A general chemical definition for acids and bases was proposed by Svante
Arrhe-Table 13.1
General Properties of Acids and Bases
nius in 1887. Arrhenius defined an acid as a substance that dissociates in water to give hydrogen ions (H) and defined a base as a substance that dissociates in water to give hydroxide ions (OH). A hydrogen ion is simply a hydrogen atom minus its electron.
Because a hydrogen atom consists of a sin-gle proton and a sinsin-gle electron, removing an electron leaves just the proton. Therefore, a hydrogen ion is equivalent to a proton, and both can be symbolized as H. If the gen-eral formula of an acid is represented as HA and a base as BOH, Arrhenius’ definitions for an acid and base can be represented by the following general reactions:
Acid HA H2Op H(aq) A(aq) Base BOH H2Op OH(aq) B(aq) When an acid dissociates to produce hydro-gen ions in water, the hydrohydro-gen ions do not remain as individual ions but are attracted to the polar water molecules represented by the following reaction:
H H2Op H3O
The H3Oion is called the hydronium ion.
To be technically correct, Arrhenius’ defi-nition for an acid should state an acid is a substance that produces hydronium ions in solution. Several water molecules may actu-ally be associated with a single hydrogen ion to produce ions such as H5O2or H7O3in acids. The important thing to remember is that hydrogen ions do not exist as individual ions in water, but become hydrated. While you will often see the symbol Hin reac-tions involving acids, the H should be interpreted as a hydronium ion rather than a hydrogen ion. Throughout this chapter hydrogen ions and hydronium ions should be considered synonymous.
The Arrhenius definition implies that acids contain hydrogen ions and bases
con-tain the hydroxide ion. This is illustrated in the chemical formulas for some common acids and bases displayed in Table 13.2.
Table 13.2 illustrates the presence of hydrogen in acids. It is also apparent that bases contain hydroxide ions, but the weak base ammonia seems to be an exception.
Ammonia illustrates one of the short-comings of the Arrhenius definition of acids and bases; specifically, bases do not have to contain the hydroxide ion to produce hydroxide in aqueous solution. When ammonia dissolves in water, the reaction is represented by:
NH3(g) H2O(l)m NH4(aq) OH(aq) Table 13.2
Some Common Acids and Bases
This reaction shows that the hydroxide ions come from ammonia pulling a hydrogen away from water resulting in the formation of OH. Therefore, a compound does not have to contain hydroxide to be a base. In addition to this limitation, the Arrhenius definition limits acids and bases to aqueous solutions.
Recognizing these limitations, the Dan-ish chemist Johannes Brønsted (1879–1947) and the English chemist Thomas Lowry (1874–1936) independently proposed a new definition for acids and bases in 1923.
Brønsted and Lowry noted that in certain reactions substances acted like acids even though hydrogen ions were not present.
Brønsted and Lowry viewed acids and bases not as isolated substances but as substances that interacted together as a pair. According to the definition, an acid is a proton donor and a base is a proton acceptor. An anal-ogy helps to explain the idea behind the Brønsted-Lowry theory. Acids and bases can be compared to throwing a football. A quarterback throws a football to a receiver.
In this example, the quarterback represents the acid, the football the proton, and the receiver the base. Applying the Brønsted-Lowry definition to nitric acid, it is seen that HNO3 donates a proton to water, Water in this case acts as a base.
When ammonia dissolves in water, water acts as an acid when it donates a proton to the base ammonia:
Just as a quarterback and receiver act as a pair, Brønsted-Lowry acid and bases act as pairs (Figure 13.1). Acids and bases are
defined in terms of reactions and the inter-action of two substances, rather than as indi-vidual substances. In the two examples it is seen that water, which is not commonly con-sidered an acid or base, is an acid when it reacts with ammonia and a base when it reacts with nitric acid.
To complete our discussion of the Brønsted-Lowry theory, we see that when a proton is transferred, two new species are formed. In the nitric acid reaction, water is transformed into the hydronium ion and nitric acid into the nitrate ion, NO3. The hydronium and nitrate ions formed are themselves an acid and base, respectfully.
Again, using our quarterback and receiver example, once the receiver has the football, he or she can act as an acid and pass it on to someone else. Likewise, the quarterback can act as a base by receiving the football.
The HNO3–NO3 and H2O–H3O are referred to as conjugate acid-base pairs.
Notice that the only difference between the substances making up a conjugate acid-base pair is the presence of a proton. In the ammonia reaction, the conjugate acid base pairs are NH3(base)–NH4(acid) and H2O (acid)–OH(base).
In the same year that Brønsted and Lowry proposed their definition for acids and bases, the American G. N. Lewis pro-posed an alternative definition based on the
Acid Base
Football Analogy (Rae Déjur)
valence electron structure of substances.
Brønsted-Lowry bases must contain at least one unshared electron pair to accept a pro-ton. We can see this if we look at the Lewis dot structure of several bases:
Lewis defined a base as an electron pair donor and an acid as an electron pair accep-tor. Lewis’ electron pair donor was the same as Brønsted-Lowry’s proton acceptor, and therefore, was an equivalent way of defining a base. Lewis’ acids were defined as a sub-stance with an empty valence shell that could accommodate a pair of electrons. This definition broadened the Brønsted-Lowry definition of an acid. The three definitions of acids and bases are summarized in Table 13.3.
Each of the three definitions expands our concept of acids and bases. Arrhenius’
basic definition is adequate for understand-ing many of the properties of acids and