This chapter focuses on the ways in which a changing CO2 concentration might alter the pH of a solution, particularly that of your precious bodily fluids. The physiological consequences of acidaemia and alkalaemia are discussed in dedicated chapters, as are the various effects of having an excessively high or precipitously low blood CO2 level (independent of pH changes). A chapter which summarises the bedside rules and equations used in the interpretation of blood gases is also available as a brief overview of the empirically derived formulae which describe acute and chronic compensation for acidosis and alkalosis.
Generally speaking, when asked "what happens in an acute respiratory acidosis", most people would spout something inane like "CO2 turns into acid", or equivalent. A more detailed explanation is expected from ICU staff. The best foundations for such an explanation can be found in Chapter 4.5 from "Acid-Base Physiology" by Brandis. The following "footnotes to Brandis" represent my attempt to explain these concepts to myself, slowly and patiently, as one might explain something to a toddler.
Any discussion of acute respiratory acidosis should begin with an exploration of CO2 transport.
Rather than to address this topic in detail here, one might summarise by saying that CO2 in the bloodstream is transported in three major forms: as dissolved gas, as carbamate (bound to haemoglobin and other proteins) and as bicarbonate. Bicarbonate is by far the most important form in terms of volume, and one gets their bicarbonate by either spontaneous or catalysed hydration of CO2, which becomes carbonic acid, and degenerates rapidly into HCO3- and H+.
The reaction can be summarised in this manner:
CO2 + H2O ⇌ H2CO3 ⇌ HCO3- + H+
So, now you have performed the synthesis of CO2 and H2O into carbonic acid, H2CO3. This begs the question: how much of it is there? And if it were so important, then why don't we measure carbonic acid levels?
Well. Carbonic acid is a fairly fragile thing. Indeed, for a long time its existence was viewed assomething of a thought experiment like the hydrogen ion. Even such eminent bodies as PubChemrefer to it as a "hypothetical" acid. The reason for this sort of dismissive treatment is the extreme instability of carbonic acid under most physiological circumstances, and the consistent historical failures of people to identify and analyse a measurable quantity of it. When somebody did finally manage to make a small amount of it in 1993, its existence was confined to a thin frozen film in a methanol ice matrix, in a 77°K environment. Serious macroscopic amounts of carbonic acid, probably presenting in the form of amorphous lumps, might be found scattered liberally across the bleak tundra-like surface of Kuiper Belt objects (where near absolute zero temperatures protect it from degradation). The fact that it can only remain stable in such extreme environments adequately conveys the utter ridiculousness of looking for measurable quantities of H2CO3 in the body fluids of living human organisms.
However, there must be some amount of it present, because it is a necessary intermediate between the two sides of the CO2-HCO3- equation. So, how much H2CO3 is actually in the body at any given moment, even theoretically? No satisfactory consensus exists. Nunn's Respiratory Physiology (2005) quotes some unreferenced "published work" as a source for a figure of 1%. In other words, 1% of all transported CO2 is in the form of H2CO3 at any given moment, either in the process of being converted to bicarbonate, or back again. This seems believable, given the ubiquity and industriousness of carbonic anhydrase.
With this knowledge, the effects of carbonic anhydrase on the raw substrates (CO2 and H2O) can be summarised in the following manner:
This diagram, while acknowledging that H2CO3 is an intermediate product, appropriately trivialises its role in the process of converting CO2 and H2O into HCO3- and H+.
However, H2CO3 does exist in the bloodstream, albeit in miniscule concentrations, and its contribution to physiology is non-trivial.
Its importance lies in its being a weak acid. And at physiological pH, it is present in such minute concentrations precisely because most of it has dissociated. The result is a surplus of hydrogen ion activity, recognisable by ICU staff as a drop in pH.
Carbonic acid is a polyprotic acid in the old Brønsted acid sense; it is able to donate either one or two protons, with each dissociation having its own pKa value. In doing so, it generates a conjugate base - either HCO3- or CO32-. The former is clearly more likely: observe this Bjerrum plot for the dissociation of these three chemical species.
Weirdly, when one searches for such an image, one finds an exhausting array of plots specific to oceanic sea water and salt-water swimming pools, but very few specific to the human blood. Fortunately, a detailed examination of bicarbonate kinetics in the human bloodstream is available from Ole Siggaard-Andersen (1962). Apparently, the pKa of the first dissociation is about 6.1-6.3, and of the second somewhere around 9.3. Therefore at a normal physiological pH, the equilibrium equation is shifted heavily in favour of HCO3- and there is very little of the other species present.
So, how much H+ will be produced per change in pCO2? Well. One can use a stripped-down Henderson-Hasselbalch equation (simply, the Henderson equation) to assess the change in pH which is to be expected from a change in CO2.
Thus, the ratio of CO2 and HCO3 multiplied by a dissociation constant gives us an impression of the increase in hydrogen ion activity we can expect from the change in pCO2.
That "apparent" first dissociation constant is derived empirically, from experiments - there is no theoretical way to determine what that value should be (Nunn's reports that its value is 24, if the αpCO2 is represented in terms of mmHg). The [H+] output of this equation is in nmol/L. Thus, if one plugs in normal values, one ends up with a [H+] of around 40 nmol/L, which corresponds to a normal human blood pH of 7.40.
Because that would be insane.
A buffer system cannot buffer itself. Consider: H2CO3, the product of H2O and CO2, dissociates into H+ and HCO3-. What point would there be in HCO3- binding H+ again? That would not be buffering; you would just get H2CO3 again, which would again dissociate and still generate H+, and the process would continue ad infinitum as an equilibrium reaction. Clearly, another system needs to soak up the extra H+ activity generated as a result of this process in order for true buffering to occur.
That system consists of intracellular phosphate and intracellular proteins, of which haemoglobin is the key player.
As discussed far above (without the sillyness of diagrams) this concept can also be explained in terms of a nice adult-looking equilibrium equation:
CO2 + H2O ⇌ H2CO3 ⇌ HCO3- + H+
Inside the red cells, intracellular carbonic anhydrase catalyses the conversion of CO2 and H2O into H2CO3, which exists for mere moments before decomposing spontaneously into HCO3- and H+.
If substrate is added to the left side of the equation, the equilibrium will shift right, favouring the production of H+ and bicarbonate. The H+ is then immediately absorbed by the vast supplies of intracellular phosphate, and the extensive collection of proteins (of which the chief is haemoglobin, because it is the most numerous - available at a concentration of 330g/L in the red cells).
Now, most people at this stage would complain that this reaction in the cells would grind to a halt once the intracellular HCO3- and H+ concentrations increase to a certain level. However, this equilibrium state is never reached.
The products of the reaction are constantly being removed:
These two processes are of sufficient importance to merit the following prolonged digression.
A substantial temptation exists to go off on an extensive tangent exploring the buffer properties of proteins; however, in the interest of brevity this discussion will be limited to whatever seems relevant to acute respiratory acid-base disturbances.
Buffering by proteins is performed by amino acid side chains and various other loose groups which hang off the main protein body and don't participate in bonds which are critical to maintaining the protein structure. Generally speaking, all proteins have these sorts of groups hanging off them, but in practice the only amino acid groups worth a damn as buffers at physiological pH are histidine aromatic rings (the imidazole groups). All the others have a pKa which is well outside the normal survivable range.
The aromatic imidazole ring of the histidine molecule has a protonation pKa of 6.8. Well, that generalisation is not entirely accurate, as quaternary protein structure plays a major role in the variation in imidazole ring pKa values, and those histidine molecules which are "buried" deep within proteins have a substantially different pKa to those molecules which are sticking out of the surface. And to add a layer of complexity, protein quaternary structure is plastic and can undergo various weird changes in response to pH and temperature fluctuations. The protein might curl up in a ball or turn itself inside out, resulting in changes in the number of exposed and buried histidine molecules. Several good examples of this exits, but the most notable is probably haemoglobin, which has asomewhat increased buffering capacity in its deoxygenated state. But that is a deep rabbit hole, and earlier the author promised not to get carried away.
So, for the sake of sanity, let's just say the imidazole ring has a protonation pKa of 6.8, which means that at a physiologically normal intracellular pH of the erythrocyte (around 7.20) it is almost completely unprotonated. In other words, of the total number of imidazole groups in any given protein, the majority will contain a nitrogen atom with a negative charge. As the activity of the hydrogen ion increases with decreasing pH, the imidazole rings will encounter more and more hydrogen ions to bind; as the hydrogen ions are bound, they are removed from the solution, and the pH of the solution increases. This underlies the principle of histidine group buffering.
Given that most other amino acid side chains are quite useless at buffering anything, the "buffering power" of all proteins is totally related to the number of histidine side chains they have.
Thus, the plasma proteins are rather weak participants in buffering, owing to the fact that they are present in small concentration (only around 60-70g/L) and their relatively poor histidine content.
Intracellular proteins are more effective as buffers for several reasons:
The awesome buffering power of intracellular contents has been demonstrated many times in many different ways. Perhaps the earliest and most relevant paper to quote here would be the 1958 work by Ellison Straumjord and Hullen. These guys titrated the whole blood and centrifuged plasma of volunteers with dilute hydrochloric acid, comparing their buffer capacities. The chemistry nerd will immediately recall that the definition of buffer capacity is "quantity of strong acid or base that must be added to change the pH of one liter of solution by one pH unit". In their experiments, the authors found that the quantity of HCl required to decrease the pH of human blood by 1.0 was 38.5 mEq/L for whole blood, and 16.1 mEq/L for plasma.
Generally speaking, the bicarbonate content of intracellular fluid is remarkably low, and the content of extracellular fluid is remarkably high. Carbonic anhydrase is scarce in the bloodstream, and the spontaneous generation of H2CO3 is slow and inefficient; thus the higher bicarbonate content of extracellular fluid is accounted for by the fact that HCO3- is actively pumped out of erythrocytes. This is called the Hamburger shift, and it is mediated by the Band 3 transport protein.
Band 3 is unimaginatively named after the third strip of proteinaceous goop which develops when an electrophoresis of erythrocyte membranes is carried out. The Hamburger shift is also unimaginatively named after Hartog Jacob Hamburger, a man who was not only distinguished by a splendid handlebar moustache, but also - little it is known - is credited as the inventer of normal saline (!), which was briefly known as "Hamburger's Solution". This chapter is probably the wrong forum for any more saline-bashing; and in any case the chapter on the contents and properties of normal saline covers this sordid history in satisfying detail.
Back on topic. Band 3 is not merely an ion exchanger; it appears to be a huge complex molecule, with numerous other functions within the erythrocyte. For one, it acts as an anchor for the red cell cytoskeleton. Carbonic anhydrase seems to associate with it in a manner resembling the streamlining of a factory production line, ensuring that the new bicarbonate anions are shuttled out of the cell immediately. In short, Band 3 is critically important in the erythrocyte contribution to acid-base buffering.
At the end of such a long discussion of buffering mechanisms, the pragmatic intensivist- enraged by academic bullshit- would demand to know what happens to the acid-base balance. Rightly so.
Whole-body buffering experiments (which have formed the basis of the Boston Bedside Rules) have revealed that in acute repsiratory disorders, for every 10mmHg rise in CO2, the bicarbonate increases by 1mmol/L. Using this as a source of bicarbonate values, one can calculate the Henderson equation and estimate that this 10mmHg increase will drop the pH from 7.40 to about 7.33.