Acids and Bases - Definitions

A lot of the behaviors and properties of acids and bases can be described by equilibrium.  Before we get into that, we need to establish some definitions.  In Gen Chem 1, we tried to recognize acids by looking for H⁺ and bases by looking for OH^-.  That's a good start, and is the basis for the Arrhenius definitions of acids and bases:

Acid = H⁺-donor

Base = OH^--donor

These very brief definitions work, but we quickly run into a problem.  Aqueous ammonia is a base.  Aqueous ammonia {NH₃(aq)} does not have a OH^- to donate.  Arg.  Fortunately, there's a qualifier there, this is aqueous ammonia, and if there's water around, we can use it:

NH₃(aq) + H₂O(l)  <=>  NH₄⁺(aq) + OH^-(aq)

So NH₃(aq) can be forced to fit the simple definition above by using water.  In fact, a more proper and complete version of the Arrhenius definitions of acids and bases is:

Acid = a substance that, when dissolved in water, increases the concentration of H⁺(aq)

Base = a substance that, when dissolved in water, increases the concentration of OH^-(aq)

These definitions are a little better, but they still seem a bit restrictive.  To make definitions that are a little more general, let's look at that ammonia equation again.  In that equilibrium, ammonia is "increasing the concentration of OH^-(aq)" by removing H⁺ from water.  If we're trying to understand the function of acids and bases, it might be nice to follow one consistent thing around, so maybe a "better" definition of acids and bases could be:

Acid = H⁺-donor

Base = H⁺-acceptor

These are the Bronsted-Lowry definitions of acids and bases, and they are the definitions we will use most often in Gen Chem 2.  It's important to note that the Bronsted-Lowry definitions and the Arrhenius definitions are (and must be) consistent with each other.  It wouldn't be very helpful if 1 definition called something a base while the other definition called that exact same substance an acid.

Since we're dealing with definitions here, what should we call "H⁺(aq)"?  "H plus" works, but we're also going to run across a few other descriptions.  If we think about the subatomic particles present in "H⁺(aq)", there's 1 proton (the atomic number of hydrogen, all hydrogens have to have 1 and only 1 proton), the "+1" charge means that the single electron of a hydrogen atom has been removed, and in the most common isotope of hydrogen there are zero neutrons.  This means that, in terms of subatomic particles, "H⁺(aq)" is just a proton, and that is what it is very often called in discussions of acids and bases.  Acids are "proton donors" and bases are "proton acceptors.
If we think in the other direction, a proton is a pretty concentrated little lump of positive charge.  If this little lump of positive charge is floating around in a polar solvent like water, it's probably going to attract (or be attracted to) the negative end of the water dipole.  We can write a chemical equation to describe this:

H⁺(aq) + H₂O(l)  <=>  H₃O⁺(aq)

"H₃O⁺(aq)" is called a "hydronium ion".  {For practice, draw Lewis structures of all the species in that equation.}
So when we're talking about acids and bases, we're likely to encounter a number of different descriptions of the same thing.  To keep them straight, remember that:

"proton" = H⁺(aq) = H₃O⁺(aq)

Now that we have a few definitions in place, what else can we do to set the table for acids and bases...