Understanding Acids: Protonic & Lewis Acids, Conjugates, pKa Scale, Study Guides, Projects, Research of Chemistry

Download Understanding Acids: Protonic & Lewis Acids, Conjugates, pKa Scale and more Study Guides, Projects, Research Chemistry in PDF only on Docsity! 43 Acids and Bases Reference: P. Bruice, Organic Chemistry, 6th Edition, Chapters 1.16-1.26, 7.9, 16.5. Definitions Bases (general definition) - All substances that contain unshared electron pairs are bases. Examples: NH2 H O .... – H N H H CH3 C O O .. .... – Cl .. .. – CH3 O .... – OH H .... .. Acids 1. Proton or protonic acids - Proton acids (Brønsted acids) are substances that can transfer a proton to a base. They are proton donors. They are usually substances that have a hydrogen atom bonded to an electronegative atom. Examples: CH3 C O O H .. .. .. ....OH H HCl NH3R + 2. Lewis Acids - According to the concept of G. N. Lewis, acids are not limited to proton donors, but an acid is any substance that contains an element having a vacant orbital that can accept a pair of electrons in forming a bond. According to the Lewis concept it is the bare proton with its vacant s orbital that is the acidic entity in protonic acids. But bare unsolvated protons do not exist in solution and it is now customary to differentiate between proton acids and Lewis acids. The term 'Lewis acid' is used to designate substances having a vacant orbital, usually substances containing an element that is two electrons short of having a complete valence shell. F3B– O Et Et + O Et Et F3B + Lewis Acid Base Addition Compound B NF F F H H H – + B F F F N H H H + Lewis Acid Base Addition Compound Note that the resulting addition compounds formed from electrically neutral molecules have a formal negative charge on the boron (the electron acceptor) and a formal plus charge on the electron donor. 44 Proton Acids and Bases (Conjugate Acid-Base Pairs) - The ionization of a proton acid involves the transfer of a proton from the acid to a base, or more correctly, the removal of a proton from the acid by a base. Strong acids are able to transfer their proton to weak bases, but weak acids may require a very strong base to bring about the proton transfer of ionization. When an acid ionizes in water, the water molecules, with their unshared pairs of electrons, serve as the base. It is important to realize that bare unsolvated protons do not exist in solution. Ionization of a proton acid always requires a base to remove or accept the proton. Often the solvent molecules serve as the base. The process can be shown by following general equation, where B represents a base and H–A represents a proton acid. H A+ B H + A–B + Note that the anion (:A_) that results from the removal of the proton from the acid H–A, is itself a base. It is called the conjugate base of acid H–A. The protonated species B+–H is now an acid, and it is called the conjugate acid of base B:. The acid-base reaction is therefore a competition reaction between two bases for a single proton. The reaction is an equilibrium process. The position of the equilibrium is affected by the relative basicities of the competing bases B: and :A_. 47 The pKa - a single scale for reporting and comparing the ionization capacity of acids and bases. Since every base has its conjugate acid, it is possible to compare all acids and bases on a single scale. The scale commonly used is the acidity constant of the conjugate acid of any conjugate pair, expressed in logarithmic units. This is called the pKa. By definition, the pKa is the negative logarithm of Ka value. pKa = –log Ka For acetic acid, Ka = 1.75 x 10-5 and pKa = –log (1.75 x 10-5) = 5 –log 1.75 = 4.76. It is very important to recognize that the pKa of bases, such as ammonia or the organic amines, is a measure of the acidity of the conjugate acid of the base, a measure of the acidity of the salts of ammonia or amines. The pKa of methylamine [CH3NH3] is given as 10.6. This means that a salt of methylamine, such as CH3NH3Cl–, has an acidity constant Ka = 10-10.6, or 2.51 x 10-11. H2O CH3NH2 H3O+ Cl–CH3NH3 Cl– + +++ = 2.51 x 10-11= [H3O+][CH3NH2] Ka [CH3NH3] + pKa = –log (2.51 x 10-11) = 10.6 The ionization of water is Kw. 2 H2O H3O+ + OH– Kw = [H3O+] [OH–] = 10-14 pKw = –log Kw = 14 Here again, the constant concentration term of the water is incorporated in the constant Kw. In pure water the concentration of hydronium and hydroxide ions must of course be equal. At 25° C, the concentration of each is 10-7 mole/liter. + + [H3O+] [OH–] [H2O] Ka = pKa = 15.7 = = 10–15.7 55.5 Kw 48 pH - By definition, the pH of a water solution is the negative log of the hydronium ion concentration in moles/liter. pH = –log [H3O+] For pure water, the pH is 7. The water is said to be 'neutral', neither acidic nor basic. If the concentration of hydronium ion is increased above [10-7], by the addition of some acid that is more acidic than water the solution is said to be acidic. The addition of a base increases the concentration of hydroxide ions and correspondingly decreases the concentration of the hydronium ion and the solution is said to be basic. The product of the concentration of H3O+ and OH– is the Kw and retains the constant value 10-14. Water solutions with pH values less than 7 are acidic and those having pH values larger than 7 are basic. More on pKa scale - The experimental measurement of pKa values in water are, of course limited to acids that are stronger acids than water, acids that will ionize to some degree in water (where the conjugate base of the acid is a weaker base than the hydroxyl anion), but at the same time are not acidic enough to be completely ionized in water, acids with pKa values between 0 and 14. The relationship Ka = [H3O+] [A H] [A –] requires that there be measurable concentrations of both the conjugate acid and conjugate base at equilibrium in order to determine Ka. For very strong acids (pKa < 0) and very weak acids (pKa > 14), indirect nonaqueous competition experiments using acids and bases of known pKa's must be used in order to compare them on the same scale as those measurable in water. The pKa values at the two extremes of the scale (beyond 0 and 14 each way) are less accurate than those between 0 and 14. 49 Examples of Acids and Acidity Scale Conjugate Acid pKa Conjugate Base Strong Acids 1. H–I ~ –10 :I:– 2. H–Cl ~ –7 :Cl:– 3. H2SO4 ~ –3 HSO4– The above acids cannot exist as such in aqueous solution. In each case the conjugate base is a weaker base than the water molecule. The proton would be completely transferred to the water molecules. The standard "strong" acids are completely ionized in water. These are HClO4, HI, HCl and H2SO4 (first proton). 4. F3C C O OH ~ 0 F3C C O O– 5. Cl3C C O OH ~ 0.8 Cl3C C O O– 6. NH3O2N + 1.0 NH2O2N .. 7. Cl2CH C O OH 1.3 Cl2CH C O O– 8. ClCH2 C O OH 2.8 ClCH2 C O O– 9. OH C OH O 3.0 OH C O– O 10. O2N C O OH 3.4 O2N C O O– 11. Cl C OH O 4.0 Cl C O– O 12. Cl NH3 + 4.0 Cl NH2 .. .. .. .. .. 52 For acids of intermediate strengths (pKa between 0 and 14) which are soluble in water, the degree of water protonation will depend on the pKa of the acid. For a 0.1 N solution of acetic acid most of the acid is in the nonionized form. Taking the pKa as 5 would give: = [AcOH] [H3O+][AcO–] Ka 10-5 = x2 0.1 - x AcOH H2O AcO– H3O++ + 0.1 molar x x Since x is small (compared to 0.1M), it can be neglected in the denominator and x2 = 10-6, x = 10-3. The pH of the solution is about 3 and the ratio of ionized to nonionized acetic acid is about 1/100. For dilute aqueous solutions of weak acids the approximate pH can be calculated as the negative log of the square root of the product of the Ka and the molar concentration. − log Ka [molar concentration of acid]pH =~ Ka = x2 [AcOH] ; x2 = Ka [AcOH] x; = Ka[AcOH] 53 Drugs which are amines (weak bases) are sometimes used as their hydrochloride salts, procaine hydrochloride for example. Salts of amines are the conjugate acids of amines and are of course acidic. The pH of a 0.1 N solution of procaine hydrochloride will be about 5. + H2OH2N C O O CH2 CH2 N Et H Et + H2N C O O CH2 CH2 N Et Et .. + H3O+ [0.1 molar] pKa = 8.9 ~ 9 Ka = , Ka [BH+] = x2 = (10-9) (10-1) = 10-10, x = 10-5, pH = 5. It is important to remember that pKa values are logarithmic values. A difference of one pKa unit means that the ionization constant varies by a factor of ten. Compare the acidity of the salts of aromatic and aliphatic amines. pH of a 0.1 Molar solution, Ka [BH+] = x2 = (10-4.6) (10-1) = 10-5.6, x = 10-2.8, pH = 2.8. pH of a 0.1 Molar solution, Ka [BH+] = x2 = (10-10.6) (10-1) = 10-11.6, x = 10-5.8, pH = 5.8. The aniline hydrochloride is 106 (one million times) more acidic than methylamine hydrochloride. The aliphatic methylamine is one million times more basic than the aromatic amine, aniline. [B:] [H3O+] [BH+] NH3 + Cl– , pKa = 10.6 Cl– , pKa = 4.6 CH3 NH3 + 54 Degree of Ionization of Acids and Bases at Controlled pH in Buffered Solutions. The Henderson-Hasselbalch equation. It is obvious that when an acid (more acidic than water) is added to water it will ionize by transferring its protons to water and thus increase the hydronium ion concentration and lower the pH of the water solution. A H H2O H3O+A – ++ The pH of the resulting solution will depend on the concentration and the pKa of the acid. In the pharmaceutical sciences, a very important question regarding acids and bases is the degree of ionization of acids and bases of known pKa values at a given, constant pH, in buffered systems such as the blood and plasma. Many important drugs are weak acids or weak bases. A knowledge of the degree of ionization of such drugs at physiological pH is very important to problems of absorption, distribution in body tissues, and excretion. Such a knowledge is also basic to problems of solubilization of drugs in the preparation of suitable dosage forms. The Henderson-Hasselbalch equation allows the calculation of the ratio of conjugate acid and conjugate base for a conjugate acid-conjugate base pair of known pKa at any given pH. pH = pKa + log The equation is readily derived: memory not required for its application. Conjugate acid + H2O H3O+ + Conjugate base Ka = Taking the log of each side of the equation gives: log Ka = log [H3O+] + log Rearranging, – log [H3O+] = – log Ka + log [Conjugate base] [Conjugate acid] [H3O+] [Conjugate base] [Conjugate acid] [Conjugate base] [Conjugate acid] [Conjugate base] [Conjugate acid] 57 Drugs as Acids and Bases A. Guanidines and Amidines Basicity of Guanidine Guanidine, pKa ~ 14, is almost completely ionized in water. Some substituted guanidines are useful drugs. BH+ + H2O B: + H3O+ .. C NH H2N H2N C NH H2N H2N.. .. ..C NH H2N H2N .. .. + H3O+ .. + .. C NH2 H2N H2N C NH2 H2N H2N.. .. + + .. .. C NH2 H2N H2N + H2O Three identical resonance structures and no separation of charges in the guanidinium cation. Resonance stabilizes the protonated species. The plus charge is equally distributed on the three nitrogen atoms. The reason for the very strong basicity of guanidine is the great stabilization of the cation resulting from the dispersal of the positive charge, which is equally distributed between the three nitrogen atoms, and the gain in resonance upon protonation. Several substituted amidines and guanidines are useful drugs. These agents have effects in the periphery, and appear not to be transported in significant amounts to the central nervous system, probably due to their polarity. Some are also not absorbed well for the same reason. C N H N NH C NH NH2 CH3 CH3 Metformin pKa = 12.4 N NH2 NH C H NCH2CH2 Guanethidine pKas = 11.9, 8.3 O O N H NH2 NH Guanoxan pKa = 12.3 pKa ~ 14 Ka = 10-14 58 Clonidine pKa = 8.4 Naphazoline pKa = 10.9 Guanidines substituted with electron withdrawing groups have greatly reduced basicity, e.g. cimetidine. N NH CH2S CH2CH2 N H C N NHCH3 CN H3C neutral C NH(insulin) O CH C CH2 NH O CH COO CH2 CH2 CH2 H N C NH2 NH CH2 CH2 H N C NH NH2 MW5800 glargine insulin Soluble HCl salt administered subcutaneously. It precipitates and the slow rate of subsequent dissolution and then absorption into the blood stream provides a long duration of effect. (Two arginines added to the carboxyl end of the B-chain of insulin.) N H Cl Cl N N H N H NCH2 59 B. Amines - Tertiary, Secondary, and Primary Many drugs are tertiary or secondary aliphatic amines, a few are primary amines. Tertiary amines usually are metabolized by fewer independent pathways than those with fewer alkyl groups on nitrogen, i.e. secondary and primary aliphatic amines may be metabolized by different processes. Most amine drugs are basic (pKa ~ 9). Even though they are highly ionized at physiological pH, most are readily absorbed after oral administration (rapid equilibrium between ionized and non-ionized forms to penetrate membranes), and many lipophilic ones reach the CNS. Examples CH OCH2 CH2 N CH3 CH3 Diphenhydramine pKa = 9.1 F3C O CH CH2CH2NHCH3 Fluoxetine pKa = 8.7 O NCH2CH2C F NC H3C H3C H2 Citalopram pKa = 9.5 OCH2 CH CH2NHCH(CH3)2 OH Propranolol pKa = 9.5 62 C. Phenols, Enols, Imides, Sulfonamides and Related Compounds Increased acidity (dissociation of acid) is associated with stabilization of resulting structures by resonance (anions especially), to a greater extent than in the non-dissociated acid. Resonance rules: 1. Only electrons move. The nuclei of the atoms never move. 2. Only π electrons or non-bonding electrons contribute. 3. Electrons move toward a positive charge or toward a π bond. 4. Total number of electrons does not change. The number of paired and unpaired electrons does not change. Examples HO HO CH2CH2NH2 Dopamine pKa = 10.6 (phenol) 8.9 (RNH3) O H3C H H H HO Estrone pKa = 10.8 O OH CH3 H H3C CH3 H C5H11 Δ9-THC pKa = 10.6 + cyclohexanol pKa ~ 15-16 phenolate anion phenol pKa ~ 10 OHO + H3O+ .... ..OH + H2O .... O .... O .... O ....O .... .. 63 N S C H N N HO O O O CH3 Piroxicam pKa = 4.6 O OH O CH2 CH Ph C O CH3 Warfarin pKa = 5.1 S H2N NH2 O O Sulfanilamide pKa = 10.6 Cl S N S NH H2N O O O O Chlorthiazide pKa = 6.7, 9.5 64 Imides and Related Compounds The imides, where the nitrogen is flanked by two carbonyl groups, are much stronger acids (pKa ~ 9) than amides, which have no detectable basicity or acidity in water. Salts of imides can be formed in water. The acidity of imides results from a greater resonance stabilization of the conjugate base than the conjugate acid and stabilization of the conjugate base by a dispersal of the negative change between the two oxygens and the nitrogen. The oxygen, being more electronegative than nitrogen, is a better accommodator of a negative charge than nitrogen. The structure of the anion is better written with the negative charge on the oxygen. C N C O O H + H2O H3O+ + C N C O O – C N C O O – C N C O O– Succinimide pKa = 9.6 (Resonance structures require separation of charges) Resonance in succinimide anion. No separation of charges. Good delocalization of the negative charge. H3C C4H9NH O C O O NHS Tolbutamide pKa = 5.4 N H NH O O Diphenylhydantoin pKa = 8.3 N H N O O CH2 O P O O O 2 Na+ 67 From the general equation for ionization of acids, Conjugate acid + H2O H3O+ + Conjugate base It is obvious that any structural change that stabilizes the conjugate base more than the conjugate acid will shift the equilibrium to the right and will therefore have an acid strengthening effect (lower the pKa). Conversely, any factor that stabilizes the conjugate acid more than the conjugate base will shift the equilibrium to the left and such factors have an acid weakening effect (a base strengthening effect of the conjugate base). Acidity of Organic Acids Most organic acids (R–COOH, OH , N N N N C HR , some enols, R–OH, etc.) are neutral molecules, the conjugate bases of which are anions. A H H2O H3O+ A –+ + Structural changes that stabilize the anion A:– to a greater extent than the conjugate acid A–H will have an acid strengthening effect. 68 Factors that Affect the Stability of Ions Factors that can bring about a delocalization of the charge on an ion have a stabilizing effect. For any given ion the greater the localization of the charge on a single atom, the greater the energy level of the ion. Factors that contribute to a delocalization of the charge, to a greater dispersal of the charge over several atoms of the molecule, have a stabilizing effect on the ion. Stabilization of anions is brought about by factors that cause a delocalization of electrons (solvation of the ions also plays an important role, but contribution from solvation is more difficult to assess.) Important factors: 1) Resonance 2) Inductive effects 3) Intramolecular hydrogen bonding 4) Steric factors Comparison of acidity of carboxylic acids, phenols, and alcohols (all with acidic proton on an oxygen): H2O H3O++ pKa~5+ ....– .... CR O – O CR O O .. .... CR O O H .. .... .... CR O – O H + .. Less resonance in free acid than in the anion because of separation of + and – charges. The negative charge is equally distributed on the two oxygens. In the carboxylate anion (the conjugate base of carboxylic acids) the negative charge is distributed equally between the two oxygens. Resonance allows a great delocalization of nonbonded electrons which results in the dispersal of the negative charge and stabilization of the anion. Resonance also contributes in another way to the acidity of carboxylic acids, because there is greater resonance in the carboxylate anion than in the undissociated acid. Resonance in the free acid requires a separation of plus and negative charges, which is not the case in the anion where the resonance forms are identical. Increases in resonance are always a stabilizing factor because resonance involves a delocalization of electrons. The gain in resonance in the anion contributes to lowering the energy of the anion in relationship to the free acid and, therefore, contributes to the shifting of the equilibrium to the right. 69 .. O ..– O .. – O .. – .. O – – O H O H – O H – O H .. + + + + H2O + H3O+ pKa~10 There is a greater contribution of resonance in the phenolate anion than in the free phenol, since the resonance structure in the nonionized form requires separation of plus and negative charges and there is no separation of charge in the anion. Resonance also contributes to the stabilization of the anion by contributing to a dispersal of the negative charge through a delocalization of nonbonded electrons. Through resonance, the negative charge is partially localized in the aromatic ring but since the oxygen atom is much more electronegative than the carbon atom, the overall localization of the negative charge is much greater on the oxygen than on the aromatic ring. There is less stabilization of the anion through dispersal of the negative charge in the phenolate than in the carboxylate anion. pKa~16-17 There is no measurable ionization of alcohols in water. A stronger base is needed. The negative charge is localized on the oxygen. B ++ B+ HR CH2 O–R CH2 O H . . . . . . . .: 72 Examples of Electron Perturbing Effects in Carboxylic Acids By Inductive Effect pKa By Resonance and Inductive Effect pKa CH3COOH 4.76 COOH 4.21 HCOOH 3.74 COOHCl inductive effect 3.99 Cl–CH2COOH 2.81 COOHCH3O .... 4.47 Cl2CHCOOH 1.37 COOHN O O– + 3.44 Cl3CCOOH Cl–CH2CH2COOH Cl–CH2CH2CH2COOH 0.65 4.00 4.51 COOHCH3 4.34 Halogens on the α–carbon of carboxylic acids have a significant acid strengthening effect. The halogens are more electronegative than the carbon and there is a permanent dipole moment in the Cl–C bond. This gives a positive character to the α–carbon. δ+ δ+ δ- CCCl H H O O H CC O O Cl H H ........ –δ- δ+ Any dipole effect on the free acid will be a destabilizing effect because of repulsing dipoles. The partial positive charge on the α– carbon attracts electrons from the anion and stabilizes the conjugate base. 73 The influence of the inductive effect falls off rapidly through a saturated aliphatic chain, but is still measurable in 3-chloropropionic acid and 4-chlorobutanoic acid; compare the acidity of these compounds with chloroacetic acid and acetic acid. In general, substituents that cause a decrease in electron density in the vicinity of the carboxyl groups have an acid strengthening effect and those that increase the electron density have an acid weakening effect, as expected. In the aromatic acids, the inductive effect can be transmitted across the ring. Resonance effects are more pronounced when the substituent is either ortho or para to the acidic functional group. – – + CN O O O O – +CN O O O O + – – This resonance hybrid contributes to the stabilization of the anion. Substituent Effects on the Acidity of Phenols pKa pKa OH 9.95 OHO2N 7.14 OHCH3 10.19 OH NO2 8.35 OHCl 9.38 OHO2N NO2 4.04 OHCl Cl 8.05 OHO2N CH3 CH3 8.20 74 The effects of ring substituents on the acidity of phenols are analogous to those seen for aromatic carboxylic acids but the effects are more pronounced for phenols, especially with electron withdrawing groups which participate in resonance with the ring, when these groups are para or ortho to the OH group. Such groups increase the resonance in the anion and increase the delocalization of the negative charge on the anionic conjugate base. ON O O ....+ – – – + .. ON O O – A resonance structure can be written where the nitro group and the phenol oxygen participate in the same molecular orbital and where the negative charge is acquired by the nitro group. This delocalization of electrons and negative charge contributes to a resonance stabilization of the anion. Note that the transfer of the negative charge to the nitro group is only possible when the groups are ortho or para to each other but not when they are meta. Note also that for maximum resonance pariticpation of the nitro group with the ring, it must become coplanar with the ring. Steric factors that prevent this coplanarity will reduce the acid strengthening effect of the nitro group in aromatic system. The lower acidity of 3,5-dimethyl-4-nitrophenol compared to 4– nitrophenol is attributed in part to steric repulsion of the methyl groups which prevents coplanarity of the nitro group and the aromatic ring. ON O O H3C H3C ....+ – – – + .. ON O O H3C H3C – 77 At most pH's the dipolar structure exists: Substituent Effects on pKa of Amino Group pKa1 pKa2 X = H phenylalanine 2.20 9.31 X = OH tyrosine 2.20 9.21 X = Cl p-Cl-phenylalanine 2.18 9.11 Inductive effect of substituent in aromatic ring influences pKa of the α-NH3+ group. Acidic and Basic Groups in Enzyme Catalysis Enzymes are proteins or have a protein component. Proteins are macromolecules made up of amino acids linked by amide or peptide linkages. The amide linkages involve the α–amino and carboxyl groups of the amino acids making up the macromolecules. With the exception of glycine, all amino acids have groups on the α–carbon which result in side chains on the protein molecule. Certain functional groups on these side chains play a role in the catalytic action of enzymes. Certain amino acids have acidic and some have basic groups in the side chain. These groups can play a role in enzyme catalysis, acting as acids and/or bases. X C OH H O H3N+ X C OH H O H2O H2O X C O H O H3O+ + X C O H O H3N+ H3N+ H2N H3O++ 78 Amino Acid Residues with Hydrocarbon Side Chains The side chains of several of the amino acids in proteins are simple hydrocarbons. Thus, when they are incorporated into a protein chain, no additional functional groups are introduced into the molecule. The non-polar nature of these residues, which precludes hydrogen bonding, plays a critical role in the structure of proteins: these residues are hydrophobic. Amino Acid Residues With Carboxylic Acid Side Chains The side chains of two of the amino acids, aspartic and glutamic acids, contain carboxylic acid functional groups linked by a hydrocarbon spacer, of one or two methylene groups respectively, to the α–carbon. At neutral pH, these groups will be present in the anionic conjugate base form (aspartate and glutamate). Amino Acid Residues with Amide Side Chains A further two amino acids, asparagine and glutamine, are closely related to aspartic and glutamic acids. In these, instead of carboxylic acid, the side chain contains an amide group. Amides can participate in hydrogen bonding, but they are neither strong acids nor bases, and do not affect the acid–base chemistry of proteins. 79 Acyclic Amino Acid Residues with Basic Nitrogen-Containing Side Chains Two of the protein amino acids have side chains consisting of a linear carbon chain terminating in a basic nitrogen functional group. The side chain of lysine is a four–carbon chain ending in an amino group. This primary amine bears a non–bonding electron pair and is of similar basicity to the amines considered previously. The pKa of the corresponding ammonium ion is 10.5 and, at neutral pH, this group is present in solution as a cation. In the case of arginine, protonation of the basic nitrogen leads to a cation in which the positive charge is dispersed over three nitrogen atoms. This factor ensures and enhanced stability to the protonated form of arginine which has a pKa of 12.5 and is present as a cation under physiological conditions. Amino Acid Residues with Hydroxyl Functional Groups The side chains of three amino acids contain hydroxyl groups. Serine and threonine are simple alcohols. For each of these residues, the hydroxyl group is attached to a carbon adjacent to the α–carbon. Threonine is distinguished from serine by an extra methyl group that makes it a secondary alcohol. An isolated hydroxyl group can act as an acid or a base, but neither process is especially favorable (the pKa of the hydroxyl or serine is approximately 16). In tyrosine, the hydroxyl function is attached to an aromatic ring. Here the functional group is a phenol. The aromatic ring stabilizes the charge on the deprotonated form. This enhances the stability of the conjugate base and lowers the pKa (to ca. 10) facilitating acid–base chemistry. Tyrosine is usually found in the hydroxyl form, but it is occasionally found to act as an acid under physiological conditions.