Buffering Agent

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 living 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 a chemical equilibrium between the weak acid HA and its conjugate base 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 Chatelier'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. In figure 1, the effect is illustrated by the simulated titration of a weak acid with p''K''_a = 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 = p''K''_a ± 1, centered at pH = 4.7, where Anbsp;=  ⁻ The hydrogen ion concentration decreases by less than the amount expected because most of the added hydroxide ion is consumed in the reaction

and only a little is consumed in the neutralization reaction (which is the reaction that results in an increase in pH)

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

Buffer capacity

Buffer capacity is a quantitative measure of the resistance to change of pH of a solution containing a buffering agent with respect to a change of acid or alkali concentration. It can be defined as follows:
$\beta = \frac,$
where $dC_b$ is an infinitesimal amount of added base, or
$\beta = -\frac,$
where $dC_a$ is an infinitesimal amount of added acid. pH is defined as −log₁₀ ⁺ and ''d''(pH) is an infinitesimal change in pH.

With either definition the buffer capacity for a weak acid HA with dissociation constant ''K''_a can be expressed as
\beta = 2.303 \left( ce+ \frac + \frac\right),
where ⁺is the concentration of hydrogen ions, and $T_\text$ is the total concentration of added acid. ''K''_w is the equilibrium constant for self-ionization of water, equal to 1.0. Note that in solution H⁺ exists as the hydronium ion H₃O⁺, and further aquation of the hydronium ion has negligible effect on the dissociation equilibrium, except at very high acid concentration.

This equation shows that there are three regions of raised buffer capacity (see figure 2).
* In the central region of the curve (coloured green on the plot), the second term is dominant, and $\beta \approx 2.303 \frac.$ Buffer capacity rises to a local maximum at pH = ''pK''_a. The height of this peak depends on the value of pK_a. Buffer capacity is negligible when the concentration Aof buffering agent is very small and increases with increasing concentration of the buffering agent. Some authors show only this region in graphs of buffer capacity. Buffer capacity falls to 33% of the maximum value at pH = p''K''_a ± 1, to 10% at pH = p''K''_a ± 1.5 and to 1% at pH = p''K''_a ± 2. For this reason the most useful range is approximately p''K''_a ± 1. When choosing a buffer for use at a specific pH, it should have a p''K''_a value as close as possible to that pH.
* With strongly acidic solutions, pH less than about 2 (coloured red on the plot), the first term in the equation dominates, and buffer capacity rises exponentially with decreasing pH: $\beta \approx 10^.$ This results from the fact that the second and third terms become negligible at very low pH. This term is independent of the presence or absence of a buffering agent.
* With strongly alkaline solutions, pH more than about 12 (coloured blue on the plot), the third term in the equation dominates, and buffer capacity rises exponentially with increasing pH: $\beta \approx 10^.$ This results from the fact that the first and second terms become negligible at very high pH. This term is also independent of the presence or absence of a buffering agent.

Applications of buffers

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 (HCO) and bicarbonate (HCO) 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. 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 colouring fabrics. They are also used in chemical analysis and calibration of pH meters.

Simple buffering agents

:
For buffers in acid regions, the pH may be adjusted to a desired value by adding a strong acid such as hydrochloric acid to the particular buffering agent. For alkaline buffers, a strong base such as sodium hydroxide may be added. Alternatively, a buffer mixture can be made from a mixture of an acid and its conjugate base. For example, an acetate buffer can be made from a mixture of acetic acid and sodium acetate. Similarly, an alkaline buffer can be made from a mixture of the base and its conjugate acid.

"Universal" buffer mixtures

By combining substances with p''K''_a values differing by only two or less and adjusting the pH, a wide range of buffers can be obtained. Citric acid is a useful component of a buffer mixture because it has three p''K''_a values, separated by less than two. The buffer range can be extended by adding other buffering agents. The following mixtures (