Opening the book…
\(\mathrm{pH}\) is a measure of the acidity or alkalinity of a solution. It is defined as the negative logarithm to the base 10 of the hydrogen ion concentration in a given aqueous solution.
\[ \begin{align*}
\mathrm{pH} &= \log 1/[\mathrm{H}^+]\\
&= -\log [\mathrm{H}^+]\\
&= -\log [0.00000004]\\
&= 7.4
\end{align*} \nonumber \]
Hydrogen ions in the body have several sources. These include dietary intake through food, digestive secretions, and various metabolic activities. Metabolism generates both volatile and nonvolatile or fixed acids. Volatile acid, primarily carbonic acid, contributes significantly to the acid load, producing over 10,000 mEq per day, which is expelled from the body through the exhalation of carbon dioxide. In contrast, nonvolatile acids, amounting to about 100 mEq per day, are managed by the kidneys, which help to maintain acid-base balance by excreting these acids. One of the physiological factors rigorously regulated within the body is the maintenance of hydrogen ion concentration. This parameter significantly impacts the three-dimensional structure of tissue proteins, thereby influencing their functionalities. Enzymes, in particular, exhibit extreme sensitivity to \(\mathrm{pH}\) variations, performing optimally within a specific \(\mathrm{pH}\) range. Any deviation from this optimal \(\mathrm{pH}\) drastically diminishes their activity.
The \(\mathrm{pH}\) scale, representing the power of hydrogen ions, operates logarithmically. In the dissociation of water, hydrogen and hydroxyl ions form with concentrations of 10-7 moles per liter at 20°C. Their product, the ionic product of water, remains constant at 10-14. To conveniently express the vast range of laboratory values without negative figures, the negative logarithm is employed, denoted by \(\mathrm{pH} = -\log [\mathrm{H}^+]\). Hence, neutral \(\mathrm{pH}\) is set at 7, equivalent to 0.0000001 mol per liter. A tenfold increase in hydrogen ion concentration decreases the \(\mathrm{pH}\) by one unit (\(\mathrm{pH}\)=6), while a tenfold decrease elevates the \(\mathrm{pH}\) by one unit (\(\mathrm{pH}\)=8). Arterial blood typically maintains a normal \(\mathrm{pH}\) of 7.4 (40 nmol per liter), with a healthy range of arterial plasma \(\mathrm{pH}\) spanning from 7.38 to 7.42. The survivable boundaries for \(\mathrm{pH}\) are between 7.7 (maximum) and 6.9 (minimum), within which life can be sustained. Intracellular \(\mathrm{pH}\) differs from extracellular \(\mathrm{pH}\), usually ranging from 6 to 7.4. \(\mathrm{pH}\) of urine is 4.5 to 8.
Despite its significance, the hydrogen ion concentration in plasma is minimal compared to other ions. For instance, while the concentration of sodium ion is 140 mmol per liter, the hydrogen ion concentration remains notably low at 40 nmol per liter. Due to this million-fold discrepancy in concentration, the osmotic influence of hydrogen ion is typically omitted in blood osmolality calculations.25
Acids, Bases, and Buffers
Universally, an acid is a provider of protons or hydrogen ions, whereas a base is recognized as a receptor of protons. Strong acids like hydrochloric acid rapidly dissociate into hydrogen and chloride ions, whereas weak acids like carbonic acid exhibit less complete dissociation into hydrogen and bicarbonate ions. Various buffering systems within the body enable blood and bodily fluids to uphold a tightly regulated \(\mathrm{pH}\) range even when faced with disturbances. At a \(\mathrm{pH}\) of 6.1, carbonic acid and bicarbonate ions exist in solutions at nearly equivalent concentrations. In a normal state at \(\mathrm{pH}\) 7.4, there are 20 bicarbonate ions for every molecule of carbonic acid.
Acidosis refers to a condition where the blood becomes acidic, while alkalosis is when the blood is alkaline. In clinical terms, acidemia is when arterial blood \(\mathrm{pH}\) falls below 7.35, and alkalemia is when it rises above 7.45.
A buffer acts as a chemical system that helps mitigate sudden shifts in fluid \(\mathrm{pH}\) by moderating the changes in hydrogen ion concentrations in cases of excessive acid or base presence. Typically, the substance responsible for absorbing ions is a weak acid, which absorbs hydroxyl ion, or a weak base, which absorbs hydrogen ions. Despite their effectiveness, buffer systems have their limitations. When faced with excessive acid, a buffer may become overwhelmed, resulting in a substantial rise in hydrogen ion concentration due to its exhaustion.
The body maintains \(\mathrm{pH}\) balance through several mechanisms. Chemical buffer systems in body fluids act immediately but temporarily, trapping hydrogen ions without eliminating or creating them. Examples of these buffers include bicarbonate, ammonia, proteins, and phosphate. The respiratory center provides a second line of defense by eliminating carbon dioxide within seconds to minutes. An increase in hydrogen ion concentration stimulates increased ventilation, leading to carbon dioxide loss, and vice versa. The kidneys, a more powerful regulator, manage non-volatile acids by secreting hydrogen ions, reabsorbing bicarbonate, and generating new bicarbonate over hours to days.
In bodily fluids, buffer systems serve as the primary line of defense against alterations in hydrogen ion concentration. These systems are remarkably efficient, functioning at varying rates. The chemical buffers in the blood can swiftly compensate for any changes in \(\mathrm{pH}\) within seconds. Chemical buffer systems in blood plasma mainly consist of bicarbonate ions and plasma proteins, while buffers within cells predominantly involve hemoglobin and phosphate. Specifically, bicarbonate ions account for 53 percent, hemoglobin 35 percent, and the remaining 12 percent is distributed between phosphates and plasma proteins.26
Bicarbonate-Carbonic Acid Buffer System
The bicarbonate-carbonic acid buffer system serves as the primary buffer in interstitial fluid. In this system, carbon dioxide produced by cells combines with water to form carbonic acid, which then dissociates into hydrogen and bicarbonate ions. This crucial reaction is facilitated by the enzyme carbonic anhydrase, which is present in locations such as kidney tubules, red blood cells, endothelial cells, and parietal cells.
\[\mathrm{H}_2 \mathrm{O}+\mathrm{CO}_2 \longleftrightarrow \mathrm{H}_2 \mathrm{CO}_3 \longleftrightarrow \mathrm{H}^{+}+\mathrm{HCO}_3^{-} \nonumber \]
The Henderson-Hasselbalch equation explains the relationship between the \(\mathrm{pH}\) of a solution containing a weak acid and its conjugate base, and the acid dissociation constant of that acid (PKa). In the bicarbonate buffer system, \(\mathrm{pH}\) depends on the balance between carbonic acid and bicarbonate in the blood. At normal body temperature, carbon dioxide dissolves in blood with a constant of 0.03, allowing carbonic acid and dissolved carbon dioxide to stay balanced. With a usual PCO2 of 40 mm Hg and bicarbonate level of 24 mmol/L, the blood maintains a 1:20 ratio of carbonic acid to bicarbonate, helping to buffer against increases in acidity (Figure 4).27
The Henderson-Hasselbalch equation is written as:
\[\mathrm{pH}=\mathrm{pK}_a+\log \left(\frac{\left[\mathrm{A}^{-}\right]}{[\mathrm{HA}]}\right) \nonumber \]
In the case of the bicarbonate buffer system:
\[\mathrm{pH}=6.1+\log \left(\frac{\left[\mathrm{HCO}_3^{-}\right]}{0.03 \times \mathrm{PCO}_2}\right) \nonumber \]
where:
- \(\mathrm{pH}\) = the acidity or alkalinity of the solution,
- pKₐ (6.1 for bicarbonate) is the dissociation constant of carbonic acid,
- [\(\ce{HCO3-}\)] = bicarbonate ion concentration,
- \(\ce{PCO2}\) = partial pressure of carbon dioxide in mm Hg,
- 0.03 = solubility coefficient of \(\ce{CO2}\) at body temperature.
The bicarbonate buffer system primarily buffers against metabolic wastes, such as lactic acid and ketones. Control over carbonic acid levels in the blood is regulated by carbon dioxide exhalation through the lungs. Inside RBCs, carbonic anhydrase prompts acid dissociation, decreasing blood acidity and leading to carbon dioxide exhalation. Typical reference points for assessing the acid-base balance in the blood and can be used to evaluate a person's respiratory and metabolic status. [\(\ce{H+}\)] = 40 nmol/L for a \(\mathrm{pH}\) of 7.4 (range: 35 – 45 nmol/L); \(\ce{PaCO2}\) = 40 mm Hg; [\(\ce{HCO3-}\)] = 24 mEq/L (range: 22-26 mEq/L).5
The kidneys regulate blood bicarbonate by reabsorbing it from the renal filtrate back into the bloodstream. About one-third of the buffering reaction also occurs within red blood cells, making this system effective beyond just the plasma.
When sodium bicarbonate (\(\ce{NaHCO3}\)) encounters strong acids or bases, it reacts in a reversible way, forming either carbonic acid (a weak acid) and salt or producing sodium bicarbonate and water.28
\[\begin{gathered}
\mathrm{NaHCO}_3+\mathrm{HCl} \rightarrow \mathrm{H}_2 \mathrm{CO}_3+\mathrm{NaCl} \\
\text { (Sodium bicarbonate) }+(\text { Strong acid }) \rightarrow(\text { Weak acid })+(\text { Salt }) \\
\mathrm{H}_2 \mathrm{CO}_3+\mathrm{NaOH} \rightarrow \mathrm{NaHCO}_3+\mathrm{H}_2 \mathrm{O} \\
\text { (Weak acid) }+(\text { Strong base }) \rightarrow(\text { Sodium bicarbonate })+(\text { Water })
\end{gathered} \nonumber\]
Figure 4 | Relationship between Carbon Dioxide (\(\ce{CO2}\)) and Blood \(\mathrm{pH}\) in Respiratory Acidosis and Alkalosis |29 This figure illustrates the effects of carbon dioxide (\(\ce{CO2}\)) levels on blood \(\mathrm{pH}\) through the mechanisms of respiratory acidosis and alkalosis. In respiratory acidosis (left side of the diagram), \(\ce{CO2}\) levels increase due to hypoventilation or other factors that impair gas exchange, leading to an increase in hydrogen ion (\(\ce{H+}\)) concentration and a decrease in blood \(\mathrm{pH}\), resulting in a more acidic environment. In respiratory alkalosis (right side of the diagram), \(\ce{CO2}\) levels decrease due to hyperventilation, which reduces \(\ce{H+}\) concentration and raises blood \(\mathrm{pH}\), creating a more alkaline environment. The balance of \(\ce{CO2}\) is a primary factor influencing blood \(\mathrm{pH}\) in respiratory conditions, with homeostatic mechanisms working to maintain optimal \(\mathrm{pH}\) levels for cellular function.
Protein Buffer System
Proteins exhibit remarkable buffering properties within the body due to their complex amino acid structure. The primary mechanism by which proteins buffer changes in \(\mathrm{pH}\) involves their amino and carboxyl groups, as well as side chains of certain amino acids. Amino groups (\(\ce{–NH2}\)) can accept \(\ce{H+}\) ions to form \(\ce{–NH3+}\), while carboxyl groups (\(\ce{–COOH}\)) can donate \(\ce{H+}\) ions to form \(\ce{–COO-}\). When the \(\mathrm{pH}\) in the surrounding environment becomes too acidic, the protein’s amino groups act as bases, binding with excess \(\ce{H+}\) ions to help reduce acidity. Conversely, when the environment becomes too basic, carboxyl groups can release \(\ce{H+}\) ions to increase the concentration of \(\ce{H+}\), thus lowering \(\mathrm{pH}\). Proteins serve as two-thirds of the buffering power in blood and are the primary source of cellular buffering.
Hemoglobin is a key buffering agent within red blood cells. When blood is slightly alkaline, proteins, especially the histidine residues in hemoglobin, act as weak acids. This buffering effect is stronger in deoxygenated blood compared to oxygenated blood due to its higher affinity for hydrogen ions, the effects of increased carbon dioxide concentration and the resulting \(\mathrm{pH}\) changes in tissues, as well as structural changes that occur during oxygen binding and release. Hemoglobin is much more concentrated than plasma proteins and has a greater ability to act as a buffer (about three times more). Therefore, it plays a big role in helping to maintain the \(\mathrm{pH}\) of blood. Oxyhemoglobin acts as a stronger acid than deoxyhemoglobin as it can release more hydrogen ions.5
\[\mathrm{H}^{+}+\mathrm{Hb} \longleftrightarrow \mathrm{HHb} \nonumber \]
Phosphate Buffer System
Phosphate, another significant buffer system, presents a distinct advantage due to its more balanced nature at a normal blood \(\mathrm{pH}\). Yet, despite its theoretically higher effectiveness than the bicarbonate buffer system, phosphate contributes only a small fraction, approximately 1 to 3 percent, to the blood's buffering capacity. Phosphates exist in two forms in blood: sodium dihydrogen phosphate and sodium monohydrogen phosphate. These compounds can transform into weaker acids or bases when reacting with strong acids or bases. The phosphate buffer system becomes far more efficient intracellularly, where phosphate concentrations are notably higher, especially in bones. These intracellular concentrations are crucial in buffering the acids originating from metabolic processes. In the intracellular compartment, proteins rich in histidine groups and compounds like adenosine phosphate and glucose-1-phosphate effectively buffer excess hydrogen ions. Any surplus hydrogen ions then diffuse into the extracellular spaces.
\[\mathrm{HPO}_4^{-}+\mathrm{H}^{+} \longleftrightarrow \mathrm{H}_2 \mathrm{PO}_4^{-}\nonumber \]
The extracellular fluid interacts with blood plasma and has a similar ability to maintain \(\mathrm{pH}\) balance. This includes using bicarbonate and phosphate systems to help buffer acids and bases. Intracellular \(\mathrm{pH}\) remains closely regulated within the narrow physiological range, whereas during acute disturbances or chronic acid-base alterations, intracellular buffering efficiency may fluctuate. Nonetheless, intracellular buffering becomes more effective over time, particularly in chronic acid-base disturbances.5
Regulation of Acid-Base Balance
The body deploys three primary lines of defense to achieve this precise control: buffer systems, which are immediate, first-line responders comprising weak acids and their conjugate bases that rapidly bind to or release hydrogen ions; the respiratory system, acting as a rapid, second-line defense by controlling the exhalation of carbon dioxide to manage carbonic acid levels; and the renal system, serving as the long-term, most powerful, and definitive regulator, primarily managing blood bicarbonate levels and hydrogen ion excretion through reabsorption, new bicarbonate generation, and direct hydrogen ion removal. These three systems work in a highly integrated and compensatory fashion, ensuring that if a metabolic disturbance pushes the \(\mathrm{pH}\) in one direction, the respiratory system will attempt rapid compensation, followed by the kidneys providing a more sustained and powerful metabolic adjustment, thereby maintaining overall \(\mathrm{pH}\) homeostasis.
Respiratory Regulation of Acid-Base Balance
The respiratory system plays a crucial role in maintaining the body's acid-base balance by managing the levels of carbon dioxide and carbonic acid in the blood. When carbon dioxide levels rise in the blood, such as during breath-holding, excess carbon dioxide reacts with water to form additional carbonic acid, leading to a decrease in blood \(\mathrm{pH}\). By increasing the rate or depth of respiration, the body can expel excess carbon dioxide, thereby lowering the levels of carbonic acid and readjusting the blood \(\mathrm{pH}\) towards normal. Conversely, excessive rapid breathing, like hyperventilation, causes the body to eliminate carbon dioxide excessively, leading to a reduction in carbonic acid levels and an overly alkaline state in the blood (Figure 4). This transient alkalosis can be corrected by rebreathing exhaled air contained in a paper bag, which helps rapidly restore blood \(\mathrm{pH}\) to normal levels. These chemical reactions occur in the pulmonary capillaries of the lungs when blood is oxygenated. Minor adjustments in breathing, altering the amount of exhaled carbon dioxide, typically suffice to regulate blood \(\mathrm{pH}\). For instance, a brief doubling of the respiratory rate for less than a minute results in a 0.2 increase in blood \(\mathrm{pH}\). This adjustment often occurs during strenuous exercise, where increased energy production leads to excess carbon dioxide and lactic acid production. The body, therefore, elevates respiration to expel the surplus carbon dioxide and maintain acid-base balance to prevent acidosis (Figure 5).
Figure 5 | Relationship between \(\mathrm{pH}\) and Alveolar Ventilation |5 This graph illustrates the relationship between arterial blood \(\mathrm{pH}\) and the rate of alveolar ventilation, with normal ventilation represented as 1. As blood \(\mathrm{pH}\) decreases (becomes more acidic), alveolar ventilation increases significantly. Conversely, as blood \(\mathrm{pH}\) increases (becomes more alkaline), alveolar ventilation decreases. This demonstrates the body's homeostatic mechanism to regulate \(\mathrm{pH}\) through adjustments in breathing rate.
The respiratory rate is regulated by chemoreceptors, primarily responding to changes in carbon dioxide levels. Peripheral sensors in major arteries, such as the aorta and carotids, alert the brain to prompt immediate adjustments in respiratory rate upon carbon dioxide level fluctuations (Figure 6). Furthermore, the brain houses sensors that detect \(\mathrm{pH}\) changes in the CSF, affecting the respiratory center in the medulla oblongata, thereby modulating breathing rates to maintain \(\mathrm{pH}\) within normal ranges.
Figure 6 | Respiratory Regulation of Blood \(\mathrm{pH}\) | 30 This diagram illustrates the body's respiratory response to disturbances in acid-base homeostasis. When acid-base balance is disturbed, leading to either acidosis (decreased \(\mathrm{pH}\)) or alkalosis (increased \(\mathrm{pH}\)), chemoreceptors in the brain and arterial system are stimulated. In acidosis, this stimulation results in an increased respiration rate, leading to a decrease in blood carbon dioxide (\(\ce{CO2}\)) levels and a subsequent decrease in hydrogen ion (\(\ce{H+}\)) concentration, ultimately raising the \(\mathrm{pH}\) back towards normal. In alkalosis, the opposite occurs; the respiration rate decreases, leading to an increase in blood \(\ce{CO2}\) levels and a subsequent increase in H⁺ concentration, lowering the \(\mathrm{pH}\) back towards normal. Both pathways lead to restore acid-base homeostasis.
Abnormal elevation of blood carbon dioxide, known as hypercapnia, occurs due to conditions that impair respiratory functions, like pneumonia or congestive heart failure. Conversely, abnormally low carbon dioxide levels, called hypocapnia, are witnessed in situations of hyperventilation, whether due to factors like salicylate toxicity, elevated room temperatures, fever, or even hysteria.
Renal Regulation of Acid-Base Balance
While the respiratory system controls carbonic acid levels by managing carbon dioxide exhalation, the renal system oversees blood bicarbonate levels. Various factors can lead to a decrease in blood bicarbonate levels, including inhibition of carbonic anhydrase by certain diuretics, excessive bicarbonate ion loss from diarrhea, conditions like Addison’s disease, and renal damage, resulting in reduced aldosterone levels.
The kidneys play a primary metabolic role in maintaining the body's acid-base balance by adjusting urine \(\mathrm{pH}\), either acidifying or alkalizing it, and crucially managing blood bicarbonate levels. Since renal tubule cells are impermeable to bicarbonate, a specialized mechanism, largely occurring in the proximal tubular cells and aided by carbonic anhydrase, is employed both to reabsorb existing bicarbonate and to generate new bicarbonate. As Figure 7 illustrates, \(\ce{CO2}\) diffuses from the renal interstitial fluid into the tubular cell, where carbonic anhydrase catalyzes its reaction with water to form carbonic acid, which then dissociates into bicarbonate and hydrogen ions. To sustain the blood's bicarbonate buffering system, bicarbonate ions are transported out of the cell into the renal interstitial fluid and returned to the bloodstream, while hydrogen ions are simultaneously secreted into the tubular lumen, often via an antiport mechanism in the apical membranes that exchanges them for sodium ions. To maintain electrochemical neutrality during bicarbonate reabsorption, chloride ions are also exchanged across the cell membrane in opposite directions—moving from the renal interstitial fluid into the tubular cell and from the tubular lumen into the cell—a process known as the "chloride shift." All these transport processes are powered by ATP, ensuring that the kidneys continuously process bicarbonate and hydrogen ions to maintain effective \(\mathrm{pH}\) balance.5
Figure 7 | Production of Bicarbonate in the Proximal Tubule | This diagram illustrates the exchange of chloride (\(\ce{Cl-}\)) and bicarbonate (\(\ce{HCO3-}\)) ions across the renal tubule cell membrane, a key process in bicarbonate reabsorption and acid-base regulation. Carbon dioxide (\(\ce{CO2}\)) diffuses from the renal interstitial fluid into the tubular cell. Inside the cell, carbonic anhydrase catalyzes the reaction between \(\ce{CO2}\) and water (\(\ce{H2O}\)) to form carbonic acid (\(\ce{H2CO3}\)). \(\ce{H2CO3}\) then dissociates into bicarbonate (\(\ce{HCO3-}\)) and hydrogen ions (\(\ce{H+}\)). Bicarbonate is transported out of the cell into the renal interstitial fluid, while \(\ce{H+}\) are secreted into the tubular lumen. To maintain electrochemical neutrality, \(\ce{Cl-}\) are exchanged across the cell membrane in opposite directions: \(\ce{Cl-}\) moves from the renal interstitial fluid into the tubular cell and from the tubular lumen into the cell. This "chloride shift" ensures the proper balance of ions during bicarbonate reabsorption. ATP is utilized to power the transport processes involved.
Figure 8 provides a quantitative overview of this bicarbonate reabsorption process along the nephron. Specifically, it illustrates that approximately 4320 mEq/day of bicarbonate enters the proximal tubule. A significant 85% (3672 mEq/day) of this bicarbonate is reabsorbed in the proximal tubule, highlighting its critical role in bicarbonate conservation. A further 10% (432 mEq/day) is reabsorbed in the loop of Henle. The distal tubule and collecting duct contribute to the fine-tuning of bicarbonate levels, with a variable amount, indicated as greater than 4.9% (215 mEq/day), being reabsorbed. Only a minimal amount, approximately 1 mEq/day, is excreted in the final urine.
Figure 8 | Reabsorption of Bicarbonate at the Renal Tubule |5 This diagram illustrates the reabsorption of bicarbonate (\(\ce{HCO3-}\)) along different segments of the nephron. Starting with 4320 mEq/day of bicarbonate entering the proximal tubule, approximately 85% (3672 mEq/day) is reabsorbed in this section. A further 10% (432 mEq/day) is reabsorbed in the thick limb of the loop of Henle. In the distal tubule and collecting duct, a variable amount, indicated as greater than 4.9% (215 mEq/day), is reabsorbed, demonstrating the fine-tuning of bicarbonate balance. A minimal amount, approximately 1 mEq/day, is excreted in the final urine. The arrows indicate the direction of bicarbonate movement, with red arrows signifying reabsorption from the tubular fluid back into the bloodstream.
The ammonia buffer system facilitates the excretion of excess hydrogen ions by the kidneys and the generation of new bicarbonate. The generation of new bicarbonate ions by the ammonia buffer system is a key renal process that helps the body maintain acid-base balance, particularly in conditions of acidosis when there is an excess of hydrogen ions. This mechanism begins in the proximal tubule, distal tubule, and the thick segment of the ascending limb of the loop of Henle tubular cells, where the amino acid glutamine is broken down into ammonia (\(\ce{NH3}\)) and bicarbonate. The ammonia is then secreted into the tubular fluid. In this environment, \(\ce{NH3}\) binds with hydrogen ions to form ammonium (\(\ce{NH4+}\)), which is subsequently excreted in urine. This reaction not only removes hydrogen ions but also helps generate new bicarbonate ions. Buffering hydrogen ion within the tubule helps maintain acidity above 4.5 in order to protect it from being exposed to very low \(\mathrm{pH}\). For each H⁺ ion neutralized and excreted as \(\ce{NH4+}\), a new \(\ce{HCO3-}\) ion is produced and returned to the bloodstream. Importantly, this new bicarbonate is freshly generated rather than coming from any existing bicarbonate pool, thus enhancing the blood’s buffering capacity. During acidosis, the kidneys ramp up glutamine breakdown, increasing \(\ce{NH3}\) production and allowing more \(\ce{HCO3-}\) generation to counteract the low \(\mathrm{pH}\). Approximately 80-90 percent of the newly generated bicarbonate is reabsorbed into the peritubular blood.
Hydrogen ion excretion occurs mainly in the form of titratable acid such as phosphate or ammonium ion, equating to the amount of bicarbonate reabsorbed into the circulation. By continuously generating new bicarbonate ions, the ammonia buffer system serves as a powerful mechanism for long-term \(\mathrm{pH}\) stabilization and acid-base regulation.
Phosphate ions act as a buffer, combining with the excess hydrogen ions to form dihydrogen phosphate ions, \(\mathrm{H}_2 \mathrm{PO}_4^{-}\) (Figure 9). Monohydrogen phosphate is the stronger base compared to dihydrogen phosphate, which is a weak acid. This difference in strength is due to the fact that monohydrogen phosphate can accept a proton (acting as a base), while dihydrogen phosphate can donate a proton (acting as an acid) in solution.
\[\mathrm{H}^{+}+\mathrm{HPO}_4^{-2} \rightleftharpoons \mathrm{H}_2 \mathrm{PO}_4^{-} \nonumber \]
Figure 9 | Phosphate Buffering in the Renal Tubule | This diagram illustrates the phosphate buffering system within the renal tubular cells, contributing to acid-base balance. Carbon dioxide (\(\ce{CO2}\)) diffuses from the renal interstitial fluid into the tubular cells. Within the cell, \(\ce{CO2}\) combines with water (\(\ce{H2O}\)) to form carbonic acid (\(\ce{H2CO3}\)) through the action of carbonic anhydrase. Carbonic acid dissociates into bicarbonate (\(\ce{HCO3-}\)) and hydrogen ions (\(\ce{H+}\)). Bicarbonate is transported out of the cell into the renal interstitial fluid, while hydrogen ions are secreted into the tubular lumen via a sodium-hydrogen antiport mechanism. In the lumen, hydrogen ions combine with monohydrogen phosphate (\(\ce{NaHPO4-}\)) to form dihydrogen phosphate (\(\ce{NaH2PO4}\)), which is then excreted in the urine. This process effectively buffers excess hydrogen ions in the tubular fluid and contributes to the regulation of blood \(\mathrm{pH}\). Additionally, sodium ions are actively transported out of the tubular cells into the renal interstitial fluid using ATP, maintaining the electrochemical gradient necessary for these transport processes.
The dihydrogen phosphate reacts with bicarbonate, resulting in the formation of phosphate ions and water.
\[\mathrm{H}_2 \mathrm{PO}_4^{-}+\mathrm{HCO}_3^{-} \rightleftharpoons \mathrm{HPO}_4^{-2}+\mathrm{H}_2 \mathrm{O} \nonumber\]
Extracellular volume significantly influences bicarbonate reabsorption. In situations like congestive heart failure, expanded extracellular space inhibits sodium and bicarbonate ion reabsorption. Conversely, decreased extracellular volume increases sodium ion concentration and bicarbonate reabsorption. Furthermore, potassium ions also influence bicarbonate reabsorption, affecting it similarly to sodium ions.
When chloride is lost, the body relies on bicarbonate to balance the positive ion charges and maintain neutrality. The typical human diet leans acidic due to nonvolatile acids like phosphoric and sulfuric acids that result from metabolic processes. To regulate these acid levels, the kidneys excrete an amount of acid equivalent to the dietary acid load, although this amount is much smaller than the pulmonary acid load from carbon dioxide produced by tissues.5
Disorders of Acid-Base Balance
The human body maintains arterial blood \(\mathrm{pH}\) within an incredibly narrow range of 7.35 to 7.45. Deviation outside this range results in either acidosis or alkalosis, both of which can have severe health implications. Acidosis, identified by a blood \(\mathrm{pH}\) below 7.35, signifies a critical condition. A continuous decrease below 7.0 can be fatal. Symptoms of acidosis include headaches, confusion, lethargy, and persistent fatigue. Conversely, alkalosis, indicated by a blood \(\mathrm{pH}\) above 7.45, is equally serious, with a \(\mathrm{pH}\) above 7.8 considered fatal. Symptoms include cognitive impairment, which can progress to unconsciousness, tingling or numbness in the extremities, muscle twitching, and even nausea and vomiting (Figure 5).
Both acidosis and alkalosis can arise from either metabolic or respiratory disorders, complicating the body’s delicate acid-base balance. The concentration of carbonic acid in the blood, closely linked to carbon dioxide levels and exhalation from the lungs, underscores the vital respiratory contribution to maintaining this balance. Understanding the significance of carbon dioxide - for every molecule of exhaled carbon dioxide, a molecule of carbonic acid is lost, while retention of carbon dioxide forms a molecule of carbonic acid - illustrates the critical role of respiration in managing the body's acid-base equilibrium.
Figure 10 | Symptoms of Acidosis and Alkalosis | 31 This illustration depicts the diverse symptoms associated with acidosis (left) and alkalosis (right) across various body systems.
Metabolic Acidosis: Primary Bicarbonate Deficiency
Metabolic acidosis manifests when the blood becomes overly acidic, exhibiting a \(\mathrm{pH}\) below the normal range of 7.35. This reduction in \(\mathrm{pH}\) often stems from a decrease in bicarbonate, altering the typical ratio of bicarbonate to the carbonic acid buffer from 20:1 at the standard \(\mathrm{pH}\) of 7.40. Various factors lead to this condition, with the most common cause attributed to the presence of organic acids or an excess of ketones in the blood.
In addition to the predominant causes, several other conditions can induce metabolic acidosis. Intense physical exertion may provoke temporary metabolic acidosis due to the generation of lactic acid. Certain substances, when ingested in large quantities, can also contribute to this acid-base imbalance. For instance, an overdose of aspirin results in acidosis due to the acidity of its metabolite, salicylate. Uremia, characterized by the retention of urea and uric acid, is another condition that can lead to metabolic acidosis. Moreover, diabetic ketoacidosis, characterized by an excessive presence of ketones in the blood, is a significant contributor to metabolic acidosis. Other factors that can lead to this condition include a decrease in hydrogen ion excretion, hindering bicarbonate conservation, and excessive loss of bicarbonate ions due to diarrhea. Table 7 provides a comprehensive list of the multiple causes of metabolic acidosis, encompassing both medical conditions and instances induced by certain physiological activities or substance ingestion.
| Cause | Metabolic |
|---|---|
| Diarrhea | Bicarbonate |
| Uremia | Phosphoric, Sulfuric, and Lactic Acids |
| Diabetic Ketoacidosis | Increased Ketones |
| Strenuous Exercise | Lactic Acid |
| Salicylate/aspirin | Sulfasalicylic Acid (SSA)* |
* acid metabolites from an ingested chemical.
Metabolic Alkalosis: Primary Bicarbonate Excess
Metabolic alkalosis presents the converse of metabolic acidosis, arising when the blood becomes excessively alkaline, with a \(\mathrm{pH}\) above 7.45, often due to an overabundance of bicarbonate (referred to as primary bicarbonate excess).
An increase in bicarbonate ions in the blood can occur temporarily following the ingestion of excessive quantities of bicarbonate, citrate, or antacids, frequently taken for conditions like stomach acid reflux. Chronic metabolic alkalosis may result from Cushing's disease, a condition marked by the chronic hypersecretion of ACTH from the anterior pituitary. The excessive ACTH secretion elevates aldosterone levels, leading to heightened urinary potassium excretion which in turn trigger increased secretion of hydrogen ions. Other factors that contribute to metabolic alkalosis include the loss of hydrochloric acid from the stomach due to vomiting, depletion of potassium ions due to the use of diuretics for hypertension, and the excessive use of laxatives. Laxative overuse disrupts the electrolyte and fluid balance, leading to chloride, potassium, and sodium losses that drive the body toward metabolic alkalosis by increasing bicarbonate retention and hydrogen ion excretion.These factors contribute to an increased bicarbonate level in the blood, resulting in the \(\mathrm{pH}\) shifting to an alkaline state.
Respiratory Disorders
Respiratory Acidosis: Primary Carbonic Acid/Carbon Dioxide Excess
Respiratory acidosis manifests when the blood becomes excessively acidic due to heightened levels of carbonic acid, which results from an excess of carbon dioxide in the bloodstream. This condition can be prompted by various factors that disrupt normal breathing patterns and carbon dioxide elimination mechanisms. Underlying causes may include respiratory ailments like pneumonia, emphysema, or congestive heart failure, all of which impede normal respiratory processes. These disorders can negatively impact the lungs' ability to exchange gasses efficiently, leading to the accumulation of carbon dioxide and the subsequent elevation of carbonic acid levels in the blood, ultimately resulting in respiratory acidosis.
Respiratory Alkalosis: Primary Carbonic Acid/Carbon Dioxide Deficiency
Respiratory alkalosis arises when the blood becomes excessively alkaline due to reduced carbonic acid and carbon dioxide levels. This condition typically occurs when an excessive amount of carbon dioxide is exhaled from the lungs, commonly observed during hyperventilation. The escalated respiratory rate inducing hyperventilation can stem from various triggers such as intense emotional distress, fear, fever, infections, oxygen deprivation (hypoxia), or abnormally high levels of stress hormones like epinephrine and norepinephrine. Unexpectedly, an overdose of aspirin or salicylate toxicity can lead to respiratory alkalosis, as the body endeavors to counterbalance an initial acidotic state.32
Compensatory Mechanisms
Several compensatory mechanisms are in place to finely regulate blood \(\mathrm{pH}\), employing buffers, respiratory adjustments, and renal functions. While these mechanisms typically function effectively, their effectiveness may be hindered if one of these systems malfunctions, such as in the case of kidney failure or respiratory disorders. When the \(\mathrm{pH}\) and the ratio of bicarbonate to carbonic acid are altered significantly, the body's ability to compensate may be limited. Moreover, substantial deviations in \(\mathrm{pH}\) levels have the potential to denature proteins. Extensive protein damage can lead to the disturbance of regular metabolic processes, severe tissue impairments, and potentially fatal outcomes.
Respiratory Compensation
In response to metabolic acidosis, the respiratory system ramps up the breathing rate to eliminate excess carbon dioxide and restore the bicarbonate to carbonic acid ratio to the optimal 20:1 level. This adjustment can occur rapidly, often within minutes. Conversely, when dealing with metabolic alkalosis, the respiratory system's ability to compensate is less efficient than in the case of acidosis. Typically, when faced with elevated \(\mathrm{pH}\) levels, the respiratory system seeks to elevate carbon dioxide levels in the blood by reducing the respiratory rate to retain carbon dioxide. However, there is a limit as to how much the body can tolerate reduced respiration. As a result, the respiratory mechanism is less effective at compensating for metabolic alkalosis compared to its response to acidosis.
Metabolic Compensation
In response to respiratory diseases causing acidosis, the body employs metabolic and renal compensation strategies centered on the preservation of bicarbonate. In instances of respiratory acidosis, the kidneys augment the conservation of bicarbonate and the excretion of hydrogen ions through the exchange mechanism previously described. These actions work to elevate the concentration of bicarbonate in the blood, thereby restoring the appropriate balance between bicarbonate and carbonic acid. Conversely, in situations of respiratory alkalosis, the kidneys reduce the generation of bicarbonate and reabsorb hydrogen ions from the tubular fluid. However, these mechanisms may face limitations due to the exchange of potassium ions by renal cells, which employ a potassium-hydrogen ion exchange mechanism (antiporter).33
Diagnosing Acidosis and Alkalosis
Laboratory tests analyzing \(\mathrm{pH}\), carbon dioxide, PCO2, and bicarbonate are critical in identifying acidosis and alkalosis, shedding light on whether the imbalance is of respiratory or metabolic origin and assessing the efficiency of compensatory mechanisms. These metrics help gauge the condition of the blood—whether it is in an acidotic state, within the normal range, or exhibiting signs of alkalosis. Moreover, the values of PCO2 and total bicarbonate play a significant role in distinguishing between metabolic or respiratory imbalances, providing insights into the body's compensatory efforts. Table 8 enumerates the distinct conditions and their corresponding laboratory results used to categorize these disorders. Typically, metabolic acid-base imbalances stem from kidney-related issues, and the body's respiratory system works to counteract these effects.
Metabolic acidosis arises from lower-than-normal bicarbonate levels in the blood. Initially, the PCO2 may register as normal, but following compensation, it will likely decrease as the body strives to restore the appropriate equilibrium between bicarbonate and carbonic acid/carbon dioxide.
In contrast, respiratory acidosis manifests when an excessive amount of carbon dioxide is present in the bloodstream. Initially, bicarbonate levels may remain normal, but compensatory actions might cause them to increase as an attempt to re-establish the proper balance of bicarbonate and carbonic acid/carbon dioxide.
Metabolic alkalosis indicates an elevated \(\mathrm{pH}\) due to an excess of bicarbonate in the blood. At first, PCO2 might appear normal, yet if compensation occurs, it may elevate as the body endeavors to restore the proper ratios of bicarbonate and carbonic acid/carbon dioxide.
Respiratory alkalosis is characterized by a deficit of carbon dioxide in the bloodstream. Initially, the bicarbonate concentration could remain within the normal range. Subsequent renal compensation might lead to a decrease in the blood's bicarbonate concentration, reflecting the kidneys' efforts to restore the proper bicarbonate and carbonic acid/carbon dioxide ratios, thereby aiming to regulate \(\mathrm{pH}\) within the physiological range.34
| Type | \(\mathrm{pH}\) | PCO2 | \(\ce{HCO3-}\) |
|---|---|---|---|
| Metabolic acidosis | ↓ | N, then ↓ | ↓ |
| Respiratory acidosis | ↓ | ↑ | N, then ↑ |
| Metabolic alkalosis | ↑ | N, then ↑ | ↑ |
| Respiratory alkalosis | ↑ | ↓ | N, then ↓ |
13.8: Acid-Base Balance is shared under a CC BY-NC-SA license and was authored, remixed, and/or curated by LibreTexts.
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