Acid dissociation constant - Wikipedia, the free encyclopedia Page 1 Help us provide free content to the world by donating today ! Acid dissociation constant From Wikipedia, the free encyclopedia An acid dissociation constant (aka acidity constant, acid-ionization constant) is an equilibrium constant for the dissociation of an acid. It is denoted by Ka. For an equilibrium between a generic acid, HA, and − its conjugate base, A , The weak acid acetic acid donates a proton to water in an equilibrium reaction to give the acetate ion and − + HA A + H the hydronium ion. Key: Hydrogen is white, oxygen is red, carbon is gray. Lines are chemical bonds. K is defined, subject to certain conditions, as a where [HA], [A−] and [H+] are equilibrium concentrations of the reactants. The term acid dissociation constant is also used for pKa, which is equal to −log 10 Ka. The term pKb is used in relation to bases, though pKb has faded from modern use due to the easy relationship available between the strength of an acid and the strength of its conjugate base. Though discussions of this topic typically assume water as the solvent, particularly at introductory levels, the Brønsted–Lowry acid-base theory is versatile enough that acidic behavior can now be characterized even in non-aqueous solutions. The value of pK indicates the strength of an acid: the larger the value the weaker the acid. In aqueous a solution, simple acids are partially dissociated to an appreciable extent in in the pH range pK ± 2. The a actual extent of the dissociation can be calculated if the acid concentration and pH are known. A knowledge of pKa values is essential for the understanding of the behaviour of acids and bases in solution. For example, many compounds used for medication are weak acids or bases, so a knowledge of the pKa and log p values is essential for an understanding of how the compound enters (or does not enter) the blood stream. Other applications include aquatic chemistry, chemical oceanography, buffer solutions, acid-base homeostasis and certain kinds of enzyme kinetics, such as Michaelis–Menten kinetics, which involve a pre-equilibrium step. Also, knowledge of pK values is a prerequisite for a quantitative a understanding of the interaction between acids or bases and metal ions to form complexes in solution. Acids and bases: Contents Acid dissociation constant Acid-base extraction Acid-base reaction 1 Definitions Acid-base catalysis 2 Equilibrium Constant Acid-base physiology Acid-base homeostasis 2.1 Monoprotic acids Acidity function 2.2 Polyprotic acids Buffer solution Dissociation constant 2.3 Water self-ionization Non-nucleophilic base 2.4 Bases pH Proton affinity 2.5 Temperature dependence Self-ionization of water 3 Acidity in nonaqueous solutions http://en.wikipedia.org/wiki/Acid_dissociation_constant 2008-09-24 18:35 Acid dissociation constant - Wikipedia, the free encyclopedia Page 2 3.1 Mixed solvents Lewis acid/base Mineral acid/base 4 Factors that determine the relative strengths of acids Organic acid/base 4.1 Thermodynamics Weak acid/base Strong acid/base 5 Experimental determination of pK a values Super acid/base 6 Importance of pK a values 7 pK a of some common substances 8 See also 9 References 10 Further reading 11 External links Definitions According to Arrhenius's original definition, an acid is a substance Concepts in [1] which dissociates in aqueous solution, releasing the hydrogen ion. Chemical Equilibria Acid dissociation constant − + HA A + H Binding constant The equilibrium constant for this "dissociation" reaction is known as a Buffer solution dissociation constant. However, since the liberated proton combines Chemical equilibrium with a water molecule to give an hydronium ion, Arrhenius proposed Chemical stability that the "dissociation" reaction should be written as an acid-base Dissociation constant reaction. Distribution coefficient Distribution ratio − + HA + H2O A + H3O Equilibrium constant Equilibrium unfolding Brønsted and Lowry generalized this definition as a proton exchange Equilibrium stage [1] reaction, as follows. Liquid-liquid extraction Phase diagram acid + base conjugate base + conjugate acid Phase rule The acid donates a proton to the base. The conjugate base is what is left Reaction quotient after the acid has lost a proton and the conjugate acid is created when Relative volatility the base gains a proton. For aqueous solutions an acid, HA, reacts with Solubility equilibrium the base, water, donating a proton to it, creating the conjugate base, A−, Stability constant and the conjugate acid, the hydronium ion. The Brønsted–Lowry Thermodynamic equilibrium definition is particularly useful when the solvent is a substance other Theoretical plate than water, such as DMSO; in that case the solvent, S, acts as a base, Vapor-liquid equilibrium accepting a proton and forming the conjugate acid SH +. It also puts acids and bases on the same footing as being, respectively, donors or acceptors of protons. The conjugate acid of a base, B, "dissociates" according to BH + + OH − B + H O 2 For example: − + H2CO 3 + H2O HCO 3 + H3O The bicarbonate ion is the conjugate base of carbonic acid. http://en.wikipedia.org/wiki/Acid_dissociation_constant 2008-09-24 18:35 Acid dissociation constant - Wikipedia, the free encyclopedia Page 3 HCO − + OH − CO 2− + H O 3 3 2 and the bicarbonate ion is also the conjugate acid of the base, the carbonate ion. In fact the bicarbonate ion is amphiprotic. These reactions are important for acid-base homeostasis in the human body (see carbonic acid). Any compound subject to an hydrolysis equilibrium can also be classed as a weak acid since, in hydrolysis, protons are produced by the splitting of water molecules. For example, the equilibrium - + B(OH)3 + 2 H2O B(OH)4 + H3O shows why boric acid behaves as a weak acid even though it is not, itself, a proton donor. In a similar way, metal ion hydrolysis causes ions such as [Al(H O) ]3+ to behave as weak acids.[2] 2 6 It is important to note that, in the context of solution chemistry, a "proton" is understood to mean a solvated hydrogen ion. In aqueous solution the "proton" is a solvated hydronium ion.[3][4] Equilibrium Constant An acid dissociation constant is a particular example of an equilibrium constant. For the specific equilibrium betwen a monoprotic acid, HA and its conjugate base A−, in water, − + HA + H2O A + H3O the thermodynamic equilibrium constant, Kt can be defined by [5] where {A} is the activity of the chemical species A etc (activity is a dimensionless quantity). Activities of the products are placed in the numerator, activities of the reactants are placed in the denominator. See Chemical equilibrium for a derivation of this expression. Since activity is the product of concentration and activity coefficient the definition could also be written as Variation of pKa of acetic acid with ionic strength where [HA] represents the concentration of HA and Γ is a quotient of activity coefficients. http://en.wikipedia.org/wiki/Acid_dissociation_constant 2008-09-24 18:35 Acid dissociation constant - Wikipedia, the free encyclopedia Page 4 In order to avoid the complications involved in using activities, dissociation constants are determined, where possible, in a medium of high ionic strength, that is, under conditions in which Γ can be assumed to be always constant.[5] For example, the medium might be a solution of 0.1 M sodium nitrate or 3 M potassium perchlorate. Furthermore, in all but the most concentrated solutions it can be assumed that the −3 concentration of water, [H2O], is constant, approximately 55 mol dm , and that the hydration of the proton can also be assumed to be constant. Leaving out the constant terms, the acid dissociation constant can be defined as a concentration quotient. This is the definition in common use. pK is defined as −log K . Note, however, that all published a 10 a dissociation constant values refer to the specific ionic medium used in their determination and that different values are obtained with different conditions. When operating under the assumption that Γ is constant, the equilibrium constant does not change upon the addition of other chemicals to the solution. This assumption holds true when the concentration of spectator ions is low relative to the concentrations of other ions in the system. This allows, for example, for the behaviour of various ions to be explored at various pH values without worry that the equilibrium constant will also change. By exploiting this property, it is possible to obtain very complicated buffer solutions composed of many protonations of the same anion. This is accomplished with the addition of a strong acid to a solution of the anion. The conjugate base of the strong acid will act as a spectator ion, and the weak-base anion will be free to react with the proton as the equilibrium constant dictates. Monoprotic acids After rearranging the expression defining Ka, and putting pH = + −log 10 [H ], one obtains pH = pK – log ( [AH]/[A−]) a This is a form of the Henderson–Hasselbalch equation. It shows how if the pH is known the ratio [AH]:[A−] may be calculated. This ratio is independent of the analytical concentration of the acid. Variation of the % formation of a if the ratio [AH]:[A−] is known the pH may be calculated. monoprotic acid, AH, and its Thus, at 50% neutralization pH =pK ([AH]:[A−] = 1). The conjugate base, A−, with the a difference between the pH and the buffer region extends over the range pK ± 2, though buffering a pKa of the acid is weak outside the range pKa ± 1. In water, measurable pKa values range from about –2 for a strong acid to about 12 for a very weak acid (or strong base).