You have no doubt heard of acids and can probably name a few just from reading food labels: Citric acid. Acetic acid. At the same time, you know that at least some acids can be harmful if you so much as handle them, so different acids clearly have different properties, including different strengths.
Bases are everywhere in the world, too, although they seem to get less publicity for some reason. Like acids, bases can be damaging to biological and other materials. You've encountered a strong base in the form of household laundry bleach (NaClO, or sodium hypochlorite).
Acids and bases are complementary in almost every way, and one can even be used to "neutralize" the other, as with taking oral antacid tablets to combat stomach acid. Part of this is in the nomenclature; when acids actually behave like acids, they become bases, and ditto for the behavior of bases. Understanding conjugate acids and bases is essential to mastering chemical reactions.
History of Acid-Base Chemistry
As far back as the mid-1600s, Robert Boyle, who seemed to be involved in just about every chemistry experiment in those days, figured out that certain solutions had properties such as the ability to damage immersed substances or change their colors, and that these effects could be prevented or negated by the addition of alkali compounds, which today are known to be basic.
In 1923, Johannes Brønsted and Thomas Lowry formally defined acids and bases in terms of the transfer of hydrogen ions (H+).
The conjugate base of an acid is the compound remaining after a hydrogen ion is donated by the acid, and the conjugate acid of an base is the compound remaining after a hydrogen ion is accepted by the base.
A Brønsted-Lowry acid is therefore simply a molecule that can donate a hydrogen ion (which is a positively charged atom) to another molecule; the remnant of that acid is called its conjugate base. For example, when hydrochloric acid donates a proton, the chloride ion left behind is the conjugate base:
HCl → H++ Cl−
Sometimes, an acid will be positively charged before donating its hydrogen ion, rather than neutral as in the instance of HCl. This can be observed with the ammonium ion donating a proton to become the conjugate base ammonia:
NH4+ → H++ NH3
H2PO4− : Acid or Base?
So far, you have seen examples of compounds with formulas that make it obvious whether the molecule functions as an acid or as a base (or, for that matter, as neither). If you see an ion with no hydrogen atoms included, such as Cl−, you know that it cannot be an acid, since it has no protons, but that it could be a base, since it is an anion with a charge of −1 and "eager" to take on a proton.
But what about compounds with multiple hydrogen atoms available for exchange? In the right environment, a compound that functions as a base in the presence of a strong enough acid can also act as an acid in the presence of a strong enough base. (Think of bases as "hydrogen-ion-pullers." Such a compound is called amphoteric or amphiprotic.
A classic example is the dihydrogen phosphate ion H2PO4− . In the presence of the strong acid HBr, this molecule readily accepts the hydrogen ion from the acid to become phosphoric acid (H3PO4). Yet in the presence of basic hydroxide (OH−) ions, dihydrogen phosphate instead donates a proton to become monohydrogen phosphate (HPO42−).
The conjugate base of H2PO4−
is therefore HPO42−, and the conjugate acid of
H2PO4− is H3PO4.
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