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Buffer solution:

A buffer solution (more precisely, pH buffer or hydrogen ion buffer) is an aqueous solution consisting of a mixture of a weak acid and its conjugate base, or vice versa. Its pH changes very little when a small amount of strong acid or base is added to it. Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications. In nature, there are many systems that use buffering for pH regulation. For example, the bicarbonate buffering system is used to regulate the pH of blood, and bicarbonate also acts as a buffer in the ocean.

Principles of buffering:

Buffer solutions resist pH change because of an equilibrium between the weak acid HA and its conjugate base A−: HA ⇌ H+ + A−

When some strong acid is added to an equilibrium mixture of the weak acid and its conjugate base, hydrogen ions (H+) are added, and the equilibrium is shifted to the left, in accordance with Le Châtelier's principle. Because of this, the hydrogen ion concentration increases by less than the amount expected for the quantity of strong acid added. Similarly, if strong alkali is added to the mixture, the hydrogen ion concentration decreases by less than the amount expected for the quantity of alkali added. The effect is illustrated by the simulated titration of a weak acid with pKa = 4.7. The relative concentration of undissociated acid is shown in blue, and of its conjugate base in red. The pH changes relatively slowly in the buffer region, pH = pKa ± 1, centered at pH = 4.7, where [HA] = [A−]. The hydrogen ion concentration decreases by less than the amount expected because most of the added hydroxide ion is consumed in the reaction

OH− + HA → H2O + A−

And only a little is consumed in the neutralization reaction (which is the reaction that results in an increase in pH) OH− + H+ → H2O.

Once the acid is more than 95% deprotonated, the pH rises rapidly because most of the added alkali is consumed in the neutralization reaction.

Applications:

The pH of a solution containing a buffering agent can only vary within a narrow range, regardless of what else may be present in the solution. In biological systems this is an essential condition for enzymes to function correctly. For example, in human blood a mixture of carbonic acid (H2CO3) and bicarbonate (HCO−3) Is present in the plasma fraction; this constitutes the major mechanism for maintaining the pH of blood between 7.35 and 7.45. Outside this narrow range (7.40 ± 0.05 pH unit), acidosis and alkalosis metabolic conditions rapidly develop, ultimately leading to death if the correct buffering capacity is not rapidly restored.

If the pH value of a solution rises or falls too much, the effectiveness of an enzyme decreases in a process, known as denaturation, which is usually irreversible.[5] The majority of biological samples that are used in research are kept in a buffer solution, often phosphate buffered saline (PBS) at pH 7.4.

In industry, buffering agents are used in fermentation processes and in setting the correct conditions for dyes used in coloring fabrics. They are also used in chemical analysis[4] and calibration of pH meters.

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