Formal charges have their plusses and minuses. Har har.
One one hand, they’re an indispensable accounting tool. If a molecule bears a charge, it would drive us nuts (for nomenclature reasons) if we didn’t adopt some kind of system where a charge was unambiguously assigned to one atom.
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. It 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.
For instance in the bottom two examples, the curved arrows, as drawn, would be showing the formation of an oxygen-oxygen bond. This doesn’t make sense.
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.
Here are some other common molecules where formal charge can be a misleading indicator of electron density.
Keep this in mind, and you’ll have a much easier time of properly understanding how reactions work.