Acid dissociation constant

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In chemistry and biochemistry, acid dissociation constant, the acidity constant, or the acid-ionization constant (Ka) is a specific type of equilibrium constant that indicates the extent of dissociation of hydrogen ions from an acid. The equilibrium is that of a proton transfer from an acid, HA, to water, H2O. The term for the concentration of water, [H2O], is omitted from the general equilibrium constant expression.

HA(aq) + H2O(l) ⇌ H3O+(aq) + A–(aq)
<math>K_a = \frac{[\mbox{H}_3\mbox{O}^+][\mbox{A}^- ]} {[\mbox{HA}]}</math>

The equilibrium is often written in terms of "H+(aq)", which reflects the Bronsted-Lowry Theory of acids.

HA(aq) ⇌ H+(aq) + A–(aq)

Because this constant differs for each acid and varies over many degrees of magnitude, the acidity constant is often represented by the additive inverse of its common logarithm, represented by the symbol pKa (using the same mathematical relationship as [H+] is to pH).

pKa = −log10 Ka

In general, a larger value of Ka (or a smaller value of pKa) indicates a stronger acid, since the extent of dissociation is larger at the same concentration.

Using the acid dissociation constants, the concentration of acid, its conjugate base, protons and hydroxide can be easily determined. If an acid is partly neutralized, the Ka can also be used to find the pH of the resulting buffer. This same information is summarized in the Henderson-Hasselbalch equation.

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Basicity constant of the conjugate base

By analogy, one can define the basicity constant Kb and the pKb of the conjugate base A–:

<math>K_b = \frac{[\mbox{HA}][\mbox{OH}^-]} {[\mbox{A}^-]}</math>
pKb = −log10 Kb

This is the dissociation constant for the equilibrium

A–(aq) + H2O(l) ⇌ HA(aq) + OH–(aq)

Analogously to Ka, an increasing value of Kb indicates a stronger base, since the number of protons accepted is larger at an identical concentration.

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Relationship between acidity and basicity constants

There exists a relationship between the value of Ka for an acid HA and the value of Kb for its conjugate base A–. Since adding the ionization reaction for HA and the ionization reaction of A– always gives the reaction for the self-ionization of water, the product of the acidity and basicity constants gives the dissociation constant of water (Kw), which is 1.0 × 10-14 at 25°C. In other words,

KaKb = Kw
pKa + pKb = pKw

As the product of Ka and Kb remains constant, it follows that STRONGER acids have WEAKER conjugate bases, while WEAKER acids have STRONGER conjugate bases.

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Factors that determine the relative strengths of acids and bases


to dissociate the hydrogen atom: a large, diffuse ion (such as chloride or iodide) will have a weaker ionic bond with a hydrogen ion than would, for example, a fluoride ion. This principle is based on the electronegativity of the ions involved. Positively-charged ions are typically acids, because they are able to readily donate protons, while negatively-charged ions are classified as bases due to their ability to receive free protons. Yet another factor that influences the dissociation of acids and bases is the oxidation number on the central atom in the molecule. Higher oxidation numbers yield stronger acids, as can be shown with the sequence of acids (in order of ascending Ka): HClO < HClO2 < HClO3 < HClO4. The difference in values of Ka between perchloric acid and hypochlorous acid is approximately 11 orders of magnitude.

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pKa of some common substances

Measurements are at 25ºC in water:

* Listed values for ammonia and amines are the pKa values for the corresponding ammonium ions.

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External links

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Further reading

Atkins, Peter, and Loretta Jones. Chemical Principles: The Quest for Insight. 3rd ed. New York: W. H. Freeman and Company, 2005de:Säurekonstante nl:Zuurconstante ru:Константа диссоциации кислоты fi:Happovakio sv:Syrakonstanten uk:PKa

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