Buffer Systems and Henderson-Hasselbalch

A buffer is a solution that resists changes in pH when acid or base is added by reversibly binding or releasing hydrogen ions. Body fluids contain several such systems, and the relationship between a buffer pair and the resulting pH is described by the Henderson-Hasselbalch equation, which underlies the standard description of acid-base status.

Definition

A physiological buffer is a conjugate weak-acid and weak-base pair that minimises change in hydrogen-ion concentration; the Henderson-Hasselbalch equation expresses the pH of such a system as a function of the ratio of the base to the acid form and the system's dissociation constant.

Scope

The topic covers the chemical principle of buffering, the major physiological buffer systems (bicarbonate, phosphate, proteins, and haemoglobin), the Henderson-Hasselbalch equation applied to the bicarbonate system, and why the open bicarbonate-carbon dioxide system is the body's most important buffer. It is presented as foundational physiology rather than clinical guidance.

Core questions

• What makes a solution act as a buffer, and what determines its capacity?
• Why is the bicarbonate-carbon dioxide pair the dominant extracellular buffer despite an unfavourable dissociation constant?
• How does the Henderson-Hasselbalch equation relate pH to bicarbonate and carbon dioxide tension?
• What roles do phosphate, protein, and haemoglobin buffers play?

Key concepts

• Conjugate acid-base pair
• Buffering capacity
• Dissociation constant (pKa)
• Open versus closed buffer systems
• Bicarbonate-carbon dioxide buffer
• Phosphate and protein buffers
• Isohydric principle

Key theories

Henderson-Hasselbalch relation

Expresses pH as the buffer's dissociation constant plus the logarithm of the ratio of conjugate base to acid; applied to the bicarbonate system it links arterial pH to the ratio of bicarbonate concentration to carbon dioxide tension.

Mechanisms

A buffer pair resists pH change because added hydrogen ions are taken up by the base form and removed by combination with the acid form, with the greatest capacity near the dissociation constant. Although the bicarbonate system's dissociation constant lies well below normal blood pH, it dominates physiologically because it is an open system: the lungs continuously remove the carbon dioxide produced and the kidneys regenerate bicarbonate, so the two components are independently regulated and the system is not exhausted. Phosphate buffering is more important inside cells and in tubular fluid, while plasma proteins and haemoglobin provide substantial intracellular and blood buffering; by the isohydric principle all buffers in a compartment share the same hydrogen-ion concentration, so measuring one pair reflects the whole.

Clinical relevance

The Henderson-Hasselbalch relationship is the basis for interpreting arterial blood-gas results, and understanding buffering explains why pH can be defended despite large acid or base loads. This entry describes the underlying chemistry and physiology and is not a basis for clinical decisions.

Evidence & guidelines

The buffering principles and the Henderson-Hasselbalch equation are standard, well-established physiology described consistently across reviews and texts (Hamm and colleagues, 2015; Berend and colleagues, 2014); the Stewart physicochemical framework offers an alternative, strong-ion-based account of what determines hydrogen-ion concentration (Story, 2016).

History

Lawrence Henderson described the carbonic-acid equilibrium of blood in the early twentieth century, and Karl Hasselbalch recast it in logarithmic (pH) form, giving the equation its joint name. The bicarbonate-centred buffer model became the standard account; Peter Stewart later argued that hydrogen-ion concentration is better understood as a dependent variable set by strong ions, weak acids, and carbon dioxide.

Debates

Is bicarbonate an independent determinant of pH or a dependent variable?

The traditional view treats bicarbonate as a regulated buffer that helps set pH, whereas the Stewart framework treats bicarbonate (like pH) as dependent on the strong-ion difference, total weak acid, and carbon dioxide tension; both describe the same measurements differently.

Key figures

• Lawrence J. Henderson
• Karl Albert Hasselbalch
• Peter A. Stewart

Seminal works

• hamm-2015
• berend-2014

Frequently asked questions

Why is the bicarbonate buffer so important if its pKa is far from blood pH?

Because it is an open system: the lungs continuously remove carbon dioxide and the kidneys regenerate bicarbonate, so its two components are independently controlled and it is far more effective in the body than its dissociation constant alone would suggest.

What does the Henderson-Hasselbalch equation tell us?

It relates pH to the ratio of the base form of a buffer to its acid form; for blood, pH depends on the ratio of bicarbonate concentration to the partial pressure of carbon dioxide.