# Henderson–Hasselbalch equation

In chemistry and biochemistry, the Henderson–Hasselbalch equation can be used to estimate the pH of a buffer solution containing given concentrations of an acid and its conjugate base (or a base and its conjugate acid). The numerical value of the acid dissociation constant of the acid must also be known.

## History

Lawrence Joseph Henderson derived an equation with which the pH of a buffer solution may be calculated.[1] Later, Karl Albert Hasselbalch re-expressed that formula in logarithmic terms,[2] resulting in the Henderson–Hasselbalch equation.

## Theory

A simple buffer solution consists of a solution of an acid and a salt of the conjugate base of the acid. For example, the acid may be acetic acid and the salt may be sodium acetate. The Henderson–Hasselbalch equation relates the pH of a solution containing a mixture of the two components to the acid dissociation constant, Ka, and the concentrations of the species in solution.[3] To derive the equation a number of simplifying assumptions have to be made.

Assumption 1: The acid is monobasic and dissociates according to the equation

${\displaystyle {\ce {HA}}\leftrightharpoons {\ce {H^+}}+{\ce {A^-}}}$

It is understood that the symbol H+ stands for the hydrated hydronium ion. The Henderson–Hasselbalch equation can be applied to a polybasic acid only if its consecutive pK values differ by at least 3. Phosphoric acid is such an acid.

Assumption 2. The self-ionization of water can be ignored.

This assumption is not valid with pH values more than about 10. For such instances the mass-balance equation for hydrogen must be extended to take account of the self-ionization of water.

CH = [H+] + Ka[H+][A]- Kw[H+]−1
CA = [A] + Ka[H+][A]

and the pH will have to be found by solving the two mass-balance equations simultaneously for the two unknowns, [H+] and [A].

Assumption 3: The quotient of activity coefficients, ${\displaystyle \Gamma }$, is a constant under the experimental conditions covered by the calculations.

The thermodynamic equilibrium constant, ${\displaystyle K^{*}}$,

${\displaystyle K^{*}={\frac {[{\ce {H+}}][{\ce {A^-}}]}{[{\ce {HA}}]}}\times {\frac {\gamma _{{\ce {H+}}}\gamma _{{\ce {A^-}}}}{\gamma _{HA}}}}$

is a product of a quotient of concentrations ${\displaystyle {\frac {[{\ce {H+}}][{\ce {A^-}}]}{[{\ce {HA}}]}}}$ and a quotient, ${\displaystyle \Gamma }$, of activity coefficients ${\displaystyle {\frac {\gamma _{{\ce {H+}}}\gamma _{{\ce {A^-}}}}{\gamma _{HA}}}}$. In these expressions, the quantities in square brackets signify the concentration of the undissociated acid, HA, of the hydrogen ion H+, and of the anion A; the quantities ${\displaystyle \gamma }$ are the corresponding activity coefficients. If the quotient of activity coefficients can be assumed to be a constant which is independent of concentrations and pH , the dissociation constant, Ka can be expressed as a quotient of concentrations.

${\displaystyle K_{a}=K^{*}/\Gamma ={\frac {[{\ce {H+}}][{\ce {A^-}}]}{[{\ce {HA}}]}}}$

## Application

Rearrangement of the above expression for Ka and taking logarithms provides the Henderson-Hasselbalch equation

${\displaystyle {\ce {pH}}={\ce {p}}K_{{\ce {a}}}+\log _{10}\left({\frac {[{\ce {A^-}}]}{[{\ce {HA}}]}}\right)}$

This equation can be used to calculate the pH of a solution containing the acid and one of its salts, that is, of a buffer solution. With bases, if the value of an equilibrium constant is known in the form of a base association constant, Kb the dissociation constant of the conjugate acid may be calculated from

pKa + pKb = pKw

where Kw is the self-dissociation constant of water. pKw has a value of approximately 14 at 25C.

If the "free acid" concentration, [HA], can be taken to be equal to the analytical concentration of the acid, TAH (sometimes denoted as CAH) an approximation is possible, which is widely used in biochemistry; it is valid for very dilute solutions.

${\displaystyle {\ce {pH}}={\ce {p}}K_{{\ce {a}}}+\log _{10}\left({\frac {[{\ce {A^-}}]}{T_{AH}}}\right)}$

The effect of this approximation is to introduce an error in the calculated pH, which becomes significant at low pH and high acid concentration. With bases the error becomes significant at high pH and high base concentration.[4] (pdf)

## References

1. ^ Lawrence J. Henderson (1908). "Concerning the relationship between the strength of acids and their capacity to preserve neutrality". Am. J. Physiol. 21: 173–179.
2. ^ Hasselbalch, K. A. (1917). "Die Berechnung der Wasserstoffzahl des Blutes aus der freien und gebundenen Kohlensäure desselben, und die Sauerstoffbindung des Blutes als Funktion der Wasserstoffzahl". Biochemische Zeitschrift. 78: 112–144.
3. ^ For details and worked examples see, for instance,Skoog, Douglas A.; West, Donald M.; Holler, F. James; Crouch, Stanley R. (2004). Fundamentals of Analytical Chemistry (8th ed.). Belmont, Ca (USA): Brooks/Cole. pp. 251–263. ISBN 0-03035523-0.
4. ^ Po, Henry N.; Senozan, N. M. (2001). "Henderson–Hasselbalch Equation: Its History and Limitations". J. Chem. Educ. 78 (11): 1499–1503. Bibcode:2001JChEd..78.1499P. doi:10.1021/ed078p1499.