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Summary Control of body pH
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  • Course
  • Physiology
  • Institution
  • Humanitas University

Delve deep into the fundamental principles of acid-base balance and explore the intricate mechanisms by which the body maintains optimal pH levels in blood and tissues. Discover the pivotal roles played by both respiratory and metabolic processes in pH regulation, unraveling the complex interpla...

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Document information

  • April 3, 2024
  • 11
  • 2022/2023
  • Summary

Subjects

  • acids and bases
  • ph buffers
  • bicarbonate
  • carbon dioxide
  • buffers
  • blood curves
  • henderson hasselbalch equation
  • a and b intercalated cells
  • metabolic acidosis and alkalosis
  • respiratory alkalosis and acidosis

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  • Humanitas University
  • Physiology
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Enrico Tiepolo


Control of body pH
Molecules that can release or bind a proton behave as acids and bases.
- An acid is a molecule that can release a proton, becoming its own correlated base (alkali), which can
instead accept a proton.
- A substance is referred to as an acid if its native, uncharged form is the acidic one (i.e., its native
form in water tends to release protons and lower the pH, e.g., H2CO3 à HCO3- + H+), while we
refer to a substance as a base if its native form in water tends to bind protons (e.g., NH3 + H+ à
NH4+).
Weak acids and bases work as pH buffers because in water they tend to be partially dissociated
and therefore they can both bind or release a certain quantity of protons thus buffering the changes in
H+ concentration. Thus, a substance displays its maximal buffering capacity when it is in a
solution that has pH equal to its own pKA, because in that situation half the molecules are in
acid form and half in base form.



Apart from any computational aspect, the pH of the blood depends on the ratio between
bicarbonate and carbon dioxide in the plasma. If this ratio is maintained constant, the pH is
constant.
- The pH is defined as the negative logarithm of the H+ concentration. It indicates whether a system
is acidic or alkaline and therefore gives details about the concentration of H+ in the system.


- The pKa is defined as negative logarithm of the dissociation constant of a solution (Ka). The smaller
the value of pKa, the stronger the acid, because the smaller the tendency to dissociate. So it indicates
whether an acid is strong or weak and gives details about the dissociation of an acid in a solution.



The intracellular buffer
The intracellular pH is buffered by a relevant concentration of phosphate (a multiple buffer as it can
exist in 4 forms: ), in addition to sulphate, amine and COOH lateral
groups of proteins, and nucleic acids.
The main reaction is from biphosphate to triphosphate, as it is the reaction with a pKa most similar to
the intracellular pH, and only a minor fraction will be in the form of PO43− or H3PO4.

Intracellularly, in addition to phosphate, the cell is full of proteins with a lot of carboxy and amino groups
which may behave as weak acids and weak bases, but also nucleic acids (both in the nucleus and in the
cytosol). There are therefore various possible substrates that contribute to regulating the situation.

The extracellular buffer
Conversely, the main extracellular buffer is bicarbonate. The system bicarbonate/carbonic acid does
not appear to be optimal, for the extracellular pH=7.4, as the pKA is 3.5, which means that at pH
7.4 the fraction in acid form, carbonic acid, will be as low as ≈ 10−4, i.e. the capacity of the buffer to
release protons is essentially null.
However, the system is an efficient buffer because it is an open system: carbon dioxide is
always present in the organism, as a by-product of cellular respiration, so the pool of carbonic acid is
easily refilled as necessary; dissolved CO2 is in equilibrium with carbonic acid in a ratio 400:1; so, if one
considers CO2 as the acid, in place of carbonic acid, the pKA of the system becomes 6.1.




138 Body At Work II

, Enrico Tiepolo




CO2 can be added very easily, and even if all CO2 we have was used to produce protons, it would
continuously be produced thanks to the metabolic activity of cells (glucose is burnt in water and carbon
dioxide). Also, in the other direction, the system is quick in buffering if protons have to be absorbed, as
it is sufficient to breathe a lot to eliminate a lot of CO2, driving the reaction from left to right and
“sucking” the protons by expelling them in the form of CO2.


The reaction above, as already said, has an open end (CO2 can continuously be produced and
eliminated) so that the reaction from right to left can proceed indefinitely. This is not so true in the other
direction, as the system cannot eliminate or reabsorb bicarbonate with the same efficiency.
Bicarbonate can be eliminated by reducing its reabsorption in the kidneys and can be produced by
intercalated cells type A: the system is open on the right too, but the kidney is not as efficient as the lungs
(it can produce enormous amounts of CO2 by reducing the respiratory rate or keeping our breath, while
bicarbonate can be synthesized and eliminated too but over a longer time).
Therefore, in case of an abrupt change in pH, the organism will very rapidly react by regulating the left
side of the system (increasing or decreasing the respiration: the more we breathe the more protons we
consume, the less we breathe the more we accumulate CO2 and we accumulate protons): rapid
changes in pH are countered by regulating the respiratory rate.

Sooner or later, however, the organism must equilibrate. The pH depends on the ratio
bicarbonate/CO2, and if CO2 has risen to push the reaction towards the right, bicarbonate too will
have to be increased to bring back the normal ratio of the two. Vice versa, if CO2 has been lost to change
the pH, some bicarbonate will have to be eliminated to reach the correct balance once again, which will
be achieved when the production/elimination of bicarbonate will have changed adequately. Therefore,
the long-term equilibrium of pH depends on the kidney, while the rapid compensation on
the pulmonary system.

The concentration of a gas in blood depends on its solubility and on its partial pressure; in particular,
solubility is expressed in terms of mmol/l/mmHg.
The solubility of carbon dioxide in water is about 0.03, so the above relations can be




Important parameters:
The organism has elaborated mechanisms that maintain the hematic concentration of H+ between
an interval that typically goes from 40 to 45 nEq/L (from 37 to 43 nmol/L) with a pH of 7,35-7,45.
Ideally, H+ is 40 nEq/L (40 nmol/L) and pH = 7,40 (we usually say 7.4 but it changes a little between
arterial and venous, of about 0.3).
Typically, arterial blood is 7.4 and venous blood 7.37. It is still acceptable if the pH goes down to 7.35
or up to 7.45, but the range is very narrow.




139 Body At Work II

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