14.1 Brønsted-lowry acids and bases

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By the end of this section, you will be able to:
• Identify acids, bases, and conjugate acid-base pairs according to the Brønsted-Lowry definition
• Write equations for acid and base ionization reactions
• Use the ion-product constant for water to calculate hydronium and hydroxide ion concentrations
• Describe the acid-base behavior of amphiprotic substances

Acids and bases have been known for a long time. When Robert Boyle characterized them in 1680, he noted that acids dissolve many substances, change the color of certain natural dyes (for example, they change litmus from blue to red), and lose these characteristic properties after coming into contact with alkalis (bases). In the eighteenth century, it was recognized that acids have a sour taste, react with limestone to liberate a gaseous substance (now known to be CO 2 ), and interact with alkalis to form neutral substances. In 1815, Humphry Davy contributed greatly to the development of the modern acid-base concept by demonstrating that hydrogen is the essential constituent of acids. Around that same time, Joseph Louis Gay-Lussac concluded that acids are substances that can neutralize bases and that these two classes of substances can be defined only in terms of each other. The significance of hydrogen was reemphasized in 1884 when Carl Axel Arrhenius defined an acid as a compound that dissolves in water to yield hydrogen cations (now recognized to be hydronium ions) and a base as a compound that dissolves in water to yield hydroxide anions.

In an earlier chapter on chemical reactions, we defined acids and bases as Arrhenius did: We identified an acid as a compound that dissolves in water to yield hydronium ions (H 3 O + ) and a base as a compound that dissolves in water to yield hydroxide ions (OH ). This definition is not wrong; it is simply limited.

Later, we extended the definition of an acid or a base using the more general definition proposed in 1923 by the Danish chemist Johannes Brønsted and the English chemist Thomas Lowry. Their definition centers on the proton, H + . A proton is what remains when a normal hydrogen atom, ${}_{1}^{1}\text{H},$ loses an electron. A compound that donates a proton to another compound is called a Brønsted-Lowry acid    , and a compound that accepts a proton is called a Brønsted-Lowry base    . An acid-base reaction is the transfer of a proton from a proton donor (acid) to a proton acceptor (base). In a subsequent chapter of this text we will introduce the most general model of acid-base behavior introduced by the American chemist G. N. Lewis.

Acids may be compounds such as HCl or H 2 SO 4 , organic acids like acetic acid (CH 3 COOH) or ascorbic acid (vitamin C), or H 2 O. Anions (such as ${\text{HSO}}_{4}{}^{\text{−}},$ ${\text{H}}_{2}{\text{PO}}_{4}{}^{\text{−}},$ HS , and ${\text{HCO}}_{3}{}^{\text{−}}\right)$ and cations (such as H 3 O + , ${\text{NH}}_{4}{}^{\text{+}},$ and ${\left[\text{Al}{\left({\text{H}}_{2}\text{O}\right)}_{6}\right]}^{3+}\right)$ may also act as acids. Bases fall into the same three categories. Bases may be neutral molecules (such as H 2 O, NH 3 , and CH 3 NH 2 ), anions (such as OH , HS , ${\text{HCO}}_{3}{}^{\text{−}},$ ${\text{CO}}_{3}{}^{\text{2−}},$ F , and ${\text{PO}}_{4}{}^{\text{3−}}\right),$ or cations (such as ${\left[{\text{Al(H}}_{2}{\text{O)}}_{5}\text{OH}\right]}^{2+}\right).$ The most familiar bases are ionic compounds such as NaOH and Ca(OH) 2 , which contain the hydroxide ion, OH . The hydroxide ion in these compounds accepts a proton from acids to form water:

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