Buffers are solutions that resist change to pH upon addition of small amounts of either an acid or a base.
Many systems need to be at a constant pH, especially biological systems. Normally, this is to control reaction rates - hydronium ion and/or hydroxide ion can act as catalysts. By controlling the pH, the amount of catalyst is kept constant and the reaction rate is correct for the particular application.
In order for a solution to behave as a buffer, there must be both an acid and a base present.
The acid can react with added bases while the base can react with added acids. Further, weak acids and bases are required because these are only slightly ionized, thereby contributing little to the hydronium ion or hydroxide ion concentration.
The problem is that acids react with bases, so preparing a buffer requires careful consideration of the components. The best way to prepare a buffer is to use a conjugate acid/base pair and rely on the common ion effect to maintain the pH.
How does this work?
Consider a solution composed of a weak acid, HA, and the sodium salt of the conjugate base, NaA
The salt completely ionizes in water to give ions:
NaA(aq) → Na+(aq) + A–(aq)
The weak acid is in equilibrium with hydronium ion and the conjugate base:
HA(aq) + H2O(l) → ← H3O+(aq) + A–(aq)
The solution is composed mainly of HA (a weak acid), A– (a weak base), and Na+ (a neutral cation). There is also a little hydronium ion and even less hydroxide ion.
What happens when a small amount of a strong acid, HX, is added the solution?
The acid will react with the base in the solution:
HX(aq) + A–(aq) → HA(aq) + X–(aq)
There is no explicit change in the hydronium on concentration because H3O+ is neither a reactant nor a product in the acid/base reaction!
The change in concentration of HA and A– will cause a small shift of the weak acid equilibrium, so the pH will drop a little.
Now suppose a strong base, MOH, is added to the buffer solution. The base will react with the acid in the solution, HA:
MOH(aq) + HA(aq) → M+(aq) + A–(aq) + H2O(l)
Again, there is no explicit change in the hydronium ion concentration so the pH does not change drastically. As above there is a small change in the weak acid equilibrium that drives the pH up a little.
A 1.00 L solution is prepared that is 0.150 M nitrous acid and 0.200 M sodium nitrite. What is the pH of the solution at 25 °C? What is the pH after adding 1.0 g of hydrogen bromide (at 25 °C)? Instead of hydrogen bromide, suppose 1.0 g of potassium hydroxide was added to the solution; what would the pH be in this case (at 25 °C)?
This is a 3-part problem:
1. The initial buffer solution.
2. The buffer plus some strong base.
3. The buffer plus some strong acid.
Notice that the buffer was effective because we only added small amounts of the strong base or strong acid. If the added acid or base had exceeded the concentration of the buffer components, then the weak acid equilibrium would no longer be valid and the strong acid or weak acid chemistry would dominate. This is known as the buffer capacity.