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Acid-base map and compensatory mechanisms
Buffering and Henderson-Hasselbalch equation
Physiologic pH and buffers
The role of the kidney in acid-base balance
Plasma anion gap
Renal system anatomy and physiology
Body fluid compartments
Movement of water between body compartments
Measuring renal plasma flow and renal blood flow
Regulation of renal blood flow
TF/Px ratio and TF/Pinulin
Phosphate, calcium and magnesium homeostasis
Free water clearance
Kidney countercurrent multiplication
Distal convoluted tubule
Loop of Henle
Proximal convoluted tubule
Tubular reabsorption and secretion
Tubular reabsorption and secretion of weak acids and bases
Tubular reabsorption of glucose
Tubular secretion of PAH
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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.
The Henderson-Hasselbalch equation is a mathematical expression that is used to predict how much of a given acid or base is required to produce a desired pH in a given solution. The equation is named after British chemists Louis Hodgkin and Frederick Gowland Hopkins, who developed it in 1898, and German chemist Wilhelm Hasselbalch, who published it in 1909.
The Henderson-Hasselbalch equation can be expressed as follows:
pH = pKa + log [base]/[acid]
The pH is the desired pH of the solution, pKa is the dissociation constant of the acid, [base] is the concentration of the base, and [acid] is the concentration of the acid.
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