Carbon Dioxide Hydration Reaction

The presence of CO₂ in the blood creates a special problem because CO₂ reacts with water (i.e., CO₂ is hydrated) to form carbonic acid (H₂CO₃). The blood must possess effective buffering mechanisms to prevent harmful increases in acidity as it transports CO₂. As the following reaction shows, when CO₂ diffuses from tissues into capillary blood, the CO₂ hydration reaction occurs:¹

Hydrolysis is a chemical reaction in which H₂O splits another molecule; the term is incorrect in reference to the reaction between H₂O and CO₂. Reaction 1 shows that CO₂ in the blood is associated with the formation of H₂CO₃ and the production of hydrogen ions (H⁺). A fundamental characteristic of an acid is the release of H⁺ into solution. Although CO₂ is technically not an acid, its immediate formation of H₂CO₃ in physiological fluids allows it to be treated as though it were an acid.

Hydration Reaction: Chemical Equilibrium (Le Chatelier’s Principle)

The CO₂ hydration reaction (reaction 1) can proceed in either direction (i.e., it is reversible). Chemical equilibrium exists when the reaction velocities to the left and right are equal. A state of equilibrium does not mean the concentrations of substances on both sides of the equation are equal, and it does not imply a static state of affairs. At equilibrium, the reaction continues to move in both directions at equal rates but in such a way that no net change occurs in the concentrations of constituents on either side of the reaction. This state is called dynamic equilibrium. The reaction between CO₂ and H₂O in plasma occurs only to a slight extent; in this state of the chemical equilibrium, CO₂ and H₂O molecules far outnumber H₂CO₃ molecules. Therefore, at equilibrium, the CO₂ hydration reaction is shifted to the left, as the following reaction shows:

The short arrows pointing to the right and the long arrows pointing to the left indicate that at equilibrium the reaction is left shifted. The numbers in parentheses indicate relative concentrations of the constituents.² At equilibrium, concentrations to the left of the arrows are greater than concentrations to the right of the arrows. Despite these concentration differences, the velocities of right-directed and left-directed reactions are identical at equilibrium.

According to Le Chatelier’s principle, if a system at equilibrium is placed under stress, it will react in such a way that counteracts the stress and creates a new equilibrium at a different point.³ For example, in reaction 2, if we add a certain amount of CO₂ molecules to the left side of the reaction, we disrupt the equilibrium and drive the reaction to the right until a new equilibrium is established. Similarly, removing CO₂ molecules from the left side “pulls” the reaction to the left. In other words, adding or subtracting CO₂ molecules drives the reaction either toward the H⁺ and HCO3− side or toward the CO₂ and H₂O side of the reaction.

Relationship between Dissolved Carbon Dioxide and Carbonic Acid Concentrations

All gases dissolve in water to some extent, as discussed in previous chapters. According to Henry’s law, the amount of CO₂ dissolving in plasma is proportional to the PCO₂ to which the plasma is exposed. In other words, CO₂ in the gaseous phase establishes equilibrium with CO₂ in the aqueous phase. In the lung, gaseous CO₂ in the alveoli establishes equilibrium with dissolved CO₂ in capillary blood plasma. The concentration of CO₂ in alveolar gas is expressed in terms of its partial pressure (mm Hg). The concentration of dissolved CO₂ in plasma is proportional to P_ACO₂ and is measured in millimoles per liter (mmol/L) rather than in milliliters per deciliter (mL/dL), as it is for oxygen content.

At 37° C, 0.03 mmol of CO₂ dissolves in 1 L of plasma for each mm Hg PCO₂; that is, mmol/L (CO₂) = 0.03 mmol/L/mm Hg × PCO₂. (The derivation of the conversion factor [0.03] to change PCO₂ to millimoles per liter of dissolved CO₂ is shown in Box 9-1.) Thus, if plasma PCO₂ is 40 mm Hg, the concentration of dissolved CO₂ in plasma is as follows: 40 mm Hg × 0.03 mmol/L/mm Hg = 1.2 mmol/L, which is the normal amount of dissolved CO₂ in arterial blood. Figure 9-1 summarizes the relationship between gaseous CO₂, dissolved CO₂, and H₂CO₃. The slow uncatalyzed reaction between CO₂ and H₂O theoretically requires about 400 molecules of dissolved CO₂ to produce 1 molecule of H₂CO₃.² In other words, an extremely small amount of H₂CO₃ in plasma is in equilibrium with a relatively large pool of dissolved CO₂. The reaction in Figure 9-1 shows that an increased alveolar PCO₂ increases plasma PCO₂, which increases the dissolved CO₂ and H₂CO₃ concentration (law of mass action). Conversely, a decrease in alveolar PCO₂ forces the hydration reaction to the left, decreasing the plasma H₂CO₃ concentration. Because it is in equilibrium with CO₂ gas, H₂CO₃ is called a volatile acid. Plasma PCO₂ is an important marker of the blood’s volatile acid content and the adequacy of alveolar ventilation.

BOX 9-1   Conversion Factors for Dissolved Carbon Dioxide: Millimoles per Liter and Milliliters per Deciliter

1. 1 mole CO₂ = 22.3 L at STP; 1 mmol CO₂ = 22.3 mL at STP

2. CO₂ sol coeff = 0.592 mL CO₂/mL plasma/760 mm Hg PCO₂

3. mL CO₂ dissolving per mL plasma per mm Hg PCO₂ is as follows: 0.592/760 mm Hg = 0.00078 mL CO₂/mL plasma/mm Hg

4. mL CO₂ dissolving per 100 mL plasma (mL/dL) = 0.078 mL

5. mL CO₂ dissolving per L (1000 mL) plasma per mm Hg PCO₂ is as follows:

0.00078× 1000 = 0.78 mL CO₂/L plasma/mm Hg

6. mmol CO₂ dissolving per L plasma per mm Hg PCO₂ is as follows: 0.78 mL CO₂/22.3 mL CO₂ per mmol = 0.03498 mmol CO₂/L/mm Hg

PCO2×0.03=mmol/LdissolvedCO2

7. Factor to convert mmol/L to mL/dL (from steps 4 and 6) = 0.078/0.03498 = 2.23:

mmol/LCO2×2.23=mL/dLCO2

mL/dLCO2/2.23=mmol/LCO2

STP, Standard temperature and pressure; sol coeff, soluble coefficient.

Although the concentration of dissolved CO₂ in plasma is not equal to the concentration of plasma H₂CO₃, its concentration is in direct proportion to the concentration of H₂CO₃; this makes it possible to treat CO₂ as if it were the acid instead of H₂CO₃. Because dissolved CO₂ and H₂CO₃ are indistinguishable from each other by chemical analysis,⁴ dissolved CO₂, which can be easily measured, is commonly substituted for the concentration of H₂CO₃ in clinical acid-base equations (see Chapter 10).

Role of Ventilation in Regulating Arterial Carbon Dioxide Pressure and Volatile Acid

Changes in alveolar PCO₂ ultimately change the plasma H₂CO₃ concentration (see Figure 9-1). (Chapter 4 discusses the reciprocal relationship between V˙A and PaCO₂.) Figure 9-2 shows the chemical events that occur during hypoventilation. In Figure 9-2, V˙A is halved, causing P_ACO₂ to double. Subsequently, PaCO₂ and dissolved CO₂ double, pushing the hydration reaction to the right. The doubling of CO₂ molecules on the left side of the reaction drives the reaction to the right to a new equilibrium point (Le Chatelier’s principle). Figure 9-2 shows that at equilibrium, the amount of H₂CO₃ molecules doubles, but the ratio of CO₂ to H₂CO₃ molecules stays the same.

At the outset of hypoventilation, the rate of CO₂ production momentarily exceeds the rate of alveolar CO₂ elimination, causing alveolar and blood CO₂ levels to increase. However, CO₂ levels cannot continually increase; instead, a new equilibrium point (i.e., a new steady state) is reached at which CO₂ production and elimination rates are again equal, but at higher P_ACO₂ and PaCO₂ levels. In Figure 9-1, a normal V˙A of 4 L per minute maintains a P_ACO₂ of 40 mm Hg and eliminates 200 mL of CO₂ per minute (an amount equal to its production rate). If the V˙A suddenly decreases to 2 L per minute with no change in CO₂ production (see Figure 9-2), the PaCO₂ and dissolved CO₂ in the blood increase temporarily until a new steady state is reached, at which time CO₂ production and elimination rates are again equal. However, for this steady state to occur at a V˙A of 2 L per minute, P_ACO₂ and PaCO₂ must increase to 80 mm Hg, and dissolved plasma CO₂ must increase to 2.4 mmol/L (80 × 0.03) (see Figure 9-2). A proportionate increase in H₂CO₃ levels is produced, increasing the blood’s acidity. This new steady state requires less ventilatory work to eliminate the same amount of CO₂; for this reason, people with limited ventilatory capacity adopt this ventilatory strategy, but the “tradeoff” is increased blood acid levels.

Hyperventilation also disrupts the chemical equilibrium of the hydration reaction (Figure 9-3). Hyperventilation pulls the reaction leftward to a new equilibrium point. At the new equilibrium, the ratio between dissolved CO₂ and H₂CO₃ is preserved, while plasma H₂CO₃ concentration is decreased. In Figure 9-3, V˙A suddenly doubles from 4 L per minute to 8 L per minute, whereas CO₂ production stays constant at 200 mL per minute. The CO₂ elimination rate temporarily exceeds the CO₂ production rate until a new equilibrium is established in which production and elimination rates are equal again. To eliminate only 200 mL of CO₂ per minute at a V˙A of 8 L per minute, the P_ACO₂ must be at 20 mm Hg, which is associated

with a dissolved plasma CO₂ of 0.6 mmol/L. A proportionate reduction in plasma H₂CO₃ concentration is produced, making the plasma less acidic or more alkaline. In this situation, more ventilatory work is required to eliminate the same amount of CO₂ that is eliminated under normal conditions.

It is important that normal lungs eliminate CO₂ at a rate equal to its production by the body, but it is equally important for V˙A to keep the P_ACO₂ near the 40-mm Hg range; otherwise, the blood becomes acidic or alkaline. For example, during heavy exercise, V˙A normally increases such that it eliminates CO₂ at a rate proportional to its production; this maintains the P_ACO₂ near 40 mm Hg and maintains normal blood acid (H₂CO₃) levels. An elaborate neurochemical control system (see Chapter 11) is responsible for regulating V˙A in this manner.

In some circumstances, the body does not maintain P_ACO₂ at 40 mm Hg, even when lung airways are normal. The lack of oxygen at high altitudes creates a chemical stimulus for ventilation, overriding the regulatory system that normally maintains a P_ACO₂ of 40 mm Hg. In such circumstances, the increased ventilation drives P_ACO₂ and PaCO₂ below 40 mm Hg, reducing the blood H₂CO₃ concentration; this creates a state of arterial alkalemia.

Ventilation also increases in response to abnormally high levels of nonvolatile (noncarbonic) acids in the blood (acidemia). Acids increase the blood’s H⁺ concentration, which chemically stimulates ventilation, eliminating more CO₂ and decreasing PaCO₂. These events pull the hydration reaction to the left, which decreases blood H₂CO₃ levels and offsets the blood’s high levels of nonvolatile acids. These events constitute an important compensatory mechanism. The opposite compensatory response occurs for abnormally low blood levels of nonvolatile acids; the normal acidic stimulus to ventilation is reduced, decreasing the V˙A. These events drive the hydration reaction to the right, increasing blood H₂CO₃ concentration; this counteracts the effects of an abnormally low nonvolatile acid level. (These important compensatory ventilatory responses to nonrespiratory acid-base disturbances are discussed in detail in Chapter 10.)

Dissolved Carbon Dioxide

As illustrated in Figure 9-4, when CO₂ diffuses out of the tissues into the blood, a small portion physically dissolves in the plasma, but most diffuses into the erythrocyte. The small amount of dissolved CO₂ in the plasma is in equilibrium with the plasma PCO₂, which determines the direction and rate of CO₂ diffusion at both tissue and alveolar levels. CO₂ dissolves according to its solubility coefficient (PCO₂ × 0.03 = mmol/L dissolved CO₂). Thus, venous blood, with a PCO₂ of about 46 mm Hg, contains 1.38 mmol/L of dissolved CO₂ (46 × 0.03 = 1.38). This converts to about 3 mL/dL of dissolved CO₂ in venous blood.

As the following equation shows, water hydrates a minuscule amount of plasma CO₂ to form H₂CO₃:

CO2+H2O→H2CO3

The rate of this reaction is quite slow in the plasma; no enzyme is present, as it is in the erythrocyte, to catalyze and speed up the reaction. The amount of CO₂ carried as H₂CO₃ in the plasma is infinitesimally small because at the pH of normal body fluids, H₂CO₃ instantly dissociates into HCO3− and H⁺.⁴ Because dissolved CO₂ reacts with H₂O to form H₂CO₃, dissolved CO₂ plays a key role in determining the blood pH (see Chapter 10).

Bicarbonate

Most CO₂ diffusing into the blood from tissue cells enters the erythrocyte rather than remaining in the plasma. Several reactions occur in the erythrocyte, ultimately generating large amounts of HCO3−, which diffuse into the plasma (see Figure 9-4).

The reaction between CO₂ and H₂O occurs about 13,000 times faster in the erythrocyte than in the plasma because of the intracellular enzyme carbonic anhydrase.⁴ When CO₂ enters the red blood cell, it is instantly converted into H₂CO₃ and its products, HCO3− and H⁺. An extremely small amount of CO₂ remains dissolved in the erythrocyte fluid; this explains why most of the CO₂ molecules diffusing into the blood enter the erythrocyte rather than linger in the plasma. The rapid conversion of dissolved CO₂ to HCO3− and H⁺ creates a concentration gradient that continually draws more plasma CO₂ into the erythrocyte.

Chloride Shift

When the catalyzed hydration reaction forms H⁺ and HCO3− ions in the erythrocyte, hemoglobin immediately buffers the H⁺ ions, removing them from solution (see Figure 9-4). H⁺ buffering occurs at the same time hemoglobin releases oxygen; deoxygenated hemoglobin is a more effective buffer than oxygenated hemoglobin. Removal of H⁺ ions from solution by hemoglobin prevents their accumulation in the erythrocyte, producing the following three consequences: (1) the hydration reaction accelerates even further, (2) the reaction continually moves to the right, and (3) HCO3− ions build up in the erythrocyte. Eventually, the erythrocyte’s HCO3− concentration builds up to a level higher than that of the plasma. Consequently, HCO3− (a negatively charged ion) diffuses down its concentration gradient, out of the erythrocytes, and into plasma (see Figure 9-4). The outward movement of HCO3− creates an electropositive environment inside the erythrocyte; to maintain electroneutrality, negatively charged plasma chloride (Cl⁻) ions diffuse into the erythrocyte. This event is known as the chloride shift.

The chloride shift occurs for several reasons. Positively charged ions in the erythrocyte, such as Na⁺ or K⁺, do not accompany exit of HCO3− because the erythrocyte membrane is impermeable to positive ions. Before it buffers H⁺, the negative charges of the hemoglobin molecule are exactly balanced by the positive ions in the erythrocyte. Diffusion of CO₂ into the erythrocyte generates H⁺ ions, which immediately combine with hemoglobin, reducing hemoglobin’s net negative charge. However, this reduction in negative charge is matched by newly generated HCO3− ions, maintaining electrical neutrality inside the erythrocyte. Because the HCO3− concentration increases above that of the surrounding plasma, HCO3− diffuses out of the erythrocyte, leaving the erythrocyte’s interior positive with respect to the plasma. In response to this electrostatic gradient, the most abundant negatively charged ion in the plasma, Cl⁻, moves into the erythrocytein exchange for HCO3− (see Figure 9-4). Hamburger, a Dutch physiologist, first described the chloride shift, which explains why it is also known as the Hamburger phenomenon.

Most CO₂ is transported from the tissues to the lungs in the form of plasma HCO3−, most of which is generated in the erythrocyte; this helps explain why the plasma HCO3− concentration is so much greater than the plasma H⁺ concentration. The plasma HCO3− does not arise from the plasma hydration reaction, which creates equal amounts of H⁺ and HCO3− ions; rather, it originates from the erythrocyte hydration reaction, where buffering of H⁺ by hemoglobin allows HCO3− to accumulate and diffuse into the plasma.

All the reactions involving CO₂ are reversed in the lungs (Figure 9-5). CO₂ diffuses out of the blood into alveoli because of its partial pressure gradient. According to the law of mass action, as CO₂ leaves the blood, the hydration reaction in the erythrocyte shifts to the left. This shift causes HCO3− and H⁺ to rapidly form CO₂ and H₂O, quickly depleting the erythrocyte HCO3− and H⁺ concentrations. Consequently, plasma HCO3− diffuses into the cell and hemoglobin releases H⁺. The CO₂ produced from these reactions diffuses into ventilated alveoli and is exhaled to the atmosphere.

CLINICAL FOCUS 9-2   Consequences of Intravenous Bicarbonate Ion Infusion

You are called to the emergency department for resuscitation of a 63-year-old man who does not have a pulse or respirations. On arrival, you note that cardiopulmonary resuscitation (CPR) is in progress and an endotracheal tube has been placed into the patient’s trachea. You are asked to ventilate the patient’s lungs with a self-inflating resuscitation bag attached to the endotracheal tube. About 5 minutes after the start of CPR, blood gases reveal a severe metabolic acidosis. The physician administers sodium bicarbonate intravenously to counteract the acidosis. You ventilate the patient more vigorously, increasing the minute ventilation. Why is this necessary?

Discussion

When a person has cardiopulmonary arrest, a lack of blood flow and oxygen to the tissues causes anaerobic metabolism and lactic acid formation. As the following equation shows, blood bicarbonate buffers this lactic acid, generating CO₂:

VentilationH+(lacticacid)+HCO3−→H2CO3→H2O+CO2↑

Increased V˙A is needed to eliminate the additional CO₂ generated by lactic acid

buffering. Otherwise, the buildup of CO₂ counteracts the beneficial effects of HCO3−.^∗

---

^∗The use of sodium bicarbonate in this situation is not recommended, given the current understanding of emergency resuscitation. Nevertheless, this example is given here to illustrate the important concept that the buffering action of bicarbonate generates CO₂.

As illustrated in Figure 9-4, a small amount of CO₂ reacts with the free amino groups (NH₂) of plasma proteins to form carbamino compounds. This rapid reaction requires no catalyst (CO₂ + prot-NH₂ → prot-NHCOO⁻ + H⁺). Prot-NH₂ represents a protein molecule with a free amino group. When CO₂ combines with the amino group, it displaces a H⁺ ion; a millimole of H⁺ is released for each millimole of carbamino compound (prot-NHCOO⁻) formed. These H⁺ ions and the H⁺ ions produced by the plasma hydration reaction are buffered by other plasma proteins (see Figure 9-4). (By convention, the carbamino compound amino group–CO₂ combination is written as NHCOO⁻ rather than NHCO2−.)

Erythrocyte Carbamino Compounds: Carbaminohemoglobin

As with plasma proteins, CO₂ combines with the hemoglobin protein molecule at free amino sites. This reaction is rapid, requiring no catalyst (CO₂ + HbO₂-NH₂ → Hb-NHCOO⁻ + H⁺). The resulting compound of hemoglobin and CO₂ is called carbaminohemoglobin (Hb-NHCOO⁻). A millimole of H⁺ is released for each millimole of Hb-NHCOO⁻ formed. These H⁺ ions are buffered by the hemoglobin molecule, along with H⁺ ions formed by the rapid hydration reaction in the erythrocyte (see Figure 9-4). The hemoglobin molecule carries oxygen and CO₂ but not at the same binding site. Oxygen combines with the heme groups (see Chapter 8), whereas CO₂ combines with the amino groups of α and β globin (protein) chains. However, the affinity of hemoglobin for CO₂ is greater when it is not combined with oxygen, a phenomenon known as the Haldane effect. The presence of oxygen on the heme portion hinders carbaminohemoglobin formation. Conversely, carbaminohemoglobin has a decreased affinity for oxygen (Bohr effect). Oxygenated blood carries less CO₂ for a given PCO₂ than deoxygenated blood. Figure 9-6 shows the CO₂ dissociation curve at different PO₂ values. For a given PCO₂, more CO₂ is carried at lower PO₂ values (lower saturations).

The CO₂ dissociation curve is essentially linear over the physiological range (see Figure 9-6), in contrast to the sigmoid-shaped oxygen-hemoglobin dissociation curve (see Chapter 8). Thus, a change in alveolar ventilation is much more effective in changing arterial CO₂ content than O₂ content; that is, a doubling of the alveolar ventilation in a healthy person cuts the CO₂ content in half but changes arterial O₂ content very little because the hemoglobin is already nearly 100% saturated with normal ventilation.

Relative Contribution of Carbon Dioxide Transport Mechanisms

In discussing how much each blood mechanism contributes to CO₂ transport, a distinction should be made between the amount of tissue CO₂ taken up by venous blood and the total venous CO₂ content. The amount of CO₂ taken up at the tissues is equal to the mixed venous CO₂ content minus the arterial CO₂ content; this is the amount of CO₂ actually transported to and eliminated by the lungs. Dissolved CO₂, HCO3−, and carbamino compounds are responsible for 8%, 80%, and 12% of CO₂ transport.⁴ Table 9-1 shows the percentage contribution of each major blood mechanism to CO₂ content and CO₂ transport.⁴ (The percentage H₂CO₃ contributes is so small that it is not included in Table 9-1.) The milliliters per deciliter concentrations of CO₂ are calculated by multiplying millimoles per liter by the conversion factor, 2.23. The CO₂ concentration in plasma alone (see Table 9-1) is greater than that of whole blood (plasma plus erythrocytes) because the concentration of CO₂ in erythrocytes is considerably less than that in the plasma. The erythrocyte volume (about 40% of the whole blood volume) represents a dilute solution of CO₂, which when mixed with plasma, produces a lower CO₂ concentration in whole blood (millimoles per liter) than exists in plasma alone.

Complementary Interaction between Oxygen and Carbon Dioxide Transport: Bohr and Haldane Effects

Haldane and Bohr effects are mutually enhancing. While CO₂ is taken up at the tissues by various blood components, oxygen is diffusing down its partial pressure gradient from the blood plasma to the tissue cells (see Figure 9-4). Oxygen molecules dissociate from hemoglobin in response to decreasing plasma PO₂. The dissociation of oxygen from hemoglobin is augmented by the following two simultaneous processes: (1) CO₂ enters the erythrocyte and binds with the globin chains’ amino acid groups, and (2) H⁺ ions are formed by the rapid hydration of CO₂ (see Figure 9-4). Both processes shift the oxyhemoglobin equilibrium curve to the right, decreasing hemoglobin’s affinity for oxygen (Bohr effect). Hemoglobin releases more oxygen than it otherwise would for a given PO₂. The reaction illustrating the Bohr effect is shown in Figure 9-4 (H⁺ + HbO₂ ⇌ HHb + oxygen).

Because the presence of H⁺ and CO₂ decreases hemoglobin’s affinity for oxygen, the presence of oxygen on the hemoglobin molecule must decrease hemoglobin’s affinity for H⁺ and CO₂. As it releases oxygen at the tissue level, hemoglobin more effectively takes up (buffers) the H⁺ produced by the CO₂ hydration reaction. This buffering action promotes the production of HCO3−, increasing this mode of CO₂ transport. Because deoxygenated hemoglobin combines with CO₂ more readily than oxygenated hemoglobin, CO₂ transport by this mode (carbaminohemoglobin) is also enhanced. Thus, decreased oxygenation of the hemoglobin increases the blood’s capacity for CO₂ at any given PCO₂ (Haldane effect). Haldane and Bohr effects enhance one another in a similar fashion in the lungs.