Physiologic pH and buffers
AssessmentsPhysiologic pH and buffers
Physiologic pH and buffers
Physiologic pH and buffers exam links
Content Reviewers:Rishi Desai, MD, MPH
Contributors:Tanner Marshall, MS, Evan Debevec-McKenney, Antonia Syrnioti, MD
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 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 nEq/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 to quickly.
The reason the body needs buffers is that acids - molecules that readily give up their hydrogen ion - are being generated by the body all the time. So the body needs a way to handle the extra hydrogen ions that are released without having a major shift in the overall pH.
To accomplish this, 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-.
CO2 + H2O H2CO3 H+ + HCO3-
In fact, because carbonic acid is such a weak acid, when there are few hydrogen ions around, it will drop off its hydrogen ion and turn into bicarbonate, 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 that the extracellular fluid is “attacked” by sodium hydroxide NaOH, 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, and the decrease in hydrogen H+ ions 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.
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.
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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- "Physiology" Elsevier (2017)
- "Human Anatomy & Physiology" Pearson (2018)
- "Principles of Anatomy and Physiology" Wiley (2014)
- "Acid-Base Homeostasis" Clinical Journal of the American Society of Nephrology (2015)
- "Role of Acid-Base Homeostasis in Diabetic Kidney Disease" Current Diabetes Reports (2017)
- "Electrolyte and Acid-Base Disorders in Chronic Kidney Disease and End-Stage Kidney Failure" Blood Purification (2017)