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Physiologic pH and buffers
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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.
The human body has a narrow range of pH that is considered physiologically normal, which is between 7.37 and 7.42. The body uses buffers to maintain pH within this range. Buffers are substances that help maintain the pH of a solution within a specific range. They work by absorbing or releasing hydrogen ions (H+) to counteract changes in pH. ate, phosphate, and plasma proteins, and the most important buffers within cells are proteins like hemoglobin, and organic phosphates, like ATP.
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