New chm 151 unit 6 power points

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New chm 151 unit 6 power points

  1. 1. CHEMICAL BONDING Chapters 9.1-9.3, 9.5-9.6 & 10.1 • IONIC BONDS • COVALENT BONDS • LEWIS ELECTRON DOT STRUCTURES 9-1
  2. 2. Chapter 9 Models of Chemical Bonding 9-2
  3. 3. Models of Chemical Bonding 9.1 Atomic Properties and Chemical Bonds 9.2 The Ionic Bonding Model 9.3 The Covalent Bonding Model 9.5 Between the Extremes: Electronegativity and Bond Polarity 9.6 An Introduction to Metallic Bonding 9-3
  4. 4. Goals & Objectives • See the following Learning Objectives on pages 356 and 388. • Understand these Concepts: • 9.1-3, 6-10, 13-15; 10.1-4. • Master these Skills: 9.1-2,6; • 10.1-5. 9-4
  5. 5. Figure 9.1 9-5 A comparison of metals and nonmetals.
  6. 6. Types of Chemical Bonding Ionic bonding involves the transfer of electrons and is usually observed when a metal bonds to a nonmetal. Covalent bonding involves the sharing of electrons and is usually observed when a nonmetal bonds to a nonmetal. Metallic bonding involves electron pooling and occurs when a metal bonds to another metal. 9-6
  7. 7. Figure 9.2 9-7 Three models of chemical bonding.
  8. 8. CHEMICAL BONDS • IONIC BONDING – process involving the complete transfer of electrons from one atom to another – generally between metals and nonmetals – for example Li metal and F a nonmetal • • • • • 9-8 Li + F --> Li+ + F1s22s1 1s22s22p5 [He] [Ne] Li+ is isoelectronic with the He atom F- is isoelectronic with the Ne atom
  9. 9. Lewis Electron-Dot Symbols To draw the Lewis symbol for any main-group element: • Note the group number, which gives the number of valence electrons. • Place one dot at a time on each of the four sides of the element symbol. • Keep adding dots, pairing them, until all are used up. Example: Nitrogen, N, is in Group 5A and therefore has 5 valence electrons. • or • N • or • N• • •• 9-9 or • • N• •• •• •• •N• •
  10. 10. Lewis Symbols and Bonding For a metal, the total number of dots in the Lewis symbol is the number of electrons the atom loses to form a cation. For a nonmetal, the number of unpaired dots equals - the number of electrons the atom gains to form an anion - or the number it shares to form covalent bonds. The octet rule states that when atoms bond, they lose, gain, or share electrons to attain a filled outer level of 8 electrons (or 2, for H and Li). 9-10
  11. 11. Figure 9.4 Lewis electron-dot symbols for elements in Periods 2 and 3. 9-11
  12. 12. Lewis Electron Dot Structures • includes only the valence shell or outermost shell electrons • Draw the Lewis Electron Dot Structures for the following atoms: – Na – Mg –S – Cl 9-12
  13. 13. 9-13
  14. 14. The Ionic Bonding Model An ionic bond is formed when a metal transfers electrons to a nonmetal to form ions, which attract each other to give a solid compound. The total number of electrons lost by the metal atom(s) equals the total number of electrons gained by the nonmetal atoms. 9-14
  15. 15. Figure 9.5 Three ways to depict electron transfer in the formation of Li+ and F-. Electron configurations Li 1s22s1 + F 1s22p5 → Li+ 1s2 + F- 1s22s22p6 Orbital diagrams Li ↑↓ ↑ + 1s 2s 2p F ↑↓ ↑↓ ↑↓ ↑↓ ↑ 1s 2s 2p Li+ ↑↓ 1s 2p ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 1s F- 2s 2s 2p Lewis electron-dot symbols Li+ + •• F •• •• •• 9-15 •• Li• •• •F ••
  16. 16. Sample Problem 9.1 Depicting Ion Formation PROBLEM: Use partial orbital diagrams and Lewis symbols to depict the formation of Na+ and O2− ions from the atoms, and determine the formula of the compound formed. PLAN: Draw orbital diagrams and Lewis symbols for Na and O atoms. To attain filled outer levels, Na loses one electron and O gains two. Two Na atoms are needed for each O atom so that the number of electrons lost equals the number of electrons gained. SOLUTION: •• •• •O • Na• 9-16 2Na+ •• 2+ O •• •• •• Na•
  17. 17. Sample Problem 9.1 Na ↑ 3s Na + 3p O ↑ 3s ↑↓ 2s 3p ↑↓ ↑↓ ↑↓ ↑↓ 2s 2Na+ + O2- 2p The formula is Na2O 9-17 ↑↓ ↑ 2p ↑
  18. 18. Figure 9.6 Na(s) 9-18 The exothermic formation of sodium bromide. Br2(l) NaBr(l)
  19. 19. Figure 9.7 9-19 The Born-Haber cycle for lithium fluoride.
  20. 20. Properties of Ionic Compounds • Ionic compounds tend to be hard, rigid, and brittle, with high melting points. • Ionic compounds do not conduct electricity in the solid state. – In the solid state, the ions are fixed in place in the lattice and do not move. • Ionic compounds conduct electricity when melted or dissolved. – In the liquid state or in solution, the ions are free to move and carry a current. 9-20
  21. 21. Figure 9.10 Solid ionic compound 9-21 Electrical conductance and ion mobility. Molten ionic compound Ionic compound dissolved in water
  22. 22. Table 9.1 Melting and Boiling Points of Some Ionic Compounds Compound mp (°C) bp (°C) CsBr 636 1300 NaI 661 1304 MgCl2 714 1412 KBr 734 1435 CaCl2 782 >1600 NaCl 801 1413 LiF 845 1676 KF 858 1505 2852 3600 MgO 9-22
  23. 23. Lewis Electron Dot Structures • Draw the Lewis electron dot structures for the reaction between Li and F • Draw the Lewis electron dot structures for the reaction between Mg and Cl 9-23 • Draw the Lewis electron dot structures for the reaction between Li
  24. 24. 9-24
  25. 25. 9-25
  26. 26. CHEMICAL BONDS • COVALENT BONDING – involves the sharing of one or more pairs of electrons between atoms – usually occurs when two nonmetals bond to each other – Represent the formation of an H2 molecule from two H atoms by Lewis electron dot formulas 9-26
  27. 27. Figure 9.12 9-27 Covalent bond formation in H2.
  28. 28. COVALENT BONDS • Other simple compounds with covalent bonds – Halogens-F2, Cl2, Br2, I2 – Halogen acids-HF, HCl, HBr, HI 9-28
  29. 29. Figure 9.13 Distribution of electron density in H2. At some distance (bond length), attractions balance repulsions. 9-29 Electron density is high around and between the nuclei.
  30. 30. Bonding Pairs and Lone Pairs Atoms share electrons to achieve a full outer level of electrons. The shared electrons are called a shared pair or bonding pair. The shared pair is represented as a pair of dots or a line: •• H H or H–H An outer-level electron pair that is not involved in bonding is called a lone pair, or unshared pair. •• 9-30 or •• •• F–F •• •• •• •• •• •• •• •• F F •• ••
  31. 31. Properties of a Covalent Bond The bond order is the number of electron pairs being shared by a given pair of atoms. A single bond consists of one bonding pair and has a bond order of 1. The bond energy (BE) is the energy needed to overcome the attraction between the nuclei and the shared electrons. The stronger the bond the higher the bond energy. The bond length is the distance between the nuclei of the bonded atoms. 9-31
  32. 32. Electronegativity and Bond Polarity A covalent bond in which the shared electron pair is not shared equally, but remains closer to one atom than the other, is a polar covalent bond. The ability of an atom in a covalent bond to attract the shared electron pair is called its electronegativity. Unequal sharing of electrons causes the more electronegative atom of the bond to be partially negative and the less electronegative atom to be partially positive. 9-32
  33. 33. POLAR AND NONPOLAR COVALENT BONDS • NONPOLAR COVALENT BONDS – involve the equal sharing of the electron pair between the two atoms – H:H nonpolar bond • POLAR COVALENT BONDS – involve the unequal sharing of the electron pair between the two atoms – H:Cl polar bond 9-33
  34. 34. Figure 9.20 Bonding between the models. Polar covalent bonds are much more common than either pure ionic or pure covalent bonds. 9-34
  35. 35. Figure 9.22 Boundary ranges for classifying ionic character of chemical bonds. 3.0 E N 2.0 0. 0 9-35
  36. 36. Figure 9.21 9-36 The Pauling electronegativity (EN) scale.
  37. 37. Trends in Electronegativity The most electronegative element is fluorine. In general electronegativity decreases down a group as atomic size increases. In general electronegativity increases across a period as atomic size decreases. Nonmetals are more electronegative than metals. 9-37
  38. 38. Figure 9.22 9-38 Electronegativity and atomic size.
  39. 39. Depicting Polar Bonds The unequal sharing of electrons can be depicted by a polar arrow. The head of the arrow points to the more electronegative element. A polar bond can also be marked using δ+ and δ- symbols. 9-39
  40. 40. Figure 9.23 Electron density distributions in H2, F2, and HF. In HF, the electron density shifts from H to F. The H–F bond has partial ionic character. 9-40
  41. 41. Sample Problem 9.4 Determining Bond Polarity from EN Values PROBLEM: (a) Use a polar arrow to indicate the polarity of each bond: N–H, F–N, I–Cl. (b) Rank the following bonds in order of increasing polarity: H–N, H–O, H–C. PLAN: (a) We use Figure 9.21 to find the EN values for each element. The polar arrow points toward the more electronegative element. (b) The greater the EN between the atoms, the more polar the bond. SOLUTION: (a) The EN values are: N = 3.0, H = 2.1; F = 4.0; I = 2.5, Cl = 3.0 N–H 9-41 F–N I–Cl
  42. 42. Sample Problem 9.4 (b) The EN values are: N = 3.0, H = 2.1; O = 3.5; C = 2.5 EN for H–N = 3.0 – 2.1 = 0.9 EN for H–O = 3.5 – 2.1 = 1.4 EN for H–C = 2.5 – 2.1 = 0.4 H-C < H-N < H-O 9-42
  43. 43. Predicting Polarity • Predict the polarity of each of the following compounds: • LiCl • CO • N2 9-43
  44. 44. 9-44
  45. 45. Metallic Bonding The electron sea model of metallic bonding proposes that: • All metal atoms in the sample contribute their valence electrons to form a delocalized electron “sea”. • The metal “ions” (nuclei with core electrons) lie in an orderly array within this mobile sea. • All the atoms in the sample share the electrons. • The metal is held together by the attraction between the metal “cations” and the “sea” of valence electrons. 9-45
  46. 46. Properties of Metals • Metals are generally solids with moderate to high melting points and much higher boiling points. – Melting points decrease down a group and increase across a period. • Metals can be shaped without breaking. – The electron sea allows the metal ions to slide past each other. • Metals are good conductors of electricity in both the solid and liquid states. – The electron sea is mobile in both phases. • Metals are good conductors of heat. 9-46
  47. 47. Table 9.5 Melting and Boiling Points of Some Metals Element bp ( C) Lithium (Li) 180 1347 Tin (Sn) 232 2623 Aluminum (Al) 660 2467 Barium (Ba) 727 1850 Silver (Ag) 961 2155 Copper (Cu) 1083 2570 Uranium (U) 9-47 mp ( C) 1130 3930
  48. 48. Figure 9.28 Melting points of the Group 1A(1) and Group 2A(2) metals. 9-48
  49. 49. Figure 9.29 9-49 Why metals dent and bend rather than crack.
  50. 50. WRITING LEWIS ELECTRON DOT STRUCTURES • N - A = S Rule • N = total number of electrons needed to achieve a rare gas configuration • A = total number of electrons available • S = total number of electrons shared – must be an even number – must be less than or equal to the number of electrons available 9-50
  51. 51. WRITING LEWIS ELECTRON DOT STRUCTURES 9-51 • The central atom in a molecule or polyatomic ion is the atom requiring the greatest number of additional electrons to achieve a rare gas configuration. • When deciding between two elements in the same group, the less electronegative atom is usually chosen. • H is never the central atom.
  52. 52. WRITING LEWIS ELECTRON DOT STRUCTURES • Write Lewis electron dot structures and dash formulas for water, H2O. • Write Lewis electron dot structures and dash formulas for ammonia, NH3. • Write Lewis electron dot structures and dash formulas for methane, CH4. 9-52
  53. 53. 9-53
  54. 54. 9-54
  55. 55. 9-55
  56. 56. Sample Problem 10.3 Writing Lewis Structures for Molecules with Multiple Bonds PROBLEM: Write Lewis structures for the following: (a) Ethylene (C2H4), the most important reactant in the manufacture of polymers (b) Nitrogen (N2), the most abundant atmospheric gas PLAN: After following steps 1 to 4 we see that the central atom does not have a full octet. We must therefore add step 5, which involves changing a lone pair to a bonding pair. SOLUTION: (a) C2H4 has 2(4) + 4(1) = 12 valence e-. H can have only one bond per atom. 9-56
  57. 57. Sample Problem 10.3 (b) N2 has 2(5) = 10 valence e-. 9-57
  58. 58. WRITING LEWIS ELECTRON DOT STRUCTURES • Write a Lewis electron dot structure for oxygen, O2. • Write a Lewis electron dot structure for nitrogen, N2. • Write a Lewis electron dot structure for carbon dioxide, CO2. 9-58
  59. 59. 9-59
  60. 60. 9-60
  61. 61. 9-61
  62. 62. Resonance Structures O3 can be drawn in 2 ways: These are two different reasonable Lewis structures for the same molecule. Neither structure depicts O3 accurately, because in reality the O-O bonds are identical in length and energy. 9-62
  63. 63. Resonance Structures The structure of O3 is shown more correctly using both Lewis structures, called resonance structures. A two-headed resonance arrow is placed between them. Resonance structures have the same relative placement of atoms but different locations of bonding and lone electron pairs. 9-63
  64. 64. The Resonance Hybrid A species like O3, which can be depicted by more than one valid Lewis structure, is called a resonance hybrid. Resonance forms are not real bonding depictions. O3 does not change back and forth between its two resonance forms. The real structure of a resonance hybrid is an average of its contributing resonance forms. 9-64
  65. 65. Sample Problem 10.4 Writing Resonance Structures PROBLEM: Write resonance structures for the nitrate ion, NO3− and find the bond order. PLAN: Write the Lewis structure, remembering to add 1e- to the total number of valence e- for the -1 charge. We apply Step 5 to form multiple bonds. Since multiple bonds can be formed in more than one location, there are resonance forms. SOLUTION: Nitrate has [1 x N(5e-)] + [3 x O(6e-)] + 1e-] = 24 valence eAfter Steps 1-4: 9-65
  66. 66. Sample Problem 10.4 Step 5. Since N does not have a full octet, we change a lone pair from O to a bonding pair to form a double bond. 4 shared electron pairs Bond order = = 1⅓ 3 bonded-atom pairs 9-66
  67. 67. Formal Charge Formal charge is the charge an atom would have if all electrons were shared equally. Formal charge of atom = # of valence e- - (# of unshared valence e- + ½ # of shared valence e-) For OA in resonance form I, the formal charge is given by 6 valence e- - (4 unshared e- + ½(4 shared e-) = 6 – 4 – 2 = 0 9-67
  68. 68. Formal Charge Formal charges must sum to the actual charge on the species for all resonance forms. OA [6 – 4 – ½(4)] = 0 OB [6 – 2 – ½(6)] = +1 OC [6 – 6 – ½(2)] = -1 OA [6 – 6 – ½(2)] = -1 OB [6 – 2 – ½(6)] = +1 OC [6 – 4 – ½(4)] = 0 For both these resonance forms the formal charges sum to zero, since O3 is a neutral molecule. 9-68
  69. 69. Choosing the More Important Resonance Form • Smaller formal charges (positive or negative) are preferable to larger ones. • The same nonzero formal charges on adjacent atoms are not preferred. Avoid like charges on adjacent atoms. • A more negative formal charge should reside on a more electronegative atom. 9-69
  70. 70. Example: NCO− has 3 possible resonance forms: +2 0 -1 -1 0 0 0 0 -1 Resonance forms with smaller formal charges are preferred. Resonance form I is therefore not an important contributor. A negative formal charge should be placed on a more electronegative atoms, so resonance form III is preferred to resonance form II. The overall structure of the NCO- ion is still an average of all three forms, but resonance form III contributes most to the average. 9-70
  71. 71. WRITING LEWIS ELECTRON DOT STRUCTURES • RESONANCE STRUCTURES are used to represent molecules when two or more dot representations are needed to describe the molecule. • Write Lewis electron dot structures for the resonance forms expected for sulfur trioxide, SO3. • Write dash formulas for sulfur trioxide. 9-71
  72. 72. 9-72
  73. 73. 9-73
  74. 74. Exceptions to the Octet Rule Molecules with Electron-Deficient Atoms B and Be are commonly electron-deficient. Odd-Electron Species A molecule with an odd number of electrons is called a free radical. 9-74
  75. 75. Exceptions to the Octet Rule Expanded Valence Shells An expanded valence shell is only possible for nonmetals from Period 3 or higher because these elements have available d orbitals. 9-75
  76. 76. MOLECULES THAT DO NOT OBEY THE N-A=S RULE • COVALENT COMPOUNDS OF Be • COVALENT COMPOUNDS OF B • Draw the Lewis electron dot structure for the BBr3 molecule. • Draw the Lewis electron dot structure for the BeCl2 molecule. 9-76
  77. 77. 9-77
  78. 78. 9-78

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