Nitrous Oxide Formation in the Colne Estuary in England: the Central Role of Nitrite
Dong et al. (2) recently examined the relationship between N2O emission and nitrite and nitrate concentrations in the Colne estuary in England. From calculations of the free-energy changes of nitrate- and nitrite-based N2O formation, the authors conclude that N2O formation using nitrite as the electron acceptor is favored in the Colne estuary. Unfortunately, there were a few glitches in those calculations. But, more importantly, the use of thermodynamics to explain the observed importance of nitrite as a critical factor regulating the formation of N2O in high-nutrient-load estuaries is questionable.
The changes in free-energy (ΔG°) values for nitrite- and nitrate-based N2O formation are given in Table 1. Important here are the choice and definition of the standard conditions. The ΔG° values in the first column are for standard conditions where N2O is at a partial pressure of 101 kPa (1 atm). These were the values used by Dong et al., but in subsequent calculations for nonstandard conditions, these authors worked with concentrations in solution. Hence, the standard values that they should have used are the ones in units of 1 M concentrations, i.e., the values given in the third column of Table 1. The standard conditions in the first column are for pH 7. For the pH occurring in the estuary, i.e., for pH 8, the values for nitrite- and nitrate-based N2O formation are, respectively, 5.7 and 5.7/2 kJ/mol of H2 higher than those at pH 7 (Table 1, columns 2 and 4). In their calculations, Dong et al. (2) did not take into consideration that their standard condition was already pH 7 but corrected with eight times the factors indicated above.
The objective of Dong et al. was to evaluate which of the two was energetically more favorable under the conditions occurring in the estuary: nitrate- or nitrite-based N2O formation (2). To this end, they derived the equation [NO2−]/[NO3−]0.5 = 0.02758[N2O]0.25, which should (after the corrections indicated above) read [NO2−]/[NO3−]0.5 = 0.000021[N2O]0.25 in molar units at pH 8. Using N2O concentrations of 1 and 1,000 nM, which are the extremes observed in the Colne estuary, as constraints, this equation implies that at [NO2−]/[NO3−]0.5 above 0.00000012 for [N2O] = 1 nM and above 0.00000070 for [N2O] = 1,000 nM, N2O production from nitrite is energetically more favorable than N2O production from nitrate. Dong et al. (2) report for the Colne estuary at, e.g., site 1 a value of 2.76 for [NO2−]/[NO3−]0.5. This value, however, is for micromolar concentrations. When expressed in molar concentrations, this value is 1,000-fold lower at 0.00276, which still leads to the conclusion that N2O formation from nitrite is energetically more favorable than N2O formation from nitrate in the Colne estuary.
The approach, with a nonintuitive component like [NO2−]/[NO3−]0.5 presented as an equation and then setting extremes as constraints, is correct but difficult to follow. It would be easier to directly make the appropriate ΔG calculations with the actual data points. A major advantage of the direct approach is also that it forces one to take the source and concentration of the reducing equivalents into account. If the concentration of reactants is such that the reaction is endergonic, this will show up in the outcome of the ΔG calculation while this remains hidden in the approach taken by Dong et al. Figure 1 illustrates that nitrite-based N2O formation was energetically more favorable than nitrate-based N2O formation in the Colne estuary. However, it is important to note that conversion of N2O to N2 was even more favorable than its formation. Hence, it is, from a thermodynamic point of view, not logical that N2O was emitted from the Colne estuary. This conclusion is hardly affected by the assumed hydrogen concentration in the system. In the calculations for Fig. 1, the assumed hydrogen partial pressure was 0.1 Pa but since the ΔG value changes by 5.7 kJ/mol of H2 for every 10-fold change in the partial pressure (or concentration) of H2, the same conclusion would also be reached at a 1010-times lower hydrogen partial pressure.
It is questionable whether the formation of N2O in the Colne estuary should be evaluated by the energetics of the reactions involved. Additions of nitrate and nitrite stimulated N2O formation (2), indicating that N2O formation was not limited by the availability of reducing equivalents. Hence, the question of whether nitrate or nitrite was the energetically more favorable electron acceptor for N2O production is not an issue. Under conditions of electron acceptor limitation, the organisms will tend to use every electron acceptor available, with the qualification that the energetically most favorable electron acceptor will be used first. Obviously, this criticism does not affect the conclusions of the authors that nitrite may be the limiting factor regulating the formation of N2O in highly nutrient-loaded estuaries and that control of nitrite input would be an effective way of reducing the production and release of N2O.
FIG. 1.
TABLE 1.
| Reaction | ΔG° (kJ/mol of H2) | ||||||
|---|---|---|---|---|---|---|---|
| N2O at 1 atm | [N2O] = 1 M | ||||||
| pH 7 | pH 8 | pH 7 | pH 8 | ||||
| NO2− + H2 + H+ → 0.5 N2O + 1.5 H2O | −226.6 | −220.9 | −222.0 | −216.3 | |||
| NO3− + 2 H2 + H+ → 0.5 N2O + 2.5 H2O | −194.8 | −192.0 | −192.0 | −189.7 | |||
| N2O + H2 → N2 + H2O | −341.4 | −341.4 | −350.6 | −350.6 | |||
Authors' Reply
We have carefully considered Dr. Dolfing's comments on the issue of free-energy calculations. We agree that we made an error in the calculations of the ΔG° by using 1 atm rather than molar concentrations and that the ΔG° should have been corrected to pH 8. We thank him for his corrections. However, the consequences of these corrections, with much lower values of [NO2−]/[NO3−]0.5, are to support even more strongly our conclusions on the relative energy yields of nitrite and nitrate as substrates for conversion to N2O.
We attempted to use a thermodynamic argument to indicate whether, under certain conditions, nitrite conversion to N2O might be energetically advantageous to nitrate conversion to N2O and, hence, in an energy-limited environment to favor those organisms using this pathway. However, although thermodynamics can indicate whether a particular reaction should proceed under given conditions of temperature, pressure, concentration, etc., it tells us nothing about the reaction mechanisms or the kinetics. The overall reactions from nitrite or nitrate to N2O or to N2 are catalyzed by different enzymes, such as nitrate reductase, nitrite reductase, nitrous oxide reductase, etc., which may vary phylogenetically. It is true to say that, as Dolfing puts it, “it is, from a thermodynamic point of view, not logical that N2O was emitted from the Colne estuary” because of higher free-energy production from conversion of N2O to N2 than its formation from nitrate or nitrite. However, whether or not N2O is emitted largely depends on the balance of the rates of its formation and removal. If formation is faster than removal, N2O may accumulate. Conversion of N2O to N2 by nitrous oxide reductase can be inhibited by a number of compounds (3, 6). For instance, sulfide, commonly present in high-sulfate marine and estuarine sediments, inhibits nitrous oxide reductase, so that N2O accumulates. This can be a major factor in the production of N2O in some estuarine sediments (7) and operates over and above any thermodynamic considerations. Also, some denitrifying bacteria may be genetically able only to use nitrite (not nitrate) as the electron acceptor and produce N2O (not N2) as the sole end product (1). Thus, factors other than thermodynamics may also influence whether N2O is formed in estuaries.
Dr. Dolfing argues for a more direct thermodynamic model for nitrate/nitrite reduction that explicitly considers concentrations of electron donors and acceptors. However, nitrate reducers can use a wide array of different electron donors, whose in situ concentrations are unknown. An explicit model would have to make assumptions about these unknown electron donor concentrations. The advantage of our approach is that, by using nitrite/nitrate ratios ([NO2−]/[NO3−]0.5), we can make conclusions about the relative free energy without having to make any assumptions about the unknown electron donor concentrations.
Dr. Dolfing (paragraph 4, sentence 4) makes the comment that if the concentrations of all reactants, including electron donors, is such that the reaction is endergonic, this might be missed by our approach. Considering the high standard free-energy release from nitrite or nitrate reduction to N2O and the actual concentrations of the reactants and products in the sediment, the reduction of nitrate or nitrite to N2O is extremely unlikely to be endergonic. For a given sediment, the concentrations of reducing substrates present in a sediment to drive nitrate/nitrite reduction is unknown but may be assumed to be relatively constant, at least compared to electron acceptor concentrations, which may vary tidally. So, evaluating energy yields by using nitrite or nitrate as an electron acceptor becomes a matter of evaluating the relative concentrations of nitrite and nitrate, i.e., the ratio [NO2−]/[NO3−]0.5. At pH 8 and 25°C, if the ratio [NO2−]/[NO3−]0.5 is >0.00000012 at [N2O] = 1 nM or >0.00000070 at [N2O] = 1,000 nM, the reaction using nitrite as the electron acceptor should be energetically more favorable. This approach has a wider justification and can be easily used for other sediments.
Dr. Dolfing comments in his last paragraph that, as addition of nitrate/nitrite stimulated N2O formation, N2O formation was not limited by the availability of reducing equivalents (i.e., the electron donor concentration) and, hence, that there is no merit in looking at the energetics of the reactions involved in its formation from nitrate/nitrite. This assumes that there is a single limiting factor (e.g., the nitrate or nitrite concentration) controlling such reactions in sediments, whereas in many sediments there is substantial evidence of simultaneous regulation of microbial activities by a number of different environmental variables, including both electron donor and electron acceptor concentrations, and physical factors such as temperature also (5). Egli et al. (2) have discussed models for simultaneous regulation of growth by a variety of organic substrates under oligotrophic conditions (and see reference 4 also), and the same models can be applied to simultaneous regulation by electron acceptors also. Addition of nitrate may stimulate denitrification in sediment, but addition of nitrate plus acetate or glucose usually stimulates rates even more. Thus, in many sediments, both electron donors and acceptors are simultaneously at limiting concentrations and there will be interaction between multiple factors that limit rates of denitrification and nitrous oxide production. Measurements of potentials for denitrification, when both electron donors and acceptors are provided in excess, gives rates of denitrification much greater than measured in situ rates. Thus, that addition of nitrate and nitrite stimulated N2O formation may not necessarily mean that N2O formation was not also limited by the availability of organic substrates.
The further assumption by Dr. Dolfing in his last paragraph that “whether nitrate or nitrite was the energetically more favorable electron acceptor for N2O production is not an issue” because all electron acceptors possible will be used is not correct. In a closed system, it is true to say “under conditions of electron acceptor limitation, the organisms will tend to use every electron acceptor available, with the qualification that the energetically most favorable electron acceptor will be used first.” A sediment is an open system; nitrate and nitrite can be used by denitrifying bacteria but be continuously resupplied by inputs of nitrate from overlying water or by internal nitrification and would never be completely exhausted by the microorganisms. Therefore, the question of which electron acceptor, nitrite or nitrate, would be energetically more favorable in the process of nitrous oxide formation remains important.
REFERENCES
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Dong, L. F., D. B. Nedwell. G. J. C. Underwood, D. C. O. Thornton, and I. Rusmana.2002. Nitrous oxide formation in the Colne estuary, England: the central role of nitrite. Appl. Environ. Microbiol.68:1240-1249.
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Egli, T., U. Lendenmann, and M. Snozzi.1993. Kinetics of microbial growth with mixtures of carbon sources. Antonie van Leeuwenhoek63:289-298.
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Firestone, M. K., and J. M. Tiedje.1979. Temporal changes in nitrous oxide and dinitrogen from denitrification following onset of anaerobiosis. Appl. Environ. Microbiol.38:673-679.
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Gottschal, J. C.1993. Growth kinetics and competition—some contemporary comments. Antonie van Leeuwenhoek63:299-313.
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Nedwell, D. B., and J. W. Abram.1979. Relative influence of temperature and electron donor and electron acceptor concentrations on bacterial sulfate reduction in saltmarsh sediment. Microbiol. Ecol.5:67-72.
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Sørensen, J., J. M. Tiedje, and R. B. Firestone.1980. Inhibition by sulfide of nitric and nitrous oxide reduction by denitrifying Pseudomonas fluorescens. Appl. Environ. Microbiol.39:105-108.
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Trimmer, M., D. B. Nedwell., D. B. Sivyer, and S. J. Malcolm.2000. Seasonal organic mineralisation and denitrification in intertidal sediments and their relationship to the abundance of Enteromorpha sp. and Ulva sp. Mar. Ecol. Prog. Ser.203:67-80.
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Amend, J. P., and E. L. Shock.2001. Energetics of overall metabolic reactions of thermophilic and hyperthermophilic Archaea and Bacteria. FEMS Microbiol. Rev.25:175-243.
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Dong, L. F., D. B. Nedwell, G. J. C. Underwood, D. C. O. Thornton, and I. Rusmana.2002. Nitrous oxide formation in the Colne estuary, England: the central role of nitrite. Appl. Environ. Microbiol.68:1240-1249.
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Applied and Environmental Microbiology
Volume 68 • Number 10 • October 2002
Pages: 5202 - 5204
PubMed: 12324378
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© 2002.
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Published online: 1 October 2002
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