Acids and Bases

I

Introduction

Acids and Bases, two classes of chemical compounds that display generally opposite characteristics. Acids taste sour, turn litmus (a pink dye derived from lichens) red, and often react with some metals to produce hydrogen gas. Bases taste bitter, turn litmus blue, and feel slippery. When aqueous (water) solutions of an acid and a base are combined, a neutralization reaction occurs. This reaction is characteristically very rapid and generally produces water and a salt. For example, sulfuric acid and sodium hydroxide, NaOH, yield water and sodium sulfate:

H2SO4 + 2NaOH⇄2H2O + Na2SO4

II

Early Theories

Modern understanding of acids and bases began with the discovery in 1834 by the English physicist Michael Faraday that acids, bases, and salts are electrolytes. That is, when they are dissolved in water, they produce a solution that contains charged particles, or ions, and can conduct an electric current Ionization. In 1884 the Swedish chemist Svante Arrhenius (and later Wilhelm Ostwald, a German chemist) proposed that an acid be defined as a hydrogen-containing compound that, when dissolved in water, produces a concentration of hydrogen ions, or protons, greater than that of pure water. Similarly, Arrhenius proposed that a base be defined as a substance that, when dissolved in water, produces an excess of hydroxyl ions, OH^-. The neutralization reaction then becomes:

H+ + OH-⇄H2O

A number of criticisms of the Arrhenius-Ostwald theory have been made. First, acids are restricted to hydrogen-containing species and bases to hydroxyl-containing species. Second, the theory applies to aqueous solutions exclusively, whereas many acid-base reactions are known to take place in the absence of water.

III

Brønsted-Lowry Theory

A more satisfactory theory was proposed in 1923 by the Danish chemist Johannes Brønsted and independently by Thomas Lowry, a British chemist. Their theory states that an acid is a proton (hydrogen ion, H⁺) donor and a base a proton acceptor. Although the acid must still contain hydrogen, the Brønsted-Lowry theory does not require an aqueous medium. For example, liquid ammonia, which acts as a base in aqueous solution, can act as an acid in the absence of water by transferring a proton to a base and forming the amide anion (negative ion) NH₂^-:

NH3 + base⇄NH2- + base + H+

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The Brønsted-Lowry definition of acids and bases also explains why a strong acid displaces a weak acid from its compounds (and likewise for strong and weak bases). Here acid-base reactions are viewed as a competition for protons. In terms of a general chemical equation, the reaction of Acid (1) with Base (2)

Acid (1) + Base (2)⇄Acid (2) + Base (1)

results in the transfer of a proton from Acid (1) to Base (2). In losing the proton, Acid (1) becomes its conjugate base, Base (1). In gaining a proton, Base (2) becomes its conjugate acid, Acid (2). The equilibrium represented by the equation above may be displaced either to the left or to the right, and the actual reaction will take place in the direction that produces the weaker acid-base pair. For example, hydrogen chloride (HCl) is a strong acid in water because it readily transfers a proton to water to form a hydronium ion:

HCl + H2O⇄H3O+ + Cl-

The equilibrium lies mostly to the right because the conjugate base of HCl, Cl^-, is a weak base, and H₃O⁺, the conjugate acid of H₂O, is a weak acid.

In contrast, hydrogen fluoride, HF, is a weak acid in water because it does not readily transfer a proton to water:

HF + H2O⇄H3O+ + F-

This equilibrium lies mostly to the left because H₂O is a weaker base than F^-, and because HF is a weaker acid (in water) than H₃O⁺. The Brønsted-Lowry theory also explains why water can be amphoteric, that is, why it can serve as either an acid or a base. Water serves as a base in the presence of an acid that is stronger than water (such as HCl), in other words, an acid that has a greater tendency to dissociate than does water:

HCl + H2O⇄H3O+ + Cl-

Water can also serve as an acid in the presence of a base that is stronger than water (such as ammonia):

NH3 + H2O⇄NH4+ + OH-

IV

Measuring Acid or Base Strength

The strength of an acid can be measured by the extent to which an acid transfers a proton to water to produce the hydronium ion, H₃O⁺. Conversely, the strength of a base is indicated by the extent to which the base removes a proton from water. A convenient acid-base scale is calculated from the amount of H₃O⁺ that is formed in water solutions of acids or of OH^- formed in water solutions of bases. The former is known as the pH scale and the latter as the pOH scale (pH). The value for pH is equal to the negative logarithm of the hydronium ion concentration—and for pOH, of the hydroxyl ion concentration—in an aqueous solution:

pH = -log [H3O+]

pOH = -log [OH-]

Pure water has a pH of 7.0. When an acid is added, the hydronium ion concentration [H₃O⁺] becomes larger than that in pure water, and the pH becomes less than 7.0, depending on the strength of the acid. The pOH of pure water is also 7.0, and in the presence of a base, the pOH drops to values lower than 7.0.

The American chemist Gilbert N. Lewis has offered another theory of acids and bases that has the further advantage of not requiring the acid to contain hydrogen. This theory states that acids are electron-pair acceptors and bases are electron-pair donors. This theory also has the advantages that it works when solvents other than water are involved and it does not require the formation of a salt or of acid-base conjugate pairs. Thus, ammonia is viewed as a base because it can donate an electron pair to the acid boron trifluoride, for example

H3N: + BF3⇄H3N-BF3

to form an acid-base association pair.

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