pages for my notes: a short primer on acids/bases
What *is* an acid or a base? This isn’t as easy a question as one might think. We have many different ways of using them, and defining them.
There are three categories of acids/bases that we use, but are espcially important when it comes to organic chemistry. So, I’ll go from more general to more detailed.
An important note that took some of my classc mates a while to fully realize is that H+ refers to a proton. In my gen chem classes, we had always written P as proton so the sudden introduction of H+ as a proton with no explanation made the connection hard to make.
1. Arrhenius acids and bases: You learn about these in general chemistry, but it’s rarely mentioned again later on. This is because it can be considered a subset of a different type of acid/base (Bronsted-Lowry). However, it’s still useful as a general definition, especially as a frame of reference for the Arrhenius equation.
Arrehenius acids are molecules that dissolve in WATER to give an ion and H+. So, it’s a proton donor, specifically in water.
Arrhenius bases are molecules that upon dissolving in WATER, give an OH- ion.
2. Bronsted-Lowry acids and bases: You also learn about these in genchem, I think, and if not, it’s a useful concept in ochem to look at reagents. This is an expansion on Arrhenius acids/bases.
Bronsted-Lowry acids are molecules that give an H+ in solution. Note that the solvent doesn’t have to be water! Any molecule that dissolves in any solution and gives an H+ is a Bronsted-Lowry acid. The solvent can be acetone, NH3, what have you. Note that this also means that any Arrheneus acid is also a Bronsted-Lowry acid (but not the opposite)
Bronsted-Lowry bases, on the other hand, are any molecules that can accept a proton (H+) in solution. Note that ANY molecule/ion can be a BL base, as long as it can accept protons while it is dissolved in solution. So again, an Arrhenius base is also a Bronsted-Lowry base, but not the opposite.
3. Lewis Acids and Bases: This is a very different definition; however, it’s ultimately the same as a Bronsted-Lowry acid/base but from a different perspecive. It’s the most useful in organic chemistry and the way most acids/bases are considered.
Lewis acids are atoms/molecules that tend to accept electrons. Note that it’s not just taking away the electrons, the way an ionic bond behaves. When we say “accepts electrons,” typically this indicates it’ll take the electrons and form a bond with whatever is giving those electrons. THIS IS A VERY GENERAL DEFINITION FOR A REASON--IT ENCOMPASSES MANY MANY THINGS. For example, anything with a positive charge is a Lewis acid because they want to accept electrons. Electrophiles are Lewis acids, etc. This sounds different from a BL acid, right? The trick is to understand that when the H in a molecule leaves, the atom it was connected to accepts the electrons from that bond. Understand that the H leaves as a proton, leaving its electrons for the molecule to take. It’s kind of confusing. In general, I use a mix of the two to look at organic chemistry questions--BronstedLowry to look at reagent/solvent acidity, and Lewis to look at reaction mechanisms.
Lewis bases, as you can guess, are atoms/molecules that donate electrons. Anything with a negative charge is a Lewis base. Nucleophiles are bases, etc. Donating electrons doesn’t necessarily mean it completely loses them, more that it donates them to form a bond with something that wants electrons (ie., Lewis acids). Again, reconciling it with Bronsted-Lowry bases is the same: the base accepts the proton by donating electrons in a bond with it.
Note that the positive/negative aren’t hard either! HSO4-, despite that negative sign, is also an acid. Some molecules are both acids and bases!