Definition of Acid and Base

According to the widely accepted Brønsted-Lowry theory, an acid is any substance that donates a proton (H⁺) to an aqueous solution; a base is any substance that accepts a proton, removing it from solution. By this definition, an H⁺ donor is an acid and an H⁺ acceptor is a base.

Body fluids are acidic if they have an abnormally high [H⁺]. Increased [H⁺] in the blood is called acidemia. Acidosis refers to a general condition characterized by the accumulation of H⁺ in body fluids.

The term alkali is synonymous with base. Originally, alkali referred to strong bases formed by metallic hydroxides, such as sodium hydroxide (NaOH) and potassium hydroxide (KOH). In medicine, the term alkali refers to any base (i.e., any H⁺ acceptor). Body fluids are alkaline (or basic) if they contain abnormally high amounts of base or if they contain abnormally low amounts of H⁺ compared with base. Therefore a body fluid can be alkaline even if the absolute concentration of base is normal. Increased base or reduced H⁺ in the blood is called alkalemia; alkalosis refers to an alkaline condition of body fluids.

Acids and bases dissociate or ionize in solution to form their component ions:

The double arrows mean that this process is reversible; equilibrium is rapidly established between undissociated molecules and their dissociated ions.

Acids and bases react with one another in an aqueous solution to form a salt and water, neutralizing one another (i.e., the products are neither acidic nor basic), as the following reaction shows:

In this reaction, hydrogen chloride (HCl) (hydrochloric acid) donates H⁺, which is accepted by the base component (OH⁻) of NaOH. A neutral salt, sodium chloride (NaCl), and water (H₂O) are the results of this reaction.

Strong and Weak Acids and Bases: Equilibrium Constants

Strong acids and bases ionize almost completely in an aqueous solution; weak acids and bases ionize only to a slight extent. An example of a strong acid is HCl; nearly 100% of HCl molecules dissociate to form H⁺ and Cl⁻, as the following reaction shows:

HCl→H++Cl− 1.

1.

At equilibrium, the concentration of HCl is extremely small compared with the concentration of H⁺ and Cl⁻. No arrow points to the left, emphasizing that HCl ionizes almost completely.

In contrast, a weak acid dissociates only partially as the following shows:

HCl→H++Cl− 2.

2.

In this reaction, HA is a generic representation of an undissociated weak acid; H⁺ and A⁻ represent the dissociated components of the weak acid. The short arrow pointing to the right indicates that at equilibrium the concentration of undissociated HA molecules is far greater than the concentrations of A⁻ or H⁺.

The equilibrium constant of an acid is a measure of the extent to which the acid molecules dissociate (ionize). The law of mass action states that the velocity of a reaction is proportional to the product of the concentrations of the reactants. The following dissociation of the generic weak acid HA illustrates this concept:

HCl→H++Cl− 3.

3.

Velocity 1 (V₁) is proportional to [HA], and velocity 2 (V₂) is proportional to [A⁻] multiplied by [H⁺]. At equilibrium, the number of HA molecules dissociating is equal to the number of A⁻ and H⁺ associating (i.e., V₁ and V₂ are equal). In this state, no further net change occurs in [HA], [A⁻], or [H⁺]. The following equation represents conditions at equilibrium:

[A−]×[H+][HA]=KA (small) 4.

4.

In equation 4, K_A represents the equilibrium constant for HA. (K_A is also known as the acid’s ionization or dissociation constant.) In this example, the weak acid HA has an extremely small K_A because the denominator [HA] is quite large with respect to the numerator ([H⁺] multiplied by [A⁻]). At a given temperature, the value of K_A is always the same for this particular weak acid (HA) regardless of its initial concentration. Each acid has its own unique ionization constant.

A strong acid, such as HCl, has a large K_A because the denominator [HCl] is extremely small compared with the numerator ([H⁺] multiplied by [Cl⁻]). This is shown as follows:

[H+]×[Cl−][HCl]=KA (large) 5.

5.

As shown by equations 4 and 5, K_A indicates an acid’s strength; the greater the value of K_A, the stronger the acid—that is, the more completely it dissociates.

Carbonic acid (H₂CO₃)—the product of the reaction between CO₂ and H₂O—is often described as a weak acid, when in reality it is a moderately strong acid. In the pH environment of the blood, H₂CO₃ dissociates instantly and almost completely as soon as it is formed.¹ The idea that H₂CO₃ is a weak acid comes from the fact that the reaction between dissolved CO₂ and H₂O is quite slow in blood plasma, and from the fact that the sum of dissolved CO₂ and H₂CO₃ is commonly lumped together and treated as the undissociated acid. This practice is sensible because H₂CO₃ and dissolved CO₂ are indistinguishable from one other by chemical analysis (see Chapter 9). Even though H₂CO₃ is a relatively strong acid, (1) its treatment as being synonymous with dissolved CO₂ and (2) its slow formation make it appear to be a weak acid.¹

Measuring Hydrogen Ion Concentration: The Concept of pH

[H⁺] in body fluids is extremely small (normally about 40 × 10⁻⁹ mol/L, or 40 billionths of 1 mole per liter [0.000000040 mol/L]). The prefix “nano” means billionths; therefore, [H⁺] of body fluids is about 40 nmol/L. In clinical medicine, acidity is generally expressed in terms of pH rather than nmol/L[H⁺].

The concept of pH was developed by Sørenson in 1909. The pH scale ranges from 0 to 14 pH units. The dissociation of pure water molecules (H₂O, or HOH) helps explain this concept:

[H+]×[Cl−][HCl]=KA (large) 6.

6.

The dissociation of H₂O is extremely slight. Therefore, water’s K_A is exceedingly small, as the following shows:

[H+]×[OH−][HOH]=KA 7.

7.

Equation 7 can be rearranged to yield the following:

[H+]×[OH−]=[HOH]KA 8.

8.

In practical terms, [HOH] is constant because of its extremely slight dissociation. Therefore, the right side of equation 8 can be considered constant:

[HOH]KA=KW 9.

9.

In equation 9, K_w represents the dissociation constant of water. Equation 8 then becomes the following:

[H+]×[OH−]=KW 10.

10.

The numerical value of K_w is about 10⁻¹⁴. At equilibrium, pure water contains 10⁻⁷ mol/L of H⁺ and 10⁻⁷ mol/L of OH⁻. If the value of [H⁺] is known, the value of [OH⁻] can be calculated because of the reciprocal relationship between [H⁺] and [OH⁻]. For example, if [H⁺] equals 10⁻³ mol/L, [OH⁻] equals 10⁻¹¹ mol/L.

pH is defined as the negative logarithm, or exponent (to the base 10), of [H⁺]; this is shown as follows:

pH=−log[H+] 11.

11.

Water’s pH ([H⁺] = 10⁻⁷ mol/L) is calculated as follows:

pH=−log(10−7)pH=−(−7)pH=7 12.

12.

A pH of 7 corresponds to [H⁺] of 100 nmol/L:

10−7mol/L=0.0000001 mol/L=0.000000100 mol/L

Because pH is the negative logarithm of [H⁺], a decrease in pH indicates an increase in [H⁺] (Box 10-1).

BOX 10-1   pH Scale: [H⁺] [OH^–] = 10^–14 mol/L

A chemically neutral solution (neither acidic nor basic) has a pH of 7.0. To the chemist, a solution with a pH less than 7.0 is acidic, and a solution with a pH greater than 7.0 is alkaline (see Box 10-1). By this standard, body fluids are slightly alkaline, as the following shows:

Body fluid [H+]=40×10−9 mol/L                       pH=−log (40×10−9)                       pH=−log 40+(−log 10−9)                       pH=−1.6+(−[−9])                       pH=−1.6+9                       pH=7.40

Because pH is a logarithmic scale, a change of one pH unit corresponds to a 10-fold change in [H⁺] (Figure 10-1). [H⁺] at a pH of 7.0 (100 nmol/L) is 10 times the [H⁺] at a pH of 8.0 (10 nmol/L). Figure 10-1 shows that the relationship between pH and [H⁺] is linear in the normal physiological range of body fluids (pH = 7.35 to 7.45). Table 10-1 shows that in the physiological range, a 1-nmol/L change in [H⁺] produces a 0.01 change in pH. The pH-to-[H⁺] relationship maintains some degree of linearity between pH values of 7.20 and 7.55. However, below a pH of 7.20, linearity deteriorates rapidly (see Figure 10-1).

Overview of Hydrogen Ion Regulation in Body Fluids

The body continually produces hydrogen ions, but the pH of body fluids varies minimally between 7.35 and 7.45 (45 to 35 nmol/L [H⁺]). As noted previously, this rigid control is necessary to sustain vital cellular enzyme activity. Hydrogen ions formed in the body arise from either volatile acids or nonvolatile acids, or fixed acids. Volatile acids in the blood arise from and are in equilibrium with a dissolved gas. The only volatile acid of physiological significance in the body is carbonic acid (H₂CO₃), which is in equilibrium with dissolved CO₂. Normal aerobic metabolism generates about 13,000 mM of CO₂ each day, producing an equal amount of H⁺:

In a process called isohydric buffering,¹ most of the hydrogen ions produced cause no pH change at the tissue level because newly forming deoxygenated hemoglobin immediately combines with the hydrogen ions (see Chapter 9). When blood reaches the lungs, hemoglobin releases these hydrogen ions, which combine with plasma bicarbonate ion (HCO3−) to form CO₂:

Ventilation eliminates carbonic acid (CO₂), keeping pace with its production. Isohydric buffering and ventilation are the two major mechanisms responsible for maintaining a stable pH in the face of massive CO₂ production.²

Catabolism of proteins continually produces fixed acids such as sulfuric and phosphoric acids. In addition, anaerobic metabolism produces lactic acid, another fixed acid. In contrast to carbonic acid, these fixed acids are nonvolatile and are not in equilibrium with a gaseous component. Therefore, the hydrogen ions of fixed acids must be buffered by bases in the body or eliminated in the urine by the kidneys. Compared with daily CO₂ production, fixed acid production is small, averaging only about 50 to 70 mEq per day.³ Certain diseases, such as untreated diabetes, increase fixed acid production. The resulting hydrogen ions stimulate the respiratory centers in the brain, increasing ventilation and CO₂ elimination, pulling the CO₂ hydration reaction to the left:

Thus, the respiratory system compensates for fixed acid accumulation, preventing a significant increase in [H⁺].

Bicarbonate and Nonbicarbonate Buffer Systems

Blood buffers are classified as either bicarbonate or nonbicarbonate buffer systems. The bicarbonate buffer system consists of carbonic acid (H₂CO₃) and its conjugate base, HCO3−. The nonbicarbonate buffer system consists mainly of phosphates and proteins, including hemoglobin. The blood buffer base is the sum of bicarbonate and nonbicarbonate bases, measured in millimoles per liter of blood.

The bicarbonate system is called an open buffer system because H₂CO₃ is in equilibrium with dissolved CO₂, which is readily removed by ventilation. That is, when HCO3− buffers H⁺, the product, H₂CO₃, is broken down into CO₂ and H₂O; ventilation continually removes CO₂ from the reaction, preventing it from reaching equilibrium. As long as ventilation is adequate, buffering activity continues without being slowed or stopped:

HCO3−+H+→H2CO3→H2O+CO2(exhaledgas)

Nonbicarbonate buffers are closed buffer systems because all components of acid-base reactions remain in the system. (In the following discussions, nonbicarbonate buffer systems are collectively represented as HBuf/Buf⁻, where HBuf is the weak acid, and Buf⁻ is the conjugate base.) When Buf⁻ buffers H⁺, the product, HBuf, builds up and eventually reaches equilibrium with the reactants, preventing further buffering activity:

HCO3−+H+→H2CO3→H2O+CO2(exhaledgas)

Table 10-2 summarizes the characteristics and components of bicarbonate and nonbicarbonate buffer systems.

Open and closed buffer systems differ in their ability to buffer fixed and volatile acids. They also differ in their ability to function in wide-ranging pH environments. Both systems are physiologically important, each playing a unique and essential role in maintaining pH homeostasis. Table 10-3 summarizes the approximate contributions of various blood buffers to the total buffer base. The bicarbonate buffer system has the greatest buffering capacity because it is an open system.

Bicarbonate and nonbicarbonate buffer systems do not function in isolation from one another; they are intermingled in the same solution (whole blood) and are in equilibrium with the same [H⁺] (Figure 10-2). As Figure 10-2 shows, increased ventilation increases the CO₂ removal rate, ultimately causing nonbicarbonate buffers (HBuf) to release H⁺. Conversely, cessation of ventilation reverses the reaction; CO₂ builds up, more H⁺ is generated, and, ultimately, more HBuf is formed.

pH of a Buffer System and Henderson-Hasselbalch Equation

Physiological buffer solutions consist mostly of undissociated acid molecules and only a small amount of H⁺ and conjugate base anions. The [H⁺] of a buffer solution can be calculated if the concentrations of its components and the acid’s equilibrium constant are known. Consider the bicarbonate buffer system:

HCO3−+H+→H2CO3→H2O+CO2(exhaledgas)