Buffer solution and buffering principle

Buffer Solution and Buffer Principle (1) Buffering Effect and Buffer Solution Pure water has a pH of 7.0 at 25°C. However, when it is exposed to air for some time, the pH decreases to around 5.5 due to the absorption of carbon dioxide, which forms carbonic acid in solution. Adding just one drop of concentrated hydrochloric acid (approximately 12.4 mol/L) to 1 liter of pure water can increase the [H⁺] concentration by about 5,000 times, rising from 1.0×10⁻⁷ to 5×10⁻⁴ mol/L. Similarly, adding one drop of a 12.4 mol/L sodium hydroxide solution to 1 liter of pure water causes a pH change of about 3 units. This shows that even a small amount of strong acid or base can significantly alter the pH of pure water. In contrast, if one drop of concentrated hydrochloric acid is added to 1 liter of an HOAc–NaOAc buffer solution or a NaH₂PO₄–Na₂HPO₄ buffer solution, the pH change is minimal. This is because buffer solutions resist changes in pH when small amounts of acid or base are added. The buffering capacity comes from the presence of both a weak acid and its conjugate base (or a weak base and its conjugate acid), which can neutralize added H⁺ or OH⁻ ions. This ability to maintain a stable pH makes buffer solutions essential in many biological and chemical processes, such as maintaining the pH of blood, regulating enzyme activity, and ensuring accurate results in laboratory experiments. Understanding the buffering principle helps in designing and using these solutions effectively in various applications.

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