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Buffer Solutions
Buffer solutions are solutions that resist change in hydronium ion and the hydroxide ion concentration (and consequently pH) upon addition of small amounts of acid or base, or upon dilution. Buffer solutions consist of a weak acid and its conjugate base (more common) or a weak base and its conjugate acid (less common). The resistive action is the result of the equilibrium between the weak acid (HA) and its conjugate base (A−):
HA(aq) + H2O(l) ↔ H3O+(aq) + A−(aq)
Any alkali added to the solution is consumed by the hydronium ions. These ions are mostly regenerated as the equilibrium moves to the right and some of the acid dissociates into hydronium ions and the conjugate base. If a strong acid is added, the conjugate base is protonated, and the pH is almost entirely restored. This is an example of Le Chatelier’s principle and the common ion effect. This contrasts with solutions of strong acids or strong bases, where any additional strong acid or base can greatly change the pH. This may be easier to see by comparing two graphs when an strong acid is titrated with a strong base the curve will have a large gradient throughout showing that a small addition of base/acid will have a large effect compared to a weak acid/strong base titration curve which will have a smaller gradient near the pKa. When writing about buffer systems they can be represented as salt of conjugate base/acid, or base/salt of conjugate acid. It should be noted that here buffer are presented in terms of the Brønsted-Lowry notion of acids and bases, as opposed to the Lewis acid-base theory (see acid-base reaction theories).
Preparation of buffer solutions
There are a couple of ways to prepare a buffer solution of a specific pH. In the first method, prepare a solution with an acid and its conjugate base by dissolving the acid form of the buffer in about 60% of the volume of water required to obtain the final solution volume. Then, measure the pH of the solution using a pH probe. The pH can be adjusted up to the desired value using a strong base like NaOH. If the buffer is made with a base and its conjugate acid, the pH can be adjusted using a strong acid like HCl. Once the pH is correct, dilute the solution to the final desired volume.
Alternatively, you can prepare solutions of both the acid form and base form of the solution. Both solutions must contain the same buffer concentration as the concentration of the buffer in the final solution. To get the final buffer, add one solution to the other while monitoring the pH.
In a third method, you can determine the exact amount of acid and conjugate base needed to make a buffer of a certain pH, using the Henderson-Hasselbach equation:
pH=pKa+log([A−][HA])
To prepare a buffer solution with a specific pH , we first choose a weak acid whose pka is close to the desired pH , then the pKa and the pH values are substituted to obtain the (conjugate base) /(acid ) ratio .these ratio can then be converted to molar quantities for the preparation of the buffer solution
- pH =pka +log (conjugate base)
- (acid)
- Log (conjugate base) =0
- (acid)
- pH =pKa
APPLICATION OF BUFFER SOLUTIONS Their resistance to changes in pH makes buffer solutions very useful for chemical manufacturing and essential for many biochemical processes. The ideal buffer for a particular pH has a pKa equal to the pH desired, since a solution of this buffer would contain equal amounts of acid and base and be in the middle of the range of buffering capacity.
Buffer are necessary to keep the correct pH for enzymes in many organisms to work. Many enzymes work only under very precise conditions; if the pH strays too far out of the margin, the enzymes slow or stop working and can denature, thus permanently disabling its catalytic activity.