What are the Physiological Buffers in the Human Body and its importance? The normal hydrogen ion concentration of body fluids is about 40 nmol liter (pH–7.4): during the course of one day, some 60 mmol/liter of hydrogen ion is added to it which, if not buffered, would raise the concentration of the extracellular fluid (12 liters say) to 5 mmol/liter.
This would change the hydrogen-ion concentration by over 10,000 times (to a pH of 2.3) if not buffered in any way.
This does not happen because most of these extra hydrogen ions are taken up by the various physiological buffers found in body fluids.
Ultimately, of course, these excess hydrogen ions and associated bases need to be excreted by the kidneys (making it acid), but initially, they combine with buffers to minimize the change in pH.
Bicarbonate—CO2 buffer: The most important physiological buffers in the body are the bicarbonate–CO2 system, the large anion complexes such as plasma proteins and phosphates and hemoglobin in cells. In all of these, the essential reaction is:
H+ + buffer ⇔ H-buffer
If hydrogen ion increases, then it combines with the buffer, if it decreases, some hydrogen ions are released from the.
H-buffer complex: in this way, the body fluids are stabilized against hydrogen-ion change.
Here we shall be concerned only with the bicarbonate-CO2 system, for it shows some special features which must be discussed: hemoglobin also shows special features.
Quantitatively, the bicarbonate-CO2 system is a very good buffer in the body, yet in a test-tube, the system is not exceptional: there are reasons for this:
pH = 6.1 + log 25/1.34 = 7.4
pKa of the Carbonic acid
The ratio of dissolved HCO3 to CO2 is 25 to 1.34, i.e. nearly 19 to 1, not as shown in the below picture.
Physiological Buffers in the human body
1. Mammalian body fluids
Fluids contains much-dissolved CO2 for they are in equilibrium with alveolar gas which contains 5% CO2 rather than with air which contains practically none. As a buffer, it, therefore, behaves as
H+ + Buffer ⇔ H-buffer ⇔ dissolved CO2
The acceptor of hydrogen ions in the buffer base (HCO–) ) as usual: the 3 donor is the weak acid (H2CO3) which is in equilibrium with the dissolved CO2: as the amount of CO2 dissolved far exceeds the amount of carbonic acid present and the dissolved CO2 can be considered as the proton donor.
The equation then becomes
H+ + HCO3– ⇔ dissolved CO2
Proton acceptor Proton donor
(25 m mol/1) (1.34m mol/1)
From the concentrations given, the pH is 7.4. i.e. that of the body.
2. The bicarbonate
CO2 system by itself can minimize changes in hydrogen-ion concentration, but must always allow some change to occur.
In the body, this limitation is overcome by systems which adjust the base (HCO3–) and the dissolved CO2, to keep them at a constant value; kidney function fixes the concentration of base, lung function that of dissolved CO2 (both the CO2 concentration itself and the hydrogen-ion concentration control respiration).
The whole bicarbonateCO2 system works as follows: If hydrogen ions enter the blood, then they combine with base (HCO3) to form carbonic respiration to exhale more CO2.
You will see that if a large amount of an acid is added to the blood, the quick compensation is an excretion of CO2 with a reduction in the plasma bicarbonate: the slower compensation is that the acid is excreted by the kidneys which replace bicarbonate in the blood.
Read about Bicarbonate Buffer system
Removal of hydrogen ions from blood—as for example, following HCl secretion into the stomach—is compensated for by retaining CO2 and forming more base (in the short term).
In both cases, the respiratory system makes rapid buffering possible, the renal system supplies the long-term buffering.
What do you mean by physiological buffers? What are the different types of buffers found in the blood? If possible, please state examples.
Physiological buffers are chemicals used by the body to prevent large changes in the pH of bodily fluid.
The four Examples of physiological buffers are here
- hemoglobin, and
- protein systems.
The pH of a buffer is determined by the Henderson-Hasselbalch equation:
pH=pKa + log(A−HA)
The buffer is best able to resist changes in pH when the pH of the buffer is close to the pH of blood (7.37 to 7.42), so the pKa of the acid should be close to 7.4.
The phosphate buffer system consists of H2PO−4 and HPO2−4 ions. The equilibrium is
H2PO−4(aq) + H2O ⇌ H3O(aq) + HPO2−4 (aq)
pKa = 7.21
The phosphate buffer can easily maintain a pH of 7.4.
The equilibrium is
H2CO3 (aq) + H2O (l) ⇌ HCO−3(aq) + H3O + (aq)
pKa = 6.1
This buffer functions in exactly the same way as the phosphate buffer, but it is not ideal because its pKa is too far from pH 7.4.
Perhaps more importantly, the enzyme carbonic anhydrase converts H2CO3 into CO2 that is dissolved in the blood and is then exhaled as CO2 gas.
Haemoglobin Buffer system
The general equation is
HHb+ O2 + H2O ⇌ HbO2 + H3O+
pKa = 6.8
It shows that the oxygenation of Hb promotes the formation of H3O+.
This shifts the bicarbonate buffer equilibrium towards CO2 formation, and CO2 is released from the red blood cells.
Proteins: A protein is a long chain of amino acid residues, but this long chain still has free carboxylate groups COO− and free amino groups NH2.
We could write the equation for a protein buffer system as
H3 + N-R-COO− + H2O ⇌ H2N—R—COO− + H3O+
The protein can then act as a buffer.
How do you calculate buffer pH for monoprotic acids?
You use the Henderson-Hasselbalch equation. A monoprotic acid has one ionizable hydrogen.
HA + H2O ⇌ H3O+ + A−
A solution that contains both HA and A− is a buffer solution, and its pH is given by the Henderson-Hasselbalch equation.
Q: Calculate the pH of a solution that contains 0.20 mol/L acetic acid and 0.10 mol/L sodium acetate.
The Ka for acetic acid is 1.74×10−5.
These are examples of physiological buffers in biological systems.
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