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 IO 2 − are from Schmitz (2008).

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.

The standard state ΔG° for the reaction = - 490.68 kJ/2 moles H₂O or 4 moles of electrons. The equilibrium constant ( K 4 0 ) is given in eqn. 2a where {} indicates activity for each chemical species and the activity of H₂O is defined as 1.

On expanding, eqn. 2b results.

log K 4 0 is for 4 mole of electrons or 21.50 for 1 mole of electrons.

For a one-electron half-reaction, we have

And equation 2b becomes equation 2c

or

From the Nernst Equation, pε o = 1/4  log K 4 0 = 21.50 (the standard state value), which on substitution gives equation O1 (see Table 1 ) where concentration is used for O_2(aq).

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

and at a pH of 8, pε = 12.58.

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

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.

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.

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 IO 3 − (Io5b) by NO 2 − (N1) in equation 3.

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

At a pH of 7, the pε_red values for IO 3 − and NO 3 − are 7.91 and 7.28, respectively. As NO 2 − is the reductant, it is oxidized; thus, the sign for pε_red (7.28) is reversed to become pε_oxid (-7.28).

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.

In the oxic environment, the oxidizing condition of the environment or pε is set by the 4-electron transfer reaction of the O_2(aq)/H₂O couple [reaction O1 in Table 1 ]. At a pH of 8, temperature 25⁰C and a salinity of 35, 100% O_2(aq) saturation is 211 μM, which gives a pε of 12.58 ( Figure 1 ). As the IO 3 − / I − couple has a pε of 10.56 at pH = 8 , IO 3 − is the thermodynamically favored iodine species.

Entering the pε value for a given [O_2(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 [O_2(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 [O_2(aq)] concentrations in the 1-100 nM range (Revsbech et al., 2009; Lehner et al., 2015), calculations show that IO 3 − is the thermodynamically preferred iodine species even at 1 nM O_2(aq), which gives a pε of 11.25 for the O_2(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 [O_2(aq)] concentrations ≤ 1 μM, IO 3 − , NO 3 − 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 NO 3 − / NO 2 − ( N 1 ) , 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 NO 2 − (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 O_2(aq), the NO 3 − → NO 2 − ( N 1 ) and MnO₂ → Mn²⁺ (Mn1) couples can be chosen to set the environmental pε. At pH = 8, the NO 3 − to NO 2 − pε is 6.15 and the MnO₂ to Mn²⁺ pε is 4.80. At pH = 7, the NO 3 − to NO 2 − pε is 7.28 and the MnO₂ to Mn²⁺ pε is 6.80. At these pε values, O_2(aq) is below 1 nM, and I^- is now the thermodynamically favored iodine species when comparing these data with the IO 3 − / I − couple (pε of 10.56 at pH = 8).

Figure 2 also shows that the IO 3 − / IO 2 − couple (Io5b) should be the first step in the reaction sequence of iodate to iodide (eqn. 1a). As the pε of the IO 3 − / IO 2 − couple has a more positive pε value than the N, Mn and Fe couples in Figure 2 , IO 3 − reduction is more favorable than these couples even though it is very close to the NO 3 − / NO 2 − couple. Thus, IO 3 − is predicted to reduce before NO 3 − , and biological activity (e.g., nitrate reductase activity) is not necessary to reduce IO 3 − (see section 3.2). Because of the strong pH dependence for the Mn and Fe couples, they cross the IO 3 − / IO 2 − and NO 3 − / NO 2 − couples at lower pH, which have similar slopes. Thus, NO 2 − is predicted to reduce IO 3 − to IO 2 − (eqn. 3). Interestingly, Mn²⁺ and Fe²⁺ should be poorer reductants than NO 2 − for conversion of IO 3 − to IO 2 − at a pH< 6 and pH< 1, respectively, but are more favorable to reduce IO 3 − than NO 2 − above those pH values (see sections 3.2 and 3.3).

Although the IO 3 − to I^- conversion occurs at higher pε, it is a 6-electron transfer (IO6), which is not a facile process. Thus, the intermediates ( IO 2 − and HOI) will dictate the reactivity sequence via a combination of thermodynamic and kinetic considerations.

As shown in Figure 2 , IO 2 − reduction to HOI and HOI reduction to I^- are also more favorable at higher pε values than the IO 3 − to I^- couple. At a pH = 8, the IO 2 − 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 IO 2 − 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 IO 2 − and HOI are not predicted to be stable in marine waters; thus, the conversion of IO 3 − to IO 2 − is a key step. Interestingly, Hardisty et al. (2021) found in situ IO 3 − reduction in the oxycline where [O_2(aq)] was 11 μM, but not at [O_2(aq)]< 2 μM. Lastly, the O₃ to O₂ + H₂O couple is highly oxidizing indicating that all iodine couples should lead to IO 3 − 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).

Although NO 2 − has not yet been shown to be a reductant for IO 3 − 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 IO 3 − to I^- by NO 2 − reduction ( NO 2 − oxidizes to NO 3 − ). 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 IO 3 − to IO 2 − reaction whereas the IO 2 − to HOI and HOI to I^- reactions are more favorable. Thus, the IO 3 − to IO 2 − 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 IO 3 − to I^- by NO 2 − is predicted via 1-electron or 2-electron transfer reactions at seawater pH values. The data plotted in Figure 2 indicate that once IO 2 − forms there is no thermodynamic barrier to I^- formation.

The IO 3 − reaction with NO 2 − 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 IO 3 − reduction. At seawater pH, the reaction seems to be hindered by kinetics in the transition state as each reactant ( IO 3 − and NO 2 − ) is an anion, which will repel each other.

The marine literature has many reports on the uptake of IO 3 − (with or without NO 3 − ) 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, NO 3 − reductase was presumed by some researchers to be a key process for IO 3 − 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 IO 3 − uptake and I^- release due to senescence.

NO 3 − reductase appears to reduce IO 3 − 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 IO 3 − . Thus, there is no need for nitrate reductase for IO 3 − reduction as I^- can be up taken by some phytoplankton rather than form from IO 3 − reduction. Moreover, Waite and Truesdale (2003) showed that nitrate reductase was not important for IO 3 − reduction by Isochrysis galbana. The latter study is consistent with the thermodynamics of the reduction IO 3 − to IO 2 − being more favorable than the reduction NO 3 − to NO 2 − .

Furthermore, under anaerobic conditions, dissimilatory IO 3 − 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 IO 3 − 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 IO 3 − 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 IO 3 − reduction is a major biogeochemical process. Thus, nitrate reductase (also an O atom transfer reaction) is not a requirement for bacterial IO 3 − reduction to I^-.

The interconversion of dimethylsulfoxide with dimethylsulfide during dissimilatory IO 3 − reduction is another 2-electron O-atom transfer reaction. Moreover, the reactions of DMS to reduce HOI, IO 2 − and IO 3 − are thermodynamically favorable ( Figure 3B , DMSO reduction is in S3, Table 1 and occurs at a lower pε than NO 3 − and IO 3 − 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.

Figure 4 shows the thermodynamics for the stepwise conversion of IO 3 − 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 IO 2 − to HOI reaction is favorable over the entire pH range. The IO 3 − to IO 2 − reaction is favorable only at pH > 6 whereas the other reactions are favorable at pH > 3. Thus, the IO 3 − to IO 2 − 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 IO 3 − with Mn²⁺ under suboxic conditions. The reaction has not been investigated in laboratory studies.

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 IO 3 − 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 IO 3 − (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 IO 3 − to IO 2 − conversion is the least favorable and likely controlling step in this reaction sequence.

Comparing Figures 3 , 4 indicates that the Mn²⁺ and NO 2 − reactions with iodine species have a similar range of ΔlogK_reaction values whereas the Fe²⁺ reactions with iodine species are more favorable (higher ΔlogK_reaction values).

In sulfidic waters and porewaters, IO 3 − 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 IO 3 − to IO 2 − 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 SO 4 2 − . 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 IO 2 − to HOI is more favorable as it has the larger ΔlogK_reaction values. Again, the IO 3 − to IO 2 − conversion is the least favorable and likely controlling step in this reaction sequence.

Figure 6 shows the thermodynamics for the stepwise conversion of IO 3 − to IO 2 − by NH 4 + 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 IO 3 − to form N₂ is very favorable. Thus, the IO 3 − to IO 2 − conversion is the controlling step in the reaction sequence with NH 4 + .

Figures 3 – 6 showed the reduction of IO 3 − with reductants. All values of ΔlogK_reaction > 0 indicate a favorable reaction and all values of ΔlogK_reaction< 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 ΔlogK_reaction< 0, then I^- oxidation is favorable, but when ΔlogK_reaction > 0, I^- oxidation is unfavorable.

In Figure 3 , NO 3 − is not an oxidant for I^- (reverse of the NO 2 − 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 IO 3 − 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 IO 3 − as products; above pH = 7 the reaction is very slow. Iodate was found mainly in lower pH waters. Both I₂ and IO 3 − 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 MnO_x 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.

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.

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 ( O 2 − ) is reduced to hydrogen peroxide (H₂O₂), which is reduced to hydroxyl radical (•OH). Only •OH is thermodynamically favorable over the pH range considered. O 2 − and H₂O₂ show favorable reactions at pH< 9 and pH< 6, respectively.

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 (ΔlogK_reaction< 0 in the plot) is favorable to reform H₂O₂ and I^- at pH > 6 whereas the reverse reaction of H₂O₂ with I₂ to reform O 2 − and I^- is favorable at a pH > 9 (ΔlogK_reaction< 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, O 2 − , 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 O 2 − with I₂ leads to I 2 − (Schwarz and Bielski, 1986). As to be discussed in section 4.3, I₂ reacts with organic matter to form organo-iodine compounds. Extracellular O 2 − 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 O 2 − also oxidized I^-. Considering that extracellular O 2 − 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 IO 3 − , 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 IO 3 − . 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 O 2 − 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 O 2 − 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.

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

Although disproportionation of HOI to IO 3 − and I^- (eqn. 6) is fast in strongly basic solution, it is not detectable at seawater pH (Wong, 1991).

Figure 9 shows the successive 2-electron oxidation reactions of I^-, HOI and IO 2 − 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 IO 2 − are favorable; thus, O₃ can affect the complete oxidation of I^- to IO 3 − . Also, the H₂O₂ oxidation reactions of I^-, HOI and IO 2 − are favorable and can lead to IO 3 − 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 IO 2 − over all pH. Thus, ¹O₂ can be an oxidant of I^- to IO 3 − at seawater pH. These data indicate that HOI oxidation leads to IO 3 − formation.

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).

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 O 2 − , 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 O 2 − 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 O 2 − 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 I O 3 − in seawater? Microbial processes and the oxidation of I species in the atmosphere by ROS are likely candidates. These are now discussed.

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, O 2 − , 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.

Hughes et al. (2021) report that IO 3 − 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, NH 4 + 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 NO 2 − (a 4-electron transfer) and NO 3 − . I^- likely goes through the intermediates IO 2 − and HOI to form IO 3 − , but these intermediates are reactive and not detectable by present analytical methods unlike NO 2 − .

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 IO 3 − with O₃ and •OH being the most effective.

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 IO 3 − and oxidation of I^-. Complete reduction of IO 3 − 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 IO 3 − by ammonia-oxidizing bacteria (Hughes et al., 2021). This is similar to release and oxidation of NH 4 + to NO 3 − from particulate organic matter in deep waters that results in an increase of NO 3 − concentration with depth (recycled element profile). Deep waters contain mainly IO 3 − , 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).

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 IO 2 − to form IO 3 − , 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 IO 3 − . In laboratory experiments using mass spectrometry detection, Teiwes et al. (2019) showed that hydrated iodide, I(H₂O)^-, reacts with gaseous O₃ to form IO 2 − directly without formation of gaseous HOI or IO^-; thus, HIO 3 / IO 3 − can form in a two-step reaction sequence in the atmosphere.

Using mass spectrometry to evaluate atmospheric I_xO_y 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 ΔH_reaction 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 IO 3 − . In mass spectrometry laboratory studies, Martín et al. (2022) showed that new iodine containing (nano)particles and IO 3 − also form in the presence of NO 3 − and provide ΔH_reaction 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 IO 3 − 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 IO 3 − should be the dominant species in the atmosphere. Although reduction of IO 3 − 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 IO 3 − 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 IO 3 − 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 IO 3 − 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 IO 3 − in aerosols resulting in gaseous I₂, which can be reoxidized to IO 3 − (catalysis) or react to form organic-I compounds. Lastly, Cuevas et al. (2022) reported that photolysis of IO 3 − 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 IO 3 − in the atmosphere.