Difference between revisions of "Buffer solution"
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| − | A '''buffer solution''' (more precisely, '''pH buffer''' or '''hydrogen ion buffer''') is an aqueous solution consisting of a mixture of a weak acid and its conjugate base, or vice versa. Its pH changes very little when a small amount of strong acid or base is added to it. Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications. | + | A '''buffer solution''' (more precisely, '''pH buffer''' or '''hydrogen ion buffer''') is an aqueous solution consisting of a mixture of a weak acid and its conjugate base, or vice versa. Its [[pH]] changes very little when a small amount of strong acid or base is added to it. Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications. |
==Principles of buffering== | ==Principles of buffering== | ||
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:HA ⇌ H<sup>+</sup> + A<sup>−</sup> | :HA ⇌ H<sup>+</sup> + A<sup>−</sup> | ||
| − | When some | + | When some strong acid is added to an equilibrium mixture of the weak acid and its conjugate base, the equilibrium is shifted to the left, in accordance with [[Le Chatelier's principle]]. Because of this, the hydrogen ion concentration increases by less than the amount expected for the quantity of strong acid added. |
Similarly, if strong alkali is added to the mixture the hydrogen ion concentration decreases by less than the amount expected for the quantity of alkali added. The effect is illustrated by the simulated titration of a weak acid with p''K''<sub>a</sub> = 4.7. The relative concentration of undissociated acid is shown in blue and of its conjugate base in red. The pH changes relatively slowly in the buffer region, pH = p''K''<sub>a</sub> ± 1, centered at pH = 4.7 where [HA] = [A<sup>−</sup>]. The hydrogen ion concentration decreases by less than the amount expected because most of the added hydroxide ion is consumed in the reaction: | Similarly, if strong alkali is added to the mixture the hydrogen ion concentration decreases by less than the amount expected for the quantity of alkali added. The effect is illustrated by the simulated titration of a weak acid with p''K''<sub>a</sub> = 4.7. The relative concentration of undissociated acid is shown in blue and of its conjugate base in red. The pH changes relatively slowly in the buffer region, pH = p''K''<sub>a</sub> ± 1, centered at pH = 4.7 where [HA] = [A<sup>−</sup>]. The hydrogen ion concentration decreases by less than the amount expected because most of the added hydroxide ion is consumed in the reaction: | ||
Revision as of 17:31, 30 November 2022
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A buffer solution (more precisely, pH buffer or hydrogen ion buffer) is an aqueous solution consisting of a mixture of a weak acid and its conjugate base, or vice versa. Its pH changes very little when a small amount of strong acid or base is added to it. Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications.
Principles of buffering
Buffer solutions achieve their resistance to pH change because of the presence of an equilibrium between the acid HA and its conjugate base A−.
- HA ⇌ H+ + A−
When some strong acid is added to an equilibrium mixture of the weak acid and its conjugate base, the equilibrium is shifted to the left, in accordance with Le Chatelier's principle. Because of this, the hydrogen ion concentration increases by less than the amount expected for the quantity of strong acid added.
Similarly, if strong alkali is added to the mixture the hydrogen ion concentration decreases by less than the amount expected for the quantity of alkali added. The effect is illustrated by the simulated titration of a weak acid with pKa = 4.7. The relative concentration of undissociated acid is shown in blue and of its conjugate base in red. The pH changes relatively slowly in the buffer region, pH = pKa ± 1, centered at pH = 4.7 where [HA] = [A−]. The hydrogen ion concentration decreases by less than the amount expected because most of the added hydroxide ion is consumed in the reaction:
- OH− + HA → H2O + A−
and only a little is consumed in the neutralization reaction which results in an increase in pH.
- OH− + H+ → H2O
