12: Chemical Principles

Chemistry is the science of matter; it encompasses the structure, physicalproperties, and reactivities of atoms and their compounds. In many respects, toxicology is the science of the interactions of matter with living entities. Chemistry and toxicology are therefore intimately linked. The study of the principles of inorganic, organic, and biologic chemistry offers important insight into the mechanisms and clinical manifestations of xenobiotics and poisoning. This chapter reviews many of these tenets and provides relevance to the current practice of medical toxicology.

Basic Structure

Matter includes the substances of which everything is made. Elements are the foundation of matter, and all matter is made from one or more of the known elements. An atom is the smallest quantity of a given element that retains the properties of that element. Atoms consist of a nucleus, incorporating protons and neutrons, coupled with its orbiting electrons. The atomic number is the number of protons in the nucleus of an atom, and it is a whole number that is unique for each element. Thus, elements with 6 protons are always carbon, and all forms of carbon have exactly 6 protons. However, although the vast majority of carbon nuclei have 6 neutrons in addition to the protons, accounting for an atomic mass (protons plus neutrons) of 12 (¹²C), a small proportion of naturally occurring carbon nuclei, called isotopes, have 8 neutrons and a mass number of 14 (¹⁴C). This is the reason that the atomic weight of carbon displayed on the periodic table is 12.011, and not 12, as it actually represents the average atomic masses of all isotopes found in nature weighted by their frequency of occurrence. Moreover, ¹⁴C is actually a radioisotope, which is an isotope with an unstable nucleus that emits radiation (particles or rays) until it achieves a stable state (Chap. 134). The atomic weight, measured in grams per mole (g/mol), also indicates the molar mass of the element. That is, in one atomic weight (12.011 g for ­carbon) there is one mole (6.023 × 10²³) of atoms.

Elements combine chemically to form compounds, which generally have physical and chemical properties that differ from those of the constituent elements. The elements in a compound can be separated only by chemical means that destroy the original compound, as occurs during the burning (ie, oxidation) of a hydrocarbon, a process that releases the carbon principally as carbon dioxide and the hydrogen principally as water (H₂O). This important property differentiates compounds from mixtures, which are combinations of elements or compounds that can be separated by physical means such as the distillation of petroleum into its hydrocarbon components or the evaporation of seawater to separate water from sodium chloride and other substances. With notable exceptions, such as the elemental forms of many metals or halogens (eg, Cl₂), most xenobiotics are compounds or mixtures.

Dimitri Mendeleev, a Russian chemist in the mid-19th century, recognized that when all of the known elements were arranged in order of atomic weight, certain patterns of reactivity became apparent. The result of his work was the Periodic Table of the Elements (Fig. 12–1), which, with some minor alterations, is still an essential tool today. All of the currently recognized elements are represented; those heavier than uranium are known not to occur in nature. Many of the symbols used to identify the elements refer to the Latin name of the element. For example, silver is Ag, for argentum, and ­mercury is Hg, for hydrargyrum, literally “silver water.”

The reason for the periodicity of the table relates to the electrons that circle the nucleus in discrete orbitals. Although the details of quantum mechanics and electronic configuration are complex, it is important to review some aspects in order to predict chemical reactivity. Orbitals, or quantum shells, represent the energy levels in which electrons may exist around the nucleus. The orbitals are identified by various schemes, but the maximum number of electrons each orbital may contain is calculated as 2x², where x represents the numerical rank order of the orbital. Thus, the first orbital may contain 2 electrons, the second orbital may contain 8, the third may contain 18, and so on. However, the outermost shell (designated by s, p, d nomenclature) of each orbital may only contain up to 8 electrons. This is irrelevant through element 20, calcium, because there is no need to fill the third-level or d shells. Even though the third orbital may contain 18 electrons, once 8 are present the 4s electrons dip below the 3d electrons in energy and this shell begins to fill. This occurs at element 21, scandium, and accounts for its chemical properties and those of the other transition elements. Transition elements are chemically defined as elements that form at least one ion with a partially filled subshell of d electrons. Also, because the inert gas elements, which are also known as noble gases, have complete outermost orbitals, they are unreactive under standard conditions.

In general, only electrons in unfilled shells, or valence shells, are involved in chemical reactions. This property relates to the fact that the most stable form of an element occurs when the configuration of its valence shell resembles that of the nearest noble gas, found in group 0 on the periodic table. This state can be obtained through the gaining, losing, or sharing of electrons with other elements and is the basis for virtually all chemical reactions.

The Periodic Table

Chemical Reactivity.

Broadly, the periodic table is divided into metals and nonmetals. Metals, in their elemental form, are typically malleable solids that conduct electricity, whereas nonmetals are usually dull, fragile, nonconductive compounds (C, N, P, O, S, Se, halogens). The metals are found on the left side of the periodic table, and account for the majority of the elements, whereas the nonmetals are on the right side. Separating the two groups are the metalloids, which fall on a jagged line starting with boron (B, Si, Ge, As, Sb, Te, At). Metalloids have chemical properties that are intermediate between the metals and the nonmetals. Each column of elements is termed a family or group, and each row is a period. Although conceived and organized in periods, trends in the chemical reactivity, and therefore toxicity, typically exist within the groups.

The ability of any particular element to produce toxicologic effects relates directly to one or more of its many physicochemical properties. These properties may, to some extent, be predicted by their location on the periodic table. For example, arsenate may substitute for phosphate in the mitochondrial production of adenosine triphosphate (ATP), creating adenosine diphosphate monoarsenate (Chap. 13). Because this compound is unstable and not useful as an energy source, energy production by the cell fails; in this manner arsenic interferes with oxidative phosphorylation. The existence of an interrelationship between Ca²⁺ and either Mg²⁺ or Ba²⁺ is predictable, although the actual effects are not. Under most circumstances Mg²⁺ is a Ca²⁺ antagonist, and patients with hypermagnesemia present with neuromuscular weakness caused by blockade of myocyte calcium channels. Alternatively, Ba²⁺ mimics Ca²⁺ and closes Ca²⁺-dependent K⁺ channels in myocytes, producing life-threatening hypokalemia. The physiologic relationships among the ions of lithium (Li⁺), potassium (K⁺), and sodium (Na⁺) are also consistent with their chemical similarities (all alkali metals in Group IA). However, the clinical similarity between thallium (thallous) ion (Tl⁺) and K⁺ is not immediately predictable. Other than their monovalent nature (ie, +1 charge), it is difficult to predict the substitution of Tl⁺ (Group IIIA, Period 6) for K⁺ (Group IA, Period 4) in membrane ion channel functions, until the similarity of their ionic radii is known (Tl⁺, 1.47 Å; K⁺, 1.33 Å).

Alkali and Alkaline Earth Metals.

Alkali metals (Group IA: Li, Na, K, Rb, Cs, Fr) and hydrogen (not an alkali metal on earth) have a single outer valence electron and lose this electron easily to form compounds with a valence of 1+. The alkaline earth metals (Group IIA: Be, Mg, Ca, Sr, Ba, Ra) (between the alkali and rare earth, Group IIIB) readily lose two electrons, and their cations have a 2+ charge. In their metallic form, members of both of these groups react violently with water to liberate strongly basic solutions accounting for their group names (2Na⁰ + 2H₂O → 2NaOH + H₂). The soluble ionic forms of sodium, potassium, or calcium, which are critical to survival, also produce life-threatening symptoms following excessive intake (Chap. 19). Xenobiotics may interfere with the physiologic role of these key electrolytes. Li⁺ may mimic K⁺ and enter neurons through K⁺ channels, where it then serves as an inadequate substrate for the repolarizing of Na⁺-K⁺-ATPase. Thus, Li⁺ interferes with cellular K⁺ homeostasis and alters neuronal repolarization accounting for the neuroexcitability manifesting as tremor. Similarly, as noted previously, the molecular effects of Mg²⁺ and Ba²⁺ may supplant those of Ca²⁺. More commonly, though, the consequential toxicities ascribed to alkali or alkaline earth salts actually relate to the anionic component. In the case of NaOH or Ca(OH)₂, it is the hydroxide anion (not the hydroxyl radical), while it is the CN^– anion in patients poisoned with potassium cyanide (KCN).

Transition Metals.

Unlike the alkali and alkaline earth metals, most other metallic elements are neither soluble nor reactive. This includes the transition metals (Group IB to VIIIB), a large group that contains several ubiquitous metals such as iron (Fe) and copper (Cu). These elements, in their metallic form, are widely used in both industrial and household applications because of their high tensile strength, density, and melting point, which is partly a result of their ability to delocalize the electrons in the d orbital throughout the metallic lattice. Transition metals also form brightly colored salts that find widespread applications, including as pigments for paints or fireworks. The ionic forms, unlike the metallic forms, of these elements are typically highly reactive and toxicologically important. ­Chemically, the transition elements are elements that form at least one ion with a partially filled subshell of d electrons. Because the transition metals have partially filled valence shells, they are capable of obtaining several, usually positive, oxidation states. This important mechanism explains the role of transition metals in redox reactions generally as electron acceptors (see Reduction-Oxidation below). This reactivity is used by organisms in various physiologic, catalytic, and coordination roles, such as at the active sites of enzymes and in hemoglobin, respectively. As expected, the substantial reactivity of these transition metal elements is associated with cellular injury caused by several mechanisms, including the generation of reactive oxygen species (Fig. 12–2). For example, manganese ion exposure is implicated in the free radical damage of the basal ganglia causing parkinsonism.

Heavy Metals.

Heavy metal is often loosely used to describe all metals of toxicologic significance, but in reality the term should be reserved to describe only those metals in the lower period of the periodic table, particularly those with atomic masses greater than 200 Da. The chemical properties and toxicologic predilection of this group vary among the elements, but their unifying toxicologic mechanism is electrophilic interference with nucleophilic sulfhydryl-containing enzymes. Some of the heavy metals also participate in the generation of free radicals through Fenton ­chemistry (Fig. 12–2). The likely determinant of the specific toxicologic effects ­produced by each metal is the tropism for various physiologic systems, enzymes, or microenvironments; thus lipophilicity, water solubility, ionic size, and other physicochemical parameters are undoubtedly critical. Also, because the chemistry of metals varies dramatically based on the chemical form (ie, organic, inorganic, or elemental), as well as the charge on the metal ion, prediction of the clinical effects of a particular metal is often difficult.

Mercury.

Elemental mercury (Hg⁰) is unique in that it is the only metal that exists in liquid form at room temperature and as such is capable of ­creating solid solutions, or amalgams, with other metals. Although it is relatively innocuous if ingested as a liquid, it is readily volatilized (ie, high vapor pressure), transforming it into a significant pulmonary mucosal irritant upon inhalation. In addition, this change in the route of exposure increases its systemic bioavailability. Absorbed, or incorporated, Hg⁰ undergoes biotransformation in the erythrocyte and brain to the mercuric (Hg²⁺) form, which has a high affinity for sulfhydryl-containing molecules, including proteins. This causes a depletion of glutathione in organs such as the kidney and also initiates lipid peroxidation. The mercurous form (Hg⁺) is considerably less toxic than the mercuric (Hg²⁺) form, perhaps because of its reduced water solubility, which limits absorption. Organic mercurial compounds, such as methylmercury and dimethylmercury, are environmentally formed by anaerobic bacteria containing the methylating compound methylcobalamin, a vitamin B₁₂ analog (Chap. 98).

Thallium.

Another toxicologically important member of the heavy metal group is thallium. Metallic thallium (Tl⁰) is used in the production of electronic equipment and is itself minimally toxic. Thallium ions, ­however, have physicochemical properties that closely mimic potassium ions, allowing them to participate in, and potentially alter, the myriad physiologic activities related to potassium. This property is exploited during a thallium-stress test to assess for myocardial ischemia or infarction. Because ischemic myocardial cells lack adequate energy for normal Na⁺-K⁺-ATPase function, they cannot exchange sodium for potassium (or in this scenario radioactive thallium), producing a “cold spot” in the ischemic areas on cardiac scintigraphy (Chap. 102).

Lead.

Although lead is not very abundant in the Earth’s crust (only 0.002%), exposure may occur during the smelting process or from one of its diverse commercial applications. Most of the useful lead compounds are inorganic plumbous (Pb²⁺) salts, but plumbic (Pb⁴⁺) compounds are also used. The Pb²⁺ compounds are typically ionizable, releasing Pb²⁺ when ­dissolved in a solvent, such as water. Pb²⁺ ions are absorbed in place of Ca²⁺ ions by the gastrointestinal tract and replace Ca²⁺ in certain physiologic processes. This mechanism is implicated in the neurotoxic effect of lead ions. Pb⁴⁺ compounds tend to be covalent compounds that do not ionize in water. Some of the Pb⁴⁺ compounds are oxidants. Although elemental lead is not itself toxic, it rapidly develops a coating of toxic lead oxide or lead carbonate on exposure to air or water (Chap. 96).

Metalloids.

Although the metalloids (B, Si, Ge, As, Sb, Te, At) share many physical properties with the metals, they are differentiated by their propensity to form compounds with both metals and the nonmetals carbon, nitrogen, or oxygen. Thus, metalloids may be either oxidized or reduced in chemical reactions.

Arsenic.

Toxicologically important inorganic arsenic compounds exist in either the pentavalent arsenite (As⁵⁺) form or the trivalent arsenate (As³⁺) form. The reduced water solubility of the arsenate compounds, such as arsenic pentoxide, accounts for its limited clinical toxicity when compared to trivalent arsenic trioxide. The trivalent form of arsenic is primarily a nucleophilic toxin, binding sulfhydryl groups and interfering with enzymatic ­function (Chaps. 13 and 89).

Nonmetals.

The nonmetals (C, N, P, O, S, Se, halogens) are highly electronegative and, unlike the metals, may be toxic in either their compounded or their elemental form. The nonmetals with large electronegativity, such as O₂ or Cl₂, generally oxidize other elements in chemical reactions. Those with lesser electronegativity, such as C, behave as reducing agents.

Halogens.

In their highly reactive elemental form, which is a covalent dimer of halogen atoms, the halogens (F, Cl, Br, I, At) carry the ­suffix -ine (eg, Cl₂, chlorine). Halogens require the addition of one electron to complete their valence shell; thus, halogens are strong oxidizing agents. Because they are highly electronegative, they form halides (eg, Cl⁻, chloride) by abstracting electrons from less electronegative elements. Thus, the halogen ions, in their stable ionic form, generally carry a charge of −1. The halides, although much less reactive than their respective elemental forms, are reducing agents. The hydrogen halides (eg, HCl, hydrogen chloride) are gases under standard conditions, but they ionize when dissolved in aqueous solution to form the hydrohalidic acids (eg, HCl, hydrochloric acid). All hydrogen halides except HF (hydrogen fluoride) ionize nearly completely in water to release H⁺ and are considered strong acids. Because of its small ionic radius, lack of charge dispersion, and the intense electronegativity of the fluorine atom, HF ionizes poorly and is a weak acid. This specific property of HF has important toxicologic implications (Chap. 107).

Group 0: Inert Gases.

Inert gases (He, Ne, Ar, Kr, Xe, Rn), also known as noble gases, maintain completed valence shells and are thus entirely unreactive except under extreme experimental conditions. However, despite their lack of chemical reactivity, the inert gases are toxicologically important as simple asphyxiants. That is, because they displace ambient oxygen from a confined space, hypoxia may occur (Chap. 124). During high-concentration exposure, inert gases may produce anesthesia, and xenon is used as an anesthetic agent. Radon, although a chemically unreactive gas, is radioactive, and prolonged exposure is associated with the development of lung cancer.

Bonds

Electrons are not generally shared evenly between atoms when they form a compound. Instead, unless the bond is between the same elements, as in Cl₂, one of the elements exerts a larger attraction for the shared electrons. The degree to which an element draws the shared electron is determined by the electronegativity of the element (Fig. 12–3). The electronegativity of each element was catalogued by Linus Pauling and relates to the ionic radius, or the distance between the orbiting electron and the nucleus, and the shielding effects of the inner electrons. The electronegativity rises toward the right of the periodic table, corresponding with the expected charge obtained on an element when it forms a bond. Fluoride ion has the highest electronegativity of all elements, which explains many of its consequential toxicologic properties.

Several types of bonds exist between elements when they form compounds. When one element gains valence electrons and another loses them, the resulting elements are charged and attract one another in an ionic, or electrovalent, bond. An example is sodium chloride (NaCl), or table salt, in which the electronegativity difference between the elements is 1.9, or greater than the electronegativity of the sodium (Fig. 12–3). Thus, the chloride wrests control of the electrons in this bond. In solid form, ionic compounds exist in a crystalline lattice, but when put into solution, as in the serum, the elements may separate and form charged particles, or ions (Na⁺ and Cl^–). The ions are stable in solution, however, because their valence shells contain eight electrons and are complete. The properties of ions differ from both the original atom from which the ion is derived and the noble gas with which it shares electronic structure.

It is important to recognize that when a mole of a salt, such as NaCl (molecular weight 58.45 g/mol), is put in aqueous solution, two moles of particles result. This is because NaCl essentially ionizes fully in water; that is, it produces one mole of Na⁺ (23 g/mol) and one mole of Cl⁻ (35.5 g/mol). For salts that do not ionize completely, less than the intrinsic number of moles are released and the actual quantity liberated can be predicted based on the defined solubility of the compound, or the solubility product constant (K_sp). For ions that carry more than a single charge, the term equivalent is often used to denote the number of moles of other particles to which one mole of the substance will bind. Thus, an equivalent of calcium ion will typically bind two moles (or equivalents) of chloride ions (which are monovalent) because calcium ions are divalent. Alternatively stated, a 10% calcium chloride (CaCl₂) aqueous solution contains approximately 1.4 mEq/mL or 0.7 mmol/mL of Ca²⁺.

Compounds formed by two elements of similar electronegativity have little ionic character because there is little impetus for separation of charge. Instead, these elements share pairs of valence electrons, a process known as covalence. The resultant molecule contains a covalent bond, which is typically very strong and generally requires a high-energy chemical reaction to disrupt it. There is wide variation in the extent to which the electrons are shared between the participants of a covalent bond, and the physicochemical and toxicologic properties of any particular molecule are in part determined by its nature. Rarely is sharing truly symmetric, as in oxygen (O₂) or chlorine (Cl₂). If sharing is asymmetric and the electrons thus exist to a greater degree around one of the component atoms, the bond is polar. However, the presence of a polar bond does not mean that the compound itself is polar. For example, methane contains a carbon atom that shares its valence electrons with four hydrogen atoms, in which there is a small charge separation between the elements (electronegativity {EN} difference = 0.40). However, because the molecule is configured in a tetrahedral formation, there is no notable polarity to the compound; this compound is nonpolar. The lack of polarity suggests that methane ­molecules have little affinity for other methane molecules and they are held together only by weak intermolecular bonds. This explains why methane is highly volatile under standard conditions and therefore exists as a gas under standard temperature and pressure.

Because the EN differences between hydrogen (EN = 2.20) and oxygen (EN = 3.44) are greater (EN difference = 1.24), the electrons in the hydrogen-­oxygen bonds in water are drawn toward the oxygen atom, ­giving it a partial negative charge and the hydrogen a partial positive charge. Furthermore, because H₂O is angular, not linear or symmetric, water is a polar molecule. Water molecules are held together by hydrogen bonds, which are stronger than other intermolecular bonds (for example, van der Waals forces; see later). These hydrogen bonds have sufficient energy to open many ionic bonds and solvate the ions. In this process, the polar ends of the water molecule surround the charged particles of the dissolved salt. Thus, because there is little similarity between the nonpolar methane and the polar water molecules, methane is not water soluble. ­Similarly, salts cannot be solvated by nonpolar compounds, and thus a salt, such as sodium chloride, cannot dissolve in a nonpolar solvent, such as carbon tetrachloride.

Alternatively, the stability and irreversibility of the bond between an organic phosphorus insecticide and the cholinesterase enzyme are a result of covalent phosphorylation of an amino acid at the active site of the enzyme. The resulting bond is essentially irreversible in the absence of another chemical reaction (Fig. 113–3).

Compounds may share multiple pairs of electrons. For example, the two carbon atoms in acetylene (HC≡CH) share three pairs of electrons between them, and each shares one pair with a hydrogen. Carbon and nitrogen share three pairs of electrons in forming cyanide (C≡N^–), making this bond very stable and accounting for the large number of xenobiotics capable of liberating cyanide. Complex ions are covalently bonded groups of elements that behave as a single element. For example, hydroxide (OH^–) and sulfate (SO₄^2–) form sodium salts as if they were simply the ion of a single element (such as chloride).

Noncovalent bonds, such as hydrogen or ionic bonds, are important in the interaction between ligands and receptors as well as between ion channels and enzymes. These are low-energy bonds and easily broken. Van der Waals forces, also known as London dispersion forces, are intermolecular forces that arise from induced dipoles as a consequence of nonuniform distribution of the molecular electron cloud. These forces become stronger as the atom (or molecule) becomes larger because of the increased polarizability of the larger, more dispersed electron clouds. This accounts for the fact that under standard temperature and pressure, fluorine and chlorine are gases, whereas bromine is a liquid, and iodine is a solid.

Reduction-Oxidation

Reduction-oxidation (redox) reactions involve the movement of electrons from one atom or molecule to another, and actually comprise two dependent reactions: reduction and oxidation. Reduction is the gain of electrons by an atom that is thereby reduced. The electrons derive from a reducing agent, which in the process becomes oxidized. Oxidation is the loss of electrons from an atom, which is, accordingly, oxidized. An oxidizing agent accepts electrons and, in the process, is reduced. By definition, these chemical reactions involve a change in the valence of an atom. It is also important to note that acid–base and electrolyte chemical reactions involve electrical charge interactions but no change in valence of any of the involved components. The implications of redox chemistry for toxicology are significant. For example, the oxidation of ferrous (Fe²⁺) to ferric (Fe³⁺) iron within the hemoglobin molecule creates the dysfunctional methemoglobin molecule (Chap. 127).

Also, elemental lead and mercury are both intrinsically harmless metals but when oxidized to their cationic forms, both produce devastating clinical effects. Additionally, the metabolism of ethanol to acetaldehyde involves a change in the oxidation state of the molecule. In this case, an enzyme, alcohol dehydrogenase, acting as a catalyst oxidizes (ie, removes electrons from) the C-O bond and delivers the electrons to oxidized nicotinamide adenine dinucleotide (NAD⁺), reducing it to NADH. As in this last example, oxidation is occasionally used to signify the gain of oxygen by a substance. That is, when elemental iron (Fe⁰) undergoes rusting to iron oxide (Fe₂O₃), it is said to oxidize. The use of this term is consistent because in the process of oxidation, oxygen derives electrons from the atom with which it is reacting and combining.

Reactive Oxygen Species.

Free radicals are reactive molecules that contain one or more unpaired electrons and are typically neutral but may be anionic or cationic. However, because certain toxicologically important reactive molecules, such as hydrogen peroxide (H₂O₂) and ozone (O₃), do not contain unpaired electrons, the term reactive species is preferred. The reactivity of these molecules directly relates to their attempts to fill their outermost orbitals by receiving an electron; the result is oxidative stress on the biologic system. Molecular oxygen is actually a diradical with two unpaired electrons in the outer orbitals. However, its reactivity is less than that of the other radicals because the unpaired electrons have parallel spins, so catalysts (ie, enzymes or metals) are typically involved in the use of oxygen in biologic processes.

Reactive species are continuously generated as a consequence of endo­genous metabolism, and there is an efficient detoxification system for their control. However, under conditions of either excessive endogenous ­generation or exposure to exogenous reactive species, the physiologic defenses against these toxic products are overwhelmed. When this occurs, reactive species induce direct cellular damage as well as initiate a cascade of oxidative reactions that perpetuate the toxic damage.

Intracellular organelles, particularly the mitochondria, may also be disrupted by various reactive species. This causes further injury to the cell as energy failure occurs. This initial damage is exacerbated by the activation of the host inflammatory response by chemokines that are released from cells in response to reactive species-induced damage. This inflammatory response aggravates cellular damage. The resultant membrane dysfunction or damage causes cellular apoptosis or necrosis.

The most important reactive oxygen species in medical toxicology are derived from oxygen, although those derived from nitrogen are also important. Table 12–1 lists some of the important reactive oxygen and nitrogen species.

The biradical nature of oxygen explains both the physiologic and toxicologic importance of oxygen in biologic systems. Physiologically, the majority of oxygen is used by the body to serve as the ultimate electron acceptor in the mitochondrial electron transport chain. In this situation, four electrons are added to each molecule of oxygen to form two water molecules (O₂ + 4H⁺ + 4e^– → 2H₂O).

Superoxide (O₂^–) is “mutated” from oxygen within neutrophil and ­macrophage lysosomes as part of the oxidative burst, which serves to help eliminate infectious agents and damaged cells. Superoxide may subsequently be enzymatically converted, or “dismutated,” into hydrogen peroxide by superoxide dismutase (SOD). Hydrogen peroxide may be subsequently converted into hypochlorous acid by the enzymatic addition of chloride by myeloperoxidase. Both hydrogen peroxide and hypochlorite ion are more potent reactive oxygen species than superoxide. However, this lysosomal protective system may also be responsible for tissue damage following poisoning as the innate inflammatory response attacks xenobiotic-damaged cells. Examples include acetaminophen (APAP)-induced hepatotoxicity (Chap. 35), carbon monoxide neurotoxicity (Chap. 125), and chlorine-induced pulmonary toxicity (Chap. 124), each of which may be altered, at least in experimental systems, by the addition of scavengers of reactive species.

Although superoxide and hydrogen peroxide are reactive species, it is their conversion into the hydroxyl radical (OH·) that accounts for their most consequential effects. The hydroxyl radical is generated by the ­Fenton ­reaction (Fig. 12–2), in which hydrogen peroxide is decomposed in the presence of a transition metal. This catalysis typically involves Fe²⁺, Cu⁺, Cd²⁺, Cr⁵⁺, Ni²⁺, or Mn²⁺. The Haber-Weiss reaction (Fig. 12–2), in which a transition metal catalyzes the combination of superoxide and hydrogen peroxide, is the other important means of generating the hydroxyl radical. Alternatively, superoxide dismutase, within the erythrocyte, contains an ion of copper (Cu²⁺) that participates in the catalytic reduction of superoxide to hydrogen peroxide and the subsequent detoxification of hydrogen peroxide by glutathione peroxidase or catalase.

Transition metal cations may bind to the cellular nucleus and locally generate reactive oxygen species, most importantly hydroxyl radicals. This results in DNA strand breaks and modification, accounting for the mutagenic effects of many transition metals. In addition to the important role that transition metal chemistry plays following iron or copper salt poisoning, the long-term consequences of chronic transition metal poisoning are exemplified by asbestos. The iron contained in asbestos is the origin of the Fenton-generated hydroxyl radicals that are responsible for the pulmonary fibrosis and cancers associated with long-term exposure.

The most important toxicologic effects of reactive oxygen species occur on the cell membrane, and are caused by the initiation by hydroxyl radicals of the lipid peroxidative cascade. The alteration of these lipid membranes ultimately causes membrane destruction. Identification of released oxidative products such as malondialdehyde is a common method of assessing lipid peroxidation.

Under normal conditions, there is a delicate balance between the ­formation and immediate detoxification of reactive oxygen species. For example, the conversion of the superoxide radical to hydrogen peroxide via SOD is rapidly followed by the transformation of hydrogen peroxide to water by glutathione peroxidase or catalase. Furthermore, the fact that transition metals exist in “free” form in only minute quantities in biologic systems minimizes the formation of hydroxyl radicals; that is, cells have developed extensive systems by which transition metal ions can be sequestered and rendered harmless. Ferritin (binds iron), ceruloplasmin (binds copper), and metallothionein (binds cadmium) are all specialized proteins that safely sequester transition metal ions. Certain proteins and enzymes such as hemoglobin or SOD have critical biological functions associated with the transition metals at their active sites.

Detoxification of certain reactive species is difficult because of their extreme reactivity. Widespread antioxidant systems typically act to trap reactive species before tissue damage occurs. An example is the availabilityof glutathione, a reducing agent and nucleophile, which prevents both exogenous oxidants from producing hemolysis and the APAP metabolite N-acetyl-p-benzoquinoneimine (NAPQI) from damaging the hepatocyte (Chap. 35).

The key reactive nitrogen species is nitric oxide. At typical physiologic concentrations, this radical is responsible for vascular endothelial relaxation through stimulation of guanylate cyclase. However, during oxidative burst, high concentrations of nitric oxide are formed from L-arginine. At these concentrations, nitric oxide is directly damaging to tissue and also reacts with the superoxide radical to generate the peroxynitrite anion. This is particularly important because peroxynitrite may spontaneously degrade to form the hydroxyl radical. Peroxynitrite ion is implicated in both the delayed neurologic effects of carbon monoxide poisoning and the hepatic injury from APAP.

Redox cycling.

Although transition metals are an important source of reactive species, certain xenobiotics are also capable of independently generating reactive species. Most do so through a process called redox cycling, in which a molecule accepts an electron from a reducing agent and subsequently transfers that electron to oxygen, generating the superoxide radical. At the same time, this second reaction regenerates the parent molecule, which itself can gain another electron and restart the process. The toxicity of paraquat (Chap. 112) is selectively localized to pulmonary endothelial cells. Its pulmonary toxicity results from redox cycling generation of reactive oxygen species (Fig. 112–3). A similar process, localized to the heart, occurs with anthracycline antineoplastics such as doxorubicin.

Acid–Base Chemistry

Water is amphoteric, which means that it can function as either an acid or a base, much the same way as the bicarbonate ion (HCO₃^–). In fact, because of the amphoteric nature of water, H⁺, despite the nomenclature, does not ever actually exist in aqueous solution; rather, it is covalently bound to a molecule of water to form the hydronium ion (H₃O⁺). However, the term H⁺, or proton, is used for convenience.

Even in a neutral solution, a tiny proportion of water is always undergoing ionization to form both H⁺ and OH⁻ in exactly equal amounts. By convention, however, it is the quantity of H⁺ that is measured to determine the acidity or alkalinity of the solution. The units for this measurement are pH, which is –log[H⁺]. In a perfect system at equilibrium, the concentration of H⁺ ions in water is precisely 0.0000001, or 10^–7, moles/L and that of OH^– is the same. The number of H⁺ ions increases when an acid is added to the solution and falls when an alkali is added. The negative log of 10^–7 is 7, and the pH of a neutral aqueous solution is thus 7. In actuality, the pH of water is approximately 6 because of dissolution of ambient carbon dioxide to form carbonic acid (H₂O + CO₂ → H₂CO₃), which ionizes to form H⁺ and bicarbonate (HCO₃⁻).

There are many definitions of acid and base. The three commonly used definitions are those advanced by (1) Svante Arrhenius, (2) BrØnsted and Lowry, and (3) Lewis. Because the focus is on physiologic systems, which are aqueous, the original definition by the Swedish chemist Arrhenius is the most practical. In this view, an acid releases hydrogen ions, or protons (H⁺), in water. Similarly, a base produces hydroxyl ions (OH^–) in water. Thus, hydrogen chloride (HCl), a neutral gas under standard conditions, dissolves in water to liberate H⁺ and is therefore an acid.

For nonaqueous solutions, the Br⊘nsted-Lowry definition is preferable. An acid, in this schema, is a substance that donates a proton and a base is one that accepts a proton. Thus, any molecule that has a hydrogen in the 1+ oxidation state is technically an acid, and any molecule with an unbound pair of valence electrons is a base. Because most of the acids or bases of toxicologic interest have ionizable protons or available electrons, respectively, the Br⊘nsted-Lowry definition is most often considered when discussing acid–base chemistry (ie, HA + H₂O → H₃O⁺ + A⁻; B⁻ + H₂O → HB + OH⁻). However, this is not a defining property of all acids or bases.

The Lewis classification offers the least-restrictive definition of such substances. A Lewis acid is an electron acceptor and a Lewis base is an electron donor.

Simplistically, acids are sour and turn litmus paper red; bases feel slippery (due to the dissolution of skin), are bitter, and turn litmus paper blue. Of course, the tasting and tactile assessment of substances to determine their acid-base status is not recommended.

Because acidity and alkalinity are determined by the number of available H⁺ ions, it is useful to classify chemicals by their effect on the H⁺ concentration. Strong acids ionize almost completely in aqueous solution and very little of the parent compound remains. Thus, 0.001 (or 10^–3) mole of HCl, a strong acid, added to 1 L of water produces a solution with a pH of 3. Weak acids, on the other hand, reach an equilibrium between parent and ionized forms, and thus do not alter the pH to the same degree as a similar quantity of a strong acid. This chemical notation defines the strength or weakness of an acid and should not be confused with the concentration of the acid. Thus, the pH of a dilute strong acid solution may be substantially less than that of a concentrated weak acid (Table 12–2).

The degree of ionization of an acid is determined by the pK_a, or the negative log of the ionization constant, which represents the pH at which an acid is half dissociated in solution. The same relationship applies to the pK_b of an alkali, although pK_b is rarely used and the degree of ionization of bases is also expressed as pK_a (of note, pK_a = 14 − pK_b). The lower the pK_a, the stronger the acid; the higher the pK_a, the stronger the base. The pK of a strong acid is clinically irrelevant because it is nearly fully ionized under all but the most extremely acidic conditions. Importantly, knowledge of the pK_a does not itself denote the strength of an acid or an alkali. To some extent, this quality may be predicted by its chemical structure or reactivity, or it may be obtained through direct measurement or from a reference source.

Because only uncharged compounds cross lipid membranes spontaneously, the pK_a has strong clinical relevance. Salicylic acid, a weak acid with a pK_a of 3, is nonionized in the stomach (pH = 2) and passive absorption occurs (Fig. 9–3). Because it is predominantly in the ionized form (ie, salicylate) in blood, which has a pH of 7.4, little of the ionized blood-borne salicylate passively enters the tissues. However, because in overdose the serum salicylate rises considerably, enough enters the tissue to result in clinical toxicity. Salicylate is the conjugate base of a weak acid (and thus itself a strong base). It diffuses into the mitochondrial intermembrane space (between the inner and outer mitochondrial membrane), where abundant protons exist that have been transported there via the electron transport chain of this organelle (Fig. 13–3). Because salicylate is a strong base, it protonates easily in this environment. In this nonionized form, some of the salicylic acid may pass through the inner mitochondrial membrane, into the mitochondrial matrix, and again establish equilibrium by losing a proton. This is the uncoupling of oxidative phosphorylation, dispersing highly concentrated protons in the intermembrane space that are normally used to generate adenosine triphosphate (Chap. 13). Uncoupling produces a metabolic acidosis, and this shifts the blood equilibrium of salicylate toward the nonionized, protonated form, enabling salicylic acid to cross the blood–brain barrier. Presumably, once in the brain, the salicylate uncouples the metabolic activity of neurons with the subsequent development of cerebral edema. This is the rationale for serum alkalinization in patients with aspirin overdose (Chap. 39).

In a similar manner, alkalinization of the patient’s urine prevents reabsorption by ionization of the urinary salicylate. Salicylate anion excreted into acidic urine may protonate, and the uncharged moiety may then back-diffuse across the renal tubular epithelium and re-enter the body. Alkalinization of the urine “traps” salicylate in the deprotonated form, enhancing excretion. Conversely, because cyclic antidepressants are organic bases, alkalinization of the urine reduces their ionization and actually decreases the urinary elimination of the drug. However, in the management of cyclic antidepressant poisoning, because the other beneficial effects of sodium bicarbonate on the sodium channel outweigh the negative effect on drug elimination, serum alkalinization is recommended (Chap. 71).

The study of carbon-based chemistry and the interaction of inorganic molecules with carbon-containing compounds is called organic chemistry, because the chemistry of living organisms is carbon based. Biochemistry (Chap. 13) is a subdivision of organic chemistry; it is the study of organic chemistry within biologic systems. This section reviews many of the salient points of organic chemistry, focusing on those with the most applicability to medicine and the study of toxicology: nomenclature, bonding, nucleophiles and electrophiles, stereochemistry, and functional groups.

Chemical Properties of Carbon

Carbon, atomic number 6, has a molecular weight of 12.011 g/mol. With few exceptions (notably cyanide ion and carbon monoxide), carbon forms four bonds in stable organic molecules. In organic compounds, carbon is commonly bonded to other carbon atoms, as well as to hydrogen, oxygen, nitrogen, or halides. Under certain circumstances, carbon can be bonded to metals, as is the case with methylmercury.

Nomenclature

The most systematic method to name organic compounds is in accordance with standards adopted by the International Union of Pure and Applied Chemistry (IUPAC); these names are infrequently used, especially for larger molecules, and alternative names are common. The complete details of the IUPAC naming system are beyond the scope of this text and can be reviewed elsewhere (www.iupac.org), but a brief description of the fundamentals of this system is included here.

The carbon backbone of a molecule serves as the basis of its chemical name. Once the carbon backbone has been identified and named, substituents (atoms or groups of atoms that substitute for hydrogen atoms) are identified, named, and numbered. The number refers to the carbon to which the substituent is attached. Some of the common substituents in organic chemistry are −OH (hydroxy), −NH₂ (amino), −Br (bromo), −Cl (chloro), and −F (fluoro). Substituents are then alphabetized and placed as prefixes to the carbon chain.

As an example, consider the molecule 2-bromo-2-chloro-1,1,1-trifluoroethane. The molecule has a two-carbon backbone (ethane), three fluoride atoms on the first carbon, a bromine atom on the second carbon, and a chlorine atom on the second carbon (Fig. 12–4A). A basic understanding of a few simple rules of nomenclature thus allows one to quickly generate the molecular structure of a familiar compound, halothane.

Although the above mentioned rules suffice to name simple structures, they are inadequate to describe many others, such as molecules with complex branching or ring structures. The IUPAC rules for naming compounds such as [1R-(exo,exo)]-3-(benzoyloxy)-8-methyl-8-azabicyclo[3,2,1]octane-2-carboxylic acid methyl ester, for example, are too complex to include here. Fortunately, many compounds with complex chemical names have simpler names for day-to-day use; as an example, this molecule is commonly referred to as cocaine (Fig. 12–4B).

Cocaine is an example of a common or trivial name; one without a systematic basis, but which is generally accepted as an alternative to frequently unwieldy proper chemical names. Common names may refer to the origin of the substance; for example, cocaine is derived from the coca leaf, and wood alcohol (methanol) can be prepared from wood. Alternatively, a common name may refer to the way in which a compound is used; rubbing alcohol is a common name for isopropanol. Common names are often imprecise and may generate some confusion, however, as evidenced by the fact that rubbing alcohol, when commercially marketed, may be either ethanol or isopropanol.

An even less precise system of nomenclature is the use of street names. A street name is a slang term for a drug of abuse, such as “blow” (cocaine), “weed” (marijuana), or “smack” (heroin). The street name ecstasy refers to the stimulant 3,4-methylenedioxymethamphetamine (MDMA), which is most frequently consumed in pill form. It would stand to reason that liquid ecstasy might refer to a solution of MDMA, but street names are not necessarily logical. Instead, liquid ecstasy refers to the drug γ-hydroxybutyrate, a sedative–hypnotic with a completely different pharmacologic and toxicologic profile. Furthermore, there are no standards for the content of ecstasy, and many street pills contain other chemicals.

A final consideration must be given to product names. Product names are trade names under which a given compound might be marketed, and are frequently different from both the chemical name and common name. Thus, the inhalational anesthetic in Fig. 12–4A with the chemical name 2-bromo-2-chloro-1,1,1-trifluoroethane has the common name halothane and the trade name Fluothane.

Bonding in Organic Chemistry

While much of the bonding in inorganic chemistry is ionic, the vast majority of bonding in organic molecules is covalent. Whereas electrons in ionic bonds are described as predominately “belonging” to one atom or another, electrons in covalent bonds are shared between two atoms; this type of bonding occurs when the difference in electronegativity between two atoms is insufficient for one atom to wrest control of an electron from another. Bonds are represented by lines between the atoms: one for a single bond, two for a double bond, and three for a triple bond.

Nucleophiles and Electrophiles

Many organic reactions of toxicologic importance can be described as the reactions of nucleophiles with electrophiles. Nucleophiles (literally, nucleus-loving) are species with increased electron density, frequently in the form of a lone pair of electrons (as in the cases of cyanide ion and carbon monoxide). Nucleophiles, by virtue of their increased electron density, have an affinity for atoms or molecules that are electron deficient; such moieties are called electrophiles (literally, electron-loving). The electron deficiency of electrophiles can be described as absolute or relative. Absolute electron deficiency occurs when an electrophile is charged, as is the case with cations such as Pb²⁺ and Hg²⁺. Relative electron deficiency occurs when one atom or group of atoms shifts electrons away from a second atom, making the second atom relatively electron deficient. This is the case for the neurotoxin 2,5-hexanedione (Fig. 12–5); the electronegative oxygen of the carbon-oxygen double bond pulls electron density away from the second and fifth carbon atoms of this molecule, making these carbon atoms electrophilic.

The reaction of a nucleophile with an electrophile involves the movement of electrons, by forming or breaking bonds. This movement of electrons is frequently denoted by the use of curved arrows, which better demonstrates how the nucleophile and electrophile interact. The interaction of acetylcholinesterase with acetylcholine, organic phosphorus pesticides, and pralidoxime hydrochloride provides an excellent example of the way in which nucleophiles and electrophiles interact, and of how the use of curved arrow notation can lead to better understanding of the reactions involved.

Under normal circumstances, the action of acetylcholine is terminated when the serine residue in the active site of acetylcholinesterase attacks this neurotransmitter, forming a transient serine–acetyl complex and ­liberating choline. This serine–acetyl complex is then rapidly hydrolyzed, producing an acetic acid molecule and regenerating the serine residue for another round of the reaction (Fig. 12–6A). In the presence of an organic phosphorus pesticide, however, this serine residue attacks the electrophilic phosphate atom, forming a stable serine–phosphate bond, which is not hydrolyzed (Fig. 12–6B). The enzyme, thus inactivated, can no longer break down acetylcholine, leading to an increase of this neurotransmitter in the synapse and possibly a cholinergic crisis.

The enzyme can be reactivated, however, by the use of another nucleophile. Pralidoxime hydrochloride (2-PAM) is referred to as a site-directed nucleophile. Because part of its chemical structure (the charged nitrogen atom) is similar to the choline portion of acetylcholine, this antidote is directed to the active site of acetylcholinesterase. Once in position, the nucleophilic oxime moiety (−NOH) that is remote from the positive charge of 2-PAM attacks the electrophilic phosphate moiety. This displaces the serine residue, regenerating the enzyme (Fig. 12–6C). A further discussion of organic phosphorus compound toxicity and the use of 2-PAM can be found in Chap. 113 and Antidotes in Depth: A33, respectively.

A second toxicologically important electrophile is NAPQI (Fig. 12–7). NAPQI is formed when the endogenous detoxification pathways of APAP metabolism (glucuronidation and sulfation) are overwhelmed (Chap. 35). As a result of the electron configuration of NAPQI, the carbon atoms adjacent to the carbonyl carbon (a carbon that is double bonded to an oxygen) are very electrophilic; the sulfur groups of cysteine residues of hepatocyte proteins react with NAPQI to form a characteristic adduct (formed when one compound is added to another), 3-(cystein-S-yl)APAP, in a multistep process. These adducts are released as hepatocytes die, and can be found in the blood of patients with APAP-related liver toxicity. Figure 12–7 diagrams the mechanism of the protein–NAPQI reaction (Chap. 35).

Nucleophiles can be described by their strength, which by convention is related to the rate at which they react with the reference electrophile CH₃I. Of more use in pharmacology and toxicology, however, are the descriptive terms “hard” and “soft.” Although imprecise, the designations hard and soft help to predict, qualitatively, how nucleophiles and electrophiles interact with one another.

Hard species have a charge (or partial charge) that is highly localized; that is, their charge-to-radius ratio is high. Hard nucleophiles are molecules in which the electron density or lone pair is tightly held; fluoride, a small atom that cannot spread its electron density over a large area, is an example. Similarly, hard electrophiles are species in which the positive charge cannot be spread over a large area; ionized calcium, a small ion, is a hard electrophile.

Soft species, on the other hand, are capable of delocalizing their charge over a larger area. In this case the charge-to-mass ratio is low, either because the atom is large or because the charge can be spread over a number of atoms within a given molecule. Sulfur is the prototypical example of a soft nucleophile and the lead ion, Pb²⁺, is a typical soft electrophile.

The utility of this classification lies in the observation that hard nucleophiles tend to react with hard electrophiles, and soft nucleophiles with soft electrophiles. For example, a principal toxicity of fluoride ion poisoning (Chap. 107) is hypocalcemia; this is because the fluoride ions (hard nucleophiles) readily react with calcium ions (hard electrophiles). On the other hand, the soft nucleophile lead is effectively chelated by soft electrophiles such as the sulfur atoms in the chelators dimercaprol (Antidotes in Depth: A25) and succimer (Antidotes in Depth: A26).

Isomerism

Isomerism describes the different ways in which molecules with the same chemical formula (ie, the same number and types of atoms) can be arranged to form different compounds. These different compounds are called isomers. Isomers always have the same chemical formula but differ either in the way that atoms are bonded to each other (constitutional isomers) or in the spatial arrangement of these atoms (geometric isomers or stereoisomers).

Constitutional isomers are conceptually the easiest to understand, because a quick glance shows them to be very different molecules. The chemical formula C₂H₆O, for example, can refer to either dimethyl ether or ethanol (Fig. 12–8). These molecules have very different physical and chemical characteristics, and they have little in common other than the number and type of their atomic constituents.

Stereoisomerism, also referred to as geometric isomerism, refers to the different ways in which atoms of a given molecule, with the same number and types of bonds, might be arranged. The most important type of stereoisomerism in pharmacology and toxicology is the stereochemistry around a stereogenic (sometimes called chiral) carbon.

Consider the two representations of halothane shown in Fig. 12–9. In this figure the straight solid lines and the atoms to which they are bonded exist in the plane of the paper, the solid triangle and the atom to which it is bonded are coming out of the paper, and the dashed triangle and the atom to which it is bonded are receding into the paper. It is clear that, for the molecules in Figs. 12–9A and B, no amount of rotation or manipulation will make these molecules superimposable. They are, therefore, different compounds.

These molecules are enantiomers or optical isomers. They differ only in the way in which their atoms are bonded to a chiral carbon. It is important to define the stereochemical configuration of these two ­molecules, which can be done in one of two ways. In the first classification—the d(+)/l(−) system—molecules are named empirically based on the direction in which they rotate plane-polarized light. Each enantiomer will rotate plane-­polarized light in one direction; the enantiomer that rotates light clockwise (to the right) is referred to as d(+), or dextrorotatory; the l(−), or levorotatory enantiomer rotates plane-polarized light in a counterclockwise fashion (to the left).

Alternatively, enantiomers can be named using an elaborate and formal set of rules known as Cahn-Ingold-Prelog. These rules establish priority for substituents, based primarily on molecular weight, and then use the arrangement of substituents to assign a configuration. To correctly assign configuration in this system, the molecule is rotated into a projection in which the chiral carbon is in the plane of the page, the lowest priority substituent is directly behind the chiral carbon (and therefore behind the plane of the page), and the other three substituents are arranged around the chiral carbon. Figure 12–10 assigns Cahn-Ingold-Prelog priority to the halothane enantiomers of Fig. 12–9, and rearranges the molecules in the appropriate projections.

If the priority of the substituents increases as one moves clockwise (to the right), the enantiomer is R (Latin, rectus = right); if it increases as one moves counterclockwise, the enantiomer is S (Latin, sinister = left). Thus, Fig. 12–10A is the R enantiomer of halothane and Fig. 12–10B is the S enantiomer.

Enantiomers have identical physical properties, such as boiling point, melting point, and solubility in different solvents; they differ from each other in only two significant ways. The first, as mentioned before, is that enantiomers rotate plane-polarized light in opposite directions; this point has no practical toxicologic importance. The second is that enantiomers may interact in very different ways with other chiral structures (such as proteins and other cell receptors), which is of both pharmacologic and toxicologic significance.

The best analogy to explain the toxicologic and pharmacologic importance of stereochemistry is that of the way a hand (analogous to a molecule of xenobiotic) fits into a glove (analogous to the biologic site of activity). Consider the left hand as the S enantiomer and the right hand as the R enantiomer. There are, qualitatively, three different ways in which the hand can fit into (interact with) a glove.

First, if the glove is pliable and relatively fluid (such as a disposable latex glove), it can accept either the left hand or the right hand without ­difficulty; this is the case for halothane, whose R and S enantiomers interact with cell membranes rather than specific receptors and possess equal activity. Second, if a glove is constructed with greater (but imperfect) specificity, one hand will fit well and the other poorly; this is the case for many substances, such as epinephrine and norepinephrine, whose naturally occurring ­levorotatory enantiomers are 10-fold more potent than the synthetic dextrorotatory enantiomers. Finally, a glove can be made with exquisite precision, such that one hand fits perfectly, while the other hand does not fit at all. This is the case for physostigmine, in which the (−) enantiomer is biologically active, whereas the (+) enantiomer is inactive.

Even the above analogy is oversimplified, however, because one enantiomer of a drug can be an agonist, while the other enantiomer is an antagonist. Dobutamine, for example, has one stereogenic carbon and thus two stereoisomers. At the α₁ receptor, l-dobutamine is a potent agonist and d-dobutamine is a potent antagonist. Because dobutamine is marketed as a racemic mixture (a racemic mixture is a 1:1 mixture of enantiomers), however, these effects cancel each other out. Interestingly, at the β₁ receptor, d- and l-dobutamine have unequal agonist effects, with d-dobutamine approximately 10 times more potent than l-dobutamine.

Functional Groups

There is perhaps no concept in organic chemistry as powerful as that of the functional group. Functional groups are atoms or groups of atoms that confer similar reactivity on a molecule; of less importance is the molecule to which it is attached. Alcohols, carboxylic acids, and thiols are functional groups; hydrocarbons are generally not considered functional groups per se, but rather are the structural backbone to which functional groups are attached. The functional groups discussed below are included to illustrate important principles, not because they represent an exhaustive list of important functional groups in toxicology.

Hydrocarbons, as their name implies, consist of only carbon and hydrogen. Alkanes are hydrocarbons that contain no multiple bonds; they may be straight chain, usually designated by the prefix n-butane [by the prefix n- as in n-butane (Fig. 12–11A) or iso- if branched, as in isobutane] (Fig. 12–11B). Alkenes contain carbon–carbon double bonds. Alkynes, which contain carbon–carbon triple bonds, are of limited toxicologic importance. Butane (lighter fluid) is an alkane, and gasoline is a mixture of alkanes.

Hydrocarbons are of toxicologic importance for two reasons: they are widely abused as inhalational drugs for their central nervous system (CNS) depressant effects, and they can cause profound toxicity when aspirated. Although these effects are physiologically disparate, they are readily understood in the context of the chemical characteristics of the hydrocarbon functional group.

Hydrocarbons do not contain polar groups (ie, or groups that introduce full or partial charges into the molecule). As such, they interact readily­ with other nonpolar substances, such as lipids or lipophilic substances. ­Hydrocarbons readily interact with the myelin of the CNS, disrupting normal ion channel function and causing CNS depression. When aspirated, hydrocarbons interact with the fatty acid tail of surfactant, dissolving this protective substance and contributing to acute respiratory distress syndrome (Chap. 108).

Alcohols possess the hydroxyl (−OH) functional group, which adds polarity to the molecule and makes alcohols highly soluble in other polar substances, such as water. For example, ethane gas (CH₃CH₃) has negligible solubility in water, whereas ethanol (CH₃CH₂OH) is miscible, or infinitely soluble, in water. In biologic systems, alcohols are generally CNS depressants, but they can also act as nucleophiles. Ethanol may react with cocaine to form cocaethylene, a longer-acting and more vasoactive substance than cocaine itself (Fig. 12–12; Chap. 78).

Alcohols can be primary, secondary, or tertiary, in which the reference carbon is bonded to one, two, or three carbons in addition to the hydroxyl group. Methanol, in which the reference carbon is bonded to no other ­carbons, is not a primary alcohol per se but shares many of the reactivity patterns of primary alcohols. The difference between primary, secondary, and tertiary structures is important, because although the alcohol functional group imparts many qualities to the molecule, the degree of substitution can affect the chemical reactivity. Primary alcohols can undergo multistep oxidation to form carboxylic acids, whereas secondary alcohols generally undergo one-step metabolism to form ketones, and tertiary alcohols do not readily undergo oxidation. This is a point of significant toxicologic importance, and is discussed in more detail later.

Alcohols can be named in many ways; the most common is to add -ol or -yl alcohol to the appropriate prefix. If the alcohol group is bonded to an interior carbon, the number to which the carbon is bonded precedes the suffix.

Carboxylic acids contain the functional group −COOH. As their name implies, they are (weakly) acidic, and the pK_a of carboxylic acids is generally 4 or 5, depending on the substitution of the molecule. Small-­molecular-weight carboxylic acids are capable of producing an elevated anion gap metabolic acidosis, which is true whether the acids are endogenous or ­exogenous. Examples of endogenous acids are β-hydroxybutyric acid and lactic acid; examples of exogenous acids are formic acid (produced by the metabolism of methanol) and glycolic, glyoxylic, and oxalic acids (produced by the metabolism of ethylene glycol). Carboxylic acids are named by adding -oic acid to the appropriate prefix; the four-carbon straight-chain ­carboxylic acid is thus butanoic acid.

Thiols contain a sulfur atom, which usually functions as a nucleophile. The sulfur atom of N-acetylcysteine can regenerate glutathione reductase and can also react directly with NAPQI to detoxify this electrophile. The sulfur atoms of many chelators, such as dimercaprol and succimer, are nucleophiles that are very effective at chelating electrophiles such as heavy metals. Thiols are generally named by adding the word thiol to the appropriate base. Thus, a two-carbon thiol is ethane thiol.

As noted above, molecules with a given functional group often have more in common with molecules with the same functional group than they have in common with the molecules from which they were derived. The alkanes methane, ethane, and propane are straight-chain hydrocarbons with similar properties. All are gases at room temperature, have almost no solubility in water, and have similar melting and boiling points. When these molecules are substituted with one or more hydroxide functional groups, they become alcohols: examples are methanol, ethanol, ethylene glycol (a glycol is a molecule that contains two alcohol functional groups), the primary alcohol 1-propanol, and the secondary alcohol 2-propanol (isopropanol). Each of these alcohols is a liquid at room temperature, and all are very water soluble. All have boiling points that are markedly different from the alkane from which they were derived, and quite close to each other.

In addition to conferring different physical properties on the molecule, the addition of the alcohol functional group also confers different chemical properties and reactivities. For example, methane, ethane, and propane are virtually incapable of undergoing oxidation in biologic systems. The alcohols formed by the addition of one or more hydroxyl groups, however, are readily oxidized by alcohol dehydrogenase (Fig. 12–13).

As Fig. 12–13 indicates, the oxidation of the primary alcohols methanol and ethanol results in the formation of aldehydes (functional groups in which a carbon atom contains a double bond to oxygen and a single bond to hydrogen), whereas the oxidation of the secondary alcohol isopropanol results in the formation of a ketone (a functional group in which a carbon is double-bonded to an oxygen atom and single-bonded to two separate carbon atoms). Although both aldehydes and ketones contain the carbonyl group (a carbon-oxygen double bond), aldehydes and ketones are distinctly different functional groups, and they have different reactivity patterns. For instance, aldehydes can undergo enzymatic oxidation to carboxylic acids (Figs. 12–13A and B), whereas ketones cannot (Fig. 12–13C).

It is here that recognition of functional groups helps to understand the potential toxicity of an alcohol. Methanol, ethanol, and isopropanol are all alcohols; as such, their toxicity before metabolism is expected to be (and in fact is) quite similar to that of ethanol, producing CNS sedation. Because these toxins are primary and secondary alcohols, all can be metabolized to a carbonyl compound, either an aldehyde or a ketone. Here, however, the functional groups on the molecules have changed; whereas aldehydes can be metabolized to carboxylic acids (which can, in turn, cause an anion gap acidosis), ketones cannot. It is for this reason that methanol and ­ethylene glycol can cause an anion gap acidosis, and isopropanol cannot (Chap. 109).

The concept of functional groups, however useful, has limitations. For example, although both formic acid and oxalic acid are organic acids, they cause different patterns of organ system toxicity. Formic acid is a mitochondrial toxin and exerts effects primarily in areas (such as the retina or basal ganglia) that poorly tolerate an interruption in the energy supplied by oxidative phosphorylation. Conversely, oxalic acid readily precipitates calcium and is toxic to renal tubular cells, which accounts for the hypocalcemia and nephrotoxicity that are characteristic of severe ethylene glycol poisoning. The concept of the functional group is thus an aid to understanding chemical reactivity, but not a substitute for a working knowledge of the toxicokinetic or toxicodynamic effects of xenobiotics in living systems.

• Understanding key principles of inorganic and organic chemistry ­provides insight into the mechanisms by which xenobiotics act.

• The periodic table forms the basis for inorganic chemistry and provides insight into the expected reactivity and to a large extent the clinical effects, of any element.

• A growing understanding of how reactive species are formed and how they interfere with physiologic processes has led to new insights in the pathogenesis and treatment of toxin-mediated diseases.

• Organic chemistry forms the basis of life, and is essential to an understanding of biochemistry and pharmacology.

References