Buffering and Henderson-Hasselbalch equation

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Buffering and Henderson-Hasselbalch equation

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Buffering and Henderson-Hasselbalch equation

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The for a buffer is the ratio of the concentration of H+ multiplied by the concentration of A- , divided by the concentration of HA. 

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Content Reviewers:

Rishi Desai, MD, MPH

Every single moment, there are trillions of biochemical reactions occurring throughout the human body that are mediated by enzymes. Enzymes are types of proteins, and they’re generally sensitive to even slight changes in the environment - in particular things like the hydrogen ion concentration.

For this reason, the blood pH which corresponds to the hydrogen ion concentration needs to stay in a very narrow range---between 7.37 and 7.42.

If the blood pH rises or falls by more than a few tenths of a unit, it can lead to death.

Now, acids and bases are generated by cells all the time. So, the body has a few mechanisms to deal with these molecules and keep blood pH within normal range.

The first scientist who studied one of these mechanisms was Robert Pitts.

Pitts injected 150 mEq of hydrochloric acid HCl into his dog.

He calculated this his dog’s body contained a total 11.4 liters of water, so separately, Pitts put 150mEq of hydrochloric acid HCl in a volume of 11.4 liters of water.

The dog’s blood pH dropped from 7.44 to 7.14, which is very low, but not fatal.

In the water, the pH dropped from 7 to 1.84, and that would have killed the dog instantly.

Based on this, Pitts concluded that his dog had a buffer contained in its body fluids, and the dog concluded that he could no longer trust Pitts to take care of him.

Physiologic buffers shield the pH from rising or falling too 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 weak acid could be symbolized as HA, where A represents molecules like fluorine or acetate. And the fact that it’s weak means that it has a “weak” effect on pH, because it doesn’t fully dissociate in water.

For example, if we were to add 100 HA molecules in 1 ml of water, only a tiny fraction, let’s say 5 of the 100 HA molecules would dissociate or break down into hydrogen H+ and their conjugate base A-: 10 HA ⇄ 9 HA + 1 H+ + 1 A-.

They like to maintain this balance. So if we remove some hydrogen ions, HA will dissociate and release a hydrogen ion and A-, maintaining that equilibrium. And if we add hydrogen ions, the A- will bind to a hydrogen ion to form HA.

Since it maintains this equilibrium, we call it an equilibrium reaction, and HA ⇄ H+ + A-.

By definition, this equilibrium reaction can move forward or backward depending on the concentrations of the molecules, and this is known as the Le Chatelier’s principle.

So, let’s imagine you’ve got this tank of water, which at any one time has some hydrogen ions floating around as well as some hydroxide ions that are in equilibrium with the water. Imagine that we add a strong base like sodium hydroxide NaOH to water.

Since it’s a strong base, the sodium hydroxide would dissociate almost completely into sodium Na+ ions and hydroxide OH- ions, and the hydroxide ions would bind to any available hydrogen H+ ions, and this rapid loss of hydrogen increases the pH really fast.

Sources
  1. "Medical Physiology" Elsevier (2016)
  2. "Physiology" Elsevier (2017)
  3. "Human Anatomy & Physiology" Pearson (2018)
  4. "Principles of Anatomy and Physiology" Wiley (2014)
  5. "Understanding Acid Base Disorders" Critical Care Clinics (2015)
  6. "Acid-Base Assessment" Veterinary Clinics of North America: Food Animal Practice (2014)
  7. "pH and the Henderson-Hasselbalch equation" The American Journal of Medicine (1973)