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Acid dissociation constant
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. It is important to remember that the term [H2O] is omitted from the general equilibrium constant expression. While strong acids dissociate practically completely in solution and consequently have large acidity constants, weak acids do not fully dissociate and generally have acidity constants far less than 1. 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 (similar to the concept of pH, though not related directly).
pKa = −log10Ka
Given a weak acid HA, its dissociation in water is subject to the following chemical equilibrium:
HA + H2O ↔ H3O+ + A– (it is also acceptable to write this as: HA ↔ H+ + A–, the difference being only what theory of acids/bases you are applying. See the Bronsted-Lowry Theory and the Arrhenius Theory for more information)
The acidity constant for the acid HA is the dissociation constant for this equilibrium. In other words,
, where [A] denotes the molar concentration of A in the solution
Using this definition, chemists can quickly and easily determine the concentrations of various chemicals in an equilibrium. For example, to determine the pH of a solution of known molar concentrations of sodium hydroxide and hydrofluoric acid, if the Ka of the acid at the given temperature (which is easily attainable information) is known, the concentration of hydrogen ions can be calculated, which will allow the determination of the pH after taking into account the neutralization due to the base. This procedure is summed up in the Henderson-Hasselbalch equation.
Acid dissociation constant - Basicity constant of the conjugate base
By analogy, one can define the basicity constant (Kb) and the pKb of the conjugate base A–:
pKb = −log10Kb
This is the dissociation constant for the equilibrium
A– + H2O ↔ HA + OH–
Analogously to Ka, the magnitude of Kb indicates the relative strength of the base, with Kb < < 1 indicating a strong base.
Acid dissociation constant - 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 M2 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, and vice versa.
Acid dissociation constant - The Relative Strengths of Acids and Bases
The strengths of acids and bases in aqueous solutions are determined by the values of the dissociation constants Ka and Kb for acids and bases, respectively, and by their molar concentrations in solution. Outside of an aqueous solution, it is much more difficult to define the strengths of acids and bases (see Acid-base reaction theories for more information on Brønsted-Lowry, Arrhenius, and Lewis acids and bases.) For the purposes of this article, it is best to assume that any acids and bases are in solution, as then the Brønsted-Lowry definition is prevalent.
In terms of weak acids and weak bases, the relative strengths between solutions of the same molarity can be determined by inspecting the dissociation constants. A high acid (base) dissociation constant describes a chemical that more readily donates (accepts) a proton to the solvent, and hence forms a stronger acid. In the reaction between acetic acid and water,
approximately 1 in every 10,000 acetic acid molecules dissociate into the acetate ion and a hydronium ion. However, in the strong acid reaction
the hydrochloric acid dissociates almost completely (since it has a Ka value that is very large), and hence more hydronium ions appear in solution per mole of hydrochloric acid than did for acetic acid.
The same discussion above can be applied to bases as well. Sodium hydroxide (NaOH), for example, is a strong base, since it almost completely dissociates in water into a sodium ion (Na + ) and a hydroxide ion (OH − ), while ammonia is a weak base with Kb = 1.8 × 10 - 5 (coincidentally, the same value as the Ka for acetic acid above).
Other factors that influence the strength of the acid or base are the polarity of the ionic bond holding the hydrogen atom to the parent (conjugate base or acid, respectively) ion, the size of the parent ion, the total charge on the molecule, and the oxidation state of the central atom. A more polar bond is easier to dissociate, and so highly polar molecules, such as hydrochloric acid (HCl) are very strong acids (or bases), while relatively non-polar molecules (such as ammonia) are weaker. The size of the central atom also defines the energy required to dissociate the hydrogen atom: a large, diffuse ion (such as chlorine or iodine) will have a weaker ionic bond with a hydrogen ion than would, for example, a flourine ion. This principle is based on the electronegativity of the ions involves. 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.
Acid dissociation constant - pKa of some common substances
Measurements are at 25ºC in water:
- 3.60: Carbonic acid
- 3.75: Formic acid
- 4.19: Succinic acid
- 4.20: Benzoic acid
- 4.63: Aniline*
- 4.74: Acetic acid
- 4.76: Dihydrogencitrate ion (Citrate)
- 5.21: Pyridine*
- 6.40: Monohydrogencitrate ion (Citrate)
- 6.99: Ethylenediamine*
- 7.00: Hydrogen sulfide and Imidazole*
- 9.25: Ammonia*
- 9.33: Benzylamine*
- 9.81: Trimethylamine*
- 9.99: Phenol
- 10.08: Ethylenediamine*
- 10.66: Methylamine*
- 10.73: Dimethylamine*
- 10.81: Ethylamine*
- 11.01: Triethylamine*
- 11.09: Diethylamine*
- 12.67: Monohydrogenphosphate ion (Phosphate)
* Listed values for ammonia and amines are the pKa values for the corresponding ammonium ions.
Other related archivesAcetic acid, Acid-base reaction theories, Ammonia, Aniline, Arrhenius Theory, Benzoic acid, Bronsted-Lowry Theory, Carbonic acid, Citrate, Dimethylamine, Ethylenediamine, Formic acid, Henderson-Hasselbalch equation, Hydrogen sulfide, Imidazole, M, Methylamine, Phenol, Phosphate, Pyridine, Sodium hydroxide, Succinic acid, Triethylamine, Trimethylamine, acetic acid, acid, acids, additive inverse, ammonia, aqueous, bases, biochemistry, charge, chemical equilibrium, chemistry, chlorine, common logarithm, concentration, conjugate base, dissociation, dissociation constant, electronegativity, equilibrium constant, flourine, hydrochloric acid, hydrofluoric acid, hydrogen ions, hypochlorous acid, iodine, ionic bond, molar, molarity, oxidation number, oxidation state, pH, perchloric acid, polarity, self-ionization of water, sodium hydroxide, strong acid, strong acids, strong base, water, weak acids, weak bases
 Adapted from the Wikipedia article "Acid dissociation constant", under the G.N U Free Docmentation License. Please also see http://en.wikipedia.org/wiki |
