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Exercise Physiology: Theory and Application to Fitness and Performance, 10e

Chapter 11: Acid-Base Balance during Exercise

OBJECTIVES

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By studying this chapter, you should be able to do the following:

Define the terms acid, base, and pH.
Discuss the principal ways that hydrogen ions are produced during exercise.
Discuss the importance of acid-base regulation to exercise performance.
List the principal intracellular and extracellular buffers.
Explain the role of respiration in the regulation of acid-base status during exercise.
Outline the interaction between intracellular/ extracellular buffers and the respiratory system in acid-base regulation during exercise.

OUTLINE

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Acids, Bases, and pH 257
Hydrogen Ion Production during Exercise 258
Importance of Acid-Base Regulation during Exercise 260
Acid-Base Buffer Systems 260
- Intracellular Buffers 260
- Influence of Muscle Fiber Type and Exercise Training on Intracellular Buffer Capacity 261
Extracellular Buffers 261
Respiratory Influence on Acid-Base Balance 262
Regulation of Acid-Base Balance via the Kidneys 264
Regulation of Acid-Base Balance during Exercise 265

KEY TERMS

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acid

acidosis

alkalosis

base

buffer

hydrogen ion

ion

respiratory compensation

strong acids

strong bases

An ion is any atom that is missing electrons or has gained electrons. A hydrogen ion is formed by the loss of the electron; molecules that release hydrogen ions are called acids. Substances that readily combine with hydrogen ions are termed bases. In physiology, the concentration of hydrogen ions in body fluids is expressed in pH units. The pH of body fluids must be regulated (i.e., normal arterial blood pH = 7.40 ±.02) in order to maintain homeostasis. This regulation of the pH of body fluids is important because changes in hydrogen ion concentrations can alter the rates of enzyme-controlled metabolic reactions and modify numerous other normal body functions. Therefore, acid-base balance is primarily concerned with the regulation of hydrogen ion concentrations. High-intensity exercise can present a serious challenge to hydrogen ion control systems due to hydrogen ion production, and hydrogen ions may limit performance in some types of intense activities (7, 14, 16, 20, 40). Therefore, given the potential detrimental influence of hydrogen ion accumulation on exercise performance, it is important to have an understanding of acid-base regulation.

ACIDS, BASES, AND pH

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In biological systems, one of the simplest but most important ions is the hydrogen ion. The concentration of hydrogen ions influences the rates of chemical reactions, the shape and function of enzymes as well as other cellular proteins, and the integrity of the cell itself (10, 46).

An acid is defined as a molecule that releases hydrogen ions and thus can raise the hydrogen ion concentration of a solution above that of pure water. In contrast, a base is a molecule that is capable of combining with hydrogen ions, which would lower the hydrogen ion concentration of the solution.

Acids that give up hydrogen ions (ionize) more completely are called strong acids. For example, sulfuric acid is produced by the metabolism of sulfur-containing amino acids (e.g., cysteine) and is a strong acid. At normal body pH, sulfuric acid liberates almost all of its hydrogen ions and therefore elevates the hydrogen ion concentration of the body.

Bases that ionize more completely are defined as strong bases. The bicarbonate ion ${HCO}_{3}^{-}$ is an example of a biologically important strong base. Bicarbonate ions are found in relatively large concentrations in the body. Importantly, bicarbonate ions are capable of combining with hydrogen ions to form a weak acid called carbonic acid. The role of ${HCO}_{3}^{-}$ in the regulation of acid-base balance during exercise will be discussed later in the chapter.

As stated earlier, the concentration of hydrogen ions is expressed in pH units on a scale that runs from 0 to 14. Solutions that contain pH values below 7 are acidic, whereas pH values above 7 are basic. The pH of a solution is defined as the negative logarithm of the hydrogen ion concentration (H⁺). Recall that a logarithm is the exponent that indicates the power to which one number must be raised to obtain another number. For instance, the logarithm of 100 to the base 10 is 2. Thus, the definition of pH can be written mathematically as:

pH = - {log}_{10} [H^{+}]

Asanexample, if the [H⁺] = 40 nM (0.000000040 M), then the pH would be 7.40. A solution is considered neutral in terms of acid-base status if the concentration of H⁺ and hydroxyl ions (OH⁻) are equal. This is the case for pure water, in which the concentrations for both H⁺ and OH⁻ are 0.00000010 M. Thus, the pH of pure water is:

\begin{array}{l} p H (p u r e w a t e r) & = & - l o g_{10} [H^{+}] \\ = & 7 . 0 \end{array}

Figure 11.1 illustrates the pH scale and highlights that the normal pH of arterial blood is 7.4. Note that as the hydrogen ion concentration increases, pH declines and the acidity of the blood increases, resulting in a condition termed acidosis. Conversely, as the hydrogen ion concentration decreases, pH increases and the solution becomes more basic (alkalotic). This condition is termed alkalosis. Conditions leading to acidosis or alkalosis are summarized in Figure 11.2.

Figure 11.1

The pH scale. If the pH of arterial blood drops below the normal value of 7.4, the resulting condition is termed acidosis. In contrast, if the pH increases above 7.4, blood alkalosis occurs.

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Figure 11.2

Acidosis results from an accumulation of acids or a loss of bases. Alkalosis results from a loss of acids or an accumulation of bases.

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Again, the normal arterial pH is 7.4, and in healthy people, this value varies less than 0.05 pH unit (42). Failure to maintain acid-base homeostasis in the body can have serious consequences. Indeed, even small changes in blood pH can have negative effects on the function of organ systems. In particular, acidosis can cause dysfunction of both the central nervous system and the cardiovascular system. For example, acidosis can lead to mental confusion and a feeling of fatigue. Further, both increases and decreases in arterial pH can promote abnormal electrical activity in the heart, resulting in rhythm disturbances, and arterial pH values below 7.0 and above 7.8 can have lethal consequences (12, 42). Numerous disease states can result in acid-base disturbances in the body and are introduced in Clinical Applications 11.1.

CLINICAL APPLICATIONS 11.1

Conditions and Diseases That Promote Metabolic Acidosis or Alkalosis

As mentioned previously, failure to maintain acid-base homeostasis in the body can have serious consequences and can lead to dysfunction of essential organs. Indeed, even relatively small changes in arterial pH (i.e., 0.1–0.2 pH units) can have a significant negative affect on organ function (12).

Metabolic acidosis occurs due to a gain in the amount of acid in the body. A number of conditions and disease states can promote metabolic acidosis. For example, long-term starvation (i.e., several days without food) can result in metabolic acidosis due to the production of ketoacids in the body as a by-product of high levels of fat metabolism. In extreme circumstances, the type of metabolic acidosis can result in death.

Diabetes is a common metabolic disease that promotes metabolic acidosis. Uncontrolled diabetes can result in a form of metabolic acidosis called diabetic ketoacidosis. Similar to starvation-induced acidosis, this form of acidosis is also due to the overproduction of ketoacids due to high levels of fat metabolism. Worldwide, numerous deaths occur each year from this form of acidosis (12, 42).

Metabolic alkalosis results from a loss of acids from the body. Conditions leading to metabolic alkalosis include severe vomiting and diseases such as kidney disorders that result in a loss of acids (12, 42). In both of these circumstances, the loss of acids results in an overabundance of bases in the body, leading to metabolic alkalosis.

IN SUMMARY

Acids are defined as molecules that release hydrogen ions, which increases the hydrogen ion concentration of an aqueous solution (e.g., body fluids).
Bases are molecules that are capable of combining with hydrogen ions; this results in a lower hydrogen ion concentration and an increase in pH.
The concentration of hydrogen ions in a solution is quantified by pH units. The pH of a solution is defined as the negative logarithm of the hydrogen ion concentration:

pH = – log₁₀ [H⁺]

HYDROGEN ION PRODUCTION DURING EXERCISE

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Much debate exists regarding the primary sites of hydrogen ion production during exercise (3, 18, 28). Nonetheless, it is clear that high-intensity exercise results in a marked decrease in both muscle and blood pH. Current evidence indicates that the exercise-induced decrease in muscle pH is due to multiple factors. Three important contributors to exercise-induced muscle acidosis are outlined here (11, 25):

Exercise-induced production of carbon dioxide and carbonic acid in the working skeletal muscles. Carbon dioxide, an end product in the oxidation of carbohydrates, fats, and proteins, is regarded as an acid by virtue of its ability to react with water to form carbonic acid (H₂CO₃), which in turn dissociates to form H⁺ and bicarbonate ( ${HCO}_{3}^{-}$ ):
${CO}_{2} + H_{2} O \leftrightarrow H^{+} + {HCO}_{3}^{-}$

Because CO₂ is a gas and can be eliminated by the lungs, it is often referred to as a volatile acid. During the course of a day, the body produces large amounts of CO₂ due to normal metabolism. During exercise, metabolic production of CO₂ increases and therefore adds a “volatile acid” load on the body.
Exercise-induced production of lactic acid in the working muscle. Although controversy exists, it is likely that production of lactic acid (lactate) in the muscle during heavy exercise is a key factor that causes the decrease in muscle pH (3, 28).
Exercise-induced ATP breakdown in the working muscles. The breakdown of ATP for energy during muscle contraction results in the release of H⁺ ions (36). For example, the breakdown of ATP results in the following reaction:
$ATP + H_{2} O \to ADP + {HPO}_{4} + H^{+}$

Therefore, the breakdown of ATP alone during exercise can be an important source of H⁺ ions in contracting muscles.

To summarize, debate exists about the primary causes of exercise-induced acidosis. Nonetheless, it is clear that contracting muscles can produce H⁺ from several sites; therefore, the cause of exercise-induced acidosis is likely due to the production of H⁺ ions from several different sources. Figure 11.3 provides a summary of the three major metabolic processes that serve as primary sources of hydrogen ions in working skeletal muscle. A discussion of which sports or exercise activities pose the greatest risk of acid-base disturbances is contained in A Closer Look 11.1. Table 11.1 provides an estimate of the risk of developing acidosis in several popular sports.

Figure 11.3

Primary sources of hydrogen ions in contracting skeletal muscles.

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A CLOSER LOOK 11.1

Sport and Exercise-Induced Disturbances in Muscle Acid-Base Balance

What types of sports or exercise promote acid-base disturbances in skeletal muscle? In general, any sport or exercise activity requiring high-intensity muscular contractions that persist for 45 seconds or longer can produce significant amounts of hydrogen ions, resulting in a decrease in muscle and blood pH. Table 11.1 provides a list of popular sports and the risk of developing muscle acid-base disturbances in each sport. Note that the risk of acid-base disturbances is classified as being high, moderate, or low. For sports classified in the low-to-moderate risk category (e.g., soccer), the risk of muscle acid-base disturbance is often connected to effort of the competitor. That is, an aggressive athlete that is constantly playing at 100% effort is more likely to develop an acid-base imbalance compared to the athlete that plays at “half speed” during the game. Also, note that track and field races like the 5,000- and 10,000-meter runs are listed as moderate and low to moderate risks for acid-base disturbances, respectively. In these types of track races, athletes generally produce small amounts of hydrogen ions during most of the race, but a sustained sprint to the finish during the last lap of these events could result in production of relatively large amounts of hydrogen ions and, therefore, result in acid-base disturbances in the muscle and blood.

Because acid-base disturbances can contribute to muscle fatigue and limit exercise performance, it is not surprising that scientists have investigated the possibility that increasing the buffering capacity of the blood results in improved exercise performance. One approach to this problem is to ingest large quantities of a buffer prior to exercise. A discussion of this topic can be found in The Winning Edge 11.1.

IN SUMMARY

High-intensity exercise results in a marked decrease in both muscle and blood pH.
This exercise-induced decrease in muscle pH is due to multiple factors, including (1) increased production of carbon dioxide; (2) increased production of lactic acid; and (3) the release of H⁺ ions during the breakdown of ATP.
Only sporting events that involve high-intensity exercise present a threat to acid- base disturbances during exercise (see Table 11.1).

TABLE 11.1

Risk of Developing an Acid-Base Disturbance in Popular Sports

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TABLE 11.1

Risk of Developing an Acid-Base Disturbance in Popular Sports

Sport

Risk of Acid-Base Disturbance

Baseball

Low

Basketball

Low-to-moderate

Boxing

Low-to-moderate

Cross-country skiing

Low

Football (American)

Low

100-meter sprint

Low

100-meter swim

High

400-meter run

High

800-meter run

High

1,500-meter run

Moderate-to-high

5,000-meter run

Moderate

10,000-meter run

Low-to-moderate

Marathon run

Low

Soccer

Low-to-moderate

Weight lifting (low repetitions)

Low

Volleyball

Low

IMPORTANCE OF ACID-BASE REGULATION DURING EXERCISE

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As discussed earlier, high-intensity exercise results in the production of large amounts of hydrogen ions. These hydrogen ions can exert a powerful effect on other molecules by interacting with molecules and thus altering their shape and function (14, 16, 33, 41). For example, high levels of hydrogen ions can alter the function of enzymes by decreasing their activity; this could have a negative effect on normal metabolism.

How can changes in muscle pH affect exercise performance? An increase in the intramuscular hydrogen ion concentration can impair exercise performance in at least two ways. First, an increase in the hydrogen ion concentration reduces the muscle cell’s ability to produce ATP by inhibiting key enzymes involved in both anaerobic (glycolysis) and aerobic production of ATP (14, 19). Second, hydrogen ions compete with calcium ions for binding sites on troponin, thereby hindering the contractile process (14, 45). This is discussed again in Chap. 19.

IN SUMMARY

Hydrogen ions can interact with molecules and alter their original size and function.
Failure to maintain acid-base homeostasis during exercise can impair performance by inhibiting metabolic pathways responsible for the production of ATP or by interfering with the contractile process in the working muscle.

ACID-BASE BUFFER SYSTEMS

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From the preceding discussion, it is clear that a rapid accumulation of hydrogen ions during intense exercise can negatively influence muscular performance. Therefore, it is important that the body has control systems capable of regulating acid-base status to prevent drastic decreases or increases in pH. One of the most important means of regulating hydrogen ion concentrations in body fluids is by the aid of buffers. A buffer resists pH change by removing hydrogen ions when the hydrogen ion concentration increases, and releasing hydrogen ions when the hydrogen ion concentration falls.

Buffers often consist of a weak acid and its associated base (called a conjugate base). The ability of individual buffers to resist pH change is dependent upon two factors. First, individual buffers differ in their intrinsic physiochemical ability to act as buffers. Simply stated, some buffers are better than others. A second factor influencing buffering capacity is the concentration of the buffer present (22, 39). The greater the concentration of a particular buffer, the more effective the buffer can be in preventing pH change.

Intracellular Buffers

The first line of defense in protecting against exercise-induced decreases in pH resides within the muscle fiber itself. Indeed, muscle fibers can protect against the accumulation of hydrogen ions and a decrease in cellular pH in two different ways. First, muscle fibers contain numerous classes of chemical buffers that can eliminate hydrogen ions. Second, the muscle fiber membrane (sarcolemma) contains two major types of hydrogen ion transporters that carry hydrogen ions from inside the muscle fiber into the interstitial space. Let’s discuss each of these buffering systems in more detail.

Four major classes of intracellular chemical buffer systems exist in the cytosol of muscle fibers: (1) bicarbonate, (2) phosphates, (3) cellular proteins, and (4) histidine-dipeptides (primarily carnosine) (1, 25, 39). The presence of bicarbonate in skeletal muscle fibers is a useful buffer during exercise (4, 16). Further, several phosphate-containing compounds also serve as intracellular buffers in skeletal muscle fibers, and phosphate buffers are of particular importance at the beginning of exercise (22). Numerous cellular proteins contain the amino acid histidine, which possesses an ionizable group that can accept (i.e., buffer) hydrogen ions. This combination of a hydrogen ion with this cellular protein results in the formation of a weak acid, which protects against a decrease in cellular pH. Finally, muscle fibers also contain several histidine-dipeptides (dipeptides are two linked amino acids) that are capable of buffering hydrogen ions. One of the major histidine-dipeptides found in skeletal muscle is carnosine, and growing evidence indicates that carnosine is an important buffer in muscle fibers. Collectively, these four intracellular chemical buffer systems work as unit to provide protection against decreases in muscle pH during intense exercise. A summary of these buffer systems is provided in Table 11.2.

TABLE 11.2

Chemical Acid-Base Buffer Systems

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TABLE 11.2

Chemical Acid-Base Buffer Systems

Buffer

Constituents

Actions

Bicarbonate

Bicarbonate ${HCO}_{3}^{-}$

Converts strong acid into weak acid

Phosphates

Phosphates ${HCO}_{4}^{-}$

Converts strong acid into weak acid

Proteins

Proteins containing histidine groups

Accepts hydrogens

Histidine-dipeptides

Primary component is carnosine

Accepts hydrogens

Muscle pH homeostasis is also regulated by the transport of hydrogen ions from muscle fibers into the interstitial space; these hydrogen ions are then buffered by extracellular fluid and blood buffer systems. Two primary transporters that move hydrogen ions across the sarcolemma are the sodium-hydrogen exchanger (NHE) and monocarboxlate transporters (MCTs). The NHE moves sodium ions into the muscle and hydrogen ions out of the muscle into the interstitial space (Fig. 11.4). Specifically, this transporter moves one hydrogen ion out the cell in exchange for one sodium ion. The second hydrogen ion transporter is the MCT. Human skeletal muscles contain two different MCTs that are labeled as MCT1 and MCT4. Both of these molecules mediate a one-to-one co-transport of lactate and hydrogen ions out of the muscle fiber. In other words, these MCTs carry one lactate molecule and one hydrogen ion across the sarcolemma. Research reveals that these transporters are important in regulating muscle pH during high-intensity exercise.

Figure 11.4

Illustration of the two important hydrogen ion (H⁺) transporters in skeletal muscle fibers. The sodium-hydrogen exchanger (NHE) moves one sodium (Na⁺) molecule into the fiber in exchange for transporting one H⁺ outward. The monocarboxlate transporters (MCTs) co-transport one lactate molecule and one H⁺ out of the muscle fiber.

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Influence of Muscle Fiber Type and Exercise Training on Intracellular Buffer Capacity

As discussed in the previous section, muscle fibers can protect against the accumulation of hydrogen ions and a decrease in cellular pH in two different ways: (1) intracellular buffers and (2) hydrogen ion transporters that move hydrogen ions from inside the muscle fiber into the interstitial space. Studies reveal that, compared to slow (i.e., type I) muscle fibers, the intracellular buffering capacity is higher in fast (i.e., type II) muscle fibers (1). Obviously, this higher buffer capacity in fast fibers is advantageous for performance during high-intensity exercise because fast muscle fibers produce high levels of both lactate and hydrogen ions during intense exercise.

Several studies show that high-intensity exercise training improves muscle buffer capacity in both untrained and trained individuals (8, 15). The precise mechanism(s) to explain exercise training-induced improvements in muscle buffer capacity remains a topic of debate. Nonetheless, training has been shown to increase the intracellular content of both carnosine and hydrogen ion transporters (i.e., MCT) in muscle fibers (8, 15). Therefore, it seems likely that a training-induced improvement in muscle buffer capacity is due to increases in both carnosine and hydrogen ion transporters in skeletal muscle.

Extracellular Buffers

The blood contains three principal buffer systems (4, 12, 16, 25): (1) proteins, (2) hemoglobin, and (3) bicarbonate. Blood proteins act as buffers in the extracellular compartment. Like intracellular proteins, these blood proteins contain ionizable groups that are weak acids and therefore act as buffers. However, because blood proteins are found in small quantities, their usefulness as buffers during heavy exercise is limited.

In contrast, hemoglobin is a particularly important protein buffer and is a major blood buffer during resting conditions. In fact, hemoglobin has approximately six times the buffering capacity of plasma proteins due to its high concentration (12, 25). Also contributing to the effectiveness of hemoglobin as a buffer is the fact that deoxygenated hemoglobin is a better buffer than oxygenated hemoglobin. As a result, after hemoglobin becomes deoxygenated in the capillaries, it is better able to bind to hydrogen ions formed when CO₂ enters the blood from the tissues. Thus, hemoglobin helps to minimize pH changes caused by loading of CO₂ into the blood (12).

The bicarbonate buffer system is probably the most important buffer system in the body (4, 25). This fact has been exploited by some investigators who have demonstrated that an increase in blood bicarbonate concentration (ingestion of bicarbonate) results in an improvement in performance in some types of exercise (4, 7, 20, 26) (see The Winning Edge 11.1).

THE WINNING EDGE 11.1

Exercise Physiology Applied to Sports

Nutritional Supplements to Buffer Exercise-Induced Acid-Base Disturbances and Improve Performance

Because intramuscular acidosis is associated with muscle fatigue, numerous studies have explored nutritional supplements to increase buffering capacity in hopes of improving athletic performance during high-intensity exercise. Indeed, it appears that supplements including sodium bicarbonate, sodium citrate, and beta-alanine have the potential to improve buffering capacity and enhance exercise performance during, high-intensity exercise. Let’s discuss these supplement strategies to improve muscle buffering capacity in more detail.

Sodium bicarbonate. Bicarbonate is a buffer that plays an important role in maintaining both extracellular and intracellular pH, despite its inability to freely cross the muscle membrane (i.e., sarcolemma). Although controversy exists (2), many studies conclude that performance during high-intensity exercise is improved when athletes ingest sodium bicarbonate prior to exercise (6, 7, 26, 30–32, 34, 37, 40). Specifically, results from numerous studies reveal that boosting the blood-buffering capacity by ingestion of sodium bicarbonate increases time to exhaustion during high-intensity exercise (e.g., 80% to 120% ${\dot{V} O}_{2}$ max). For example, a recent survey of the scientific literature reveals that sodium bicarbonate is effective in improving a 60-second “all out” exercise bout by approximately 2% (5). Further, laboratory studies employing repeated bouts of high-intensity exercise (i.e., >100% ${\dot{V} O}_{2}$ max) have reported that ingestion of sodium bicarbonate prior to exercise can enhance performance by more than 8% (18). In addition to these laboratory studies, evidence exists that sodium bicarbonate is also beneficial to sport performance in activities where the metabolic demands are primarily anaerobic, such as judo, swimming, and water polo (18).

It appears that sodium bicarbonate improves physical performance by increasing the extracellular buffering capacity, which, in turn, increases the transport of hydrogen ions out of the muscle fibers (38). This would reduce the interference of hydrogen ions on muscle ATP production and/or the contractile process itself.

In deciding whether to use sodium bicarbonate prior to a sporting event, an athlete should understand the risks associated with this decision. Ingestion of sodium bicarbonate in the doses required to improve blood-buffering capacity can cause gastrointestinal problems, including diarrhea and vomiting (7, 37).

Sodium Citrate. Similar to sodium bicarbonate, sodium citrate is another agent capable of increasing extracellular buffering capacity (18). The question of whether ingestion of sodium citrate can improve exercise performance during high-intensity exercise remains controversial because experimental results are often inconsistent. Nonetheless, a review of the research literature suggests that although low doses of sodium citrate does not improve performance, ingestion of high doses of sodium citrate (i.e., >0.5 grams/kilogram body weight) improves performance during high-intensity cycling exercise lasting 120 to 240 seconds (18).

Unfortunately, similar to sodium bicarbonate, ingestion of high doses of sodium citrate can produce undesired side effects such as nausea, gastrointestinal discomfort, and headaches. Therefore, before deciding whether to use sodium citrate prior to competition, athletes should consider the negative side effects associated with the use of sodium citrate.

Beta-alanine. Recent evidence suggests that supplementation with beta-alanine can play a beneficial role in protecting against exercise-induced acidosis and improve performance during short, high-intensity exercise (39). Beta-alanine is a non-essential amino acid produced in the liver, gut, and kidney. However, fasting blood levels of beta-alanine are low indicating that endogenous synthesis of this amino is limited.

The link between beta-alanine and protection against acidosis is linked to the fact that beta-alanine is an important precursor for the synthesis of carnosine in skeletal muscle. As discussed in the text, carnosine is a small molecule (dipeptide) found in the cytoplasm of excitable cells (i.e., neurons, skeletal and cardiac muscle fibers) (18). Carnosine has several important physiological functions including the ability to buffer hydrogen ions and protect against exercise-induced decreases in cellular pH (18).

The availability of beta-alanine is the rate limiting factor for carnosine synthesis in muscle fibers. However, supplementation (2 to 3 grams/day) with beta-alanine for >2 weeks results in a 60% to 80% increase in muscle carnosine levels. Importantly, this increase in muscle carnosine levels is associated with a 3% to 5% increase in muscle buffering capacity (18). Theoretically, this increase in intracellular buffering capacity could translate into improvements in performance during high-intensity exercise. In this regard, growing evidence suggests that beta-alanine supplementation improves high-intensity exercise performance in both running and cycling events lasting 1 to 4 minutes (18). Interestingly, some of these studies have recorded performance improvements of 12% to 14% (18).

The only known side effect of beta-alanine supplementation is paraesthesia (tingling of the skin); this sensation begins within 20 minutes after ingestion and lasts up to 60 minutes (18). Although harmless, paraesthesia is unpleasant and, several investigations have reported that paraesthesia can be avoided by staggering dosing throughout the day (18).

Final words of caution on use of “supplement buffers” to improve exercise performance. Regardless of the type of buffer ingested, extremely large doses of any buffer can result in severe alkalosis and pose negative health consequences. Another important consideration in the use of any ergogenic aid is the legality of the drug. In regard to the use of acid-base buffers, some sports regulatory agencies have banned the use of sodium buffers during competition. See Sahlin (2014) and Junior et al. (2015) in Suggested Readings for detailed information about the possible ergogenic effects of sodium bicarbonate, sodium citrate, and beta-alanine.

The bicarbonate buffer system involves carbonic acid (H₂CO₃), which undergoes the following dissociation reaction to form bicarbonate ( ${HCO}_{3}^{-}$ ):

{CO}_{2} + H_{2} O \leftrightarrow H_{2} {CO}_{3} \leftrightarrow H^{+} + {HCO}_{3}^{-}

The ability of bicarbonate $({HCO}_{3}^{-})$ and carbonic acid (H₂CO₃) to act as a buffer system is described mathematically by a relationship known as the Henderson-Hasselbalch equation:

p Η = pKa + {log}_{10} (\frac{{HCO}_{3}^{-}}{H_{2} {CO}_{3}})

where pKa is the dissociation constant for H₂CO₃ and has a constant value of 6.1. In short, the Henderson-Hasselbalch equation states that the pH of a weak acid solution is determined by the ratio of the concentration of base (i.e., bicarbonate, ${HCO}_{3}^{-}$ ) in solution to the concentration of acid (i.e., carbonic acid). The normal pH of arterial blood is 7.4, and the ratio of bicarbonate to carbonic acid is 20 to 1.

Let’s consider an example using the Henderson-Hasselbalch equation to calculate arterial blood pH. Normally the concentration of blood bicarbonate is 24 mEq/l and the concentration of carbonic acid is 1.2 mEq/l. Note that mEq/l is an abbreviation for milli-equivalent per liter, which is a measure of concentration. Therefore, the blood pH can be calculated as follows:

\begin{matrix} pH & = & pKa + {log}_{10} (\frac{2.4}{1.1}) \\ = & 6.1 + {log}_{10} 20 \\ = & 6.1 + 1.3 \\ = & 7.4 \end{matrix}

Although the traditional view of acid-base chemistry has considered the levels of bicarbonate and hydrogen ions as two primary determinants of pH, new ideas were advanced in the early 1980s that exposed the complexity of regulating pH balance in the body. These new concepts were launched by Peter Stewart and are introduced in A Look Back—Important People in Science.

A LOOK BACK—IMPORTANT PEOPLE IN SCIENCE

Peter Stewart Challenged Acid-Base Research by Proposing the Concept of “Strong Ion Difference”

Peter Stewart (1921–1993) was born and raised in Winnipeg, Canada. He earned both his Master’s and Ph.D. degrees from the University of Minnesota. During a long and productive research career, Dr. Stewart served on the faculty of the University of Illinois, Emory University, and Brown University.

In the 1970s, Dr. Stewart developed a strong interest in acid-base regulation and started to mathematically analyze the variables involved in controlling the pH of body fluids. In 1981, he challenged the traditional concepts of acid-base control by publishing his book titled How to Understand Acid-Base Balance—A Quantitative Acid-Base Primer for Biology and Medicine. This book quickly stirred both controversy and debate among researchers in acid-base balance. A brief synopsis of this controversy follows. Historically, it has been believed that the balance between the levels of hydrogen ions and bicarbonate ions determines the pH of body fluids. Dr. Stewart challenged this concept and argued that hydrogen ions and bicarbonate ions are not the independent variables that control acid-base. Instead, he suggested that these variables are dependent variables that are regulated by other factors, including the strong ion difference,^* carbon dioxide levels, and the amount of weak and nonvolatile acids in body fluids. So, does this mean that calculating pH using the Henderson-Hasselbalch equation is of no value? The short answer to this question is no. However, Dr. Stewart’s work demonstrates that the factors that regulate acid-base balance in the body are probably more complex than originally believed.

Although some physiologists have accepted the science supporting Dr. Stewart’s proposal of how acid-base balance is maintained, criticism of his concept of strong ion difference remains. A concern with the strong ion balance concept of acid-base control is the complexity of the chemistry and mathematics behind this model. Moreover, in practice, it is often difficult to measure all the variables required to calculate pH using Dr. Stewart’s strong ion difference method. Therefore, in the foreseeable future, the traditional approach to acid-base balance seems certain to prevail.

IN SUMMARY

The body maintains acid-base homeostasis by buffer-control systems. A buffer resists pH change by removing hydrogen ions when the pH declines and by releasing hydrogen ions when the pH increases.
The principal intracellular buffers are proteins, phosphates, bicarbonates, and histidine-dipeptides (i.e., carnosine).
Muscle fibers possess two important hydrogen ion transporters that carry hydrogen ions from inside the muscle fiber into the interstitial space: (1) sodium/hydrogen exchanger and (2) monocarboxlate transporters.
Primary extracellular buffers include bicarbonates, hemoglobin, and blood proteins.

RESPIRATORY INFLUENCE ON ACID-BASE BALANCE

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The fact that the respiratory system contributes to acid-base balance during exercise was introduced in Chap. 10. However, more details about how the respiratory system contributes to the control of pH will be presented here. Recall that CO₂ is considered a volatile acid because it can be readily changed from CO₂ to carbonic acid (H₂CO₃). Also, remember that it is the partial pressure of CO₂ in the blood that determines the concentration of carbonic acid. For example, according to Henry’s law, the concentration of a gas in solution is directly proportional to its partial pressure. That is, as the partial pressure increases, the concentration of the gas in solution increases, and vice versa.

Because CO₂ is a gas, it can be eliminated by the lungs. Therefore, the respiratory system is an important regulator of blood carbonic acid and pH. To better understand the role of the lungs in acid-base balance, let’s reexamine the carbonic acid dissociation equation:

{CO}_{2} + H_{2} O \leftrightarrow H_{2} {CO}_{3} \leftrightarrow H^{+} + {HCO}_{3}^{-}

This relationship demonstrates that when the amount of CO₂ in the blood increases, the amount of H₂CO₃ increases, which lowers pH by elevating the acid concentration of the blood (i.e., the reaction moves to the right). In contrast, when the CO₂ content of the blood is lowered (i.e., CO₂ is eliminated by the lungs), the pH of the blood increases because less acid is present (reaction moves to the left). Therefore, the respiratory system provides the body with a rapid means of regulating blood pH by controlling the amount of CO₂ present in the blood (17, 27, 44).

IN SUMMARY

Respiratory control of acid-base balance involves the regulation of blood PCO₂.
An increase in blood PCO₂ lowers pH, whereas a decrease in blood PCO₂ increases pH.
Increased pulmonary ventilation can remove CO₂ from the body and thus eliminate hydrogen ions and increase pH.

REGULATION OF ACID-BASE BALANCE VIA THE KIDNEYS

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Because the kidneys do not play an important part in acid-base regulation during short-term exercise, only a brief overview of the kidney’s role in acid-base balance will be presented here. The principal means by which the kidneys regulate hydrogen ion concentration is by increasing or decreasing the bicarbonate concentration (9–11, 23). When the pH decreases (hydrogen ion concentration increases) in body fluids, the kidney responds by reducing the rate of bicarbonate excretion. This results in an increase in the blood bicarbonate concentration, which assists in buffering the increase in hydrogen ions. Conversely, when the pH of body fluids rises (hydrogen ion concentration decreases), the kidneys increase the rate of bicarbonate excretion. Therefore, by changing the amount of buffer present in body fluids, the kidneys aid in the regulation of the hydrogen ion concentration. The kidney mechanism involved in regulating the bicarbonate concentration is located in a portion of the kidney called the tubule, and acts through a series of complicated reactions and active transport across the tubular wall.

Why is the kidney not an important regulator of acid-base balance during exercise? The answer lies in the amount of time required for the kidney to respond to an acid-base disturbance. It takes several hours for the kidneys to react effectively in response to an increase in blood hydrogen ions (10, 46). Therefore, the kidneys respond too slowly to be of major benefit in the regulation of hydrogen ion concentration during exercise.

IN SUMMARY

Although the kidneys play an important role in the long-term regulation of acid-base balance, the kidneys do not play a major role in the regulation of acid-base balance during exercise.

REGULATION OF ACID-BASE BALANCE DURING EXERCISE

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During the final stages of an incremental exercise test or during near-maximal exercise of short duration, there is a decrease in both muscle and blood pH primarily due to the increase in the production of hydrogen ions by the muscle (13, 29). This point is illustrated in Figure 11.5, where the changes in blood and muscle pH during an incremental exercise test are graphed as a function of exercise (i.e., % ${\dot{V} O}_{2} max$ ). Notice that muscle and blood pH follow similar trends during this type of exercise, but that muscle pH is always 0.4 to 0.6 pH unit lower than blood pH. This ocurrs because muscle hydrogen ion concentration is higher than that of blood, and muscle buffering capacity is lower than that of blood.

Figure 11.5

Changes in arterial blood pH and muscle pH during incremental exercise. Notice that arterial and muscle pH begin to fall together at work rates above 50% V˙O2max.

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The amount of hydrogen ions produced during exercise is dependent on (1) the exercise intensity, (2) the amount of muscle mass involved, and (3) the duration of the exercise (13). Exercise involving high-intensity leg work (e.g., running) may reduce arterial pH from 7.4 to a value of 7.0 within a few minutes (13, 27, 35, 43). Further, repeated bouts of this type of exercise can cause blood pH to decline even further to a value of 6.8 (13). This level of blood pH is the lowest ever recorded and would present a life-threatening situation if it were not corrected within a few minutes.

How does the body regulate acid-base balance during exercise? Because working muscles are the primary source of hydrogen ions during exercise, it is logical that the first line of defense against a rise in acid production resides in the individual muscle fibers. It is estimated that intracellular proteins and histidine dipeptides (i.e., carnosine) contribute as much as 60% of the cell’s buffering capacity, with an additional 20% to 30% of the total buffering capacity coming from muscle bicarbonate (25). The final 10% to 20% of muscle buffering capacity comes from intracellular phosphate groups.

Because the muscle’s buffering capacity is limited, extracellular fluid (principally the blood) must possess a means of buffering hydrogen ions as well. Therefore, the blood buffering systems become the second line of defense against exercise-induced acidosis. The principal extracellular buffer is blood bicarbonate (25, 39). Hemoglobin and blood proteins assist in this buffer process, but play only a minor role in blood buffering of hydrogen ions during exercise (25). Figure 11.6 illustrates the role of blood bicarbonate as a buffer during incremental exercise (44). Note that at approximately 50% to 60% of ${\dot{V} O}_{2} max$ , blood levels of lactate begin to increase and blood pH declines due to the rise in hydrogen ions in the blood. This increase in blood hydrogen ion concentration stimulates the carotid bodies, which then signal the respiratory control center to increase alveolar ventilation (i.e., ventilatory threshold; see Chap. 10). An increase in alveolar ventilation results in the reduction of blood PCO₂ and therefore acts to reduce the acid load produced by exercise (17). The overall process of respiratory assistance in buffering lactic acid during exercise is referred to as respiratory compensation for metabolic acidosis.

Figure 11.6

Changes in blood concentrations of bicarbonate, lactate, and pH as a function of work rate.

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In summary, the control of acid-base balance during exercise is important. During high-intensity exercise (i.e., work above the lactate threshold), the contracting skeletal muscle produces significant amounts of hydrogen ions. The first line of defense against exercise-induced acidosis resides in the muscle fiber (i.e., bicarbonate, phosphate, and protein buffers). However, because the muscle fiber’s buffering capacity is limited, additional buffer systems are required to protect the body against exercise-induced acidosis. In this regard, the second line of defense against pH shifts during exercise is the blood buffer systems (i.e., bicarbonate, phosphate, and protein buffers). Importantly, an increase in pulmonary during intense exercise assists in eliminating carbonic acid by “blowing off carbon dioxide.” This respiratory compensation to exercise-induced acidosis plays an important role in the second line of defense against pH change during intense exercise. Together, these first and second lines of defense protect the body against exercise-induced acidosis (Fig. 11.7).

Figure 11.7

The two major lines of defense against pH change during intense exercise.

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IN SUMMARY

Figure 11.7 summarizes the “two-stage” process of buffering exercise-induced production of hydrogen ions.
The first line of defense against hydrogen ions produced in the muscle resides inside the muscle fiber itself. During exercise, muscle buffer systems act rapidly to buffer hydrogen ions and prevent a significant decline in muscle pH.
The second line of defense against exercise-induced acidosis is the blood buffer systems, which include bicarbonate, phosphate, and protein buffers.
Importantly, increased pulmonary ventilation during intense exercise eliminates carbonic acid by “blowing off” carbon dioxide. This process is termed “respiratory compensation for metabolic acidosis” and plays an important role in the second line of defense against exercise-induced acidosis.

STUDY ACTIVITIES

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Define the terms acid, base, buffer, acidosis, alkalosis, and pH.
Graph the pH scale. Label the pH values that represent normal arterial and intracellular pH.
List and briefly discuss the major sources of hydrogen ions produced in the muscle during exercise.
Why is the maintenance of acid-base homeostasis important to physical performance?
Discuss both the intracellular and extracellular buffer mechanisms that protect against exercise-induced acidosis. What are the principal buffer systems and hydrogen ion transporters found in each of these lines of defense?
Discuss respiratory compensation for metabolic acidosis. What would happen to blood pH if an individual began to hyperventilate at rest? Why?
Briefly outline how the body resists pH change during exercise. Include in your outline both the cellular and blood buffer systems.

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Chapter 11: Acid-Base Balance during Exercise

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OBJECTIVES

OUTLINE

KEY TERMS

INTRODUCTION

ACIDS, BASES, AND pH

Figure 11.1

Figure 11.2

HYDROGEN ION PRODUCTION DURING EXERCISE

Figure 11.3

IMPORTANCE OF ACID-BASE REGULATION DURING EXERCISE

ACID-BASE BUFFER SYSTEMS

Intracellular Buffers

Figure 11.4

Influence of Muscle Fiber Type and Exercise Training on Intracellular Buffer Capacity

Extracellular Buffers

RESPIRATORY INFLUENCE ON ACID-BASE BALANCE

REGULATION OF ACID-BASE BALANCE VIA THE KIDNEYS

REGULATION OF ACID-BASE BALANCE DURING EXERCISE

Figure 11.5

Figure 11.6

Figure 11.7

STUDY ACTIVITIES

SUGGESTED READINGS

REFERENCES

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