Bases are considered the chemical opposite of acids because of their ability to neutralize acids. In 1887 the Swedish physicist and chemist Svante Arrhenius defined a base as the chemical substance that produces hydroxide ions (OH − ) and cations. A typical base, according to the Arrhenius definition, is sodium hydroxide (NaOH). The neutralization of an acid with a base to yield salt and water may be represented as
HCl ( aq ) + KOH ( aq ) ⇆ H 2 O ( l ) + KCl ( aq ) (1)
A major problem with Arrhenius's definition of bases is that several chemical compounds, such as NaHCO 3 , Na 2 CO 3 , Na 3 PO 4 , which produce basic solutions when dissolved in water, do not contain hydroxide ions. The Brønsted-Lowry theory, which was proposed independently by Danish chemist Johannes Brønsted and English chemist Thomas Lowry in 1923, states that a base accepts hydrogen ions and an acid donates hydrogen ions. This theory not only includes all bases containing hydroxide ions, but also covers any chemical species that are able to accept hydrogen ions in aqueous solution . For example, when sodium carbonate is dissolved in solution, the carbonate ion accepts a hydrogen ion from water to form the bicarbonate ion and hydroxide ion.
The Brønsted-Lowry theory includes water as a reactant and considers its acidity or basicity. In reaction (2) a new acid and base are formed, which are called the conjugate acid and conjugate base, respectively.
The strength of a base is determined by the extent of its ionization in aqueous solution. Strong bases, such as NaOH, are 100 percent ionized in aqueous solution and weak bases, such as ammonia, are only partially ionized in aqueous solution.
The partial ionization is a dynamic equilibrium , as indicated by the double arrow in equation (3).
The strength of acids and bases also determines the strength of their conjugate bases and conjugate acids, respectively. Weak acids and bases have strong conjugate bases and acids. For example, when ammonium chloride is dissolved in water, it gives an acidic solution because ammonium ion is a strong conjugate acid of the weak base ammonia, but chloride ion is a weak conjugate base of the strong acid hydrochloric acid.
NH 4 + ( aq ) + H 2 O ( l ) → NH 3 ( aq ) + H 3 O + ( aq ) (4)
The carbonate ion in equation (2) yields a basic solution because it is the strong conjugate base of the weak acid HCO 3 − .
When NaHCO 3 is dissolved in water, it gives a basic solution, even though a hydrogen ion is available. Predicting this requires one to consider the strength of carbonic acid, H 2 CO 3 , which is a very weak acid.
H 2 CO 3 ( aq ) + H 2 O ( l ) ⇆ HCO 3 − ( aq ) + H 3 O + ( aq ) (5)
However, HCO 3 − will act as an acid if a strong base is added.
HCO 3 − ( aq ) + OH − ( aq ) → H 2 O ( l ) + CO 3 2− ( aq ) (6)
This ability to act as a base or an acid is called amphoterism. Any anions of polyprotic acids, such as HCO 3 − , H 2 PO 4 − , and HPO 4 2− , which contain replaceable hydrogen ions, are amphoteric. Some hydroxides, such as Al(OH) 3 and Zn(OH) 2 , are also amphoteric, reacting with a base or acid, as illustrated by the following equations:
Al(OH) 3 ( s ) + OH − ( aq ) → Al(OH) 4 − ( aq ) (7)
Al(OH) 3 ( s ) + 3 H 3 O + ( aq ) → Al 3+ ( aq ) + 6 H 2 O ( l ) (8)
Equations (7) and (8) can also be explained by American chemist Gilbert Lewis's acid-base theory. A Lewis acid is a substance that can accept a pair of electrons to form a new bond, and a Lewis base is a substance that can donate a pair of electrons to form a new bond.
All Arrhenius and Brønsted-Lowry bases are also Lewis bases. All metal cations are potential Lewis acids. Complexes of metal ions with water, ammonia, and hydroxide ion are examples of Lewis acid-base reactions. For example, [Al(H 2 O) 6 ] 3+ may be regarded as a combination of the Lewis acid, Al 3+ , with six electron pairs from six H 2 O molecules.
Buffer solutions contain a base and an acid that can react with an added acid or base, respectively, and they maintain a pH very close to the original value. Buffers usually consist of approximately equal quantities of a weak acid and its conjugate base, or a weak base and its conjugate acid. For example, one of the buffers used to keep the pH of the blood near 7.45 is the H 2 PO 4 − /HPO 4 2− acid/conjugate base system. Small amounts of an acid or base react with one of the components of the buffer mixture to produce the other component as follows:
H 2 PO 4 − ( aq ) + OH − ( aq ) → H 2 O ( l ) + HPO 4 2− ( aq ) (10)
HPO 4 2− ( aq ) + H 3 O + ( aq ) → H 2 O ( l ) + H 2 PO 4 − ( aq ) (11)
Melvin D. Joesten
Joesten, Melvin D., and Wood, James L. (1996). The World of Chemistry , 2nd edition. Fort Worth, TX: Saunders College.
Moore, John W.; Stanitski, Conrad L.; Wood, James L.; Kotz, John C.; and Joesten, Melvin D. (1998). The Chemical World , 2nd edition. Philadelphia: Saunders.
Carpi, Anthony. "Acids and Bases: An Introduction." Visionlearning. Available from http://www.visionlearning.com/library/science/chemistry-2/CHE2.2-acid_base.htm .
"CHEMystery: An Interactive Guide to Chemistry." Available from http://library.thinkquest.org/3659/acidbase/ .