Book Cover Organic Chemistry 4e Carey
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Chapter 1: Chemical Bonding



Summary | Atoms, Electrons and Orbitals | Bond types: Ionic, Covalent, Polar Covalent | Lewis Structures and Formal Charge | The Shapes of Molecules | Hybridization | Self Assessment | Quiz |


Lewis Structures and Formal Charge

Chapter 1: Chemical Bonding

Lewis Diagrams give the connectivity between atoms in a molecule, they tell you nothing about the arrangement of the atoms of a molecule within 3D space (that is left to VSEPR or Hybridisation).

You should be able to draw a Lewis Structures for both organic and inorganic molecules as the latter are often used as reagents in organic syntheses. Examples.   If you are not yet comfortable with determining a Lewis Structures, take a look at an outline of how to determine a Lewis Structures.
 

In reality, electrons in a molecule are being delocalized within that molecule.  Consequently, more than one Lewis Structures can be used to describe molecules which possess multiple bonds.  Collectively these Lewis Structures are then known as resonance structures and their sum is often represented as the resonance hybrid (note a resonance hybrid is not a Lewis Structures!).
 


You should be able to draw all reasonable resonance structures for a given organic molecule. Examples.   If you are not yet comfortable with examining the resonance in a molecule you should review the concept of resonance.

 

Going Deeper into Lewis Diagrams:

Given below is an outline of how to detemine the "best" Lewis structure for NO3-.

1.  Determine the total number of valence electrons in a molecule
2.  Draw a skeleton for the molecule which connects all atoms using only single bonds.  In simple molecules, the atom with the most available sites for bondng is usually placed central.  The number of bonding sites is determined by considering the number of valence electrons and the ability of an atom to expand its octet.  As you become better, you will be able to recognize that certain groups of atoms prefer to bond together in a certain way.
3.  Of the 24 valence electrons in NO3-, 6 were required to make the skeleton. Consider the remaining 18 electrons and place them so as to fill the ocets of as many atoms as possible (start with the most electronegative atoms first then proceed to the more electropositive atoms).
4.  Are the octets of all the atoms filled?   If not then fill the remaining octets by making multiple bonds (make a lone pair of electrons, located on a more electronegative atom, into a bonding pair of electrons that is shared with the atom that is electron deficient).
5.  Check that you have the lowest FORMAL CHARGES possible for all the atoms, without violating the octet rule;       (valence e-) - (1/2 bonding e-) - (lone electrons).
This is very important, no Lewis diagram is complete without formal charges.  Lewis diagrams are to be used to examine mechanisms so that knowing which parts of a molecule are electron deficient and which are electron rich are vital.  It is ideal to have a formal charge of 0 for as many of the atoms as possible.  If a formal charge of 1- is located next to a formal charge of 1+, the formal charges can be minimized by having a lone pair of electrons, located on the atom with the 1- charge become a bonding pair of electrons that is shared with the atom that has the 1+ formal charge (this can be visualized in the same way as the formation of multiple bonds were above).
CAUTION, octets are able to be expanded to minimize formal charge but only atoms with n=3 or greater are able to do this!   For instance in our example, N cannot expand its octet so keeps a formal charge of 1+ and both singly bonded oxygens a formal charge of 1-.  If our molecule were SO3 , however, it would be possible to minimize all formal charges by having the sulfur expand it's octet.
6.  You may find that the best Lewis diagram (the one with the lowest formal charges and all octets satisfied) is given in a number of different ways.  For NO3-, three different diagrams are given below.  From left to right they start with the most complete Lewis diagram to the most simplified.
Why so many different ways?  Depends on the need of the chemist. For instance, complete structures are more useful for the novice organic chemist learning to appreciate the mechanism of a reaction while simplified versions may be preferred by inorganic chemists.
Going Deeper into Resonance Structures:

In order to determine all possible resonance structures one must learn to "push" curved arrows.  These arrows indicate how to move the appropriate electrons within one resonance structure to obtain another.

It can never be stressed enough - curved arrows represent the movement of electrons not atoms, and the electrons are most often moving towards an electron deficient center!

Rules to remember when considering resonance:

  • Atoms never move.
  • You can only move electrons in multiple bonds or lone pairs.
  • The overall charge of the system must remain the same.
  • The bonding framework of a molecule must remain intact.
Ranking the importance of resonance structures:
  • Those with an electron count around an atom of at least an octet, barring exceptions, and no charge separation are favored over
  • Those that exhibit a charge separation, which are favored over
  • Those that exhibit a charge separation against that predicted by electronegativity.
Here is an organic example. Note how the possible resonance structures are derived from the starting point by pushing curved arrows.
Resonance of Propanone

Now let's look at an inorganic example.  Again the possible resonance structures are derived from the starting point by pushing curved arrows.
 

 



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