In this article, the buffer capacity concept is revisited, particularly concerning its behavior in hydroorganic mobile phases. The buffer capacity of a polyprotic acid, or a mixture of monoprotic acids, depends upon the concentration of each weak acid–conjugate base pair, and the pH of its maximum value mainly fits to the acid–base pK
, but it is shifted to a certain degree according to the ionic strength of the buffered solution. Consequently, when an organic solvent is added to an aqueous buffer to prepare a particular mobile phase, the buffer capacity of the hydroorganic mixture is reduced due to the dilution effect, and the maximum buffer capacity is shifted to lower or higher pH values according to the nature of the buffering acid–base pair.
In reversed-phase high performance liquid chromatography (HPLC), the pH of the mobile phase is a fundamental parameter that significantly affects the retention of ionizable analytes. The buffer capacity is a very important quality of a buffer solution, because it gives information about the resistance of a solution to pH change. Between December 2002 and February 2003, a series of interesting and informative articles from Tindall and Dolan dealing with mobile-phase buffers was published in LCGC
Europe (1–3). These articles were about pH interpretation in partially aqueous mobile phases and buffer selection and preparation. They also included a brief introduction about buffer capacity in aqueous solutions. The objective of this article is to describe buffer capacity more fully, particularly in hydroorganic mobile phases.
Some Essential Foundations for the Concept of Buffer Capacity
In 1922, Van Slyke (4) proposed the following differential ratio as a quantitative measure of the buffer action in aqueous solution:
expressing the relationship between the increment of the concentration of a strong base B, C
B (or strong acid A, C
A) added to a buffer solution and the resultant increment in pH. Buffer capacity (β) is always a positive numerical value: if base is added to a buffered solution the pH is increased, so both dC
b and dpH are positive and the ratio is also positive; if acid is added, dC
A is positive but dpH is negative, thus the minus sign before the dC
A/dpH turns the negative ratio into a positive function. β is an additive quantity that depends upon the buffer capacities of the acid–base pair that make up the particular buffer system, and the unavoidable acid–base pairs of water (H3O+ /H2O and H2O/OH- ).
In the case of a polyprotic acid buffer (for example, phosphoric) or a mixture of monoprotic acids (for example, ammonium carbonate or ammonium formate), the total buffer capacity of the system can be expressed as follows (5):
2. . . refers to the total concentration of each weak acid-conjugate base pair, K'ai is the concentration acidity constant at the working ionic strength, and K
is the autoprotolysis constant of water. This equation is correctly applicable to acids presenting a ratio of successive ionization constants (K
a1) lower than 0.05, a condition fulfilled by most of the common polyprotic acids. When this condition is not satisfied (for example, with tartaric and adipic acids), more complex equations are required. For example, the exact expression for buffer capacity of a diprotic acid is as follows (6):
Buffer capacity is proportional to the concentration of the buffering species and depends upon its pK'a. It must be noticed that activity correction has not been carried out in equations 3 and 4. For an acid–base equilibrium like HAz € Az-1 + H+ the concentration acidity constant (K'ai) is defined as follows: