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Ap chem unit 7


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Ap chem unit 7

  1. 1. Atomic Structure and Periodicity<br />AP Chem: Unit 7<br />
  2. 2. Electromagnetic Radiation<br />
  3. 3. Electromagnetic Radiation<br />One of the ways that energy travels through space is by electromagnetic radiation. <br />light from the sun<br />X-rays<br />microwaves<br />
  4. 4. Wave Characteristics<br />Wavelength (λ) – is the distance between two consecutive peaks or troughs in a wave.<br />Frequency (ν) – is the number of waves (cycles) per second that pass a given point in space. <br />units – hertz or waves/sec (s-1)<br />Speed (c) – all types of electromagnetic radiation travel at the speed of light.<br />2.9979 x 108m/s<br />c = λν<br />
  5. 5. Electromagnetic Radiation<br />
  6. 6. Electromagnetic Radiation<br />
  7. 7. The Nature of Matter<br />
  8. 8. Wave and Particle Duality<br />Planck found that matter could only absorb or emit energy in whole number multiples of the quantity hν. <br />h is Planck’s constant = 6.626 x10-34Js<br />ΔE = hν<br />Transfer of energy is not continuous but is quantized and can occur only in discrete amounts called quantum. Thus energy has particle properties as well as wave properties.<br />
  9. 9. Einstein<br />
  10. 10. Wave and Particle Duality<br />Einstein proposed that electromagnetic radiation was also quantized and could be viewed as a stream of “particles” called photons. <br />Ephoton = hv = hc/λ<br />
  11. 11. The Photoelectric Effect<br />The photoelectric effect refers to the phenomenon in which electrons are emitted from the surface of a metal when light strikes it. <br />No electrons are emitted by a metal below a specific threshold frequency (vo)<br />For light with frequency lower than the threshold frequency, no electrons are emitted regardless of intensity of the light.<br />
  12. 12. The Photoelectric Effect<br />For light with frequency greater than the threshold frequency, the number of electrons emitted increases with the intensity of the light.<br />For light with frequency greater than the threshold frequency, the kinetic energy of the emitted electrons increases directly with frequency of the light.<br />
  13. 13. The Photoelectric Effect<br />These observations can be explained by assuming that electromagnetic radiation is quantized (consists of photons), and that the threshold frequency represents the minimum energy required to remove the electron from the metal’s surface.<br />Minimum energy required to remove an electron = Eo = hvo<br />KEelectron = ½ mv2 = hv – hvo<br />
  14. 14. Planck and Einstein Conclusions<br />Energy is quantized. It can occur only in discrete units called quanta.<br />Electromagnetic radiation, which was previously thought to exhibit only wave properties, seems to show certain characteristics of particulate matter as well. This phenomenon is sometimes referred to as the dual nature of light.<br />
  15. 15. Wave Particle Duality<br />The main significance of the equation E = mc2 is that energy has mass. <br />m = E/c2<br />
  16. 16. Louis de Broglie (1892-1987)<br />Since light which previously was thought to be purely wavelike, was found to have certain characteristics of particulate matter. But is the opposite also true? Does matter have that is normally assumed to be particulate exhibit wave properties?<br />
  17. 17. Louis de Broglie (1892-1987)<br />de Broglie’s equation allows us to calculate the wavelength for a particle:<br />
  18. 18. de Broglie’s Proof<br />
  19. 19. Louis de Broglie (1892-1987)<br />Conclusion: Energy is really a form of matter, and all matter shows the same types of properties. All matter exhibits both particulate and wave properties. <br />
  20. 20. The Atomic Spectrum of Hydrogen<br />
  21. 21. Spectrum<br />A continuous spectrum results when white light passes through a prism and all wavelengths (colors) are shown.<br />An emission spectrum produces only a few lines of color that is limited to discrete wavelengths produced by an atom. This is called a line spectrum and is specific to each atom.<br />
  22. 22. Hydrogen Line Spectrum<br />The significance of the line spectrum is that it indicates that only certain energies are allowed for the electron in the hydrogen atom. In other words the energy of the electron in the hydrogen atom is quantized <br />
  23. 23. Hydrogen Line Spectrum<br />
  24. 24. The Bohr Model<br />
  25. 25. Niels Bohr<br />Bohr developed a quantum model for the hydrogen atom that allowed for only specific energy levels around the atom that corresponded with specific radii.<br />
  26. 26. Niels Bohr (1885-1962)<br />The most important equation to come from Bohr’s model is the expression for the energy levels available to the electron in the hydrogen atom.<br />Z is the nuclear charge, n is the energy level. <br />
  27. 27. Niels Bohr (1885-1962)<br />The most important equation to come from Bohr’s model is the expression for the energy levels available to the electron in the hydrogen atom.<br />the negative sign calculates a lower energy closer to the atom, not the radiation of negative energy.<br />
  28. 28. Example<br />What is the change in energy if an electron in level 6 (excited state) returns to level 1 (ground state) in a hydrogen atom?<br />ni=6; nf=1; Z=1 (hydrogen nucleus contains a single proton)<br />
  29. 29. Example<br />What is the change in energy if an electron in level 6 (excited state) returns to level 1 (ground state) in a hydrogen atom?<br />
  30. 30. Example<br />ΔE=Ef – Ei= E1 – E6=-2.117 x 10-18J<br />The negative sign for the change in energy indicates that the atom has lost energy and is now more stable. This loss of energy produces a photon.<br />
  31. 31. Example<br />What is the corresponding wavelength for the energy produced from the electron jump?<br />E = -2.117 x 10-18J<br />9.383x10-8 m<br />
  32. 32. Bohr Model Conclusions<br />The model correctly fits the quantized energy levels of the hydrogen atom and postulates only certain allowed circular orbits for the electrons.<br />As the electron becomes more tightly bound, its energy becomes more negative relative to the zero-energy reference state. As the electron is brought closer to the nucleus, energy is released from the system. <br />
  33. 33. Bohr Model Conclusions<br />
  34. 34. Bohr Model Conclusions<br />The energy levels calculated by Bohr closely agreed with the values obtained from the hydrogen emission spectrum but does not apply well to other atoms. The Bohr’s model is fundamentally incorrect but is very important historically because it paved the way for our current theory of atomic structure.<br />
  35. 35. The Quantum Mechanical Model of the Atom<br />
  36. 36. Quantum Mechanics<br />Quantum Mechanics or Wave Mechanics were developed by three physicists: Heisenberg, de Broglie, and Schrodinger.<br />Emphasis was given to the wave properties of the electron.<br />The electron bound to the nucleus behaves similar to a standing wave.<br />
  37. 37. Quantum Mechanics<br />Like a standing wave, electrons can travel in patterns that allow for a common node. In other words, wave patterns around the nucleus must be in whole number wave patterns. But their exact movement is not known.<br />
  38. 38. Heisenberg Uncertainty Principle<br />There is a fundamental limitation to just how precisely we can know both the position and momentum of a particle at a given time. This limitation is small for large particles but substantial for electrons. <br />
  39. 39. Probability Distribution<br />A probability distribution is used to indicate the probability of finding an electron in a specific position.<br />Electron density map<br />Radial probability distribution <br />
  40. 40. Probability Distribution<br />For the hydrogen 1s orbital, the maximum radial probability occurs at a distance of 5.29x10-2nm or .529Å from the nucleus. This is the exact radius of the innermost orbit calculated in the Bohr Model.<br />The definition most often used by chemists to describe the size of the hydrogen 1s orbital is the radius of the sphere that encloses 90% of the total electron probability<br />
  41. 41. Quantum Numbers<br />
  42. 42. Quantum Numbers<br />Each orbital is characterized by a series of numbers called quantum numbers, which describe various properties of an orbital:<br />Principal quantum number (n)- has integral values : 1,2,3,4. It describes the size and energy of the orbital. Energy Level<br />
  43. 43. Quantum Numbers<br />Angular momentum quantum number (l) – has integral values from 0 to n-1. This is related to shape of the atomic orbitals. Sublevel<br /><ul><li>l=0 is s
  44. 44. l=1 is p
  45. 45. l=2 is d
  46. 46. l=3 is f
  47. 47. l=4 is g</li></li></ul><li>Quantum Numbers<br />Magnetic quantum number (ml) has values between l and –l , including 0. The value of ml is related to the orientation of the orbital in space. Axis designation <br />
  48. 48. Quantum Numbers<br />Electron spin quantum number (ms)- can only have one of two values, +½, -½. Electrons can spin in one of two opposite directions. <br />
  49. 49. Quantum Numbers<br />In a given atom no two electrons can have the same set of four quantum numbers (n, l, ml , ms). This is called the Pauli exclusion principle; an orbital can only hold two electrons, and they must have opposite spins. <br />
  50. 50. Quantum Numbers<br />
  51. 51.
  52. 52. Orbital Shapes and Energies<br />
  53. 53. S orbitals<br />The s orbitals have a characteristic spherical shape and contain areas of high probability separated by areas of zero probability. These areas are called nodal surfaces, or nodes.<br />
  54. 54. S orbitals<br />The number of nodes increases as n increases. The number of nodes = n - 1. <br />
  55. 55. P Orbitals<br />P orbitals each have two lobes separated by a node at the nucleus. The p orbitals are labeled according to the axis of the xyz coordinate system along which the lobes lie. <br />
  56. 56. P Orbitals<br />Cross section of electron probability of a p orbital<br />
  57. 57. D Orbitals<br />The five d orbitals first occur in energy level 3. They have two fundamental shapes. Four of the orbitals (dxz, dyz, dxy, and dx2-y2) have four lobes centered in the plane indicated in the orbital label. dx2-y2 lie along the x and y axes and dxylie between the axes. The fifth orbital dz2 has a unique shape with two lobes along the z axis and a belt centered in the xy plane. <br />
  58. 58. D Orbitals<br />
  59. 59. F Orbitals<br />The f orbitals first occur in level 4 and have shapes more complex than those of the d orbitals. These orbitals are not involved in the bonding in any of the compounds that we will consider. <br />
  60. 60. F Orbitals<br />
  61. 61. Orbital Energies <br />For the hydrogen atom, the energy of a particular orbital is determined by its value of n. Thus all orbitals with the same value of n have the same energy – they are said to be degenerate.<br />
  62. 62. Polyelectronic Atoms<br />
  63. 63. Polyelectronic Atoms<br />Polyelectronic atoms are atoms with more than one electron. To look at these atoms, three energy contributions must be considered:<br />Kinetic energy of the electrons as they move around the nucleus.<br />The potential energy of attraction between the nucleus and the electrons.<br />The potential energy of repulsion between the two electrons.<br />
  64. 64. Polyelectronic Atoms<br />Since electron pathways are unknown, dealing with the repulsions between electrons cannot be calculated exactly.<br />This is called the electron correlation problem.<br />
  65. 65. Polyelectronic Atoms<br />The electron correlation problem occurs with all polyelectronic atoms. To deal with this, we assume each electron is moving in a field of charge that is the net result of the nuclear attraction and the average repulsions of all the other electrons.<br />In other words,…..<br />
  66. 66. Polyelectronic Atoms<br />A valence electron is attracted to the highly charged nucleus and still repelled by the other ‘inner’ electrons. The net effect is that the electron is not bound nearly as tightly to the nucleus as it would be if it were alone. <br /> This is a screened or shielded affect. <br />
  67. 67. Polyelectronic Atoms<br />Because of this shielded affect. orbitals within a principal energy level do not have the same energy (degenerate). Sublevels vary in energy within a principal quantum level.<br />s<p<d<f<br />
  68. 68. Polyelectronic Atoms<br />Hydrogen vs. Polyelectronic<br />
  69. 69. History of the Periodic Table<br />
  70. 70. Early Greeks<br />Earth <br />Air<br />Fire <br />Water<br />
  71. 71. Dobereiner (1780-1849)<br />Johann Dobereiner was the first chemist to recognize patterns and found several groups of three elements that have similar properties.<br />chlorine, bromine and iodine<br />called triads.<br />
  72. 72. Newlands<br />John Newlands suggested that elements should be arranged in octaves, based on the idea that certain properties seemed to repeat for every eighth element in a way similar to the musical scale.<br />
  73. 73. Meyer and Mendeleev<br />The present form of the periodic table was conceived independently by two chemists: Meyer and Mendeleev. Usually Mendeleev is given most of the credit, because it was he who emphasized how useful the table could be in predicting the existence and properties of still unknown elements. <br />
  74. 74. Meyer and Mendeleev<br />In 1872 when Mendeleev first published his table, the elements gallium, scandium, and germanium were unknown. Mendeleev correctly predicted the existence and properties of these elements from the gaps in his periodic table. Mendeleev also corrected the atomic masses of several elements. <br />
  75. 75. Mendeleev’s Table<br />
  76. 76. The Aufbau Principle and the Periodic Table<br />
  77. 77. Three rules for Orbital Configuration<br />Aufbau principle – As protons are added, so are electrons, and fill in orbitals in order of energy levels.<br />Pauli Exclusion – Two electrons with opposite spins can occupy an orbital.<br />Hund’s rule – The lowest energy configuration for an atom is the one with one unpaired electrons in each degenerate orbital. (Electrons don’t like roommates)<br />
  78. 78. Valence Electrons<br />Valence electrons are the electrons in the outermost principal quantum level of an atom. These are the most important electrons because they are involved in bonding. <br />The inner electrons are known as core electrons.<br />
  79. 79. Valence Electrons<br />The elements in the same group have the same valence electron configuration. Elements with the same valence electron configuration show similar chemical behavior. <br />
  80. 80. Transition Metals<br />Transition metals have electron configurations that fill in the order of 4s before 3d. Copper and Chromium have a configuration that is observed different than what is expected.<br />Expected: Cr: 1s22s22p63s23p64s23d4<br />Observed: 1s22s22p63s23p64s13d5<br />Expected: Cu:1s22s22p63s23p64s23d9<br />Observed: 1s22s22p63s23p64s13d10<br />
  81. 81. Transition Metals<br />
  82. 82. Additional Orbital Rules<br />The (n+1)s orbital always fills before the nd orbitals. The s orbitals fill prior to the d orbitals due to the vicinity of the nucleus.<br />After lanthanum, which has the configuration of [Xe] 6s25d1, a group of 14 elements called the lanthanide series, or the lanthanides occurs. This seris of elements corresponds to the filling of the seven 4f orbitals. <br />
  83. 83. Additional Orbital Rules<br />
  84. 84. Additional Orbital Rules<br />After actinium, a group of 14 elements called the actinide series or actinides occurs. <br />The groups 1A, 2A, 3A…, the group numbers indicate the total number of valence electrons for the atoms in these groups.<br />
  85. 85. Additional Orbital Rules<br />After actinium, a group of 14 elements called the actinide series or actinides occurs. <br />The groups 1A, 2A, 3A…, the group numbers indicate the total number of valence electrons for the atoms in these groups.<br />
  86. 86. Additional Orbital Rules<br />
  87. 87. Periodic Trends in Atomic Properties<br />
  88. 88. Ionization Energy<br />Ionization energy is the energy required to remove an electron from a gaseous atom or ion when the atom or ion is assumed to be in its ground state:<br />X(g)X+(g) + e-<br />
  89. 89. Ionization Energy<br />It is always the highest-energy electron (the one bound least tightly) that is removed first. The first ionization energy (I1) is the energy required to remove that first electron. The second ionization energy (I2) is considerably larger. <br />
  90. 90. Ionization Energy<br />The first electron is removed from a neutral atom, the second from a +1 cation. The increase in positive charge binds the electrons more firmly and the ionization energy increases. The trend continues for consecutive electrons removed.<br />Core electrons are always held tighter than valence. <br />
  91. 91. Ionization Energy<br />First ionization energy increases from left to right across a period.<br />First ionization energy decreases in going down a group.<br />
  92. 92. Ionization Energy<br />
  93. 93. Electron Affinity<br />Electron Affinity is the change in energy change associated with the addition of an electron to a gaseous atom:<br />X(g) + e- X-(g)<br />
  94. 94. Electron Affinity<br />If the addition of the electron is exothermic the corresponding value for electron affinity will carry a negative sign.<br />The more negative the energy, the greater the quantity of energy released.<br />
  95. 95. Electron Affinity<br />Electron affinities generally become more negative from left to right across a period and becomes more positive down a group.<br />As with Ionization energy. Some exceptions occur due to repulsions and electron configuration. <br />
  96. 96. Electron Affinity<br />
  97. 97. Electron Affinity<br />
  98. 98. Atomic Radius<br />Atomic radii are measured by the distances between atoms in chemical compounds. <br />Covalent atomic radii are assumed to be half the distance between atoms in covalent bonds.<br />For metallic atoms, the metallic radii are obtained from half the distance between metal atoms in a solid metal crystal<br />
  99. 99. Atomic Radius<br />Atomic radii decrease in going from left to right across a period because of increasing nuclear charge and decreasing shielding.<br />Atomic radius increases down a group, because of the increases in the orbitals sizes associated with principal quantum numbers. <br />
  100. 100. Atomic Radius<br />
  101. 101. Atomic Radius<br />
  102. 102. The Properties of a Group: The Alkali Metals<br />
  103. 103. Information and the Periodic Table<br />It is the number and type of valence electrons that primarily determine an atom’s chemistry<br />The organization of the period table allows the prediction of electron configuration without memorization.<br />
  104. 104. Information and the Periodic Table<br />Groups on the periodic table have specialized names: Alkali metals, Alkaline earth metals, Halogens, …etc.<br />The most basic division of elements in the periodic table is into metals and non-metals. This division affects chemical properties.<br />Metals tend to give up electrons and have low ionization energies. The opposite is true for non-metals. <br />
  105. 105. Information and the Periodic Table<br />
  106. 106. Information and the Periodic Table<br />Metalloids are elements along the division line and exhibit both metallic and nonmetallic properties under certain circumstances. These elements are sometimes called semimetals. <br />
  107. 107. The Alkali Metals<br />Lithium, sodium, potassium, rubidium, cesium, and francium are the most chemically reactive of the metals. Hydrogen is found in group 1 but behaves as a nonmetal because its very small and the electron is bound tightly to the nucleus. <br />
  108. 108. The Alkali Metals<br />Going down the group the first ionization energy decreases and the atomic radius increases. The overall density increases due to the increase of atomic mass relative to atomic size (therefore more mass per unit volume). <br />
  109. 109. The Alkali Metals<br />There is a smooth decrease in melting point and boiling points in Group 1 that is not typical for other groups. <br />The most important chemical property of Group 1 is its ability to lose its valence electrons. Group 1 are very reactive. <br />
  110. 110. The Alkali Metals<br />Hydration energy of an ion represents the change in energy that occurs when water molecules attach to the metal cation. <br />The hydration energy is greatest with Li+ because it has the most charge density (charge per unit volume). This means that polar water molecules are more strongly attracted to the small Li+ ions <br />The order of reducing abilities in an aqueous reaction is Li > K > Na<br />
  111. 111. The Alkali Metals<br />
  112. 112. The Alkali Metals<br />
  113. 113. The END<br />