Time is precious. Drawing the full structures of chemical formulae out is a gigantic pain. From a distance, maybe an extra 10-20 seconds here or there might not sound like a lot to you, but you might be surprised by what people will do to shave 5 seconds of boredom off their journey.
The point is, shortcuts are inevitable. This post is about the shortcuts you have to know about when looking at chemical structures.
There are 3 common types of shortcuts that organic chemists make when drawing structures. They will omit drawing the following:
• hydrogen atoms
• lone pairs
The key is learning to recognize the things that “you’re just supposed to know are there”.
Line drawings are the chemists’ method of choice for depicting structures. And it’s not hard to see why: they’re a snap to draw. Why? Because we omit drawing in all the pesky little hydrogens. The key thing in recognizing hidden hydrogens is that wherever you see less than 4 bonds to carbon, that carbon is going to have hydrogens attached such that there’s a full octet. Here are some examples.
Hidden lone pairs.
Lone pairs are often not drawn on structures either. Again, this is a time-saving measure, but also they tend to get in the way and clutter up the structure. Like hydrogens, you’re “just supposed to know that they’re there”. How do you know how many to draw? If you understand formal charge, this should be a snap. (If you don’t understand formal charge by this point, tsk tsk : this is a core skill).
The atom that causes the most confusion on this point is probably carbon when it is drawn with a negative charge: the negative charge implies the existence of a lone pair.
This also comes up in reactions. Using the curved arrow notation, you’ll often see the tail of an arrow coming from a negative charge on a given atom. You’re just supposed to know that the negative charge represents a pair of electrons on the atom.
Finally, if you see a positive charge on carbon, you’re supposed to know that not only are there zero lone pairs on the carbon, there are only 3 bonds.
The subject of counterions also causes a lot of confusion. Students coming out of gen chem where they have had to balance these massive redox equations often get confused when they notice their instructor (or the textbook) not balancing charges anymore. It’s a signal that “suddenly the charges aren’t that important”, or even worse, they aren’t there.
The charge IS there, it just hasn’t been drawn in. Charge is always conserved, and balanced: the presence of a charge implies the existence of an equal and opposite charge somewhere else. Always.
The reason why the balancing ions are often left off, I think, is not due to simple laziness but also for a desire to avoid confusion by focusing on specifics. It’s a bit of a Catch-22.
If the counterion is omitted, one risks confusing the students by implying that the charges are unbalanced. However, if one puts in a specific counterion such as Cl(-) , then one will have to eventually explain that Cl(-) is not crucial to the reaction and any number of other counterions would work just as well.
My compromise is to use X(-) or M(+) for this purpose. The exact counterion for a given reaction will depend on how the ion is made in the first place.
A related issue is not hidden counterions but hidden charges. Specifically, ionic compounds. Again, it’s easier to just omit drawing in charges rather than drawing them in. A classic example is NaOH. From general chemistry, it’s well understood that NaOH is an ionic compound and is better represented as Na(+)OH(-). However often times the charges will be omitted.
Hopefully now you’ll have a better idea of the things that you’re “just supposed to know are there”.