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Chapter: Biochemistry: Living Cell


A buffer is a mixture of a weak acid and its salt with a strong base (eg) A mixture of acetic acid and sodium acetate.


A buffer is a mixture of a weak acid and its salt with a strong base (eg) A mixture of acetic acid and sodium acetate.

HAC + NaAC -----> Na+ +H+ + 2AC-

where HAC = Acetic acid; NaAC = Sodium acetate.

A buffer solution is one which resists a change in its pH value (hydrogen ion concentration) on dilution or on addition of an acid or alkali. The process by which added H+ and OH- ions are removed so that pH remains constant is known as buffer action.

For (eg) if alkali is added to the above mentioned buffer it forms NaAC and no free H+ or OH- will be available.


[Na+ + H+ + 2AC-] + NaOH ------> 2NaAC + H2O

If an acid is added to the buffer it will form NaCl and no free H+  or OH-  will be available.

[Na+ + H+ + 2AC-] + HCl --------> NaCl + 2 HAC

In either cases there is no change in hydrogen ion concentration i.e. it resist the change in pH of the solution


Examples of buffer

A mixture of

·           Glycine and HCl

·           Potassium dihydrogen phosphate and dipotassium hydrogen phosphate.

·           Sodium bicarbonate and sodium carbonate.


Uses of buffer

·           Bufers are used for preparing standard solutions in which it is always desired to maintain a constant pH.

·           These are used to maintain H+ concentration which is necessary for optimal activity of enzymes.

·           Buffers regulate acid-base balance by restricting pH change in body fluids and tissues, since they are capable of absorbing H+ ions and OH- ions when an acid or an alkali is formed during metabolic activities.


Buffers of blood

The important buffers present in blood are

·           Bicarbonate buffer

·           Phosphate buffer

·           Protein buffer

·           Hemoglobin buffer

Bicarbonate buffer

It is the most important buffer in blood plasma and consist of bicarbonate [HCO3-] and carbonic acid [H2CO3] This buffer is efficient in maintaining the pH of blood plasma to 7.4 against the acids produced in tissue metabolism (eg) phosphoric acid, lactic acid, aceto acetic acid and b-hydroxy butyric acid. These acids are converted to their anions and the bicarbonate is converted to carbonic acid a weak acid.

CO2 thus formed is expirated through lungs during respiration.

Phosphate buffer

The phosphate buffer consists of dibasic phosphate [HPO42-] and monobasic phosphate (H2PO4-). Its pKa value is about 6.8. It is more effective in the pH range 5.8 to 7.8. Plasma has a ratio of 4 between [HPO42-] : [H2PO4-].

Therefore pH = pKa + log { [HPO42-] /[H2PO4-] }

pH = 6.8 + log 4 = 7.4                                                                                           [7.4 is the normal pH of blood]

Protein buffer

The protein buffers are very important in the plasma and in the intracellular fluids but their concentration is very low in CSF, lymph and interstitial fluids.

They exist as anions serving as conjugate bases (Pr-) at the blood pH 7.4 and form conjugate acids (HPr) accepting H+. They have the capacity to buffer some H2CO3 in the blood.

H2CO3 + Pr-  -- >  HCO3- + HPr  - - > H2CO3 - - > CO2

Hemoglobin buffer

They are involved in buffering CO2 inside erythrocytes. The buffering capacity of hemoglobin depends on its oxygenation and deoxygenation. Inside the erythrocytes, CO2 combines with H2O to form H2CO3 under the action of carbonic anhydrase. At the blood pH 7.4, H2CO3 dissociates into H+ and HCO3- and needs immediate buffering. Oxyhemoglobin (HbO2-) on the other side loses O2 to form deoxyhemoglobin (Hb-) which remains undissociated (HHb) by accepting H+ from the ionization of H2CO3. Thus, Hb- buffers H2CO3 in erythrocytes.


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