Uses Of Buffer Solutions

When a modest amount of strong acid or base is added to it, the pH of the solution changes very little. When it comes to a wide range of chemical applications, buffer solutions are employed to maintain the pH at a practically constant value. There are numerous pH-regulating mechanisms in nature that rely on buffering to function properly. Bicarbonate buffering is used to control the pH of blood, and bicarbonate also works as a buffer in the ocean, as previously mentioned.

Titration of an acidified solution of a weak acid (pKa = 4.7) with an alkali in a simulated environment

Because of an equilibrium between the weak acid HA and its conjugate base A, buffer solutions are resistant to pH change.

Following the application of Le Chatelier’s principle, when a strong acid is added to an equilibrium mixture of a weak acid and its conjugate base, the addition of hydrogen ions (H+) causes the equilibrium to move to the left. Consequently, the hydrogen ion concentration increases by a smaller amount than would be predicted given the amount of strong acid that has been injected. As with strong alkali, when strong alkali is added to the mixture, the hydrogen ion concentration reduces by a smaller percentage than would be predicted given the amount of alkali added. As seen in Figure 1, the impact is demonstrated by a simulated titration of a weak acid with a pKa = 4.7 concentration. The relative concentration of undissociated acid is depicted in blue, and the relative concentration of its conjugate base is depicted in red. Throughout the buffer zone, pH = pKa 1 is maintained, with the pH = 4.7 being the center of the range. Here, [HA] = [A] is maintained at a reasonably constant rate. Because the majority of the additional hydroxide ion is consumed in the reaction, the hydrogen ion concentration reduces by less than the amount predicted by the equation.

Once the acid has been deprotonated to greater than 95 percent, the pH rises rapidly due to the fact that the majority of the additional alkali is spent in the neutralization reaction.

Capacity for buffering

Buffer capacity is a quantitative measure of the resistance of a solution containing a buffering agent to change in pH when the concentration of an acid or alkali is changed in relation to the concentration of the buffering agent.

The buffer capacity increases until it reaches a local maximum at pH = pKa. In this case, the value of pKa determines the height of the peak. When the concentration [HA] of the buffering agent is very low, the buffer capacity is minimal; nevertheless, as the concentration of the buffering agent grows, the buffer capacity increases. Some authors just display this region in graphs of buffer capacity, while others show the entire graph.

At pH = pKa 1, buffer capacity drops to 33 percent of its highest value, 10 percent at pH = pKa 1.5, and 1 percent at pH = pKa 2. At pH = pKa 2, buffer capacity drops to 1 percent. In order to make use of this range, pKa should be about equal to 1. When selecting a buffer for usage at a given pH, it is important to select one with a pKa value that is as close as feasible to the pH being used. 

This is due to the fact that the second and third terms become trivial at extremely low pH values, as previously stated. This word is not affected by the presence or absence of a buffering agent in the solution.

When dealing with strongly alkaline solutions with pH values more than roughly 12 (seen in blue on the plot), the third element in the equation takes precedence, and buffer capacity increases exponentially as pH increases. This is due to the fact that the first and second terms become negligible at extremely high pH values, as previously stated. The presence or absence of a buffering agent has no effect on the meaning of this phrase.

Buffers have a variety of uses

1. Because of the presence of a buffering agent, the pH of a solution can only vary within a narrow range regardless of the other components of the solution. In biological systems, this is a necessary state for enzymes to perform their functions properly. For example, in human blood, a combination of carbonic acid (HCO₃) and other acids can be found.
2. It is present in the plasma fraction, and it is this that is responsible for the majority of the pH regulation in the blood, which occurs between 7.35 and 7.45. Acute acidosis and alkalosis metabolic conditions occur rapidly outside of this restricted pH range (7.40 0.05 pH unit), and if the correct buffering capacity is not quickly restored, the patient will die.
3. When the pH value of a solution fluctuates excessively, the efficacy of an enzyme declines, a process known as denaturation occurs. Denaturation is normally irreversible, but it can occur in some cases.
4. For the most part, biological samples that are utilized in scientific research are maintained in a buffer solution, most commonly phosphate buffered saline (PBS) at pH 7.4.
5. In the manufacturing business, buffering agents are used in fermentation processes as well as in the establishment of the proper conditions for dyes used in the coloring of fabrics. Their applications include chemical analysis[4] and the calibration of pH meters, amongst others.

Agents that are simple to use as buffers

When using buffers in acidic environments, the pH can be changed to a desirable value by mixing the buffering agent with a strong acid such as hydrochloric acid. Strong bases such as sodium hydroxide can be used to adjust the pH of alkaline buffers. Alternatively, a buffer mixture can be created by combining an acid and its corresponding base in equal parts. For example, an acetate buffer can be created by mixing acetic acid and sodium acetate together in equal parts. Similarly, an alkaline buffer can be created by mixing a base with its corresponding acid in the same proportion.

Conclusion

The Carmody buffer and the Britton–Robinson buffer, both of which were invented in 1931, are examples of universal buffers. A wide range of buffers can be created by combining compounds with pKa values that differ by only two or fewer units and altering the pH of the mixture. Due to the fact that it has three pH values that are separated by less than two, citric acid is a helpful component of a buffer combination. Additional buffering agents can be used to broaden the range of the buffer range. These solutions (McIlvaine’s buffer solutions) have a pH range of 3 to 8, which is within the buffer range.