Acid-Base Balance

Acid-Base Balance

Will Beachey

Even small changes in hydrogen ion concentration [H+] can cause vital metabolic processes in the body to fail. Normal metabolism continually generates H+, and H+ regulation is of utmost biologic importance. Various physiologic mechanisms work together to keep [H+] of body fluids in a range compatible with life. This chapter helps the clinician understand how these mechanisms work and how to detect abnormalities in their function. With this knowledge, the clinician can make informed decisions about treating the underlying causes of acid-base disturbances.

Hydrogen Ion Regulation in Body Fluids

Acid-base balance refers to physiologic mechanisms that keep [H+] of body fluids in a range compatible with life. Hydrogen ions react readily with the protein molecules of vital cellular catalytic enzymes. Such reactions change the physical contour of the protein molecule and may render the enzyme inactive. To sustain life, the body must maintain the pH of fluids within a narrow range, from 7.35 to 7.45 (corresponding to a [H+] of 45 to 35 nmol/L).

Hydrogen ions formed in the body come from either volatile or fixed (nonvolatile) acids. A volatile acid is one that is in equilibrium with a dissolved gas. The only volatile acid of physiologic significance in the body is carbonic acid (H2CO3), which is in equilibrium with dissolved carbon dioxide (CO2). Normal aerobic metabolism generates approximately 13,000 mmol/L of CO2 each day, producing an equal amount of H+:


As CO2 diffuses into the blood at the tissue level, this reaction occurs primarily in the erythrocyte where it is catalyzed by carbonic anhydrase, an intracellular enzyme. In a process called isohydric buffering,1 most H+ produced in this fashion causes no change in pH because hemoglobin (Hb) in the erythrocyte immediately buffers the H+. When blood reaches the lungs, Hb releases H+ to form CO2 as shown:


In this way, ventilation eliminates carbonic acid, 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.

Catabolism of proteins continually produces fixed (nonvolatile) acids such as sulfuric and phosphoric acids. In addition, anaerobic metabolism produces lactic acid. In contrast to carbonic acid, these nonvolatile acids are not in equilibrium with a gaseous component. However, H+ of fixed acids can be buffered by bicarbonate ions (HCO3) and converted to CO2 and water (H2O) (see the previous reaction); the CO2 formed is eliminated in exhaled gas. Compared with daily CO2 production, fixed acid production is small, averaging only about 50 to 70 mEq/day.2 Certain diseases, such as untreated diabetes, increase fixed acid production. Hydrogen ions produced in this way stimulate respiratory centers in the brain. The resulting increase in ventilation eliminates more CO2, pulling the hydration reaction to the left:


In this way, the respiratory system compensates for fixed acid production, preventing a significant increase in [H+].

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 small extent. An example of a strong acid is hydrochloric acid (HCl). Nearly 100% of the HCl molecules dissociate to form H+ and Cl:

HClH++Cl (1)

image (1)

At equilibrium, the concentration of HCl is extremely small compared with either [H+] or [Cl]. There is no arrow pointing to the left in Reaction 1, emphasizing that HCl ionizes almost completely in solution. In contrast, carbonic acid is an example of a relatively weak acid:

H2CO3HCO3+H+ (2)

image (2)

The long arrow pointing to the left indicates that at equilibrium, the concentration of undissociated H2CO3 molecules is far greater than the concentration of HCO3 or H+.

The equilibrium constant of an acid is a measure of the extent to which the acid molecules dissociate (ionize). At equilibrium, the number of dissociating H2CO3 molecules in Reaction 2 is equal to the number of associating HCO3 and H+, even though the concentrations of reactants and products are unequal. In this state, no further change occurs in [H2CO3], [HCO3], or [H+]. At equilibrium, the following is true:

[H+]×[HCO3][H2CO3]=KA(Small) (3)

image (3)

Where KA is the equilibrium constant for H2CO3. (KA is also known as the acid’s ionization or dissociation constant.)

KA is small because the H2CO3 concentration is quite large with respect to the numerator of Reaction 3. The value of KA is always the same for H2CO3 at equilibrium, regardless of the initial concentration of H2CO3.

A strong acid, such as HCl, has a large KA because the denominator [HCl] is extremely small compared with the numerator ([H+] × [Cl]):

[H+]×[Cl][HCl]=KA(Large) (4)

image (4)

As shown by Equations 3 and 4, KA indicates the strength of an acid.

Buffer Solution Characteristics

A buffer solution resists changes in pH when an acid or a base is added to it. Buffer solutions are mixtures of acids and bases. The acid component is the H+ cation, formed when a weak acid dissociates in solution. The base component is the remaining anion portion of the acid molecule, known as the conjugate base. An important blood buffer system is a solution of carbonic acid and its conjugate base, HCO3:

H2CO3(Acid)HCO3(Conjugate base)+H+


In the blood, HCO3 combines with sodium ions (Na+) to form sodium bicarbonate (NaHCO3). If hydrogen chloride, a strong acid, is added to the H2CO3/NaHCO3 buffer solution, HCO3 reacts with the added H+ to form weaker carbonic acid molecules and a neutral salt:



The strong acidity of HCl is converted to the relatively weak acidity of H2CO3, preventing a large decrease in pH.

Similarly, if sodium hydroxide, a strong base, is added to this buffer solution, it reacts with the carbonic acid molecule to form the weak base, NaHCO3, and H2O:



The strong alkalinity of NaOH is changed to the relatively weak alkalinity of NaHCO3. pH change is minimized.

Bicarbonate and Nonbicarbonate Buffer Systems

Blood buffers are classified as bicarbonate or nonbicarbonate buffer systems. The bicarbonate buffer system consists of H2CO3 and its conjugate base, HCO3. The nonbicarbonate buffer system consists mainly of phosphates and proteins, including Hb. 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 H+ is buffered by HCO3, the product, H2CO3, is broken down into H2O and CO2 as long as ventilation removes CO2. The removal of CO2 from the reaction prevents it from reaching equilibrium with the reactants. For this reason, buffering activity can continue without being slowed or stopped:

HCO3+H+H2CO3H2O+CO2(Exhaled gas)


A nonbicarbonate buffer system is called a closed buffer system because all the 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 H+ is buffered by Buf, the product, HBuf, accumulates and eventually reaches equilibrium with the reactants, preventing further buffering activity:



Box 13-1 summarizes the characteristics and components of bicarbonate and nonbicarbonate buffer systems.

Open and closed buffer systems play different roles in buffering fixed and volatile acids, and they differ in their ability to function in wide-ranging pH environments. Volatile acid (H2CO3) accumulates only if ventilation cannot eliminate CO2 fast enough to keep up with the body’s CO2 production. In such a case, the reaction between CO2 and H2O moves continually to the right, creating more H2CO3 and, ultimately, more H+ and HCO3. The HCO3 produced in this way is incapable of buffering the H+ with which it was coproduced. The only buffer system that can buffer the H+ of volatile acid is the nonbicarbonate buffer system. Both nonbicarbonate and bicarbonate buffer systems can buffer the H+ produced by fixed acids; this is true of the bicarbonate buffer system only if ventilation is not impaired and CO2 can be adequately eliminated. Both systems are physiologically important, each playing a unique and essential role in maintaining pH homeostasis. Table 13-1 summarizes the approximate contributions of various blood buffers to the total buffer base. Bicarbonate buffers have the greatest buffering capacity because they function in an open system.

Bicarbonate and nonbicarbonate buffer systems do not function in isolation from one another but are intermingled in the same solution (whole blood), in equilibrium with the same [H+] (Figure 13-1). Increased ventilation increases the CO2 removal rate, causing nonbicarbonate buffers (Hbuf) to release H+. Decreased ventilation ultimately causes Hbuf to accept more H+.

pH of a Buffer System: Henderson-Hasselbalch Equation

Buffer solutions in body fluids consist of mostly 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 the buffer’s components and the acid’s equilibrium constant are known. Consider the bicarbonate buffer system. As described earlier, the equilibrium constant (KA) for H2CO3 is as follows:



[H+] can be calculated by algebraic rearrangement of this equation, as follows:



[H+] is determined by the ratio between undissociated acid molecules [H2CO3] and base anions [HCO3]. This equation is the basis for deriving the Henderson-Hasselbalch (H-H) equation:



pH is a logarithmic expression of [H+], and the term 6.1 is the logarithmic expression of the H2CO3 equilibrium constant. Because dissolved carbon dioxide (Pco2 × 0.03) is in equilibrium with and directly proportional to blood [H2CO3], and because blood Pco2 is more easily measured than [H2CO3], dissolved CO2 is used in the denominator of the H-H equation. The H-H equation is specific for calculating the pH of the bicarbonate buffer system of the blood. The calculation of this pH is important because it equals the pH of blood plasma; because all buffer systems in the blood are in equilibrium with the same pH, the pH of one buffer system is the same as the pH of the entire plasma solution (the isohydric principle).1

Clinical Use of Henderson-Hasselbalch Equation

The H-H equation allows the pH, [HCO3], or Pco2 to be computed if two of these three variables are known (shown as follows for Pco2 and HCO3):





Blood gas analyzers measure pH and PCO2 but compute [HCO3]. Assuming a normal arterial pH of 7.40 and a PaCO2 of 40 mm Hg, arterial [HCO3] can be calculated as follows:







Solving for [HCO3]:

[HCO3]=antilog (7.406.1)×1.2=antilog(1.3)×1.2=20×1.2=24 mEq/L


The H-H equation is useful for checking a clinical blood gas report to see if the pH, PCO2, and [HCO3] values are compatible with one another. In this way, transcription errors and analyzer inaccuracies can be detected. It is also clinically useful to predict what effect changing one H-H equation component will have on the other components. For example, a clinician may want to know how low the arterial blood pH will fall for a given increase in PaCO2.

Physiologic Roles of Bicarbonate and Nonbicarbonate Buffer Systems

The functions of bicarbonate and nonbicarbonate buffer systems are summarized in Table 13-2.

Bicarbonate Buffer System

The bicarbonate buffer system is particularly effective in the body because it is an open system—that is, one of its components (CO2) is continually removed through ventilation:

(Exhaled gas)CO2+H2OH2CO3HCO3+H+


In this way, HCO3 continues to buffer H+ as long as ventilation continues. Hypothetically, this buffering activity can continue until all body sources of HCO3 are used up in binding H+.

The bicarbonate buffer system can buffer only fixed acid. An increased fixed acid load in the body (e.g., lactic acid) reacts with HCO3 of the bicarbonate buffer system:


As shown, the process of buffering fixed acid produces CO2, which is eliminated in exhaled gas. Large amounts of acid are buffered in this fashion. If the ability to ventilate is impaired, this type of buffering cannot occur.

The bicarbonate buffer system cannot buffer carbonic (volatile) acid, which accumulates in the blood whenever ventilation fails to eliminate CO2 as fast as it is produced (hypoventilation). The resulting accumulation of CO2 drives the hydration reaction in the direction that produces more carbonic acid, H+, and HCO3, as shown:


H+ produced by dissociating H2CO3 molecules cannot be buffered by the simultaneously produced HCO3 because hypoventilation prevents the reaction from reversing its direction. The closed nonbicarbonate buffer systems are the only buffers that can buffer carbonic acid.

Nonbicarbonate Buffer System

Table 13-1 lists the nonbicarbonate buffers in the blood. Of these, Hb is the most important because it is the most abundant. As mentioned, these buffers are the only ones available to buffer carbonic acid. However, they can buffer H+ produced by any acid, fixed or volatile. Because nonbicarbonate buffers (Buf/HBuf) function in closed systems, the products of their buffering activity eventually accumulate, slowing or stopping further buffering activity:



This slowing or stopping of buffering activity means that not all of the Buf is available for buffering activity. At equilibrium (denoted by the double arrow), Buf still exists in solution but cannot combine further with H+. In contrast, most of the HCO3 in the bicarbonate buffer system is available for buffering activity because it functions in an open system where equilibrium between reactants and products does not occur. Both open and closed systems function in a common fluid compartment (blood plasma) as illustrated in the following equation:


Most of the added fixed acid is buffered by HCO3 because ventilation continually pulls the reaction to the left. Smaller amounts of H+ react with Buf because equilibrium is approached, slowing the reaction.

Acid Excretion

Bicarbonate and nonbicarbonate buffer systems are the immediate defense against the accumulation of H+. However, if the body fails to eliminate the remaining acids, these buffers are soon exhausted, and the pH of body fluids quickly decreases to life-threatening levels.

The lungs and kidneys are the primary acid-excreting organs. The lungs can excrete only volatile acid (i.e., the CO2 from dissociating H2CO3). However, as discussed previously, bicarbonate buffers effectively buffer the H+ originating from fixed acid, converting it to H2CO3 and to CO2 and H2O. By eliminating the CO2, the lungs can rapidly remove large quantities of fixed acid from the blood. The kidneys also remove fixed acids but at a slow pace. In healthy individuals, the acid excretion mechanisms of lungs and kidneys are delicately balanced. In individuals affected by disease, failure of one system can be partially offset by a compensatory response of the other.


Because the volatile acid H2CO3 is in equilibrium with dissolved CO2, the lungs can decrease blood H2CO3 concentration through ventilation. The elimination of CO2 is crucial because normal aerobic metabolism produces large quantities of CO2, which reacts with H2O to form large quantities of H2CO3. The reaction between fixed acids and bicarbonate buffers also produces H2CO3. H2CO3 generated by both pathways is eliminated as CO2 through the lungs. Approximately 24,000 mmol/L of CO2 is removed from the body daily through normal ventilation. CO2 excretion of the lungs does not remove H+ from the body. Instead, the chemical reaction that breaks down H2CO3 to form CO2 binds H+ in the harmless H2O molecule:


Jun 12, 2016 | Posted by in RESPIRATORY | Comments Off on Acid-Base Balance
Premium Wordpress Themes by UFO Themes