Acid-Base Definitions

Do you know what acids and bases are? Seriously? I know I used to think I did and I totally didn't. I mean, I knew some examples of acids and bases, but the chemistry behind them was completely unknown to me even after I took chemistry in high school and even when I started taking it in college. But finally, in one class I was introduced to three acid-base definitions. I knew there were more, but I had no idea how many. This post will introduce some of them. Really, I've only ever used two of them and those two, which happen to coincide with each other a lot, are the only two that will be important here, but I think the historical definitions are interesting and this is my blog or whatever and so I get to make a post including them.

Lavoisier definition
In case you didn't know who Lavoisier was, he was one of the founders of chemistry and you are not worthy. Through a series of experiments, he determined that oxygen combining with other elements could lead to "acidic" (the word comes from the Greek word for "sharp") properties. He extrapolated from this that oxygen was the element that contributed the acidity. This is where oxygen got its name, which means "acid-generating." There was a small problem with this: Lavoisier was wrong. Considering that he basically invented chemistry, I think he's allowed to be wrong every once in a while.

Liebig definition
You know who Liebig was too, right? Because he was another great chemist. Anyway, this was sort of the first real definition. A Liebig acid is a molecule that contains at least one hydrogen atom that can be replaced by a metal. This was only a definition for acids. Back then, bases were loosely defined as the opposites of acids. Known acids were generally liquids, so a base was whatever compound would react with an acid to neutralize it into a solid salt. Such reactions are known as acid-base reactions and we'll deal with them other posts pretty soon.

Arrhenius definition
This was one of the three definitions I originally learned and it was the main definition everyone used for a long time. But that was in the past. The distant past. Like before I was born, even. Probably before you were born too. Arrhenius acids are molecules that, in water, lose hydrogen atoms, generating hydrogen ions in the solution. Well, we now know that they're actually hydronium ions. Oh, hydronium is important. You should know what it looks like. Here, I'll paint one for you.
If you weren't stupid, you'd remember what a water molecule looks like. Man, I'm not posting water again here just for you. Go back and find the post where I did show it or something. Anyway, this is like water, but with another hydrogen. Oxygen normally only forms two bonds. I guess I never talked about formal charges or whatever, but the oxygen is positively charged now. Really. We'll talk about it later if you want. Hydronium is properly written as H₃O⁺. But you should be aware that a common shorthand is just to just write H⁺.

So those are Arrhenius acids, but there are also Arrhenius bases. They are molecules that, in water, generate hydroxide ions. You want a hydroxide ion? Here you go.
Note that when it has oxygen has one too many bonds, it is positively charged, but when it has one too few bonds, it is negatively charged. Unlike H⁺, hydroxide ions actually can and do exist in water. The shorthand for hydroxide is ⁻OH. That's how I would notate it if I were writing stuff down by hand, but superscripts and subscripts, although necessary for chemical notation, can be annoying to do on Blogger, so I'll probably stick to just typing "hydroxide" most of the time.

Now might be a good time to mention that in aqueous systems (systems with water as the solvent—I'm not going to elaborate on this for now), hydronium and hydroxide act as a sort of currency of acidity and bacisity. When a reaction takes place, the actual atoms from the acid/base aren't the ones that are participating. I'll illustrate this with a classic acid-base reaction...

HCl + NaOH → NaCl + H₂O.

So, hydrochloric acid and sodium hydroxide yield sodium chloride and water. Sodium chloride is the salt in this case. In other reactions, other salts would be formed. A salt and water are the products of Arrhenius acid-base reactions. But keep in mind that water is the solvent in which this whole thing is taking place. HCl is a gas and NaOH is a solid. When we do this reaction in a lab, we're likely to just mix samples of water that have the compounds dissolved in them. The bond between the chlorine and the hydrogen breaks and the hydrogen reacts with a water molecule to form hydronium while the chlorine atom becomes a chloride ion and sits there dissolved in the water. The bond between sodium and oxygen likewise breaks and we're left with hydroxide and a sodium ion, which also sits there dissolved in the water.

The hydronium ion produced by dissolving the acid in water can and will react with another water molecule. But the product of that reaction leaves the original hydronium ion turned into an ordinary water molecule and the water molecule attacked by hydronium as the new hydronium ion. This process occurs repeatedly, but without really changing anything because the number of hydronium ions remains constant. The same is true with the base. If hydroxide reacts with a water molecule, the hydroxide gains a proton and is now a water molecule while the water molecule it reacted with lost a proton and is now a hydroxide ion. The charges are rapidly transferred from one water molecule to the next, but they remain there. This tranfer would continue for a long time, but when we mix the two solutions, in addition to reacting with water, the hydroniums and hydroxides can react with each other. Whenever this reaction takes place, the two ions neutralize each other and we're left with only water (H₃O⁺ + ⁻OH → 2H₂O). This reaction also produces heat, so if you do it at home for some reason, keep that in mind so that you don't die or whatever. If you were wondering, the sodium and chloride ions stay dissolved. You have seen what happens when you add sodium chloride to water, right? Seriously, you'd better have. If not, go do it right now.

So that's the Arrhenius definition. There are several other definitions, some of them relevant to certain fields, but two definitions are by far the most popular and important, so I'll introduce them now.

Brønsted-Lowry definition
This is the definition that seems to be used the most in my textbook and introductory chemistry courses. It is more inclusive than the Arrhenius definition. A Brønsted-Lowry acid is a proton donor. A Brønsted-Lowry base is a proton acceptor. One big difference between this and the older Arrhenius definition is that water is not necessarily present as a solvent. It can be and often is, but it's not necessary. The most important difference to keep in mind might be that not every Brønsted-Lowry base will lose hydroxide. Every Brønsted-Lowry acid must have a hydrogen atom that can be lost in order to give off a proton, but the bases need not look anything like Arrhenius bases. What a Brønsted-Lowry base does need is the ability to form a bond to a proton. That means it needs free valence electrons that can be recruited to form the bond. Lone pairs of electrons work best, but electron pairs in π-bonds also have basic potential.

My next few posts will be dealing with Brønsted-Lowry acids and reactions with them, so this is the definition to pay the most attention to. Hopefully, you'll get a feel for what chemicals act as good acids and bases and how acid-base reactions work. For now, definitely remember that an acid is a proton donor and a base is a proton acceptor.

You might have deduced that once an acid has donated a proton, what's left is capable, under the right circumstances, of accepting a proton and recreating the original acid. Likewise, a base that accepts a proton now has a proton that could be donated to something else, recreating the original base. These are conjugate acid/bases. That is, whenever an acid-base reaction occurs, the acid becomes its conjugate base and the base becomes its conjugate acid. The reaction could potentially reverse. Here's one reaction...

H—A (acid) + :B (base) → :A⁻ (conjugate base) + H—B⁺ (conjugate acid)

Now we'll reverse the reaction...

:A⁻ (conjugate base) + H—B⁺ (conjugate acid) → H—A (acid) + :B (base)

The system can bounce back and forth between these two. So where does that leave us? The rule I learned in general chemistry and used all the time in my organic chemistry classes was that the equilibrium favors the side with the weaker acid. Unhelpful if you don't know which acid is weaker, but we have ways of figuring that out.

Under the Brønsted-Lowry definition, acidity and basicity are relative to what something is reacting with. Water, for example, is an acid when it's reacting with hydroxide (it donates a proton) and a base when reacting with hydronium (it accepts a proton). Also, you might have deduced that the conjugate acid of water is hydronium and the conjugate base of water is hydroxide. If you deduced that, good job. You get a gold star.

Lewis definition
This definition was formulated by Gilbert N. Lewis. You do know who he was, right? A Lewis acid is an electron pair acceptor. A Lewis base is an electron pair donor. This definition is more inclusive than the Brønsted-Lowry definition. A base still needs free valence electrons to donate, but the acid in the reaction no longer needs to provide hydrogen. So while all Lewis bases are also potential Brønsted-Lowry bases, not all Lewis acids are Brønsted-Lowry acids.

I'll say more about Lewis acids later. This material is in the last section of the acids and bases chapter in my textbook and it's time to move on to some specific Brønsted-Lowry acid-base reactions.