## 2.2.6. Ionogenic organic chemicals

Authors: Steven Droge

Reviewer: John Parsons, Satoshi Endo

Learning objectives:

You should be able to:

- understand that IOCs are abundantly present in many types of chemical classes, but all share common features such as dissociation of a proton from a polar moiety (acids to form anions) and association of a proton onto a polar moiety (bases to form cations).
- calculate the fraction of neutral/ionic species for each chemical at a given pKa and pH.
- make some rough predictions about pKa values from the IOC molecular structure.

Keywords: pKa, dissociation constant, speciation, drugs, surfactants, solubility

Introduction

Ionogenic organic chemicals (IOCs) are widely used in industry and daily life, but also abundantly present as chemicals of emerging concern. For environmental risk assessment purposes, IOCs may be defined as organic acids, bases, and zwitterionic chemicals that under common environmental pH conditions exist for a large part as charged (ionic) species, with only modest fraction of neutral species. The environmentally relevant pH range could be argued to lie between 4 (acidic creeks, even lower in polluted streams from volcanic regions or mine drainage systems) to 9 (sewage treatment effluents). The environmental behaviour of IOC pollutants of concern are different from neutral chemicals of concern, because the aqueous pH controls the neutral fraction of dissolved IOCs and the ionic form is highly soluble and interacts partly via electrostatic interactions with environmental susbtrates. Note that the major fraction of an IOC can also be neutral in a certain environmental system, and in that case it is often the neutral form that dominates the chemical’s behavior.

Figure 1. A survey on the 1999 World Drug Index (WDI) database revealed that >63% of the 51,596 chemicals had acidic or basic functionalities. Shown in the circle diagram A is the distribution of different ionizable groups in this fraction of 63% drugs, and in bar diagram B the distribution in pKa values for basic drugs (CNS = drugs acting on central nervous system). Adapted from Manallack (2007) by Wilma IJzerman.

IOCs are common in many different types of pollutant classes. A random subset analysis of all EU (pre-)registered industrial chemicals indicated that large fractions of the total list of >100,000 chemicals are IOCs (51% neutral; 27% acids; 14% bases; 8% zwitterions/amphoterics). In another source, It has been estimated that >60% of all prescription drugs (Section 2.3.3) are IOCs (Manallack, 2007), and even higher fractions for illicit drugs (Section 2.3.4) (Figure 1). Well known examples are basic beta-blockers (e.g. propranolol), basic antidepressants (e.g. fluoxetine and sertraline), acidic non-steroidal anti-inflammatory drugs (NSAID such as diclofenac), and basic opioids (e.g. morphine, cocaine, heroin) and basic designer drugs (e.g. MDMA). The majority of surfactants and polyfluorinated chemicals (e.g. PFOS and GenX) are IOCs (Section 2.3.8), as well as wide variety of important pesticides (e.g. zwitterionic glyphosate) (Section 2.3.1) and (natural) toxins (Section 2.1) (e.g. peptide based multi-ionic cyanobacterial toxins).

Environmental behavior of IOCs

The release into the environment is specific for each of these types of IOCs with a different use, but in many cases happens via sewage treatment systems. If sorption to sewage sludge is very strong, application of sludge onto terrestrial (agricultural) systems is a key entry in many countries. However, many IOCs are rather hydrophilic and will mainly be present in wastewater effluent released into aquatic systems. As they are hydrophilic, they are considered rather mobile which allows for rapid transport through e.g. groundwater plumes, soil aquifers, and (drinking water) filtration steps. The distinction between (mostly) neutral chemicals and IOCs is important because the ionic molecular form generally behaves very differently from the corresponding neutral molecular form. For example, in many aspects ionic molecules are non-volatile compared to the corresponding neutral molecules, while neutral molecules are more hydrophobic than the corresponding ionic molecules. As a result of their lower “hydrophobicity”, ionic molecules often bind with lower affinity to soils and are therefore more mobile in the environment. The ionic forms bioaccumulate to a lower extent and can therefore be less toxic than the corresponding neutral form (though not necessarily). However, there are various important exceptions for these rules. For example, clay minerals sorb cationic IOCs fairly strongly via ion exchange mechanisms. Certain proteins (e.g. the blood serum protein albumin) tightly bind anionic chemicals because of cationic subdomains on specific (enzymatic) pockets, which allows for effective transport throughout our systems and over cell membranes.

Calculating and predicting the dissociation constant (pKa)

The critical chemical parameter describing the ability to ionize is the acid dissociation constant (pKa). The pKa defines at which pH 50% of the IOC is in either the neutral or ionic form by releasing an H+ from the neutral molecule acids (AH to anion A-), or accepting an H+ onto the neutral molecule base (B to cation BH+). The equilibrium between neutral acid and dissociated form can thus be defined as:

[ HA ] ↔ [H+] + [A- ] (eq.1)

where the chemical’s equilibrium speciation is defined as:

(eq.2)

which gives the pKa as:

(eq.3)

and as a function of pH, the ratio of the acid and anion is defined as:

for acids and for bases (eq.4)

Although the term pKb is also used to denote the base association constant, it is conventional we consider [BH+] as acid and use ‘pKa’ and other relationships for bases as well. The fraction of neutral species (fN) for simple IOCs (one acidic or basic site) can be readily calculated with a derivatization of the Henderson-Hasselbalch equation:

(eq.5)

in which α = 1 for bases, and -1 for acids.

in which α = 1 for bases, and -1 for acids. Using equation 5, Figure 2 presents a typical speciation profile for an acid (shown with pKa 5, so perhaps a carboxylic acid) and a base (shown with pKa 9, so perhaps a beta-blocker drug). Following the curve of equation 5, it is interesting to see some simple rules: if the pH is 1 unit lower than the pKa, the deprotonated species fraction is present at 10%. If the pH is 2 units lower than the pKa, the deminished species fraction is present at 1%, 3 units lower would give 0.1%, etc. From this, it is easy to make a good estimate for the protonation of a strong basic drug like MDMA (pKa reported 9.9-10.4) in blood (pH 7.4): with a maximum of 3 units lower pH, up to 99.9% of the MDMA will be in the protonated form, and only 0.1% neutral. For toxicological modeling studies, e.g. in terms of permeation through the blood-brain barrier membrane, this is highly relevant.

Figure 2. pH dependent speciation of an acid (with a pKa of 5) and a base (with a pKa of 9). As shown by the arrows, if pH = pKa, the IOCs exist 50% in the neutral form, and 50% in the ionic form. For the acid at pH = pKa -1 (pH 4), 90% is in the neutral form (AH), and 10% is in the negatively charged form (A-). Drawn by Steven Droge.

Boxes 1-3. Extended learning: calculating the dissociation constant for multiprotic chemicals: see end of this module

Acidic IOCs:

Figure 3. Different types of anionic moieties: carboxylate (anionic form of carboxylic acid), sulfonate (anionic form of sulfonic acid), sulfate (anionic form of sulfuric acid), phenolate (anionic form of phenol). (Source: Steven Droge)

For example, the painkiller (or non-steroidal anti-inflammatory drug, or NSAID) diclofenac is a carboxylic acid with a pKa of 4.1. This means that at pH 4.1, 50% of the dissolved diclofenac is in the dissociated (anionic) form (so, (1 - fN) from equation 5). At pH 5.1 (1 unit higher than the pKa) this is roughly 90% (90.91% to be more precise, but simply remembering 90% helps), at pH 6.1 (2 units higher than the pKa) this is 99%. This stepwise increase in 50-90-99% with each pH unit works for all acids, and for bases the other way round. Test for yourself that at physiological pH of 7.4 (e.g. in blood) diclofenac is calculated to be 99.95% anionic.

Many carboxylic acids have a pKa in the range of 4-5, but the neighboring molecular groups can affect the pKa. Particularly electronegative atoms such as chlorine, fluorine, or oxygen may lower the pKa, as they reduce the forces holding the dissociating proton to the oxygen atom. For example, trichloroacetic acid (CCl3-COOH) has a pKa of 0.77, while acetic acid (CH3-COOH) has a pKa of 4.75. For the same reason, perfluorinated carboxylic acids have a strongly reduced acidic pKa compared to the analogous non-perfluorinated carboxylic acids.

Sulfate acids (see figure 3) are very strong acids, with a pKa <0. These acids almost always occur in their pure form as a salt, for example the common soap ingredient sodium dodecylsulfate (“SDS” or “SLS”, Na.C12H25-SO4). Other common detergents are sulfonates, such as linear alkylbenzenesulfonate (“LAS”, C10-14-(benzyl-SO3)), where the SO3 anionic moiety is attached to a benzene ring, which can be positioned to different carbon atoms of a long alkyl chain. Even at the lower environmental pH range of about 4 these soap chemicals are fully in the anionic form. Such very strong acids, but also many weaker acids, are thus often sold in pure form as salts with sodium, potassium, or ammonium, which causes them to have different names and CAS numbers (e.g. Na.C12H25-SO4 or K.C12H25-SO4) than the neutral form.

Many phenols have a pKa > 8, and are therefore mostly neutral under environmental pH. Electron-withdrawing groups on the aromatic ring of the phenol group, such as Cl, Br and I, can lower the dissociation constant. For example, the dinitrophenol-based pesticide dinoseb has a pKa of 4.6, and is thus mostly anionic in the aquatic environment. Note that a hydroxyl group (-OH) not connected to an aromatic ring, such as -OH of alcohols, can in most cases for risk assessment be considered permanently neutral.

To help interpret the differences in pKa between molecules, it sometimes helps to remember that in more acidic solutions, there are simply higher H+ concentrations, in a logarithmic manner on the pH scale. At pH 3, the concentration H+ protons in solution is 1 mM, while at pH 9 the H+ concentration is 1 nM (6 pH units equals 106 times lower concentrations). The affinity (“Ka”) of H+ to associate with a negatively charged molecular group, is so low for strong acids that even at very high dissolved H+ concentrations (low pH) only very few AH bonds (neutral acid fraction) are actually formed. In other words, for chemicals with a low Ka, even at low pH the neutral fraction is still low. For weak acids such as phenols, already at a very low dissolved H+ concentrations (high pH) many AH bonds (neutral acid fraction) are formed. So it can be reasoned that the affinity of common acidic groups to hold on to a proton is in the order:

SO4 < SO3 < CO2 < amide ( C(=O)NH ) < phenol < hydroxyl

Basic IOCs:

Figure 4. Different types of cationic moieties: primary amine, secondary amine, tertiary amine, quaternary ammonium (permanently charged), pyridinium (cationic form of pyridine), alkylpyridinium (permanently charged). (Source: Steven Droge).

For bases, it is mostly a nitrogen atom that can accept a proton to form an organic cation, because of the lone electron pair in nitrogen. Neutral nitrogen atoms have opportunities for 3 bonds. A primary amine group has the nitrogen atom bonding to only 1 carbon atom (represented here as part of a molecular fragment R), and two additional bonds with hydrogen atoms. The remaining electron pair readily accepts another proton to form a cationic molecule [ R-NH3+ ]. Neutral secondary amines have one bond to hydrogen and two bonds to carbon atoms and can accept a proton to form [ R-NH2+-R’ ], whereas neutral tertiary amines have no bonds to hydrogen but only to carbon atoms and can form [ R-NH+-(R’)(R’’) ]. Of course, each R group may be the same (e.g. a methyl unit).

Many basic chemicals have complex functionalities that can influence the pKa of the nitrogen moiety. However, as shown in the examples of figure 5, as long as there are at least two carbon atoms between the amine and the polar molecular fragment (for example OH, but much stronger for =O), the pKa of the basic nitrogen groups in all three types of bases (primary, secondary and tertiary amines) is high, often above 9 (dissolved H+ concentration <10-9). So even at very low H+ concentrations, dissolved protons like to be associated to such amine groups. As a result, amines such as most beta-blockers and amphetamine based drugs are predominantly positively charged molecules (organic cationic amines) in the common environmental pH range of 4-9, as well as in the pH of most biotic tissues that are useful for toxicological assessments. As soon as a polar group with oxygen (e.g. ketones or hydroxyl groups) is connected to the second carbon away from the nitrogen (e.g. R-CH(OH)-CH2-NH2) the pKa is considerably lowered. Also nitrogen atoms as part of an aromatic ring, or connected to an aromatic ring, have much lower pKa’s: protons have rather low affinity to bind to these N-atoms and only start doing so if the proton concentration becomes relatively high (solution becomes more acidic).

Figure 5. Different proton dissociation constants for amine groups: pKa is influenced by other functional structures. (Source: Steven Droge)

Relevance of accounting for electrostatic interactions

Most classical pollutants, such as DDT, PCBs and dioxins, are neutral hydrophobic chemicals. On the other hand, most metals are almost always cationic species (e.g. Cd2+). Consequently, their environmental distribution and biological exposure are influenced by quite distinct processes. Obviously, predominantly charged IOCs behave somewhat in between these two extremes. The charged positive or negative groups cause a strong effect of electrostatic interactions between the IOC and environmental substrates (sorption or ligand/receptor binding). While also metals speciate into different forms, pH difference between environmental compartments can strongly influence the IOCs chemical fate and effect if ionizable group is relatively weak. An important difference to metals is that the nonionic molecular part still influences the IOC’s hydrophobicity even in charged state for several processes.

As will be discussed in other chapters regarding chemical processes (see Chapter 3), it needs to be taken into account for IOCs that many environmental substrates (DOC, soil organic matter, clay minerals) are mostly negatively charged in the common range of environmental pH, but also that proteins involved in biotic uptake-distribution-effects are rich in ionogenic peptides that are part of binding pockets and reactive centers.

References

Manallack, D.T. (2007). The pKa distribution of drugs: application to drug discovery. Perspectives in Medical Chemistry 1, 25-38.

Box 1. Extended learning: calculating the dissociation constant for multiprotic chemicals: Several common inorganic acids are multiprotic: they have multiple protons that can dissociate. Multiples species can occur at a certain pH, such as for phosphoric acid (H3PO4, H2PO4-, HPO42- , HPO42-), and carbonic acid (H2CO3, HCO3-, CO32-). It is important to realize that there are actually two micro-species of HCO3-, because two hydroxylgroups can dissociate: HO-C(=O)-OH A polyprotc acid HnA can undergo n dissociations to form n+1 species. Each dissociation has a pKa. But how to calculate the fraction of each species of multiprotic chemicals? The charge of a polyprotic acid can be described as Hn-jAj-. A useful variable, v, can be defined for each general polyprotic acid: for each dissociation reaction: Hn-j+1A-j+1 ↔ H+ + Hn-jAj- the dissociation constant Kj is: The degree of dissociation of the acid (η) is equal to the ratio of the total charge (TC) to the total mol acid (TM). For a diprotic acid, a plot of η as a function of pH provides the dissociation curve: which can be a fitted curve to experimental data. The degree of protolysis for the jth species, αj, can be calculated from the ratio of [Hn-jAj-]/TM You can set up such a calculation in MS Excel, with calculations of α0, α1, α2, at a range of different pH values ([H+] concentrations), for a given K1 and K2, and plot the speciation against pH. More details are described by King et al. (1990) J. Chem. Educ. 67 (11), p. 932; DOI: 10.1021/ed067p932

Box 2. Example 1 for multiprotic chemicals: Carbonic acid Let’s try carbonic acid (bicarbonate, or H2CO3) as an example first. H2CO3 is the product of carbon dioxide dissolved in water. In pure water/seawater the hydration equilibrium constant Kh = [H2CO3]/[CO2] ≈ 1.7×10−3 / ≈ 1.2×10−3 respectively, indicating that only 0.1% of dissolved CO2 equilibrates to H2CO3. The dissolved concentration of CO2 depends on atmospheric CO2 levels according to the air-water distribution coefficient (Henry constant kH = pCO2/[CO2]= 29.76 atm/(mol/L)). Because of the relevance of CO2 in e.g. ocean acidification, and gas exchange in our lungs, it is interesting to see how H2CO3 speciates depending on pH. As in the formula HnA, n= 2 for H2CO3. With pK1*=6.5 (in equilibrium with atmospheric CO2) and pK2=10.33, K1=10-6.5 and K2=10-10.3. At pH 7, [H+]=10-7 , so at pH 7 with these dissociation constants or from a series of Excel calculations at different pH: you can copy/paste the following cells into cell A1 of a new Excel sheet, and extend the range of pH:
K1 | 3.16E-07
K2 | 5.01E-11
pH | 5
[H+] | =10^-B3
a0 = H2CO3 | =(B4*B4)/((B4*B4)+(B4*$B$1)+($B$1*$B$2))
a1 = HCO3- | =(B4*$B$1)/((B4*B4)+(B4*$B$1)+($B$1*$B$2))
a2 = CO32- | =($B$1*$B$2)/((B4*B4)+(B4*$B$1)+($B$1*$B$2))
ETA | =((B4*$B$1)+2*($B$1*$B$2))/((B4*B4)+(B4*$B$1)+($B$1*$B$2))

Box 3. Example 2 for multiprotic chemicals: Zwitterions Many organic pH buffers are zwitterionic chemicals, that contain both an acidic and a basic group. Norman Good and colleagues described a set of 20 of such buffers for biochemical and biological research (see for example www.interchim.fr/ft/0/062000.pdf, or www.applichem.com/fileadmin/Broschueren/BioBuffer.pdf). Examples are MES, MOPS, HEPPS (Figure A). These buffers are selected to: have a buffering pKa in the range of pH6-8 where most biochemical tests are performed; be readily soluble in water, be stable in test solutions, so resistant to (non)enzymatic degradation, not forming precipitates with salts ideally be impermeable to cell membranes so that they don’t accumulate or reach active intracellular sites readily available, reasonably cheap. the zwitterionic chemicals with sulfate groups are actually always have the sulfate group charged, making it highly soluble and impermeable to cell membranes, while the amine group protonates between pH6-10, depending on neighbouring functional groups. The speciation of the amine groups in MES and MOPS simply follows the single pKa calculation of equations 1-5. In HEPPS, either of the two amines is protonated, the second pKa is 3, so the doubly charged molecules only occurs at much lower pH, but can still be used as a buffer. Figure A. MES, MOPS and HEPPS in charged form. A zwitterionic chemical with two apparent pKa values relatively close is p-amino-benzoic acid. If chemicals have not one ionisable group, but N ionisable groups that speciation in a relevant pH range, than the amount of possible species is 2N. So the zwitterion p-amino benzoic acid has 4 species, each with a separate pKa (pH where both species are present in equal concentrations). Let’s formulate the benzyl group in p-amino benzoic acid as X, the neutral amine base as B, and the neutral carboxylic acid as AH, so that the fully neutral species is BXAH. Compared to the carboxylic acid, we now have under the most acidic conditions (BHXAH)+, as , the neutral species BXAH and the zwitterionic intermediate (BHXA)0, as , and anionic species at most alkaline conditions (BXA)-, as . The calculation of the fraction of each species can be calculated according to similar rules as for carbonic acid if the two dissociation constants are known (K1 = 2.4, K2=4.88). However, this does not inform us on the ratio between the zwitterionic form and the fully neutral form. To do this, the speciation constants of the 4 microspecies are required. [BH+XAH] ↔ [BXAH] + [H+] for which the pk1 is calculated to be 2.72 K1=10^-2,72 = [BXAH]*[H+]/ [BH+XAH] [BH+XAH]*10^-2.72 = [BXAH]*[H+] which rearranges to [BXAH]= 10^-2.72 *[BH+XAH] /[ H+] [BH+XAH] ↔ [BH+XA-] + [H+] for which the pk2 is calculated to be 3.93 K2=10^-3.93 = [BH+XA-]*[H+] / [BH+XAH] [BH+XAH]*10^-3.93 = [BH+XA-]*[H+] which rearranges to [BH+XA-] =10^-3.93*[BH+XAH]/[H+] [BXAH] ↔ [BXA-] + [H+] for which the pk3 is calculated to be 4.74 K3= 10^-4,74 = ([BXA-] * [H+]) / [BXAH] [BXAH]*10^-4,74 = [BXA-] * [H+]) [BH+XA-] ↔ [BXA-] + [H+] for which the pk4 is calculated to be 4.31 K4=10^-4,31 = [BXA-]*[H+] / [BH+XA-] [BH+XA-]*10^-4,31 = [BXA-]*[H+] So the ratio between zwitterionic form [BH+XA-] and neutral form [BXAH] equals to: [BH+XA-] / [BXAH] = 10^-pK2 / 10^-pK1 [BH+XA-] / [BXAH] = 10^-3.93 / 10^-2.72 = 0.06: so only 6% zwitterionic vs 94% neutral species. The macro pK1 and pK2 are then calculated as: K1 = ([BXAH] *[H+] + [BH+XA-]*[H+] )/ [BH+AXH] K1 = [BXAH] *[H+]/[BH+XAH] + [BH+XA-]*[H+] / [BH+XAH] K1 = 10^-pK1 + 10^- pK2 K1 = 10^-2.72 + 10^-3.93 =10^-2.69 K1=10^-2.69 K2 = ([BXA-]) *[H+] / ( [BH+XA-] + [BXAH] ) 1/K2 = [BH+XA-]/([BXA-]) *[H+] + [BXAH] /([BXA-]) *[H+] 1/K2 = 1/10^-pK3 + 1/10^-pK4 1/K2 = 1/10^-4,31 + 1/10^-4,74 = 1/10^-4.87 K2 =10^-4.87