Chapter 9 Pptrevised


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Chapter 9 Pptrevised

  1. 1. Chapter 9 Notes Part I Covalent Bonding
  2. 2. Obj. 1…Octet Rule <ul><li>Recall that all elements want 8 valence e-. </li></ul><ul><li>metals will lose their valence e-. </li></ul><ul><li>non-metals will gain more valence e-. </li></ul>- (+) cations - (-) anions <ul><li>noble gases do not gain or lose e-. </li></ul>- happy already! *** atoms need other atoms in order to lose/gain e-. - electronegativity differences determine bond. - metals = low / non-metals = high electroneg.
  3. 3. Obj. 2…Ionic Bonds <ul><li>transfer of e- b/n (+) cation and (-) anion. </li></ul>- compounds created are neutral - 10% of all bonds - strongest bond - forms crystalline solids - highest melting/boiling pts. - can conduct electricity in molten states.
  4. 4. Obj. 3…Pure Covalent Bonds <ul><li>equal sharing of e- b/n non-metals ( molecule ) . </li></ul>- compounds created are neutral - 2% of all bonds - weakest bond - typically form gases - lowest melting/boiling pts. - do NOT conduct electricity!
  5. 5. Obj. 4…Polar Covalent Bonds <ul><li>un-equal sharing of e- b/n non-metals ( molecule ) . </li></ul>- compounds created are neutral - 88% of all bonds - intermediate bond strength - typically form soft solids and liquids - intermediate melting/boiling pts. - some conduct electricity in molten states.
  6. 6. Obj. 6…Metallic Bonds <ul><li>cations (+) in metals are surrounded by a sea of </li></ul>mobile e-…(‘ delocalized e- ’). <ul><li>metals held together by attraction b/n e- and </li></ul>- cushion b/n cations cations.
  7. 7. Obj. 6 cont… <ul><li>good conductors of electricity </li></ul>- as e- enter one end, an = # of e- leave the other <ul><li>ductile </li></ul>- can be drawn into a wire <ul><li>Because of mobile ‘delocalized’ e-, metals are… </li></ul><ul><li>malleable </li></ul>- can be hammered
  8. 8. Obj. 7…Bonding Vocabulary <ul><li>ion : </li></ul>atom w/ a charge. - has lost or gained e- <ul><li>valence electron : </li></ul>all e- in outer ‘s’ and ‘p’ orbitals. <ul><li>dipole : </li></ul>molecule w/ one slightly (-) end and one slightly (+) end. - occurs in polar covalent bonds <ul><li>polyatomic ion : </li></ul>a group of atoms (covalently bonded), that act as a single atom - NH 4 +1 , SO 4 -2 , ClO 3 -1 <ul><li>molecule : </li></ul>atoms covalently bonded together.
  9. 9. Covalent Bonding <ul><li>A covalent bond occurs between two non-metals </li></ul><ul><li>Electrostatic bonding does not occur—in other words, there is no “give and take” of electrons </li></ul>
  10. 10. It ends up being a “tug of war” of electrons Where the electrons end up somewhere in the middle.
  11. 11. Single Bonds <ul><li>A single bond occurs when one pair of electrons is shared by two atoms. </li></ul><ul><li>This pair of bonded electrons is called a shared pair . </li></ul>
  12. 12. Double and Triple Bonds <ul><li>Double bonds occur when two atoms have two shared pair of electrons </li></ul><ul><li>Triple bonds occur when two atoms share three pair of electrons </li></ul>
  13. 13. Sigma and Pi Bonds <ul><li>How do covalent bonds form between two elements? </li></ul><ul><li>By the combining of their p orbitals. </li></ul>
  14. 14. Sigma and Pi Bonds <ul><li>The first bond between two atoms is a sigma bond (  ) and it forms because of end-to-end overlap of p orbitals. </li></ul>
  15. 15. Sigma and Pi Bonds <ul><li>Any other bonds between the same atoms would be a pi bond (  ) and they form because of side-to-side overlap of p orbitals. </li></ul>
  16. 16. Sigma and Pi Bonds <ul><li>Single bond = 1 sigma bond </li></ul><ul><li>Double bond = 1 sigma+1 pi bond </li></ul><ul><li>Triple bond = 1 sigma + 2 pi bonds </li></ul>
  17. 17. Sigma and Pi Bonds <ul><li>How many sigma and pi bonds do the following compounds have? </li></ul>NH 3 CH 2 O SiO 2
  18. 18. Strength of Covalent Bonds <ul><ul><li>The strength of a covalent bond depends on how much distance separates the bonded nuclei. </li></ul></ul><ul><ul><li>This is known as bond length. </li></ul></ul><ul><ul><li>Determined by size of the atoms and how many electrons pairs are shared. </li></ul></ul>
  19. 19. <ul><ul><li>The shorter the bond length, the stronger the bond. </li></ul></ul><ul><ul><li>Single bonds, such as F 2 , are weaker than double bonds, such as those in O 2. </li></ul></ul><ul><ul><li>Double bonds are weaker than triple bonds, such as N 2 . </li></ul></ul>
  20. 20. Naming Molecules 9.2 <ul><li>Rules to name binary molecular compounds. </li></ul><ul><li>1. The first element in the formula is always named first, using the entire element name. </li></ul><ul><li>2. The second element in the formula is named using the root of the element and adding the suffix –ide . </li></ul><ul><li>3. Prefixes are used to indicate the number of atoms of each type that are present in the compound. </li></ul>
  21. 21. Prefixes in Covalent Compounds <ul><li>1 mono- 2 di- 3 tri- </li></ul><ul><li>4 tetra- 5 penta- 6 hexa- </li></ul><ul><li>7 hepta- 8 octa- 9 non- </li></ul><ul><li>10 deca </li></ul><ul><li>One exception – the first element in the formula never uses the prefix mono-. </li></ul><ul><li>When the element name begins with a vowel drop the final letter in the prefix. CO carbon monoxide not monocarbon monooxide </li></ul>
  22. 22. Naming Acids binary acids <ul><li>A binary acid contains hydrogen and one other element. </li></ul><ul><li>Use the prefix hydro- to name the hydrogen. </li></ul><ul><li>Use the root of the second element plus the suffix -ic acid. HBr hydrobromic acid </li></ul>
  23. 23. Naming Oxyacids <ul><li>Contains an oxyanion referred to as oxyacids. </li></ul><ul><li>If anion suffix is -ate, it is replaced with the suffix -ic. HNO 3 (nitrate ion) nitric acid </li></ul><ul><li>If anion suffix is -ite, it is replaced with the suffix -ous. HNO 2 (nitrite ion) nitrous acid </li></ul>
  24. 24. Writing Formulas from Names <ul><li>The name of any binary molecule allows you to write the correct formula with ease. </li></ul><ul><li>Subscripts are determined from the prefixes used in the name because the name indicates the exact number of each atom present in the molecule. Carbon Monoxide CO </li></ul>
  25. 25. Coordinate Covalent Bonding <ul><li>A coordinate covalent bond occurs when one atom gives both electrons to a shared pair between them. </li></ul>
  26. 26. Coordinate Covalent Bonding <ul><li>In a line diagram, instead of a straight line showing a shared pair, the coordinate covalent bond is shown as an arrow. </li></ul>
  27. 27. Resonance <ul><li>If you draw the covalent bonding diagram of some compounds there is not only one way to draw it. </li></ul>
  28. 28. Resonance <ul><li>There are two valid ways to show the structure of nitrite. The real structure is an average of the two, and this is called a resonance structure . </li></ul>
  29. 29. Resonance, cont. <ul><li>Earlier chemists thought the compound just quickly flipped back and forth between the two structures (or resonated). </li></ul><ul><li>This is proven to be untrue now using bond lengths. </li></ul>
  30. 31. An Addendum to Lewis Structures <ul><li>Carbon and silicon are exceptions to the pattern of how to place electrons in a Lewis Dot Structure. </li></ul>
  31. 32. An Addendum to Lewis Structures <ul><li>This is because they have hybrid orbitals (where the s and p sublevels blend together and have four equal energy orbitals.) </li></ul>
  32. 33. Drawing Dot Diagrams <ul><li>A dot diagram shows the valence electrons for an atom </li></ul><ul><li>First, write the symbol for the element. </li></ul><ul><li>Then, put a dot to show each electron in the outer shell, in the following order: </li></ul>X 1 2 8 5 7 4 6 3
  33. 34. The reason for this order, is the first two spots for electrons represent the one orbital in the s sublevel, the last three spots represent the three orbitals of the p sublevel.
  34. 35. Your book’s way to Draw Lewis Dot Structures <ul><li>1. Predict the location of certain atoms. </li></ul><ul><li>a. Hydrogen is always a terminal. </li></ul><ul><li>b. The atom with the least attraction for shared electrons in the molecule is the central atom. </li></ul><ul><li>2. Find the total number of electrons available for bonding. This total is the number of valence electrons in the atoms in the molecule. </li></ul>
  35. 36. <ul><li>3. Determine the number of bonding pairs by dividing the number of electrons available for bonding by two. </li></ul><ul><li>4. Place one bonding pair (single bond) between the central atom and each of the terminal atoms. </li></ul><ul><li>5. Subtract the number of pairs you used in step 4 from the number of bonding pairs you determined in step 3. </li></ul>
  36. 37. <ul><li>6. If the central atom is not surrounded by four electron pairs, it does not have an octet. You must convert one or two of the lone pairs on the terminal atoms to a double bond or a triple bond between the terminal atom and the central atom </li></ul>
  37. 38. Exceptions to the Octet Rule <ul><li>Some molecules and ions do not obey the octet rule. Three reasons exist for these exceptions. </li></ul><ul><li>1. A small group of molecules has an odd number of valence electrons and cannot form an octet around each atom. </li></ul><ul><li>2. Coordinate covalent bond. </li></ul>
  38. 39. <ul><li>3. The third group of compounds that does not follow the octet rule has central atoms that contain more than eight valence. This electron arrangement is referred to as an expanded octet . </li></ul>
  39. 40. Covalent Bonding and Polyatomic Ions <ul><li>A polyatomic ion is really just a charged molecule that bonds ionically. </li></ul><ul><li>The charge signifies how many electrons are given/taken away. </li></ul>
  40. 41. Practice <ul><li>CN - </li></ul><ul><li>SiF 4 </li></ul><ul><li>C 2 H 4 </li></ul><ul><li>NBr 3 </li></ul>
  41. 42. Chapter 16 Notes Part II VSEPR Theory
  42. 43. The VSEPR Theory <ul><li>VSEPR stands for Valence Shell Electron Pair Repulsion theory </li></ul><ul><li>All atoms have electrons orbiting the nucleus </li></ul>
  43. 44. The VSEPR Theory <ul><li>This creates several different shapes that molecules can be in. </li></ul><ul><li>These like charges repel, causing the atoms bonded to the central atom to move as far away from each other as possible. </li></ul>
  44. 45. Shapes not involving unshared pairs: <ul><li>The following shapes come about from compounds where the central atom does not have unshared pairs of electrons: </li></ul>
  45. 46. Shapes not involving unshared pairs: <ul><li>Linear </li></ul><ul><li>Trigonal Planar </li></ul><ul><li>Tetrahedral </li></ul><ul><li>Trigonal Bipyramidal </li></ul>
  46. 47. Linear <ul><li>2 atoms off central, no unshared pairs </li></ul><ul><li>Ex: CO 2 </li></ul><ul><li>Bond Angle: 180 o </li></ul>
  47. 48. Trigonal Planar <ul><li>3 Atoms off central, no unshared pair </li></ul><ul><li>Ex: BF 3 </li></ul><ul><ul><ul><li>Bond Angle-120 o </li></ul></ul></ul>
  48. 49. Tetrahedral <ul><li>4 Atoms off central, no unshared pair </li></ul><ul><li>Ex: CH 4 </li></ul><ul><li>Bond Angles—109.5 o </li></ul>
  49. 50. Trigonal Bipyramidal <ul><li>5 Atoms off central, no unshared pair </li></ul><ul><li>Ex: PCl 5 </li></ul><ul><li>2 Bond Angles—120 o and 90 o </li></ul>
  50. 51. Shapes that involve an unshared pair: <ul><li>These shapes are altered by one or more unshared pair on the central atom. </li></ul><ul><ul><li>Bent </li></ul></ul><ul><ul><li>Pyramidal </li></ul></ul>
  51. 52. Bent <ul><li>2 Atoms off central, 1 or 2 unshared pair </li></ul><ul><li>Ex: H 2 O </li></ul><ul><li>Bond Angles—105 o </li></ul>
  52. 53. Pyramidal <ul><li>3 Atoms off central, 1 unshared pair </li></ul><ul><li>Ex: NH 3 </li></ul><ul><li>Bond Angles—107 o </li></ul>
  53. 54. Obj. 8… Intra molecular Bonds vs. Inter molecular Attraction <ul><li>Intramolecular bonds : </li></ul>- bonds w/in a molecule - i.e. holds H and O together in water (H 2 O) - ionic, covalent, polar covalent and metallic <ul><li>Intermolecular attraction : </li></ul>- attraction b/n two or more different molecules. - i.e. holds H 2 O molecules together in a pond. ** always weaker than intramolecular bonds. - strength indicates solid, liquid or gas.
  54. 55. Obj. 9-10…Intermolecular Attractions <ul><li>three types… </li></ul>- dispersion forces - dipole interactions - Hydrogen bonds van der Waals forces ( weakest ) (strongest) <ul><li>dispersion forces (DF) … </li></ul>- seen in Halogen diatomic molecules (Z 2 ) - strength of force as # of e- in molecules - F and Cl have few e- = weak DF (gases) - Br has more e- = stronger DF (liquid) - I has most e- = strongest DF (solid)
  55. 56. Obj. 9-10 cont… <ul><li>dipole interactions … </li></ul>- slightly (-) pole attracted to (+) pole of another molecule. - similar to but much weaker than ionic bonds
  56. 57. Obj. 9-10 cont… <ul><li>Hydrogen bonds … </li></ul>- strong attraction b/n H on one polar molecule and the e- of an electronegative atom (N,O, or F) from a different molecule. - H atom is forced to share its only e- - H is therefore very attracted to pair of e- in neighboring molecule.
  57. 58. Obj. 9-10 cont… <ul><li>bond strengths… </li></ul>dispersion forces dipole interactions Hydrogen bonds Covalent bonds Metallic bonds Ionic bonds van der Waals inter molecular intra molecular WEAKEST STRONGEST (Pure and Polar)
  58. 59. Chapter 16 Notes, part III Bond Polarity and Molecular Polarity
  59. 60. Types of Bonds <ul><li>Up until now, we have assumed that there are two types of bonds: Covalent and Ionic. </li></ul><ul><li>This is true, but covalent bonds can be broken into two categories </li></ul>
  60. 61. Nonpolar Covalent <ul><li>In nonpolar covalent bonding electrons are shared equally . </li></ul><ul><li>Electrons spend an equal amount of time with both elements in the bond. </li></ul><ul><li>Typical in diatomic elements: </li></ul><ul><li> Br 2 , I 2 , N 2 , Cl 2 , H 2 O 2 , F 2 , </li></ul>
  61. 62. Polar Covalent <ul><li>In polar covalent bonding, electrons are still shared, but they are shared unequally . </li></ul><ul><li>This is due to one nucleus pulling the shared pair harder than the other. </li></ul><ul><li>This creates a dipole —a bond where one side is slightly positive and the other is slightly negative. </li></ul>
  62. 63. <ul><li>A dipole is caused because the electron spends more time on one side than the other. </li></ul><ul><li>The polarity of the bond is shown like: </li></ul>Polar Covalent H—Cl H—Cl  +  - OR
  63. 64. Ionic Bonds <ul><li>In an ionic bond, electrons are transferred . </li></ul><ul><li>The nucleus of one element pulls hard enough to take electrons away completely. </li></ul>
  64. 65. How can you tell what kind of bond there is? <ul><li>By looking at the difference in electronegativity! </li></ul><ul><li>Remember, electronegativity is the tendency of an atom to attract an electron when bonding. </li></ul><ul><li>The more electronegative, the more it will pull electrons. </li></ul>
  65. 66. Ionic Bonds - highest electronegativity diff. (ED  1.67) Ex… NaCl: Na + = 0.9 Cl - = 3.0 3.0 – 0.9 = KBr: K + = 0.8 Br - = 2.8 2.8 – 0.8 = 2.1 2.0
  66. 67. Polar Covalent Bonds - lowest electronegativity diff. (ED  0.4) Ex… H 2 : H + = 2.1 H + = 2.1 2.1 – 2.1 = 0 CH 4 : C = 2.5 H + = 2.1 2.5 – 2.1 = 0.4
  67. 68. Nonpolar Covalent Bonds - int. electronegativity diff. (ED 0.41 - 1.66) Ex… H 2 O: H = 2.1 O = 3.5 3.5 – 2.1 = 1.4 CO 2 : C = 2.5 O = 3.5 3.5 – 2.5 = 1.0
  68. 69. Nonpolar Bonds: EN Diff ≤ 0.4 Polar Bonds: 0.4 < EN Diff ≤ 1.7 Ionic Bonds: EN Diff > 1.7
  69. 70. Practice <ul><li>Identify bond types… </li></ul>HBr KCl CO 2 2.8 – 2.1 = 0.7 PC 3.0 – 0.8 = 2.2 I 3.5 – 2.5 = 1.0 PC Li 2 O 3.5 – 1.0 = 2.5 I Br 2 2.8 – 2.8 = 0 <ul><li>Rank polarities… 1 = least polar… 5 = most polar </li></ul>1 2 3 4 5 C
  70. 71. What type of bond is between: <ul><li>H and Cl </li></ul><ul><li>Li and Cl </li></ul><ul><li>C and S </li></ul><ul><li>F and O </li></ul>
  71. 72. What type of bond is between: <ul><li>N and Br </li></ul><ul><li>Na and F </li></ul><ul><li>C and O </li></ul>
  72. 73. Polar Molecules <ul><li>If the slightly positive and slightly negative ends of polar bonds can collect on two different sides of a molecule, it can make an entire molecule polar. </li></ul>
  73. 74. Polar Molecules <ul><li>If bonds are nonpolar, a molecule will always be nonpolar. </li></ul>
  74. 75. Polar Molecules <ul><li>If bonds are polar and the shape of a molecule is symmetrical, the molecule will be nonpolar because the charges cancel out. </li></ul><ul><li>(Linear, trigonal planar, tetrahedral and trigonal bipyramidal are the symmetrical shapes we talked about.) </li></ul>
  75. 76. Polar Molecules <ul><li>If bonds are polar and the molecule is asymmetrical, the molecule will be polar! </li></ul><ul><li>(Bent and pyramidal molecules are asymmetrical) </li></ul>