Formal charges in organic chemistry is, perhaps, one of the most fundamental bookkeeping devices which is often misunderstood or neglected by students.
Knowing formal charges can help us understand the reactivity patterns in reactions, find reactive centers, and make sense out of electron flow in the mechanisms. For instance, negatively charged species tend to be the sources of electron density in reactions, while the positively charged species—accept those electrons.
If you’ve started your journey in organic chemistry, you’re well aware that molecules are often depicted with neutral atoms. However, not all atoms remain neutral.
When we think about organic molecules, the first step is to visualize their bonding patterns. For instance:
In these classic examples, the charge on each of these elements is zero.
Now, things get a bit more intriguing when atoms have a negative charge. In this scenario, one bond morphs into an electron pair.
Notably, negatively charged halide and hydride anions lack any bonds. Thus, they exist solely as free ions.
Positive charges add an extra layer of complexity:
Here’s an important fact: All heteroatoms (like nitrogen, oxygen, and halogens) will always have a full octet. Even though it’s theoretically possible for heteroatoms to be open-shelled with just 6 electrons, such scenarios are too unstable in organic molecules. The only exception? Carbon with a positive charge. These are termed as ‘carbocations’. Despite their existence, they’re quite unstable.
Drawing bond line structures in chemistry can sometimes allow us to implicitly display hydrogens and electron pairs. But charges? They’re never implicit. A common oversight many students make is overlooking these charges. Thus, always double-check your drawings!
To calculate an atom’s charge, you can use the following formulas:
This latter method, with its clear emphasis on subtracting lines and dots from valence electrons, can help you determine the charge on any atom in your molecule effortlessly.
A word of caution: While organic molecules can technically bear charges greater than +1 or -1, encountering such is a rarity. If you find an atom with a charge like +3 or -4, double-check your work. Having multiple charges spread across different atoms is feasible, but a single large charge on an individual atom is uncommon.
Now, I’ve mentioned earlier that there’s a difference between the formal and the actual charge. Formal charge is a bookkeeping tool that is important to help us keep track of the electron flow in the reaction.
The actual charge, however, is the actual electron density that is present on the atom. For instance, let’s take a look at borohydride anion:
The electronegativity of boron is 2.0 while electronegativity of hydrogen is 2.2. So, the hydrogen is more electronegative (not by much but still) and will polarize the bond. This means that hydrogen actually “pulls” the electron density towards itself. Thus:
While formal charges are merely a “formality,” they are very important for the reactions mechanisms understanding. Thus you need to make sure you master the skill of quickly finding the formal charge.
You also notice that I’ve indicated my real electron densities with the delta-minus (𝛿-) symbol. That denotes that I only have a partial negative charge on each of the hydrogens. How much of that partial charge we have on them? Well, we could calculate it using fancy quantum chemical calculations, but that’s utterly unnecessary for the purposes of a typical organic chemistry course. What’s more important, is to realize that the boron is not actually negatively charged in this molecule. So, when we write a reaction with a borohydride anion, we won’t be showing electrons coming from boron!
Notice how in the example above my arrow starts at the H-B bond and not at the boron atom! That’s because it’s not the boron that is a source of electron density here! You remember from earlier in this post, it’s the hydrogens are the atoms with the 𝛿-. So now it makes sense that the arrow doesn’t start at boron.