A buffer solution contains relatively equal amounts of a weak acid and its conjugate base, HA(aq)/A^-(aq).

A buffer system maintains a nearly constant pH when small quantities of strong acid or strong base are added. This is because most of the added hydronium ions or hydroxide ions react with a component of the buffer system, as shown in the example below.

To better understand how a buffer functions, first consider unbuffered acetic acid solution. Acetic acid is a weak acid. In solution, the acetic acid molecules are in equilibrium with the acetate ions and hydrogen ions as shown below. However, the reactants are favoured highly because acetic acid is a weak acid.

According to Le Chatelier's Principle, adding hydronium ions to this system should shift the equilibrium to the left. However, this shift to the left is limited by the lack of acetate ions. As soon as there are no more acetate ions to react with added hydronium ions, the pH of the solution falls dramatically as H₃O⁺ ions build up in the solution. In other words, acetate ions are the limiting reagent.

What if you could increase the concentration of acetate ions in the original solution? Would not that greatly increase the ability of the system to "soak up" added hydronium ions? The answer is yes.

\( \mathrm { CH_3COOH(aq) + H_2O(l) \leftrightharpoons CH_3COO^- (aq) + H_3O^+(aq) } \)

If a small amount of strong acid is added to the acetic acid- acetate ion system, the base part of the buffer (CH₃COO^-) will react with these additional hydronium ions to form water and CH₃COOH molecules. In other words, the system will alleviate the added stress by favouring the reverse reaction. Because the equilibrium shift to the left causes hydronium ions to be consumed, the pH of the solution does not change appreciably. In fact, the buffer system will continue to resist any significant pH change until all the acetate ions have been used up.

Conversely, if a small amount of a strong base is added to this buffer solution, the additional hydroxide ions will react with any available hydronium ions to produce water. Because this reaction removes hydronium ions from the system, the system will alleviate the added stress by favouring the forward reaction. Because the equilibrium shift to the right causes hydronium ions to be replaced, the pH of the solution does not change appreciably.

Designing a Buffer As you noticed, there are many weak acids on the Table of Relative Strengths of Acids and Bases. How do we know which weak acid/conjugate base pair will create the best buffer system at a specific pH? Step One - Calculate the concentration of hydronium ions in the solution that requires buffering. You can determine this concentration from the pH of the solution. Use the formula: \( \mathrm { [H_3O^+(aq)] = 10^{-pH} } \) Remember, for logarithmic values, such as pH, any digit to the left of the decimal point is not significant. For example, a pH of 4.3 has one significant digit; therefore, the hydronium ion concentration is 5 x 10-5 mol/L. Significant digits sometimes will be a factor when you are choosing the best weak acid for your buffer. Step Two - Using the Table of Relative Strengths of Acids and Bases, choose the weak acid that has a Ka closest to the hydronium ion concentration of the solution that requires buffering. Step Three - Mix equal quantities of the weak acid and its conjugate base (salt). Buffers work best if the ratio of weak acid to conjugate base is 1:1. An ideal buffer contains these two entities in equal concentration.

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