pH

Most formal definitions of acids or bases when applied to biologic solutions utilize the Brønsted–Lowry concept, which classifies acids as proton donors and bases as proton acceptors. A more appropriate working definition, however, may be that acids are substances that increase H⁺ ion activity (aH⁺), the preferred chemical term for protons in aqueous solutions [9]. Hydrogen ion activity is preferred to the term H⁺ ion concentration ([H⁺]) for a number of reasons. First, a proton unaccompanied by an electron shell is much smaller than elements such as Na⁺ or Cl⁻ or any molecule. The proton is therefore more appropriately viewed as a highly reactive electromotive force or free energy source in solution rather than a concentration. Second, the strength of an acid and resultant acidity of a solution are determined, in part, by its activity coefficient, a factor influenced by temperature and ionic strength that determines the degree of dissociation. The glass electrode that is used to determine pH provides a measure of the hydrogen ion activity referenced to that of standard buffers, and theoretically [H⁺] can be calculated by multiplying aH⁺ by the activity coefficient. However, the activity coefficient for the proton in water cannot be accurately estimated and is < 1 in vertebrate fluids because the ionic strength approximates 0.15 M and the fluids are not infinitely dilute (ionic strength ≈ 0 M, the value where the activity concentration = 1). In other words, use of the term [H⁺] is based on an incorrect assumption that the activity coefficient is 1. Third, the concept of pH is historically based as a symbol for the “power of hydrogen” and defined as the negative value of the log to the base 10 of hydrogen ion activity, –log₁₀(aH⁺). The logarithmic expression was developed to simplify the notation necessary to describe the large amounts of H⁺ produced and the changes of aH⁺ observed in nature and chemical experiments. This means of expression, although cumbersome mathematically, defines the nonlinear relationship between pH and aH⁺ (Fig. 30.2):

(30.4)

Fourth, the variability in pH is much less than that of aH⁺, and the distribution of aH⁺ is usually skewed in vertebrate fluids. Consequently, pH is more accurately expressed as mean ± SD than aH⁺. Finally, because the proton is so reactive in water, it exists fleetingly in many forms, with the hydronium ion (H₃O⁺) being the predominant form. As such, use of the term “hydrogen ion concentration,” which suggests a constant state over time, is misleading and inaccurate when applied to protons in water.

Because a base is defined as a H⁺ (proton) acceptor, each acid dissociates into H⁺ and a potential H⁺ acceptor, or conjugate base. For example, H₂CO₃ in aqueous solution dissociates into H⁺ and its conjugate base HCO₃⁻ (eqn 30.2). Substances that are strong acids have weak conjugate bases and vice versa. Interestingly, water, the most abundant solvent in the body, can function as both an acid (H₃O⁺; proton donor) or a base (H₂O; proton acceptor) depending on local conditions (H₃O⁺ ↔ H⁺ + H₂O). At normal pH (7.40), temperature (37–38 °C), and ionic strength (0.15 M), water is the most abundant base in the body.

Physiologically and clinically, the formation of acids and therefore H⁺ production are emphasized in acid–base regulation because the end‐products of ingestion, absorption, and metabolism, as well as that of many pathophysiologic processes, are protons or CO₂.

Regardless of conversion issues and the relatively narrow range (20–150 nEq/L) over which changes in aH⁺ occur in biologic fluids, the concept of pH has persisted and is routinely reported on all pH and blood gas analyzers.

pK_a and the Henderson–Hasselbalch equation

The development of methods to measure pH in blood by Hasselbalch (1912) and CO₂ content and cHCO₃⁻ in plasma or serum by Van Slyke (1917) led to the widespread adoption of the Henderson–Hasselbalch equation and subsequent characterization of acid–base disturbances as being either respiratory or non‐respiratory (metabolic) in origin. The presence of many chemical equilibriums in blood (e.g., phosphates and sulfates) and the law of mass action produce many potential equilibrium equations that could be used to explain acid–base balance. The reasons why the carbonic acid equilibrium equation (eqn 30.3) was chosen to describe acid–base balance were (1) historical (methods were available to determine CO₂ content); (2) HCO₃⁻ is quantitatively the most important buffer in extracellular fluid; and (3) the carbonic acid equation includes a volatile substance (CO₂). The respiratory system in air‐breathing vertebrates is an open system in that CO₂ can leave the system (body) rapidly by an increase in minute ventilation. As a consequence, the buffering ability of HCO₃⁻ in an open system is much greater than that in a closed system, such as a test tube that is capped. This is best appreciated by expressing eqn 30.3 in the reverse direction, such that:

(30.5)

Continuous removal of CO₂ by ventilation in an open system drives the reaction to the right and therefore greatly assists the rapid removal of excess protons from the body.

The law of mass action states that the rate (velocity) of a reaction is dependent on the concentration of the reactants, the dissociation constant (K) for the reaction, as well as the temperature and ionic strength of the solution. For mammalian fluids such as blood, plasma, and serum, changes in temperature and ionic strength are relatively small, and consequently, their effects can be ignored for clinical purposes. It is important to note that this is not the case for ectothermic (“cold‐blooded” animals).

The rate of dissociation (r) for an uncharged monoprotic acid (HA) can be characterized as:

(30.6)

using the concentration of HA (cHA) and dissociation constant K₁,

(30.7)

Similarly:

(30.8)

and using the hydrogen ion activity and the concentration of A⁻ (cA⁻):

(30.9)

which at equilibrium results in r₁ = r₂, or the following mass action equation:

(30.10)

where K_a is the dissociation constant for the uncharged monoprotic acid HA. Applying this general approach to carbonic acid, an uncharged diprotic acid (H₂CO₃) using the dissociation reaction described in eqn 30.3, Henderson derived:

(30.11)

Henderson used the concentration of dissolved molecular CO₂ instead of H₂CO₃ because H₂CO₃ could not be measured. Hasselbalch then introduced PCO₂ into Henderson’s equation and applied Sørensen’s logarithmic format, producing the Henderson–Hasselbalch equation for carbonic acid, arguably the most famous equation in biology:

(30.12)

where pH is –log₁₀(aH⁺), pK_a is = −log₁₀(K_a), and S (0.0307 {mmol/L}/mmHg) is the solubility of CO₂ in the solution. In blood, plasma, and serum, pK_a is termed the apparent dissociation constant (pK₁´) for carbonic acid, with a value at 37 ^oC of 6.095 when blood is analyzed and 6.105 when plasma or serum is analyzed.

The Henderson–Hasselbalch equation has been rewritten for clinical explanatory purposes as follows:

(30.13)

This equation is an oversimplification of acid–base regulation because it ignores the role of non‐bicarbonate buffers such as hemoglobin, albumin, globulins, and phosphate, ignores the role of the liver in amino‐acid metabolism and bicarbonate formation, and implies that the time required for renal adaptation to an acute acid–base disturbance (typically hours to days) is similar to the time required for respiratory adaptation to an acute acid–base disturbance (typically minutes). It is therefore time to retire eqn 30.13.

Chemical buffers

Chemical buffers are substances that minimize changes in the pH of a solution when an acid or base is added. A buffer solution consists of a weak acid and its conjugate base and is most effective when the pH is within 1.0 to 1.5 pH units of its dissociation constant (pK_a) [16] (Table 30.3). Alterations in blood, interstitial, and intracellular fluid pH are immediately buffered (modified) by chemical buffer systems, as described by the mass action equation (eqn 30.10). For the weak acid HA, as pH decreases (aH⁺ increases), cA⁻ decreases, and cHA increases by equal amounts, keeping the total amount of A_TOT (A_TOT = cHA + cA⁻) the same.

The principal chemical buffers are the bicarbonate (HCO₃⁻/H₂CO₃), protein (Prot/H‐Prot⁺), and phosphate (HPO₄^2–/H₂PO₄^–) systems. In rapidly growing animals, bone can readily contribute calcium carbonate and calcium phosphate to the extracellular fluid, thereby increasing the buffering capacity. Functionally, anytime there is a decreased pH in the body, the buffer (HCO₃⁻, Prot, HPO₄^2–) at physiologic pH accepts the excess proton, converting the buffer to its conjugate acid (H₂CO₃, H‐Prot⁺, H₂PO₄^–‐) [17]. Because the body can have only one pH in a well‐mixed fluid compartment, the ratio of the acid to salt forms of the various buffer pairs in solution can always be predicted by the relevant mass action equation (isohydric principle), providing their concentration and dissociation constant (pK_a) are known [18]. In other words, with knowledge of the behavior of one buffer pair, one can predict the behavior of all the other buffer pairs in solution. As pointed out above, the HCO₃⁻/H₂CO₃ buffer pair is most frequently used to determine acid–base status in clinical practice because it is the most prominent chemical buffer in the extracellular fluid and in the presence of carbonic anhydrase, carbonic acid forms CO₂, which is eliminated by alveolar ventilation. Thus, during normal conditions, the body can be considered an “open” system.

Approximately, 60% of the body’s chemical buffering capacity occurs by intracellular phosphates and proteins. Inorganic and organic (ATP, ADP, and 2,3‐DPG) phosphates possess pK_a values that range from 6.0 to 7.5, making them ideal chemical buffers over a wide range of potential intracellular pH values. The most important intracellular protein‐dissociable group is the imidazole ring of the amino‐acid histidine (pK_a 6.4–6.7). The α‐amino groups of proteins (pK_a 7.4–7.9) play a secondary but important role in intracellular buffering.

Hemoglobin contributes approximately 80% of the non‐bicarbonate buffering capacity of whole blood and with other intracellular proteins is responsible for three‐fourths of the chemical buffering power of the body. Plasma proteins, particularly albumin, typically contain 16 histidine and 1 α‐amino groups, with some species differences, and collectively are responsible for approximately 20% of the non‐bicarbonate buffering capacity of whole blood.

Inorganic phosphate (pK_a 6.8) and creatinine (pK_a 5.0) are the major buffers in carnivore and omnivore urine because renal tubular pH (typically 5.5–6.5) is within 1.0 to 1.5 units of their pK_a [19].

Respiratory system

The respiratory system provides a route by which pH can be regulated by varying the partial pressure of carbon dioxide (PCO₂) in blood (Fig. 30.1). Chemoreceptors throughout the body, but particularly those located in the carotid body and medulla oblongata, monitor changes in pH and PCO₂ and adjust breathing (tidal volume and frequency) to maintain a normal pH. The association of protons with cHCO₃⁻ and the subsequent formation of CO₂ and H₂O is an example of very rapid chemical buffering (closed system), whereas subsequent elimination of CO₂ by the lung via increased ventilation (open physiologic system) requires a longer response time (usually minutes). Changes in blood PCO₂ also have important consequences for hemoglobin’s affinity for oxygen and its buffering capacity. Increases in PCO₂ decrease blood pH and decrease hemoglobin affinity for oxygen (Bohr effect: shifting the oxygen–hemoglobin curve to the right). This decrease in the oxygen affinity of hemoglobin is advantageous in tissues, allowing hemoglobin to release more oxygen for metabolism. Non‐oxygenated hemoglobin in turn can transport more CO₂ in the form of hemoglobin carbamino compounds to the lungs (Haldane effect).

Renal system

The synthesis of new HCO₃⁻ and excretion of excess H⁺ emphasizes the role of the kidneys in both chemical and physiologic buffering (Fig. 30.1). Although relatively slow (hours to days), compared with the lungs (minutes) and chemical buffering (seconds), the kidney serves as the principal means by which acids produced by metabolic processes (not owing to CO₂ production) are ultimately eliminated (Fig. 30.3). Protons produced by metabolic processes are excreted in the urine in combination with weak anions (titratable acidity), primarily phosphate and to a lesser extent creatinine, and in a much smaller amount as free protons [19,20].

The term “titratable acidity” may be considered synonymous with urinary phosphate concentration but actually represents the buffering effect of all weak acids in urine, including creatinine and urate. Although titratable acidity is frequently emphasized when the renal role in acid–base balance is discussed, the much more important mechanism for acid excretion during severe metabolic acidosis is urinary ammonium excretion (Fig. 30.4); however, it can take up to five days for the maximal rate of urinary ammonium excretion to be obtained [21]. Ammonium (NH₄⁺) is produced in the proximal tubule primarily from glutamine metabolism to α‐ketoglutarate and NH₃, a process that simultaneously generates HCO₃⁻ that is returned to the plasma.

Traditional (bicarbonate‐centric) approach

Traditional descriptions of acid–base balance and acid–base abnormalities are based on the Brønsted–Lowry definition (proton donor) of an acid and base and the Henderson–Hasselbalch equation to determine pH. The terms “acidemia” and “alkalemia” are used to describe whether the blood pH is acid or alkaline, respectively. The terms “acidosis” and “alkalosis” are used to describe the abnormal or pathologic (‐osis) state due to the accumulation or decrease of protons in the body. Importantly, the terms “acidosis” and “alkalosis” must always be preceded by the descriptor respiratory or non‐respiratory (metabolic). The Henderson–Hasselbalch equation characterizes all acid–base disturbances as being either respiratory or non‐respiratory because of the body’s production and elimination of volatile (dissolved CO₂, H₂CO₃) and non‐volatile or fixed (e.g., lactic, ketoacid, and uremic) acids, respectively.

Clinically, the terms “respiratory” and “non‐respiratory” (metabolic) have been used to imply the respective roles of the lung and kidney in acid–base regulation. The term “non‐respiratory” frequently replaces metabolic in many discussions of acid–base imbalance because it incorporates all mechanisms responsible for acid–base imbalance other than the production of CO₂ and carbonic acid (H₂CO₃). Therefore, only four primary acid–base abnormalities are possible when using the traditional approach to acid–base evaluation: respiratory acidosis, non‐respiratory acidosis, respiratory alkalosis, and non‐respiratory alkalosis (Table 30.4; Fig. 30.6).

The traditional approach posits that primary acid–base abnormalities arise from altered PCO₂ or cHCO₃⁻ (Fig. 30.5). A fundamental flaw with using the traditional approach and the Henderson–Hasselbalch equation is that blood cHCO₃⁻ cannot be an independent predictor of blood pH because cHCO₃⁻ is calculated from blood pH and PCO₂!

In the traditional approach, simple acid–base abnormalities occur only when either PCO₂ or cHCO₃⁻ are responsible for the acid–base disturbance. Mixed acid–base abnormalities are caused by disturbances in both PCO₂ or cHCO₃⁻. Mixed acid–base abnormalities may be additive (respiratory and non‐respiratory acidosis) or offsetting (respiratory alkalosis and non‐respiratory acidosis) with regard to their ability to influence pH (Table 30.5) [25]. Offsetting mixed acid–base abnormalities occur when two primary acid–base abnormalities produce opposite effects on plasma pH. Animals with offsetting mixed acid–base abnormalities have both acidosis and alkalosis but do not necessarily demonstrate acidemia or alkalemia because the offsetting change might be partially (i.e., returns pH toward normal) or fully compensatory in which case blood pH will be normal. Observations that should lead to suspicions of a mixed acid–base disturbance when evaluating blood gas and pH values include [25]:

• the presence of a normal pH with abnormal PCO₂ or cHCO₃⁻
  
• a pH change in a direction opposite to that predicted for the known primary disorder
  
• PCO₂ and cHCO₃⁻ changing in opposite directions.

Mixed acid–base disorders can be classified based on the origin of the primary disturbances as mixed respiratory disturbances, mixed non‐respiratory and respiratory disturbances, mixed non‐respiratory disturbances, and triple disorders. They also can be classified based on their effect on an animal’s pH in additive combinations, offsetting combinations, and occasionally triple disorders (Table 30.5). In additive combinations, both primary disorders tend to change pH in the same direction (e.g., respiratory acidosis and non‐respiratory acidosis), whereas in offsetting combinations, the primary disorders tend to change the pH in opposite directions (e.g., respiratory alkalosis and non‐respiratory acidosis). The final pH reflects the dominant of the two offsetting disorders in offsetting combinations.

Secondary or compensatory (adaptive) acid–base changes frequently occur in response to most primary acid–base abnormalities and aid in buffering or minimizing changes in plasma pH. Respiratory acid–base abnormalities, for example, are generally compensated for by controlled, oppositely directed changes in non‐respiratory function (Table 30.6). In simple acid–base abnormalities such as primary respiratory acidosis caused by hypoventilation, the kidney compensates by producing non‐respiratory alkalosis (Table 30.4).

Respiratory compensation for primary non‐respiratory acidosis is accomplished by increasing alveolar ventilation and CO₂ excretion by the lungs. Non‐respiratory acidosis is characterized by a decrease in blood cHCO₃⁻ and pH, and a decrease in PCO₂ (respiratory alkalosis), caused by secondary hyperventilation. Non‐respiratory alkalosis is characterized by an increase in blood cHCO₃⁻ and pH, and an increase in PCO₂, owing to compensatory hypoventilation (respiratory acidosis) (Table 30.4).

In primary respiratory acid–base disorders, the compensation occurs in two phases. The first phase consists of titration of protons by non‐bicarbonate buffers, and the second phase reflects renal compensation of the acid–base disorder, by increasing or decreasing HCO₃⁻ and Cl⁻ excretion in the urine (Fig. 30.7). Primary respiratory acidosis is characterized by increased PCO₂ due to alveolar hypoventilation, leading directly to decreased pH and an associated increase in blood cHCO₃⁻. Renal compensation occurs by titration of non‐bicarbonate buffers, an increase in net acid and Cl⁻ excretion by the kidneys, and increased HCO₃⁻ reabsorption and production by the kidneys. Primary respiratory alkalosis is characterized by decreased PCO₂ due to alveolar hyperventilation, leading directly to increased pH and a compensatory decrease in blood cHCO₃⁻. The initial compensation in respiratory alkalosis is caused by release of H⁺ from non‐bicarbonate buffers within cells. The second phase is mediated by a compensatory decrease in net acid excretion by the kidneys.

When analyzing secondary changes in a given acid‐base disorder, it is important to remember the following three points:

• With the exception of chronic respiratory alkalosis, compensation does not return the pH to normal. The pH almost always trends toward the primary condition.
  
• Overcompensation does not occur.
  
• Sufficient time must elapse for compensation to reach a steady state, at which time the expected compensation can be estimated (see Table 30.6).

The question that often arises when analyzing simple acid–base abnormalities that demonstrate both respiratory acidosis and non‐respiratory alkalosis is, “Which is the primary problem and is there a secondary and/or compensatory event?” The answer is not always obvious, although simple primary acid–base abnormalities generally change pH in the direction of the primary disorder. For example, an animal with respiratory acidosis and compensatory non‐respiratory alkalosis would have a pH that tended to be acidemic (below the reference range; e.g., a pH of 7.31). Mixed respiratory and non‐respiratory acid–base abnormalities are much more difficult to decipher, and like simple acid–base abnormalities must be carefully evaluated in the context of an animal’s disease, and other available diagnostic information.

Because non‐respiratory acidosis is so frequently associated with disease processes in animals, two indices of acid–base balance have been developed to enable quantitative evaluation of the non‐respiratory component of an acid–base abnormality: standard bicarbonate and base excess. Standard bicarbonate was the first index to be developed, and it represents the concentration of bicarbonate in plasma after the whole‐blood sample has been equilibrated to a PCO₂ of 40 mmHg at 37 °C. This index purportedly quantifies the non‐respiratory component of an acid–base abnormality; however, standard bicarbonate is no longer used because it is increased in acute respiratory acidosis and initially required two additional measurements using a tonometer. Methods to estimate standard bicarbonate using a nomogram or calculate standard bicarbonate from other acid–base values have been developed. These methods have the advantage that tonometry is not required; however, nomograms are cumbersome and slow to use, and the equation is based on the calculated value for base excess. Consequently, base excess has become widely used to evaluate the non‐respiratory component of an acid–base abnormality.

Base excess (buffer‐centric) approach

Buffering and buffer values are foundational concepts when acid–base balance is interpreted using the base excess approach, and consequently, the approach is best described as buffer‐centric or buffer‐centered. The base excess of blood (BE_b) quantitates the mmols of strong acid (such as HCl) or strong base (such as NaOH) required to titrate 1 L of blood to pH 7.40 at 37 °C while the PCO₂ is held constant at 40 mmHg. BE_b from healthy humans is defined as 0 mmol/L with a reference range of –2.3 to +2.3 mmol/L. A positive BE_b value indicates base excess (non‐respiratory alkalosis), whereas a negative BE_b value indicates a base deficit (non‐respiratory acidosis) [4].

Values for BE_b were initially read from a graphical calculating device called a “nomogram,” which is a two‐dimensional diagram that provides an approximate graphical computation of a relationship between two or more entities. Since 1977, BE_b has been calculated using the Van Slyke equation [26], and the nomogram is no longer used to provide the value for BE_b. The Van Slyke equation takes the following general form for human arterial (oxygenated) blood when temperature = 37 °C, pH = 7.40, and PCO₂ = 40 mmHg, with Z providing an adjustment for the effect of changes in hematocrit assuming a fixed value for the mean hemoglobin concentration in erythrocytes [27], such that:

(30.14)

Stay updated, free articles. Join our Telegram channel

Join