Henderson–Hasselbalch equation

In chemistry, the Henderson–Hasselbalch equation describes the derivation of pH as a measure of acidity (using $p K a$ , the negative log of the acid dissociation constant) in biological and chemical systems. The equation is also useful for estimating the pH of a buffer solution and finding the equilibrium pH in acid-base reactions (it is widely used to calculate the isoelectric point of proteins).

The equation is given by:

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

Here, $[HA]$ is the molar concentration of the undissociated weak acid, $[A⁻]$ is the molar concentration (molarity, M) of this acid's conjugate base and $p K a$ is $-log 10 K a$ where $K a$ is the acid dissociation constant, that is:

{\mathrm {p}}K_{{\mathrm {a}}}=-\log _{{10}}(K_{{\mathrm {a}}})=-\log _{{10}}\left({\frac {[{\mathrm {H}}_{{3}}{\mathrm {O}}^{+}][{\mathrm {A}}^{-}]}{[{\mathrm {HA}}]}}\right)

for the non-specific Brønsted acid-base reaction:

{\mathrm {HA}}+{\mathrm {H}}_{{2}}{\mathrm {O}}\rightleftharpoons {\mathrm {A}}^{-}+{\mathrm {H}}_{{3}}{\mathrm {O}}^{+}

In these equations, $A⁻$ denotes the ionic form of the relevant acid. Bracketed quantities such as [base] and [acid] denote the molar concentration of the quantity enclosed.

For bases

For the standard base equation:^[1]

{\mathrm {B}}+{\mathrm {H}}^{{+}}\rightleftharpoons {\mathrm {BH}}^{{+}}

A second form of the equation expressed in terms of $K_{{\mathrm {b}}}$ where $K_{{\mathrm {b}}}$ is the base dissociation constant:

{\mathrm {p}}K_{{\mathrm {b}}}=-\log _{{10}}(K_{{\mathrm {b}}})=-\log _{{10}}\left({\frac {[{\mathrm {O}}{\mathrm {H}}^{-}][{\mathrm {HA}}]}{[{\mathrm {A}}^{-}]}}\right)

In analogy to the above equations, the following equation is valid:

{\mathrm {pOH}}={\mathrm {p}}K_{{\mathrm {b}}}+\log _{{10}}\left({\frac {[{\mathrm {BH}}^{+}]}{[{\mathrm {B}}]}}\right)

where BH⁺ denotes the conjugate acid of the corresponding base B. Using the properties of these terms at 25 degrees Celsius one can synthesise an equation for pH of basic solutions in terms of pK_a and pH:

{\mathrm {pH}}={\mathrm {p}}K_{{\mathrm {a}}}+\log _{{10}}\left({\frac {[{\mathrm {B}}]}{[{\mathrm {BH}}^{{+}}]}}\right)

Derivation

The Henderson–Hasselbalch equation is derived from the acid dissociation constant equation by the following steps:^[2]

K_{{\mathrm {a}}}={\frac {[{\mathrm {H}}^{+}][{\mathrm {A}}^{-}]}{[{\mathrm {HA}}]}}

Taking the log, to base ten, of both sides gives:

\log _{{10}}K_{{\mathrm {a}}}=\log _{{10}}\left({\frac {[{\mathrm {H}}^{+}][{\mathrm {A}}^{-}]}{[{\mathrm {HA}}]}}\right)

Then, using the properties of logarithms:

\log _{{10}}K_{{\mathrm {a}}}=\log _{{10}}[{\mathrm {H}}^{+}]+\log _{{10}}\left({\frac {[{\mathrm {A}}^{-}]}{[{\mathrm {HA}}]}}\right)

Identifying the left-hand side of this equation as -pK_a and the $\log _{{10}}[{\mathrm {H}}^{{+}}]$ as -pH:

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

Adding pH and pK_a to both sides:

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

The ratio $[{\mathrm {A}}^{-}]/[{\mathrm {HA}}]$ is unitless, and as such, other ratios with other units may be used. For example, the mole ratio of the components, $n_{A^-}/n_{HA}$ or the fractional concentrations $\alpha_{A^-}/\alpha_{HA}$ where $\alpha_{A^-}+\alpha_{HA}=1$ will yield the same answer. Sometimes these other units are more convenient to use.

History

Lawrence Joseph Henderson wrote an equation, in 1908, describing the use of carbonic acid as a buffer solution. Karl Albert Hasselbalch later re-expressed that formula in logarithmic terms, resulting in the Henderson–Hasselbalch equation.^[3] Hasselbalch was using the formula to study metabolic acidosis.

Limitations

There are some significant approximations implicit in the Henderson–Hasselbalch equation. The most significant is the assumption that the concentration of the acid and its conjugate base at equilibrium will remain the same as the formal concentration. This neglects the dissociation of the acid and the binding of H+ to the base. The dissociation of water and relative water concentration itself is neglected as well. These approximations will fail when dealing with relatively strong acids or bases (pKa more than a couple units away from 7), dilute or very concentrated solutions (less than 1 mM or greater than 1M), or heavily skewed acid/base ratios (more than 100 to 1). In high buffer dilutions, where the concentration of protons arising from water become equally or more prevalent than the buffer species themselves (at pH 7, this means buffer component concentrations of <10⁻⁵ M formally, but practically much higher), the pKa of the 'buffer' system will tend towards neutrality.

Conventions

Note that it is conventional for basic compounds to quote a pK_a which refers to the pK_a of the conjugate acid rather than that of the basic compound itself, likewise for pK_b when discussing acids. Thus a low pK_a quoted for a basic compound means that the conjugate acid readily dissociates and therefore that the basic compound is likely to be neutral at physiological pH. This can be contrasted with a low pK_a for an acid, which means that the acid readily dissociates and the molecule is therefore likely to exist as its conjugate base at physiological pH. To interpret literature pK_a and pK_b values an inspection of the molecule to determine whether it contains acid or base functionality (or both) is important.

Estimating blood pH

The Henderson–Hasselbalch equation can be applied to relate the pH of blood to constituents of the bicarbonate buffering system:^[4]

{\mathrm {pH}}={\mathrm {p}}K_{{{\mathrm {a}}~{\mathrm {H}}_{2}{\mathrm {CO}}_{3}}}+\log _{{10}}\left({\frac {[{\mathrm {HCO}}_{3}^{-}]}{[{\mathrm {H}}_{2}{\mathrm {CO}}_{3}]}}\right),

where:

$p K a H 2 CO 3$ is the cologarithm of the acid dissociation constant of carbonic acid. It is equal to 6.1.
$[HCO 3 -]$ is the concentration of bicarbonate in the blood
$[H 2 CO 3]$ is the concentration of carbonic acid in the blood

This is useful in arterial blood gas, but these usually state $p CO 2$ , that is, the partial pressure of carbon dioxide, rather than $H 2 CO 3$ . However, these are related by the equation:^[4]

[{\mathrm {H}}_{2}{\mathrm {CO}}_{3}]=k_{{{\rm {H~CO_{2}}}}}\,\times p_{{{\mathrm {CO}}_{2}}},

where:

$[H 2 CO 3]$ is the concentration of carbonic acid in the blood
$k H CO 2$ is the Henry's law constant for the solubility of carbon dioxide in blood. k_{H CO₂} is approximately 0.0307 mmol/(L-torr)
$p CO 2$ is the partial pressure of carbon dioxide in the blood

Taken together, the following equation can be used to relate the pH of blood to the concentration of bicarbonate and the partial pressure of carbon dioxide:^[4]

{\mathrm {pH}}=6.1+\log _{{10}}\left({\frac {[{\mathrm {HCO}}_{3}^{-}]}{0.0307\times p_{{{\mathrm {CO}}_{2}}}}}\right),

where:

$pH$ is the acidity in the blood
$[HCO 3 -]$ is the concentration of bicarbonate in the blood
$p CO 2$ is the partial pressure of carbon dioxide in the arterial blood

References

↑ Larsen, D. "Henderson-Hasselbalch Approximation". Chemwiki. University of California. Retrieved 27 March 2014.
↑ Henderson Hasselbalch Equation: Derivation of pK_a and pK_b
↑ "Acid-Base Tutorial - History". www.acid-base.com. Retrieved 2016-04-04.
1 2 3 page 556, section "Estimating plasma pH" in: Bray, John J. (1999). Lecture notes on human physiolog. Malden, Mass.: Blackwell Science. ISBN 978-0-86542-775-4.

External links

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Henderson–Hasselbalch equation

For bases

Derivation

History

Limitations

Conventions

Estimating blood pH

See also

References

Further reading

External links