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Overview of Physiologic pH and Buffers
Overview of Physiologic pH and Buffers
Overview of Physiologic pH and Buffers

    Physiologic pH, where ‘pH’ stands for ‘potential of hydrogen’, is a way of quantifying the balance between acids and bases in the body. In fact, the pH depends on the concentration of hydrogen ions and can be described with this equation.

    The cells and enzymes in our tissues and organs work best when the concentration of hydrogen ions is 40 x 10^-9 Eq/L, otherwise known as 40 nEq/L. Small changes to that number matter a lot, and because it can get annoying working with such small numbers, scientists converted this concentration into a logarithmic function and expressed it as pH. In this case, a hydrogen ion concentration of 40 x 10^-9 Eq/L, works out to a pH of 7.4. Now, there are two important aspects to remember when using this logarithmic function. First, as hydrogen concentrations increase, the pH decreases, because of the negative sign in front of the log. Second, since it’s a logarithmic function, pH and the hydrogen ion concentration don’t have a linear relationship. For example, an increase in pH from 7.4 to 7.6 means a decrease in the hydrogen concentration of 15 nEq/L. Whereas a decrease in pH from 7.4 to 7.2 means an increase in the hydrogen concentration of 23 nEq/L. That’s why the graph of hydrogen ion concentration versus pH has a curve to it, rather than being a straight line. For simplicity sake, when the body’s pH drops below 7.4 it’s considered acidemia, and when it goes above 7.4, it’s considered alkalemia. So due to this logarithmic relationship, a change in pH in the acidic range, pH < 7.4, will show a larger change in hydrogen concentrations than if the same change occurred in pH in the alkaline range, pH > 7.4.

    Maintaining a pH between 7.37 and 7.42 is essential for the human body. This is accomplished with buffers. In everyday language a "buffer," is something that acts like a protective cushion or shield, and the same is true of physiologic buffers - they shield the pH from rising or falling too quickly. The reason the body needs buffers is that acids - or molecules that readily give up their hydrogen ion - are being generated by the body all the time. These extra hydrogen ions would shift the pH into the acidic zone, so the body needs a way to handle these without having a major shift in the overall pH. To accomplish this, buffers can essentially take on some of these extra hydrogen ions, and therefore keep the pH from dropping too much. buffers are usually a weak acid with its conjugate base form, or a weak base with its conjugate acid form.

    The most important buffer in the body is the weak acid carbonic acid H2CO3 and its conjugate base, bicarbonate ion HCO3-. Carbonic acid forms when carbon dioxide combines with water, with the help of the enzyme carbonic anhydrase. As a weak acid, carbonic acid H2CO3 easily dissociates into hydrogen H+ ions and bicarbonate ions HCO3-. Of course, these reactions are reversible, and can happen in the opposite direction as well.

    In fact, because carbonic acid is such a weak acid, when hydrogen ion concentration gets low, it will drop off it’s hydrogen ion and the equation moves to the right, producing more bicarbonate and a hydrogen ion, and when there are lots of hydrogen ions around, the bicarbonate will bind to one and form carbonic acid, which can go the other way and split into water and carbon dioxide. And the carbon dioxide can be breathed out through the lungs.

    So, imagine you’ve got some extracellular fluid, and it’s a normal amount of H+ ions that puts it in the normal physiologic pH range. Now we drop some NaOH in, or sodium hydroxide, which is a strong base. That means that in water, it completely dissociates into sodium Na+ ions and hydroxide OH- ions. The hydroxide OH- ions bind to hydrogen H+ ions, forming water, making the fluid more basic, since there’s a decrease in hydrogen H+ ions, which increases the pH a lot. But when there’s plenty of CO2 around, it reacts with the water to form more H2CO3 carbonic acid, which splits into bicarbonate ions and hydrogen ions, rapidly replacing the hydrogen H+ ions, and buffering the pH and keeping it in a normal range.

    On the flip side, imagine tossing HCl into the fluid, or hydrochloric acid, which is a strong acid. It would fully dissociate into hydrogen H+ ions and a bunch of chloride Cl- ions. Without a buffer this would cause it to become acidic, since there’s more H+, and therefore the pH would go down. The new hydrogen H+ ions, would get instantly grabbed by bicarbonate ions HCO3- and converted to H2CO3 carbonic acid, which would then dissociate into carbon dioxide CO2 and H2O. Once again, this buffer normalizes the pH!

    The great thing is that there’s a huge supply of carbonic acid because it’s formed from carbon dioxide CO2 and H2O, which are in abundant supply in the body. Finally, if there’s too much carbon dioxide CO2, it is blown off by the lungs, and if there are too many bicarbonate ions HCO3-, they get eliminated by the kidneys.

    Apart from the bicarbonate buffer system, there’s also the phosphate buffer system. Again, that’s a mixture of the weak acid dihydrogen phosphate H2PO4- and its conjugate base monohydrogen phosphate HPO42-. Dihydrogen phosphate has two hydrogens and it’s always ready to give off one of them and get converted to monohydrogen phosphate HPO42- which has a single hydrogen.

    Now, the extracellular fluid and especially, the plasma, is also full of proteins, like albumin, which serve as another extremely important buffer system. That’s because some of the proteins’ amino acids have exposed carboxyl groups —COOH, that act as weak acids, meaning that they are ready to release hydrogen H+ ions when the pH begins to rise while others have exposed —NH2 amine groups that act as weak bases, meaning that they can bind a hydrogen ion, preventing it from decreasing the pH.

    In other words, a single protein molecule can function both as an acid as well as a base depending on the pH it has to deal with. But, unlike the bicarbonate buffering system, there are only a limited number of proteins in our plasma, which limits how much they can buffer.

    Alright, until now we’ve explored the extracellular fluid, but intracellular pH has to stay balanced as well, and this is the pH inside cells. Cells are full of proteins, and they’re the most significant intracellular buffer. An example is hemoglobin within red blood cells. Hemoglobin can reversibly bind either hydrogen ions H+, which binds to the protein itself, or O2 which binds to the iron of the heme group, and the interesting thing is that when one of these is bound, the other is released. So when red blood cells are in capillaries of various tissues, the O2 levels are low and CO2 levels are high. The CO2 diffuses into red blood cells, where carbonic anhydrase enzyme combines H2O with CO2 to form carbonic acid, H2CO3, which then dissociates to H+ and HCO3-. Because there’s so much CO2 around, the reaction continues in the direction of generating more H+ and HCO3-, and over time the buildup of hydrogen ions can cause the pH to fall. To prevent the hydrogen ion concentration from rising too quickly, each deoxygenated hemoglobin binds hydrogen ions. The bicarbonate HCO3- on the other hand, moves into the plasma in exchange for chloride Cl- ions. This exchange, called the chloride shift, keeps positive and negative charges balanced.

    Now, once the red blood cell gets into the capillaries of the lungs, there’s a low level of carbon dioxide and high level of O2, so the process reverses. Oxygen binds to hemoglobin, and hydrogen ions H+ get released. Bicarbonate reenters the red blood cell and combines with the hydrogen ions to form carbonic acid, which dissociates into H2O and CO2, and the CO2 gets breathed out by the lungs.

    Intracellular fluid also has large amounts of organic phosphates, like ATP and glucose-6-phosphate which can serve as intracellular buffers. The phosphate group of these organic molecules can serve as a source or sink for excess hydrogen H+ ions.

    All right, as a quick recap, between 7.37 and 7.42 is physiologic pH. The body uses buffering systems to maintain pH within this range. Buffers are pairs of a weak acid and its conjugate base or a weak base and its conjugate acid and their physiologic role in our body is to resist pH changes. The most important buffers in the extracellular fluid, are bicarbonate, phosphate, and plasma proteins, and the most important buffers within cells are proteins like hemoglobin, and organic phosphates, like ATP.

Physiologic pH and Buffers (Acid-Base Physiology) (https://www.youtube.com/watch?v=6EnIPG3WRRo&index=9&list=PLY33uf2n4e6PT53f0Z5LmFHo7Vb0ljn5b) by Osmosis (https://open.osmosis.org/) is licensed under CC-BY-SA 4.0 (https://creativecommons.org/licenses/by-sa/4.0/).