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From Wikipedia, the free encyclopedia

An acid dissociation constant (aka acidity constant, acid-ionization constant) is an for the dissociation of an acid. It is denoted by Ka. For an equilibrium between a generic acid, HA, and − its conjugate , A , The weak acid donates a to in an equilibrium reaction to give the and − + HA A + H the ion. Key: is white, is red, is gray. Lines are chemical bonds. K is defined, subject to certain conditions, as a

where [HA], [A−] and [H+] are equilibrium 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 , particularly at introductory levels, the Brønsted–Lowry acid-base theory is versatile enough that acidic behavior can now be characterized even in non-aqueous .

The value of pK indicates the strength of an acid: the larger the value the weaker the acid. In aqueous a , simple 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 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 , chemical oceanography, buffer solutions, acid-base homeostasis and certain kinds of 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 to form complexes in solution.

Acids and bases:

Contents Acid dissociation constant Acid-base extraction Acid-base reaction 1 Definitions Acid-base 2 Equilibrium Constant Acid-base physiology Acid-base homeostasis 2.1 Monoprotic acids 2.2 Polyprotic acids Dissociation constant 2.3 Water self-ionization Non-nucleophilic base 2.4 Bases pH 2.5 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 Lewis acid/base /base 4 Factors that determine the relative strengths of acids /base 4.1 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 , releasing the hydrogen ion. Chemical Equilibria Acid dissociation constant − + HA A + H The equilibrium constant for this "dissociation" reaction is known as a Buffer solution dissociation constant. However, since the liberated proton combines with a water to give an hydronium ion, Arrhenius proposed that the "dissociation" reaction should be written as an acid-base Dissociation constant reaction. Distribution coefficient Distribution ratio − + HA + H2O A + H3O Equilibrium constant Brønsted and Lowry generalized this definition as a proton exchange Equilibrium stage [1] reaction, as follows. -liquid extraction diagram acid + base conjugate base + 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 Relative the base gains a proton. For aqueous solutions an acid, HA, reacts with 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 than water, such as DMSO; in that case the solvent, S, acts as a base, -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 . 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 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 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 equilibrium can also be classed as a weak acid since, in hydrolysis, protons are produced by the splitting of water . For example, the equilibrium

- + B(OH)3 + 2 H2O B(OH)4 + H3O

shows why 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 of concentration and the definition could also be written as

Variation of pKa of acetic acid with

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 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 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). 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]

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An example of a strong acid is , HCl, which has has a pK value, estimated from a 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 pK and analytical concentration of the acid are known. See table for details. a

Polyprotic acids

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 of either H PO − or HPO 2− may be crystallized from solution by adjustment 2 4 4 ofpH toeither 4 or10. % species' formation as a function of pH 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, H A, the two equilibria are 2

− + H2A HA + H HA − A2− + H+

it can be seen that the second proton is removed from a negatively % species formation calculated charged species. Since the proton carries a positive charge extra with the program HySS (http:// work is needed to remove it; that is the cause of the trend noted www.hyperquad.co.uk/hyss.htm) for a 10mM solution of citric acid. above. Phosphoric acid, H3PO 4, (values below), illustrates this rule, pKa1=3.13, pKa2 = 4.76, as does vanadic acid. When an exception to the rule is found it pKa3=6.40. indicates that a major change in structure is ocurring. In the case of + VO 2 (aq), the vanadium is octahedral, 6-coordinate, whereas all the other species are tetrahedral, 4-coordinate. This explans why pKa1 > pKa2 for vanadium(V) oxoacids.

+ + pK = 4.2 VO 2 H3VO 4 + H a1 H PO H PO − + H+ pK = 2.15 H VO H VO − + H+ pK = 2.60 3 4 2 4 a1 3 4 2 4 a2 − 2− + pK = 7.20 − 2− + pK = 7.92 H2PO 4 HPO 4 + H a2 H2VO 4 HVO 4 + H a3 HPO 2− PO 3− + H+ pK = 12.37 HVO 2− VO 3− + H+ pK = 13.27 4 4 a3 4 4 a4 http://en.wikipedia.org/wiki/Acid_dissociation_constant 2008-09-24 18:35 Acid dissociation constant - Wikipedia, the free encyclopedia Page 6

Water self-ionization

Water has both acidic and basic properies. The equilibrium constant for the equilibrium

H O + H O OH − + H O+ 2 2 3

is given by

Since the concentration of water can be assumed to be constant, this expression simplifies to

The self-ionization constant of water, K , can thus be seen as a special case of an acid dissociation w constant.

Bases

Historically the equilibrium constant Kb for a base was defined as the association constant for of the base, B, to form the conjugate acid, HB +.

B + H O HB + + OH − 2

Using similar reasoning to that used before

In water, the concentration of the ion, [OH −], is related to the concentration of the hydrogen ion by K = [H+][OH –], therefore w

Substitution of the expression for [OH −] into the expression for K gives b

It follows, taking cologarithms, that pK = pK – pK . In aqueous solutions at 25 °C, pK is 13.9965,[10] b w a w so pK ~ 14 – pK . b a

In effect there is no need to define pK separately from pK , but it is done here because pK values can b a b be found in the older literature.

Temperature dependence

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All equilibrium constants vary with temperature according the van 't Hoff equation [11 ]

Thus, for exothermic reactions, (HO is negative) K decreases with temperature, but for endothermic reactions (HO 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]

1. it is a , capable of forming hydrogen bonds 2. it has a high , making it a strong Lewis base. 3. it has a high constant (relative ), 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 Non-polar 2.2 Non-polar 2.3 Diethylether Non-polar,Donor 4.3 Aceticacid Proticdonor 6.1 Donor 7.6 Polardonor 21 Liquidammonia Polardonor 25at195K Polardonor 37 Dimethylsulfoxide Polardonor 47 Water Polarproticdonor78 Polarproticdonor 111 Sulphuricacid Polarprotic 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 + CH CO H Cl − + CH C(OH) + 3 2 3 2 acid + base conjugate base + conjugate acid

Compare this reaction with what happens when acetic acid is dissolved in the more acidic solvent pure sulphuric acid [13]

H SO + CH CO H HSO − + CH C(OH) + 2 4 3 2 4 3 2

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The apparently unlikely geminal diol species CH C(OH) + is stable in these environments. 3 2

pK values of organic compounds are often obtained using solvents other than water, such as dimethyl a 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.

[15][16][17] Below is a table of selected pKa values at 25 °C in acetonitrile (AN) and (DMSO).[18] Values for water are included for comparison.

BH + B + H+ AN DMSO water HA A− + H+ AN DMSO water Pyrrolidine 19.5610.8 11.4 p-Toluenesulfonicacid 8.5 0.9 strong Triethylamine 18.82 9.0 10.72 2,4-Dinitrophenol 16.66 5.1 3.9 Proton sponge 18.62 7.5 12.1 Benzoicacid 21.5111.1 4.2 12.533.4 5.2 Aceticacid 23.5112.6 4.756 10.623.6 9.4 Phenol 29.1418.0 9.99

In solvents of low dielectric constant ions tend to associate forming ion pairs and clusters, which complicates the interpretation of pKa values.

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 dimerization of a carboxylic enhanced by a factor of nearly 800.[19] acid

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 pK values in a solvent such as water/dioxane or water/, in a [20] which the compound is more soluble. 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 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 pK value for use with aqueous solutions it has to be a extrapolated to zero co-solvent concentration from values obtained from various co-solvent .

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. http://en.wikipedia.org/wiki/Acid_dissociation_constant 2008-09-24 18:35 Acid dissociation constant - Wikipedia, the free encyclopedia Page 9

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 rule [9] states that the value of the first pK for acids of the formula XO (OH) is a m n approximately independent of n and X and is approximately 8 form=0,2form=1,−3form=2and< −10 for m = 3. This correlates with the oxidation state of the central , X: the higher the oxidation state the stronger the oxyacid. For example, pKa for HClO is 7.2, for HClO 2 is 2.0, for HClO 3 is −1 and HClO 4 is a strong acid.

With organic acids inductive effects and mesomeric effects affect the pK'a values. The effects are summarised in the and subsequent extensions.[21]

Structural effects can also be important. The difference between fumaric acid and fumaric acid 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 pK values of a approximately 3.5 and 4.5. By contrast, maleic acid has pK values of approximately a [22] 1.5 and 6.5. The reason for this large difference is that when one proton is maleic acid removed from the cis- isomer (maleic acid) a strong intramolecular 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, 1,8-Bis(dimethylamino)naphthalene, has a pKa value of 12.1. It is one of the strongest 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 change for the proton sponge reaction, so for an acid dissociation constant

GO = 2.303 RT pK . a

Note that pK = –log K . At 25 °C GO /kJ mol -1 = 5.708 pK . Free energy is made up of an a a a term and an term.[23]

GO = HO – TSO

The standard enthalpy change can be determined by 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 O [23] pKa and H . The data were critically selected and refer to 25 °C and zero ionic strength, in water.

Acids K 0 −1 0 −1 Compound Equilibrium p a ∆H /kJ mol –T∆S /kJ mol http://en.wikipedia.org/wiki/Acid_dissociation_constant 2008-09-24 18:35 Acid dissociation constant - Wikipedia, the free encyclopedia Page 10

HA = Acetic acid HA H+ + A− 4.756−0.41 27.56 H A+ = GlycineH+ H A+ HA + H+ 2 2 2.3514.00 9.419 HA H+ + A− 9.78 44.20 11.6 H A = Maleic acid − + 2 H2A HA + H 1.92 1.10 9.85 HA − H+ + A2− 6.27 −3.60 39.4 H A = Citric acid − + 3 H3A H2A + H 3.1284.07 13.78 H A− HA 2− + H+ 2 4.76 2.23 24.9 HA 2− A3− + H+ 6.40 −3.38 39.9 HA = Boric acid HA H+ + A− 9.23713.80 38.92 H A = Phosphoric acid H A H A− + H+ 3 3 2 2.148−8.00 20.26 − 2− + H2A HA + H 7.20 3.60 37.5 HA 2− A3− + H+ 12.3516.00 54.49 HA − = Hydrogen sulphate HA − A2− + H+ 1.99 −22.40 33.74 H A = Oxalic acid − + 2 H2A HA + H 1.27 −3.90 11.15 HA − A2− + H+ 4.2667.00 31.35

Conjugate acid of bases K 0 −1 0 −1 Compound Equilibrium p a ∆H /kJ mol –T∆S /kJ mol B = HB + B + H+ 9.245 51.95 0.8205 B = Methylamine HB + B + H+ 10.645 55.34 5.422 B = Triethylamine HB + B + H+ 10.72 43.13 18.06

The first point to note is that when pK is positive, the standard free energy change for the dissociation a 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 HO is negative –TSO is the dominant factor which determines that GO 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

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 http://en.wikipedia.org/wiki/Acid_dissociation_constant 2008-09-24 18:35 Acid dissociation constant - Wikipedia, the free encyclopedia Page 11

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 A calculated titration curve of ensure that pKa remains invariant during the titration. oxalic acid titrated with a solution of A calculated titration curve for oxalic acid is shown at the right. Oxalic acid has pK values of 1.27 and 4.27. Therefore the buffer regions will be centered at about pH 1.3 and a pH 4.3. The buffer regions carry the information necessary to get the pK values as the concentrations of a 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 electrode, because the breaks down at such low pH values. To determine pK values of less than about 2 or more than about 11 spectrophotometric [25] or NMR [26] 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 the pK values of proteins and amino acid side chains are of major importance for a the activity of and the stability of proteins.[28] The reaction that converts adenosine triphosphate to adenosine diphosphate is very pH sensistive.

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 of strong acid or strong base. The buffer capacity [29] of a simple buffer solution (illustrative diagram (http://www.chembuddy.com/?left=pH- calculation&right=pH-buffer-capacity) ) 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 http://en.wikipedia.org/wiki/Acid_dissociation_constant 2008-09-24 18:35 Acid dissociation constant - Wikipedia, the free encyclopedia Page 12

solution with pH 7.2; others are listed in buffer solutions and Good's buffers. Buffers such as tricine are used in Gel .[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 homeostasis [33] and disorders such as Acid-base imbalance.[34][35][36]

Coordination compounds

A 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

m+ (m−1)+ + [M(H2O)n] +LH [M(H2O)n−1L] + 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 EDTA 4− can accept four protons; in that case, all pK values must be known. In addition, the a metal ion is subject to hydrolysis, that is, it behaves as a weak acid, sothe pK values for the hydrolysis reactions must also be known.[37]

Solvent 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

Acid-base equilibria are important for rivers and lakes,[39][40] and in chemical oceanography.[41]

Pharmacology

Ionization of a compound alters its physical behavior and macro properties such as solubility and (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 [42] of a compound in the blood by adjusting the pKa of an ionizable group.

pH indicators

The transition range of a pH indicator is about pK ± 1. This is the range over which the color is a intermediate between the colors of the acidic and basic forms of the indicator. Universal indicator is a mixture of indicators whose adjacent pK values differ by about two. a

pKa of some common substances

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There are multiple techniques to determine the pK of a chemical causing some discrepancy between a 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.[22][43] More values can be found in thermodynamics, above.

K Chemical Name Equilibrium p a BH 2+ BH + + H+ B = Adenine 2 4.17 BH + B + H+ 9.65 H A = Arsenic acid − + 3 H3A H2A + H 2.22 H A− HA 2− + H+ 2 6.98 HA 2− A3− + H+ 11.53 HA = Benzoic acid HA H+ + A− 4.204 HA = Butanoic acid HA H+ + A− 4.82 H A = Chromic acid − + 2 H2A HA + H 0.98 HA − A2− + H+ 6.5 B = Codeine BH + B + H+ 8.17 HA = Cresol HA H+ + A− 10.29 HA = HA H+ + A− 3.751 HA = HA H+ + A− 3.17 HA = Hydrocyanic acid HA H+ + A− 9.21 HA = Hydrogen selenide HA H+ + A− 3.89 HA = (90%) HA H+ + A− 11.7 HA = Lactic acid HA H+ + A− 3.86 HA = Propanoic acid HA H+ + A− 4.87 HA = Phenol HA H+ + A− 9.99 H A = L-(+)-Ascorbic Acid − + 2 H2A HA + H 4.17 HA − 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 http://en.wikipedia.org/wiki/Acid_dissociation_constant 2008-09-24 18:35 Acid dissociation constant - Wikipedia, the free encyclopedia Page 14

1. ^ a b Miessler, G. (1991). Inorganic Chemistry , 2nd edition, Prentice Hall, 165. ISBN 0134656598. 2. ^ Burgess, J. (1978). Metal ions in solution . Ellis Horwood. ISBN 0853120277. Section 9.1, "Acidity of solvated cations", lists many pKa values. 3. ^ 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 4. ^ 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 5. ^ a b Rossotti, F.J.C.; Rossotti, H. (1961). The Determination of Stability Constants . McGraw-Hill. 6. ^ Shriver, D.F; Atkins, P.W. (1999). Inorganic Chemistry , third edition, Oxford: Oxford Univerisy Press. ISBN 0198503318. Section 5.2 7. ^ Dasent, W.E. (1982). Inorganic energetics : an introduction . Cambridge University Press. ISBN 0521284066. 8. ^ Brown, T.E.; Lemay, H.E.; Bursten, B.E. (2009). Chemistry The Central Science , 11th Edition, Pearson Publications. ISBN 0131096869. p. 689 9. ^ a b Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements , 2nd Edition, Oxford:Butterworth- Heinemann. ISBN 0-7506-3365-4. p. 50 10. ^ Lide, D.R. (2004). CRC Handbook of Chemistry and Physics, Student Edition , 84th. ed., CRC press. ISBN 0849305977. 11. ^ Atkins, P.W.; de Paula, J. (2006). . Oxford University Press. ISBN 0198700725. p 212 12. ^ a b Loudon, G.M. (2005). Organic Chemistry , 4th Edition, New York: Oxford University Press. ISBN 0-19-511999-1. p. 317–318 13. ^ Housecroft, C.E.; Sharpe, A.G. (2008). Inorganic chemistry , 3rd. ed., Prentice Hall. ISBN 0131755536. Chapter 8 14. ^ March, J.; Smith, M. (2007). Advanced Organic Chemistry , 6th edition, New York: J. Wiley and Sons. ISBN 978-0-471-72091-1. 15. ^ 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). "Pentakis(trifluoromethyl)phenyl, a Sterically Crowded and -withdrawing Group: Synthesis and Acidity of Pentakis(trifluoromethyl)benzene, -, -phenol, and -aniline". J. Org. Chem. 73 (7): 2607 -2620.doi:10.1021/jo702513w 16. ^ 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 17. ^ 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 18. ^ Bordwell pKa Table in DMSO (http://www.chem.wisc.edu/areas/reich/pkatable/) 19. ^ 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. 20. ^ Box, K.J.; Völgyi, G. Ruiz, R. Comer, J.E. Takács–Novák, K., Bosch, E. Ràfols, 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. 21. ^ 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. 22. ^ a b Speight, J.G. (2005). Lange's handbook of chemistry , 18th. ed., McGraw-Hill. ISBN 0071432205. 23. ^ 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. 24. ^ Martell, A.E.; Motekaitis, R.J. (1992). Determination and use of stability constants . Wiley. ISBN 0471188174. 25. ^ 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/S0731-7085(98)00010-7. 26. ^ 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

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27. ^ 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. 28. ^ Onufriev, Alexey; Case D.A; Ullmann G.M. (2001). "A Novel View of pH Titration in ". Bochemistry 40 : 3413–3419. doi:10.1021/bi002740q. 29. ^ Hulanicki, A. (1987). Reactions of acids and bases in analytical chemistry . Horwood. ISBN 0853123306. (translation editor: Mary R. Masson) 30. ^ 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. 31. ^ Dunn, M.J. (1993). Gel Electrophoresis: Proteins . Bios Scientific Publishers. ISBN 187274821X. 32. ^ Martin, R. (1996). Gel Electrophoresis: Nucleic Acids . Bios Scientific Publishers. ISBN 1872748287. 33. ^ Brenner, B.M. (Editor); Stein, J.H (Editor) (1979). Acid-base and Potassium Homeostasis . Churchill Livingstone. ISBN 0443080178. 34. ^ Scorpio, R. (2000). Fundamentals of Acids, Bases, Buffers & Their Application to Biochemical Systems . ISBN 0787273740. 35. ^ Beynon, R.J.; Easterby, J.S. (1996). Buffer solutions : the basics . Oxford: Oxford University Press. ISBN 0199634424. 36. ^ Perrin, D.D.; Boyd Dempsey. (1974). Buffers for pH and metal ion control . London: Chapman & Hall. ISBN 0412117002. 37. ^ Beck, M.T.; Nagypál, I. (1990). Chemistry of complex equilibria . Horwood. ISBN 0853121435. 38. ^ Eyal, A.M (1997). "Acid Extraction by Acid-Base-Coupled Extractants". and Solvent Extraction: A Series of Advances, Volume 13 : 31–94. 39. ^ Stumm, W.; Morgan, J.J. (1996). Water chemistry . New York: Wiley. ISBN 0471051969. 40. ^ Snoeyink, V.L.; Jenkins, D. (1980). Aquatic chemistry : chemical equilibria and rates in natural waters . New York: Wiley. ISBN 0471511854. 41. ^ Millero, F.J. (2006). Chemical oceanography , 3rd. edition, London: Taylor and Francis. ISBN 0849322804. 42. ^ Avdeef, A. (2003). Absorption and drug development : solubility, permeability, and charge state . New York: Wiley. ISBN 0471423653. 43. ^ Washburn, E.W. (2003). International Critical Tables of Numerical Data, Physics, Chemistry and Technology , 1st. electronic edition, Knovel.http://knovel.com/web/portal/browse/display?_EXT_KNOVEL_ DISPLAY_bookid=735&VerticalID=0 Further reading

Atkins, P.W.; Jones, L. (2008). Chemical Principles: The Quest for Insight , 4th. edition, 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 editor: 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 edition published as Ionization constants of acids and bases . London: Methuen, 1962)

External links

Acidity-Basicity Data (pKa Values) in Nonaqueous Solvents (http://tera.chem.ut.ee/~ivo/HA_UT/) Extensive bibliography Shodor.org Acid-Base Chemistry (http://www.shodor.org/unchem/basic/ab/) Factors that Affect the Relative Strengths of Acids and Bases (http://library.thinkquest.org/ C006669/data/Chem/acidsbases/factors.html) http://en.wikipedia.org/wiki/Acid_dissociation_constant 2008-09-24 18:35 Acid dissociation constant - Wikipedia, the free encyclopedia Page 16

Purdue Chemistry (http://chemed.chem.purdue.edu/genchem/topicreview/bp/3organic/3org_ frame.html) Distribution diagrams of acids and bases (http://www2.iq.usp.br/docente/gutz/Curtipot_.html) (generation from pKa values with free spreadsheet) SPARC Physical/Chemical property calculator (http://sparc.chem.uga.edu) List of Aqueous-Equilibrium Constants (http://www.jesuitnola.org/upload/clark/Refs/aqueous.htm) Free guide to pKa & logP interpretation and measurement (http://www.raell.demon.co.uk/chem/ logp/logppka.htm)

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