Oxygen Is the High-Energy Molecule Powering Complex Multicellular Life: Fundamental Corrections to Traditional Bioenergetics

A fundamental re-assessment of the overall energetics of biochemical electron transfer chains and cycles is presented, highlighting the crucial role of the highest-energy molecule involved, O2. The chemical energy utilized by most complex multicellular organisms is not predominantly stored in glucose or fat, but rather in O2 with its relatively weak (i.e., high-energy) double bond. Accordingly, reactions of O2 with organic molecules are highly exergonic, while other reactions of glucose, fat, NAD(P)H, or ubiquinol (QH2) are not, as demonstrated in anaerobic respiration with its meager energy output. The notion that “reduced molecules” such as alkanes or fatty acids are energy-rich is shown to be incorrect; they only unlock the energy of more O2, compared to O-containing molecules of similar mass. Glucose contains a moderate amount of chemical energy per bond (<20% compared to O2), as confirmed by the relatively small energy output in glycolysis and the Krebs cycle converting glucose to CO2 and NADH. Only in the “terminal” aerobic respiration reaction with O2 does a large free energy change occur due to the release of oxygen’s stored chemical energy. The actual reaction of O2 in complex IV of the inner mitochondrial membrane does not even involve any organic fuel molecule and yet releases >1 MJ when 6 mol of O2 reacts. The traditional presentation that relegated O2 to the role of a low-energy terminal acceptor for depleted electrons has not explained these salient observations and must be abandoned. Its central notion that electrons release energy because they move from a high-energy donor to a low-energy acceptor is demonstrably false. The energies of (at least) two donor and two acceptor species come into play, and the low “terminal” negative reduction potential in aerobic respiration can be attributed to the unusually high energy of O2, the crucial reactant. This is confirmed by comparison with the corresponding half-reaction without O2, which is endergonic. In addition, the electrons are mostly not accepted by oxygen but by hydrogen. Redox energy transfer and release diagrams are introduced to provide a superior representation of the energetics of the various species in coupled half-reactions. Electron transport by movement of reduced molecules in the electron transfer chain is shown to run counter to the energy flow, which is carried by oxidized species. O2, rather than glucose, NAD(P)H, or ATP, is the molecule that provides the most energy to animals and plants and is crucial for sustaining large complex life forms. The analysis also highlights a significant discrepancy in the proposed energetics of reactions of aerobic respiration, which should be re-evaluated.


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value of -350 kJ/mol for C-C or C-O bonds, 2,3 which are the most common) is taken as an average value for the molar enthalpy of electron-pair bonds in the fuel, excluding C-H and O-H bonds. Hydrogen balance requires n CH + n OH = 2 n H 2 O (S2) which can be rewritten as The unchanged number of electron-pair bonds before and after the reaction, assuming a bond order of two for O 2 , gave us 1 n CH + n OH + n no H + 2 n O 2 = 2 n H 2 O + 4 n CO 2 (S4) If we assumed that O 2 has only one bond but gives rise to an additional bond in the products, the equivalent relation, with the 'extra-bond correction' -n O2 on the side of the products, n CH + n OH + n no H + n O 2 = 2 n H 2 O + 4 n CO 2 -n O 2 (S5) would be obtained. Generally, relation (S4) between integers characteristic of the molecular structures remains valid independent of the bond order assigned to O 2 .
According to eq.(S2), n CH + n OH and 2 n H 2 O in eq.(S4) cancel and we retain 2 n CO 2 = n no H /2 + n O 2 (S6) After regrouping terms in eq.(S1) and using eq.(S6) one obtains:  4 ) for the s-bond in O2. While the triplet p-electron system of O2 is unusual and complicated, it contributes more strongly to the bonding in O2 than does the s-bond. 4 According to Pauling, the p-electrons give rise to two three-electron bonds, whose bond energy is generally "about one-half that of a single bond". 6 With two 'half p-bonds' and a single s-bond, representing O2 with a double bond seems as justified as any simple alternative. A double bond also simplifies a generalized bond-energy analysis, which is easiest if the number of bonds does not change during the reaction. 1 Borden et al. have deduced that the bonding by the two three-electron p bonds is further stabilized by resonance between them, which contributes to the kinetic stability (persistence) of O2 in the atmosphere. 4 S5   Table S1. Enthalpies and Gibbs free energies of simple molecules and ions relevant in biochemical energetics, in kJ/mol. Left two columns: traditional enthalpies and Gibbs energies of formation, which are not energies of just the species named but also depend on the bond energies of the elements in their standard states, and in the case of ions on the hydration energy of z protons, where z is the ion charge number. Right two columns: Natural individual enthalpies and Gibbs energies, dominated by bond energies of the molecule named, and/or ionization energy in water in the case of an ion.  "bond" = bond energy; "cond." = -DvapH o "IE" = (total) ionization energy; "hydr." = hydration/solvation energy; "EA" = electron affinity "IE+hydr." = ionization energy in water "cohesive" = lattice cohesive energy = crystal bonding energy = energy released when the atoms come together to form the crystal = -DsublH o = -DfH o of the gaseous atom.
Note that in the DrG o" values given in the text, typically only three figures are significant.
Origin of data for Table S1 and Figure 1. Some data were given in ref. 7 , while others were calculated as follows, from widely tabulated free energies of formation 8 based on the formation reaction:  Table S1) Fe 2+ = 368 kJ/mol and G o Fe 3+ = 856 kJ/mol, different by 413 kJ/mol + 74.2 kJ/mol.

Evaluating free energies of reactants and products.
In the main text, the calculation of D rG o from free energies of reactants and products was mentioned as the first (i) of two approaches. This can be performed using conventional free energies of formation or meaningful individual free-energy values. We demonstrate the calculation for the reaction O2 + 4 H + (aq) + 4 Fe 2+ (aq) ® 2 H2O + 4 Fe 3+ (aq).
For conceptual simplicity, we consider ferrous and ferric ions in water, rather than in cytochrome c.

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(ib) The use of meaningful energies relative to the free atoms, G o prod,i, i.e. bond energies or ionization energies in water, provides more insight. Here, the calculation is The small difference in the energies of electron-pair bonds in reactants and products immediately shows that the reaction will be only mildly exothermic or exergonic.
The small entropy of combustion. Our previous quantitative analysis 1 focused on DcH o , the enthalpy of combustion, which equals the heat evolved in an uncontrolled reaction with O2 at constant T and P. In respiration, we are interested in the free energy change of reactions with O2, which can be used to pump protons and eventually generate ATP. To evaluate DcG o = DcH o -T DcS o we need for a general combustion of organic molecules S10 CcOoHhNn + nO 2 O2 ® nH 2 O H2O(l) + nCO 2 CO2 + nN 2 N2 Since both reactants and products contain similar numbers of linear gas molecules, there is significant cancelation. The entropic contribution to the free energy of combustion at the standard temperature of Since the last term partially cancels the second, this is usually a fairly small correction to 1 Reduced molecules". A common misconception attributes high energy to "reduced molecules", 11 see also ref. 12 , p.529 and p. 662, top. To start with, the terminology is confused. The total oxidation number of an uncharged molecule is zero and therefore, unlike an ion, a molecule is rarely reduced. What changes is the oxidation state of carbon, so the terminology "reduced molecules" actually means "reduced-carbon molecules". This 'concept' is often closely tied to an electron-transfer picture that ignores that oxidation numbers of covalently bonded atoms are mostly fictitious. For instance, in H2O the O atom does not have net +2 electrons, but only has an effective atomic charge of -0.66 (see Figure S4 below).
The supposed higher energy of reduced-carbon molecules is really the higher energy of the larger number of O2 molecules used in their combustion: Reduced-carbon molecules usually contain little O; we have pointed out and shown quantitatively that molecules containing oxygen atoms have a lower heat of combustion per mass since they pull less of the energy-providing O2 into the reaction. 1 This misconception is closely related to the faulty assumption that the heat of combustion derives from the fuel. In this context, it is useful to realize that CH4 has the same total bond energy as CO2 (actually, CH4 is slightly lower in energy (more stable) relative to the free atoms). 1 The claim that reduced molecules are high-energy molecules has a further flaw: Without O2, partially oxidized molecules like glucose (carbohydrates) can actually release more energy than S11 hydrocarbons as shown in the main text. Anaerobic respiration releases energy from less reduced glucose, but not from nearly fully reduced fatty acids.

Reaction energetics of reduced molecules.
In the main text, reactions of CH4 were analyzed and shown to be endergonic unless O2 (or another strong oxidant) is involved. As an example of a slightly larger reduced molecule, we consider n-hexane, for which the needed thermodynamic data can be found fairly easily: Most of these reactions are endergonic, confirming that reduced molecules are low-energy species. Only in combustion do they release the energy stored in the oxidant, i.e. in O2, particularly efficiently. Consider, for instance, CO2 + 2 (NADPH + H + ) ® CH2O + H2O + NADP +

NAD(P)H + H + is analogous to
The simplified reaction is CO2 + 2 "H2"NAD(P)H ® CH2O + H2O which has an only slightly (17 kJ/mol) higher DrG o than the reaction with real H2.

O2 in the last or first step of respiration?
In the conventional presentation, the reaction with O2 is usually shown as the last, and correspondingly O2 is termed the terminal electron acceptor. Based on Figure S1, one can argue that the reaction with O2 is not the last but the first in the chain. A reaction naturally proceeds from high-energy reactants to lower-energy products, so the order cannot be switched. As shown in Figure S1b, even if steps are left out, O2 is still among the high-energy reactants, which must be present in the beginning for the series of reactions to start. That most of the energy released in respiration can be attributed to O2 becomes clear when comparing with the corresponding reaction without O2 (and forming H2 instead of H2O), see Figure S1c. It can be argued that in steady state, even the high-energy species such as Fe 3+ and Q are present in significant concentration. Nevertheless, in such a steady state a "first" or a "terminal" reaction cannot be uniquely identified, and one can still choose the reaction with O2 as the first in the energy-transfer chain.
Presenting the electron transfer chain of aerobic respiration correctly. The relevance of the analysis of the electron transfer chain in this paper is not based on every step occurring as claimed, or with an exact free-energy change. Rather, the goal is to show how a series of electron-transfer processes can be described in a proper redox energy transfer and release analysis. If details of the electron-transfer process are modified, the energy analysis can be adjusted accordingly. Some of the species in the "electron transport" chains are separated by only nanometers and certain oxidation states are extremely short-lived. Their analysis, like that of radicals in combustion, does not shed light on the bioenergetics framework, which is the focus of this analysis. The interest here is in the overall energetics, not individual reaction mechanisms.
The chemical equations reviewed in the main text show that electrons are transferred through the following species: NADH à QH2 à 2 Fe 2+ (cytc) à H2O S13 Note that all of these are donor (i.e. reduced) species. These species ferry the electrons to the next reaction site. Some representations of electron transport chains incorrectly show some acceptors, in particular in the last step where O2 is involved (furthermore without its required reaction partner, H + , which is the main electron acceptor, see the main text and Figure S4 below). 12 Why should there be no acceptors in the electron-transfer chain? An acceptor is ready to accept electrons it does not have. As soon as it receives the electron, it stops being an acceptor, having become a donor for the next step. So an acceptor never contains the transferred electron and must not be shown in a simple electron transport chain. An expanded representation that shows all species involved, in particular O2, could be constructed as follows: NADH / NAD + + H + (aq) à Q + 2 H + (aq) / QH2 / Q + 2 H + (aq) à 2 Fe 3+ (cytc) / 2 Fe 2+ (cytc) / 2 Fe 3+ (cytc) à ½ O2 + 2 H + (aq) / H2O Bold font labels the species containing the transferred electron. Note that just showing donor/acceptor pairs, e.g. à QH2 / Q àdoes not properly represent the process, which includes a reaction with Q + 2 H + (aq) as products and a subsequent one with Q + 2 H + (aq) as reactants.
Finally, it has to be acknowledged that it may also be misleading to refer to H2O simply as an electron donor. H2O does not donate an electron to become H2O + . While for Fe 2+ (aq) ® Fe 3+ (aq) + e -, we can indeed consider just the electron-release or ionization energy in water, for a molecule like H2O that splits up and recombines into ions and a species with very different bonding, O=O, bond-energy changes are more crucial energetically than is electron release.

Different views of energy release in the electron transfer chain of aerobic respiration. The following
Figures S1-S3 show different representation of the sequence of reactions in aerobic respiration: The "electron waterfall" in Figure S2, which does not show the energies of reactants and products, and a true energy-level diagram that does, Figure S1, but is hampered by complexity. The redox energy transfer and release (RETAR) diagram shown in Figure 2c of the main text is superior in showing the origin of chemical energy in these reactions. Potential variants of the RETAR diagram are explained in Figure S3. "Half-reaction free-energy levels". In a correct free-energy diagram of a series of reactions as shown in Figure S1, it is inconvenient that reactants converted in later steps must be included in the analysis from the start. This is avoided in an alternative analysis in terms of half reactions, which is convenient for displaying a series of reactions, since the same half reaction (in reverse direction) can be used in subsequent reactions, see Figure S2. This is the currently accepted representation of the energetics of electron transfer chains. However, as pointed out in the main text, the levels do not correspond to energies of specific species and therefore do not reveal where chemical energy is stored.
In the main text it was pointed out that acceptors of high energy lower their half-reaction energy level. To make this point graphically, in Figure S1b the level for F2 + 2 e -® 2 Fis also shown; it is lower than all others. F2 is undoubtedly a high-energy molecule due to its weak bond (bond formation enthalpy of -155 kJ/mol). This confirms that "free-energy levels" associated with high-energy oxidants show up low in this plot of DredG o .

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These equations are related by DredG o i = -ne F E o red,i + ne c. The constant c, which has the value of -413 kJ/mol when meaningful bond energies and ionization free energies in water are used, 7,15 is always subtracted out in measurable quantities such as the free energies of full redox reactions (see. eq(17a)) and is therefore of little consequence.   The two pairs of bonding electrons are underlined. Hydrogen's effective atomic charge of +0.33 e means that 67% of an electron charge compensates the proton charge (+1 -0.67 = +0.33) and is thus associated with hydrogen. Only 33% of the two electrons taken up by the molecule in half reaction (16a) are transferred to oxygen.

Standard reduction potentials are not electron energies.
Interpreting standard reduction potentials as electron free energies in a half-reaction or half-cell, which is the underpinning of the "electron waterfall" picture of electron transfer chains, is highly problematic. The energetics of electrons in a half-cell is a difficult problem. It requires consideration of the outer potential y, inner potential f, surface dipolar potential c, contact potential (difference) between different metals, work function F, electrochemical potential, absolute electrode potential, etc. 15 The energy of the electron in its surroundings (in the simplest case a metallic electrode) must be taken into account. As a simple model system, consider a Zn/Cu galvanic cell: Zn(s) ® Zn 2+ (aq) + 2 e -(Zn) Cu 2+ (aq) + 2 e -(Cu) ® Cu(s) The free energy of the electron depends on the type of metal and on the (unmeasurable) inner (electrical) potential f. The energy of the electrons can be ignored only if the electrons end up in the same environment. Vacuum is such an environment, and for this reason one refers to vacuum absolute reduction potentials when only the difference of meaningful individual Gibbs free energies of chemical reactants and products, as we had done tacitly in the main text. This demonstrates that statements about the free energy of the electrons, e.g. in terms of standard reduction potentials, should be avoided. Fortunately, the electrons are only intermediates in the overall reaction and therefore their energies can be ignored. 7 Figure S5. Model system with two half-cells of different standard reduction potentials where the electrons in the two electrodes have the same free energy (same electrochemical potential): Zn/Cu galvanic cell with a good conductor connecting the electrodes, equalizing the electron electrochemical potential in the two electrodes, but without a salt bridge. (This set-up is not contrived but actually equivalent to two electrodes of two batteries connected in series, a common situation.) Figure S5 shows a set-up where the electrons in the electrodes of two half-cells with different standard reduction potentials have the same free energy (are at the same electrochemical potential) because they are connected by a good conductor. This disproves the claim that the standard reduction potential is proportional to the electron free energy.
The electron waterfall picture appears to represent the electron energy, or electrochemical potential, in the various steps of the transfer process. This would necessarily refer to the electron energy in the donor species, since that is where the electron exclusively resides, as pointed out earlier. This, however, is easily shown to be incorrect. For instance, the energy level of an electron in Fe 0 is not unique, being different for the half reactions Fe 0 ® Fe 2+ + 2eand Fe 0 ® Fe 3+ + 3 e -.
Reduced "electron-carrying" molecules do not carry the energy. Ubiquinol, QH2, is a molecule in the electron-transfer chain that carries electrons (from complexes I and II to complex III). From the standard reduction potential E o' = 0.045 V, we can conclude QH2 is lower in free energy than Q + 2 H + by -2 F E o' -2´413 kJ/mol = -835 kJ/mol, so QH2 is lower in energy than Q by -835 kJ/mol + 2´170 kJ/mol = -495 kJ/mol. This reduction in energy of QH2 relative to Q can be attributed to its stabilization by one extra O-H bond: Q has four double bonds, QH2 three "double" bonds and two single O-H bonds.
The analysis has shown that the electron-carrying, reduced form QH2 is of lower energy. Energy is carried by its oxidized counterpart, Q, which is less stabilized by bonding. When Q reacts with NADH + H + , some of this energy is released. On the other hand, when lower-energy QH2 reacts with a high-energy species, Fe 3+ (cytc), it gets charged up into its high-energy form, Q.
Energetics of the Calvin cycle. The CO2-fixation reaction of photosynthesis is an important biochemical process with energy implications, which goes by various different names (dark reaction (fixation of CO2), Calvin(-Benson) cycle). The net effect of the cycle is that atmospheric CO2 elongates sugar molecules, with the needed hydrogen atoms (not energy) provided by NADPH + H + produced in the primary reaction of photosynthesis, H2O + NADP + ® NADPH + H + + ½ O2. The net reaction of the Calvin cycle is 3 CO2 + 6 (NADPH + H + ) + 9 ATP + 8 H2O ® C3H5O3PO3H + 3 H2O + 6 NADP + + 9 ADP + 8 Pi Glyceraldehyde 3-phosphate The analysis and calculation is simplified if we use the shortcut introduced in eq. (13) 3 CO2 + 6 "H2"NADPH + 9 ATP + 8 H2O ® C3H5O3PO3H + 3 H2O + 9 ADP + 8 Pi The free energy change is quite minor (10% of the free energy of combustion of glucose), and its sign can be attributed to entropy (linking small molecules into a large one reduces their entropy). This disproves the notion that a lot of energy must be put in to convert a "highly oxidized molecule" like CO2 to a "more reduced" carbohydrate. When O2 is not involved, the energy changes are small. The energetics of this net reaction of the Calvin cycle have rarely been analyzed, probably since its negative enthalpy and small free-energy change cannot be rationalized in the dogmatic view that fuel molecules store the chemical energy. Instead, the canonical combustion reaction of glucose, or its reverse, is sometimes brought up in this context, 16 confusing the energetic picture. A chemical process with similar reactants as in eq.(28), the reverse of eq.(3c) with DrG o' = -804 kJ/mol, is indeed used by methanogens to produce energy, qualitatively confirming that the reaction of CO2 and H2 to form "more reduced molecules" can be energetically downhill.

Review of biochemical energy in textbooks.
A reviewer of this paper requested documentation that the misconceptions regarding biochemical energetics identified in this work are pervasive in current textbooks. The following quotes from eight biochemistry or molecular biology textbooks and three specialized books on biochemical thermodynamics provide evidence for this, in particular that O2 had not been identified as the molecule storing most biochemical energy. Particularly relevant phrases are highlighted in bold.
In Lehninger Principles of Biochemistry, 12 Figure 1-25 shows "energy transformations in living organisms": At the top, "Potential energy" is associated with sunlight and "Nutrients in environment (complex molecules such as sugars, fats)" and in the caption it is explained that "the potential energy of the complex nutrient molecules decreases." O2 is not shown at all in this figure. The text states that "Organisms obtain energy from their surroundings … they take up chemical fuels (such as glucose) from the environment and extract energy by oxidizing them." On p.22, it is asserted that "organisms … obtain the energy they need by oxidizing the energy-rich products of photosynthesis stored in plants, then passing the electrons thus acquired to atmospheric O2 …". On p. 25, it is stated that cells are "extracting energy (from nutrients such as glucose)…", on p. 357 that "Fats and oils are the principal forms of energy in many organism." and on p. 555 that: "Electron transfer from NADH to O2 in mitochondria provides the energy for synthesis of ATP…".
Alberts et al. 17 in Molecular Biology of the Cell assert very clearly (p. 54) that "All animal and plant cells are powered by energy stored in the chemical bonds of organic molecules… In both plants and animals, energy is extracted from food molecules …". The caption of Figure 2-18 refers to conversion of "…sunlight into chemical-bond energy in sugars and other organic molecules." According to p. 73, "The constant supply of energy that cells need … comes from the chemical-bond energy in food molecules." Their Figure 2-20 shows that reduced molecules like CH4 are high in energy because of more electrons on carbon. This is disproved in the present work. The crucial highenergy molecule, O2, is missing from this picture. On p. 67, the authors assert that "Activated carriers are specialized to carry electrons held at a high energy level (sometimes called "high-energy" electrons). The most important of the electron carriers are NAD + … and … NADP + " [this is extra wrong, because these forms of the molecules are obviously depleted of electrons]. "Each picks up a "packet of energy" corresponding to two electrons plus a proton." In the same vein, on p. 68 one can read that "NADPH operates … supplying the high-energy electrons needed to synthesize energy-rich biological molecules." The observation on p. 81 that "animals … cannot convert fatty acids to sugars." is explained in this work: fatty acids are low-energy molecules.
Becker's World of the Cell 18 clearly states that "Chemotrophs .. must depend completely on energy that has been packaged into oxidizable food molecules." close to Figure 5.4, whose caption ends in "Solar energy is used to reduce low-energy inorganic compounds to high-energy organic compounds, which are used by both phototrophs and autotrophs", while the main high-energy compound in the biosphere is inorganic O2.
Berg, Tymoczko, and Stryer in Biochemistry 19 , p. 444, show a diagram "Fuel (carbohydrates, fats) ® CO2 + H2O + useful energy" without O2, even though the oxygen-rich products imply its contribution. Clearly, the dominant contribution of O2 to the useful energy was not recognized here. For fats, the reaction is impossible without O2 and would not produce energy. According to p. 453: "Energy from foodstuffs is extracted in three stages", which attributes chemical energy to organic molecules, not O2. The text on p. 454: "The reduced forms of these carriers then transfer their high-potential electrons to O2." attributes energy to species other than O2, specifically reduced species. Table 15.2 asserts that NADH, NADPH, FADH2, and FMNH2 carry electrons, while they obviously carry and transfer H atoms. On p. 546, the energetics of aerobic respiration are summarized: "In oxidative phosphorylation, the electrontransfer potential of NADH or FADH2 is converted into the phosphoryl-transfer potential of ATP." without mentioning O2.
Campbell Biology 20 on p. 142 refers to "extracting the energy stored in sugars and other fuels". Further, (p.143) "… glucose and other organic fuels are broken down in the presence of oxygen … Energy that was stored in the organic molecules becomes available …" "these complex molecules, such as glucose, are in high in chemical energy. " Karp's Cell and Molecular Biology 11 states very clearly on p. 110 that "Carbohydrates are rich in chemical energy … Fats contain even greater energy per unit weight because the contain strings of more reduced CH2 units.", while it has been shown in this work that reduced molecules like fats are low in energy.
Harper's Illustrated Biochemistry 21 on p. 122 states: "Aerobic organisms are able to capture a far greater proportion of the available free energy of respiratory substrates…" attributing chemical energy not to O2 but to other reactants. Table 18-1 on p. 177 incorrectly attributes the energy of O2 to the citricacid cycle.
Voet & Voet in Biochemistry 22 initially do not seem to attribute biochemical energy to specific molecules. However, on p. 574 they write "NADH thereby functions as an energy rich electron-transfer coenzyme." This implicitly attributes the energy of the overall reaction in the respiratory chain, between NADH and O2, to NADH rather than O2. In Figure 22.9, the electron-transport chain is shown in terms of standard reduction potentials and associated "energy levels", with O2 at the bottom.
Physical chemists routinely set the energy of O2 to zero (since it is the element in its standard state). This makes the bond energy of O2 invisible. They rely on long tables of meaningless numbers (enthalpies of formation etc.). Therefore they have been unable to explain the energetics of combustion. Both fire scientists and some biochemists have empirically discovered the proportionality of the heat of combustion with O2 consumed ("oxygen combustion calorimetry"). But they attributed it to "the number of available electrons per mole of substrate" 23 or argued that "the energetic processes are the result of breaking either carbon-carbon or carbon-hydrogen bonds, and these have similar bond strengths" 24 . These arguments are all about the fuel, not O2. Correcting textbooks: an example. Can the needed dramatic revisions of biochemical energetics be incorporated into existing textbooks? The following example, a more correct rephrasing of an introductory text on chemical energy and electron-transfer chains, from Lehninger, 12 p. 528, shows that this is manageable. Significantly revised parts are underlined.
"Every time we use a motor, a refrigerator, or a battery charger, we use the flow of electrons to perform work. In the circuit that powers a motor, the source of electrons can be a battery in which a chemical reaction can only occur if electrons flow, usually through a metallic wire, from the half reaction at one pole of the battery, through the motor, to the half reaction at the other pole of the battery. The chemical force driving the electrons is manifested as the electromotive force, emf. The emf (typically a few volts) can accomplish work if an appropriate energy transducer -in this case a motor -is placed in the circuit. The motor can be coupled to a variety of mechanical devices to do useful work.
Living cells have analogous biological "circuits", with chemical reactions that convert high-energy reactants to lower-energy products driving the flow of electrons or protons. The energy released in exergonic reactions is channeled to drive the flow of electrons, avoiding loss as heat, in an electrontransfer chain involving a series of electron-carrier intermediates. The electrons provide energy to a variety of molecular energy transducers (enzymes and other proteins) that do biological work. In the mitochondrion, for example, membrane-bound enzymes couple electron flow, driven by energy that had been stored in the weak double bond of O2, to the production of a transmembrane proton concentration gradient and electrical potential, performing osmotic and electrical work. The resulting proton gradient has potential energy, sometimes called the proton-motive force by analogy with electromotive force. Another enzyme, ATP synthase in the inner mitochondrial membrane, uses the proton motive force to do mechanical work and convert it to chemical energy in ATP synthesized from ADP and Pi as protons flow S22 spontaneously across the membrane. Similarly, membrane-localized enzymes in E. Coli use chemical energy to drive electron and proton flow, which is then used to power flagellar motion. The principles of electrochemistry that relate free-energy changes in a battery to electrical energy driving a motor or charging a second battery apply with equal validity to the molecular processes accompanying electron and proton flow in living cells."