Chemical Education Today
Commentary
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Conceptual Considerations in Molecular Science by Donald T. Sawyer
With retirement comes the opportunity and time to reflect on the basic concepts of molecular science in relation to 40 years of experience in chemical education and research. Such reflections have led me to conclude that there are significant misconceptions within the chemical community, particularly in the undergraduate curriculum and the associated textbooks; for example, two expressions that occur in all textbooks and in much of the reference literature contain several misconceptions: Ce4+ + Fe2+ → Fe3+ + Ce3+ ·O2· + 4H+ + 4e− → 2H2O E ⬚ = +1.23 V vs NHE These include: (a) The existence of charged atomic centers in condensed phases, which prompts the assumption that the bonding in inorganic systems is dominated by electrostatic ion–ion interactions of charged atomic centers, rather than uncharged atomic centers with covalent bonds; (b) Most oxidation–reduction reactions occur via electron transfer, rather than atom transfer; (c) The standard reduction potentials, E ⬚, for multi-electron half-reactions are based on electrochemical electron-transfer measurements, rather than derived from thermochemical data (calorimetry) for atom-transfer reactions, for example, ·O2·(g) + 2H2(g) → 2H2O(l) Two more misconceptions result from the preceding; (d) Dioxygen, ·O2·, reacts as an electron-transfer oxidant and in its interactions with transition metals, rather than as a biradical and atom-transfer reagent; (e) The interaction of reduced iron, Fe(II), with hydrogen peroxide, HOOH, known as Fenton chemistry, produces free hydroxyl radical, HO·, via electron-transfer reduction of HOOH, rather than a nonradical reactive intermediate via nucleophilic addition by HOOH. Each of these misconceptions and suggested alternatives is summarized in the following paragraphs and discussed in detail in the Supplemental Material.W The Absence of Charged Atomic Centers and Ionic Bonds in the Solution Equilibria of Metal Complexes; Some Alternate Views The metal ions in solution equilibria, usually depicted as atomic cations (e.g., Ag+ and Na+), are atoms with their valence electrons coupled via radical–radical coupling to form covalent bonds with electron-deficient solvent molecules, for example, Ag· + (·OH2)+(OH2)5 → AgI(OH2)6+ ᎑∆GBF = 441 kJ mol−1 Na· + (·OH2)+(OH2)5 → NaI(OH2)6+ ᎑∆GBF = 631 kJ mol−1 The covalent bond(s) is delocalized among the six OH2 groups as is the positive charge among the 12H’s. Thus, a www.JCE.DivCHED.org
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more accurate conceptualization of hydrated metal ions and their reaction with anions, bases, and ligands via nucleophilic displacement (covalent-bond formation and covalent-bond breaking) provides a reasoned rationalization for experimental observation. Thermodynamic data in combination with the bond-formation energies for AgI–Cl(s) (᎑∆GBF = 502 kJ mol᎑1) and NaI–Cl(s) (᎑∆GBF = 611 kJ mol᎑1) provide a rational basis for the observed reaction of chloride ion via nucleophilic substitution with hydrated silver ion (exoergonic by 61 kJ mol᎑1), and its nonreactivity with hydrated sodium ion (endoergonic by 20 kJ mol᎑1). Similar conceptual arguments provide reasonable explanations to the questions: Given that the stability constant for HgII(EDTA)2− (log K = 22.1; ionic radius, 1.1 Å) is 11 orders of magnitude larger than that for CaII(EDTA)2− (log K = 10.7; ionic radius, 1.0 Å), why does the addition of sulfide ion to their solutions precipitate HgIIS (s), but is unreactive with CaII(EDTA)2−? Given that the radius of Al3+ is 0.5 Å and Fe3+ is 0.6 Å, why is the stability constant for FeIII(EDTA) (log K = 25.1) 9 orders of magnitude larger than that for AlIII(EDTA) (log K = 16.1)?
In contrast, these experimental observations are a major dilemma for the concept of ionic bonds in metal complexes (see the Supplemental MaterialW). Most Oxidation–Reduction Reactions Occur via Atom Transfer [Often H· or ·O·]; Not Electron Transfer The ionic-solution-equilibria discussions of first-year chemistry courses prompt most to view metal–oxygen and metal–halogen interactions as ionic rather than covalent (via radical–radical coupling). Such misconceptions give us a poor basis to understand molecular bonding and structure, and reaction mechanisms. Precipitation equilibria imply electrostatic atomic-ion interactions Ag+(aq) + Cl−(aq) AgCl(s) Fe3+(aq) + 3HO−(aq) Fe(OH)3(s) and redox equilibria infer direct electron transfer between charged atomic centers Ce4+ + Fe2+ Fe3+ + Ce3+ The discussion associated with such equilibria leads most to assume that the bonding in inorganic systems is dominated by electrostatic ion–ion interactions of charged atomic centers. However, most so-called “outer-sphere electron-transfer” reactions are atom-transfer (often H· or ·O·) or group-transfer processes, especially in biology; for example, H-atom transfer:
[(H2O)73+]CeIVOH + [(H2O)52+]FeII(OH2) [(H2O)73+]CeIII(OH2) + [(H2O)52+]FeIII–OH
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Fe Cl2 + (H3N)4Cu Cl2 Fe Cl3 + (H3N)4Cu –Cl Note that the Roman numeral superscripts associated with the uncharged metal-atom centers in the formulas for their compounds and complexes indicate their covalence (number of covalent bonds), not their charge or oxidation state. Electrochemical oxidation–reduction processes for metals, metal–solvates, and metal complexes invariably are ligandcentered (see the Supplemental MaterialW). Most Multi-Electron Standard Reduction Potentials Are Derived from Combustion Calorimetry for Atom-Transfer Reactions; Not Electrochemical Electron-Transfer Measurements Because most have assumed that the standard reduction potentials for multi-electron half-reactions are based on electrochemical measurements for electron-transfer processes, the contemporary belief that most “redox” processes of industry and biology involve electron transfer is not surprising. In fact, most of these parameters are based on calorimetric measurements of atom-transfer processes. For example, ·O2· + 4H+ + 4e− → 2H2O E ⬚ = +1.23 V vs NHE is based on the calorimetry for an atom-transfer reaction ·O2· + 2H2 → 2H2O + Q Q ⬚ = ᎑∆H ⬚ ᎑∆G⬚ = ᎑∆H ⬚+ T∆S ⬚ = nE ⬚F = (4 × 96,494)E ⬚ The E ⬚ value is a useful thermodynamic parameter, but has prompted all of us to believe that (a) ·O2· can be transformed to H2O via electron transfer (electrochemical reduction), (b) the activity of ·O2· can be sensed by potentiometry, (c) the air cathode of fuel cells can deliver the potential and energy implied by the relation, and (d) the oxidative metabolism of glucose in biology is a concerted multi-electron-transfer process. A proper electrochemical measurement of an ·O2·兾H+ electrode at standard state conditions yields an E ⬚ value for a one-electron process ·O2· + H+ + e− → HOO· E ⬚ = ᎑0.05 V vs NHE Similar contrasts exist for most multi-electron half-reactions. Because several steps are multi-electron in the so-called “electron-transfer” processes of biology, the observed energies must be due to atom-transfer processes, for example, the enzymecatalyzed “cold combustion” of glucose C6H12O6 + 6·O2· → 6 CO2 + 6H2O + Q Q⬚ = 2799 kJ mol᎑1 (see the Supplemental MaterialW).
gen (·O2·, S = 3兾2) by formation of covalent bonds via radical–radical coupling, for example, ·CoIILx + ·O2· → [LxCoIII–OO·] These adducts undergo further reaction and rearrangement to give a series of reactive intermediates, for example, [LxCoIII⫺OO⫺CoIIILx] + 2H2O → 2LxCoIIIOH + HOOH ·CoIILx + HOOH + B → [Lx−Co·II⫺OOH(BH+)] II 2Fe Lx + ·O2· → [LxFeIII⫺OO⫺FeIIILx] → 2 [LxFeIV⫽O] The intermediates have enhanced reactivity for C⫺H bond breakage (form stronger XO⫺H bonds than free ·O2·) and for oxygenation of unsaturated centers (have weaker X⫽O bonds than ·O2·). Thus, the function of catalysts is to lower energy barriers via formation of reactive intermediates with weakened reactant bonds. There is no evidence that transition metals activate ·O2· via electron transfer from metal centers (see the Supplemental MaterialW). Fenton Reagents (FeIILx/HOOH) Produce a Reactive Intermediate with a Covalently Bound Hydroxyl Group; Not Reactive Quantities of Free Hydroxyl Radical (HO·) There is a general belief that Fenton reagents produce free hydroxyl radical (HO·), but they are unreactive with methane while free HO· reacts (k = 2 × 108 M᎑1 s᎑1). In reactions with hydrocarbons, free HO· produces carbon radicals as the primary product, while Fenton reagents produce molecular products (mainly oxygenated) and no detectable carbon radicals. The one-to-one combination of hydrogen peroxide and Fe II (PA) 2 (PAH is picolinic acid; 2pyridinecarboxylic acid), a Fenton reagent, forms an adduct [(PA)2−]FeIIOOH(solH+) that is the primary reactant with (a) excess FeII(PA)2 to give two (PA)2FeIIIOH, (b) excess HOOH to give ·O2· plus two H2O, and (c) excess cyclohexane (c-C6H12) to give (c-C6H11)pyl (or c-C6H11OH in the absence of pyridine). When dioxygen is present with excess c-C6H12, the sole product is cyclohexanone, c-C6H10(O). The reactivity with hydrocarbons and their product profiles for Fenton reagents formed by several transition-metal complexes in combination with hydrogen peroxide or t-butyl hydroperoxide has been evaluated in the absence and presence of dioxygen. The variation of products, yields, and kinetic isotope effects for the different complexes, peroxides, solvents, and the presence or absence of dioxygen establish that the reactive intermediate is not a single oxidant, such as free hydroxyl radical (HO·) (see the Supplemental MaterialW). Acknowledgment This work has been supported by the Provost’s Office of Texas A&M University.
Binding and Activation of Dioxygen (·O2·) by Transition Metals [e.g., Iron(II), Cobalt(II), Copper(I)] Occurs Via Radical–Radical Coupling; Not via Electron Transfer
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The reduced forms of transition-metal complexes, for example, FeIILx (d6sp, S = 4兾2), CuILx (d10s, S = 1兾2), ·CoIILx, (d7sp, S = 1兾2), and MnIILx (d5sp, S = 5兾2) activate dioxy-
Donald Sawyer is Distinguished Professor Emeritus, Department of Chemistry, Texas A & M University. He is now retired and resides in Lexington, KY;
[email protected].
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Material Each of these misconceptions and suggested alternatives is summarized in this issue of JCE Online.
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