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Case 4
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Case 4Learning goals:
1. What is the function of a buffer and what is it?
2. What is osmolarity? (PBS)
3. Different kinds of buffers
4. Henderson Hasselbach equation
5. Acids and bases
6. Selectively permeable membranes
, 1. Buffers
Buffer = chemical systems which regulate the pH/homeostasis of the
acid-base balance by taking up or release H+-ions.
- pH high: release of H+ ions
- pH low: take up of H+ ions
Weak acids have a smaller effect on pH.
BUT, they are efficient at preventing pH changes and this allows them
to play an important role in chemical buffer systems.
A buffer is most effective when the pH is similar to the pKa (pH = 7,4). It
can range 1. So between 6,4 and 8,4.
It’s important to keep the good pH. Because otherwise, especially,
proteins don’t work anymore. (optimum, denaturate)
Two types of buffers
1. Acidic buffer: weak acid + strong base
2. Basic buffer: weak base + strong acid
Isoelectric point = the pH of an amino acid peptide at which it has a net
charge of zero.
2. Different kinds of buffers
The three major chemical buffer systems in the body:
1. The bicarbonate (HCO3-) buffer system
- Is only active in the ECF, not in the cells themselves.
- Carbonic acid (H2CO3) is a weak acid and is therefore in equilibrium
(evenwicht) with HCO3- in solution (aq). When significant amounts of both
H2CO3 and HCO3-are present, a buffer is formed. This buffer system can be
written as:
CO2 + H2O H2CO3 HCO3- + H+
CO2 is excreted by the lungs. HCO3- and the is excreted by urine
If too many H+ ions build up, the equation above will be pushed to
the right, and HCO3- -ions will absorb the H+ to form H2CO3. Similarly,
if H+ concentrations drop too low, the equation will be pulled the left
and H2CO3 will turn into HCO3-, donating H+ ions to the solution.
- Under normal circumstances there is much more HCO3- present
than H2CO3 , (ca. 20:1). As normal metabolism produces more acids than
bases, this consistent with the body’s need.
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