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This article incorporates, in modified form, material from Illustrated Guide to Home Chemistry Experiments: All Lab, No Lecture.

 

Acids and bases are two of the most important classes of compounds in chemistry. Surprisingly, it took chemists many years to agree on just what made an acid an acid and a base a base.

From experience, we understand that acids and bases have certain characteristics. For example, acid solutions taste sour and base solutions are bitter. (Obviously, you should never taste any laboratory chemical, but the sour taste of edible acids such as lemon juice and vinegar are familiar to most people, as are the bitter tastes of edible bases such as baking soda and quinine water). Acid solutions generally have an astringent feel on the skin, while base solutions feel slimy. Acid solutions turn blue litmus paper red, while base solutions turn red litmus paper blue. Acids and bases combine to form salts. And so on.

But these characteristics are insufficient to define acids and bases. Around the turn of the 19th century, French and European chemists believed that all acids must contain oxygen. (They were wrong, but echoes of that error persist; for example, the German word for oxygen is Sauerstoff, which translates literally as “sour material.”) Davy and other English chemists were much closer to the mark. They believed that all acids must contain hydrogen. And, although that statement is not absolutely true for all substances considered to be acids by modern definitions, it is true that all acids known to be acids at that time do contain hydrogen.

The first good working definition of acids and bases was proposed in the late 19th century by the Swedish chemist Svante August Arrhenius, who defined an acid as a substance that when dissolved in water increases the concentration of the hydronium (H3O+) ion and a base as a substance that when dissolved in water increases the concentration of the hydroxide (OH-) ion. Although that definition limits acids and bases to compounds that are water soluble, it was a pretty good definition for the time and remains a useful definition even today.

But not all compounds that behave chemically as acids can dissociate to produce hydrogen ions, and not all compounds that behave chemically as bases can dissociate to produce hydroxide ions. In 1923, the Danish chemist Johannes Nicolaus Brønsted and the English chemist Thomas Martin Lowry redefined an acid as a proton (hydrogen nucleus) donor, and a base as a proton acceptor. Under the Brønsted-Lowry definition of acids and bases, an acid and its corresponding base are referred to as a conjugate acid-base pair. The conjugate acid is the member of the pair that donates a proton, and the conjugate base the member of the pair that accepts a proton. For example, hydrochoric acid dissociates in water to form chloride ions and hydronium ions:

HCI + H2O ⇌ H3O+ + CI-

In the forward reaction, the acid reactant (HCl) and the base reactant (H2O) form the acid product (H3O+) and the base product (CI-). In the reverse reaction, the acid reactant (H3O+) and the base reactant (CI-) form the acid product (HCl) and the base product (H2O). If the conjugate acid (HCl on the left side of the equation and H3O+ on the right side) is strong, the conjugate base (H2 on the left side of the equation and Cl- on the right side) is weak, and vice versa. At equilibrium, the weaker acid is favored.

For aqueous solutions, the Arrhenius definition and the Brønsted-Lowry definition are essentially the same. The value of the Brønsted-Lowry definition is that it extends the concept of acids and bases to compounds that are not soluble in water.

The same year that Brønsted and Lowry defined acids and bases as proton donors and proton acceptors, respectively, the American chemist Gilbert N. Lewis extended the definition of acids and bases to include compounds that behaved chemically as acids and bases without donating or accepting a proton. Under the Lewis definition of acids and bases, a Lewis acid (also called an electrophile) is a substance that accepts an electron pair, and a Lewis base (also called a nucleophile) is a substance that donates an electron pair. (For example, the compound ferric chloride, FeCI3, behaves as an acid, although it has no proton to donate, and so is classified as a Lewis acid.) Lewis acids and bases are particularly important to organic chemists, who make use of them in many syntheses.

In this section, we’ll examine the properties of acids and bases.

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