D. D. Perrin

Metal-Ion Buffers

Chemical and biological reactions often depend critically on the presence of low concentrations of certain free metal ions. Concentrations less than 10−5M are difficult to maintain in the presence of adventitious complexing agents, hydrolytic equilibria, adsorption and, possibly, contamination. Many problems can be overcome by the use of metal-ion buffers which provide a controlled source of free metal ions in a manner similar to the regulation of hydrogen ion concentration by pH buffers.

The Theory of Buffer Action

An acid is a species such as the ammonium ion or the acetic acid molecule which has a tendency to lose a proton. A base (such as the ammonia molecule or the acetate anion) is able to accept a proton. Hence for every acid, HA, there is a conjugate base, A-, and for every base, B, there is a conjugate acid, BH+.

Practical Limitations in the use of Buffers

The suitability of a buffer system for any particular application depends on many factors, the first of which is the pKa value of the buffer acid or base. For a buffer to be effective, its pH must be within the range pKa ± 1, or preferably, within pKa ± 0.5. The former scarcity of buffer substances with a pH range 6 to 8 frequently led to the use of buffers such as phosphate or Tris in pH regions where they had little buffer capacity.

Applications of pH Buffers

The buffering ability of a weakly acidic or basic group is limited approximately to the range, pH = pKa ± 1, the greatest effect being at pH = pKa. This is clearly the most important single factor in choosing a buffer for any particular application and reference to lists of pKa values such as those given in Appendix III provides a rapid indication of possible buffer substances.

Buffers for use in Partially Aqueous and Non-Aqueous Solvents and Heavy Water

Solvent molecules are involved in acid-base equilibria as acceptors or donors of protons, so that the acidic or basic strength of a substance varies with the nature of the solvent. The lower alcohols resemble water, in that they can form the ions ROH2+ and RO- , but their dissociation is less than for water ((pKCH3OH=16.7, PKC2H5OH=19.1, cf. pKw= 14.0). Consequently, substances dissolved in alcohols are weaker acids and bases than in water. Other factors influencing acidic and basic strengths in solutions include the dielectric constant and solute-solvent interactions which, in mixed solvents, can lead to the further complication of selective ordering of solvent molecules around ionic species.

New pH-Buffer Tables and Systems

Although available buffer tables are extensive, they may fail to meet particular requirements, such as a specified ionic strength or a nominated buffer species. Where pKa values of the buffer substances are known the compositions of the solutions can be calculated as indicated below. Considerations governing the design of new types of buffer systems are also discussed.

Prediction of pK a Values of Substituted Aliphatic Acids and Bases

Free energy changes produced by inserting substituents into organic molecules are approximately additive. Hence, the simplest way to use a linear free energy relationship for predicting the pK a of an aliphatic acid or amine where the pK a of the parent compound is known, is to add to it increments of pK a (∆pK a values) corresponding to the free energy changes produced by inserting the individual substituents. Finally, the total is adjusted by applying statistical factors or other corrections such as a decrease of 0.2 in the pK a value of an amine for every methyl group bound to the nitrogen, or an increase of 0.2 if the nitrogen atom forms part of a ring.

Prediction of p K a Values for Phenols, Aromatic Carboxylic Acids and Aromatic Amines

Within any class of substituted aromatic acid or base it is usually found that the acidic dissociation constant will vary in much the same way as that of the corresponding derivative in any other class. In most cases, these changes in acid strength, ∆pK a, produced by a substituent in a meta- or para-position can be predicted by the Hammett equation.

Preparation of Buffer Solutions

Materials used in the preparation of buffer solutions should be good quality laboratory chemicals, purified if necessary as described in Chapter 8 and dried to constant composition. The distilled water used as solvent should have been recently boiled to remove dissolved carbon dioxide and have been protected from contamination by atmospheric carbon dioxide while cooling. (This precaution is unnecessary in preparing buffer solutions having pH values less than 5.) The water should have a specific conductivity of less than 2 × 10−6 ohm−1 cm−1 at 25° C. Solutions should be stored in stoppered Pyrex (or similar borosilicate glass) or pure polyethylene bottles.

Recapitulation of the Main pK a Prediction Methods

When a pK a value is required, it may be possible to find the value at the required temperature and ionic strength in one of the compilations listed in Chapter 1. If the required value is not available, the ideal procedure would be to determine the value under the required conditions using procedures described by Albert and Serjeant (1971). If time does not permit or when a less accurately known value will suffice, a prediction based on methods described in this work can be made. It should be noted that the best predictions are those based on linear free energy relationships. In general, predictions are best which use the least number of substituent sigma constants and the least number of approximations. Ideally, pK a predictions should be based on a model compound as similar as possible in structure to the required compound.