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14-September-2008 18:38:37 - Acid dissociation constant The weak acid acetic acid donates a proton to water in an equilibrium reaction to give the acetate ion and the hydronium ion. Key: Hydrogen is white, oxygen is red, carbon is gray. Lines are chemical bonds. The weak acid acetic acid donates a proton to water in an equilibrium reaction to give the acetate ion and the hydronium ion. Key: Hydrogen is white, oxygen is red, carbon is gray. Lines are chemical bonds. 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-, HA \rightleftharpoons A- + H+ Ka is defined, subject to certain conditions, as K_a = \mathrm\fracA^-H^+HA 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 -log10 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 pKa indicates the strength of an acid: the larger the value the weaker the acid. In aqueous solution, simple acids are partially dissociated to an appreciable extent in in the pH range pKa ± 2. The 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 pKa values is a prerequisite for a quantitative understanding of the interaction between acids or bases and metal ions to form complexes in solution. Acids and bases: Acid dissociation constant Acid-base extraction Acid-base reaction Acid-base catalysis Acid-base physiology Acid-base homeostasis Acidity function Buffer solution Dissociation constant Non-nucleophilic base pH Proton affinity Self-ionization of water Lewis acid/base Mineral acid/base Organic acid/base Weak acid/base Strong acid/base Super acid/base Contents 1 Definitions 2 Equilibrium Constant 2.1 Monoprotic acids 2.2 Polyprotic acids 2.3 Water self-ionization 2.4 Bases 2.5 Temperature dependence 3 Acidity in nonaqueous solutions 3.1 Mixed solvents 4 Factors that determine the relative strengths of acids 4.1 Thermodynamics 5 Experimental determination of pKa values 6 Importance of pKa values 7 pKa of some common substances 8 See also 9 References 10 Further reading 11 External links Definitions Concepts in Chemical Equilibria Acid dissociation constant Binding constant Buffer solution Chemical equilibrium Chemical stability Dissociation constant Distribution coefficient Distribution ratio Equilibrium constant Equilibrium unfolding Equilibrium stage Liquid-liquid extraction Phase diagram Phase rule Reaction quotient Relative volatility Solubility equilibrium Stability constant Thermodynamic equilibrium Theoretical plate Vapor-liquid equilibrium According to Arrhenius's original definition, an acid is a substanstance which dissociates in aqueous solution, releasing the hydrogen ion.1 HA \rightleftharpoons A- + H+ The equilibrium constant for this dissociation reaction is known as a dissociation constant. However, since the liberated proton combines with a water molecule to give an hydronium ion, Arrhenius proposed that the dissociation reaction should be written as an acid-base reaction. HA + H2O \rightleftharpoons A- + H3O+ Brønsted and Lowry generalized this definition as a proton exchange reaction, as follows.1 acid + base \rightleftharpoons conjugate base + conjugate acid The acid donates a proton to the base. The conjugate base is what is left after the acid has lost a proton and the conjugate acid is created when the base gains a proton. For aqueous solutions an acid, HA, reacts with the base, water, donating a proton to it, creating the conjugate base, A-, and the conjugate acid, the hydronium ion. The Brønsted-Lowry definition is particularly useful when the solvent is a substance other than water, such as DMSO; in that case the solvent, S, acts as a base, 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- \rightleftharpoons B + H2O For example: H2CO3 + H2O \rightleftharpoons HCO3- + H3O+ The bicarbonate ion is the conjugate base of carbonic acid. HCO3- + OH- \rightleftharpoons CO32- + H2O 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 BOH3 + 2 H2O \rightleftharpoons BOH4- + 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 AlH2O63+ to behave as weak acids.2 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.34 Equilibrium Constant Main article: 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 \rightleftharpoons A- + H3O+ the thermodynamic equilibrium constant, Kt can be defined by5 K^\mboxt=\mathrm\frac\A^-\ \H_3O^+\ \HA\ \H_2O\ 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. Variation of pKa of acetic acid with ionic strength Variation of pKa of acetic acid with ionic strength Since activity is the product of concentration and activity coefficient the definition could also be written as K^\mboxt = \mathrm\fracA^-H_3O^+HAH_2O\times \frac\gamma_A^-\gamma_H_3O^+\gamma_HA\gamma_H_2O =\mathrm\fracA^-H_3O^+HAH_2O\times\Gamma where HA represents the concentration of HA and Γ is a quotient of activity coefficients. 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 concentration of water, H2O, is constant, approximately 55 mol dm-3, 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. K_a = \mathrm\fracA^-H^+HA This is the definition in common use. pKa is defined as -log10 Ka. Note, however, that all published 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. Variation of the % formation of a monoprotic acid, AH, and its conjugate base, A-, with the difference between the pH and the pKa of the acid Variation of the % formation of a monoprotic acid, AH, and its conjugate base, A-, with the difference between the pH and the pKa of the acid Monoprotic acids After rearranging the expression defining Ka, and putting pH = -log10H+, one obtains pH = pKa - log AH/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. if the ratio AH:A- is known the pH may be calculated. Thus, at 50% neutralization pH =pKa AH:A- = 1. The buffer region extends over the range pKa ± 2, though buffering 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. Any acid with a pKa value of less than -2 is more than 99% dissociated at pH 0 1M acid. Any base with a pKa value larger than the upper limit is fully de-protonated at all attainable pH values. This is known as solvent leveling.6 An example of a strong acid is hydrochloric acid, HCl, which has has a pKa value, estimated from thermodynamic quantitities, of -9.3 in water.7 The concentration of undissociated acid in a 1 mol dm-3 solution, will be less than 10-4 mol dm-3. In common parlance this is known as complete dissociation. The extent of dissociation and pH of a solution of a monoprotic acid can be easily calculated when the pKa and analytical concentration of the acid are known. See ICE table for details. Polyprotic acids % species' formation as a function of pH % species' formation as a function of pH % species formation calculated with the program HySS for a 10mM solution of citric acid. pKa1=3.13, pKa2 = 4.76, pKa3=6.40. % species formation calculated with the program HySS for a 10mM solution of citric acid. pKa1=3.13, pKa2 = 4.76, pKa3=6.40. Polyprotic acids are acids which can lose more than one proton. The constant for dissociation of the first proton may be denoted as Ka1 and the constants for dissociation of successive protons as Ka2, etc. When the difference between succesive pK values is about four or more, each species may be considered as an acid in its own right;8 the pH range of existence of each species is about pK± 2, so there is very little overlap between the ranges for successive species. The case of phosphoric acid illustrates this point. In fact salts of either H2PO4- or HPO42- may be crystallized from solution by adjustment of pH to either 4 or 10. When the difference between succesive pK values is less than about four there is overlap between the pH range of existence of the species in equilibrium. The smaller the difference, the more the overlap. The case of citric acid is shown at the right; solutions of citric acid are buffered over the whole range of pH 2.5 to 7.5. It is generally true that successive pK values increase Pauling's first rule.9 For example, for a diprotic acid, H2A, the two equilibria are H2A \rightleftharpoons HA- + H+ HA- \rightleftharpoons A2- + H+ it can be seen that the second proton is removed from a negatively charged species. Since the proton carries a positive charge extra work is needed to remove it; that is the cause of the trend noted above. Phosphoric acid, H3PO4, values below, illustrates this rule, as does vanadic acid. When an exception to the rule is found it indicates that a major change in structure is ocurring. In the case of VO2+aq, the vanadium is octahedral, 6-coordinate, whereas all the other species are tetrahedral, 4-coordinate. This explans why pKa1 pKa2 for vanadiumV oxoacids. VO2+ \rightleftharpoons H3VO4 + H+ pKa1 = 4.2 H3PO4 \rightleftharpoons H2PO4- + H+ pKa1 = 2.15 H3VO4 \rightleftharpoons H2VO4- + H+ pKa2 = 2.60 H2PO4- \rightleftharpoons HPO42- + H+ pKa2 = 7.20 H2VO4- \rightleftharpoons HVO42- + H+ pKa3 = 7.92 HPO42- \rightleftharpoons PO43- + H+ pKa3 = 12.37 HVO42- \rightleftharpoons VO43- + H+ pKa4 = 13.27 Water self-ionization Water has both acidic and basic properies. The equilibrium constant for the equilibrium H2O + H2O \rightleftharpoons OH- + H3O+ is given by K_a=\mathrm\fracH^+OH^-H_2O^2 Since the concentration of water can be assumed to be constant, this expression simplifies to K_w =H^+OH^-\, The self-ionization constant of water, Kw, can thus be seen as a special case of an acid dissociation constant. Bases Historically the equilibrium constant Kb for a base was defined as the association constant for protonation of the base, B, to form the conjugate acid, HB+. B + H2O \rightleftharpoons HB+ + OH- Using similar reasoning to that used before K_b = \mathrm\fracHB^+OH^-B In water, the concentration of the hydroxide ion, OH-, is related to the concentration of the hydrogen ion by Kw = H+OH-, therefore \mathrmOH^- = \frac\mathitK_wH^+ Substitution of the expression for OH- into the expression for Kb gives \mathrm\mathitK_b = \fracHB^+\mathitK_wB H^+ = \frac\mathitK_w\mathitK_a It follows, taking cologarithms, that pKb = pKw - pKa. In aqueous solutions at 25 °C, pKw is 13.9965,10 so pKb ~ 14 - pKa. In effect there is no need to define pKb separately from pKa, but it is done here because pKb values can be found in the older literature. Temperature dependence All equilibrium constants vary with temperature according the van 't Hoff equation11 \frac \operatornamed \ln \mathitK \operatornamedT = \frac\Delta \mathitH_m^\ominus RT^2 Thus, for exothermic reactions, ΔHO is negative K decreases with temperature, but for endothermic reactions ΔHO is positive K increases with temperature. Acidity in nonaqueous solutions A solvent will be more likely to promote ionization of a dissolved acidic molecule if:12 it is a protic solvent, capable of forming hydrogen bonds it has a high donor number, making it a strong Lewis base. it has a high dielectric constant relative permittivity, making it a good solvent for ionic species. Solvents can be polar, protic, donor or non-polar. The data in the following table refer to a temperature at or near 25 °C, unless stated otherwise.12 Compound Solvent Class Dielectric constant 1,4-Dioxane Non-polar, Donor 2.2 Hexane Non-polar 1.9 Carbon tetrachloride Non-polar 2.2 Benzene Non-polar 2.3 Diethyl ether Non-polar, Donor 4.3 Acetic acid Protic donor 6.1 Tetrahydrofuran Donor 7.6 Acetone Polar donor 21 Liquid ammonia Polar donor 25 at 195 K Acetonitrile Polar donor 37 Dimethylsulfoxide Polar donor 47 Water Polar protic donor 78 Formamide Polar protic donor 111 Sulphuric acid Polar protic 110 Ionization of acids is less in an acidic solvent than in water. For example, hydrogen chloride is a weak acid when dissolved in acetic acid. This is because acetic acid is a much weaker base than water. HCl + CH3CO2H \rightleftharpoons Cl- + CH3COH2+ acid + base \rightleftharpoons conjugate base + conjugate acid Compare this reaction with what happens when acetic acid is dissolved in the more acidic solvent pure sulphuric acid13 H2SO4 + CH3CO2H \rightleftharpoons HSO4- + CH3COH2+ The apparently unlikely geminal diol species CH3COH2+ is stable in these environments. pKa values of organic compounds are often obtained using solvents other than water, such as dimethyl sulfoxide DMSO and acetonitrile.14 Water is more basic than DMSO so most acids dissociate to a lesser extent in DMSO than in water. DMSO is widely used as an alternative to water in evaluating acids and bases because it has a lower dielectric constant than water, it is less polar and so dissolves non-polar, hydrophobic substances more easily. Below is a table of selected pKa values at 25 °C in acetonitrile AN151617 and dimethyl sulfoxide DMSO.18 Values for water are included for comparison. HA \rightleftharpoons A- + H+ AN DMSO water p-Toluenesulfonic acid 8.5 0.9 strong 2,4-Dinitrophenol 16.66 5.1 3.9 Benzoic acid 21.51 11.1 4.2 Acetic acid 23.51 12.6 4.756 Phenol 29.14 18.0 9.99 BH+ \rightleftharpoons B + H+ AN DMSO water Pyrrolidine 19.56 10.8 11.4 Triethylamine 18.82 9.0 10.72 Proton sponge 18.62 7.5 12.1 Pyridine 12.53 3.4 5.2 Aniline 10.62 3.6 9.4 In solvents of low dielectric constant ions tend to associate forming ion pairs and clusters, which complicates the interpretation of pKa values. dimerization of a carboxylic acid dimerization of a carboxylic acid In aprotic solvents, oligomers may be formed by hydrogen bonding. An acid may also form hydrogen bonds to its conjugate base. This process is known as homoconjugation. Homoconjugation has the effect of enhancing the acidity of acids, lowering their effective pKa values, by stabilizing the conjugate base. Due to homoconjugation, the proton-donating power of toluenesulfonic acid in acetonitrile solution is enhanced by a factor of nearly 800.19 Homoconjugation does not occur in aqueous solutions because water forms stronger hydrogen bonds and prevents the oligomers from forming. Mixed solvents When a compound has limited solubility in water it is common practice in the pharmaceutical industry, for example to determine pKa values in a solvent mixture such as water/dioxane or water/methanol, in which the compound is more soluble.20 However, a pKa value obtained in a mixed solvent cannot be used directly for aqueous solutions. The reason for this is that when the solvent is in its standard state its activity is defined as one. For example, the standard state of water:dioxane 9:1 is precisely that solvent mixture, with no added solutes. To obtain the pKa value for use with aqueous solutions it has to be extrapolated to zero co-solvent concentration from values obtained from various co-solvent mixtures. These facts are obscured by the omission of the solvent from the expression which is normally used to define pKa, but pKa values obtained in a given mixed solvent can be compared to each other, giving relative acid strengths. The same is true of pKa values obtained in a particular non-aqueous solvent such a DMSO. A universal, solvent-independent, scale for acid dissociation constants has not yet been developed, since there is no known way to compare the standard states of two different solvents. Factors that determine the relative strengths of acids Pauling's second rule9 states that the value of the first pKa for acids of the formula XOmOH n is approximately independent of n and X and is approximately 8 for m = 0, 2 for m = 1, -3 for m = 2 and -10 for m = 3. This correlates with the oxidation state of the central atom, X: the higher the oxidation state the stronger the oxyacid. For example, pKa for HClO is 7.2, for HClO2 is 2.0, for HClO3 is -1 and HClO4 is a strong acid. fumaric acid fumaric acid maleic acid maleic acid With organic acids inductive effects and mesomeric effects affect the pK'a values. The effects are summarised in the Hammett equation and subsequent extensions.21 Structural effects can also be important. The difference between fumaric acid and maleic acid is a classic example. Fumaric acid is E-1,4-but-2-enedioic acid, a trans isomer, whereas maleic acid is the corresponding cis isomer, i.e. Z-1,4-but-2-enedioic acid see cis-trans isomerism. Fumaric acid has pKa values of approximately 3.5 and 4.5. By contrast, maleic acid has pKa values of approximately 1.5 and 6.5.22 The reason for this large difference is that when one proton is removed from the cis- isomer maleic acid a strong intramolecular hydrogen bond is formed with the nearby remaining carboxyl group. This favors the formation of the maleate H+, and it opposes the removal of the second proton from that species. In the trans isomer, the two carboxyl groups are always far apart, so hydrogen bonding is not observed. proton sponge proton sponge Proton sponge, 1,8-Bisdimethylaminonaphthalene, has a pKa value of 12.1. It is one of the strongest amine bases known. The high basicity is attributed to the relief of strain upon protonation and strong internal hydrogen bonding. Thermodynamics An equilibrium constant is related to the standard Gibbs free energy change for the reaction, so for an acid dissociation constant ΔGO = 2.303 RT pKa. Note that pKa= -log Ka. At 25 °C ΔGO /kJ mol-1 = 5.708 pKa. Free energy is made up of an enthalpy term and an entropy term.23 ΔGO = ΔHO - TΔSO The standard enthalpy change can be determined by calorimetry or by using the van't Hoff equation, though the calorimetric method is preferable. When both the standard enthalpy change and acid dissociation constant have been determined, the standard entropy change is easily calculated from the equation above. In the following table, the entropy terms are calculated from the experimental values of pKa and ΔHO. The data were critically selected and refer to 25 °C and zero ionic strength, in water.23 Acids Compound Equilibrium pKa ΔH0 /kJ mol-1 -TΔS0 /kJ mol-1 HA = Acetic acid HA \rightleftharpoons H+ + A- 4.756 -0.41 27.56 H2A+ = GlycineH+ H2A+ \rightleftharpoons HA + H+ 2.351 4.00 9.419 HA \rightleftharpoons H+ + A- 9.78 44.20 11.6 H2A = Maleic acid H2A \rightleftharpoons HA- + H+ 1.92 1.10 9.85 HA- \rightleftharpoons H+ + A2- 6.27 -3.60 39.4 H3A = Citric acid H3A \rightleftharpoons H2A- + H+ 3.128 4.07 13.78 H2A- \rightleftharpoons HA2- + H+ 4.76 2.23 24.9 HA2- \rightleftharpoons A3- + H+ 6.40 -3.38 39.9 HA = Boric acid HA \rightleftharpoons H+ + A- 9.237 13.80 38.92 H3A = Phosphoric acid H3A \rightleftharpoons H2A- + H+ 2.148 -8.00 20.26 H2A- \rightleftharpoons HA2- + H+ 7.20 3.60 37.5 HA2- \rightleftharpoons A3- + H+ 12.35 16.00 54.49 HA- = Hydrogen sulphate HA- \rightleftharpoons A2- + H+ 1.99 -22.40 33.74 H2A = Oxalic acid H2A \rightleftharpoons HA- + H+ 1.27 -3.90 11.15 HA- \rightleftharpoons A2- + H+ 4.266 7.00 31.35 Conjugate acid of bases Compound Equilibrium pKa ΔH0 /kJ mol-1 -TΔS0 /kJ mol-1 B = Ammonia HB+ \rightleftharpoons B + H+ 9.245 51.95 0.8205 B = Methylamine HB+ \rightleftharpoons B + H+ 10.645 55.34 5.422 B = Triethylamine HB+ \rightleftharpoons B + H+ 10.72 43.13 18.06 The first point to note is that when pKa is positive, the standard free energy change for the dissociation reaction is also positive, that is, dissociation of a weak acid is not a spontaneous process. Secondly some reactions are exothermic and some are endothermic, but when ΔHO is negative -TΔSO is the dominant factor which determines that ΔGO is positive. Lastly, the entropy contribution is always unfavourable in these reactions. Note. The standard free energy change for the reaction is for the changes from the reactants in their standard states to the products in their standard states. The free energy change at equilibrium is zero since the chemical potentials of reactants and products are equal at equilibrium. Experimental determination of pKa values Main article: Determination of equilibrium constants A calculated titration curve of oxalic acid titrated with a solution of sodium hydroxide A calculated titration curve of oxalic acid titrated with a solution of sodium hydroxide pKa values are commonly determined by means of titrations, in a medium of high ionic strength and at constant temperature.24 A typical procedure would be as follows. A solution of the compound in the medium is acidified with a strong acid to the point where the compound is fully protonated. The solution is then titrated with a strong base until all the protons have been removed. At each point in the titration pH is measured using a pH meter. The equilibrium constants are found by fitting calculated pH values to the observed values, using the method of least squares. The total volume of added strong base should be small compared to the initial volume of to keep the ionic strength nearly constant. This will ensure that pKa remains invariant during the titration. A calculated titration curve for oxalic acid is shown at the right. Oxalic acid has pKa values of 1.27 and 4.27. Therefore the buffer regions will be centered at about pH 1.3 and pH 4.3. The buffer regions carry the information necessary to get the pKa values as the concentrations of acid and conjugate base change along a buffer region. Between the two buffer regions there is an end-point, or equivalence point, where the pH rises by about two units. This end-point is not sharp and is typical of a diprotic acid whose buffer regions overlap by a small amount: pKa2 - pKa1 is about three in this example. If the difference in pK values were about two or less, the end-point would not be noticeable. The second end-point begins at about pH 6.3 and is sharp. This indicates that all the protons have been removed. When this is so, the solution is not buffered and the pH rises steeply on addition of a small amount of strong base. However, the pH does not continue to rise indefinitely. A new buffer region begins at about pH 11 pKw - 3, which is where self-ionization of water becomes important. It is very difficult to measure pH values of less than two with a glass electrode, because the Nernst equation breaks down at such low pH values. To determine pK values of less than about 2 or more than about 11 spectrophotometric25 or NMR26 measurements may be used instead of, or combined with pH measurements.27 Importance of pKa values A knowledge of pKa values is important for the quantitative treatment of systems involving acid-base equilibria in solution. Applications include: Biochemistry Further information: Protein pKa calculations 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.28 The reaction that converts adenosine triphosphate to adenosine diphosphate is very pH sensistive. Buffer solutions Main article: buffer solutions A buffer solution is made up of a mixture of an acid and its conjugate base, or a base and its conjugate acid. Compared with an aqueous solution, the pH of a buffer solution is relatively insensitive to the addition of a small amount af strong acid or strong base. The buffer capacity29 of a simple buffer solution illustrative diagram is largest when pH = pKa. Buffer solutions are used extensively in biochemistry to provide solutions at or near the physiological pH for the study of biochemical reactions.30 For example, MOPS provides a solution with pH 7.2; others are listed in buffer solutions and Good's buffers. Buffers such as tricine are used in Gel electrophoresis.31 32 Isoelectric focussing is a technique used for separation of proteins by 2-D gel polyacrylamide gel electrophoresis. Buffering is essential in Acid base physiology including Acid-base homeostasis33 and disorders such as Acid-base imbalance.343536 Coordination compounds Main article: Complex chemistry A coordination complex is formed by interaction of a metal ion, Mm+, acting as a Lewis acid, with a ligand, L, acting as a Lewis base. However, the ligand may also undergo protonation reactions, so the formation of a complex in aqueous solution could be represented, symbolically by the reaction MH2Onm+ +LH \rightleftharpoons MH2On-1Lm-1+ + H3O+ To determine the equilibrium constant for this reaction, in which the ligand loses a proton, the pKa of the protonated ligand must be known. In practice, the ligand may be polyprotic; for example EDTA4- can accept four protons; in that case, all pKa values must be known. In addition, the metal ion is subject to hydrolysis, that is, it behaves as a weak acid, so the pK values for the hydrolysis reactions must also be known.37 Solvent extraction Main article: Acid-base extraction In solvent extraction, the efficiency of extraction of a compound into an organic phase, such as ether, can be optimized by adjusting the pH of the aqueous phase using an appropriate buffer. At the optimum pH, the concentration of the electrically neutral species is maximized; such a species is more soluble in organic solvents having a low dielectric constant than it is in water. This technique is used for the purification of weak acids and bases.38 Natural waters Acid-base equilibria are important for rivers and lakes,3940 and in chemical oceanography.41 Pharmacology Ionization of a compound alters its physical behavior and macro properties such as solubility and lipophilicity log p. For example ionization of any compound will increase the solubility in water, but decrease the lipophilicity. This is exploited in drug development to increase the concentration of a compound in the blood by adjusting the pKa of an ionizable group.42 pH indicators Main article: pH indicator The transition range of a pH indicator is about pKa ± 1. This is the range over which the color is intermediate between the colors of the acidic and basic forms of the indicator. Universal indicator is a mixture of indicators whose adjacent pKa values differ by about two. pKa of some common substances There are multiple techniques to determine the pKa of a chemical causing some discrepancy between different sources. Well measured values are typically are within 0.1 units of each other. Data presented here was taken at 25 °C in water.2243 More values can be found in thermodynamics, above. Chemical Name Equilibrium pKa B = Adenine BH22+ \rightleftharpoons BH+ + H+ 4.17 BH+ \rightleftharpoons B + H+ 9.65 H3A = Arsenic acid H3A \rightleftharpoons H2A- + H+ 2.22 H2A- \rightleftharpoons HA2- + H+ 6.98 HA2- \rightleftharpoons A3- + H+ 11.53 HA = Benzoic acid HA \rightleftharpoons H+ + A- 4.204 HA = Butanoic acid HA \rightleftharpoons H+ + A- 4.82 H2A = Chromic acid H2A \rightleftharpoons HA- + H+ 0.98 HA- \rightleftharpoons A2- + H+ 6.5 B = Codeine BH+ \rightleftharpoons B + H+ 8.17 HA = Cresol HA \rightleftharpoons H+ + A- 10.29 HA = Formic acid HA \rightleftharpoons H+ + A- 3.751 HA = Hydrofluoric acid HA \rightleftharpoons H+ + A- 3.17 HA = Hydrocyanic acid HA \rightleftharpoons H+ + A- 9.21 HA = Hydrogen selenide HA \rightleftharpoons H+ + A- 3.89 HA = Hydrogen peroxide 90% HA \rightleftharpoons H+ + A- 11.7 HA = Lactic acid HA \rightleftharpoons H+ + A- 3.86 HA = Propanoic acid HA \rightleftharpoons H+ + A- 4.87 HA = Phenol HA \rightleftharpoons H+ + A- 9.99 H2A = L-+-Ascorbic Acid H2A \rightleftharpoons HA- + H+ 4.17 HA- \rightleftharpoons A2- + H+ 11.57 See also Determination of equilibrium constants Dissociation constant Henderson-Hasselbalch equation Hammett equation Isoelectric point Hydrolysis of metal salts QSAR References ^ a b Miessler, G. 1991. Inorganic Chemistry, 2nd ion, Prentice Hall, 165. ISBN 0134656598. ^ Burgess, J. 1978. Metal ions in solution. Ellis Horwood. ISBN 0853120277. Section 9.1, Acidity of solvated cations, lists many pKa values. ^ Headrick, Jeffrey M.; Eric G. Diken, Richard S. Walters, Nathan I. Hammer, Richard A. Christie, Jun Cui, Evgeniy M. Myshakin, Michael A. Duncan, Mark A. Johnson, Kenneth D. Jordan 2005. Spectral Signatures of Hydrated Proton Vibrations in Water Clusters. Science 308: 1765 - 1769. DOI: 10.1126/science.1113094 ^ Smiechowski, M.; Stangret J. 2006. Proton hydration in aqueous solution: Fourier transform infrared studies of HDO spectra. J. Chem. Phys.: 204508-204522. DOI:10.1063/1.2374891 ^ a b Rossotti, F.J.C.; Rossotti, H. 1961. The Determination of Stability Constants. McGraw-Hill. ^ Shriver, D.F; Atkins, P.W. 1999. Inorganic Chemistry, third ion, Oxford: Oxford Univerisy Press. ISBN 0198503318. Section 5.2 ^ Dasent, W.E. 1982. Inorganic energetics : an introduction. Cambridge University Press. ISBN 0521284066. ^ Brown, T.E.; Lemay, H.E.; Bursten, B.E. 2009. Chemistry The Central Science, 11th ion, Pearson Publications. ISBN 0131096869. p. 689 ^ a b Greenwood, N. N.; Earnshaw, A. 1997. Chemistry of the Elements, 2nd ion, Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4. p. 50 ^ Lide, D.R. 2004. CRC Handbook of Chemistry and Physics, Student ion, 84th. ed., CRC press. ISBN 0849305977. ^ Atkins, P.W.; de Paula, J. 2006. Physical chemistry. Oxford University Press. ISBN 0198700725. p 212 ^ a b Loudon, G.M. 2005. Organic Chemistry, 4th ion, New York: Oxford University Press. ISBN 0-19-511999-1. p. 317-318 ^ Housecroft, C.E.; Sharpe, A.G. 2008. Inorganic chemistry, 3rd. ed., Prentice Hall. ISBN 0131755536. Chapter 8 ^ March, J.; Smith, M. 2007. Advanced Organic Chemistry, 6th ion, New York: J. Wiley and Sons. ISBN 978-0-471-72091-1. ^ Kütt, Agnes; Valeria Movchun, Toomas Rodima, Timo Dansauer, Eduard B. Rusanov, Ivo Leito, Ivari Kaljurand, Juta Koppel, Viljar Pihl, Ivar Koppel, Gea Ovsjannikov, Lauri Toom, Masaaki Mishima, Maurice Medebielle, Enno Lork, Gerd-Volker Röschenthaler, Ilmar A. Koppel, and Alexander A. Kolomeitsev 2008. Pentakistrifluoromethylphenyl, a Sterically Crowded and Electron-withdrawing Group: Synthesis and Acidity of Pentakistrifluoromethylbenzene, -toluene, -phenol, and -aniline. J. Org. Chem. 73 7: 2607 -2620. doi:10.1021/jo702513w ^ Kütt, Agnes; Ivo Leito, Ivari Kaljurand, Lilli Sooväli, Vladislav M. Vlasov, Lev M. Yagupolskii, and Ilmar A. Koppel 2006. A Comprehensive Self-Consistent Spectrophotometric Acidity Scale of Neutral Brønsted Acids in Acetonitrile. J. Org. Chem. 71 7: 2829 -2838. doi:10.1021/jo060031y ^ Kaljurand, I.; Kütt, A.; Sooväli, L.; Rodima, T.; Mäemets, V. Leito, I; Koppel, I.A. 2005. Extension of the Self-Consistent Spectrophotometric Basicity Scale in Acetonitrile to a Full Span of 28 pKa Units: Unification of Different Basicity Scales. J. Org. Chem. 70 3: 1019 -1028. doi:10.1021/jo048252w ^ Bordwell pKa Table in DMSO ^ Coetzee, J. F. and Padmanabhan, G. R. 1965. Proton Acceptor Power and Homoconjugation of Mono- and Diamines. J. Amer. Chem. Soc. 87: 5005-5010. doi:10.1021/ja00950a006. ^ Box, K.J.; Völgyi, G. Ruiz, R. Comer, J.E. Takács-Novák, K., Bosch, E. Rà fols, C. Rosés, M. 2007. Physicochemical Properties of a New Multicomponent Cosolvent System for the pKa Determination of Poorly Soluble Pharmaceutical Compounds. Helv. Chim. Acta 90 8: 1538-1553. doi:10.1002/hlca.200790161. ^ Hammett, L.P. 1937. The Effect of Structure upon the Reactions of Organic Compounds. Benzene Derivatives. J. Am. Chem. Soc. 59 1: 96-103. doi:10.1021/ja01280a022. ^ a b Speight, J.G. 2005. Lange's handbook of chemistry, 18th. ed., McGraw-Hill. ISBN 0071432205. ^ a b R. Goldberg, N. Kishore, R. Lennen 2002. Thermodynamic Quantities for the Ionization Reactions of Buffers reprinted at NIST. J. Phys. Chem. Ref. Data 31: 231-370. doi:10.1063/1.1416902. ^ Martell, A.E.; Motekaitis, R.J. 1992. Determination and use of stability constants. Wiley. ISBN 0471188174. ^ Allen, R.I.; Box,K.J., Comer, J.E.A., Peake, C., Tam, K.Y 1998. Multiwavelength spectrophotometric determination of acid dissociation constants of ionizable drugs. J. Pharm. Biomed. Anal. 17 4-5: 699-641. doi:10.1016/S0731-70859800010-7. ^ Szakács, Zoltán; Hägele,Gerhard 2004. Accurate determination of low pK values by 1H NMR titration. Talanta 62: 819-825. doi:10.1016/j.talanta.2003.10.007 ^ Box, K.J.; Donkor, R.E. Jupp, P.A. Leader, I.P. Trew, D.F. Turner, C.H. 2008. The chemistry of multi-protic drugs Part 1: A potentiometric, multi-wavelength UV and NMR pH titrimetric study of the micro-speciation of SKI-606. J. Pharm. Biomed. Anal. 47 2: 303-311. doi:10.1016/j.jpba.2008.01.015. ^ Onufriev, Alexey; Case D.A; Ullmann G.M. 2001. A Novel View of pH Titration in Biomolecules. Bochemistry 40: 3413-3419. doi:10.1021/bi002740q. ^ Hulanicki, A. 1987. Reactions of acids and bases in analytical chemistry. Horwood. ISBN 0853123306. translation or: Mary R. Masson ^ N. E. Good, G. D. Winget, W. Winter, T. N. Connolly, S. Izawa and R. M. M. Singh 1966. Hydrogen Ion Buffers for Biological Research. Biochemistry 5 2: 467-477. doi:10.1021/bi00866a011. ^ Dunn, M.J. 1993. Gel Electrophoresis: Proteins. Bios Scientific Publishers. ISBN 187274821X. ^ Martin, R. 1996. Gel Electrophoresis: Nucleic Acids. Bios Scientific Publishers. ISBN 1872748287. ^ Brenner, B.M. or; Stein, J.H or 1979. Acid-base and Potassium Homeostasis. Churchill Livingstone. ISBN 0443080178. ^ Scorpio, R. 2000. Fundamentals of Acids, Bases, Buffers Their Application to Biochemical Systems. ISBN 0787273740. ^ Beynon, R.J.; Easterby, J.S. 1996. Buffer solutions : the basics. Oxford: Oxford University Press. ISBN 0199634424. ^ Perrin, D.D.; Boyd Dempsey. 1974. Buffers for pH and metal ion control. London: Chapman Hall. ISBN 0412117002. ^ Beck, M.T.; Nagypál, I. 1990. Chemistry of complex equilibria. Horwood. ISBN 0853121435. ^ Eyal, A.M 1997. Acid Extraction by Acid-Base-Coupled Extractants. Ion Exchange and Solvent Extraction: A Series of Advances, Volume 13: 31-94. ^ Stumm, W.; Morgan, J.J. 1996. Water chemistry. New York: Wiley. ISBN 0471051969. ^ Snoeyink, V.L.; Jenkins, D. 1980. Aquatic chemistry : chemical equilibria and rates in natural waters. New York: Wiley. ISBN 0471511854. ^ Millero, F.J. 2006. Chemical oceanography, 3rd. ion, London: Taylor and Francis. ISBN 0849322804. ^ Avdeef, A. 2003. Absorption and drug development : solubility, permeability, and charge state. New York: Wiley. ISBN 0471423653. ^ Washburn, E.W. 2003. International Critical Tables of Numerical Data, Physics, Chemistry and Technology, 1st. electronic ion, Knovel. http://knovel.com/web/portal/browse/display?_EXT_KNOVEL_DISPLAY_bookid=735VerticalID=0 Further reading Atkins, P.W.; Jones, L. 2008. Chemical Principles: The Quest for Insight, 4th. ion, W.H. Freeman. ISBN 1-4292-0965-8. Housecroft, C.E.; Sharpe, A.G. 2008. Inorganic chemistry, 3rd. ed., Prentice Hall. ISBN 0131755536. Non-aqueous solvents Hulanicki, A. 1987. Reactions of acids and bases in analytical chemistry. Horwood. ISBN 0853123306. translation or: Mary R. Masson Leggett, D.J. 1985. Computational methods for the determination of formation constants. Plenum. ISBN 0306419572. Perrin, D. D.; Dempsey, B. and Serjeant, E.P. 1981. pKa prediction for organic acids and bases. Chapman and Hall. ISBN 041222190x. Albert, A.; Serjeant, E.P. 1971. The determination of ionization constants : a laboratory manual. Chapman and Hall. ISBN 0412103001. Previous ion published as Ionization constants of acids and bases. London: Methuen, 1962 External links Acidity-Basicity Data pKa Values in Nonaqueous Solvents Extensive bibliography Shodor.org Acid-Base Chemistry Factors that Affect the Relative Strengths of Acids and Bases Purdue Chemistry Distribution diagrams of acids and bases generation from pKa values with free spreadsheet SPARC Physical/Chemical property calculator List of Aqueous-Equilibrium Constants Free guide to pKa logP interpretation and measurement v d e Articles related to solutions Solution Ideal solution Aqueous solution Solid solution Flory-Huggins Mixture Suspension chemistry Colloid Phase diagram Eutectic point Alloy Concentration Saturation chemistry Supersaturation Molar solution Percentage solution Serial dilution Solubility Solubility equilibrium Total dissolved solids Solvation Solvation shell Enthalpy change of solution Lattice energy Raoult's law Henry's law Solubility table data Solubility chart Solvent category Acid dissociation constant Protic solvent Inorganic nonaqueous solvent Solvation List of boiling and freezing information of solvents Partition coefficient Polarity Hydrophobe Hydrophile Lipophilic Amphiphile Retrieved from http://en..org/wiki/Acid_dissociation_constant Categories: Acids | Analytical chemistry | Thermodynamics Views Article Discussion this page History Personal tools Log in / create account Navigation Main page Contents Featured content Current events Random article Search Go Search Interaction Community portal Recent changes Contact Donate to Help Toolbox What links here Related changes Upload file Special pages Printable version Permanent link Cite this page Languages Dansk Deutsch Español Français Bahasa Indonesia Italiano Lietuvių Nederlands 日本語 ‪Norsk bokmÃ¥l‬ Português РуÑ?Ñ?кий Suomi Svenska УкраїнÑ?ька ä¸æ–‡ This page was last modified on 7 September 2008, at 21:50
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