This document provides an overview of buffers in biological systems. It discusses how biological buffers help maintain pH within a narrow physiological range through reversible acid-base reactions. The key types of biological buffers mentioned are bicarbonate, phosphate, proteins, and hemoglobin. Bicarbonate buffers blood pH through reactions involving carbonic acid and carbon dioxide transport. Phosphate buffers intracellular and renal fluid. Proteins and hemoglobin buffer through amino acid side chains that react with hydrogen ions. Maintaining pH is critical for biochemical processes and homeostasis in living organisms.

1. BUFFERS IN
BIOLOGICAL
SYSTEMS
BY JWALA
1ST MSc MARINE MICROBIOLOGY

2. INTRODUCTION
• A biological buffer is an organic substance that has a
neutralizing effect on hydrogen ions.
• Biological buffers helps maintaining the body at the correct pH
so that biochemical processes continue to run optimally.
• Most buffers consist of a weak acid and a weak base. They
help maintain a given pH even after the addition of an acid or
a base.
• Biological buffers help maintain a steady pH around the
physiological pH.

3. BUFFER
• Buffer is a mixture of weak acid and salt of conjugate base
that resist the change in pH upon the addition of acid or base.
• Buffering results two reversible reaction equilibrium in a
solution wherein the concentration of proton donor and its
conjugate proton acceptor are equal.
• Buffers have both intensive and extensive properties.
• Buffer capacity generally depends on the concentration of
buffer solutions.
• Buffer capacity is a term used to describe the ability of a given
buffer to resist changes in pH on addition of acid and base.

4. How do Buffers Act
• Buffers acts as “Shock absorbers” against sudden changes of
pH by converting injurious strong acids and bases into
harmless weak acid salts.
• If a buffer solution is composed of weak acid HA and its salt
BA, they ionize as follows:-
HA  H+ + A-
BA  B+ + A-

5. Handerson-Hasselbalch Equation
• The pH of a solution containing a weak acid is related to its
acid dissociation constant. The relationship can be stated in
the convenient form of the “ Handerson- Hasselbalch
equation”
pH = Pka+ log10 [A-]
[HA]
• The Handerson- Hasselbalch equation is an expression of
great predictive value in protonic equilibria.

6. BIOLOGICAL BUFFERS
Biological buffers should meet the following criteria:
• Their pKa should reside between 6.0 to 8.0
• They should exhibit high water solubility and minimal solubilty in
organic solvents.
• They should not permeate cell membranes.
• They should not exhibit any toxicity towards cells.
• They should not interfere with any biological process.
• The salt effect should be minimal, however, salts can be added as
required.
• Ionic composition of the medium and temperature should have
minimal effect of buffering capacity.
• Buffers should be stable and resistant to enzymatic degradation.
• Buffers should not absorb either in the visible or in the UV region.

7. Importance of Biological Buffer
• To maintain homeostasis.
• To regulate enzymatic function.
• To control pH in biochemical reaction.

8. Types of Biological Buffers
• Bicarbonate buffer
• Phosphate buffer
• Protein buffer
• Haemoglobin

9. BICARBONATE BUFFER
• The bicarbonate buffer system is an acid- base homeostatic
mechanism involving the balance of carbonic acid,
bicarbonate ion and carbon dioxide in order to maintain pH
in the blood and duodenum, among other tissues, to support
proper metabolic function.
• This process is catalyzed by carbonic anhydrase.
CO2 + H2O H2CO3 HCO3
-+ H+
• As with any buffer system, the pH is balanced by the presence
of both weak acid and its conjugate base so that any excess
acid or base introduced to the system is neutralised.
• Failure of this system to function properly results in acid-base
imbalance, such as acidemia and alkalemia in the blood.

10. The Bicarbonate Buffer

11. Disorders of acid-base balance in respiratory system

12. PHOSPHATE BUFFER
• It plays a major important role in buffering renal tubular fluid
and intracellular fluids.
• The main elements of phosphate buffer system are H2PO4
- and
HPO4
- . When a strong acid such as HCl is added to a mixture
of these two substances, the hydrogen is accepted by the base
HPO4
- and converted to H2PO4
-.
HCl + Na2HPO4 NaH2PO4+ NaCl
• The phosphate buffer system has a pK of 6.8. This allows the
system to operate near its maximum buffering power.
• The phosphate buffer is important in the tubular fluids of the
kidneys.
• The phosphate buffer is also important in buffering
intracellular fluid.

13. PROTEIN BUFFER
• Protein buffer system consists of the plasma proteins ( such as
albumins) and certain proteins in cells, including the
haemoglobin uses freely exposed amino acids groups of an
amino acid or protein; combines with a hydrogen ion in the
presence of excess acid.
• Proteins are made up of amino acids, which contain positively
charged amino groups and negatively charged carboxyl
groups. The charged regions of these molecules can bind
hydrogen and hydroxyl ions, and thus function as buffers.
• Protein buffers interact extensively with other buffer systems.
• They helps in regulating pH in extracellular fluid and
intracellular fluid.

14. Intracellular buffers:- proteins, phosphates and the pottassium-hydrogen ion exchange.

15. HAEMOGLOBIN BUFFER
• Transports oxygen from lungs to peripheral tissues.
• Transports carbon dioxide from tissues to lungs for exhalation.
• Haemoglobin is a buffer for both carbon dioxide and H+.
• Carbon dioxide diffuses across RBC membrane from tissues.
• The CO2 can bind directly with haemoglobin and be released
in the lungs.
• The CO2 that reacts with water forms carbonic acid that then
dissociates into bicarbonate in RBC.
• Bicarbonate ions diffuses into plasma in exchange for chloride
ions.
• H+ binds to haemoglobin and released in RBC’s in lungs to
combine with bicarbonate and reform CO2 for exhalation.

17. Buffering in Cells and Tissues
• Amino acids present in proteins in cells and tissues contain
functional groups that act as weak acid and bases.
• The phosphate and bicarbonate buffer systems are most
predominant in biological systems.
• The bicarbonate buffer system plays an important role in
buffering the blood system where in carbonic acid acts as a
weak acid (proton donor) and bicarbonate act as a conjugate
base (proton acceptor). Their relationship can be expressed as
follows:-
K1 = [H+] [HCO3
-]
[H2CO3]

18. Relationship between bicarbonate buffer system and carbon dioxide

19. Effect of temperature on pH
Generally when we consider the use of buffers we make the
following two assumptions:-
• The activity coefficients of the buffer ions is approximately
equal to 1 over the useful range of buffer concentrations.
• The value of Ka is constant over the working range of
temperature.
• The pH changes slightly with change in temperature. This is
very critical in biological systems where a precise hydrogen ion
cocncentration is required for reaction systems to operate
with maximum efficiency.
• The mathematical relationship of activty and temperature
may be complictaed, the actual change of pKa with
temperature is approximately linear.

20. Effect of Temperature on pH of Phosphate Buffer

21. pKa and ∆pKa∕oC of Selected Buffers

22. Effectsof Buffers on FactorsOther than pH
• Some weak acids and bases may interfere with the reaction
system. For example, citrate and phosphate buffers are not
recommended for systems that are highly calcium-
dependent.
• Phosphate ions in buffers can inhibit the activity of some
enzymes, such as carboxypeptidase, fumeras, carboxylase,
and phosphoglucomutase.
• Trisaminomethane can chelate copper and also acts as a
competitive inhibitor of some enzymes.
• Tris based buffers are not recommended when studying the
metabolic effects of insulin.
• Borate buffers are not suitable for gel electrophoresis of
protien, they can cause spreading of the zones if polyols are
present in the medium.

23. Effect of pH on Solubility
• Changes in pH can affect the solubility of partially soluble ionic
compounds.
• As a result of the common ion effect, the solubility of Mg(OH)2
can be increased or decreased. When a base is added to the
concentration of OH- increases and shifts the solubilty
equillibrium to the left causing a dimunition in the solubility of
Mg(OH)2. . When an acid is added to the solution, it neutralises
OH- and shifts the solubility equilibrium to the right. This
results in increased dissolution of Mg(OH)2.

24. THANK YOU