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.