Midterm Question: Lewis Structure for HOCO [ENDORSED]

[torialmquist1F]

What does this lewis structure look like? I know that it is a radical but do not know how to draw it.

[Srbui Azarapetian 2C]

HOCO is the same order as how the elements are aligned in the lewis structure.
First you count the number of valence electrons: 1 for Hydrogen, 2 x 6 for Oxygen, and 4 for Carbon. This equals 15 electrons total to be shared. Then begin by drawing a single bond between all the atoms, resulting in H-O-C-O as a start. This leaves you with 9 electrons. These must go to the oxygens, aiming for a formal charge of 0. This results in two lone pairs for the first oxygen in between the hydrogen and carbon. For the second oxygen at the end bonded to the carbon, a double bond is formed so that the oxygen can have a formal charge of 0. If there is confusion as to if the carbon should have a single unpaired electron or the oxygen, it goes to the oxygen because it is more electronegative. The radical thus ends up on the carbon, which has a formal charge of 0 with a single lone electron, one single bond with the oxygen on the left and a double bond with the oxygen on the right.

[Nora Sharp 1C]

The H atom should have a single bond with the oxygen, which in turn should have a single bond with the carbon. The carbon atom should have a double bond with the remaining oxygen. In this case, HOCO is a radical because all the valence electrons of the atoms add up to form an odd number. The remaining electron goes to the carbon atom because the carbon atom is the only one that has not achieved noble gas configuration. Because it is not favored that the electron join the hydrogens and oxygens, it should be with the carbon.

Here is what it looks like:
http://images.slideplayer.com/18/620070 ... lide_6.jpg

[juchung7]

Why can't Carbon be double bonded to both oxygens and the radical electron go on one of the oxygens? Oxygen is more electronegative than Carbon, and if Carbon had 2 double bonds then its orbitals would be filled and formal charge in most places would be the lowest possible.

[Nora Sharp 1C]

That figuration would not have the lowest formal charge possible. The oxygen on the left now bonds with carbon twice, so its formal charge is now 6 - 5 = 1 e-. The oxygen on the right now also has the formal charge 6 - 5 = 1 e-, (assuming it has the lone pair and the one radical electron) while the carbon has 0. The issue is that the most electronegative element should have the lowest formal charge, and all should be as close to zero as possible. Both of the oyxgens, the most electronegative elements in the molecule, are the only atoms that have a positive formal charge. Also, there are formations that are more favorable for this molecule, with all formal charges = 0. This is partially why you don't often see oxygen with triple bonds in a molecule like this, because that leaves the very electronegative oxygen with a positive formal charge.

Also note that an atom's electronegativity does not necessarily mean that an atom is favored to bond with more electrons.

Looking at the accepted answer for HOCO, if we check the formal charges, we can see that they all come out as 0, even carbon.
I hope this helped!

[dstemp]

Why wouldn't it work to have carbon as the central atom with a single bond to an oxygen and hydrogen, and a double bond to the other oxygen (resonance)? Put 5e- on the single bonded oxygen (making it radical) and 4 e- on the double bonded oxygen. This would make a formal charge of zero on each atom, so why can't this be a solution?

[Nora Sharp 1C]

While it is true that this formation would result in formal charges of zero for every atom, the accepted answer is still the one listed on the midterm solutions. There are two tip offs that let you know this is the case:

First, note the way the molecule is written. The molecule you bring up actually exists, although I don't know if it appears in radical form, so it usually has an extra electron. Written, it actually has the formula HCO2-. This name is a clue that carbon acts as the central atom, because the oxygen atoms are grouped together with the subscript 2. In the case of HOCO, the fact that the oxygens are apart is very telling.

Second, oxygen is not favored to be the atom that has an unfilled shell. The oxygen in the molecule you suggest would be a very electronegative atom with a nearly full electron shell. This would be a very high energy configuration for oxygen. Carbon is a better alternative because it is not overly picky with the configurations it needs to be in, meaning it is very versatile when it comes to bonding. It doesn't have a very high electronegativity and has four slots open to bond with electrons, and given that hydrogen and oxygen wouldn't want that radical electron due to their tendency to have full shells, carbon is the best choice for an unfilled shell.