APChem- Chapter 7 Lecture- Periodic Trends

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Chapter seven lecture for AP Chemistry on periodic trends

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APChem- Chapter 7 Lecture- Periodic Trends

  1. 1. Development of Periodic Table <ul><li>Elements in the same group generally have similar chemical properties. </li></ul><ul><li>Properties are not identical, however. </li></ul>
  2. 2. Development of Periodic Table <ul><li>Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped. </li></ul>
  3. 3. Development of Periodic Table <ul><li>Mendeleev, for instance, predicted the discovery of germanium (which he called eka-silicon) as an element with an atomic weight between that of zinc and arsenic, but with chemical properties similar to those of silicon. </li></ul>
  4. 6. <ul><li>The atomic number depends on the number of protons in the nucleus, while the atomic weight depends (mainly) on the number of both protons and neutrons in the nucleus. </li></ul>
  5. 7. Periodic Trends <ul><li>In this chapter, we will rationalize observed trends in </li></ul><ul><ul><li>Sizes of atoms and ions. </li></ul></ul><ul><ul><li>Ionization energy. </li></ul></ul><ul><ul><li>Electron affinity. </li></ul></ul>
  6. 8. Effective Nuclear Charge <ul><li>In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons. </li></ul><ul><li>The nuclear charge that an electron experiences depends on both factors. </li></ul>
  7. 9. <ul><li>The attraction of an electron to the nucleus depends on </li></ul><ul><ul><li>The charge of the nucleus </li></ul></ul><ul><ul><li>The distance the electron is from the nucleus. </li></ul></ul><ul><ul><li>In large atoms, the electrons in outer shells do not experience the full charge of the nucleus because inner shells of electrons “shield” them from the nucleus. </li></ul></ul>
  8. 10. Effective Nuclear Charge <ul><li>The effective nuclear charge, Z eff , is found this way: </li></ul><ul><li>Z eff = Z − S </li></ul><ul><li>where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons. </li></ul>
  9. 13. <ul><li>The 2 p electron of a Ne atom. </li></ul>
  10. 14. <ul><li>2 s orbital </li></ul><ul><li>2 p orbital </li></ul><ul><li>3 s orbital </li></ul><ul><li>3 p orbital </li></ul><ul><li>3 d orbital </li></ul>In an germanium atom, which electron will experience the greatest effective nuclear charge? An electron in a…
  11. 15. <ul><li>2 s orbital </li></ul><ul><li>2 p orbital </li></ul><ul><li>3 s orbital </li></ul><ul><li>3 p orbital </li></ul><ul><li>3 d orbital </li></ul>In an germanium atom, which electron will experience the greatest effective nuclear charge? An electron in a…
  12. 16. The effective nuclear charge experienced by 3 p electrons in phosphorus is <ul><li>+2 </li></ul><ul><li>+3 </li></ul><ul><li>+5 </li></ul><ul><li>+7 </li></ul>
  13. 17. <ul><li>+2 </li></ul><ul><li>+3 </li></ul><ul><li>+5 </li></ul><ul><li>+7 </li></ul>Correct Answer: The effective nuclear charge is given by the equation: Z eff = Z  S where Z represents number of protons in the nucleus and S represents the average number of electrons between the nucleus and the electron in question. Here: Z eff = 15  10 = +5
  14. 18. Sizes of Atoms <ul><li>It is difficult to define atomic radii for single atoms because the outside of the electron cloud is ambiguous. </li></ul>
  15. 19. Sizes of Atoms <ul><li>The non-bonding atomic radius is defined as one-half of the distance between two atoms when they collide. </li></ul>
  16. 20. Sizes of Atoms <ul><li>The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei. </li></ul>
  17. 21. Sizes of Atoms <ul><li>Adding two bonding radii together gives a useful estimate of bond length. </li></ul>
  18. 22. Sizes of Atoms <ul><li>Bonding atomic radius tends to… </li></ul><ul><ul><li>… decrease from left to right across a row </li></ul></ul><ul><ul><ul><li>due to increasing Z eff . </li></ul></ul></ul><ul><ul><li>… increase from top to bottom of column </li></ul></ul><ul><ul><ul><li>due to increasing value of n </li></ul></ul></ul>
  19. 24. <ul><li>No </li></ul>
  20. 25. Order the following according to increasing atomic radius. <ul><li>Ge < Si < Se < Cl </li></ul><ul><li>Se < Si < Ge < Cl </li></ul><ul><li>Si < Cl < Ge < Se </li></ul><ul><li>Cl < Si < Se < Ge </li></ul><ul><li>Si < Ge < Se < Cl </li></ul>Ge Si Se Cl
  21. 26. Order the following according to increasing atomic radius. <ul><li>Ge < Si < Se < Cl </li></ul><ul><li>Se < Si < Ge < Cl </li></ul><ul><li>Si < Cl < Ge < Se </li></ul><ul><li>Cl < Si < Se < Ge </li></ul><ul><li>Si < Ge < Se < Cl </li></ul>Ge Si Se Cl
  22. 27. Based on atomic radii, which of the following bonds would be expected to be the shortest? <ul><li>H  H </li></ul><ul><li>H  F </li></ul><ul><li>F  F </li></ul><ul><li>Cl  Cl </li></ul>
  23. 28. Correct Answer: The shortest bonding atomic radius belongs to the H atom, so an H 2 molecule will have the shortest H  X (where X is any atom) bond distance. <ul><li>H  H </li></ul><ul><li>H  F </li></ul><ul><li>F  F </li></ul><ul><li>Cl  Cl </li></ul>
  24. 29. Sizes of Ions <ul><li>Ionic size depends upon: </li></ul><ul><ul><li>Nuclear charge. </li></ul></ul><ul><ul><li>Number of electrons. </li></ul></ul><ul><ul><li>Orbitals in which electrons reside. </li></ul></ul>
  25. 30. Sizes of Ions <ul><li>Cations are smaller than their parent atoms. </li></ul><ul><ul><li>The outermost electron is removed and repulsions are reduced. </li></ul></ul>
  26. 31. Sizes of Ions <ul><li>Anions are larger than their parent atoms. </li></ul><ul><ul><li>Electrons are added and repulsions are increased. </li></ul></ul>
  27. 32. Sizes of Ions <ul><li>Ions increase in size as you go down a column. </li></ul><ul><ul><li>Due to increasing value of n . </li></ul></ul>
  28. 33. Sizes of Ions <ul><li>In an isoelectronic series , ions have the same number of electrons. </li></ul><ul><li>Ionic size decreases with an increasing nuclear charge. </li></ul>
  29. 35. Order the following according to increasing atomic/ionic radius. <ul><li>C < Li + < O 2 - < N 3 - </li></ul><ul><li>N 3 - < O 2 - < C < Li + </li></ul><ul><li>Li + < C < N 3 - < O 2 - </li></ul><ul><li>Li + < C < N 3 - < O 2 - </li></ul><ul><li>Li + < C < O 2 - < N 3 - </li></ul>N 3 - Li + C O 2 -
  30. 36. Order the following according to increasing atomic/ionic radius. <ul><li>C < Li + < O 2 - < N 3 - </li></ul><ul><li>N 3 - < O 2 - < C < Li + </li></ul><ul><li>Li + < C < N 3 - < O 2 - </li></ul><ul><li>Li + < C < N 3 - < O 2 - </li></ul><ul><li>Li + < C < O 2 - < N 3 - </li></ul>N 3 - Li + C O 2 -
  31. 37. Which is larger: Na + or Na? <ul><li>Na + </li></ul><ul><li>Na </li></ul>
  32. 38. <ul><li>Na + </li></ul><ul><li>Na </li></ul>Correct Answer: Both have the same number of protons, but Na with one more electron will be larger.
  33. 39. Which is larger: Cl  or Cl? <ul><li>Cl  </li></ul><ul><li>Cl </li></ul>
  34. 40. <ul><li>Cl  </li></ul><ul><li>Cl </li></ul>Correct Answer: Both have the same number of protons, but Cl − with one more electron will be larger.
  35. 41. Lecture 7b- Sections 7.4 - 7.5
  36. 42. Ionization Energy <ul><li>Amount of energy required to remove an electron from the ground state of a gaseous atom or ion. </li></ul><ul><ul><li>First ionization energy is that energy required to remove first electron. </li></ul></ul><ul><ul><li>Second ionization energy is that energy required to remove second electron, etc. </li></ul></ul>
  37. 44. <ul><li>the photoelectric effect </li></ul>
  38. 45. Ionization Energy <ul><li>It requires more energy to remove each successive electron. </li></ul><ul><li>When all valence electrons have been removed, the ionization energy shows a HUGE increase. </li></ul>
  39. 48. <ul><li>I 2 for the carbon atom is greater. </li></ul>
  40. 49. Trends in First Ionization Energies <ul><li>As one goes down a column, less energy is required to remove the first electron. </li></ul><ul><ul><li>For atoms in the same group, Z eff is essentially the same, but the valence electrons are farther from the nucleus. </li></ul></ul>
  41. 50. Trends in First Ionization Energies <ul><li>Generally, as one goes across a row, it gets harder to remove an electron. </li></ul><ul><ul><li>As you go from left to right, Z eff increases. </li></ul></ul><ul><ul><li>And the electron’s distance from the nucleus decreases. </li></ul></ul>
  42. 51. Trends in First Ionization Energies <ul><li>However, there are two apparent discontinuities in this trend. </li></ul>
  43. 52. Trends in First Ionization Energies <ul><li>The first occurs between Groups IIA and IIIA. </li></ul><ul><li>Electron removed from p -orbital rather than s -orbital </li></ul><ul><ul><li>Electron farther from nucleus </li></ul></ul><ul><ul><li>Small amount of repulsion by s electrons. </li></ul></ul>
  44. 53. Trends in First Ionization Energies <ul><li>The second occurs between Groups VA and VIA. </li></ul><ul><ul><li>Electron removed comes from doubly occupied orbital. </li></ul></ul><ul><ul><li>Repulsion from other electron in orbital helps in its removal. </li></ul></ul>
  45. 55. Which will have the highest ionization energy? <ul><li>C </li></ul><ul><li>N </li></ul><ul><li>O </li></ul><ul><li>Al </li></ul><ul><li>Si </li></ul>
  46. 56. Which will have the highest ionization energy? <ul><li>C </li></ul><ul><li>N </li></ul><ul><li>O </li></ul><ul><li>Al </li></ul><ul><li>Si </li></ul>
  47. 57. Which will be the largest? <ul><li>I 1 of Na </li></ul><ul><li>I 2 of Na </li></ul><ul><li>I 1 of Mg </li></ul><ul><li>I 2 of Mg </li></ul><ul><li>I 3 of Mg </li></ul>I = ionization energy
  48. 58. Which will be the largest? <ul><li>I 1 of Na </li></ul><ul><li>I 2 of Na </li></ul><ul><li>I 1 of Mg </li></ul><ul><li>I 2 of Mg </li></ul><ul><li>I 3 of Mg </li></ul>I = ionization energy
  49. 60. <ul><li>They have the same electron configuration: [Ar]3d 3 . </li></ul>
  50. 61. Electron Affinity <ul><li>Energy change accompanying addition of electron to gaseous atom: </li></ul><ul><li>Cl + e −  Cl − </li></ul>
  51. 64. <ul><li>They are equal in magnitude and opposite in sign. </li></ul>
  52. 65. Trends in Electron Affinity <ul><li>In general, electron affinity becomes more exothermic as you go from left to right across a row. </li></ul>
  53. 66. Trends in Electron Affinity <ul><li>There are again, however, two discontinuities in this trend. </li></ul>
  54. 67. Trends in Electron Affinity <ul><li>The first occurs between Groups IA and IIA. </li></ul><ul><ul><li>Added electron must go in p -orbital, not s -orbital. </li></ul></ul><ul><ul><li>Electron is farther from nucleus and feels repulsion from s -electrons. </li></ul></ul>
  55. 68. Trends in Electron Affinity <ul><li>The second occurs between Groups 4A and 5A. </li></ul><ul><ul><li>Added electron must go in an occupied p -orbital. </li></ul></ul>
  56. 70. Based on periodic trends, which of the following elements is expected to have the largest (i.e., most negative) electron affinity? <ul><li>K </li></ul><ul><li>Na </li></ul><ul><li>Si </li></ul><ul><li>S </li></ul>
  57. 71. Correct Answer: S has most negative electron affinity in this list. <ul><li>K </li></ul><ul><li>Na </li></ul><ul><li>Si </li></ul><ul><li>S </li></ul>
  58. 72. Lecture 7c- Sections 7.6 - 7.8
  59. 73. Properties of Metal, Nonmetals, and Metalloids
  60. 74. Metals versus Nonmetals <ul><li>Table 7.3 on page 277 </li></ul><ul><li>Differences between metals and nonmetals tend to revolve around these properties. </li></ul>
  61. 75. Metals versus Nonmetals <ul><li>Metals tend to form cations. (Low I.E.) </li></ul><ul><li>Nonmetals tend to form anions. ( exo E.A.) </li></ul>
  62. 76. Metals <ul><li>Tend to be lustrous, malleable, ductile, and good conductors of heat and electricity. </li></ul>
  63. 77. Metals <ul><li>Compounds formed between metals and nonmetals tend to be ionic. </li></ul><ul><li>Metal oxides tend to be basic. </li></ul>
  64. 78. Metal oxide + water  metal hydroxide MgO (s) + H 2 O (l)  Mg(OH) 2 (aq) but only if it dissolves a bit in water Metal oxide (a.k.a. base anhydride) reactions Metal oxide + acid  water + salt MgO (s) + HCl (aq)  MgCl 2 (aq) + H 2 O (l)
  65. 79. Nonmetals <ul><li>Dull, brittle substances that are poor conductors of heat and electricity. </li></ul><ul><li>Tend to gain electrons in reactions with metals to acquire noble gas configuration. </li></ul>
  66. 80. Nonmetals <ul><li>Substances containing only nonmetals are molecular compounds. </li></ul><ul><li>Most nonmetal oxides are acidic. </li></ul>
  67. 81. nonmetal oxide + water  acid SO 3 (g) + H 2 O (l)  H 2 SO 4 (aq) nonmetal oxide (acid anhydride) reactions nonmetal oxide + base  water + salt CO 2 (g) + 2NaOH (aq)  Na 2 CO 3 (aq) + H 2 O (l)
  68. 82. Metalloids <ul><li>Have some characteristics of metals, some of nonmetals. </li></ul><ul><li>For instance, silicon looks shiny, but is brittle and fairly poor conductor. </li></ul>
  69. 83. Group Trends
  70. 84. Alkali Metals <ul><li>Soft, metallic solids. </li></ul><ul><li>Name comes from Arabic word for ashes. </li></ul>
  71. 85. Alkali Metals <ul><li>Found only as compounds in nature. </li></ul><ul><li>Have low densities and melting points. </li></ul><ul><li>Also have low ionization energies. </li></ul>
  72. 86. Alkali Metals <ul><li>Their reactions with water are famously exothermic. </li></ul><ul><li>2M (s) + H 2 O (l)  2MOH (aq) + H 2 (g) </li></ul>
  73. 87. Alkali Metals <ul><li>Alkali metals (except Li) react with oxygen to form peroxides. </li></ul><ul><li>K, Rb, and Cs also form superoxides: </li></ul><ul><li>K + O 2  KO 2 </li></ul><ul><li>Produce bright colors when placed in flame. </li></ul>
  74. 88. Alkaline Earth Metals <ul><li>Have higher densities and melting points than alkali metals. </li></ul><ul><li>Have low ionization energies, but not as low as alkali metals. </li></ul>
  75. 89. Alkaline Earth Metals <ul><li>Be does not react with water, Mg reacts only with steam, but others react readily with water. </li></ul><ul><li>Reactivity tends to increase down the group. </li></ul>Mg (s) + 2H 2 O (g)  Mg(OH) 2 (aq) + H 2 (g)
  76. 90. Hydrogen <ul><li>Hydrogen is a nonmetal. </li></ul><ul><li>It has a high ionization energy because its single electron experiences no nuclear shielding </li></ul><ul><li>When it does lose it’s electron it isn’t really an atom anymore. It is a proton. </li></ul>
  77. 91. Group 6A <ul><li>Oxygen, sulfur, and selenium are nonmetals. </li></ul><ul><li>Tellurium is a metalloid. </li></ul><ul><li>The radioactive polonium is a metal. </li></ul>
  78. 92. Oxygen <ul><li>Two allotropes: </li></ul><ul><ul><li>O 2 </li></ul></ul><ul><ul><li>O 3 , ozone </li></ul></ul><ul><li>Three anions: </li></ul><ul><ul><li>O 2 − , oxide </li></ul></ul><ul><ul><li>O 2 2 − , peroxide </li></ul></ul><ul><ul><li>O 2 1 − , superoxide </li></ul></ul><ul><li>Tends to take electrons from other elements (oxidation) </li></ul>
  79. 93. Sulfur <ul><li>Weaker oxidizing agent than oxygen. </li></ul><ul><li>Most stable allotrope is S 8 , a ringed molecule. </li></ul>
  80. 94. Group VIIA: Halogens <ul><li>Prototypical nonmetals </li></ul><ul><li>Name comes from the Greek halos and gennao : “salt formers” </li></ul>
  81. 95. Group VIIA: Halogens <ul><li>Large, negative electron affinities </li></ul><ul><ul><li>Therefore, tend to oxidize other elements easily </li></ul></ul><ul><li>React directly with metals to form metal halides </li></ul><ul><li>Chlorine added to water supplies to serve as disinfectant </li></ul>
  82. 96. Group VIIIA: Noble Gases <ul><li>Astronomical ionization energies </li></ul><ul><li>Positive electron affinities </li></ul><ul><ul><li>Therefore, relatively unreactive </li></ul></ul><ul><li>Monatomic gases </li></ul>
  83. 97. Group VIIIA: Noble Gases <ul><li>Xe forms three compounds: </li></ul><ul><ul><li>XeF 2 </li></ul></ul><ul><ul><li>XeF 4 (at right) </li></ul></ul><ul><ul><li>XeF 6 </li></ul></ul><ul><li>Kr forms only one stable compound: </li></ul><ul><ul><li>KrF 2 </li></ul></ul><ul><li>The unstable HArF was synthesized in 2000. </li></ul>

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