What is an Acid or Base?
In Alien, a gigantic extraterrestrial monster called a Xenomorph is injured and bleeds onto the floor of a spaceship. Immediately, the blood begins to sizzle and eat through the spaceship’s metal floor. The Xenomorph’s blood is acidic and imbues it with the ability to eat through or corrode nearly any substance.
While not as fantastical as the SciFi creature’s blood we encounter acids and bases all the time, from lemon juice to baking soda and beyond.
Classifications
This quick discussion of 1970s horror films raise a foundational question. What is an acid or a base? Multiple classifications exist and it is worth getting to know each now. First up the Lewis acid and base.
Lewis
Lewis’ definition of acids and bases focuses on the behavior of electrons within each compound. Under this definition, Lewis acids are electron-pair acceptors while Lewis bases are electron-pair donors. Bases are said to attack acids and form new covalent bonds with their acid pair.
This description sounds strangely familiar and is identical to the idea of nucleophiles and electrophiles. Here nucleophiles typically carry negative charges and seek out positively charged electrophiles to form new bonds in a process called nucleophilic attack. These nucleophiles are the Lewis bases and the electrophiles are the Lewis acids.
Mentally, we can file Lewis’ definition under acids and bases but more importantly, we should think of this concept as describing coordinate covalent bonds, complex ions, and organic chemistry reactions. We already touched on this idea when going over different covalent bonds and complex ion formation.
Brønsted-Lowry
While Lewis’ definition focused on the action of electrons the Brønsted-Lowry definition of acids and bases focuses on the role of the proton (H+). Here acids are H+ donors while bases are H+ acceptors. With acids being the donors and bases the acceptors.
This swap between Lewis and Brønsted-Lowry often frustrates students since it makes keeping the two definitions straight a bit challenging.
However, both the Lewis and Brønsted-Lowry definitions are consistent with one another. In Lewis’s definition, we focused on electrons that are negatively charged while with Brønsted-Lowry we focus on protons that are positively charged. Therefore, when we will flip the charge we also have to flip the notation. So in either case acids get rid of their positive charges and bases get rid of their negative charges.
While not very kind to Gilbert Lewis I think of the L in Lewis as standing in for loser. It is a negative term so Lewis decided to describe electrons with their negative charge.
Arrhenius
Last and definitely least, the Arrhenius definition of acids and bases narrowly defines acids as molecules that dissociate into H+ ions and bases dissociate to form OH– ions. At this point, you might be thinking isn’t this the same thing as Brønsted-Lowry? Not quite.
For example, ammonia NH3 is a weak base but under the Arrhenius definition, it wouldn’t be classified as such. This makes sense if we place ammonia in water it can’t dissociate into an OH– ion since it doesn’t contain an OH in its molecular formula. Yet, ammonia is one of the main urinary buffers in our bodies that help maintain the acid-base balance of our urine. For this reason, the Arrhenius definition isn’t commonly used to define acids and bases.
Nonetheless, we need to know the definition for the MCAT and should be able to identify molecules that fall into each of the different classifications.
Some Species Can’t Decide
While some compounds are definitely bases or definitely acids some can’t quite decide and can act like either. When this occurs we call these amphoteric species meaning they are both acids and bases. Their behavior isn’t consistent and is highly dependent on the external environment.
Many polyprotic species or compounds with multiple acidic hydrogens or multiple basic spots fall into this category. Often the first hydrogen lost or gained is clearly acidic or clearly basic. However, the behavior of the second or third hydrogen completely depends on the external solution. Furthermore, as the external pH changes, these spots will switch between being acidic and basic.
While most polyprotic species are amphoteric it isn’t a strict requirement. H2O is an excellent example of a non-polyprotic amphoteric species. Although water technically has more than one hydrogen it isn’t actually able to lose both of them in an acid-base fashion.
Since water readily forms H3O+ and OH– depending on the external solution it is considered amphoteric. When forming H3O+ water acts as a base accepting an H+ from the solution when forming OH– it acts as an acid and gives up an H+ ion to the solution.
