Broadly, the periodic table is divided into metals and nonmetals. Metals, in their pure 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). These elements 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, which may, to some extent, be predicted by their location on the periodic table. For example, the substitution of arsenate for phosphate in the mitochondrial production of adenosine triphosphate (ATP) creates adenosine diphosphate monoarsenate (Chap. 12). 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. Similarly, the existence of an interrelationship between Ca2+ and either Mg2+ or Ba2+ is predictable, although the actual effects are not, that is, 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. Additionally, the physiologic relationship among lithium (Li+), potassium (K+), and sodium (Na+) is consistent with their chemical similarities (all alkali metals in Group IA). However, the clinical similarity between thallium (thallous) ion (Tl+) and K+ is not 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 2 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. 16). Xenobiotics may interfere with the physiologic role of these key electrolytes. Li+ may mimic K+ and enter neurons through K+ channels, following which it serves as a poor substrate for the repolarizing 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 a hydroxide anion (not the hydroxyl radical), while it is a CN− anion in patients poisoned with potassium cyanide (KCN).
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 both in 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 pigments for paints or fireworks. However, the ionic forms, unlike the metallic form, of these elements are typically highly reactive and toxicologically important. Transition elements are chemically defined as elements which 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). This reactivity is used by living organisms in various physiologic catalytic and coordination roles, such as at the active sites of enzymes and in hemoglobin, respectively. Expectedly, the substantial reactivity of these transition metal elements is highly associated with cellular injury caused by several mechanisms, including the generation of reactive oxygen species (Fig. 11–2). For example, manganese ion exposure is implicated in the free radical damage of the basal ganglia causing parkinsonism.
The Fenton and Haber-Weiss reactions, which are the two most important mechanisms to generate hydroxyl radicals, are both mediated by transition metals (TM). Iron (Fe2+) and copper (Cu+) are typical transition 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. 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. 11–2). The likely determinant of the specific toxicologic effects produced by each metal is the tropism for various physiologic systems, enzymes, or microenvironments; thus the 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.
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 raises 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 form, perhaps because of its reduced water solubility. Organic mercurial compounds, such as methylmercury and dimethylmercury, are environmentally formed by anaerobic bacteria containing the methylating agent methylcobalamin, a vitamin B12 analog (Chap. 96).
Another toxicologically important member of the heavy metal group is thallium. Metallic thallium is used in the production of electronic equipment and is itself minimally toxic. Thallium ions, however, have physicochemical properties that most closely mimic potassium ions, allowing them to participate in, and potentially alter, the various physiologic activities related to potassium. This property is clinically used 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. 100).
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. Pb2+ compounds tend to be covalent compounds that do not ionize in water. However, 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. 94).
Although the metalloids (B, Si, Ge, As, Sb, Te, At) share many physical properties with the metals, they are differentiated because of 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.
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. 12 and 88).
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.
In their highly reactive elemental form, which contains 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 intense electronegativity, HF ionizes poorly and is a weak acid. This specific property of HF has important toxicologic implications (Chap. 105).
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, consequential hypoxia may occur, and the expected warning signs may be completely absent (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.
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. 11–3). 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 serious toxicologic properties.
Electronegativity of the common elements. Note that the inert gases are not reactive and thus do not have electronegativity.
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 NaCl, or table salt, in which the electronegativity difference between the elements is 1.9, or greater than the electronegativity of the sodium (see subsequent paragraphs and Fig. 11–3). 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 8 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, 2 moles of particles result. This is because NaCl essentially ionizes fully in water; that is, it produces 1 mole of Na+ (23 g/mol) and 1 mole of Cl− (35.45 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 (Ksp). 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 2 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 2 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 is polar. For example, methane contains a carbon atom that shares its valence electrons with 4 hydrogen atoms, in which there is a small charge separation between the elements (electronegativity [EN] difference = 0.40). Furthermore, 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.
Because the EN differences between hydrogen (EN = 2.20) and oxygen (EN = 3.44) are greater (EN difference = 1.24), the electrons in the HO 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.
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 its own 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 (SO42−) 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, and 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 (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 medical toxicology are profound. For example, the oxidation of ferrous (Fe2+) to ferric (Fe3+) iron within the hemoglobin molecule creates the dysfunctional methemoglobin molecule.
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 to which it is binding.
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 do not contain unpaired electrons, such as hydrogen peroxide (H2O2) and ozone (O3), 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 continually generated as a consequence of endogenous metabolism and there is an efficient system for their control. Under conditions of either excessive endogenous generation or exposure to exogenous reactive species, the physiologic defense against these toxic products is 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 compounded 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 11–1 lists some of the important reactive oxygen and nitrogen species.
Table 11–1. Structure of Important Reactive Species |Favorite Table|Download (.pdf)
Table 11–1. Structure of Important Reactive Species
|Reactive Oxygen Species|
|Superoxide radical||O2·− or O2−|
|Singlet oxygen||[O] or 1O2|
|Reactive Nitrogen Species|
This 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 (Fig. 11–3). In this situation, 4 electrons are added to each molecule of oxygen to form 2 water molecules (O2 + 4H+ + 4e- → 2H2O).
Superoxide is generated within neutrophil and macrophage lysosomes as part of the oxidative burst, a method of eliminating 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-induced hepatotoxicity (Chap. 34), 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. 11–2), 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. 11–2), 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 dismutation (reduction) of superoxide to hydrogen peroxide (SOD was originally called erythrocuprein) and the subsequent detoxification of hydrogen peroxide by glutathione peroxidase or catalase.
Transition metal cations may bind to the cellular nucleus where they locally generate reactive oxygen species, most importantly hydroxyl radical. This results in DNA strand breaks and modification, accounting for the promutagenic 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 consequential toxicologic effects of reactive oxygen species occur on the cell membrane, and are caused by the initiation by hydroxyl radical 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 endogenous 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, in order to minimize the formation of hydroxyl radicals, transition metals exist in "free" form in only minute quantities in biologic systems; 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 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 to trap reactive species before tissue damage occurs. An example is the availability of glutathione, a reducing agent and nucleophile, which prevents both exogenous oxidants from producing hemolysis and the acetaminophen metabolite N-acetyl-p-benzoquinoneimine (NAPQI) from damaging the hepatocyte.
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 has primarily both damaging effects and 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 acetaminophen.
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 is selectively localized to pulmonary endothelial cells. Its pulmonary toxicity results from redox cycling generation of reactive oxygen species (Fig. 115-3). A similar process, localized to the heart, occurs with anthracycline antineoplastic agents such as doxorubicin.
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 neutral solution, a tiny proportion of water is always undergoing ionization to form both H+ and OH− in exactly equal amounts. It is, however, the quantity of H+ that is of concern, and this is the basis of using the pH to characterize a solution. In a perfect system at equilibrium, the concentration of H+ ions in water is precisely 0.0000001, or 10−7, moles per liter 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. In an attempt to make this quantity more practical, the negative log of the H+ concentration is calculated, which defines the pH. Thus, the negative log of 10−7 is 7, and the pH of a neutral aqueous solution is 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 Brønsted-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 Brønsted-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. Thus, Lewis offered 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, whereas bases are slippery and bitter and turn litmus paper blue.
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 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, obtain 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 11–2).
Table 11–2. pH of 0.10 M Solutions of Common Acids and Bases Represents the Strength of the Acid or Base |Favorite Table|Download (.pdf)
Table 11–2. pH of 0.10 M Solutions of Common Acids and Bases Represents the Strength of the Acid or Base
|HCl (hydrochloric acid)||1.1|
|H2SO4 (sulfuric acid)||1.2|
|H2SO3 (sulfurous acid)||1.5|
|H3PO4 (phosphoric acid)||1.5|
|HF (hydrofluoric acid)||2.1|
|CH3CO2H (acetic acid)||2.9|
|H2CO3 (carbonic acid)||3.8|
|H2S (hydrogen sulfide)||4.1|
|HCN (hydrocyanic acid)||5.1|
|NaHCO3 (sodium bicarbonate)||8.3|
|NH4Cl (ammonium chloride)||4.6|
|NaCH3CO2 (sodium acetate)||8.9|
|Na2HPO4 (sodium hydrogen phosphate)||9.3|
|Na2SO3 (sodium sulfite)||9.8|
|NaCN (sodium cyanide)||11.0|
|NH4OH (aqueous ammonia)||11.1|
|Na2CO3 (sodium carbonate)||11.6|
|Na3PO4 (sodium phosphate)||12.0|
|NaOH (sodium hydroxide)||13.0|
The degree of ionization of a weak acid is determined by the pKa, 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 pKb of an alkali, although by convention the pKb is expressed as the pKa (pKa = 14 − pKb). The lower the pKa, the stronger the acid; the converse is true for bases. The pK of a strong acid is clinically irrelevant because it is fully ionized under all but the most extreme acid conditions. Knowledge of the pKa does not itself denote whether a substance is an acid or an alkali. To some extent, this quality may be predicted by its chemical structure or reactivity, or obtained through direct measurement or from a reference source.
Because only uncharged compounds cross lipid membranes spontaneously, the pKa has clinical relevance. Salicylic acid, a weak acid with a pKa of 3, is nonionized in the stomach (pH 2) and passive absorption occurs (Fig. 8-3). 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 have devastating clinical effects. Salicylate, a conjugate base of a weak acid and thus a strong base, equilibrates within the various tissues across the outer mitochondrial membrane. In this intermembrane space (between the inner and outer mitochondrial membrane) abundant protons exist, which are transported there via the electron transport chain of this organelle (Fig. 12-3). 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. The process just described uncouples oxidative phosphorylation, by dispersing the highly concentrated protons in the intermembrane space that are normally used to generate adenosine triphosphate (Chap. 12). Uncoupling in the skeletal muscle, for example, 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. 35).
In a similar manner, alkalinization of the patient's urine prevents reabsorption by ionization of the urinary salicylate. Conversely, because cyclic antidepressants are organic bases, alkalinization of the urine reduces their ionization and actually decreases the drug's urinary elimination. 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 (see Chap. 73).