Bonding

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Bonding

  1. 1. Bonds <ul><li>Forces that hold groups of atoms together and make them function as a unit. </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  2. 2. Bond Energy <ul><li>It is the energy required to break a bond. </li></ul><ul><li>It gives us information about the strength of a bonding interaction. </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  3. 3. Ionic Bonds <ul><li>Formed from electrostatic attractions of closely packed, oppositely charged ions. </li></ul><ul><li>Metal + nonmetal </li></ul><ul><li>Formed when an atom that easily loses electrons reacts with one that attracts electrons </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  4. 4. Achieving Noble Gas Electron Configurations (NGEC) <ul><li>Two nonmetals react: They share electrons to achieve NGEC. </li></ul><ul><li>A nonmetal and a representative group metal react (ionic compound): The valence orbitals of the metal are emptied to achieve NGEC. The valence electron configuration of the nonmetal achieves NGEC. </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  5. 5. Isoelectronic Ions <ul><li>Ions containing the the same number of electrons </li></ul><ul><li>(O 2  , F  , Na + , Mg 2+ , Al 3+ ) </li></ul><ul><li>O 2  > F  > Na + > Mg 2+ > Al 3+ </li></ul><ul><ul><li>largest smallest </li></ul></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  6. 6. Figure 12.8: Ions as packed spheres. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  7. 7. Figure 12.8: Positions (centers) of the ions. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  8. 8. Figure 12.9: Relative sizes of some ions and their parent atoms. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  9. 9. Covalent Bonds <ul><li>Formed when electrons are shared by nuclei </li></ul><ul><li>Overlapping of orbitals to share electrons </li></ul><ul><li>Only diatomics are truly covalent! They share the pair of electrons equally </li></ul><ul><li>All other covalent bonds are polar covalent – not an equal sharing of pair of electrons </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  10. 10. Polarity <ul><li>Unequal sharing of electrons: A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar , or to have a dipole moment . </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  11. 11. Figure 12.2: Probability representations of the electron sharing in HF. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  12. 12. Electronegativity <ul><li>The ability of an atom in a molecule to attract shared electrons to itself. </li></ul><ul><li>One atom pulls the electron in closer to its nucleus than the other. </li></ul><ul><li>Trend: increase across a period due to increased effective charge; decreases down a period due to distance from nucleus </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  13. 13. Figure 12.3: Electronegativity values for selected elements. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  14. 14. Table 12.1 Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  15. 15. Figure 12.4: The three possible types of bonds . Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  16. 16. Polar bonds vs. Polar molecules <ul><li>Molecules with polar bonds can be nonpolar overall when the polarity of the bonds cancels out </li></ul><ul><li>Methane CH 4 </li></ul><ul><li>Carbon Dioxide CO 2 </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  17. 17. Figure 12.5: Charge distribution in the water molecule. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  18. 18. Figure 12.5: Water molecule behaves as if it had a positive and negative end. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  19. 19. Figure 12.6: Polar water molecules are strongly attracted to positive ions by their negative ends. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  20. 20. Figure 12.6: Polar water molecules are strongly attracted to negative ions by their positive ends. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  21. 21. Lewis Structure <ul><li>Shows how valence electrons are arranged among atoms in a molecule. </li></ul><ul><li>Reflects central idea that stability of a compound relates to noble gas electron configuration. </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  22. 22. Octet Rule <ul><li>To be stable, an atom in a molecule needs 8 total electrons </li></ul><ul><li>These electrons come in pairs </li></ul><ul><li>The pairs can be shared with another atom, or unshared </li></ul><ul><li>H and He have their own rule – the Duet Rule </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  23. 23. Comments About the Octet Rule <ul><li>2nd row elements C, N, O, F observe the octet rule . </li></ul><ul><li>2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive. </li></ul><ul><li>3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals. </li></ul><ul><li>When writing Lewis structures, satisfy octets first , then place electrons around elements having available d orbitals . </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  24. 24. Drawing Lewis Structures <ul><li>Count the total number of valence electrons of all the elements in the compound. Add electrons for negative ions; subtract electrons for positive ions. </li></ul><ul><li>Use pairs of electrons to attach each element to the central atom (usually first atom listed) </li></ul><ul><li>Give the central atom unshared pairs of electrons to satisfy the octet rule </li></ul><ul><li>Assign the rest of the electron pairs to the other atoms (divide up evenly) </li></ul><ul><li>If there are not enough electrons, pull pairs from the central atom to create double or triple bonds </li></ul><ul><li>If there are too many electrons, place the pairs on the central atom </li></ul>
  25. 25. Pairs of electrons… <ul><li>Shared – between atoms (bonding pairs) </li></ul><ul><li>Unshared – not between atoms (lone pairs) </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  26. 26. Multiple Bonds <ul><li>When you don’t have enough valence electrons to give all atoms a NGEC, then you might need to have a double or triple bond between atoms. </li></ul><ul><li>These atoms CANNOT double bond: H, F, Cl, Br, I. This is due to only one available space in the shared orbital. </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  27. 27. Resonance <ul><li>Occurs when more than one valid Lewis structure can be written for a particular molecule. </li></ul><ul><li>These are resonance structures . The actual structure is an average of the resonance structures. </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  28. 28. VSEPR Model <ul><li>The structure around a given atom is determined principally by minimizing electron pair repulsions. </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  29. 29. Predicting a VSEPR Structure <ul><li>Draw Lewis structure. </li></ul><ul><li>Put pairs as far apart as possible. </li></ul><ul><li>Determine positions of atoms from the way electron pairs are shared. </li></ul><ul><li>Determine the name of molecular structure from positions of the atoms. </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  30. 30. Figure 12.12: Molecular structure of methane. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  31. 31. Figure 12.13: Tetrahedral arrangement of electron pairs. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  32. 32. Figure 12.13: Hydrogen atoms occupy only three corners of the tetrahedron. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  33. 33. Figure 12.13: The NH 3 molecule has the trigonal pyramid structure. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  34. 34. Figure 12.14: Tetrahedral arrangement of four electron pairs around oxygen. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  35. 35. Figure 12.14: Two electron pairs shared between oxygen and hydrogen atoms. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  36. 36. Figure 12.14: V-shaped molecular structure of the water molecule. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  37. 37. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  38. 38. <ul><li>Classification of shapes: # of e- pairs # of e- pairs # of e- pairs shape bond angle examples (total) (bonding) (nonbonding) ........................................................................... 1 1 0 linear 180 H-Cl ............................................................................ 2 2 0 linear 180 CO 2 ............................................................................. 3 3 0 trigonal 120 BF 3 planar 3 2 1 bent or < 120 SO 2 V-shaped ............................................................................. 4 4 0 tetrahedral 109.5 CH 4 4 3 1 trigonal < 109.5 NH 3 pyramid (107 for NH3) 4 2 2 bent or < 120 H 2 S V-shaped 4 1 3 linear 180 H-Cl </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  39. 39. <ul><li>Classification of shapes: # of e- pairs # of e- pairs # of e- pairs shape bond angle examples (total) (bonding) (nonbonding) .............................................................................................................................................. </li></ul><ul><li>5 5 0 trigonal 90 (axial) PCl 5 bipyramid 120 (equatorial) 5 4 1 seesaw or 90 (a) SF 4 teetertotter < 120 (eq) 5 3 2 T-shaped 90 ClF 3 5 2 3 linear 180 XeF 2 ............................................................................. 6 6 0 octahedral or 90 SF 6 square bipyramid 6 5 1 square-based 90 BrF 5 pyramid 6 4 2 square planar 90 XeF 4 ????????????????????????????????????????????????????????????????????????????? ? 6 ? 3 ? 3 ? T-shaped 90 ??? ? 6 ? 2 ? 4 ? linear 180 ??? </li></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.
  40. 40. Shapes of Molecule Practice <ul><li>Find this site (link on e-chalk) </li></ul><ul><ul><li>http://employees.oneonta.edu/viningwj/sims/ </li></ul></ul><ul><ul><li>Go through all simulations on Chapter 9 </li></ul></ul>Copyright©2000 by Houghton Mifflin Company. All rights reserved.

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