Defining acids and bases
It turns out that there are 3 common definitions of acids and bases (or alkalis). They all compliment each other to provide a complete picture of what acids and bases are.
First off, please note that protons and hydrogen ions are the same thing and thus will be used interchangeably. One may think strictly of protons when looking at the general behavior of acids and bases, and may think of hydrogen ions concerning solutions and reactions.
- Acids produce protons in a solution
- Bases produce (negative) hydroxide ions in a solution
- Acids donate protons
- Bases accept protons
- Acids accept electron pairs
- Bases donate electron pairs
Harmonizing the theories
Each definition expands upon the preceding one, but they all compliment each other. Let’s imagine an acid mixing with another chemical. It would donate protons to the reactant (Brønsted-Lowry), thus producing protons in the solution (Arrhenius). The acid would become negatively charged, effectively accepting an electron pair (Lewis). Looking at an alkaline reaction, the base would accept protons, producing a negative charge in hydroxide ions in the solution. Since its net charge is positive, it effectively donated an electron pair. For more on these theories, click this link.
The truth is, there are 3 ways of measuring acidity/alkalinity: pH, pOH, and pK.
This is what most people refer to when describing an acid or a base: but what does it mean?
Soren Sorensen introduced the concept of pH in 1909. He defined pH as a mathematical function of how many hydrogen ions was in a substance, with pH standing for “potential of hydrogen ions.” So pH just has to do with the amount of protons, right? Well, sort of.
He realized that pH was not just a measure of how many hydrogen ions were in a solution, but a measure of how active they are. This activity can be seen as an effective concentration of the ions. An ion’s activity is defined by much more than just its concentration, however, so pH is a rough yet highly effective approximation of a solution’s ionic activity. The pH functions are shown here:
The “p” is an operator referring to the negative log scale of what’s in front of it. The first equation is the original pH equation, while the second was updated by Sorensen to take into account the activity of the ions. The brackets in front of the log function refer to a solution’s molarity, or concentration. Concentrations have very small numbers (along the lines of 10−4). As the concentration increases, the -log value decreases, showing increasing acidity.
This is the typical pH scale, but it is possible for values to go slightly above 14 and slightly below 0 (yes, below 0!). For this to be possible, one would need a concentration of hydrogen ions higher than 1, i.e., a solution with over 100% ions! We can resolve this problem by remembering that pH depends on the activity of the hydrogen ions, not the amount. Its effective concentration may be much higher than the actual concentration.
For more on Sorensen and the invention of the pH scale, follow this link.
As one may suspect, pOH measures the concentration of hydroxide ions, or OH−. It isn’t any more complicated than that, but what is interesting is that in any given solution, its pH and pOH always add up to 14. This makes sense when thinking of liquid solutions, since the vast majority of them have water. If there is a hydrogen atom (H) separated from water (H2O), there must be a hydroxide ion (OH) left behind. If we take the concentration of each ion and add them, it should even out to 14.
Let’s use pure water as an example. Its pH and pOH are both 7. Add these together and we get 14. One could think of pOH as a lack of pH and vice versa, with a (typical) maximum of 14.
There are actually two pK’s: pKa and pKb. The K’s are dissociation constants for acids and bases. They refer to how easily a compound dissociates, or splits up into its constituent molecules. Since weak acids and bases can have very low constants, the “p” operator is used to make the numbers easier to work with.
Amphoteric means that a substance behaves as both an acid or base. We see one like this all the time: water! To show its amphoteric nature, look at this reaction:
Two water molecules react to form hydronium (which is often used in pH calculations instead of H+), and hydroxide. The hydronium acts like an acid and the hydroxide acts like a base.