VSEPR Valence Shell Electron Pair Repulsions
Covalent Bond: A Model <ul><li>Chemical bonds can be viewed as forces that cause a group of atoms to act as a unit </li></...
Example:  Methane <ul><li>1652 kJ of energy are required to break a mole of methane into separate C and H atoms </li></ul>...
Bonding Model <ul><li>Models originate from our observations of the properties of nature </li></ul><ul><li>Atoms can form ...
Covalent Bonds <ul><li>Electron pair(s)  </li></ul><ul><ul><li>shared between two atoms </li></ul></ul><ul><ul><li>attract...
Multiple covalent bonds around the same atom determine the shape <ul><li>Negative e- pairs with same charge repel each oth...
Single bonds <ul><li>Sigma bond  ө </li></ul><ul><li>Overlap of orbitals allow electron pair to be shared between the two ...
Double and Triple Bonds <ul><li>Pi bonds  π </li></ul><ul><li>Since the space between the nuclei is occupied, e- pair is s...
 
Lewis Structures <ul><li>Drawn to show the bonds between the atoms in the structure </li></ul><ul><li>Only shows whether s...
Lewis Structure <ul><li>Represents the arrangement of valence electrons among atoms in the molecule </li></ul><ul><li>Rule...
Rules for Drawing Lewis Structures <ul><li>Count the number of valence electrons </li></ul><ul><li>Draw the skeleton struc...
Determine # Valence e-  from column #
 
 
 
 
Electron Clouds repel each other, thus structure around an atom is determined principally by minimizing repulsions
2 electron pairs (2 EP) around central atom <ul><li>Two clouds pushed as far apart as possible </li></ul><ul><ul><li>Great...
3 electron pairs (3 EP)  around central atom <ul><li>Three clouds pushed as far apart as possible </li></ul><ul><ul><li>Gr...
4 electron pairs (4 EP)  around central atom <ul><li>Four clouds pushed as far apart as possible </li></ul><ul><ul><li>Gre...
Orbital Hybridization #1 <ul><li>Atomic orbitals such as s and p are not well suited for overlapping and allowing two atom...
Orbital Hybridization #2 <ul><li>Hybrid orbitals (cross of atomic orbitals) </li></ul><ul><li>Remember: The pink flower hy...
Overlap of two s-orbitals NOT A GOOD LOCATION-  far from one nucleus Note:  shared in this overlap the e- pair would spend...
Overlap of two p-orbitals One atom & its p-orbital The other atom & its p-orbital represents the nucleus BAD location far ...
Hybrid Orbitals yield more favorable shape for overlap <ul><li>Atomic orbitals are not shaped to maximize attractions nor ...
Angles  and  Shape <ul><li>Atomic orbitals are not shaped to maximize attractions nor minimize repulsions </li></ul><ul><l...
Orbital Hybridization #3 <ul><li>Each e-pair requires a hybrid orbital </li></ul><ul><li>If two hybrid orbitals required t...
Electron-Pair Geometry vs Molecular Geometry <ul><li>Electron-pair geometry </li></ul><ul><ul><li>Where are the electron p...
Examples of 3 EP <ul><li>3 BP + 0 NBP = 3 EP </li></ul><ul><ul><li>3 EP = EP geom is trigonal planar </li></ul></ul><ul><u...
Carbonate Ion (CO 3 2- )
Nitrate Ion (NO 3 - )
Nitrite Ion (NO 2 - )
Examples of 4 EP <ul><li>4 BP + 0 NBP = 4 EP </li></ul><ul><ul><li>Both EP geom and molecular geom </li></ul></ul><ul><ul>...
4 BP + 0 NBP = 4 EP TETRAHEDRAL Cl Cl Cl Cl S
3 BP + 1 NBP = 4 EP TRIGONAL PYRAMIDAL N H H H ●● lone pair of e- NBP H H H N 107
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
Hydronium Ion (H 3 O + )
Ammonia Molecule (NH 3 )
 
Summary of 4 EP
Exceptions to Octet Rule <ul><li>Reduced Octet  </li></ul><ul><ul><li>H  only forms one bond- only one pair of e- </li></u...
Lewis Structures in Which the Central Atom Exceeds an Octet
5 EP <ul><li>Trigonal bipyramidal </li></ul><ul><li>Orbital hybridization </li></ul><ul><li>Requires 5 hybrid orbitals </l...
Trigonal planar shape
5 EP = trigonal pyramidal <ul><li>molecular geometry </li></ul><ul><li>5 BP + 0 NBP = 5 EP  4 BP + 1 NBP = 5 EP </li></ul>
5 EP = trigonal pyramidal <ul><li>molecular geometry </li></ul><ul><li>3 BP + 2 NBP = 5 EP  2 BP + 3 NBP = 5 EP </li></ul>
6 EP <ul><li>Octahedral </li></ul><ul><li>Orbital hybridization </li></ul><ul><li>Requires 5 hybrid orbitals </li></ul><ul...
6 EP = octahedral 6 BP +0 NBP 6 EP
6 EP = octahedral 5 BP +1 NBP 6 EP
6 EP = octahedral 4 BP +2 NBP 6 EP
 
 
Summary: Molecular Geometry of Expanded Octets
Summary of EP Geometry <ul><li>2 EP 3 EP 4 EP  5 EP  6 EP </li></ul>
Summary of EP Geometry
Predict the geometry, angles  and orbital hybridization Predict the geometry, angles  and orbital hybridization
 
 
Covalent Bond: A Model <ul><li>Chemical bonds can be viewed as forces that cause a group of atoms to act as a unit </li></...
Example:  Methane <ul><li>1652 kJ of energy are required to break a mole of methane into separate C and H atoms </li></ul>...
Bonding Model <ul><li>Models originate from our observations of the properties of nature </li></ul><ul><li>Atoms can form ...
 
Bond Energy and Enthalpy
 
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Vsepr

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Vsepr

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