Table of Contents
Introduction
What is buffering
Buffering is the process by which significant changes in [H+], the concentration of hydrogen ions, are minimized and kept within the normal range, despite the addition or removal of hydrogen ions.
What is a buffer
A buffer is a substance capable, in solution, of neutralizing both acids and bases and so maintaining pH within the normal range.
A buffer can receive and release H+ in a solution to keep the pH of the solution in which it is present, despite the addition of a considerable amount of acid or base.
A buffer in a solution consists of a weak acid and its conjugate base.
When an acid is added to a body’s solution, it tends to increase the solution’s acidity. Still, the body’s buffer system becomes active and reduces it to maintain it within the normal range.
Likewise, when an alkali is added to the body’s solution, it tends to decrease acidity. Still, the body’s buffer system becomes active, increasing acidity to maintain it within the normal range.
The acidity of a solution is expressed on the pH scale. It is from 0 to 14. When it is 7, the solution is neutral; below 7, it is acidic; above 7, it
0 14
Normal range
The normal H ion concentration in the ECF is 40 nmol/L.
The normal acid-base ratio is 1: 20
PH is an indicator of acidity, and the negative logarithm of H.
An increase in pH by one fold requires a 10-fold decrease in H ion concentration,
pH of blood is 7.35 to 7.45. This is essential for life, as this range is vital for proper cellular function.
When it is below 7.35, the condition is known as acidosis.
When it is above 7.45, the condition is known as alkalosis.
Both these conditions are serious.
Although the pH of certain body fluids is different, for example, that of urine is about 4.5 to 8, and that of gastric juice is up to 1.5.
General equation for a buffer system
Weak undissociated acid==hydrogen ion + anion (negative charge) HA= H +A
If a stronger acid than HA is added to a solution containing this buffer system, the sequence of events is
A stronger acid will dissociate into H ions rapidly and provide more H ions, which combine with A of the weak acid and reduce the H ion concentration in the solution, restoring the pH of the solution. So, the increase in H+ concentration is much smaller than it would otherwise be.
Now consider the addition of a strong alkali to the solution; more HA will dissociate to offer H ions that will react with OH of the alkali to form H2O. So, the pH is maintained.
The most effective buffers in the body are expected to be those with pKs close to the pH at which they operate.
The pK of the bicarbonate system is 6.1, that of the diabase system is 6.8, and that of the ammonium system is 9.
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