Medical Physiology, 3rd Edition

pH and Buffers

pH values vary enormously among different intracellular and extracellular compartments

According to Brønsted's definition, an acidimageN28-1 is any chemical substance (e.g., CH3COOH, image) that can donate an H+. A base is any chemical substance (e.g., CH3COO, NH3) that can accept an H+. The term alkali can be used interchangeably with base.


Hydrogen Ions in Aqueous Solutions

Contributed by Emile Boulpaep, Walter Boron

Hydrogen ions do not truly exist as “free” protons in aqueous solutions. Instead, a shell of water molecules surrounds a proton, forming an extended complex sometimes denoted as H3O+ (hydronium ion) or H9O4+. Nevertheless, for practical purposes, we will treat the proton as if it were free. Also—as we do elsewhere in this book—we shall refer to concentrations of H+, bicarbonate (image), and other ions. Bear in mind, however, that it is more precise to work with ion activities (i.e., the effective concentrations of ions in realistic, nonideal solutions).

[H+] varies over a large range in biological solutions, from >100 mM in gastric secretions to <10 nM in pancreatic secretions. In 1909, the chemist Sørensen introduced the pH scale in an effort to simplify the notation in experiments in which he was examining the influence of [H+] on enzymatic reactions. He based the pH scale on powers of 10: imageN28-2




Origin of the p of pH

Contributed by Walter Boron

Although the conventional wisdom is that the p of pH stands for the power of 10, Norby's analysis of Sørensen's original papers reveals a far more accidental explanation.

Sørensen used p and q to represent two solutions in an electrometric experiment. He arbitrarily assigned the standard q solution an [H+] of 1 N (normal). That is, his standard solution had an [H+] of 1 N, which is Cq = 10−q. His unknown, therefore, had an [H+] of Cp = 10−p. Using this approach, Sørensen proposed the nomenclature p+H … and nowadays, we use simply pH.

We thank Christian Aalkjaer of the University of Aarhus in Denmark for bringing Norby's paper to our attention.


Norby JG. The origin and the meaning of the little p in pH. Trends Biochem Sci. 2000;25:36–37.

Sørensen SPL. Enzymstudien. II. Mitteilung. Über die Messung und die Bedeutung der Wasserstoffionenkoncentration bei enzymatischen Prozessen. Biochem Z. 1909;21:131–304 [22:352–356].

Sørensen SPL. Études enzymatiques. II. Sur la mesure et l'importance de la concentration des ions hydrogène dans les réactions enzymatiques. C R Lab Carlsberg. 1909;8:1–168.

Thus, when [H+] is 10−7 M, the pH is 7.0. The higher the [H+], the lower the pH (Table 28-1). It is worth remembering that a 10-fold change in [H+] corresponds to a pH shift of 1, whereas a 2-fold change in [H+] corresponds to a pH shift of ~0.3.

TABLE 28-1

Relationship between [H+] and pH Values


[H+] (m)




1 × 10−6


1 pH unit

1 × 10−7


1 × 10−8



8 × 10−8


0.3 pH unit

4 × 10−8


2 × 10−8


1 × 10−8


Even small changes in pH can have substantial physiological consequences because most biologically important molecules contain chemical groups that can either donate an H+ (e.g., R–COOH → R–COO + H+) and thereby act as a weak acid, or accept an H+ (e.g., R–NH2 + H+ → image) and thus behave as a weak base. To the extent that these groups donate or accept protons, a pH shift causes a change in net electrical charge (or valence) that can, in turn, alter biological activity either directly (e.g., by altering the affinity for a charged ligand) or indirectly (e.g., by altering molecular conformation).

pH-sensitive molecules include a variety of enzymes, receptors and their ligands, ion channels, transporters, and structural proteins. For most proteins, pH sensitivity is modest. The activity of the Na-K pump (see pp. 115–117), for example, falls by about half when the pH shifts by ~1 pH unit from the optimum pH, which is near the resting pH of the typical cell. However, the activity of phosphofructokinase, a key glycolytic enzyme (see p. 1176), falls by ~90% when pH falls by only 0.1. The overall impact of pH changes on cellular processes can be impressive. For example, cell proliferation in response to mitogenic activation is maximal at the normal, resting intracellular pH but may fall as much as 85% when intracellular pH falls by only 0.4.

Table 28-2 lists the pH values in several body fluids. Because the pH of neutral water at 37°C is 6.81, most major body compartments are alkaline.

TABLE 28-2

Approximate pH Values of Various Body Fluids



Gastric secretions (under conditions of maximal acidity)




Chromaffin granule


Neutral H2O at 37°C


Cytosol of a typical cell


Cerebral spinal fluid (CSF)


Arterial blood plasma


Mitochondrial inner matrix


Secreted pancreatic fluid


Buffers minimize the size of the pH changes produced by adding acid or alkali to a solution

buffer is any substance that reversibly consumes or releases H+. In this way, buffers help to stabilize pH. Buffers do not prevent pH changes, they only help to minimize them.

Consider a hypothetical buffer B for which the protonated form HB(n+1), with a valence of n + 1, is in equilibrium with its deprotonated form B(n), which has the valence of n:



Here, HB(n+1) is a weak acid because it does not fully dissociate; B(n) is its conjugate weak base. Conversely, B(n) is a weak base and HB(n+1) is its conjugate weak acid. The total buffer concentration, [TB], is the sum of the concentrations of the protonated and unprotonated forms:



The valence of the acidic (i.e., protonated) form can be positive, zero, or negative:



In these examples, image (ammonium), H2CO3 (carbonic acid), and image (“monobasic” inorganic phosphate) are all weak acids, whereas NH3 (ammonia), image (bicarbonate), and image (“dibasic” inorganic phosphate) are the respective conjugate weak bases. Each buffer reaction is governed by a dissociation constant, K



If we add to a physiological solution a small amount of HCl—which is a strong acid because it fully dissociates—the buffers in the solution consume almost all added H+:



For each H+ buffered, one B(n) is consumed. The tiny amount of H+ that is not buffered remains free in solution and is responsible for a decrease in pH.

If we instead titrate this same solution with a strong base such as NaOH, H+ derived from HB(n+1) neutralizes almost all the added OH:



For each OH buffered, one B(n) is formed. The tiny amount of added OH that is not neutralized by the buffer equilibrates with H+ and H2O and is responsible for an increase in pH. imageN28-3


Buffering of H+

Contributed by Walter Boron

For each H+ buffered in Equation 28-6, one B(n) is consumed. For each OH buffered in Equation 28-7, one B(n) is formed. Because almost all of the added H+ or OH is buffered, the change in the concentration of the unprotonated form of the buffer (i.e., Δ[B(n)]) is a good index of the amount of strong acid or base added per liter of solution.

A useful measure of the strength of a buffer is its buffering power (β), which is the number of moles of strong base (e.g., NaOH) that one must add to a liter of solution to increase pH by 1 pH unit. This value is equivalent to the amount of strong acid (e.g., HCl) that one must add to decrease the pH by 1 pH unit. Thus, buffering power is



In the absence of image, the buffering power of whole blood (which contains erythrocytes, leukocytes, and platelets) is ~25 mM/pH unit. This value is known as the image buffering power (image). In other words, we would have to add 25 mmol of NaOH to a liter of whole blood to increase the pH by 1 unit, assuming that β is constant over this wide pH range. For blood plasma, which lacks the cellular elements of whole blood, image is only ~5 mM/pH unit, which means that only about one fifth as much strong base would be needed to produce the same pH increase.

According to the Henderson-Hasselbalch equation, pH depends on the ratio [CO2]/[image]

The most important physiological buffer pair is CO2 and image. The impressive strength of this buffer pair is due to the volatility of CO2, which allows the lungs to maintain stable CO2 concentrations in the blood plasma despite ongoing metabolic and buffer reactions that produce or consume CO2. Imagine that a beaker contains an aqueous solution of 145 mM NaCl (pH = 6.81), but no buffers. We now expose this solution to an atmosphere containing CO2 (Fig. 28-1). The concentration of dissolved CO2 ([CO2]Dis) is governed by Henry's law (see Box 26-2):




FIGURE 28-1 Interaction of CO2 with water.

At the temperature (37°C) and ionic strength of mammalian blood plasma, the solubility coefficient, s, is ~0.03 mM/mm Hg. Because the alveolar air with which arterial blood equilibrates has a image of ~40 mm Hg, or torr, [CO2]Dis in arterial blood is



So far, the entry of CO2 from the atmosphere into the aqueous solution has had no effect on pH. The reason is that we have neither generated nor consumed H+. CO2 itself is neither an acid nor a base. If we were considering dissolved N2 or O2, our analysis would end here because these gases do not further interact with simple aqueous solutions. The aqueous chemistry of CO2, however, is more complicated, because CO2reacts with the solvent (i.e., H2O) to form carbonic acid:



This CO2 hydration reaction is very slow. imageN28-4 In fact, it is far too slow to meet certain physiological needs. The enzyme carbonic anhydrase,imageN18-3 present in erythrocytes and elsewhere, catalyzes a reaction that effectively bypasses this slow hydration reaction. Carbonic acid is a weak acid that rapidly dissociates into H+ and image:




Out-of-Equilibrium image Solutions

Contributed by Walter Boron

As described on page 629, adding CO2 to an aqueous solution initiates a series of two reactions that, in the end, produce image and protons:


(NE 28-1)

In the absence of carbonic anhydrase (CA), the first reaction—the hydration of CO2 to form carbonic acid—is extremely slow. The second is extremely fast. At equilibrium, the overall rate of the two forward reactions must be the same as the overall rate of the reverse reactions:


(NE 28-2)

The first of these reverse reactions is extremely rapid, whereas the second—the dehydration of H2CO3 to form CO2 and water—is extremely slow in the absence of CA. As noted in the text, we can treat the system as if only one reaction were involved:


(NE 28-3)

Moreover, we can describe this pseudoreaction by a single equilibrium constant. As shown by Equation 28-15, we can describe the equilibrium condition in logarithmic form by the following equation:


(NE 28-4)

In other words, considering only equilibrium conditions, it is impossible to change pH or [image] or [CO2] one at a time. If we change one of the three parameters (e.g., pH), we must change at least one of the other two (i.e., [image] or [CO2]). Unfortunately, there are many cases in which the addition of CO2 and image markedly enhances some process. For example, at identical values of intracellular pH, the activation of quiescent cells by mitogens is far more robust in the presence than in the absence of the physiological image buffer. Which buffer component is critical in this case, CO2 or image? There are also many cases in which a stress such as respiratory acidosis or metabolic acidosis triggers a particular response. For example, respiratory acidosis stimulates image reabsorption by the kidney—a metabolic compensation to a respiratory acidosis, as discussed beginning on pages 641–642 in the text. Again, we can ask which altered parameter signals the kidney to increase its reabsorption of image, the rise in [CO2] or the fall in pH?

In the 1990s, the laboratory of Walter Boron realized that it could exploit the slow equilibrium CO2 + H2O ⇄ H2CO3 to create image solutions that are temporarily out of equilibrium. eFigure 28-1A illustrates how one can make an out-of-equilibrium (OOE) “pure” CO2 solution that contains a physiological level of CO2, has a physiological pH, but contains virtually no image. The approach is to use a dual syringe pump to rapidly mix the contents of two syringes, each flowing at the same rate. One syringe contains a double dose of CO2 (e.g., 10% CO2) at a pH that is so low (e.g., pH 5.40) that, given a pK of ~6.1, very little image is present. The other syringe contains a well-buffered, relatively alkaline solution that contains no CO2 or image. The pH of this second solution is chosen so that at the instant of mixing at the T connection, the solution has a pH of 7.40. Of course, the [CO2] after mixing is 5% (which corresponds to a image of ~37 mm Hg), and the [image] is virtually zero. The solutions flow so rapidly that they reach the cell of interest before any significant re-equilibration of the CO2 and image. Moreover, a suction device continuously removes the solution. As a result, the cells are continuously exposed to a freshly generated OOE solution.

eFigure 28-1B illustrates how one could make the opposite solution—a “pure” image solution that has a physiological [image] and pH but virtually no [CO2]. The OOE approach can be used to make solutions with virtually any combination of [CO2], [image], and pH—at least for moderate pH values. At extremely alkaline pH values, the reaction CO2 + OH → image generates image so fast that it effectively short-circuits the OOE approach. Conversely, at extremely acid pH values, the rapid reaction H+ + image ⇄ H2CO3 creates relatively high levels of H2CO3 so that even the uncatalyzed reaction H2CO3 → CO2 + H2O is high enough to short-circuit the OOE approach. Nevertheless, at almost any pH of interest to physiologists, OOE technology allows one to change [CO2], [image], and pH one at a time.

Work on renal proximal tubules with OOE solutions has shown that proximal tubules have the ability to sense rapid shifts in the CO2 concentration of the basolateral (i.e., blood-side) solution that surrounds the outside of the tubule. This work suggests that the tubule has a sensor for CO2 that is independent of any changes in pH or image. The tubule uses this CO2-sensing mechanism in its response to respiratory acidosis (see pp. 637–638). This metabolic compensation involves a rapid increase in the rate at which it transports image from the tubule lumen to the blood, and thus a partial correction of the acidosis.

Work on neurons cultured from the hippocampus suggests that certain neurons can detect rapid decreases in the extracellular image concentration (image), independent of any changes in pHo or [CO2]o. The cell may use this detection system to stabilize intracellular pH (pHi) during extracellular metabolic acidosis, which would otherwise lower pHi.


EFIGURE 28-1 Generation of out-of-equilibrium solutions. (Adapted from Zhao J, Hogan EM, Bevensee MO, et al: Out-of-equilibrium image solutions and their use in characterizing a novel K/HCO3 cotransporter. Nature 374:636–639, 1995.)


Ganz MB, Boyarsky G, Sterzel RB, et al. Arginine vasopressin enhances pHi regulation in the presence of image by stimulating three acid–base transport systems. Nature. 1989;337:648–651.

Ganz MB, Perfetto MC, Boron WF. Effects of mitogens and other agents on mesangial cell proliferation, pH and Ca2Am J Physiol. 1990;259:F269–F278.

Zhao J, Hogan EM, Bevensee MO, et al. Out-of-equilibrium image solutions and their use in characterizing a novel K/HCO3 cotransporter. Nature. 1995;374:636–639.

Zhao J, Zhou Y, Boron WF. Effect of isolated removal of either basolateral image or basolateral CO2 on image reabsorption by rabbit S2 proximal tubule. Am J Physiol Renal Physiol. 2003;285:F359–F369.

This dissociation reaction is the first point at which pH falls. Note that the formation of image (the conjugate weak base of H2CO3) necessarily accompanies the formation of H+ in a stoichiometry of 1 : 1. The observation that pH decreases, even though the above reaction produces the weak base image, is sometimes confusing. A safe way to reason through such an apparent paradox is to focus always on the fate of the proton: if the reaction forms H+, pH falls. Thus, even though the dissociation of H2CO3 leads to generation of a weak base, pH falls because H+ forms along with the weak base.

Unlike the hydration of CO2, the dissociation of H2CO3 is extremely fast. Thus, in the absence of carbonic anhydrase, the slow CO2 hydration reaction limits the speed at which increased [CO2]Dis leads to the production of H+image can accept a proton to form its conjugate weak acid (i.e., H2CO3) or release a second proton to form its conjugate weak base (i.e., image). Because this latter reaction generally is of only minor physiological significance for buffering in mammals, we will not discuss it further.

We may treat the hydration and dissociation reactions that occur when we expose water to CO2 as if only one reaction were involved:



Moreover, we can define a dissociation constant for this pseudoequilibrium:



In logarithmic form, this equation becomes imageN28-5


Derivation of the Henderson-Hasselbalch Equation

Contributed by Emile Boulpaep, Walter Boron

As shown in Equation 28-13,


We can define a dissociation constant for this pseudoequilibrium:


(NE 28-5)

Factoring out [H2O], we can define an apparent equilibrium constant:


(NE 28-6)

The equation immediately above is Equation 28-14 in the text. K at 37°C is ~10−6.1 M or 10−3.1 mM. Taking the log (to the base 10) of each side of this equation, and remembering that log(a × b) = log(a) + log(b), we have


(NE 28-7)

Remembering that log [H+] ≡ −pH and log K ≡ −pK, we may insert these expressions into Equation NE 28-7 and rearrange to obtain


(NE 28-8)

This expression is the same as Equation 28-15 in the text. Finally, we may express [CO2] in terms of image, recalling from Henry's law that [CO2] = s × image:


(NE 28-9)

This is the Henderson-Hasselbalch equation, a logarithmic restatement of the image equilibrium in Equation NE 28-5 above.



Finally, we may express [CO2] in terms of image, recalling from Henry's law that [CO2] = s · image:



This is the Henderson-Hasselbalch equation, a logarithmic restatement of the image equilibrium in Equation 28-14. Its central message is that pH depends not on [image] or image per se, but on their ratio. Human arterial blood has a image of ~40 mm Hg and an [image] of ~24 mM. If we assume that the pK governing the image equilibrium is 6.1 at 37°C, then



Thus, the Henderson-Hasselbalch equation correctly predicts the normal pH of arterial blood.

image has a far higher buffering power in an open than in a closed system

The buffering power of a buffer pair such as image depends on three factors:

1. Total concentration of the buffer pair, [TB]. Other things being equal, β is proportional to [TB].

2. The pH of the solution. The precise dependence on pH will become clear below.

3. Whether the system is open or closed. That is, can one member of the buffer pair equilibrate between the “system” (the solution in which the buffer is dissolved) and the “environment” (everything else)?

If neither member of the buffer pair can enter or leave the system, then HB(n+1) can become B(n), and vice versa, but [TB] is fixed. This is a closed system. An example of a closed-system buffer is inorganic phosphate in a beaker of water, or a titratable group on a protein in blood plasma. In a closed system, the buffering power of a buffer pair is imageN28-6


Derivation of Expressions for Buffering Power in Closed and Open Systems

Contributed by Emile Boulpaep, Walter Boron

How does buffering power, in either a closed or open system, depend on pH?

Closed System

We start with a restatement of Equation 28-8, in which we define buffering power as the amount of strong base that we need to add (per liter of solution) in order to increase the pH by 1 pH unit. In differential form, this definition becomes the following:


(NE 28-10)

Second, because the change in the concentration of the unprotonated form of the buffer, [B(n)], is very nearly the same as the amount of strong base added (see Equation 28-7), Equation NE 28-10 becomes


(NE 28-11)

The third step is to combine Equation 28-3, reproduced below,


(NE 28-12)

and Equation 28-5, reproduced below,


(NE 28-13)

to obtain an expression that describes how [B(n)] depends on the concentration of total buffer, [TB], and [H+]:


(NE 28-14)

Finally, we obtain the closed-system buffering power by taking the derivative of [B(n)] in Equation NE 28-14 with respect to pH. If we hold [TB] constant while taking this derivative (i.e., if we assume that the buffer can neither enter nor leave the system), we obtain the following expression for βclosed:


(NE 28-15)

Open System

How does βopen depend on image and pH? The first two steps of this derivation are the same as for the open system, except that we recognize that, for the image buffer pair, [B(n)] is [image]:


(NE 28-16)

For the open system, the third step is to rearrange the Henderson-Hasselbalch equation (see Equation 28-16, reproduced below),


(NE 28-17)

and rearrange it to solve for [image]:


(NE 28-18)

If we take the derivative of [image] in the above equation with respect to pH, holding image constant, the result is βopen:


(NE 28-19)

Because Equation NE 28-18 tells us that everything after “2.3” in the above equation is, in fact, “[image],” we obtain the final expression for βopen:


(NE 28-20)

This is Equation 28-20 in the text.



Two aspects of Equation 28-18 are of interest. First, at a given [H+], βclosed is proportional to [TB]. Second, at a given [TB], βclosed has a bell-shaped dependence on pH (green curve in Fig. 28-2A). βclosed is maximal when [H+] = K (i.e., when pH = pK). Most image buffers in biological fluids behave as if they are in a closed system. Although many fluids are actually mixtures of several image buffers, the total image in such a mixture is the sum of their βclosed values, each described by Equation 28-18. The red curve in Figure 28-2B shows how total image varies with pH for a solution containing a mixture of nine buffers (including the one described by the green curve), each present at a [TB] of 12.6 mM, with pK values evenly spaced at intervals of 0.5 pH unit. In this example, total image is remarkably stable over a broad pH range and has a peak value that is the same as that of whole blood: 25 mM/pH unit. Indeed, whole blood is a complex mixture of many image buffers. The most important of these are titratable groups on hemoglobin and, to a far lesser extent, other proteins. Even less important than the “other proteins” are small molecules such as inorganic phosphate. The [TB] values for these many buffers are not identical, and the pK values are not evenly spaced. Nevertheless, the buffering power of whole blood is nearly constant near the physiological pH.


FIGURE 28-2 Buffering power in a closed system. In B, the solution contains nine buffers, each at a concentration of 12.6 mM and with pK values evenly spaced 0.5 pH unit apart.

The other physiologically important condition under which a buffer can function is in an open system. Here, one buffer species (e.g., CO2) equilibrates between the system and the environment. A laboratory example is a solution containing CO2 and image in which dissolved CO2 equilibrates with gaseous CO2 in the atmosphere (see Fig. 28-1). A physiological example is blood plasma, in which dissolved CO2equilibrates with gaseous CO2 in the alveoli. In either case, [CO2]Dis is fixed during buffering reactions. However, the total CO2—[CO2] + [image]—can vary widely. Because total CO2 can rise to very high values, image in an open system can be an extremely powerful buffer. Consider, for example, a liter of a solution having a pH of 7.4, a image of 40 mm Hg (1.2 mM CO2), and an [image] of 24 mM—but no other buffers (Fig. 28-3, stage 1). What happens when we add 10 mmol of HCl? Available [image] neutralizes almost all of the added H+, forming nearly 10 mmol H2CO3 and then nearly 10 mmol CO2 plus nearly 10 mmol H2O (see Fig. 28-3, stage 2A). The CO2 that forms does not accumulate, but evolves to the atmosphere so that [CO2]Dis is constant. What is the final pH? If [CO2]Dis remains at 1.2 mM in our open system, and if [image] decreases by almost exactly 10 mM (i.e., the amount of added H+), from 24 to 14 mM, the Henderson-Hasselbalch equation predicts a fall in pH from 7.40 to 7.17, corresponding to an increase in free [H+] of




FIGURE 28-3 Buffering of strong acids and bases by image in an open system.

Even though we have added 10 millimoles of HCl to 1 L, [H+] increased by only 28 nanomolar (see Fig. 28-3, stage 3A). Therefore, the open-system buffer has neutralized 9.999,972 mmol of the added 10 mmol H+. The buffering provided by image in an open system (βopen) is so powerful because only depletion of image limits neutralization of H+. The buildup of CO2 is not a limiting factor because the atmosphere is an infinite sink for newly produced CO2.

The opposite acid-base disturbance occurs when we add a strong base such as NaOH. Here, the buffering reactions are just the reverse of those in the previous example. If we add 10 mmol of NaOH to 1 L of solution, almost all added OH combines with H+ derived from CO2 that enters from the atmosphere (see Fig. 28-3, stage 2B). One image ion forms for each OH neutralized. Buffering power in this open system is far higher than that in a closed system because CO2 availability does not limit neutralization of added OH. Only the buildup of image limits neutralization of more OH. In this example, even though we added 10 millimoles NaOH, free [H+] decreased by only 12 nanomolars (see Fig. 28-3, stage 3B) as pH rose from 7.40 to 7.55.

Whether image neutralizes an acid or a base, the open-system buffering power is imageN28-6



Notice that βopen does not have a maximum. Because βopen is proportional to [image], βopen rises exponentially with pH when image is fixed (Fig. 28-4, blue curve). In normal arterial blood (i.e., [image] = 24 mM), βopen is ~55 mM/pH unit. As we have already noted, the buffering power of all image buffers (image) in whole blood is ~25 mM/pH unit. Thus, in whole blood, βopen represents more than two thirds of the total buffering power. The relative contribution of βopen is far more striking in interstitial fluid, which lacks the cellular elements of blood and also has a lower protein concentration.


FIGURE 28-4 Buffering power of the image system. At pH 7.4 on the blue curve, the solution has the same composition as does arterial blood plasma: a image of 40 mm Hg and an [image] of 24 mM. If the system is “open” (i.e., CO2 equilibrates with atmosphere), βopen is ~55 mM/pH unit at pH 7.40 and rises exponentially with pH (see Equation 28-20). If the system is closed, βclosed is only ~2.6 mM/pH unit at pH 7.4 and is maximal at the pK of the buffer (see Equation 28-18).

image does not necessarily behave as an open-system buffer. In the previous example, we could have added NaOH to a image solution in a capped syringe. In such a closed system, not only does accumulation of image limit neutralization of OH, but the availability of CO2 is limiting as well. Indeed, for fluid having the composition of normal arterial blood, the closed-system image buffering power is only 2.6 mM/pH unit (see Fig. 28-4, black curve), <5% of the βopen value of 55 mM/pH unit. You might think that the only reason the closed-system buffering power was so low in this example is that the pH of 7.4 was 1.3 pH units above the pK. However, even if pH were equal to the pK of 6.1, the closed-system image buffering power in our example would be only ~14 mM/pH unit, about one quarter of the open-system value at pH 7.4. A physiological example in which the image system is “poorly open” is ischemia, wherein a lack of blood flow minimizes the equilibration of tissue CO2 with blood CO2. Thus, ischemic tissues are especially susceptible to large pH shifts.