Buffers

Buffers are solutions that work towards maintaining a constant pH in a system regardless of the addition of strong acids or bases. In the absence of buffers, the addition of a strong acid or a base to a solution changes the pH significantly. Buffers are crucial in everyday life. For example, blood acts as a buffer and is able to maintain a pH between 7.35 and 7.45 since a pH above 7.8 or below 6.8 can cause death.

So, how does a buffer work? A buffer is essentially a weak acid or base and its conjugate base or acid in equilibrium with each other. Take, for example, a buffer made of acetic acid and sodium acetate. The buffer solution contains a weak acid and its conjugate salt, which dissociates to form the conjugate base acetate. The solution is acidic because the K_a of acetic acid is higher than the K_b of the conjugate base.

Remember that the dissociation constant, K_a or K_b, defines the strength of the acid or base, respectively, as it is the equilibrium constant for the dissociation of the compound in water. The dissociation constant can also be represented as pK_a, or the negative log of K_a, where the smaller the pK_a, the stronger the acid. The same follows for pK_b, where the smaller the pK_b, the stronger the base.

Returning to the acetate buffer, if we add a strong base, like sodium hydroxide, the hydroxide ions react with the hydronium ions present in the solution. This causes more acetic acid to react with water in order to return to equilibrium, thus forming more hydronium ions. If we add a strong acid, like hydrochloric acid, more hydronium ions are formed in solution. Those ions react with the acetate ions to form more acetic acid. In both cases, there is very little change in the solution pH.

Keep in mind that each buffer has a specific pH range where it can buffer a solution. Here we show a titration curve for our acetic acid buffer. The buffering range is highlighted. This region is determined by the Henderson-Hasselbalch equation, which tells us that the pH of the buffer is determined by the K_a of the weak acid and the ratio of the conjugate base to the weak acid.

Far away from this pH, the weak acid or base of the buffer is depleted, and it cannot buffer pH. So, when selecting a buffer for your application, choose one whose pK_a is close to the pH desired.

Finally, all buffers have a limit to their buffering capacity, meaning the amount of acid or a base that can be added to the buffer solution before the pH changes significantly. For example, if we add too much acid to our acetic acid and sodium acetate buffer solution, we will protonate all of the acetate ions and accumulate a lot of hydronium ions, thereby lowering the pH.

Similarly, if we add too much base, we will deprotonate all of the acetic acid and accumulate a lot of hydroxide ions, thereby increasing the pH. Thus, the buffering capacity is influenced by the concentration of the weak acid and conjugate salt. So, a 1 molar acetate buffer has a higher buffering capacity than a 0.1 molar acetate buffer.

In this lab, you'll prepare and examine buffers with a wide variety of pH ranges. You'll then use your buffers to determine the pK_a of a pH indicator, called neutral red, and explore the change in neutral red's pK_a when riboflavin binding protein is introduced.