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A Primer On Organic Reactions

By James Ashenhurst

Common Mistakes: Formal Charges Can Mislead

Last updated: November 6th, 2019 |

Sometimes The “Formal Charge” Does Not Accurately Represent The Electron Density Around An Atom! 

Formal charges have their plusses and minuses. Har har.

In many instances, the formal charge on an atom is an “honest” expression of its electron density. We’re all familiar with the ions Cl(-), HO(-), CH3O(-),Br(-), Li(+) and so on.

The formal charge assigned to these atoms truly reflects that these molecules bear additional positive or negative charge.

However! then there are the outlier cases. And these cause problems. From someone who preaches “opposite charges attract, like charges repel“, it’s important to know when to pay attention to formal charge, and when to ignore it.

“Formal” charge is called “formal” because it’s ultimately an accounting issue. If a molecule bears a charge, it makes things a lot easier (for nomenclature reasons) if we adopt some kind of system where a charge was unambiguously assigned to one atom.

Just like  the rules in baseball  sometimes assign a “Win” or “Loss” to a pitcher who didn’t contribute much to the team’s overall performance, formal charge doesn’t take into account the true electron densities of a molecule, which are based on a combination of electronegativity and resonance.

When trying to understand a new reaction, apply electronegativity to understand electron densities , not formal charge.

Where Formal Charge Can Lead One Astray

For instance in the two examples below-left, the curved arrows, as drawn, would be showing the formation of an oxygen-oxygen bond. This doesn’t make sense given the weakness of the oxygen-oxygen bond (about 35 kcal/mol).

When you apply electronegativities, however, you get a much better picture of the true electron density of a molecule. And this can help you figure out how a reaction might proceed.

Several Examples Of Species Where The Formal Charge Does Not Accurately Represent Electron-Density (And Therefore, Reactivity)

Here are some other common species where formal charge can be a misleading indicator of electron density.

  • AlH4 (–) and BH4 (–) are hydride donors (sources of H-). The nucleophilic atom is actually hydrogen, not Al or B.
  • NH4 (+) is a weak acid (source of H+). Bases react with NH4 at H, not N.
  • The species on the right is called an “iminium ion”. Nucleophiles react with the iminium ion at carbon, not at nitrogen.

Keep this in mind, and you’ll have a much easier time of properly understanding how reactions work.

Next Post: Seven Factors that Stabilize Negative Charge in Organic Chemistry

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6 thoughts on “Common Mistakes: Formal Charges Can Mislead

  1. Kind of an esoteric application, but some organic chemistry researchers have built entire research programs on the idea of misleading formal charges. See “Lewis Base Activation of Lewis Acids” here:

    http://www.scs.illinois.edu/denmark/research/index.html

    The idea is that for certain acid-base pairs, coordination increases the partial charge on the coordinating atoms, in contrast to the intuitive “acid + base = neutralization” idea. Frustrated Lewis acid/base pairs are kind of the same idea, taken to an extreme:

    http://en.wikipedia.org/wiki/Frustrated_Lewis_pair

    Great post! Using electronegativity to think about formal charge is GREAT advice.

    1. You’re right, he has based an entire research program out of that! Scott Denmark gives great talks. I love it when you can come away from a lecture not only having a good idea of what the researcher does, but also having learned something: he really dug into the physical organic chemistry aspect of it. I remember vividly him asking the audience (via a story about Albert Eschenmoser) “where are the electrons on the borohydride ion”, and when you think about it, the boron actually has a partial positive charge. Then he extended it to silicon. Great stuff.

  2. Hi,

    The question I have is very basic (the problem is, over time, I keep forgetting or getting confused with the basic facts that help understand/figure out why reactions happen the way they do). My question is: If we have an electronegative atom, rich in electron density, perhaps with a negative charge on it, why would it attack another atom that is slightly deficient in electrons (like in the examples above). Don’t electronegative atoms like to “keep” their negative charges? I know that electrons flow from higher density to lower density, and it makes sense. But I can never shake off the feeling that the electronegative atom is “compromising”.

    Thank you!

    1. Let’s take the example of an electron-rich, electronegative species with a negative charge like F- . Yes, fluoride has a full octet, and yes, the fluorine nucleus is highly electronegative. However, having a negative charge (i.e. high charge density) is somewhat destabilizing.
      If the fluoride ion were to combine with H+ [for example] yes fluoride would go from “owning” a pair of electrons to “sharing” a pair with hydrogen, but this is not really a “compromise”. This is like a happy marriage where the partners are better off together than they were apart! The new H-F bond is worth about 136 kcal/mol and the resulting species is neutral.
      What’s happening here is that in the process of creating the bond, that pair of electrons has gone from being localized around the fluorine nucleus to being in a larger volume (the H-F sigma bond) which significantly lowers the overall energy.

      In practice there are no point charges in nature so a real-life example would be NaF plus HCl [so we’d also obtain NaCl in the bargain.] Still a highly exothermic process.

      It’s only [neutral] neon and its relatives who can obtain a full octet of electrons and not have to “compromise” through bonding.
      Thanks for the question, Mallika!

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