Buffer and pH Adjustment

A buffer resists pH change because a weak acid and its conjugate base are both present in quantity, so added acid or alkali is consumed by shifting the ratio between them rather than by changing free hydrogen ion concentration. The Henderson-Hasselbalch relationship states it directly: pH equals the acid's pKa plus the logarithm of the base-to-acid ratio, which is why buffering capacity peaks within about one pH unit of the pKa and collapses outside it.

The usual parenteral buffers are acetate, citrate, phosphate, histidine and tris, and the choice is rarely arbitrary. Sodium phosphate is poor for frozen or lyophilised products because it undergoes a large pH shift on freezing as the dibasic salt crystallises out. Citrate is known for stinging subcutaneously, which drove citrate-free reformulations of several biologics. Insulin glargine shows pH as a design tool rather than housekeeping: formulated acidic where it is soluble, it precipitates as a microcrystalline depot on meeting neutral tissue.

pH selection is a compromise, not an optimum. Solubility is lowest near the isoelectric point, so formulations are pushed away from it, but degradation chemistry moves too. Asparagine deamidation and disulfide scrambling accelerate at neutral to alkaline pH, while peptide bond hydrolysis and aspartate isomerisation are favoured under acidic conditions, so the stability optimum and the solubility optimum are frequently different numbers.

The error worth naming is treating pH as an inert specification. Introducing a vehicle of different pH and no buffering capacity can move a formulation into a region where the peptide precipitates or a rapid degradation pathway opens. A related misreading is assuming a stated pH implies buffering: an unbuffered solution can sit at exactly the right pH with no capacity whatsoever to stay there.