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One of the earliest and most important skills needed in organic chemistry is the ability to represent molecules in 2D (on paper). To do this, you have to master the shared electron pair (covalent) Lewis bonding model. Before starting to practice, review the table below to see what we already know about the number of valence electrons, bonds, and lone pairs associated with common neutral atoms in organic molecules.
Below is a systematic method to draw the "best" Lewis structure using NO3- as the example:
1. Determine the total number of valence electrons in a molecule (notice the extra valence electron due to the negative charge)
2. Draw a skeleton for the molecule which connects all atoms using only single bonds. In simple molecules, the atom that forms the most bonds (see table above) is usually placed centrally. The number of bonding sites is detemined by considering the number of valence electrons and the ability of an atom to expand it's octet (only n=3 or below atoms can have more than 8 valence electrons, like sulfur and phosphorus). As you progress in your understanding of organic chemistry, you will be able to recognize that certain groups of atoms prefer to bond together in a certain ways called functional groups.
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? Not on nitrogen! 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. Also, remember (valence e-) - (1/2 bonding e-) - (lone pair electrons) = Formal Charge.
IMPORTANT : no Lewis diagram is complete without formal charges. Lewis diagrams are also drawn to examine reaction mechanisms so knowing which parts of a molecule are electron deficient (+) and which are electron rich (-) is vital. It is best to have a formal charge of 0 for as many of the atoms in a structure as possible.
If a formal charge of 1- is located next to a formal charge of 1+, the formal charges can usually 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 visualised in the same way as the formation of multiple bonds were above).
CAUTION : octets can be expanded to minimize formal charges but only for atoms in the third row of the periodic table (where n=3 or greater). 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 its 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 other chemists.
While this systematic process will help you draw most any Lewis structure, you will eventually abandon this rigid step-wise process for a more intuitive style (see diagram below).