Buffering agent
A buffering agent is a weak acid or base used to maintain the acidity (pH) of a solution near a chosen value after the addition of another acid or base. That is, the function of a buffering agent is to prevent a rapid change in pH when acids or bases are added to the solution. Buffering agents have variable properties—some are more soluble than others; some are acidic while others are basic. As pH managers, they are important in many chemical applications, including agriculture, food processing, biochemistry, medicine and photography.
A buffering agent can be either a weak acid or weak base. Buffering agents are usually added to water to form a buffer solution, which only slightly changes its pH in response to other acids and bases being combined with it, particularly a strong acid or a strong base. Another example is buffered aspirin which has a buffering agent, such as MgO or CaCO3, that will help maintain the balance between the acid H-A (protonated) and the salt A− (deprotonated) forms of aspirin as it passes through the acidic stomach of the patient. The H-A form of aspirin is a covalent molecule and is more absorbed by the stomach lining, leading to irritation. The antacid properties of these buffering agents help to maintain the balance toward the salt form by reducing the amount of stomach acid (HCl) which protonates the salt form, A−.
The way buffering agents work can be seen by calculating how little the pH of buffer solutions will change after addition of a strong acid or a strong base, whereas the same addition would significantly change the pH of a non-buffered solution. Using the Henderson-Hasselbalch equation we get an equilibrium expression between the acid and conjugate base in terms of the log of the ratio of the acid to conjugate base (the salt of the acid). The concentrations of the weak acid and its salt can change significantly, but the log of their ratio will not. The resulting pH of this combination can be found by using Le Chatelier's principle. For a simple numerical example, take the case where the concentrations of the weak acid and its salt are equal. If an added strong base halves the [HA] ([HA] decreases to [HA]-0.5[HA]), then [A−] will increase by [A−]+0.5[A−] since every molecule of HA that dissociates forms one molecule of [A−]. Thus the pH will be raised by a factor of log3 or 0.5 pH units (when the original [HA] and [A−] are equal):
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