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.
Titration of a weak acid with a strong base the flat region at pKa is the buffering region.
Titration of a weak acid with a strong base the flat region at pKa is the buffering region.
Titration of a strong acid with a strong base. Note the sharp rise in pH: this solution can not buffer.
Titration of a strong acid with a strong base. Note the sharp rise in pH: this solution can not buffer.
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 solutions 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). Omitted here are buffer solutions prepared with solvents other than water.
Applications
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 solutions 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. A buffer of carbonic acid (H2CO3) and bicarbonate (HCO3¥) is present in blood plasma, to maintain a pH between
7.35 and 7.45.
Industrially, buffer solutions 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
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