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REVIEW article

Front. Mar. Sci., 15 February 2023
Sec. Marine Biogeochemistry
This article is part of the Research Topic The Marine Iodine Cycle, Past, Present and Future View all 11 articles

Review on the physical chemistry of iodine transformations in the oceans

  • School of Marine Science and Policy, University of Delaware, Lewes, DE, United States

The transformation between iodate (IO3), the thermodynamically stable form of iodine, and iodide (I-), the kinetically stable form of iodine, has received much attention because these species are often dependent on the oxygen concentration, which ranges from saturation to non-detectable in the ocean. As suboxic conditions in the ocean’s major oxygen minimum zones indicate that IO3 is minimal or non-detectable, the incorporation of IO3 into carbonate minerals has been used as a redox proxy to determine the O2 state of the ocean. Here, I look at the one and two electron transfers between iodine species with a variety of oxidants and reductants to show thermodynamics of these transformations. The IO3 to IO2 conversion is shown to be the controlling step in the reduction reaction sequence due to thermodynamic considerations. As IO3 reduction to IO2 is more favorable than NO3 reduction to NO2 at oceanic pH values, there is no need for nitrate reductase for IO3 reduction as other reductants (e.g. Fe2+, Mn2+) and dissimilatory IO3 reduction by microbes during organic matter decomposition can affect the transformation. Unfortunately, there is a dearth of information on the kinetics of reductants with IO3; thus, the thermodynamic calculations suggest avenues for research. Conversely, there is significant information on the kinetics of I- oxidation with various oxygen species. In the environment, I- oxidation is the controlling step for oxidation. The oxidants that can lead to IO3 are reactive oxygen species with O3 and •OH being the most potent as well as sedimentary oxidized Mn, which occurs at lower pH than ocean waters. Recent work has shown that iodide oxidizing bacteria can also form IO3. I- oxidation is more facile at the sea surface microlayer and in the atmosphere due to O3.

1 Introduction

The thermodynamically favorable form of iodine in seawater is iodate (IO3). However, iodide (I-) is present in oxic, suboxic and anoxic waters. The one electron transfer reaction of I- with molecular oxygen, 3O2, to form the iodine atom (I•) and superoxide (O2-) is thermodynamically unfavorable as is the reaction of two I- with 3O2 to form I2 and H2O2 (Luther et al., 1995; Luther, 2011). Thus, other oxidants are required to initiate abiotic iodide oxidation, and I- is a known sink for O3. Biotic iodide oxidation has received much interest with one report showing conversion of iodide to iodate (Hughes et al., 2021). Iodate reduction can occur with common reductants (e.g., sulfide, Fe2+), and various organisms that decompose organic matter using iodate as the electron acceptor. Nitrate reductase and dimethyl sulfoxide reductase enzymes from these and planktonic microbes are considered important mediators for biotic iodate reduction (e.g., Hung et al., 2005; Amachi, 2008). Thus, there has been extensive interest in the chemistry of these two iodine species and the possible intermediates that form during their 6-electron redox interconversion ever since the element, iodine, was first discovered as I2 during the study of brown kelp algae of the Laminariales (kelps/seaweeds) by Courtois in the early 1800s (Wong, 1991; Küpper et al., 2008; Küpper et al., 2011).

Possible chemical species that form during the IO3I interconversion are given in eqn. (1a, b). The loss of an O atom is equivalent to a two-electron transfer resulting in the reduction of the iodine from +5 in IO3 to +3 for IO2 (HOIO) to +1 for HOI (IO-) and to -1 for I-. The acid-base species in parentheses are minor species at seawater pH as the pKa values for HOI and HOIO are 10.7 (Ka = 2 x 10-11) and 4.49 (Ka = 3.2 x 10-5), respectively.

IO3IO2(HOIO)HOI(IO)I(1a)

Equation 1b shows the interconversion between HOI and I- as HOI undergoes one-electron transfer to I2 followed by another one-electron reduction per I atom to I-.

HOI(IO)I2I(1b)

This work considers the thermodynamics of these transformations during the reduction of IO3, which occurs in anoxic systems, during organic matter decomposition and by phytoplankton, as well as the oxidation of I-, which occurs by the direct oxidation of iodide by iodide oxidizing bacteria, oxidized metals and reactive oxygen species (ROS) that are produced by certain microbes, (macro)algae and abiotic processes including photochemistry. In a previous work (Luther, 2011), the chemistry and thermodynamics of chloride, bromide and iodide oxidation were compared; however, I(+3) species (HOIO, iodous acid, and IO2, iodite) and stepwise iodate reduction were not considered. Here, stepwise reactions of the iodine species in equations 1a and 1b with environmentally important reactants (including transient ROS species) are considered for both the oxidation of I- and the reduction of IO3 to affect their interconversion. The kinetics of these stepwise reactions are also considered. Kinetic data for the first step(s) in iodide oxidation are available, but less kinetic information is available for iodate reduction.

2 Methods

2.1 Calculations of aqueous redox potentials from half-reactions

Table 1 gives several equations for redox half-reactions that include the pH dependence for the reaction considered. These are pϵ(pH) relationships based on the balanced chemical equations and the thermodynamics of each chemical species. The basic mathematical approach has been fully developed in standard textbooks (Stumm and Morgan, 1996; Luther, 2016) and used in previous publications (Luther, 2010; Luther, 2011). Aqueous thermodynamic data to calculate the pϵ(pH) or logK(pH) relationships in Table 1 (at 25°C and 1 atm) are from Stumm and Morgan (1996) and other sources (Bard et al., 1985; Stanbury, 1989). The value used for the Gibbs free energy for Fe2+ (-90.53 kJ/mole) is that discussed in Rickard and Luther (2007). Values of the free energy for HOIO and IO2 are from Schmitz (2008).

TABLE 1
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Table 1 Reduction half-reactions for relevant species of oxygen, nitrogen, sulfur, manganese, iron and iodine normalized to one electron.

The calculated pϵ value from each half-reaction is given as a function of pH as in the examples in Table 1, and these half reactions can be used for simple calculations of the pE values of full reactions (see next section). When H+ or OH- is not in a balanced equation for a half-reaction, there is no pH dependence on the half-reaction. The pε calculated is termed pϵ(pH) which provides a log K for each half-reaction at a given pH. Concentration dependence for the other reactants are not considered in the calculation; thus, these are considered standard state calculations. When concentration dependence is considered, the calculated pϵ value can vary as in the following example for the O2/H2O couple (O1 in Table 1).

Using the balanced half reaction and the Gibbs free energy of formation of each species at 25°C, the Gibbs free energy of the reaction and the equilibrium constant are calculated.

O2(aq)+ 4 H++ 4e  2H2O16.32     0       0       2(237.18) ΔG°formation(kJ/mol)

The standard state ΔG° for the reaction = - 490.68 kJ/2 moles H2O or 4 moles of electrons. The equilibrium constant (K40) is given in eqn. 2a where {} indicates activity for each chemical species and the activity of H2O is defined as 1.

K40={H2O}2{O2} {H+}4 {e}4 (2a)

On expanding, eqn. 2b results.

logK40=log{O2(aq)}log{H+}4log{e}4 = ΔGreaction02.303 RT=86.00  (2b)

logK40 is for 4 mole of electrons or 21.50 for 1 mole of electrons.

For a one-electron half-reaction, we have

14O2(aq)+H++e12H2O

And equation 2b becomes equation 2c

14logK40= 14log{O2(aq)}log{H+}log{e}(2c)

or

14logK40=14log{O2(aq)}+pH+ and on rearranging
 = 14logK40+14log{O2(aq)}pH

From the Nernst Equation, o=1/4 log K40=21.50 (the standard state value), which on substitution gives equation O1 (see Table 1) where concentration is used for O2(aq).

=o+14log[O2(aq)]pH=21.50+14log{O2(aq)}pH(O1)

At ocean surface conditions of 211 µM O2 (211 x 10-6 M; 100% saturation at 25°C and salinity of 35), this expression becomes

=21.50+14 log[211×106M]pH=20.58pH

and at a pH of 8, pε = 12.58.

At 1µM O2 (10-6 M) which occurs in oxygen minimum zones, this expression becomes

=21.50+14log[106M]pH =20.00pH

and at a pH of 7.5, pε = 12.50.

At unit activity for all reagents including H+, pε = pε°. At unit activity of all reagents other than the H+, equation O1a results, which is used for many calculations in this paper.

=opH=21.50pH(O1a)

Note that the above equations show a 1.50 log unit change for an O2 concentration range from 1 µM to unity activity (O1a) so the calculations could vary an order of magnitude or more in either direction when concentration dependence is included. However, comparisons can be more easily made when combining different half-reactions at a given pH. This permits an assessment of which combined half-reactions are thermodynamically favorable and thus more likely to occur in a given environmental setting.

2.2 Coupling half-reactions

As an example of coupling two half reactions to determine whether a reaction is favorable, I use the data in Table 1 for the reduction of IO3 (Io5b) by NO2 (N1) in equation 3.

NO2+IO3NO3+IO2(3)

Equation 4 is used to calculate a complete reaction’s pε or ΔlogKreaction value. All values of ΔlogKreaction > 0 indicate a favorable reaction and all values of ΔlogKreaction< 0 indicate an unfavorable reaction.

reaction=red+oxid=ΔlogKreaction(4)

At a pH of 7, the pεred values for IO3 and NO3 are 7.91 and 7.28, respectively. As NO2 is the reductant, it is oxidized; thus, the sign for pεred (7.28) is reversed to become pεoxid (-7.28).

ΔlogKreaction=red(IO3)+oxid(NO3)=7.91+(7.28)=0.63

For reaction 3, there is no pH dependence as the pH dependence of each half-reaction is similar so cancels.

For this work, Table 1 lists the pϵ(pH) values for Mn, Fe, oxygen, nitrogen, sulfur and iodine species for the relevant iodine redox reactions considered. Dissolved Fe(II) and Mn(II) are primarily hexaaquo species until the pH is > 7, where hydroxo complexes start to become important. As most reactions occur via one and two-electron transfers, the calculations will permit assessment of a thermodynamically unfavorable step along a reaction coordinate of six-electrons as in the reduction of iodate to iodide and the oxidation of iodide to iodate. From surface waters to decomposition zones, seawater pH values range from 8 down to 7; thus, the following discussion will emphasize this pH range.

3 Results and discussion: Iodate reduction

3.1 Iodate and iodide speciation at different seawater oxygen conditions

In the oxic environment, the oxidizing condition of the environment or pε is set by the 4-electron transfer reaction of the O2(aq)/H2O couple [reaction O1 in Table 1]. At a pH of 8, temperature 250C and a salinity of 35, 100% O2(aq) saturation is 211 μM, which gives a pε of 12.58 (Figure 1). As the IO3/I couple has a pε of 10.56 at pH=8,IO3 is the thermodynamically favored iodine species.

FIGURE 1
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Figure 1 IO3 and I- concentrations for a total iodine concentration of 450 nM at different environmental pε values assuming that O2(aq) reduction to H2O sets the pε value of ocean waters.

Entering the pε value for a given [O2(aq)] into equation Io6 allows the determination of the iodide to iodate ratio and the actual concentration of each assuming a total iodine concentration of 450-470 nM (Elderfield and Truesdale, 1980). Figure 1 shows the iodate and iodide concentrations are equivalent at a pε of 10.56. The vertical lines indicate the environmental pε for [O2(aq)] of 1, 10, 100 nM, and 1, 50 and 211 μM. As oxygen minimum zones (OMZ) of the Arabian Sea and the equatorial Pacific Ocean have [O2(aq)] concentrations in the 1-100 nM range (Revsbech et al., 2009; Lehner et al., 2015), calculations show that IO3 is the thermodynamically preferred iodine species even at 1 nM O2(aq), which gives a pε of 11.25 for the O2(aq)/H2O couple. However, I- is the dominant iodine species detected in OMZ waters (Wong and Brewer, 1977; Luther and Campbell, 1991; Rue et al., 1997; Farrenkopf and Luther, 2002; Cutter et al., 2018). At [O2(aq)] concentrations ≤ 1 μM, IO3, NO3 and Mn2+ concentrations are now similar or higher in concentration and should determine the pε of the water.

As most reactions occur by 1- or 2-electron transfers, Figure 2 shows the redox sequence for two electron transfer redox couples NO3/NO2(N1), MnO2/Mn2+ (Mn1) and the one-electron redox couple Fe(OH)3/Fe2+ (Fe3) over a wide range of pH. The redox sequence at pH = 8 is as expected for the N, Mn and Fe systems. As NO2 (up to 12 μM) and Mn2+ (up to 8 μM) are formed at OMZ oxic-anoxic transition zones (e.g., Arabian Sea, Black Sea, Equatiorial Pacific, see Lewis and Luther, 2000; Trouwborst et al., 2006; Cutter et al., 2018, respectively) and are in higher concentration than O2(aq), the NO3NO2(N1) and MnO2 → Mn2+ (Mn1) couples can be chosen to set the environmental pε. At pH = 8, the NO3 to NO2 pε is 6.15 and the MnO2 to Mn2+ pε is 4.80. At pH = 7, the NO3 to NO2 pε is 7.28 and the MnO2 to Mn2+ pε is 6.80. At these pε values, O2(aq) is below 1 nM, and I- is now the thermodynamically favored iodine species when comparing these data with the IO3/I couple (pε of 10.56 at pH = 8).

FIGURE 2
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Figure 2 Two electron transfer redox couples for O3 (O4), N (N1), Mn (Mn1) and I (Io2a, Io4b, Io5b), and one electron transfer redox couple for Fe (Fe3). The oxidized species is always above the line and the reduced below the line as in the O2/H2O and the H2O/H2 couples, which dictate the water stability field.

Figure 2 also shows that the IO3/IO2 couple (Io5b) should be the first step in the reaction sequence of iodate to iodide (eqn. 1a). As the pε of the IO3/IO2 couple has a more positive pε value than the N, Mn and Fe couples in Figure 2, IO3 reduction is more favorable than these couples even though it is very close to the NO3/NO2 couple. Thus, IO3 is predicted to reduce before NO3, and biological activity (e.g., nitrate reductase activity) is not necessary to reduce IO3 (see section 3.2). Because of the strong pH dependence for the Mn and Fe couples, they cross the IO3/IO2 and NO3/NO2 couples at lower pH, which have similar slopes. Thus, NO2 is predicted to reduce IO3 to IO2 (eqn. 3). Interestingly, Mn2+ and Fe2+ should be poorer reductants than NO2 for conversion of IO3 to IO2 at a pH< 6 and pH< 1, respectively, but are more favorable to reduce IO3 than NO2 above those pH values (see sections 3.2 and 3.3).

Although the IO3 to I- conversion occurs at higher pε, it is a 6-electron transfer (IO6), which is not a facile process. Thus, the intermediates (IO2 and HOI) will dictate the reactivity sequence via a combination of thermodynamic and kinetic considerations.

As shown in Figure 2, IO2 reduction to HOI and HOI reduction to I- are also more favorable at higher pε values than the IO3 to I- couple. At a pH = 8, the IO2 to HOI couple has a pε value of 12.06 corresponding to 2 μM O2 (see Figure 1). Similarly, the HOI to I- couple has a pε value of 12.66 corresponding to 250 μM O2. At a pH of 7, both couples have pε values greater than 13 indicating that, even at O2 saturation, I- is the dominant species predicted when these intermediates form. At a pH = 7.5 (that is found in many OMZ waters), both IO2 to HOI and HOI to I- couples have pε values greater than 12.6; also indicating that at O2 saturation, I- is the dominant species predicted. Thus, the intermediates IO2 and HOI are not predicted to be stable in marine waters; thus, the conversion of IO3 to IO2 is a key step. Interestingly, Hardisty et al. (2021) found in situ IO3 reduction in the oxycline where [O2(aq)] was 11 μM, but not at [O2(aq)]< 2 μM. Lastly, the O3 to O2 + H2O couple is highly oxidizing indicating that all iodine couples should lead to IO3 formation. O3 reactions will be discussed in more detail below (sections 4.2, 4.7).

In the next sections (3.2 – 3.5), the thermodynamics for the conversion of iodate to iodide via the intermediates outlined in equations 1a and 1b by environmental reductants are considered to show what step, if any, in the reduction of iodate to iodide may be unfavorable over a wide range of pH. Iodate reduction is well known in the marine environment (e.g., Wong and Brewer, 1977; Wong et al., 1985; Luther and Campbell, 1991; Rue et al., 1997; Farrenkopf and Luther, 2002; Cutter et al., 2018) and occurs via chemical reductants like sulfide (Zhang and Whitfield, 1986) and via microbes like Shewanella putrfaciens (Farrenkopf et al., 1997) and Shewanella oneidensis (Mok et al., 2018) during dissimilatory reduction coupled with decomposition (oxidation) of organic matter as well as phytoplankton mediated processes (e.g., Chance et al., 2007).

3.2 Iodate reduction by NO2-

Although NO2 has not yet been shown to be a reductant for IO3 in aqueous lab studies (eqn. 3), HNO2 is a reductant for MnO2 (Luther and Popp, 2002) and Mn(III)-pyrophosphate (Luther et al., 2021). Figure 3A shows the thermodynamic calculations for the stepwise conversion of IO3 to I- by NO2 reduction (NO2 oxidizes to NO3). All 2-electron transfer reactions, which involve O atom loss for iodine, are favorable over the pH range. For seawater pH (7-8), the least favorable reaction is the IO3 to IO2 reaction whereas the IO2 to HOI and HOI to I- reactions are more favorable. Thus, the IO3 to IO2 conversion appears to be the controlling step in the reaction sequence. The 1-electron transfer reaction of HOI to I2 is the most favorable, but the second 1-electron transfer reaction of I2 to I- is only favorable at pH > 4. Thus, reduction of IO3 to I- by NO2 is predicted via 1-electron or 2-electron transfer reactions at seawater pH values. The data plotted in Figure 2 indicate that once IO2 forms there is no thermodynamic barrier to I- formation.

FIGURE 3
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Figure 3 Thermodynamics for (A) the 1-electron and 2-electron reductions of IO3 (Io5a, Io5b, Io6), IO2 (Io4b), HOI (Io2, Io3), and I2 (Io1) by NO2 (N1) and (B) 2-electron reductions of IO3 (Io5b), IO2 (Io4a, Io4b), HOI (Io2) by DMS (S3). The vertical line represents the pKa value of 4.49 for HO2I. Data above the horizontal line at ΔlogK (ΔlogKreaction) = 0 indicate a favorable reaction and data below the horizontal line indicate an unfavorable reaction.

The IO3 reaction with NO2 has been reported to produce I2 in ice by Kim et al. (2019), but not in solution. The pH in the ice was 3 where HNO2 and H2ONO+ exist and are the likely reductants. Thus, polar areas may be locales for IO3 reduction. At seawater pH, the reaction seems to be hindered by kinetics in the transition state as each reactant (IO3 and NO2) is an anion, which will repel each other.

3.3 Biological iodate reduction

The marine literature has many reports on the uptake of IO3 (with or without NO3) by phytoplankton with the iodine released as I- (e.g., Elderfield and Truesdale, 1980; Wong, 2001; Wong et al., 2002; Chance et al., 2007; Bluhm et al., 2010). As a result of this iodate uptake, NO3 reductase was presumed by some researchers to be a key process for IO3 reduction to I-. Also, Bluhm et al. (2010) and Carrano et al. (2020) showed that algal senescence enhanced I- release, and Hepach et al. (2020) showed that there is a considerable lag between IO3 uptake and I- release due to senescence.

NO3 reductase appears to reduce IO3 in some phytoplankton (Hung et al., 2005). However, de la Cuesta and Manley (2009) showed that I- can be up taken by phytoplankton, and that different phytoplankton uptake I- whereas other phytoplankton uptake IO3. Thus, there is no need for nitrate reductase for IO3 reduction as I- can be up taken by some phytoplankton rather than form from IO3 reduction. Moreover, Waite and Truesdale (2003) showed that nitrate reductase was not important for IO3 reduction by Isochrysis galbana. The latter study is consistent with the thermodynamics of the reduction IO3 to IO2 being more favorable than the reduction NO3 to NO2.

Furthermore, under anaerobic conditions, dissimilatory IO3 reduction occurs without nitrate reductase for the denitrifying bacterium, Pseudomonas stutzeri, (Amachi et al., 2007; Amachi, 2008). Reyes-Umana et al. (2022) and Yamazaki et al. (2020) showed that iodate reductase is in the periplasmic space of Pseudomonas sp SCT. Also, Mok et al. (2018) showed that dissimilatory IO3 reduction by Shewanella oneidensis does not involve nitrate reductase. Recently, Shin et al. (2022) showed that Shewanella oneidensis requires extracellular dimethylsulfoxide (DMSO) reductase involving a molybdenum enzyme center for IO3 reduction. Guo et al. (2022) studied bacterial genomes in a variety of environments and documented that Shewanella oneidensis are ubiquitous in all fresh and marine waters; they concluded that IO3 reduction is a major biogeochemical process. Thus, nitrate reductase (also an O atom transfer reaction) is not a requirement for bacterial IO3 reduction to I-.

The interconversion of dimethylsulfoxide with dimethylsulfide during dissimilatory IO3 reduction is another 2-electron O-atom transfer reaction. Moreover, the reactions of DMS to reduce HOI, IO2 and IO3 are thermodynamically favorable (Figure 3B, DMSO reduction is in S3, Table 1 and occurs at a lower pε than NO3 and IO3 reduction). The reaction of DMS with HOI has been suggested by Müller et al. (2021) to be a sink for DMS based on the rapid reaction of DMS with HOBr.

3.4 Iodate reduction by Mn2+ and Fe2+

Figure 4 shows the thermodynamics for the stepwise conversion of IO3 to I- by reduction with Mn2+ and Fe2+. Concentrations of Mn2+ and Fe2+ range from several nM to μM in OMZs (e.g., Trouwborst et al., 2006; Moffett and German, 2020) and in suboxic porewaters (e.g., Oldham et al., 2019; Owings et al., 2021) to mM in waters emanating from hydrothermal vents (e.g., Estes et al., 2022); in these cases, Mn2+ and Fe2+ are normally higher in concentration than the total iodine concentration. For the 2-electron transfer reactions with Mn2+, only the IO2 to HOI reaction is favorable over the entire pH range. The IO3 to IO2 reaction is favorable only at pH > 6 whereas the other reactions are favorable at pH > 3. Thus, the IO3 to IO2 conversion is the controlling step in the reaction sequence when Mn2+ is the reductant. Using high resolution porewater profiles of I- and Mn2+ obtained by voltammetric microelectrodes, Anschutz et al. (2000) showed that a I- maximum occurred at the depth where upward diffusing Mn(II) was being removed and proposed that I- formed by the reaction of IO3 with Mn2+ under suboxic conditions. The reaction has not been investigated in laboratory studies.

FIGURE 4
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Figure 4 Thermodynamics for the reduction of IO3 (Io5a, Io5b, Io6), IO2 (Io4b), HOI (Io2, Io3), and I2 (Io1) by Mn2+ (Mn1) and Fe2+ (Fe3). The vertical line represents the pKa value of 4.49 for HO2I. Data above the horizontal line at ΔlogK (ΔlogKreaction) = 0 indicate a favorable reaction and data below the horizontal line indicate an unfavorable reaction.

For the reaction sequence with Fe2+, all iodine species reductions are favorable over the entire pH range except for the I2 to I- reaction, which is favorable at pH > 2.5. Thus, there is no thermodynamic inhibition to IO3 reduction to I- by Fe2+, and this abiotic reaction at a pH of 7 was reported to be 92% complete after 2 hours using initial concentrations of 2 mM Fe2+ and 0.1 mM IO3 (Councell et al., 1997). Because the Fe(OH)3 to Fe2+ couple is a 1-electron transfer, two Fe2+ are required in each step of the sequence. Again, the IO3 to IO2 conversion is the least favorable and likely controlling step in this reaction sequence.

Comparing Figures 3, 4 indicates that the Mn2+ and NO2 reactions with iodine species have a similar range of ΔlogKreaction values whereas the Fe2+ reactions with iodine species are more favorable (higher ΔlogKreaction values).

3.5 Iodate reduction by sulfide

In sulfidic waters and porewaters, IO3 does not exist as sulfide reacts readily with it (Zhang and Whitfield, 1986), and S(0) forms as the initial sulfur product. Figure 5 shows the thermodynamics for the stepwise conversion of IO3 to IO2 and to HOI by sulfide where S(0) forms as an intermediate leading to S8. As the Gibbs free energy of formation for HSOH is unknown, HSOH could not be evaluated as an intermediate, which on continued oxidation would form SO42. The reaction of sulfide with I2 and HOI is well known as the iodometric titration, so calculations were not performed. The only unfavorable iodine reduction reactions are the 1-electron reductions that lead to the formation of the HS radical (HS• or HS rad). The conversion of IO2 to HOI is more favorable as it has the larger ΔlogKreaction values. Again, the IO3 to IO2 conversion is the least favorable and likely controlling step in this reaction sequence.

FIGURE 5
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Figure 5 Thermodynamics for the 2-electron transfer reductions of (A) IO3 (Io5b) and (B) IO2 (Io4b) by sulfide species (S1, S2, S4, S5). The vertical line represents the pKa1 value for H2S. Data above the horizontal line at ΔlogK (ΔlogKreaction) = 0 indicate a favorable reaction and data below the horizontal line indicate an unfavorable reaction.

3.6 Iodate reduction by NH4+

Figure 6 shows the thermodynamics for the stepwise conversion of IO3 to IO2 by NH4+ where hydrazine (N2H4) and hydroxylamine (NH2OH) as well as their protonated forms could form as the first N intermediates. The thermodynamic calculations for these 2 electron transfers indicate that these reactions are not favorable. However, the reaction of the intermediates, if they could form by other processes, with IO3 to form N2 is very favorable. Thus, the IO3 to IO2 conversion is the controlling step in the reaction sequence with NH4+.

FIGURE 6
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Figure 6 Thermodynamics for the reduction of IO3 to IO2 (IO5b) by NH4+ to hydrazine species (N3a, N3b), hydroxylamine species (N4a, N4b) and by N2H5+ (N5) and NH2OH (N6). The vertical lines represent the pKa values for NH3OH+ (5.82) and N2H5+ (7.93), respectively. Data above the horizontal line at ΔlogK (ΔlogKreaction) = 0 indicate a favorable reaction and data below the horizontal line indicate an unfavorable reaction.

4 Results and discussion: Iodide oxidation

4.1 Iodide oxidation by NO3-, MnO2 and Fe(OH)3

Figures 36 showed the reduction of IO3 with reductants. All values of ΔlogKreaction > 0 indicate a favorable reaction and all values of ΔlogKreaction< 0 indicate an unfavorable reaction. These figures can be used to discuss the reverse reaction of I- oxidation with oxidants. For reverse reactions, when a ΔlogKreaction< 0, then I- oxidation is favorable, but when ΔlogKreaction > 0, I- oxidation is unfavorable.

In Figure 3, NO3 is not an oxidant for I- (reverse of the NO2 and I2 reaction) except for the formation of I2 at a pH< 4.

In Figure 4, MnO2 oxidizes I- to HOI (reverse of the Mn2+ and HOI reaction) at a pH< 3 and I- to I2 (reverse of the Mn2+ and I2 reaction) at a pH< 5. A couple of laboratory studies showed I- oxidation with synthetic birnessite (δ-MnO2). First, Fox et al. (2009) showed that I2 was produced over the pH range 4.50 – 6.25, and that IO3 formed in smaller amounts. The kinetics of the reaction were slower at higher pH by 1.5 log units (> 30-fold) and were slower when smaller amounts of MnO2 were added (Table 2). Allard et al. (2009) investigated the same reactants to a pH of 7.5 and found I2 and IO3 as products; above pH = 7 the reaction is very slow. Iodate was found mainly in lower pH waters. Both I2 and IO3 adsorb to the birnessite surface. Similar results have been found over the pH range 4-6 for Mn(III) solids (Szlamkowicz et al., 2022). These MnOx reactions with I- are much slower that the reactions with reactive oxygen species (Table 2). Nevertheless, these are important as Kennedy and Elderfield (1987a, 1987b) showed that the conversion of iodide to iodate occurred in marine sediments.

TABLE 2
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Table 2 Kinetic rate constants for the reaction of oxidants with Iodide and I2.

Figure 4 also shows that I- oxidation by Fe(OH)3 to I2 (reverse of the Fe2+ and I2 reaction) should occur only at a pH< 2.5. The I- to HOI conversion (reverse of the Fe2+ and HOI reaction) is favorable at pH ≤ 0.5.

4.2 Iodide oxidation by oxygen species

Figure 7 shows the thermodynamics of I- oxidation to I2 by oxygen species. Figure 7A shows that the one-electron process for I- oxidation with 3O2 is thermodynamically unfavorable over all pH whereas Figure 7B shows that the two-electron process is favorable at a pH< 3. Figure 7A shows that the successive 1-electron oxidations of I- where superoxide (O2) is reduced to hydrogen peroxide (H2O2), which is reduced to hydroxyl radical (•OH). Only •OH is thermodynamically favorable over the pH range considered. O2 and H2O2 show favorable reactions at pH< 9 and pH< 6, respectively.

FIGURE 7
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Figure 7 Thermodynamics for I2 formation via the oxidation of I- (Io1) with (A) the successive 1-electron oxidants 3O2 (O6), O2 (O7), H2O2 (O8) and •OH (O9b); (B) the 2-electron oxidants 3O2 (O2), 1O2 (O5), H2O2 (O3), and O3 (O4). Data above the horizontal line at ΔlogK (ΔlogKreaction) = 0 indicate a favorable reaction and data below the horizontal line indicate an unfavorable reaction.

By contrast, Figure 7B shows that the reactions of I- with the 2-electron oxidants H2O2, 1O2 and O3 are all thermodynamically favorable. The likely reaction pathway is the loss of 2-electrons to produce I+, which then reacts with I- to form I2. Note that H2O2 reacts to form H2O not •OH in Figure 7B. The 2-electron reaction with O3 (Figure 7B) is more favorable than the 1-electron reaction (Figure 7A).

Wong and Zhang (2008) showed that H2O2 oxidizes I- in artificial seawater from pH 7-9, which is consistent with Figure 7B. However, I- oxidation does not lead to iodate. In fact, I- reforms. They proposed that I2 formed and was reduced back to I-, but they did not provide a mechanism. The reverse reaction of I2 with •OH (ΔlogKreaction< 0 in the plot) is favorable to reform H2O2 and I- at pH > 6 whereas the reverse reaction of H2O2 with I2 to reform O2 and I- is favorable at a pH > 9 (ΔlogKreaction< 0 in the plot). These thermodynamic data indicate that H2O2 can form I2 in a 2-electron transfer (Figure 7B) and then reduce I2 to I- in a 1-electron transfer (Figure 7A).

At seawater pH, superoxide, O2, can oxidize I- to I2 and the reaction occurs with a rate constant of 108 M-1s-1 (Bielski et al., 1985; Table 2). Because I2 is a good electron acceptor, the subsequent reaction of O2 with I2 leads to I2 (Schwarz and Bielski, 1986). As to be discussed in section 4.3, I2 reacts with organic matter to form organo-iodine compounds. Extracellular O2 is generated by Roseobacter sp. AzwK-3b (Li et al., 2014) and results in the oxidation of Mn2+ to Mn(III,IV) oxides. However, Li et al. (2014) found that O2 also oxidized I-. Considering that extracellular O2 formation is a widespread phenomenon among marine and terrestrial bacteria, this could represent an important first step in the pathway for iodide oxidation in some environments. The Mn oxides formed by Roseobacter sp. AzwK-3b are not the oxidant as MnO2 kinetics is slower (Table 2).

To obtain IO3, further oxidation of I2 to HOI must occur, and •OH is one candidate with a rate constant of 1.2 x 1010 (Buxton et al., 1988; Table 2). Also, O3 has a rate constant of 1.2 x 109 (Liu et al., 2001).

I2 is a prominent intermediate in I- oxidation yet HOI is needed to form IO3. HOI can form directly from I- and I2 oxidation or from hydrolysis of I2 (reverse of eqn. 5), which is fast at basic pH (Wong, 1991). Figure 8A shows that of the successive 1-electron oxidants (starting from O2) for I2 oxidation, only •OH is thermodynamically favorable over all pH to form HOI whereas O3 is favorable at pH > 6, and O2 is favorable at pH< 6. H2O2 as a 1-electron oxidant cannot oxidize I2 to form HOI, but H2O2 can reduce HOI to I2 (reverse of the O2 and I2 reaction). Figure 8B indicates that, as 2-electron oxidants, H2O2 and O3 oxidation can lead to HOI formation. Comparing Δlog K values in Figures 7, 8 indicates that oxidation of I2 to HOI is less favorable than the oxidation of I- to I2.

FIGURE 8
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Figure 8 Thermodynamics for the formation of HOI via the oxidation of I2 (Io3) with (A) the successive 1-electron oxidants 3O2 (O6), O2 (O7), H2O2 (O8), •OH (O9b) and O3 (O11); (B) the 2-electron oxidants 3O2 (O2), 1O2 (O5), H2O2 (O3), and O3 (O4). Data above the horizontal line at ΔlogK (ΔlogKreaction) = 0 indicate a favorable reaction and data below the horizontal line indicate an unfavorable reaction.

These data also indicate why the comproportionation reaction of HOI with I- to form I2 can occur (eqn. 5, Carpenter et al., 2013).

H++HOI+II2+H2O(5)

Although disproportionation of HOI to IO3 and I- (eqn. 6) is fast in strongly basic solution, it is not detectable at seawater pH (Wong, 1991).

3HOI+3OHIO3+2I+3H2O(6)

Figure 9 shows the successive 2-electron oxidation reactions of I-, HOI and IO2 with 3O2, 1O2, H2O2 and O3. 3O2 cannot affect the oxidation at any pH. Figure 9 shows that O3 oxidation reactions with I-, HOI and IO2 are favorable; thus, O3 can affect the complete oxidation of I- to IO3. Also, the H2O2 oxidation reactions of I-, HOI and IO2 are favorable and can lead to IO3 formation; however, the kinetics of H2O2 oxidation can be slow. Haloperoxidase enzymes from organisms enhance the kinetics (Butler and Sandy, 2009) as does the reaction of H2O2 with carboxylic acids secreted by microbes to form peroxy carboxylic acids, which in turn oxidize I- to I2 (Li et al., 2012). The reactive oxygen species 1O2 can oxidize I- at pH< 10, oxidize HOI at pH > 5, and IO2 over all pH. Thus, 1O2 can be an oxidant of I- to IO3 at seawater pH. These data indicate that HOI oxidation leads to IO3 formation.

FIGURE 9
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Figure 9 Thermodynamics for the reaction of 3O2 (O2), 1O2 (O5), H2O2 (O3), and O3 (O4) as 2-electron oxidants with (A) I- to form HOI (Io2), (B) HOI to form IO2 (Io4b) and (C) IO2 to form IO3 (Io5b). The vertical lines represent the pKa value of 4.49 for HO2I dissociation to IO2. Data above the horizontal line at ΔlogK (ΔlogKreaction) = 0 indicate a favorable reaction and data below the horizontal line indicate an unfavorable reaction.

Interestingly, ΔlogK values in Figure 9A show that the thermodynamics of I- oxidation by the 2-electron oxidants O3 and H2O2 to form HOI is slightly less favorable than I2 formation (Figure 7B). Conversely, thermodynamics of I- oxidation by H2O2 as a 2-electron oxidant to form HOI (Figure 9A) is more favorable than I2 formation (Figure 7A).

As shown in Figure 10, a potentially potent oxidant for I- is N2O, which is an O atom transfer oxidant like O3. However, the N2O concentration in seawater is minor, but the largest reported values are 90 and 250 nmol kg-1 for the OMZs of the Arabian Sea (Freing et al., 2012) and the Eastern Tropical North Pacific (Damgaard et al., 2020), respectively. These values are smaller than the total iodine concentration in seawater. The N2O concentration in the atmosphere is 335 ppbv (August 2022, https://www.n2olevels.org), which is equivalent to 0.0331 Pa or 7.8 nM dissolved in surface seawater (salinity of 35) at 20 0C using the solubility data from Weiss and Price (1980).

FIGURE 10
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Figure 10 Thermodynamics for the 1- and 2- electron reductions of IO3 (Io5a, Io5b, Io6), IO2 (Io4b), HOI (Io2, Io3), and I2 (Io1) by N2O (N2). Data above the horizontal line at ΔlogK (ΔlogKreaction) = 0 indicate a favorable reaction and data below the horizontal line indicate an unfavorable reaction.

4.3 ROS in seawater

Reactive oxygen species exist in marine waters, but at low concentrations. O3 penetrates a few micrometers through the water-air interface at surface iodide concentrations (Carpenter et al., 2013). Powers and Miller (2014) showed that solar-induced processes with organic matter in freshwater and seawater are a major source of ROS (as O2, H2O2, and •OH) with the inventory and production rates for H2O2 in surface seawater being highest of the ROS. Also, Sutherland et al. (2020) report that dark, extracellular O2 production is prolific among marine heterotrophic bacteria, cyanobacteria, and eukaryotes. In surface ocean waters, the concentration of H2O2 ranges from 20 - 80 nM (Yuan and Shiller, 2001), biological O2 production gives a total concentration of ~ 0.07 to 0.30 nM (Sutherland et al., 2020), •OH concentration is ~10-18 M (Mopper and Zhou, 1990), and 1O2 concentration ranges from 10-13 to 10-14 M (Sunday et al., 2020). However, these ROS concentrations are typically smaller than I- concentrations, which range from 10 to 200 nM (Chance et al., 2019). Thus, I- oxidation in seawater samples should be difficult to observe experimentally. Hardisty et al. (2020) tracked the addition of stable isotopes of iodide in sample incubations and report the rate of I- oxidation to be 118–189 nM yr-1, which is similar to rates reported by mass balance approaches (Campos et al., 1996a; Truesdale et al., 2001; Žic and Branica, 2006; Žic et al., 2008). Hardisty et al. (2020) report that the product is likely HOI that results in the formation of organic-iodine compounds (see section 4.6) which on decomposition can release I-.

As the surface concentrations of ROS are smaller than the I- concentration, the question is how does I- get oxidized to IO3 in seawater? Microbial processes and the oxidation of I species in the atmosphere by ROS are likely candidates. These are now discussed.

4.4 Iodide oxidation in brown kelp

Brown kelp are the strongest accumulators of iodine as I- among living organisms (up to 100 mM, Küpper et al., 2008). The element iodine was discovered by the formation of I2 during exposure of brown kelp to concentrated sulfuric acid, which oxidized I- to I2. Kelp releases I- on the thallus surface and in the apoplast when undergoing oxidative stress during the partial emersion of the brown kelp forest at low tide; e.g., by exposure to high irradiance, desiccation, and atmospheric O3. Kelp contain vanadium haloperoxidases (Colin et al., 2003; Küpper et al., 2008) that enhance I- oxidation by H2O2. Whereas the nonenzymatic reaction of I- with H2O2 is slow, the reactions with O3, O2, 1O2, and •OH are very fast (> 108 M-1s-1, Table 2); they are also faster than MnO2 oxidation of I-. Küpper et al. (2008) consider I- as the simplest antioxidant known.

4.5 Iodide oxidation by microbes

Hughes et al. (2021) report that IO3 production occurs in cultures of the ammonia-oxidizing bacteria Nitrosomonas sp. and Nitrosococcus oceani supplied with I-, but not in cultures of three different nitrite oxidizing bacteria. Information on the enzymes mediating the oxidation were not studied. Nevertheless, NH4+ oxidation via nitrification occurs via NH2OH formation which is a 2-electron reaction. Further reaction of NH2OH via metalloenzymes (e.g., Mo and W oxidases that transfer O atoms) leads to NO2 (a 4-electron transfer) and NO3. I- likely goes through the intermediates IO2 and HOI to form IO3, but these intermediates are reactive and not detectable by present analytical methods unlike NO2.

Consistent with the Hughes et al. (2021) report, Kennedy and Elderfield (1987a, 1987b) showed that the conversion of iodide to iodate occurred in marine sediments. Microbial intervention is likely, but reaction of I- with oxidized Mn is possible depending on the pH.

Amachi and Iino (2022) reviewed the genus Iodidimonas, which was originally found in brines, but was also cultured from seawater enriched with I-. I2 is the first oxidation product. Iodidimonas contains the iodide oxidizing enzyme (IOX), which is an extracellular protein that contains multicopper oxidases. Iodidimonas requires O2, not H2O2, as the electron acceptor. Other oxidants are required to oxidize I2 to IO3 with O3 and •OH being the most effective.

4.6 HOI and I2 formation leads to organic iodine

Competing with the inorganic interconversion between iodide and iodate is the formation of organic iodine compounds. Formation of C-I bonds can occur during the reduction of IO3 and oxidation of I-. Complete reduction of IO3 to I- does not need to occur intercellularly and can lead to HOI and I2 formation as in Figure 4. The first step in I- oxidation also leads to I2 and HOI. I2 is neutral and adds to organic compounds such as olefins, which are not very reactive in seawater, whereas I+ in HOI reacts with α-keto compounds and peptides through keto-enol isomerization (Truesdale and Luther, 1995). Both I2 and HOI lead to volatile and nonvolatile organic-iodine (R-I) compounds with C-I or N-I bonds, and Harvey (1980) showed that N-iodo amides were the main organic iodine components in marine sediments. On decay of organic compounds, the C(N)-I bond breaks leading to I- release, which mimics the senescence pathway outlined by Bluhm et al. (2010) and Hepach et al. (2020). Recently, Ooki et al. (2022) showed that CH3I and CH3CH2I formed in sediments from polar and subpolar seas and was related to increased phytodetritus at the seafloor after the spring bloom.

Allard and Gallard (2013) showed that the oxidation of I- by birnessite in the presence of organic matter also led to CH3I over the pH range 4-5.

As total iodine in surface ocean waters is lower by a few percent compared to deep waters (Wong, 1991), the decomposition of organic-iodine leads to some I- release, which may be oxidized to IO3 by ammonia-oxidizing bacteria (Hughes et al., 2021). This is similar to release and oxidation of NH4+ to NO3 from particulate organic matter in deep waters that results in an increase of NO3 concentration with depth (recycled element profile). Deep waters contain mainly IO3, so not much I- is released to the deep-water column by in situ water column processes, and most organic-iodine gets to the sediments where it is released as I- (Kennedy and Elderfield, 1987a, Kennedy and Elderfields 1987b; Luther et al., 1995). Kennedy and Elderfield (1987a, 1987b) and Shimmield and Pedersen (1990) report that the molar I/C ratio in planktonic organisms is 10-4 whereas it is typically >10-3 in sediments. Decomposition of sedimentary organic-I releases I- to porewaters and the overlying water column where it can be transported hundreds of kilometers offshore along isopycnal surfaces in OMZs (Farrenkopf and Luther, 2002; Cutter et al., 2018).

4.7 Surface seawater and atmospheric formation of IO3-, and iodine speciation in the atmosphere

There is significant literature showing that coastal and oceanic regions are sources of iodine emissions to the atmosphere, and I note some important aspects of this air-sea connection. I- reacts with O3 to form IO-, which at seawater pH forms HOI. Carpenter et al. (2013) showed that this reaction occurs in the first few micrometers below the air-water interface and that HOI is ten-fold greater than I2 above the sea surface. HOI contributes 75% of the observed iodine oxide aerosol levels over the tropical Atlantic Ocean, and these iodine emissions to the atmosphere have increased 3-fold over the last century due to the increase in anthropogenic O3 (Carpenter et al., 2021). O3 reacts stepwise with this gaseous HOI (IO-) and gaseous IO2 to form IO3, which can attach to aerosols.

Formation and release of gaseous I2 from seawater to air permits photochemical breaking of the I-I bond to form gaseous I atoms, •I, which are reactive radicals. Similarly, release of volatile organic-iodine compounds leads to the homolytic cleavage of the C-I bond to form I•. O3 reacts readily with I• to form gaseous IO• in the marine boundary layer (Whalley et al., 2010). Further stepwise oxidation of gaseous IO•/HOI leads to IO3. In laboratory experiments using mass spectrometry detection, Teiwes et al. (2019) showed that hydrated iodide, I(H2O)-, reacts with gaseous O3 to form IO2 directly without formation of gaseous HOI or IO-; thus, HIO3/IO3 can form in a two-step reaction sequence in the atmosphere.

Using mass spectrometry to evaluate atmospheric IxOy cluster and (nano)particle formation above seabed macroalgae, Sipilä et al. (2016) showed the stepwise formation of HIO3 via HOI and IO•, which leads to (I2O5)x clusters (x=2-5) containing HIO3 that result in iodine rich aerosol particles. These data on the formation of I2O5 aerosols agree with the exothermic ΔHreaction values of iodine oxide species reacting with O3 and each other calculated using quantum mechanics (Kaltsoyannis and Plane, 2008). Sipilä et al. (2016) also showed that cluster formation increased as a burst at low tide indicating significant I2 release from the macroalgae (and subsequent oxidation) as found by Küpper et al. (2008). Hydration of I2O5 leads to two IO3. In mass spectrometry laboratory studies, Martín et al. (2022) showed that new iodine containing (nano)particles and IO3 also form in the presence of NO3 and provide ΔHreaction data for the gas phase reactions involved. Experiments using the CERN CLOUD (Cosmics Leaving Outdoor Droplets) chamber documented the formation of HIO3 via iodooxy hypoiodite, IOIO, as an intermediate (Finkenzeller et al., 2022) and the fast growth of HIO3 as (nano)particles (He et al., 2021).

In recent atmospheric campaigns, Koenig et al. (2020) showed that IO3 is the main iodine reservoir as it forms on aerosols in the stratosphere with iodine being responsible for 32% of the halogen induced O3 loss. Cuevas et al. (2022) also showed that iodine can dominate (∼73%) the halogen-mediated lower stratospheric ozone loss during summer and early fall, when the heterogeneous reactivation of inorganic chlorine and bromine reservoirs is reduced.

The information in the preceding paragraphs along with the thermodynamic data from Martín et al. (2022), Figure 2 (the half reaction for O3 to O2 and H2O) and Figure 9 predict that IO3 should be the dominant species in the atmosphere. Although reduction of IO3 is not predicted in an oxidizing atmosphere, analyses of rainwater (Campos et al., 1996b; Truesdale and Jones, 1996; Baker et al., 2001; Hou et al., 2009), aerosols (Gilfedder et al., 2008; Droste et al., 2021) and snow (Gilfedder et al., 2008) in the marine boundary layer indicate that aqueous iodide and iodate coexist. Hou et al. (2009) reviewed wet iodine speciation data and reported that IO3 predominates over I- from marine sources/air masses whereas I- predominates from continental air masses.

There are several ways that I- (or reduced I) can form in rainwater and aerosols. The interconversion between IO3 and I- at the pH of wet deposition also leads to HOI and I2, which can react with organic material forming C-I bonds that can release I- (section 4.6). This material has been given the term soluble organically bound iodine and can be larger than the sum of the concentrations of IO3 and I- in aerosols (Gilfedder et al., 2008; Droste et al., 2021). Soluble organically bound iodine can form from release of natural organic iodine from land and sea (a primary source) or from the reaction of natural organic material with HOI or I2 in the atmosphere (a secondary source). On photolysis of C-I, I• forms and reacts with O3, and on C-I reaction with nucleophiles, I- forms. During a study on the formation of cloud condensation nuclei, Huang et al. (2022) also showed that natural gaseous organic material in the marine boundary layer reacts with IO3 in aerosols resulting in gaseous I2, which can be reoxidized to IO3 (catalysis) or react to form organic-I compounds. Lastly, Cuevas et al. (2022) reported that photolysis of IO3 particles in the stratosphere at a wavelength of about 260 nm can lead to gaseous I• and O2 during transport from the tropics to the Antarctic region. Thus, there are several pathways for reduction of IO3 in the atmosphere.

5 Conclusions

The reduction of IO3 to I- in solution is a facile process by biotic and abiotic reactions. The intermediates IO2 and HOI dictate the reactivity sequence via a combination of thermodynamic and kinetic considerations. The IO3 to IO2 conversion is the least favorable and likely controlling step in this reaction sequence, but there is no need for nitrate reductase for IO3 reduction based on numerous studies. The data from this study indicate that once IO2 forms there is no thermodynamic barrier to I- formation. Chemical reduction of all iodine species (not iodide) by sulfide, Fe2+ and Mn2+ are favorable at seawater and sedimentary pH values, but only sulfide has been studied in the laboratory at oceanic pH values. Dissimilatory IO3 reduction during organic matter decomposition seems to be a key process as the IO3/IO2 couple is more favorable than the NO3/NO2 couple.

However, the oxidation of I- back to IO3 via 3O2 has a major thermodynamic barrier in solution, and the disproportionation of HOI at seawater pH values is not measurable. Thus, ROS, oxidized Mn and microbes are important for I- oxidation to IO3 due to favorable thermodynamics and kinetics (Table 2). Recent reports of microbial oxidation have not documented the entire six-electron oxidation in a stepwise manner so further work on this topic is necessary. Oxidation of I- by oxidized Mn is a pH dependent reaction and less likely at seawater pH values but could occur in sedimentary environments. The reactions of O3 and •OH with iodine species (not IO3) are thermodynamically favorable over all pH. However, ROS are not normally in significant concentration in seawater to influence IO3 formation. Notable exceptions are for (1) sea surface microlayer, which adsorbs atmospheric O3, and (2) the reaction of Fe2+ with O2 that leads to Fenton chemistry with •OH production. Systems where Fenton chemistry can occur are at/near hydrothermal vents (Shaw et al., 2021), submarine groundwaters (Burns et al., 2010), and sediments or water columns where O2 and Fe2+ concentration profiles overlap including ancient earth (Chan et al., 2016).

I- is a major sink for O3 in the sea surface microlayer and the atmosphere. IO3 formation in the atmosphere and IO3 redeposition to surface seawater may be major iodine processes with the latter being similar to the deposition of trace metals from wet and dry deposition to the surface ocean (e.g., Chance et al., 2015; Meskhidze et al., 2019). Most atmospheric iodine originates from marine sources where I- oxidation to I2 and homolytic cleavage of C-I bonds occurs; thus, gaseous iodine emissions from the ocean are reduced. IO3 forms from these sources during oxidation by O3 in the atmosphere. An estimate of atmospheric deposition of IO3 to the ocean surface could be made by using the amount of IO3 in rainwater and aerosols that would be returned to the ocean surface, but more information on iodine speciation in rainwater and aerosols is needed as global spatial coverage appears limited. Despite major advances in iodine geochemistry over the last two decades, significant research is still needed on the processes that affect I- oxidation to IO3 in the atmosphere, seawater and ocean sediments.

Author contributions

The author confirms being the sole contributor of this work and has approved it for publication.

Acknowledgments

The author thanks NSF for funding his group’s research on iodine marine chemistry over his career, and Thomas Church and Timothy Ferdelman for suggesting the author’s initial foray into iodine chemistry. The author thanks the reviewers and guest editor, Rosie Chance, for their comments and constructive suggestions to improve the manuscript.

Conflict of interest

The author declares that the research was conducted in the absence of any commercial or financial relationships that could be construed as a potential conflict of interest.

Publisher’s note

All claims expressed in this article are solely those of the authors and do not necessarily represent those of their affiliated organizations, or those of the publisher, the editors and the reviewers. Any product that may be evaluated in this article, or claim that may be made by its manufacturer, is not guaranteed or endorsed by the publisher.

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Keywords: iodate, iodide, iodine intermediates, thermodynamics, oxidation, reduction

Citation: Luther GW III (2023) Review on the physical chemistry of iodine transformations in the oceans. Front. Mar. Sci. 10:1085618. doi: 10.3389/fmars.2023.1085618

Received: 31 October 2022; Accepted: 03 January 2023;
Published: 15 February 2023.

Edited by:

Rosie Chance, University of York, United Kingdom

Reviewed by:

Lucy J. Carpenter, University of York, United Kingdom
Alex Baker, University of East Anglia, United Kingdom

Copyright © 2023 Luther. This is an open-access article distributed under the terms of the Creative Commons Attribution License (CC BY). The use, distribution or reproduction in other forums is permitted, provided the original author(s) and the copyright owner(s) are credited and that the original publication in this journal is cited, in accordance with accepted academic practice. No use, distribution or reproduction is permitted which does not comply with these terms.

*Correspondence: George W. Luther III, luther@udel.edu

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