 Chapter 6
  44. 46. Prentice Hall © 2003 Chapter 6
  45. 47. Prentice Hall © 2003 Chapter 6
  46. 48. Prentice Hall © 2003 Chapter 6 <ul><li>Orbitals and Quantum Numbers </li></ul>Quantum Mechanics and Atomic Orbitals
  47. 49. Prentice Hall © 2003 Chapter 6 <ul><li>Orbitals and Quantum Numbers </li></ul><ul><li>Orbitals can be ranked in terms of energy to yield an Aufbau diagram . </li></ul><ul><li>Note that the following Aufbau diagram is for a single electron system. </li></ul><ul><li>As n increases, note that the spacing between energy levels becomes smaller. </li></ul>Quantum Mechanics and Atomic Orbitals
  48. 50. Prentice Hall © 2003 Chapter 6 <ul><li>Orbitals and Quantum Numbers </li></ul>Quantum Mechanics and Atomic Orbitals
  49. 51. Prentice Hall © 2003 Chapter 6 <ul><li>The s-Orbitals </li></ul><ul><li>All s -orbitals are spherical. </li></ul><ul><li>As n increases, the s -orbitals get larger. </li></ul><ul><li>As n increases, the number of nodes increase. </li></ul><ul><li>A node is a region in space where the probability of finding an electron is zero. </li></ul><ul><li>At a node,  2 = 0 </li></ul><ul><li>For an s -orbital, the number of nodes is ( n - 1). </li></ul>Representations of Orbitals
  50. 52. Representations of Orbitals
  51. 53. Prentice Hall © 2003 Chapter 6 The s-Orbitals Representations of Orbitals
  52. 54. Prentice Hall © 2003 Chapter 6 <ul><li>The p-Orbitals </li></ul><ul><li>There are three p -orbitals p x , p y , and p z . </li></ul><ul><li>The three p -orbitals lie along the x -, y - and z - axes of a Cartesian system. </li></ul><ul><li>The letters correspond to allowed values of m l of -1, 0, and +1. </li></ul><ul><li>The orbitals are dumbbell shaped. </li></ul><ul><li>As n increases, the p -orbitals get larger. </li></ul><ul><li>All p -orbitals have a node at the nucleus. </li></ul>Representations of Orbitals
  53. 55. Prentice Hall © 2003 Chapter 6 The p-Orbitals Representations of Orbitals
  54. 56. Prentice Hall © 2003 Chapter 6 <ul><li>The d and f-Orbitals </li></ul><ul><li>There are five d and seven f -orbitals. </li></ul><ul><li>Three of the d -orbitals lie in a plane bisecting the x -, y - and z -axes. </li></ul><ul><li>Two of the d -orbitals lie in a plane aligned along the x -, y - and z -axes. </li></ul><ul><li>Four of the d -orbitals have four lobes each. </li></ul><ul><li>One d -orbital has two lobes and a collar. </li></ul>Representations of Orbitals
  55. 58. Prentice Hall © 2003 Chapter 6 <ul><li>Electron Spin and the Pauli Exclusion Principle </li></ul><ul><li>Line spectra of many electron atoms show each line as a closely spaced pair of lines. </li></ul><ul><li>Stern and Gerlach designed an experiment to determine why. </li></ul><ul><li>A beam of atoms was passed through a slit and into a magnetic field and the atoms were then detected. </li></ul><ul><li>Two spots were found: one with the electrons spinning in one direction and one with the electrons spinning in the opposite direction. </li></ul>
  56. 60. Electron Spin and the Pauli Exclusion Principle
  57. 61. Prentice Hall © 2003 Chapter 6 <ul><li>Electron Spin and the Pauli Exclusion Principle </li></ul><ul><li>Since electron spin is quantized, we define m s = spin quantum number =  ½. </li></ul><ul><li>Pauli’s Exclusions Principle : no two electrons can have the same set of 4 quantum numbers. </li></ul><ul><ul><li>Therefore, two electrons in the same orbital must have opposite spins. </li></ul></ul>
  58. 62. Prentice Hall © 2003 Chapter 6 <ul><li>Electron Spin and the Pauli Exclusion Principle </li></ul><ul><li>In the presence of a magnetic field, we can lift the degeneracy of the electrons. </li></ul>
  59. 63. Prentice Hall © 2003 Chapter 6 <ul><li>Orbitals and Their Energies </li></ul><ul><li>Orbitals of the same energy are said to be degenerate. </li></ul><ul><li>For n  2, the s - and p- orbitals are no longer degenerate because the electrons interact with each other. </li></ul><ul><li>Therefore, the Aufbau diagram looks slightly different for many-electron systems. </li></ul>
  60. 64. Prentice Hall © 2003 Chapter 6
  61. 65. Prentice Hall © 2003 Chapter 6
  62. 66. Prentice Hall © 2003 Chapter 6
  63. 67. Many-Electron Atoms Orbitals and Their Energies
  64. 68. Prentice Hall © 2003 Chapter 6
  65. 69. Prentice Hall © 2003 Chapter 6
  66. 70. Prentice Hall © 2003 Chapter 6 <ul><li>Hund’s Rule </li></ul><ul><li>Electron configurations tells us in which orbitals the electrons for an element are located. </li></ul><ul><li>Three rules: </li></ul><ul><ul><li>electrons fill orbitals starting with lowest n and moving upwards; </li></ul></ul><ul><ul><li>no two electrons can fill one orbital with the same spin ( Pauli ); </li></ul></ul><ul><ul><li>for degenerate orbitals, electrons fill each orbital singly before any orbital gets a second electron ( Hund’s rule ). </li></ul></ul>Electron Configurations
  67. 71. Prentice Hall © 2003 Chapter 6 <ul><li>Condensed Electron Configurations </li></ul><ul><li>Neon completes the 2 p subshell. </li></ul><ul><li>Sodium marks the beginning of a new row. </li></ul><ul><li>So, we write the condensed electron configuration for sodium as </li></ul><ul><li>Na: [Ne] 3 s 1 </li></ul><ul><li>[Ne] represents the electron configuration of neon. </li></ul><ul><li>Core electrons: electrons in [Noble Gas]. </li></ul><ul><li>Valence electrons: electrons outside of [Noble Gas]. </li></ul>Electron Configurations
  68. 72. Prentice Hall © 2003 Chapter 6 <ul><li>Transition Metals </li></ul><ul><li>After Ar the d orbitals begin to fill. </li></ul><ul><li>After the 3 d orbitals are full, the 4 p orbitals being to fill. </li></ul><ul><li>Transition metals : elements in which the d electrons are the valence electrons. </li></ul>Electron Configurations
  69. 73. Prentice Hall © 2003 Chapter 6 <ul><li>Lanthanides and Actinides </li></ul><ul><li>From Ce onwards the 4 f orbitals begin to fill. </li></ul><ul><li>Note: La: [Xe]6 s 2 5 d 1 4 f 0 </li></ul><ul><li>Elements Ce - Lu have the 4 f orbitals filled and are called lanthanides or rare earth elements. </li></ul><ul><li>Elements Th - Lr have the 5 f orbitals filled and are called actinides . </li></ul><ul><li>Most actinides are not found in nature. </li></ul>Electron Configurations
  70. 74. Prentice Hall © 2003 Chapter 6 <ul><li>The periodic table can be used as a guide for electron configurations. </li></ul><ul><li>The period number is the value of n . </li></ul><ul><li>Groups 1A and 2A have the s -orbital filled. </li></ul><ul><li>Groups 3A - 8A have the p -orbital filled. </li></ul><ul><li>Groups 3B - 2B have the d -orbital filled. </li></ul><ul><li>The lanthanides and actinides have the f -orbital filled. </li></ul>Electron Configurations and the Periodic Table
  71. 77. End of Chapter 6: Electronic Structure of Atoms Prentice Hall © 2003 Chapter 6

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