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How To Use Electronegativity To Determine Electron Density (and why NOT to trust formal charge)


Last time we talked about how electrons are the “currency” of chemistry and every reaction is a transaction of electrons between atoms. That means that if we really want to understand a reaction, we have to understand where the electrons are (and aren’t).

There’s two factors to employ when doing this. The first is electronegativity. That’s what today’s post is about: using electronegativity to determine electron densities. (The second is resonance BTW – more on that in future posts)

I’m assuming you know how to draw Lewis structures and understand the concepts of electronegativity and formal charge, as well as being able to interpret line drawings, but that’s it.  If you’re not at that level, back up and read up on those concepts.

OK. Let’s start with the first question: how do we tell where the electrons are in a molecule?

The first skill is being able to draw proper Lewis structures for a molecule. To succeed in organic chemistry you absolutely need to be able to do this in your sleep. The Lewis structure should account for all the electrons around an atom, including the often-hidden lone pairs of electrons if applicable. Let’s have a look.

The second skill lies in being able to apply electronegativity to determine partial charges in bonds.

See, our drawings of chemical structures can sometimes get in the way of what is really going on with the electrons. 

If we just paid attention to the drawings themselves, the lines we draw between atoms – “covalent bonds” – are electron pairs that are shared equally between the two.

However, the difference between the idealized sharing of electrons in a covalent bond versus the reality of different electron densities is, to use an analogy, not unlike the difference between a utopian socialist worker-state, and Soviet Russia.

Remember electronegativity – a ranking of an atom’s “greed” for electrons, in other words? In a bond, the more electronegative element will have a greater share of the electrons, and a partial negative charge to reflect this greater electron density.  The less electronegative element will have a partial positive charge to reflect the lack of electron density. 

Let’s look at a few simple molecules and analyze the dipoles in their bonds by comparing relative electronegativities. Remember the mnemonic for electronegativity? “Phone call bro…. CSI is on, please hold…. bye” .



Why is it important to go through all of this? Because in chemical reactions, electrons will flow from areas of high electron density to areas of low electron density. Knowing where the partial charges are is an important first step in determining where the molecule will react. Covalent bonds with large dipoles (i.e. large differences in electronegativity) are worth looking paying attention to: frequently, this will be where the “action” is.

“Hold on”, you might say. “I thought electron densities were reflected by their charges, like in H3O+, BF4-, and NH2- !” No no no no no. This is one of the first real curveballs that gets thrown at you in organic chemistry, and one that continually gives students fits.

Formal charge is NOT the same as electron density. 

“Formal charge” gives us a bit of a dilemma. When a molecule bears a charge (either positive or negative),  for bookkeeping purposes, we have to denote one atom as “bearing” that charge. However it’s important to realize that this “bookkeeping” is NOT the same as electron density, which is  the real sources of reactivity. Sure, there are lots of examples of molecules where the charge actually *does* represent the electron density. But then again, there are a lot of counter examples too – like BH4(-), NH4(+), H3O(+) and others.

Make sure you understand this because it might be hard to get your head around the first time.

The O in H3O+ might have a “formal” charge of +1, but it is actually the most electron-rich (i.e. negatively charged) atom on the molecule!

So although the molecule H3O(+) has a charge of +1, that positive charge  is actually not on the oxygen itself, but is “spread out” – dispersed – around the molecule, but particularly around the hydrogens.( Click here to see a potential energy map of H3O(+) ). However, for “book-keeping” we assign a charge of +1 on the oxygen, because of the underlying assumption in bond diagrams that electrons are shared equally between atoms.

Bottom line: if H3O(+) is reacting with an electron rich species, those electrons will go to the hydrogens (since they are electron poor), NOT the oxygen!!

If you use formal charge to determine electron densities you will get screwed over on a regular basis. 

I apologize for writing such a long post but let’s finish by applying this concept toward some potential attractive and repulsive interactions between molecules. These are not meant to depict actual reactions, although as we will later see, reactions will begin with an attractive interaction between two molecules. Conversely, where we line up two like charges (positive-positive and negative-negative) these are sites where repulsion will occur (and reactions between these atoms will not occur).
Bottom line #2: If you learn electronegativities and how to assign partial charges, you will be always be able to figure out where the electron rich and electron poor areas of a molecule are, and we can apply this toward reactions. 
Things will get a bit more complicated when resonance is possible.  More on that in the next post.

Note how on the last part of the 3rd slide I threw in some new words – “electrophilic” and “nucleophilic” to denote “electron-poor” and “electron-rich” respectively. Much more on these later.

Here’s an excellent online resource on the topic of electrostatic potential in organic chemistry from Reed College.

Next Post: Introduction to Resonance

 

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