Normal Acid-Base Regulation – Understanding ABGs – Lecture 2
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[Normal Acid-Base Regulation – Understanding ABGs – Lecture 2]
Eric Strong, MD
Hospitalist, Palo Alto VA Hospital
Affiliated Assistant Professor of Medicine, Stanford University
[Eric Strong] Source: LYBIO.net
Hello, this is Eric Strong from the Palo Alto Veterans Hospital, and Stanford University. This is the second lecture on Understanding ABGs, and the topic is Normal Acid-Base Regulation. Although the topic of acid-base physiology is quite complex, I will be limiting this discussion to just information that will be necessary and clinically relevant at the bedside for the diagnosis and management of common acid-base disorders.
The specific learning objectives of this talk are as follows. First, to know the role of acid-base balance in normal physiology. Second, to understand the three stages of how the body produces, transports, and ultimately eliminates acid. Next, to be familiar with the concept of compensation, and finally, to be familiar with the interaction between acid-base abnormalities, and potassium balance.
The primary role of acid-base balance in the body is to maintain a stable pH within a narrow window centered at 7.40. This is necessary as most enzymes in the human body optimally function here, and relatively small perturbations in pH can lead to their denaturing and even irreversible inactivation. There are few notable exceptions to this. For example, the enzyme pepsin, whose function is to break down ingested proteins optimally operates at a pH of about two, which is perfect for its acidic environment in the stomach.
However, the vast majority of enzymes in the human body prefer a pH at or very close to 7.40. The clinical consequences of dysregulation of acid-balance are many. The most important are poor vascular tone, myocardial pump failure, increased risk of arrhythmias, skeletal muscle weakness, secondary electrolyte abnormalities, delirium or coma and in its most profound stages impairment of cellular respiration leading to death.
The single most important principle of acid-base balance in the body is that the rate of acid production must equal the rate of acid elimination. This process can be broken down into three separate stages. First, acid is produced as a consequence of normal metabolism, and occasionally as a consequence of pathologic processes. Second, this acid has to be transported via the blood, finally, the acid is eliminated via the lungs and kidneys. Let’s take a brief look at each stage one at a time.
First is the production of acid. Most of us have a gut response that acid being produced in the body sounds bad, this is largely due to the fact that the word acid contains a lot of negative connotations outside of medicine, such as acid rain or it’s occasionally used as a horrific weapon. However, in the body, the production of acid is actually a natural and necessary consequence of dual metabolism. The physiologic metabolism of different types of macro molecules needs different acids as byproducts.
Breakdown of carbohydrates, and fats yield carbon dioxide, a volatile acid. This term volatile means that it vaporizes easily, which is a critical characteristic in order for acid-base balance to be possible as it allows carbon dioxide, which carries the major acid burden of the body to be exhaled as a gas by the lungs. Breakdown of protein yields sulphuric acid, and breakdown of phospholipids yields phosphoric acid. These acids are created in much smaller amounts than that of carbon dioxide for the metabolism of carbohydrates and fats.
[Eric Strong] Source: LYBIO.net
However, because both sulphuric and phosphuric acid are non-volatile acids, they can’t be eliminated from the body via the lungs and require a significantly more complicated mechanism in the kidneys for their elimination. Although it is of limited clinical relevance, I’ll quickly point out that metabolism of certain foods can actually result in generation of alkali, however, the rate of alkali production is much smaller than that of acid, which neutralizes its effect on acid-base panels.
In addition to the normal physiologic formation of acids by the body, there are numerous clinical situations in which pathologic production of acid can occur. This is a rough schematic of how the more common of these pathologic acids can be formed. I won’t say anything more about these specific processes at this point as they will be discussed in more detail in future lectures. Once the body has produced acid, either as part of normal physiologic metabolism or as part of an abnormal pathologic process, it needs to be transported in the blood stream to lungs and the kidneys where it can be eliminated.
Acids could simply be transported by having their dissociated hydrogen ions carried freely in the blood. Unfortunately, this would result in sudden and dramatic swings in pH that would not be compatible with life. Therefore, the use of buffers is necessary. In chemistry, a buffer can be defined as such by the presence of two major characteristics. First, it has to consist of either a weak acid and its conjugate base, or a weak base and its conjugate acid.
Second, a buffer has a property of resisting changes in pH. In other words, hydrogen ions could be added or removed from a buffer solution with only minimal change in pH. In the human body, by far the most important buffer is actually the system shown here, which consists of three components in equilibrium with one another. On the far right, there are the bicarbonate and hydrogen ions, which can spontaneously combine into H2Co3 more commonly known as carbonic acid.
Thankfully speaking, the bicarbonate ion, and carbonic acid pair is the true buffer. Carbonic acid can be split spontaneously into carbon dioxide and water, but this reaction is very slow in the absence of the enzyme carbonic anhydrase. Note that the sequence of reactions is completely reversible. Because carbonic acid is in much smaller quantities as carbon dioxide, for simplicity, this reaction is often summarized as this. What is the most important, the bicarbonate system isn’t the only buffer, other extra cellular buffers include phosphoric acid and albumin.
Well, intracellular buffers include organic phosphates such as 2,3-DPG and ATP as well as hemoglobin. While these secondary buffers are physiologically important. In my experience, an awareness of them is usually not of great clinical relevance at the bedside. I hate to bring up the Henderson-Hasselbalch Equation because I’m concerned it may evoke traumatic memories of high school chemistry. However, it is necessary to spend a few minutes on it in order to demonstrate a critical principle in acid-base physiology.
To remind you, this equation states that the pH of a solution is equal to the pKa plus the log the ratio of the concentration of base to the concentration of conjugate acid. The pKa is the negative log of the equilibrium constants, which is an empirically measured constant that describes the kinetics of a specific chemical reaction. If you’re concerned that that you don’t have a particularly strong background in chemistry or for any reason about to become lost, don’t worry, now, this will all make sense in a minute.
[Eric Strong] Source: LYBIO.net
Now, as I just discussed the major buffer in the body is the bicarbonate and carbonic acid pair. So, let’s substitute some numbers into our equation to make it specific for human acid-base regulation. The pKa is an empiric constant, which for this reaction is known to be 6.1. In the event that there are any physical chemists listening you may remember Henry’s Law, which relates to the concentration of a gas and liquid to its partial pressure.
Thus the concentration of carbonic acid can be substituted with 0.03 times the partial pressure of carbon dioxide with which it is an equilibrium with the help of carbonic anhydrase, with which can be easily measured with the ABG. The 0.03 here like the 6.1 is an empirically measured constant. This is an interesting result. We now have an equation, which links three easily measured values, the pH, bicarbonate concentration, and partial pressure of carbon dioxide. Thus, measurement of just two of these in any given situation should be sufficient to calculate the third.
Let’s see how this works out for a person with normal-acid base status. A normal bicarbonate level is approximately 24 millimoles per litre. Of course, it’s actually a range of normal values but since we can’t plug a range into the equation very easily, just trust me that 24 is the single most reliable number to use for a “normal bicarb”. The normal partial pressure of carbon dioxide in arterial blood is 40 mm of mercury, 24 divided by the product of 0.03, and 40 is equal of 20. The log of 20 is 1.3, and finally, 6.1 plus 1.3 equals 7.4, which as it turns out is a normal value of arterial pH.
I love it when physiology and math works so well together. You might now ask, so what. Well, here is the clinical relevance of the Henderson-Hasselbalch Equation. Since, this equation defines the relationship between arterial pH, like carbonic concentration, and the partial pressure of co2 in arterial blood, it is impossible to change one of these values without altering the others. In other words, if there is some shift in either the bicarbonate or the PCO2, the pH must necessarily shift as well.
Our bodies actually make critical use of this linkage to further help blunt changes in pH due to pathologic processes. This is known generally as compensation. From the equation, you can see that in order for the pH to remain constant, any change in bicarb or PCO2 is to be accompanied by a proportional change in the same direction as the other variable. For example, if the initial pathologic process is too much bicarb, the body responds in such a way as to increase PCO2. Low bicarb triggers a decrease in PCO2.
Conversely, if the initial problem is a PCO2 that it’s too high, the body increases bicarb, whereas a low PCO2 triggers a decrease in bicarb, although it would seem to be physiologically advantageous to do so, mechanisms of compensation never work perfectly, that is they never bring the pH completely back to normal. From a physician’s standpoint, it’s actually good that the compensation doesn’t return pH all the way to normal, because if it did, it would be practically impossible to diagnose acid-base disorders.
[Eric Strong] Source: LYBIO.net
That was a little more intuitive when we discuss compensation in detail in the future lecture. Now, that I’ve talked a little bit about how the body transports acid in the blood, I will finally talk about how acid is eliminated from the body altogether. As already mentioned, acid is eliminated from the body at two sites. First, is the lungs. The mechanism for acid elimination here is the expiration of carbon dioxide. Of course, co2 itself isn’t technically an acid.
However, because of the rapid equilibrium it’s in with the bicarbonate and hydrogen ion pair, as co2 is eliminated from the body, formation of new co2 from the ion pair of your carbolic anhydrase becomes kinetically favorable. The second, and more complicated site of acid elimination is the kidneys where acid is partly eliminated in both the proximal convoluted tubule of the nephron, as well as the distal tubule and collecting duct.
In the former, the major mechanism is through reabsorption of bicarbonate, in the latter, two mechanisms are seen. The distal tubule and collecting duct excrete hydrogen ions into the tubule lumen where it combines with HPO4 to form phosphoric acid. Phosphoric acid acts as a buffer here in order to prevent the concentration of hydrogen ions in the tubule lumen from becoming so great that the chemical gradient exceeds that, which the tubular cells can pump against.
The other mechanism present in the distal tubule and collecting duct is the direct excretion of ammonium ion through a complicated series of steps, which starts with a metabolism of the amino acid glutamine by the cells of the proximal tubule. Let’s take a brief look at each organ. Compared to the kidneys, the lungs’ role in acid-base regulation is relatively straight forward. With each inspiration, as oxygen diffuses across the alveolar capillary membrane and into the pulmonary capillary bed, carbon dioxide brought to the lungs by the pulmonary artery diffuses out of the pulmonary capillaries into the alveoli.
From there, the lungs will expel the carbon dioxide with the next expiration. In addition to this role in normal physiology, they can also respond to pathologic changes, and pH as mentioned previously when I talked briefly about the concept of compensation. There are two general pathways of which the lungs can respond to changes in pH. The peripheral pathway, and the central pathway. Let’s take a look at how the lungs respond during a period of acidemia. Either low arterial pH, or high PCO2, which could be sensed independently of one another, stimulate chemoreceptors in the carotid body, which lies right at the bifurcation of the carotid arteries.
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This is the same location that senses low arterial oxygen tension as well. These chemoreceptors then send neural input to the medullary respiratory center, which responds by increasing minute ventilation. Minute ventilation could be increased by a fast respiratory rates, a larger tidal volume or both. The net result is a lower PCO2 and thus a higher pH. The primary trigger of the central pathway is low pH in the CSF itself, which stimulates chemoreceptors on the ventral surface of the medulla. These chemoreceptors then stimulate the medullary respiratory center with the same end result.
Note that in the presence of alkalemia, the opposite process occurs with the result of increased PCO2 and decreased pH. Moving on to the kidneys, first let me orient you to my perfectly rendered nephron, up here is the glomerulus into which flows the afferent arteriole, and away from which flows the efferent arteriole. Immediately off the glomerulus is the proximal convoluted tubule leading to the Loop of Henle, and finally ending with the distal tubule, and collecting duct, which will eventually lead to the ureters and bladder.
After fluid is filtered across the glomerular membrane, the first half is the proximal tubule at which point sodium bicarbonate can be reabsorbed more or less together. Stimuli for this process are low serum pH, hypokalemia, and angiotensin too. While the Loop of Henle plays a major role in electrolyte balance and determining the ultimate concentration of urine, its role in specifically acid-base balance is relatively minor, and of little clinical importance.
Next, we have the collecting duct, where under the stimulus of either aldosterone or low pH, sodium is reabsorbed while potassium and hydrogen is excreted into the lumen. At the same location, under control of hypokalemia, potassium is reabsorbed in exchange for excretion of hydrogen. There is one final aspect of the kidney’s role in acid-base balance. Phosphate is usually reabsorbed predominantly in the proximal tubule. However, this reabsorption is blocked by high levels of PTH or parathyroid hormone, somewhat irrespective what PTH is complicated on calcium phosphate metabolism, excretion can also be stimulated by low pH, to what possible advantage is this? Well, if you remember one of the key components that allows the kidneys to excrete acid in significant quantities is the presence of phosphoric acid in the urine, which acts as a buffer to prevent the urine from becoming too acidic and thus preventing it from having too great a concentration of hydrogen ions for the collecting duct to pump against under the stimulus of either aldosterone, acidemia or hypokalemia.
[Eric Strong] Source: LYBIO.net
In other words, blocking reabsorption of phosphate in the proximate tubule provides a greater amount of buffer in the distal tubule, and collecting duct. I want you to understand that a professor of renal physiology would likely look at this diagram and be upset with what he or she might perceive as extreme over simplicity. However, this is all you need to know about the kidney’s contribution to acid-base balance at the bedside.
The final topic I will discuss in this lecture is the relatively unique relationship between pH and potassium balance. You just saw a little bit about this in the last slide, where I mentioned that hypokalemia can serve as a stimulus for the excretion of hydrogen ions into the urine in exchange for reabsorption of potassium ions. Another interaction between these two ions is related to their transcellular shifts in the presence of acid-base disorders.
In the presence of acid-base disturbances, hydrogen ions move across the cell membrane in exchange for potassium. This effect is most pronounced with metabolic acidoses. Once intracellular, hydrogen ions are buffered by organic phosphates, and hemoglobin. The consequence is that there may be an apparent abnormality of potassium balance that may not represent the body’s total potassium stores. That is in the presence of acidemia, a patient can be net potassium depleted, but still hyperkalemic based on a serum potassium level.
Let’s take a closer look at how this works. This blue box will represent a typical cell, and here are a bunch of potassium ions scattered inside and out. The concentration of potassium is greater within the cell, but in this diagram, there isn’t intended to be any accurate scale of this display per se. Then, here are some hydrogen ions, so, what happens during acidemia? Hydrogen ions start accumulating in the extracellular space. There is then an inward movement of hydrogen ions, and in exchange, there is an outward movement of potassium ions in order to conserve charge.
Some of the intracellular hydrogen ions get buffered and no longer directly contribute to the hydrogen ion concentration, thus the concentration gradient favors more inward movement of hydrogen and outward movement of potassium, while the newly intracellular hydrogen ions can be buffered, there is no analogous buffering system for newly extracellular potassium ions. Thus, the net result is hyperkalemia. In reality, this process doesn’t occur in discrete steps like I just outlined, but rather as a continuous and simultaneous flow.
[Eric Strong] Source: LYBIO.net
In summary, acidemia leads to hyperkalemia as a rough guide for every 0.1 drop in pH. Serum potassium can increase in average of 0.6 milliequivalents per liter. Unfortunately there is a big range seen from patient to patient, which is why potassium levels need to be monitored very closely when correcting severe acid-base disorders. Alkalemia can conversely need to hypokalemia by the reverse process. There is not a well described rule of thumb however in – how to predict the scale of this effect, but it is usually quantitatively smaller than that seen in acidemia.
That concludes this lecture on normal acid-base physiology. In the next lecture, I will discuss how to use the ABG to identify simple and uncomplicated acid-base disorders.
Normal Acid-Base Regulation – Understanding ABGs – Lecture 2. Hello, this is Eric Strong from the Palo Alto Veterans Hospital, and Stanford University. This is the second lecture on Understanding ABGs, and the topic is Normal Acid-Base Regulation. Complete Full Transcript, Dialogue, Remarks, Saying, Quotes, Words And Text.