A buffer is a solution that can resist pH change upon the addition of an acidic or basic components. It is able to neutralize small amounts of added acid or base, thus maintaining the pH of the solution relatively stable.
Acidic buffer : Acidic buffers are solutions that have a pH below 7 and contain a weak acid and one of its salts. For example, a mixture of acetic acid and sodium acetate acts as a buffer . Alkaline buffers, on the other hand, have a pH above 7 and contain a weak base and one of its salts. For example, a mixture of ammonium chloride and ammonium hydroxide acts as a buffer solution with a pH of about 9.25
How does buffer work ? buffer is able to resist pH change because the two components (conjugate acid and conjugate base) are both present in appreciable amounts at equilibrium and are able to neutralize small amounts of other acids and bases (in the form of H3O+ and OH–) when the are added to the solution.
Example :
Hydrofluoric Acid (HF) and Sodium Fluoride (NaF) buffer. Hydrofluoric acid is a weak acid due to the strong attraction between the relatively small F– ion and solvated protons (H3O+), which does not allow it to dissociate completely in water. Therefore, if we obtain HF in an aqueous solution, we establish the following equilibrium with only slight dissociation (Ka(HF) = 6.6×10-4, strongly favors reactants):
HF(aq)+H2O(l)⇌F−(aq)+H3O+(aq)(1)(1)HF(aq)+H2O(l)⇌F(aq)−+H3O(aq)+
We can then add and dissolve sodium fluoride into the solution and mix the two until we reach the desired volume and pH at which we want to buffer. When Sodium Fluoride dissolves in water, the reaction goes to completion, thus we obtain:
NaF(aq)+H2O(l)→Na+(aq)+F−(aq)(2)(2)NaF(aq)+H2O(l)→Na(aq)++F(aq)−
Since Na+ is the conjugate of a strong base, it will have no effect on the pH or reactivity of the buffer. The addition of NaFNaF to the solution will, however, increase the concentration of F– in the buffer solution and, consequently, by Le Chatelier’s Principle, lead to slightly less dissociation of the HF in the previous equilibrium, as well.
If acid were added:
F−(aq)+H3O+(aq)⇌HF(aq)+H2O(l)(3)
In this reaction, the conjugate base, F-, will neutralize the added acid, H3O+, and this reaction goes to completion, because the reaction of F- with H3O+ has an equilibrium constant much greater than one. (In fact, the equilibrium constant the reaction as written is just the inverse of the Ka for HF: 1/Ka(HF) = 1/(6.6×10-4) = 1.5×10+3.) So long as there is more F- than H3O+, almost all of the H3O+ will be consumed and the equilibrium will shift to the right, slightly increasing the concentration of HF and slightly decreasing the concentration of F-, but resulting in hardly any change in the amount of H3O+ present once equilibrium is re-established.
If base were added:
HF(aq)+OH−(aq)⇌F−(aq)+H2O(l)(4)
In this reaction, the conjugate acid, HF, will neutralize added amounts of base, OH-, and the equilibrium will again shift to the right, slightly increasing the concentration of F- in the solution and decreasing the amount of HF slightly. Again, since most of the OH- is neutralized, little pH change will occur.
These two reactions can continue to alternate back and forth with little pH change.
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