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In chemistry and biochemistry, the 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) 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. Basicity constant of the conjugate base By analogy, one can define the basicity constant Kb and the pKb of the conjugate base A–: 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. 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. Factors that determine the relative strengths of acids and bases Being an equilibrium constant, the acid dissociation constant Ka is determined by the difference in free energies ΔG° between the reactants and products, specifically, between the protonated (AH) and de-protonated (A–) states. Molecular interactions that favor the deprotonated (A–) state over the protonated (AH) state will increase Ka (because the ratio A–/AH increases) or, equivalently, decrease pKa. Conversely, molecular interactions that favor the protonated (AH) state over the de-protonated (A–) state will decrease Ka (because the ratio A–]/AH is lower) or, equivalently, increase pKa. For example, suppose that the protonated (AH) form donates a hydrogen bond AHX to another atom X, which the de-protonated form cannot do (since it has no hydrogen left). The protonated form is favored by having a hydrogen bond, so the pKa increases (the Ka decreases). The magnitude of the pKa shift can even be determined from the change in ΔG° using the equation . Other molecular interactions can also shift the pKa. Adding an electron-withdrawing chemical group (such as oxygen, a halide, a cyano group or even a phenyl ring) to the molecule near the titrating hydrogen will favor the deprotonated state (by stabilizing the electron left behind when the proton dissociates) and thus decrease pKa (increase Ka). For example, successive oxidation of hypochlorous acid leads to ever-increasing Ka: HClO < HClO2 < HClO3 < HClO4. The difference in values of Ka between hypochlorous acid HClO and perchloric acid HClO4 is approximately 11 orders of magnitude (pKa shift of ~11). Electrostatic interactions can affect the equilibrium as well. The presence of surrounding negative charges would disfavor the formation of a negatively charged, de-protonated species and thus increase pKa. In particular, the ionization of one group on a molecule can affect the pKa of another. Fumaric and maleic acid are classic examples of pKa shifts. Both molecules have the same composition, being two carboxylic acid groups separated by two double-bonded carbon atoms; fumaric acid is the trans isomer, whereas maleic acid is the cis isomer. By symmetry, one might imagine that the two carboxylic acids had the same pKa, which is typically ~4 for carboxylic acids. This is almost true for fumaric acid, which has pKa's of roughly 3.5 and 4.5. By contrast, maleic acid has pKa's of roughly 1.5 and 6.5. When one of its carboxylic acids de-protonates, the other can form a strong hydrogen bond to it; overall, the effect is to favor the de-protonated state of the hydrogen-bond-accepting group (lowering its pKa from ~4 to 1.5) and to favor the protonated state of the hydrogen-bond-donating group (raising its pKa from ~4 to 6.5). Importance of pK The pKa value(s) of a compound influence many characteristics of the compound such as its reactivity, solubility and spectral properties (colour). In biochemistry the pKa values of proteins and amino acid side chains are of major importance for the activity of enzymes and the stability of proteins. See Methods for calculating protein pKa values pK Measurements are at 25ºC in water, except for those with a pKa below -1.76: Further reading Atkins, Peter, and Loretta Jones. Chemical Principles: The Quest for Insight. 3rd ed. New York: W. H. Freeman and Company, 2005 | |||||||
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