Acids, Bases, and pH

ACIDS AND BASES

Chemical definitions of acid and base

Acids and bases can be defined by their chemical properties in several different ways. However, because acid–base reactions important for biochemistry occur in aqueous solution (that is, in water), the Bronsted–Lowry definitions of acids and bases are generally used. In the Bronsted–Lowry scheme:

A proton is a hydrogen atom that has lost its electron and is designated H⁺. Thus, acids are chemical compounds that tend to give up H⁺ ions when placed in water. For example, when hydrochloric acid (HCl) is placed in water, the reaction that occurs is

HCl	+	H2O		H3O+	+	Cl–
proton donor (acid)		proton acceptor (base)

The HCl has donated its H⁺ ion, which was accepted by the water molecule. Notice how the H⁺ ion, once donated by the HCl molecule, does not simply float around in the water freely, but instead associates strongly with water (the proton acceptor) to form a hydronium ion (H₃O⁺). Thus, using the Bronsted–Lowry convention, the acid proton H⁺ is often written as H₃O⁺ to reflect this fact.

Similarly, ammonia (NH₃) is a base because it acts as a proton acceptor:

NH3	+	H2O		NH4+	+	OH–
proton acceptor (base)		proton donor (acid)

As you can see in the examples above, water can act not only as a proton acceptor, but also as a proton donor. In fact, water actually reacts with itself. The result is that even in a sample of pure water, not every molecule is in the form of H₂O. A very small proportion of molecules undergo this reaction, called autoionization:

H2O	+	H2O		H3O+	+	OH–
proton donor (acid)		proton acceptor (base)

Autoionization of water

Just how common is this reaction? Is a cup of water out of the tap really just a collection of charged molecules rather than H₂O?

Let's look closer. The equilibrium constant for the reaction of water molecules is

By convention, when a reaction involves water, the concentration of water is factored into the value of K_eq. For the reaction of water molecules, the resulting equilibrium constant is known as K_w, the ion product constant of water, and its value is 10^–14.

Scientists determined experimentally that the equilibrium constant for this reaction of water molecules is

The value of K_w is known as the ion product constant of water. Note that by convention, water molecules are not included in the equilibrium constant expression, because the reaction is happening in water, and the concentration of water is already factored in to the K_w constant. Therefore K_w is just the K_eq for the ionization of water, with the concentration of water factored into the constant (K_w = K_eq/[H₂O]).

Pure water is neutral, having neither an overall positive nor an overall negative charge. This is because each time a water molecule donates a proton, another accepts it, so that the concentrations of OH^– and H₃O⁺ are always equal:

Since the K_w expression for pure water states that the product of these two must equal 10^–14:

These are very minute concentrations, especially since pure water is 55.5 M. In other words, the fraction of H₃O⁺ and OH^– ions is very tiny compared to the overall H₂O concentration. In fact, there are over 500 million times more H₂O molecules than either H₃O⁺ or OH^– ions in water!

The K_w value is still very important for biochemists, however, since it aids in the understanding of acid–base chemistry. The K_w value is useful because it says that in any aqueous solution, the product of [OH^–] and [H₃O⁺] ions is a constant value. Thus, an acid solution contains not only H₃O⁺ ions, but also a lesser amount of OH^– ions.

Example 1

If a 10^–1 M solution of HCl was prepared, what would be the equilibrium concentrations of H₃O⁺ and OH^–?

Solution

Returning to the hydrochloric acid dissociation reaction from above, we can describe the reaction of the acid with water as:

HCl	+	H2O		H3O+	+	Cl–
proton donor		proton acceptor

Because HCl is a strong acid, essentially all of the acid will dissociate, and the concentration of H₃O⁺ ions would be equal to 10^–1 M. Plugging this into the K_w expression gives the following:

Kw = [OH–][H3O+] = 10–14 [OH–](10–1) = 10–14 [OH–] = 10–13

Similarly, a basic solution contains not only OH^– ions, but also some H₃O⁺ ions. To further grasp this concept, interact with the graph below. Note that while the ratio of hydronium ion to hydroxide ion changes, the K_w (the product of these two) does not change.