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  • 1. VSEPR Valence Shell Electron Pair Repulsions
  • 2. Covalent Bond: A Model
    • Chemical bonds can be viewed as forces that cause a group of atoms to act as a unit
    • They result from the tendency of a system to seek its lowest possible energy
    • Bonds occur when collections of atoms are more stable (lower in energy) than the separate atoms
    • Note: The next three slides will repeat at the end. This is preliminary intro info that may make more sense at the end.
  • 3. Example: Methane
    • 1652 kJ of energy are required to break a mole of methane into separate C and H atoms
    • OR 1652 kJ of energy are released when one mole of methane is formed from one mole of C atoms and four moles of H atoms
    • Methane is therefore a stable molecule relative to its stable atoms
    • Since there are four H atoms arranged around the central C, it is natural to envision four individual attractions between C and H (bonds)
    • Each bond has an associated bond energy, found by dividing the total energy by four (1652/4 = 413 kJ)
    • The positive Bond Energy value indicates the energy required to break the bond between C and H atoms
  • 4. Bonding Model
    • Models originate from our observations of the properties of nature
    • Atoms can form stable groups by sharing electrons, shared electrons give a lower energy state because simultaneously attracted to two nuclei
    • Remember: Models are human inventions that allow us to explain and predict. A model is a useful way of thinking; they include simplifications and assumptions. A model does not equal reality.
  • 5. Covalent Bonds
    • Electron pair(s)
      • shared between two atoms
      • attracted to both nuclei
    • Location of a single shared pair
      • Directly between two nuclei
      • Maximizes attractions with shortest distance between two positive nuclei
      • Minimizes repulsions with negative electrons between positive nuclei that would repel one another
  • 6. Multiple covalent bonds around the same atom determine the shape
    • Negative e- pairs with same charge repel each other
    • Repulsions push the pairs as far apart as possible
  • 7. Single bonds
    • Sigma bond ө
    • Overlap of orbitals allow electron pair to be shared between the two atoms
    • Electron pair shared directly between two nuclei
    • Only one pair may be shared in this space - just as only one pair of electrons may occupy a single atomic orbital
  • 8. Double and Triple Bonds
    • Pi bonds π
    • Since the space between the nuclei is occupied, e- pair is shared above and below the plane or front and back
    • Overlap of p-orbital lobes allow for this sharing above and below OR front and back
  • 9.  
  • 10. Lewis Structures
    • Drawn to show the bonds between the atoms in the structure
    • Only shows whether single, double or triple bonds
    • Does not show the shape
  • 11. Lewis Structure
    • Represents the arrangement of valence electrons among atoms in the molecule
    • Rules based upon observations of thousands of molecules, which show that in most stable compounds the atoms achieve noble gas configurations
    • Duet Rule
      • hydrogen stable with only a pair of e-
    • Octet Rule
      • other atoms stable with 4 pairs of e-
  • 12. Rules for Drawing Lewis Structures
    • Count the number of valence electrons
    • Draw the skeleton structure- the central is generally listed first in formula
    • Distribute electrons to give each atom a stable octet
    • Reconcile # e-
      • Do you have enough electrons? You may need to use double or triple bonds.
      • Do have too many electrons? You may need to explain the octet, but only if empty d-orbital available
  • 13. Determine # Valence e- from column #
  • 14.  
  • 15.  
  • 16.  
  • 17.  
  • 18. Electron Clouds repel each other, thus structure around an atom is determined principally by minimizing repulsions
  • 19. 2 electron pairs (2 EP) around central atom
    • Two clouds pushed as far apart as possible
      • Greatest angle possible 180 º
      • LINEAR shape
  • 20. 3 electron pairs (3 EP) around central atom
    • Three clouds pushed as far apart as possible
      • Greatest angle possible 120 º
      • TRIGONAL PLANAR shape
      • (3) (flat)
  • 21. 4 electron pairs (4 EP) around central atom
    • Four clouds pushed as far apart as possible
      • Greatest angle no longer possible
      • in two dimensions
      • Requires three-dimensional
      • TETRAHEDRAL shape
  • 22. Orbital Hybridization #1
    • Atomic orbitals such as s and p are not well suited for overlapping and allowing two atoms to share a pair of electrons
    • Remember: best location of shared pair is directly between two atoms
    • e- pair spends little time in best location
      • With overlap of two s-orbital
      • With overlap of two p-orbitals
  • 23. Orbital Hybridization #2
    • Hybrid orbitals (cross of atomic orbitals)
    • Remember: The pink flower hybrid cross of the red and white flower
    • Hybrid orbitals
      • Shape more suitable for bonding
        • One large lobe and one very small lobe
        • Large lobe oriented towards other nucleus
      • Angles more suitable for bonding
        • Angles predicted from VSEPR
  • 24. Overlap of two s-orbitals NOT A GOOD LOCATION- far from one nucleus Note: shared in this overlap the e- pair would spend most of the time in an unfavorable location GOOD SPOT between both nuclei
  • 25. Overlap of two p-orbitals One atom & its p-orbital The other atom & its p-orbital represents the nucleus BAD location far from other nucleus GOOD SPOT between both nuclei
  • 26. Hybrid Orbitals yield more favorable shape for overlap
    • Atomic orbitals are not shaped to maximize attractions nor minimize repulsions
    • Hybrid orbital shape
      • One large lobe oriented towards other atom
      • Notice the difference in this shape compared to p-orbital shape
  • 27. Angles and Shape
    • Atomic orbitals are not shaped to maximize attractions nor minimize repulsions
    • BUT the angles are also not favorable
      • p-orbitals are oriented at 90 º to
      • each other
      • Other angles are required 180º,
      • 120º or 109.5º
  • 28. Orbital Hybridization #3
    • Each e-pair requires a hybrid orbital
    • If two hybrid orbitals required than two atomic orbitals must be hybridized, an s and a p orbital forming two sp orbitals at 180 º
    sp hybrids 2 EP 4 EP 3 EP sp 2 hybrids sp 3 hybrids
  • 29. Electron-Pair Geometry vs Molecular Geometry
    • Electron-pair geometry
      • Where are the electron pairs
      • Includes
        • bonding pairs (BP) – shared between 2 atoms
        • nonbonding pairs (NBP) – lone pair
    • Molecular geometry
      • Where are the atoms
      • Includes only the bonding pairs
  • 30. Examples of 3 EP
    • 3 BP + 0 NBP = 3 EP
      • 3 EP = EP geom is trigonal planar
      • All locations occupy by an atom,
      • so molecular geometry is also trigonal planar
    • 2 BP + 1 NBP = 3 EP
      • 3 EP = EP geom is trigonal planar
      • Only two bonding pairs
      • One of the locations is only lone pair of e-
      • so molecular geometry is bent
    O O O N O O O
  • 31. Carbonate Ion (CO 3 2- )
  • 32. Nitrate Ion (NO 3 - )
  • 33. Nitrite Ion (NO 2 - )
  • 34. Examples of 4 EP
    • 4 BP + 0 NBP = 4 EP
      • Both EP geom and molecular geom
      • tetrahedral
    • 3 BP + 1 NBP = 4 EP
      • 4 EP so EP geom is tetrahedral
      • Molecular geom is TRIGONAL PYRAMIDAL
      • No atom at top location
    • 2 BP + 2 NBP = 4 EP
      • 4 EP so EP geom is tetrahedral
      • Molecular geom is BENT
      • no atoms at two locations
  • 35. 4 BP + 0 NBP = 4 EP TETRAHEDRAL Cl Cl Cl Cl S
  • 36. 3 BP + 1 NBP = 4 EP TRIGONAL PYRAMIDAL N H H H ●● lone pair of e- NBP H H H N 107
  • 37. 2 BP + 2 NBP = 4 EP BENT O H H ●● lone pair of e- NBP H H O ●● lone pair of e- NBP 104.5 H
  • 38. Hydronium Ion (H 3 O + )
  • 39. Ammonia Molecule (NH 3 )
  • 40.  
  • 41. Summary of 4 EP
  • 42. Exceptions to Octet Rule
    • Reduced Octet
      • H only forms one bond- only one pair of e-
      • Be tends to only form two bonds
        • only two pair of e-
      • B tends to only form three bonds
        • only three pair of e-
    • Expanded Octet
      • Empty d-orbitals can be used to accommodate extra e-
      • Elements in the third row and lower can expand
      • Up to 6 pairs of e- are possible
  • 43. Lewis Structures in Which the Central Atom Exceeds an Octet
  • 44. 5 EP
    • Trigonal bipyramidal
    • Orbital hybridization
    • Requires 5 hybrid orbitals
    • So, 5 atomic orbitals required
    • sp 3 d
  • 45. Trigonal planar shape
  • 46. 5 EP = trigonal pyramidal
    • molecular geometry
    • 5 BP + 0 NBP = 5 EP 4 BP + 1 NBP = 5 EP
  • 47. 5 EP = trigonal pyramidal
    • molecular geometry
    • 3 BP + 2 NBP = 5 EP 2 BP + 3 NBP = 5 EP
  • 48. 6 EP
    • Octahedral
    • Orbital hybridization
    • Requires 5 hybrid orbitals
    • So, 5 atomic orbitals required
    • sp 3 d
  • 49. 6 EP = octahedral 6 BP +0 NBP 6 EP
  • 50. 6 EP = octahedral 5 BP +1 NBP 6 EP
  • 51. 6 EP = octahedral 4 BP +2 NBP 6 EP
  • 52.  
  • 53.  
  • 54. Summary: Molecular Geometry of Expanded Octets
  • 55. Summary of EP Geometry
    • 2 EP 3 EP 4 EP 5 EP 6 EP
  • 56. Summary of EP Geometry
  • 57. Predict the geometry, angles and orbital hybridization Predict the geometry, angles and orbital hybridization
  • 58.  
  • 59.  
  • 60. Covalent Bond: A Model
    • Chemical bonds can be viewed as forces that cause a group of atoms to act as a unit
    • They result from the tendency of a system to seek its lowest possible energy
    • Bonds occur when collections of atoms are more stable (lower in energy) than the separate atoms
  • 61. Example: Methane
    • 1652 kJ of energy are required to break a mole of methane into separate C and H atoms
    • OR 1652 kJ of energy are released when one mole of methane is formed from one mole of C atoms and four moles of H atoms
    • Methane is therefore a stable molecule relative to its stable atoms
    • Since there are four H atoms arranged around the central C, it is natural to envision four individual attractions between C and H (bonds)
    • An average bond energy associated with each bond is found by dividing the total energy by four (1652/4 = 413 kJ)
    • The positive Bond Energy value indicates the energy required to break the bond between C and H atoms
  • 62. Bonding Model
    • Models originate from our observations of the properties of nature
    • Atoms can form stable groups by sharing electrons, shared electrons give a lower energy state because simultaneously attracted to two nuclei
    • Remember: Models are human inventions that allow us to explain and predict. A model is a useful way of thinking; they include simplifications and assumptions. A model does not equal reality.
  • 63.  
  • 64. Bond Energy and Enthalpy
  • 65.