3.59

[Laura Riccardelli]

In problem 3.39, how do we know that some of the Lewis structures should have double bonds such as in c and some should only have single bonds such as in a and b? Does this have to do with the Formal Charge?

[Gabriela Carrillo 1B]

For part c) H2C(NH2)COOH, the first step I did was to calculate the total number of valence electrons (30e-s). I understand after looking at the solutions manual that each atom has a full octet (or for hydrogen a full shell), but I don't understand the thought process behind connecting and centering certain atoms

[Katherine Jordak 1H]

Your goal when drawing a Lewis structure is usually to get each element to have 8 valence electrons. You know whether to use a single or double bond usually based on how close each atom is to having a full outer shell. If you draw the Lewis structure for a molecule and then realize that both the central atom and an atom attached to it both only have 7 valence electrons, you would add another bond. If the molecule has resonance, you'll have to play around with the number of bonds a bit more and try to find the structure with the lowest formal charge.

[Raymond Zhang 3H]

Adding to that, for example Oxygen, being number 6, needs 2 additional electrons to have a full shell of 8. Thus, when paired with another Oxygen, they form a double bond to both have 8 electrons.