Last time we talked about how to use electronegativity to find the electron densities in a molecule – and when to ignore formal charge. However I didn’t mention one of the factors that can sometimes complicate the analysis of electron densities: double bonds (π bonds). But let’s start with the simple stuff. In many cases, an analysis of electronegativities can give you a very accurate idea of where electron densities are in molecules with π bonds. This can tell you where the “dipoles” are.
This model is very effective for finding the reactive sites in these molecules. Here are two examples.
Experiments support our prediction that the positively charged acid (i.e. “H+”) will react at the “negative” atom bearing the electrons (oxygen in this case) while negatively charged hydroxide (HO-) will react at the atom bearing the partially positive charge (C in this case).
So far so good.
The problem with our precious little model is that it does not agree well with experiment in some cases. Things start to fall apart when we look at molecules like this.
Part of the experience of being a scientist is watching your very simple, appealing hypotheses/models get ripped to shreds by experimental evidence (or lack thereof). We’d predict from this model that the more “negative” oxygen would be more reactive than the neutral oxygen, and likewise for the “positive” and neutral carbons. However, if we test this model by using labelled compounds, we find that our simple model is inconsistent with experiment.
From the experiments above, the molecules behave like they are a 1:1 mixture of these two compounds.
Let’s introduce a new concept: Resonance.
When we can draw two (or more) forms of the same molecule that differ only in the placement of their electrons, these are called “resonance forms” (or “resonance structures”). We can also say these forms are “in resonance”. Sometimes they are also called resonance “isomers” although this will get a hand-slap in many circles, as they are not technically isomers.
It is tempting (and very wrong!) to think that these two forms are in “equilibrium” between each other. Avoid this common mistake!
Why is this wrong? Because were this true, the bond lengths of the C=O double bond and the C–O single bond in the acetate ion would be different. (1.21 vs. 1.36 Angstrom). However experimental evidence (from X-ray crystal structures) shows us that each bond length is exactly the same (1.26 Angstrom). This observation has been made for countless other resonance isomers as well, such as benzene, carbonate, nitro groups, and many more.
What this means is that the true structure of the acetate ion (and the allyl carbocation) is a “hybrid” of two resonance forms. We use a special arrow to depict the two resonance forms (the double-headed arrow).
What this means is that the formal charge of –1 in the acetate ion is distributed evenly between the two oxygens. And the formal charge of +1 in the allyl carbocation is distributed equally between the two terminal carbons. [Note 1]. As we learned last time, formal charge can be a poor guide to electron densities, so here we see another example where it’s inaccurate.
Again: electron densities are what really matters for understanding the reactivity of a molecule. Lewis structures and formal charges are imperfect guides to fully describing electron densities in molecules, but they’re so useful in many other contexts that we have to learn to live with their flaws.
Next post we’ll talk about another accounting system we will use to describe the movement of electrons: the curved arrow formalism.
[Note 1: These are simple examples. Charge distribution will not always be evenly divided, and resonance forms will not always be of equal energy – more on that in future posts].