16.5

Buffer Effectiveness

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The effectiveness of any buffer to resist pH change depends upon the concentration ratio of the weak acid and its conjugate base, or weak base and its conjugate acid, as well as their absolute concentrations.

The buffer range is the pH range that inhibits a significant increase or decrease in the pH upon the addition of an acid or base. The range is one unit higher or lower than the pK_a.

Therefore, to be an effective buffer, the ratio of weak acid to base or weak base to acid should be between 10 to 1 and 1 to 10.

The Henderson-Hasselbalch equation can be solved to support these values. If the acid’s concentration is ten times that of the base, the pH will be 1 unit less than the pK_a. In contrast, if the base’s concentration is ten times that of the acid, the pH will be 1 unit more than the pK_a.

A buffer is most effective in the middle of its buffer range when the concentration of the weak acid and conjugate base are equal, and the pH equals the pK_a. As the difference between the amounts of the weak acid and base increases, the buffer becomes less effective.

Therefore, buffer A, containing 1 molar each of acetic acid and acetate, will be more effective than buffer B, containing 0.1 molar acetic acid and 1 molar acetate.

The absolute concentration of a weak acid and the base also determines the effectiveness of the buffer. The greater the concentration of the weak acid and base, the more strong acid or base it can neutralize.

Therefore, a buffer with 1 molar each of formic acid and formate is more effective than a buffer with 0.1 molar each.

Buffer capacity is the amount of a strong acid or base a buffer can neutralize before a significant change in its pH. Therefore, the buffer capacity increases both with higher concentrations of a weak acid and its conjugate base and when the ratio of a weak acid to the base approaches one.

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Buffer solutions do not have an unlimited capacity to keep the pH relatively constant . Instead, the ability of a buffer solution to resist changes in pH relies on the presence of appreciable amounts of its conjugate weak acid-base pair. When enough strong acid or base is added to substantially lower the concentration of either member of the buffer pair, the buffering action within the solution is compromised.

The buffer capacity is the amount of acid or base that can be added to a given volume of a buffer solution before the pH changes significantly, usually by one unit. Buffer capacity depends on the amounts of the weak acid and its conjugate base that are in a buffer mixture. For example, 1 L of a solution that is 1.0 M in acetic acid and 1.0 M in sodium acetate has a greater buffer capacity than 1 L of a solution that is 0.10 M in acetic acid and 0.10 M in sodium acetate even though both solutions have the same pH. The first solution has more buffer capacity because it contains more acetic acid and acetate ion.

Selection of Suitable Buffer Mixtures

There are two useful rules of thumb for selecting buffer mixtures:

A good buffer mixture should have about equal concentrations of both of its components. A buffer solution has generally lost its usefulness when one component of the buffer pair is less than about 10% of the other.
Weak acids and their salts are better as buffers for pHs less than 7; weak bases and their salts are better as buffers for pHs greater than 7.

Blood is an important example of a buffered solution, with the principal acid and ion responsible for the buffering action being carbonic acid, H₂CO₃, and the bicarbonate ion, HCO₃⁻. When a hydronium ion is introduced to the bloodstream, it is removed primarily by the reaction:

Lewis structure diagram of PF5 molecule illustrating electron arrangement and chemical bonds.

An added hydroxide ion is removed by the reaction:

Gas chromatography diagram showing equipment setup for compound separation and analysis.

The added strong acid or base is thus effectively converted to the much weaker acid or base of the buffer pair (H₃O⁺ is converted to H₂CO₃ and OH⁻ is converted to HCO₃⁻). The pH of human blood thus remains very near the value determined by the buffer pairs pK_a, in this case, 7.35. Normal variations in blood pH are usually less than 0.1, and pH changes of 0.4 or greater are likely to be fatal.

This text is adapted from Openstax, Chemistry 2e, Section 14.6: Buffers.

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Buffer Capacity

Henderson Hasselbalch Equation

Weak Acid Base Pair

Conjugate Acid Base

Acid Base Ratio

Blood Buffer System

Carbonic Acid Bicarbonate

Acid Base Neutralization