9.1: What are acids and bases?
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Harper College Chemistry Department
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 Svante 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 the most common isotope of hydrogen, $^1_1H$ , 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 HSO 4 – , H 2 PO 4 – , HS−, and HCO 3 – ) and cations (such as H 3 O + , NH 4 + , and [ Al ( H 2 O ) 6 ] 3+ ) 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 − , HCO 3 – , CO 3 2 - , F − , and PO 4 3 - ) , or cations (such as [Al(H 2 O) 5 OH] 2+ ) . 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:
We call the product that remains after an acid donates a proton the conjugate base of the acid. This species is a base because it can accept a proton (to re-form the acid):
We call the product that results when a base accepts a proton the base’s conjugate acid. This species is an acid because it can give up a proton (and thus re-form the base):
In these two sets of equations, the behaviors of acids as proton donors and bases as proton acceptors are represented in isolation. In reality, all acid-base reactions involve the transfer of protons between acids and bases. For example, consider the acid-base reaction that takes place when ammonia is dissolved in water. A water molecule (functioning as an acid) transfers a proton to an ammonia molecule (functioning as a base), yielding the conjugate base of water, OH − , and the conjugate acid of ammonia, NH 4 + NH4+ :
The reaction between a Brønsted-Lowry acid and water is called acid ionization. For example, when hydrogen fluoride dissolves in water and ionizes, protons are transferred from hydrogen fluoride molecules to water molecules, yielding hydronium ions and fluoride ions:
When we add a base to water, a base ionization reaction occurs in which protons are transferred from water molecules to base molecules. For example, adding pyridine to water yields hydroxide ions and pyridinium ions:
Notice that both these ionization reactions are represented as equilibrium processes. The relative extent to which these acid and base ionization reactions proceed is an important topic treated in a later section of this chapter. In the preceding paragraphs we saw that water can function as either an acid or a base, depending on the nature of the solute dissolved in it. In fact, in pure water or in any aqueous solution, water acts both as an acid and a base. A very small fraction of water molecules donate protons to other water molecules to form hydronium ions and hydroxide ions:
This type of reaction, in which a substance ionizes when one molecule of the substance reacts with another molecule of the same substance, is referred to as autoionization.
Pure water undergoes autoionization to a very slight extent. Only about two out of every 10 9 molecules in a sample of pure water are ionized at 25 °C. The equilibrium constant for the ionization of water is called the ion-product constant for water ( K w ):
$\ce{H2O(l) + H2O(l) <=> H3O+(aq) + OH-(aq)}$
$K_W = [H_3O^+] [OH^-]$
The slight ionization of pure water is reflected in the small value of the equilibrium constant; at 25 °C, K w has a value of 1.0 × 10 −14 . The process is endothermic, and so the extent of ionization and the resulting concentrations of hydronium ion and hydroxide ion increase with temperature. For example, at 100 °C, the value for K w is about 5.6 × 10 −13 , roughly 50 times larger than the value at 25 °C.