Kinetic vs. Thermodynamic Variables Governing Protein Tyrosyl Radical Migration and Reduction in Heme Binding Proteins

One point that comes out of these articles [Reeder et al., 2008: (http://www.ncbi.nlm.nih.gov/pubmed/18215735); Reeder et al., 2008b: (http://www.ncbi.nlm.nih.gov/pubmed/18215736)] and that I think is important is that, in view of the fact that protonated, protein-bound perferryl heme is thought to contribute heavily to the damaging effects of acellular myoglobin or hemoglobin and that the concentration of the most damaging, protonated ferryl heme species (which is likely to consist, primarily, of protein-bound perferryl heme with a tyrosyl radical, given that ferryl heme, as in compound II, apparently is protonated under normal circumstances, constitutively) was estimated to be one-10-millionth of the concentrations of the deprotonated species (Reeder et al., 2008b), protein-bound, protonated perferryl heme may be as much as ten million times as damaging as deprotonated perferryl (and ferryl) hemes are. More importantly, the "through-protein" radical transfer reactions and also the capacities of reductants, such as those that concentrate in lysosomes as lysosomotropic amines that exhibit net charges of -3 at pH 7.4 and of +1 at the lysosomal pH range of 4-5, to reduce perferryl heme appear to be governed more by thermodynamic variables than kinetic variables (Reeder et al., 2008b), meaning that the rates of reduction are slow and plateau at concentrations that are relatively low. That seems to me to be a key point that emerges from the articles, and Reeder et al. (2008b) are basically saying that the reductions, by some of these reductants, of tyrosyl radicals on myoglobin are slow and that pH (in this case, acidic pH), which is a thermodynamic variable [Mongan and Case, 2005: (http://www.ncbi.nlm.nih.gov/pubmed/15837173)] that can affect the redox potentials for redox reactions (see Schaefer and Buettner, 2001, cited in past postings), causes a fraction of myoglobin to exist as a species in which perferryl heme is protonated and a tyrosyl radical near to the heme binding site is deprotonated. In essence, some reductants seem to be able to gain "thermodynamically-privileged" access, as reductants, to this species, given that other reductants' charge distribution or intracellular localization may hinder, in pH-dependent manners, their capacities to serve as reductants of this particular fraction of protein-bound perferryl heme. The pH and other thermodynamic variables cause tyrosyl radicals to form and be reduced slowly by reductants, perhaps in a manner that's specific to the pH-dependence of the charge distribution on the reductant. These authors [Al-Ayash and Wilson, 1979: (http://www.pubmedcentral.nih.gov/picrender.fcgi?artid=1186415&blobtype=pdf)(http://www.ncbi.nlm.nih.gov/pubmed/35158)] suggested, similarly, that the failure of ascorbate to reduce an "alkaline isomer" of cytochrome c might have been explainable in terms of thermodynamic variables and not kinetic ones. The within-protein radical transfer reactions that govern the location of a protein radical, at any given time, also tend to be driven by thermodynamics and not kinetics. For example, the thermodynamic "product" of within-protein, radical transfer reactions in equine myoglobin, in one case, was a tyrosyl radical that was closer to the site at which the radical initiated (namely, heme), and the kinetic "transfer product" was an indolyl radical on a tryptophan residue that was farther from the initiation site than the tyrosyl radical was. That's not the scenario one would expect to see if the within-protein radical transfer migration were governed by kinetics. The protein tyrosyl radicals that have been identified in many heme binding proteins are long-lived, and this contrasts with the extremely short half-life of other protein tyrosyl radicals and of free tyrosyl radicals. Additionally, some reductants are inefficient reductants of free tyrosyl radicals because of kinetic factors that make those reactions unfavorable. Anyway, it's an interesting area. Another factor, other than charge distribution per se, could be the presence of amide groups that don't serve as ligands in organometallic "complexation" reactions but that have been shown to be sites at which prokaryotic serine proteases cleave the reductants [Winkelmann et al., 1999: (http://www.bashanfoundation.org/hartmann/hartmannirakense.pdf)(http://www.ncbi.nlm.nih.gov/pubmed/10581690); Zaya et al., 1998: (http://www.ncbi.nlm.nih.gov/pubmed/9734303)]. Given that some reductants are substrates of prokaryotic serine proteases, some of the protein reductions might be lysosomal-protease-specific in some way or involve the kinds of amide-aromatic or amide-tyrosyl-specific interactions that are known to be important in determining the tertiary and quaternary structures of proteins [Toth et al., 2001: (http://www.ncbi.nlm.nih.gov/pubmed/11340654); (http://scholar.google.com/scholar?hl=en&q=%22aromatic+amide%22+OR+%22amide+aromatic%22&as_sdt=2000&as_ylo=&as_vis=0); (http://scholar.google.com/scholar?hl=en&q=amide+aromatic+interactions&as_sdt=2000&as_ylo=&as_vis=0)]. Ayala et al. (2002) [Ayala et al., 2002: (http://pubs.acs.org/doi/abs/10.1021/ja0164327)] found evidence for a sequence-specific interaction between amide groups and tyrosyl radicals, etc., as one might expect, given that tyrosine is, obviously, a major aromatic amino acid (and given that histidine is an aromatic molecule at all pH values but is not classified as being an aromatic amino acid, given that it is basic, aromatic, and carries a net charge) and, along with histidine, tends to be present near to or as part of the heme binding sites of heme-binding proteins (or as an axial ligand of heme) [(http://scholar.google.com/scholar?hl=en&q=amide+interactions+aromatic+tyrosyl+OR+tyrosine&as_sdt=2000&as_ylo=&as_vis=0); (http://scholar.google.com/scholar?hl=en&q=amide+interactions+aromatic+histidyl+OR+histidine&as_sdt=2000&as_ylo=&as_vis=0)].

Acid-Catalyzed Aldol Condensation of Pyruvate to Form Parapyruvate and Enol-Lactone Derivative

This shows a mechanism for the acid-catalyzed dimerization of pyruvate, and the dimer, which is parapyruvate, can undergo a reversible cyclization to a lactone that then tautomerizes to the enol lactone form that's the predominant tautomer. The structures of the lactone, which is supposedly the predominant species that's present initially and forms within 20 minutes or so, and parapyruvate, which predominates subsequently, at least under the conditions used by the authors, are shown here [Montgomery & Webb, 1956: (http://www.jbc.org/cgi/reprint/221/1/359.pdf)(http://www.ncbi.nlm.nih.gov/pubmed/13345826)].

Discussion of Methyl Pyruvate and Alkyl Pyruvates

I should probably not even mention this, given that "some of" these sources are fringe-ish, evidently. But I came across this web site on which people were talking about methyl pyruvate being available as a supplement, and this is not the website but is a link to a google search (http://www.google.com/search?hl=en&num=100&q=%22methyl+pyruvate%22+tablespoon+energy&aq=f&oq=&aqi=). Yeah--that's a teeth-chatterer. It literally causes shivering. Ethyl pyruvate has been much more heavily researched (http://scholar.google.com/scholar?hl=en&q=%22methyl+pyruvate%22+neuron+OR+axon+OR+protect+OR+protective+OR+protection&as_sdt=2000&as_ylo=&as_vis=0), and that'll probably become available eventually. But one issue with either ethyl pyruvate or methyl pyruvate, apart from methyl pyruvate's apparent lability (http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=cv3p0610) and apparent capacity to undergo polymerization in the presence of a metal catalyst, at least, at room temperature (http://scholar.google.com/scholar?hl=en&q=%22methyl+pyruvate%22+polymerises&as_sdt=2000&as_ylo=&as_vis=0), might be that, at least in the case of methyl pyruvate, the compound(s) can serve as substrates, prior to their deesterification, of alanine aminotransferase, lactate dehydrogenase, and presumably other enzymes that metabolize pyruvate or other alpha-keto acids [Jijakli et al., 1996: (http://www.ncbi.nlm.nih.gov/pubmed/8914921)]. The problem I could imagine would be the loss of stereoselectivity in the conversion of O-alkylpyruvate esters to lactate by lactate dehydrogenase (LDH). LDH normally forms only (S)-lactate (L-lactate) from pyruvate, which is achiral (ethyl and methyl pyruvate are also achiral), but there's some evidence that prokaryotic enzymes, such as xylose reductase, can form both (R)-ethyl lactate and (S)-ethyl lactate from ethyl pyruvate [Kratzer and Nidetzky, 2007: (http://www.rsc.org/delivery/_ArticleLinking/ArticleLinking.asp?JournalCode=CC&Year=2007&ManuscriptID=b616475g&Iss=10)]. The other product, formed by the reaction of ethyl or methyl pyruvate with alanine aminotransferase, would be ethylalanine or methylalanine, I think. But would the L- or D-alkylalanines be formed (or both)? It's possible that there wouldn't be any issues with it. D-lactate is metabolized much more slowly than L-lactate is, although "nanomolar" amounts of D-lactate are supposedly formed, under normal circumstances, by the metabolism of methylglyoxal [Khan and Garner, 2007: (http://www.ramcjournal.com/2007/jun07/khan.pdf)]. In this abstract [Kou and Guan, 2008: (http://www.ncbi.nlm.nih.gov/pubmed/18344089)], ethyl pyruvate improved intestinal barrier function and reduced sepsis-associated elevations in plasma D-lactate, which was of prokaryotic origin (produced by colonic microorganisms) [Schoorel et al., 1980: (http://adc.bmj.com/cgi/reprint/55/10/810.pdf)]. That's a separate issue but could be a source of confusion or something. And then methyl pyruvate's deesterification yields methanol, rather than ethanol (for ethyl pyruvate). Aspartame also yields methanol, upon the deesterification of its methyl ester moiety (Jijakli et al., 1996). There's all this research saying that the amounts of methanol derived from aspartame aren't damaging, etc., but I dunno. "Slurping wood alkey (wood alcohol, a.k.a. methanol) dudn't sound too good to meeeeeee. I'm not saying I won't take a nip, because I like the taste of that s#$&, but..." It's probably not a very large amount, but anyway...My main concern would be the potential for racemic products to be formed, but it's possible that the eukaryotic and mammalian enzymes retain their stereoselectivity/enantioselectivity in the utilization of alkylpyruvates as substrates. Here's a search to potentially answer that question (http://scholar.google.com/scholar?hl=en&q=NADH+%22ethyl+pyruvate%22+%22R-lactate%22+OR+%22D-lactate%22+OR+%22%28R%29-lactate%22+OR+%22ethyl+R-lactate%22+OR+%22ethyl+D-lactate%22+OR+%22ethyl+%28R%29-lactate%22&as_sdt=2000&as_ylo=&as_vis=0), but I can't address the question to an adequate degree. Then, I wonder if there could be the acyl-glucuronidation-mediated-conjugation issue, etc. These authors discuss some of their research that had shown ethyl pyruvate to be a substrate/competitive inhibitor of glyoxylase-1 [Santel et al., 2008: (http://www.ncbi.nlm.nih.gov/pmc/articles/PMC2567432/pdf/pone.0003508.pdf)], and, given that D-lactate is a product of glyoxylase enzymes and that GSH is a cofactor and that GSH has sometimes been shown to be depleted, evidently in association with therapeutic effects [Zingarelli, 2004: (http://critcaremed.org/pt/re/ccm/pdfhandler.00003246-200407000-00022.pdf)], by ethyl pyruvate, it follows that some ethyl pyruvate may well be a substrate of glyoxylase-1 and that D-ethyllactate and, by extension, D-lactate, upon the deesterification of D-ethyllactate, might be formed as products of the reaction(s). The metabolism of methanol consumes GSH, though, too, by some relatively specific mechanism and acetaldehyde could, presumably, produce some of that effect. Alkylpyruvates and other acyl esters of nutrients can, incidentally, serve as substrates for a lot of different enzymes with esterolytic activity, such as carbonic anhydrase enzymes, in the case of methyl pyruvate [Pocker et al., 1978: (http://scholar.google.com/scholar?hl=en&q=%22methyl+pyruvate%22+%22carbonic+anhydrase%22&as_sdt=2000&as_ylo=&as_vis=0), etc.]. In any case, these issues may not be issues at all, but it was something to think about.

Proposed Mechanism for the Peroxynitrite-Mediated Oxidation of Uric Acid to Allantoin

God, this "Marvin" can be a real ditty to use and can still take a little time, even though I guess it is easier than the other one for this type of thing. This shows two one-electron oxidations of uric acid by ferryl heme (just because I couldn't think of another oxidant I could show the reactions of, easily) and then a series of two-electron oxidations of the products, yielding allantoin. These are proposed mechanisms that I've tried to piece together. The last reaction is shown as a two-electron oxidation, but it wouldn't really be possible in that step. Oxygen would only have 6 electrons after that. Peroxynitrite can, actually, act as either a one- or two-electron oxidant, but not in that case. I didn't feel like showing another one-electron oxidation, assuming that peroxynitrite can even bind to that carbon. The intermediates are, however, known to be intermediates in the oxidation of uric acid by peroxynitrite.

Psychedelic Mechanism for the Liberation of Nitric Oxide From a Substituted Acetohydroxamate

These articles [Samuni et al., 2009: (http://www.ncbi.nlm.nih.gov/pubmed/19447172); Gantt et al., 2006: (http://www.ncbi.nlm.nih.gov/pubmed/16681389); King, 2004: (http://www.ncbi.nlm.nih.gov/pubmed/15304249)] are relevant to the psychedelic mechanism shown below, in which the oxidation of a substituted acetohydroxamate (a "secondary" hydroxylamine) to its nitrosonium/oxoammonium species allows for the activation of the oxoammonium species toward a "modified Cope elimination," through a hydrogen abstraction. The third step is actually consistent with the mechanism proposed by King et al. (2004) for the formation of nitric oxide from an "unsubstituted" acylhydroxamate (a.k.a. a "primary" or terminal hydroxylamine), and King et al. (2004) argued that only those compounds with a terminal hydroxylamine moiety could serve as nitric oxide or nitrate donors. But Samuni et al. (2009) found that suberoylanilide hydroxamate, a substituted acylhydroxamate (a secondary hydroxylamine) histone deacetylase inhibitor, could serve as a nitric oxide donor, albeit by an unknown mechanism. It's interesting that a major class of histone deacetylase (HDAC) inhibitors are hydroxamate compounds that are thought to inhibit HDAC enzymes by chelating the zinc ions that are coordinated by amino acid residues on or near the active sites of the enzymes. The hydroxamate then "stays" there or something like that. But Gantt et al. (2006) found evidence and cited research to indicate that iron(II), not Zn2+, is required for the catalytic activity of HDAC enzymes. There's a lot of research showing that HDAC inhibitors are neuroprotective, and so a decrease in the availability of nonheme iron(II) to HDAC enzymes could serve an indirect, HDAC inhibitory effect, even if Zn2+ is, in fact, the metal cofactor. Or, either Zn2+ or Fe(II) could serve as the cofactor, as is the case for other enzymes, etc. The activities of HDAC enzymes can be decreased by various antioxidants, and there are "fringe" interactions through which HDAC activity can regulate HIF-1alpha expression, etc. Those are some "far out" things, in any case, that are lacking in mass appeal, probably. This is one of the last obnoxious heme drawings I'll probably be able to come up with, for the time being, as this stuff winds down. I don't know what would happen to that reactive alkene product of the "modified Cope elimination." The actual nitrogen-containing species would probably be something else, or the product I've shown would react to form something else.

Use of the Nernst Equation to Estimate the Spontaneity of a Redox Reaction (or, Here, the Reduction Potential for a "Composite" Half-Reaction)

This is not likely to be of interest to many people, but it's useful as a crude way of assessing the potential for a given reductant to reduce another species (which functions as the oxidant in the forward reaction). I was just realizing that one could also use the Nernst equation to estimate the upper or lower boundary for the reduction potential (an unknown, which is E0(sub1)) of the half-reaction for a reductant, as long as one knows that the overall reaction is occurring (deltaE > 0). In this article [Schafer and Buettner, 2001: (http://www.healthcare.uiowa.edu/corefacilities/esr/education/2003/appendix/FRBM-Redox-2001.pdf)(http://www.ncbi.nlm.nih.gov/pubmed/11368918)], the authors discuss the use of the Nernst equation (see below) to calculate deltaE for the reduction of a given species (Ox2, which in this case is ferryl heme) by a given reductant (in this case, a hydroxamate-based drug). In one article, the authors found that the rate constant was half-maximal at a ratio of reductant to ferryl heme of about 20:1 (200 uM reductant/10 uM ferryl heme), and [Hirst and Goodwin, 2000: (http://www.jbc.org/content/275/12/8582.full.pdf+html)(http://www.ncbi.nlm.nih.gov/pubmed/10722697?dopt=Abstract)] found that, until the H2O2 was consumed, the "initial steady-state" ratio of ferryl to ferric heme, for an in vitro system in which a cytochrome c peroxidase enzyme was functioning catalytically (the enzyme is "cycling" and has a supply of H2O2 or O2), was, at a maximum, ~0.6 (Hirst and Goodwin, 2000). Koppenol and Liebman (1984) [Koppenol and Liebman, 1984: (http://pubs.acs.org/doi/abs/10.1021/j150645a024)] estimated that the redox potential for the ferryl heme/ferric heme ["compound II/metmyoglobin" (Koppenol and Liebman, 1984, p. 100)] redox couple is about +990 mV (+0.99 V). E0(sub1) is the reduction potential for a "composite" of these two reduction reactions:

RNO(rad) (nitroxide) + e(-) ---> RNO(-) (hydroxamate)
RN+=O (nitrosonium) + e(-) ---> RNO(rad) (nitroxide)

That's the convention for writing the half-reactions. It's oxidant + n(electrons donated by the reductant to the oxidant) ---> reductant (Schafer and Buettner, 2001). This would work more effectively if there were only one possible reaction, instead of two, but it's still potentially useful. Some people approximate deltaE and say that deltaE = E0(sub2) - E0(sub1). E0(sub2) = (the reduction potential or redox potential for the compound or species that's being reduced in the forward reaction), and E0(sub1) is the reduction potential for the species being oxidized in the forward reaction shown below. That approximation gives slightly different numbers, but the numbers are sort of similar and can sometimes be similar enough to get a sense of things. The Nernst equation takes into account the stoichiometry (numbers of species taking part in the half reactions and overall reactions), and that takes into account entropy (as in GSSG ---> 2GSH, forming two molecules from one), etc. The Nernst equation also takes into consideration the temperature (in the 59.1 number) and concentrations. High concentrations of one species can sometimes affect the likelihood that deltaE will be greater than 0, meaning that the rxn will be spontaneous (as in E greater than 0 for an electrochemical cell). The less than/greater than symbol (blogger won't let me put those symbols in without confusing the html code or whatever, sometimes) reverses directions because it's necessary to multiply by a negative 1, as in "crayon-based third-grade math." Actually, I have to go the "crayon route" when I do math, to avoid making errors. When one plugs those numbers into the equation and applies the crayon method, one gets this:

Two-Electron Reduction of Nitrosonium Species by Glutathione

This shows a proposed mechanism for the two-electron reduction of a nitrosonium (an oxoammonium cation), derived from an acetohydroxamic acid (shown protonated in the product, despite the unlikelihood of that), to an acetohydroxylamine (a.k.a. an acetohydroxamic acid) by glutathione, to form oxidized glutathione (GSSG). The overall mechanism is similar to the mechanism shown by Arends et al. (2006) [Arends et al., 2006: (http://cat.inist.fr/?aModele=afficheN&cpsidt=17844353)] for the two-electron reduction of an oxoammonium species by an alcohol, in the presence of a base. Boese et al. (1997) [Boese et al., 1997: (http://www.jbc.org/content/272/35/21767.full.pdf+html)(http://www.ncbi.nlm.nih.gov/pubmed/9268306?dopt=Abstract)] and Li et al. (2005) [Li et al., 2005: (http://www.jbc.org/content/280/17/16594.full.pdf+html)(http://www.ncbi.nlm.nih.gov/pubmed/15695823?dopt=Abstract)] also discussed mechanisms for reactions involving nitrosothiols, and those mechanisms are likely to be applicable in this case. There's evidence that the two-electron reduction of oxoammonium species, formed by the one-electron oxidation of nitroxide radicals, is going to be the predominant reduction pathway in vivo, even (and, in fact, particularly) during hypoxia or metabolic stress. In vitro, the one-electron reduction does occur, and the two-electron reduction also may occur.

Citrate and the "Lemonade Approach"

In this article [Kelleher et al., 1996: (http://www.ncbi.nlm.nih.gov/pubmed/8836557)], Kelleher et al. (1996) found that either citrate or fructose-1,6-bisphosphate protected cultured astrocytes and also neurons co-cultured with astrocytes against death due to hypoxia. The idea was that citrate would act as a phosphofructokinase inhibitor and inhibit the excessive glycolytic flux ("hyperglycolysis") that can occur after a traumatic brain injury or in cells under extreme hypoxic stress. That inhibition of phosphofructokinase could actually help explain the alkalinizing effect that citrate supplementation produces [McNaughton et al., 1990: (http://www.ncbi.nlm.nih.gov/pubmed/2079058)], given that glycolytic activity can lead to metabolic acidosis, especially after or during hypoxia. Those authors (McNaughton et al., 1990) found that citrate increased serum bicarbonate in a dose-dependent manner, and that's not necessarily good, past a certain point. That can be harmful in the context of phosphate depletion, etc., in the sense that an increase in serum bicarbonate can reduce oxygen unloading from hemoglobin. Increases in serum bicarbonate also fairly reliably decrease serum ionized calcium by increasing the binding of calcium to albumin (by deprotonating the sites on albumin that normally bind calcium or magnesium, etc.). Kelleher et al. (1996) noted that either fructose-1,6-BP or citrate had been found, by other researchers, to decrease the ionized calcium concentrations in the media of cultured astrocytes or neurons. There's some research showing that phosphocitrate is one "mediator" that can reduce calcium influx into mitochondria, I think, and potassium or sodium or magnesium citrate(s) have been used to prevent kidney stones. Citrate normally helps to prevent calcification in the tubular fluid, and there's a lot of research on the use of magnesium in combination with citrate (to prevent kidney stones, etc.) (http://scholar.google.com/scholar?hl=en&q=magnesium+citrate+nephrolithiasis+OR+urolithiasis&as_ylo=&as_vis=0). It's interesting that Kelleher et al. (1996) found that fluorocitrate, an aconitase inhibitor, also was protective, and the authors suggested that fluorocitrate had increased citrate availability by decreasing its conversion to isocitrate by aconitase. That could lend credence to the calcium-modulating effect or another mechanism. It doesn't necessarily rule out an effect on the TCA cycle, though, even in those experiments, and there's reason to think that citrate could have some usefulness as an energy substrate (or citric acid, which is sold also but that might be too acidic or something and evidently, but maybe not surprisingly, doesn't have the alkalinizing effect that citrate has) [Sakhaee et al., 1992: (http://www.ncbi.nlm.nih.gov/pubmed/1552616)]. There's research showing that 10-15 percent of the ATP production by cells in the proximal tubules comes from citrate oxidation [cited in Mandel, 1985: (http://www.ncbi.nlm.nih.gov/pubmed/3888090)]. I'm thinking that the use of magnesium citrate could be less likely than the use of alpha-ketoglutarate salts to cause GI-related problems, given that at least two, presumably, of citrate's ionizable carboxylate groups would be chelating the magnesium, somewhat. Apparently magnesium citrate's a salt and not a coordination compound but behaves more like a coordination compound [Lindberg et al., 1990: (http://www.ncbi.nlm.nih.gov/pubmed/2407766)]. Those authors found that Mg2+ citrate forms a soluble complex and that the bioavailability of Mg2+ from it is higher than the Mg2+ bioavailability from Mg2+ oxide. I should mention that the amount of citrate in that is large. Mg2+ citrate is about 11.3 percent magnesium. So 400 mg of elemental Mg2+, from Mg2+ citrate, would also provide 3,636 mg (3.636 grams) of citrate. So it's like a citrate supplement without the sodium or lousy calcium or potassium, etc. I haven't done any reading to see what the characteristics of the pharmacology of citrate are going to be. I'd watch for that bicarbonate effect, though, because it's not necessarily going to be rosy for a person in a disease state. But there's also the fact that, as Kelleher et al. (1996) allude to, citrate is not especially abundant. I think that the flux through citrate synthase is normally relatively low, and the reason is that oxaloacetate is in great demand. Citrate is also a major precursor of or substrate used in the formation of cytosolic acetyl-CoA and could be "lipogenic" to some extent, but one could say the same thing about glucose or glutamine or any number of "citrate precursors." It could be more true for citrate, though. There's actually research showing that citrate can help prevent ketoacidosis in people taking C5 ketones [Mochel et al., 2005: (http://www.ncbi.nlm.nih.gov/pubmed/15781190)]. Mochel et al. (2005) found that citrate restored ketogenesis, basically, in people who had been taking triheptanoin, a seven-carbon precursor that's oxidized to form C5 ketones (5-carbons) and acetyl-CoA (2-carbons). It's like an "adjunctive anaplerotic agent" in that context. But those types of lipid-based energy substrates have the potential to be problematic, given that acyl-CoA thioesters tend to accumulate and inhibit all sorts of mitochondrial enzymes, etc., in my opinion.

I see that "lemonade-based" therapies (and cranberry-juice-but-not-cranberry-juice-concentrate-based therapies), as sources of citrate, have been used, instead of the citrate salts, to prevent calcium oxalate kidney stones (http://scholar.google.com/scholar?hl=en&q=nephrolithiasis+cranberry+OR+lemonade&as_ylo=&as_vis=0). Someone should change the name of citrate to something that will be taken seriously in neuroscience research. Seriously. "Well, did you take the @#$%*&ng lemonade for your stroke, like I told you, or did you sip it in a straw and get little piddly, lazy-afternoon, "lazy-lemonade" doses. Did you get that great big keg of sleepy-time lemonade like I @#$---Ah, you know, what I usually tell people is to take a half gallon or don't even take a sip."

(I was going to mention that the side effects associated with alpha-ketoglutarate salts are thought to be a consequence, at least in part, of the net charge of -2 of alpha-ketoglutarate, in the physiological pH range. Mg2+ citrate could conceivably have a net charge of -1, as opposed to the -3 charge of free citrate, much as MgATP(2-) vs. ATP(4-) have different effective net charges. Mg2+ citrate could be absorbed as a "monoanion," to some extent, and then yield the trianionic citrate in the extracellular fluid or whatever, etc.)

"Quartz (Silica) Goodies"

In this article [Flythe et al., 2009: (http://linkinghub.elsevier.com/retrieve/pii/S0272638608016132)], Flythe et al. (2009) discussed a case of a person who developed silica (silicon dioxide) kidney stones after the ingestion of silica as an excipient in supplements. The stones stopped forming after the person discontinued the source of silica. There are actually lots of reports of silica urolithiasis [Ichiyanagi et al., 1998: (http://www.ncbi.nlm.nih.gov/pubmed/9792982); Joekes et al., 1973: (http://www.ncbi.nlm.nih.gov/pmc/articles/PMC1588389/pdf/brmedj01539-0034.pdf)(http://www.ncbi.nlm.nih.gov/pubmed/4350443); (http://scholar.google.com/scholar?q=related:33eEpUxra-UJ:scholar.google.com/&hl=en)]. I don't think this is likely to be a big problem for a person whose kidneys are functioning normally, but it's always possible to say things like that. "A pathological disease mechanism is likely to be more pronounced in a person in a disease state." It wouldn't be very difficult to reduce the silica intake from supplements, if one wanted to. Some supplements just use magnesium stearate as an excipient in capsules or tablets. Or one could just put the silica-containing supplements in a glass of water, let the insoluble precipitate settle (silica and calcium silicate and other silicates are not very soluble in water), and drink the water but not the silica (or pour it through a coffee machine filter--I don't know--I'm not a silica chemist, but the easier thing would just be to take some brand with no silica). From what I can tell, however, the issue is not the solubility of the silica as much as it is the fact that 1-5 percent of some forms of colloidal silica can evidently be absorbed [see last page of Jugdaohsingh et al., 2000: (http://www.ajcn.org/cgi/content/full/71/4/944)(http://www.ncbi.nlm.nih.gov/pubmed/10731501)]. That means that someone can tell a person all day that silica is insoluble, but it doesn't matter. A colloidal suspension contains, by definition, an insoluble substance that's suspended in the water or other medium (it's not called a solvent because the substance isn't being solvated). A colloid is a suspension and isn't a solution. Milk is partly a colloidal substance and is basically a lot of highly-dispersed proteins and other constituents that are not solubilized. An emulsion could be called a colloidal suspension. So if the insoluble portion can be absorbed into the bloodstream or extracellular fluid in the intestinal tract, at least, it doesn't matter, in my opinion, that it can't enter solution. It's beside the point. I looked at a few articles, and it sounds like a real mess, in the sense that one doesn't even know what "form," as in oligomeric polymers or monomers, as discussed by Jugdaohsingh et al. (2000), the silica is being supplied in.

It actually doesn't sound good to me, given that silica, in vitro, can apparently promote collagen "mineralization" (i.e. soft-tissue calcification) [see discussion by Yendt, 1973: (http://www.pubmedcentral.nih.gov/picrender.fcgi?artid=1946599&blobtype=pdf)]. Most of the reports of autoimmune diseases, in association with silica "exposure," have been in people who have inhaled silica dust (it causes a granulomatous disease, basically, in the lungs, by activating the macrophages that try to phagocytose the silica crystals and can't break them down and go wild, releasing pro-inflammatory cytokines, etc.). However, Parks et al. (2002) [Parks et al., 2002: (http://service004.hpc.ncsu.edu/toxicology/websites/courses/TOX705/papers/Parks%20et%20al%202002.pdf)(http://www.ncbi.nlm.nih.gov/pubmed/12124868)] cited research (reference 16) that had shown that silica ingestion could activate macrophages in animals (meaning that the expression and release of pro-inflammatory cytokines would be increased, etc.). It obviously can be absorbed. If it weren't absorbed, one wouldn't find cases of "pure" silica-containing kidney stones in people. Even if the issue were about solubility, the GI tract is not a system at equilibrium. If 0.5 percent or whatever of a given amount of silica is absorbed (because it's soluble) in 60 minutes, say, then 0.5 percent of the remaining amount can be absorbed over the next 60 minutes or 30 minutes or whatever. Just look at the shocking case of inulin, which is still being sold [(http://hardcorephysiologyfun.blogspot.com/2009/01/disturbing-articles-on-inulin.html); (http://hardcorephysiologyfun.blogspot.com/2009/01/note-on-inulin.html)]. Here's another one showing inulin can be absorbed through the colon in rats [Ma et al., 1995: (http://www.ncbi.nlm.nih.gov/pubmed/7806033)], not just the small intestine. I'm not sure what to say, except "that's very serious business," in my opinion. It's problematic, in my opinion, to make safety determinations that are based solely, evidently, or almost solely on the supposed insolubility of a substance or on some flimsy notions about the "inviolability" of the "brush border" of the intestinal tract. I'm joking somewhat, but oh my God. It's a nonspecific immune adjuvant (Parks et al., 2002). I used to think it was essential, too, but I'm thinking that the standard "fact" about silica essentiality (that every 100th carbon, or something similar to that, in glycosaminoglycan chains is a silicon atom and that that's evidence of essentiality) isn't evidence of essentiality. That's just my opinion, but it sounds like it might just be precipitating in the extracellular matrix. Or it might be the case that some monomeric silica is being inappropriately incorporated into some of the GAG polymers? Now that I think about it, the whole thing sounds far-fetched. I doubt silica is all that harmful, but it seems like it's one of those things, like aluminum, that may not be the sole cause of a disease (such as Alzheimer's) but just shouldn't really be in foods, in my opinion. There's research showing that aluminum can exacerbate iron-mediated lipid peroxidation, for example. The FDA should have removed aluminum from the food supply 50 years ago, but that's just one of those FDA-associated horrors.

Reduction of Ferryl Heme by a Nitroxide Radical

This shows a supposed mechanism for the reduction of ferryl heme (a "compound-II-like" species) by a nitroxide radical, thereby forming the oxoammonium cation form of the nitroxide. I was going to use a single bond alongside a dative bond to show the donation of both electrons to a p(sigma)-d(sigma) bonding or antibonding orbital by oxygen, as suggested by Everse (1998) (see yesterday's postings), but the program wouldn't let me do that. I added a partial (delta) symbol to the formal charges instead, to try to convey the uncertainty about the electron configuration.

One Way of Thinking About the Electrons in the (Controversial & Apparent) Sigma Bonding & Antibonding Molecular Orbitals in Ferryl and Perferryl Heme

These types of issues with the molecular orbitals of heme species are not all that relevant to much that's of practical value, but it really helps me to be able to think of the ways the reactions of heme with oxidants and reductants fit into the theoretical framework of bonding...theory. Then there's the fact that inorganic chemistry is interesting, and I'm a first-class nerd in having had an inorganic chemistry book for ten+ years and puttered around with it (without learning much of it in detail). I can explore the chemistry for longer periods of time than I can write for, on a given day, maybe because the chemistry is visual or something and not as much verbal, etc.

This article is interesting [Harcourt et al., 1986: (http://pubs.acs.org/doi/abs/10.1021/ja00278a005)], and Harcourt et al. (1986) argued that the lone pairs of oxygen in carboxylate-coordinated copper(II) complexes can overlap with copper's d(x^2-y^2) orbitals and interact in sigma orbitals. Even though the lone pairs wouldn't really be oriented toward the iron in ferryl or perferryl heme, as they are in the compound discussed by Harcourt et al. (1986), some researchers have found evidence for the formation of sigma bonding interactions between iron's dz2 orbital and oxygen's pz orbital in ferryl heme [Lehnert et al., 2001: (http://www.ncbi.nlm.nih.gov/pubmed/11516278); Decker et al., 2004: (http://pubs.acs.org/doi/abs/10.1021/ja0498033)] and have argued that the biradical (or "diradical") model cannot entirely explain the bonding in the Fe=O moiety of ferryl heme (or, at least, perhaps, in some high-spin, excited states of ferryl or perferryl heme). One way of looking at the sigma molecular orbitals in Fe=O, as modeled by Lehnert et al. (2001), might be to say that the dz^2-p(sigma) antibonding orbital (beta<33>, discussed on the top of column 2 of p. 8287 of Lehnert et al., 2001) would be like a cross between a lone pair and a conventional sigma bond. In the diagram on p. 8289 (Lehnert et al., 2001), the authors show that oxygen has four electrons in the pi bonds (the px-d(xz) and py-d(yz) pi antibonding orbitals I diagrammed in the previous posting), and that's not the same ferryl heme spin state discussed in much of the text (because the d(x^2-y^2) orbital is shown as being unoccupied in the diagram on p. 8289 but is described, in much of the text, as being occupied in the spin-2 state of ferryl heme). But the point is that two of the four "lone-pair" electrons of oxygen could be in a kind of intermediate state between a sigma bond and unshared pairs in different spin states. In the spin 1 state, oxygen donates both electrons to a dz^2-p(sigma) sigma antibonding molecular orbital (the p-contribution comes from the pz atomic orbital of oxygen) that has "53% metal and 21% oxo character corresponding to a strong [sigma] bond" (Lehnert et al., 2001, p. 8287), but that sigma orbital would be oriented away from the iron, largely, and would behave like a lone pair, from a bookkeeping and bonding-theory standpoint. Also in the spin 1 state, shown in the diagram on page 8289, iron would have two nonbonding electrons in the dxy orbital, as discussed by other authors, and oxygen would donate an extra two of its six electrons to the pi molecular orbitals (the px-d(xz) and py-d(yz) orbitals). Those could also be essentially "both" lone-pair electrons and pi bonding electrons, given the complexity of the electron distribution in those orbitals.

Similarly, in the spin 2 state, according to Lehnert et al. (2001), one of those electrons from the d(xy) orbital of iron would be excited into occupying the d(x^2-y^2) orbital, as discussed on p. 8289, and would interact with both the heme nitrogens and the oxygens (both of which would function as donor ligands and create a sigma bonding molecular orbital that's very delocalized and has what amounts to 1/2 of an electron contribution from oxygen and half from the nitrogens) in the d(x^2-y^2)-N(p) orbital, discussed on p. 8287 and shown in figure 8, that's also referred to as beta<32>. There would still be one electron from oxygen in the dz^2-p(sigma) molecular orbital and zero from iron in that orbital (one of iron's four electrons remains in the d(xy) orbital). Part of the reason it's confusing is that different authors use different notation to refer to essentially the same molecular orbitals. In the spin 2 state, one electron or one-half of an electron from oxygen could be in another highly-delocalized, so-called N(p)/p(sigma)-d(sigma) orbital (beta<21>), discussed in column two of p. 8287. Also, Lehnert et al. (2001) use notation in which the p(sigma)-dz^2 orbital (a.k.a. beta<25>) is a sigma bonding molecular orbital and is not the same as the dz^2-p(sigma) orbital, which is beta<33> and is an antibonding molecular orbital. That's basically what Lehnert et al. (2001) are saying.

Yeah, I think the upshot of that article is that four of oxygen's electrons in Fe=O could be like lone pairs that are also donated to the iron (to form sigma bonding and antibonding molecular orbitals that iron doesn't make a contribution to). The other main point made by the authors is that the bonding in Fe=O can be very similar in both the spin 1 and spin 2 states but that one of iron's "sigma" electrons is excited from the d(xy) to the d(x^2-y^2) orbital in the spin 2 state. One of or, from a bookkeeping standpoint, one-half of an oxygen electron is donated to that orbital in the spin 2 state, but, in each of the two spin states, four of oxygen's electrons could behave and be thought of as being "lone pairs" that are in these psychedelic d(sigma)-p(sigma) bonding and antibonding molecular orbitals.

Finally: Leaving the Shorthand Notation for Heme Species Behind--It's Freedom, "Everse-Style"

Finally. This [Everse, 1998: (http://www.ncbi.nlm.nih.gov/pubmed/9626592)] is an article that addresses some of the significant problems with the notation used to represent heme compounds. The notation just doesn't make any sense, in many cases. Everse (1998) also confirms, when viewed alongside the other articles I have, my suspicion that the Fe=O bond has two electrons. Some of my older mechanisms need to be corrected, because I show homolytic cleavage of a pi or d(pi)-p(pi) bond (the "pi double bond") in some instances. The mechanisms are valid, but it's necessary to not show all of the electron transfer occurring at once. But the diagram showing one electron in the pi*(xz) antibonding molecular orbital, in the previous posting, is accurate. There's another electron in the pi*(yz) antibonding molecular orbital, and those are the only electrons in the bond. It's not a double bond, and Everse (1998) refers to it as a biradical pi-bonding interaction. Or one could call it a d(pi)-p(pi) interaction. But the whole thing with [Fe(IV)O]2+ doesn't make any sense, and I think the whole thing with the [Fe(III)-OH(-)]+ is a convention left over from the whole "crystal field theory" concept of dsp3 and d2sp3 orbitals, etc. In that type of theoretical framework, the ligand "O" is treated as if it's "O2-" or something. It doesn't even have that electronic strucure, though, and Everse (1998) reiterates the point that the oxygen in Fe=O has an oxene (atomic oxygen) electron configuration and has six valence electrons. I also finally found a great article that analyzes the electronic structure of a ferric heme species (t-butoxo-iron(III)-heme) [Lehnert et al., 2001: (http://www.ncbi.nlm.nih.gov/pubmed/11516278)], and that article, as other articles have shown, shows that some ferric heme species do contain sigma bonds. But ferryl heme has no sigma bonding electrons in the molecular orbitals that are included in the ferryl (Fe(IV)=O) moiety of heme.

No Electrons in Sigma Bonding or Antibonding Molecular Orbitals for Fe=O in That Spin State of Ferryl Heme: "A 'Single' Bond But No Single Sigma Bond"

Incidentally, I was going to emphasize that there's no sigma bond with iron(IV) in that spin state of ferryl heme. A lot of researchers have compared the Fe=O bond to the triple bond ("formal bond order") that Fe can form with nitrogen atoms of some ligands. It's like a triple bond, with two sets of pi interactions, without the sigma bond, from what I can tell. That's one way of thinking about the more complex reality, in any case, especially given that the Fe=O bond length in ferryl heme is 1.65 angstroms (that's significantly shorter than a usual Fe-O bond). In that article by Shaik et al. (2005), cited in the last posting, even the sulfur of the cysteine residue, as one of the axial ligands, forms a pi bond, designated "pi(sub)S." For a pentacoordinated species of ferric heme, however, those authors do show an unpaired electron in a sigma antibonding molecular orbital. Anyway, part of the reason I think it's helpful to get a sense of the bookkeeping side is that it's almost impossible to tell what the species being referred to, in some articles, actually are, given the use of all of these different types of nonconventional notation and shorthand and so on. But it's interesting to note that there isn't a sigma bond between the oxygen of a water molecular and iron(II), for example (Shaik et al., 2005). There's a "single bond," but it's a pi*(yz or xz, depending on which coordinates one assigns to which "direction") antibonding orbital and not a sigma molecular orbital. I know that no one on the planet cares, but anyway...I guess I'll have to save some more of these "goodies" for tomorrow.

Electron Configuration and Abbreviated Molecular Orbital Diagram of a Low-Spin Species of Ferryl Heme: My First "Bid"

Sketching these things helps me pin down some of these things, for my own, personal leisure and meditation-like benefit. The articles on this type of thing tend to actually be totally inconsistent, and the differences can't be explained in terms of differences in the spin state of the high-spin vs. low-spin states of ferryl heme, for example. According to some articles, two of the four d-electrons of iron(IV) in ferryl heme are nonbonding and are in a dxy orbital [a.k.a. a would-be sigma*(xy) antibonding molecular orbital], but that doesn't make sense to me. Shaik et al. (2005) [Shaik et al., 2005: (http://pubs.acs.org/doi/abs/10.1021/cr030722j)] reported, as other authors have reported, that iron's two other electrons are in the d(x^2-y^2) orbital, which is unusual in the sense that the lobes of that orbital point in between the nitrogens, toward the meso carbons. Usually, that d-orbital is the highest because its four lobes are oriented directly toward the ligands. But the strange thing is that some of these articles make it sound like the double bond in the Fe=O moiety has only two electrons in it [two singly-occupied d(pi)-p(pi) antibonding orbitals, which are pi*(xz) and pi*(yz), as shown below]. That would mean that the oxygen would have four nonbonding electrons, in sp2 or sp3 or pi nonbonding (or antibonding) orbitals. That would make six total and could explain some of these more or less totally confusing uses of random charge symbols, such as the whole Fe(4+)=O(2-). The formal bond order is either 1/2(bonding-antibonding) or 1/2(bonding)-antibonding, and I can't tell because of the way the person phrased it. If that were the case, there would be a -2 bond order? That might explain the crazy "-2" value that keeps showing up, but come now. The authors of other sources (such as the terror-tome that is my advanced inorganic chemistry 12,000-reference book, among others--seriously, it lists 5-10 references per page, in footnotes, and is 1500 pages but is a compact little ditty) have reported that the bond order of the Fe=O "double bond" is greater than 2, but some experimental numbers from another group show the bond orders of the Fe-O bonds, in different intermediates of the catalytic cycles of peroxidase enzymes or CYP450 enzymes, to range from something like 0.7 to 1.34. A lot is evidently not known about "compound II," as shown below (ferryl heme), and compound I (perferryl heme), but I tend to favor the bookkeeping model in which O has, effectively, 6 valence electrons in ferryl heme and perferryl heme and 7 in ferric hemes and protonated ferryl or perferryl heme. In any case, it's possible to understand the reactions without having perfect knowledge, but I generally like to know how many valence electrons are in an atom, etc. I'll try to gradually go through some more spin states and heme species and modify those sketches to show their electron configurations. It's not that complicated. The porphyrin ring's molecular orbital is usually referred to as the a(2u) orbital and has one unpaired electron in ferryl heme. It's interesting the way the nitrogens only contribute two electrons to iron, too. The confusing part is the inconsistent use of notation and shorthand, etc., and it basically means that it's impossible to tell what a lot of the articles' authors are trying to say. It's as if there's the layer of knowledge that allows one to do most of the research very effectively and then this "seedy" underside to the research, as a whole, in which heme species are "up for grabs," like a big market with everyone haggling and using auctioneers' voices...I'm joking. "Give 'em ten-ten-ten-Spin 5/2, gimme 5/2, 5/2, 5/2, geddem 2, give 'em 2, spin 2, spin 2." It's an interesting and challenging area. The mathematical-solution-to-the-Schrodinger-equation type of electron distribution in the d(pi)-p(pi) orbital that I'm showing (a.k.a. the pi*(xy)antibonding molecular orbital) is consistent with a computer-based analysis. I don't feel like citing the article now. There's another molecular orbital that's in the yz plane and is identical to that one. I can't find any information on the type of orbital that the oxygen's lone pair or pairs (in the event, in the latter case (pairs), that the double bond contains two electrons), but I suppose that doesn't matter. Groves and Nemo (1983) [Groves and Nemo, 1983: (http://pubs.acs.org/doi/abs/10.1021/ja00358a009)] show a single lone pair along the z-axis, along with one xz or yz pi*-antibonding orbital containing one electron. I suppose that could mean that a similar situation exists in the other pi*-antibonding orbital that, along with the first, comprises the Fe=O "double bond." But that sounds strange to me. The molecular orbital I'm showing below is the result of the dxz-px overlap shown in the xz-plane, in the diagram above the molecular orbital.

In the above diagram, the t2g "subshell" of d-orbitals usually comprises the dxz, dyz, and dxy orbitals (each of which is also known as a "d(pi) orbital"), but the ligand environment in heme causes the lobes of the d(x^2-y^2) orbital to be oriented, as discussed above, toward the meso carbons. That makes it the lowest-energy orbital. It's usually in the higher-energy eg set. In an unliganded Fe(IV) ion (Fe4+, if it existed) or an Fe2+ or Fe3+ ion, all the d-orbitals are degenerate (at the same energy level), lower than the relative energy level of the t2g set. The formation of any complex, including the most basic hexaquo-iron(III) or hexaquo-iron(II) complexes, causes the first "splitting" of the energy levels of the orbitals, into the eg set and t2g set (or subshells), and then the complex interplay of the interactions of iron(IV) or iton(III) with the nitrogen ligands and various axial substrates or ligands and cysteine residue or tyrosine residue or histidine residue (as one of the two axial ligands) causes further splitting of the energy levels of the individual d-orbitals in the eg and t2g sets. The spin is the sum of the individual spin values for each unpaired electron. If there are 3, then it's 3/2, etc.

This is one of the two molecular orbitals that are visual approximations of mathematical solutions to the dxz-px overlap, shown above. The + and - have no basis in reality and signify the positive phase and negative phase of a waveform, just as a sine wave can be greater than zero or less than zero. The overlap of a lobe with a positive phase and a lobe of an orbital exhibiting a negative phase is meant to signify destructive overlap, leading to the formation of an (supposed, mathematically) antibonding orbital. Constructive overlap occurs when two lobes are in phase. The different solutions to the (nonrelativistic) Schrodinger equation provide a crude prediction of electron density but mainly explain bonding geometries/orientations.

This is the presumed-to-be-accurate case, with each pi* orbital containing two electrons:

Crude Picture of (Part of the) Electron Configuration of the Hydroxo-Ferric Heme Species

These articles [Shaik et al., 2005: (http://pubs.acs.org/doi/abs/10.1021/cr030722j); Silaghi-Dumitrescu, 2008: (http://macroheterocycles.isuct.ru/en/system/files/08MHC_79-81.pdf); Filatov et al., 1999: (http://zernike.eldoc.ub.rug.nl/FILES/root/1999/AngewChemFilatov/1999AngewChemIntEdFilatov.pdf)] are some "fun" articles I've been looking over. I wanted to clarify the electron configuration of the singly-bonded hydroxo-ferric heme species, and the oxygen apparently has, effectively, 7 valence electrons (as opposed to 6 in ferryl and perferryl hemes). In that species [Fe(III)-OH], apparently, three electrons are shared between iron and oxygen. From a bookkeeping standpoint, oxygen has a lone pair and an unpaired electron, but they're all pi electrons and aren't in the sigma*(xy) and sigma*(z^2) antibonding molecular orbitals for the Fe-O species. Two of these three electrons shared by iron and oxygen, in the Fe(III)-OH species, are in a dxz molecular orbital [one of the d(pi)-p(pi) interactions] (Shaik et al., 2005, Fig. 4). The other shared electron is in a pi*(xz) pi-antibonding molecular orbital. The electron configuration of ferryl heme is easier to understand than that, but it's necessary for me to understand what some of these authors are trying to refer to, with regard to the hardcore chemistry, when they "do" try to describe the electron configurations of these species.

Nitroxide Species Derived From Hydroxamate-Based Drugs as Reductants of Ferryl (and Perferryl) Heme, SOD-Mimetics, and "Catalase Activity Augmenters"

This article [Atamna et al., 2000: (http://www.jbc.org/content/275/10/6741.full.pdf+html)(http://www.ncbi.nlm.nih.gov/pubmed/10702229?dopt=Abstract); Krishna et al., 1996: (http://www.jbc.org/content/271/42/26018.full.pdf+html)(http://www.ncbi.nlm.nih.gov/pubmed/8824241)] is useful and shows that nitroxide compounds can exert antioxidant effects by serving as one-electron reductants of ferryl heme to ferric heme, and a key point is that the ferryl-to-ferric step can be rate-limiting in the maintenance of the catalase activity of protein-bound hemes. Drugs that contain hydroxamate groups are oxidized to their nitroxide radical species by their reactions with peroxyl radicals, and then the nitroxide form of the drug could reduce ferryl heme to ferric heme, without generating the superoxide that would usually be formed through the reduction of ferryl heme by hydrogen peroxide, and yield the oxo-ammonium cation form of the hydroxamate-containing drug. The oxo-ammonium cation species generally tend to exert superoxide-dismutase-mimetic effects by (reacting with superoxide and) forming O2, and that regenerates the nitroxide species that is likely to be the key species, in the absence of free iron, for example. But the acceleration of the rate of reduction of ferryl heme to ferric heme, via the reduction of ferryl heme by the nitroxide species, can, evidently, accelerate the two-electron catalase or "hydroperoxidase" activity of the perferryl-to-ferric heme reduction step. The redox cycling, via the reduction of the oxoammonium species by reduced glutathione, for example, of nitroxide-oxoammonium-hydroxylamine species is actually thought to be necessary for the so-called "catalytic" antioxidant mechanisms (as opposed to stoichiometric antioxidant mechanisms) by which nitroxides or hydroxamate-based drugs that are serving as "nitroxide prodrugs" can act. The nitroxide-oxoammonium cycling is "catalytic" because the species aren't consumed or inactivated through the cycling. And the SOD-mimetic effect may also be an important element, particularly since that mechanism can regenerate the nitroxide species. Anyway, it's too late to put up a sketch, but one reduction mechanism (for the reduction of ferryl heme to ferric heme by a nitroxide radical) is very similar to the reduction by H2O2.

Electrophilicity of the Oxene Oxygen of Tetravalent Hemes; Proton-Coupled Electron Transfer Mechanisms in Heme Reduction Reactions

These are some articles about, among other things, proton-coupled electron transport in hydrogen abstraction/hydrogen transfer reactions [Cheng and Li, 2007: (http://pubs.acs.org/doi/abs/10.1021/cr040077w); Sligar et al., 1980: (http://www.pnas.org/content/77/3/1240.full.pdf)(http://www.ncbi.nlm.nih.gov/pubmed/6929480); Guallar et al., 2003: (http://www.ncbi.nlm.nih.gov/pubmed/12771375)(http://www.pubmedcentral.nih.gov/articlerender.fcgi?artid=165819); Stubbe et al., 2003: (http://pubs.acs.org/doi/abs/10.1021/cr020421u); Reece et al., 2006: (http://www.pubmedcentral.nih.gov/articlerender.fcgi?artid=1647304)(http://www.ncbi.nlm.nih.gov/pubmed/16873123); Sutcliffe et al., 2000: (http://www.ncbi.nlm.nih.gov/pubmed/10973049)]. One of the reasons the articles on the cycling of ferryl and perferrryl and ferric heme are confusing is that the iron-oxygen "double" bond is depicted in a manner that makes the polarity of the bond appear to be similar to the polarity of a carbonyl C=O double bond. In fact, the oxygen can be treated as if it has a formal charge of +2, given that the oxygen, in the p(pi)-d(pi) overlap of the oxygen's pi orbitals with the d(xy) and d(yz) orbitals of iron (and also with a mess of empty or half-filled antibonding d-orbitals or hybrid orbitals of iron), only winds up, effectively, with six valence electrons (Sligar et al., 1980). The oxygen has the electron configuration of an oxene oxygen, which is "atomic oxygen." That's the number of electrons that "atomic oxygen" has, in its ground-state electron configuration. The +2 is the formal charge and is calculated [6-(1/2(4))-2] in the same way any other formal charge calculation is calculated [(ground state valence electrons) - 1/2(shared) - (unshared)], even though the reality is that it's a mess of d-orbitals and all of that and is probably 8/3rds or something crazy like that. I think another example of an "oxene" compound is that short-lived, unsaturated ethylene epoxide compound (an unsaturated, three-membered oxirane). I'm not sure of another good example of a true oxene-like compound, but it's really unusual. It means that the Fe=O bond polarity (in ferryl and perferryl hemes) is reversed, in comparison to the polarity of a C=O bond. The oxygen is electrophilic (electron-deficient), and I've drawn it as such in these "electrophilic loser" diagrams. But anyway, that's helpful to me. Many articles describe the iron-oxo moieties of perferryl and ferryl heme as being something like [Por-Fe(4+)=O(2-)]2+, but a better way to abbreviate the structure might be to write something like this: [Por-Fe=O(+)(+)]2+. It's really necessary to know that oxygen has, effectively, six valence electrons.

Anyway, it's interesting that the hydrogen abstraction reactions, such as the one shown below, are thought to be more like proton-coupled/proton-assisted electron transfer reactions than they are like run-of-the-mill homolytic cleavage reactions. I've shown the reduction of ferryl heme by an unsaturated lipid, to form an alkyl radical, and that can then rearrange and react with molecular oxygen and then epoxidize or participate in other reactions that mediate lipid peroxidation. In proton-coupled electron transport, the electron essentially is transferred to the oxidant, which is the electrophilic oxene oxygen on ferryl or perferryl heme, and the hydrogen atom (with its one electron) may not be transferred at all but is transferred "after" the electron is transferred. Sutcliffe et al. (2000) discussed that concept in a more popular forum (the "Trends" journals) and discussed the way that's thought to be the underlying, core mechanism for enzyme catalysis, in the sense that electron tunneling or "quantum tunneling" allows the enzyme to go "through" the activation free-energy barrier through vibration-associated electron tunneling. The electron can be transferred from one residue to another through the protein's single bonds, in effect, or via proton transfers, etc. The math describes or models enzyme catalysis in terms of quantum mechanical vibrational energy of different parts of the enzyme (http://scholar.google.com/scholar?hl=en&q=enzyme+activation+energy+barrier+quantum+mechanical+vibrational&as_ylo=&as_vis=0), but it's as if the enzyme as a whole, instead of just the active site and other discrete sites, participates in the catalysis and "cheats" the activation energy barrier. Here are the initiation reactions for heme-dependent lipid peroxidation [Fe(IV)=O(2+) + LH--->L(e-) + Fe(III)-OH(+)]. I actually botched the product--it's supposed to contain an Fe-O "single bond," but I'm not going to change it:

This is a mechanism, discussed by Reece et al. (2006), by which the perferryl species can undergo reduction to ferric heme by a long-range, "hard-wired" (Reece et al., 2006, p. 1355) proton transfer mechanism. The cytochrome P450 enzyme creates a "water channel" that allows for proton-coupled electron transfer. I've drawn only the deprotonations (fast proton transfers), but the overall reaction is also a two-electron reduction of perferryl heme to ferric heme. There's probably some way to draw it, but I think the idea, as discussed by Guallar et al. (2003) in another proton-coupled-electron-transfer reduction of perferryl heme, is that the proton transfer allows the electrons to be donated through the hydrogen bonds between the carboxylate groups of the alkyl substituents of the heme species (I haven't shown them on the porphyrin ring) and amino acid residues that they form the hydrogen bonds with (I've shown hydrogen bonds, between water molecules, as being dotted lines, but I haven't shown the carboxylate-containing substituents of heme or the hydrogen bonds they can form with amino acid residues on proteins, etc.). So it's like oxygen donates two electrons in its protonation, and two electrons are tunneled through the enzyme and the porphyrin ring to heme. I've drawn the two-electron transfer as if it's coming from the oxygen to the iron, and maybe Reece et al. (2006) was saying that the electrons would be transfered through the water and not the enzyme. That could be the case in the CYP450 case (Reece et al., 2006), because the water molecules are essentially "locked in" as part of the enzyme. The overall idea is that electrons can be transferred through hydrogen bonds and allow for these mind-bending, catalytic mechanisms for electron donation, such as to heme species, in enzymes. The electrons "sweep" up through....No, that's a little joke:

This is the product: