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Logo of nihpaAbout Author manuscriptsSubmit a manuscriptHHS Public Access; Author Manuscript; Accepted for publication in peer reviewed journal;
Nat Geosci. Author manuscript; available in PMC 2011 January 1.
Published in final edited form as:
Nat Geosci. 2010 July; 3(7): 473–476.
doi:  10.1038/ngeo892
PMCID: PMC2902198

Photolytic degradation of methylmercury enhanced by binding to natural organic ligands


Monomethylmercury is a neurotoxin that poses significant risks to human health1 due to its bioaccumulation in food webs. Sunlight degradation to inorganic mercury is an important component of the mercury cycle that maintains methylmercury at low concentrations in natural waters. Rates of photodecomposition, however, can vary drastically between surface waters25 for reasons that are largely unknown. Here, we show that photodegradation occurs through singlet oxygen, a highly reactive form of dissolved oxygen generated by sunlight irradiation of dissolved natural organic matter. The kinetics of degradation, however, depended on water constituents that bind methylmercury cations. Relatively fast degradation rates (similar to observations in freshwater lakes) applied only to methylmercury species bound to organic sulfur-containing thiol ligands such as glutathione, mercaptoacetate, and humics. In contrast, methylmercury-chloride complexes, which are dominant in marine systems, were unreactive. Binding by thiols lowered the excitation energy of the carbon-mercury bond on the methylmercury molecule67 and subsequently increased reactivity towards bond breakage and decomposition. Our results explain methylmercury photodecomposition rates that are relatively rapid in freshwater lakes24 and slow in marine waters5.

Demethylation of methylmercury (MeHg) in natural waters occurs through both biological89 and photochemical processes3. In sunlit waters, the photochemical process is the dominant removal pathway and has been estimated to account for as much as 80% of MeHg flux out of lakes3,10. While sunlight is known to induce MeHg degradation, rates vary among different types of water. Photodegradation is rapid in remote lakes in North America and the Arctic24, but relatively slow in seawater, especially compared to microbial processes5. The reason for such variability is unknown because the photodemethylation process is poorly understood.

MeHg degradation is likely to be induced by the ultraviolet spectrum (<400 nm) of sunlight2. Hydroxyl radicals (OH), which are produced via photo-Fenton reactions or by photolysis of nitrate, have been proposed as the reactive intermediates responsible for photodegradation1112. However, with previously reported OH demethylation rates (k ~2 to 9 × 109 M−1 s−1)1113, significant MeHg degradation by this pathway would require large steady state OH concentrations that would occur only under specific conditions: agriculturally-impacted waters (with high nitrate and low dissolved organic carbon1314) or in acidic waters (pH ≤ 6) with sufficient soluble Fe to drive photo-Fenton processes15. While production of OH is possible at neutral pH16, production rates are highly sensitive to pH shifts, dissolved natural organic matter (NOM) content, and filtration of samples (in which removal of particulate Fe-oxides would decrease OH production). Such patterns are inconsistent with previous field studies on MeHg photodegradation in which pH variations (from 6 to 7.5) and filtration of the sample did not change photodegradation rates24,10. Such discrepancies point to another photoreactive intermediate (in addition to OH) that can induce MeHg degradation in surface waters.

MeHg degradation can also occur through singlet oxygen (1O2)17, which is generated by sunlight sensitization of dissolved NOM1819. Here, we show that reaction rates between 1O2 and MeHgare sufficient to account for sunlight-induced degradation. However, these rates apply only to certain types of MeHg species in water: CH3Hg-thiol complexes with NOM.

We performed photodegradation experiments in a simulated freshwater containing MeHg, Suwannee River humic acid, and a phosphate buffer (pH 7 to 7.4). In two separate experiments, sample containers (Teflon FEP) were exposed to natural sunlight over 4 days in October and December 2008 on a building roof top in Durham, North Carolina. During the exposure period, MeHg did not degrade in appreciable amounts in samples that contained only the phosphate buffer (Figure 1a). MeHg degradation was observed only when humic acid was present and when the MeHg concentration was low (15 nM) relative to humic acid. The differences observed between the 15 nM and 1500 nM MeHg treatments could be caused by complexation between MeHg and thiols associated with the humic. In the sample with 2 mg-C/L humic, the reduced-sulfur concentration from the humic acid was estimated at 150 nM (assuming 7.3×10−5 moles reduced-S per g C20) and greater than total MeHg (15 nM). In contrast, the sample with 1.1 mg-C/L humic and 1500 nM MeHg contained 80 nM reduced-S, less than MeHg concentration.

Figure 1
Concentration dependence of MeHg photodegradation rates

We performed similar photodegradation experiments using a laboratory UV-A reactor (λ=365 nm) to show that degradation rates depended on the concentration of MeHg relative to humic acid. MeHg degradation occurred only if humic acid was present (Supplementary Figure S1). Furthermore, degradation was observed only at MeHg concentrations (0.5 nM and 125 nM) that were less than the estimated reduced-S content of the humic (200 nM) (Figure 1b). We observed similar trends demonstrating the concentration dependence for MeHg degradation in samples formulated in seawater (Supplementary Figure S1). The UV-B component of the solar spectrum may also contribute to photodegradation2 but was not tested in our study.

Sunlight irradiation of chromophoric NOM is known to generate reactive intermediates such as OH, superoxide, and 1O221. In additional photodegradation experiments, we utilized selective probes to identify the reactive intermediate that was responsible for MeHg decomposition (Figure 2). The addition of β-carotene and sodium azide (NaN3), which are probes for 1O2, resulted in slower MeHg degradation kinetics (Figure 2a). Experiments were also repeated in 98.5% deuterated water (D2O), a solvent that is ~13 times slower than H2O at quenching singlet oxygen19,22 and thus, increases steady state 1O2 levels. The use of D2O appeared to increase the rate of decomposition; however, this enhancement was smaller than expected. The addition of isopropyl alcohol (a probe for OH23) and isoprene (a probe for photosensitized triplet state NOM2324) did not appreciably change photodegradation rates (Figure 2b). MeHg degradation was not observed in dark controls: the recovery of MeHg was 92% (± 8.3%) after 120 hr. Additional dark experiments with enzyme-generated superoxide indicated that superoxide was not capable of decomposing MeHg within this time frame (Supplementary Figure S2). Collectively these results suggested that 1O2 was the reactive intermediate responsible for photodegradation in our experiments and not other types of intermediates such as OH, photosensitized NOM, and superoxide.

Figure 2
MeHg degradation via generation of 1O2 from photosensitized humic acid

The experiments with β-carotene, NaN3, and D2O demonstrated the consequence of a heterogeneous distribution of 1O2 in humic acid solutions and the binding of MeHg to NOM. According to Latch and McNeill19, singlet oxygen is generated in the light-absorbing regions of dissolved NOM and quickly quenched by water. Therefore, the steady state 1O2 concentration can be orders of magnitude greater in close proximity to dissolved NOM molecules than in bulk solution19. In our experiments, this concentration gradient may explain why hydrophilic modifiers such as NaN3 and D2O had less effect on MeHg degradation rates than would be predicted based on known reaction rates between the modifiers and 1O2 (Supplementary Table S1). In contrast, β-carotene is a hydrophobic modifier that sorbed to NOM macromolecules and subsequently inhibited MeHg degradation to an extent that was predicted (Table S1). These results indicated that MeHg was sorbed to or complexed by humic molecules and, consequently, reacted with 1O2 molecules in the hydrophobic regions of the NOM.

In our experiments water composition was an important factor controlling degradation kinetics. Rates depended on the relative concentrations of MeHg and humic acid (Figure 1). Direct interactions between MeHg and NOM (such as metal-ligand complexation) may be critical to photodemethylation kinetics.

We tested the importance of complexation by quantifying second-order rate constants k1O2,CH3HgL for reactions between 1O2 and CH3Hg+-ligand complexes. Glutathione (GSH) and mercaptoacetate (MA) were selected as surrogates of thiol-containing ligands that are expected to bind to CH3Hg+ ions in NOM-containing freshwater2526. Cl and OH ligands are relevant for saline or low-NOM water. Singlet oxygen was generated by UV-A irradiation of rose bengal. Dark controls were also performed to ensure that the reactants did not cause degradation (Supplementary Figure S3). Additional experiments with amendments (D2O or NaN3) and with Eosin-Y (instead of rose bengal) were performed to demonstrate that singlet oxygen, and not sensitized rose bengal, was the reactive intermediate for the degradation experiments (Supplementary Figure S4). Reaction rate constants were quantified by competition kinetics in which the simultaneous degradation of 2-chlorophenol was monitored in the same solution.

In water containing GSH and MA, we observed faster MeHg decomposition by 1O2 than in water containing Cl or the phosphate buffer alone (Figure 3a). We observed similar trends in experiments with seawater (Supplementary Figure S5). These results indicated that 1O2–induced degradation was faster for CH3Hg-GSH and CH3Hg-MA complexes than for CH3HgCl and CH3HgOH/CH3HgHPO4 complexes (corresponding to PO4-buffered water with no additional ligands).

Figure 3
Effect of ligand complexation on degradation of MeHg

Complexation of sulfhydryl groups by mercury was expected to slow thiol oxidation by reactive oxygen species27. However, the total thiol concentration was in excess of total Hg in the samples. Thus, the thiols were decomposing simultaneously during the reaction (Supplementary Figure S6), producing compounds such as organic sulfur radicals that may be capable of inducing demethylation. We tested the reactivity of thiol oxidation products by performing dark experiments with hydrogen peroxide (H2O2), which oxidized GSH and MA to form sulfhydryl radical intermediates28. Decomposition of CH3Hg-GSH complexes was not observed, even as GSH (which was in excess of MeHg) was oxidized by H2O2 (Supplementary Figure S7). Therefore, we concluded that sulfhydryl radicals were not the cause of MeHg degradation.

Second-order rates constants (Table 1) calculated from the 1O2 experiments indicated that k1O2,CH3HgL values for CH3Hg-GSH and CH3Hg-MA complexes were 2 to 4 times greater than the rate constant measured for the CH3HgOH/CH3HgHPO4 mixture. Degradation of CH3HgCl was not detected for the time scale of our experiments (34 hr). Thus, the CH3HgCl rate constant with 1O2 was at least an order of magnitude slower. For the OH degradation pathway, rate constants did not appear to differ between CH3Hg-thiol complexes and CH3HgOH/CH3HgHPO4 mixed speciation (Table 1).

Table 1
Degradation rate constants for methylmercury complexes. Values represent average (± 1 standard deviation) of duplicate photochemical experiments at 25–26 °C.

The increased reactivity between CH3Hg-thiol complexes and 1O2 may be caused by shifts in carbon-mercury bond energies induced by thiol complexation. In theoretical studies, Ni et al.6 demonstrated that coordination of a reduced sulfur residue (-SH) to CH3Hg+ resulted in greater electronegativity at the carbon atom (and presumably increased reactivity with electrophiles). Tossel7 also utilized quantum mechanical calculations to show that sulfhydrl coordination decreased electron transition energies to a range where irradiation by low UV wavelengths (λ<270 nm) could induce demethylation.

We confirmed these theoretical calculations by performing direct photolysis experiments with UV-C radiation (λ = 254 nm) to demonstrate that thiol complexation alters electron transition energies of the CH3Hg-ligand molecule. Degradation of MeHg by UV-C was observed only if methylmercury was in the form of CH3Hg-GSH complexes (Figure 3b). Degradation was not observed in the phosphate buffer alone, where methylmercury consisted of 90% CH3HgOH and 10% CH3HgHPO4. Moreover, the observed decomposition rates increased as the GSH:MeHg molar ratio increased from zero to 1, and remained constant for all mixtures where GSH was in excess of MeHg. GSH did not decompose during direct UV-C exposure (91.9% ± 1.4% GSH recovered after 2600 mJ/cm2 fluence). These results showed that GSH complexation altered the reactivity of the MeHg molecule to make it more susceptible to photodegradation.

Speciation-dependent MeHg degradation rates can explain the variability observed between freshwater lakes and seawater. MeHg speciation differs greatly in these two ecosystems due to ligand competition between Cl and thiolate (R-S) functional groups associated with NOM. CH3HgCl complexes are expected to be the major species in coastal marine waters with Cl concentration greater than 0.1 M (Supplementary Figure S8). In contrast, most freshwater systems contain enough dissolved NOM (>1 mg-C/L) and low Cl such that CH3Hg+ ions are bound to thiol ligands associated with NOM26,2930.

In sunlit surface water containing NOM, the steady state concentration of 1O2 can be as high as 10−13 to 10−12 M 19. Using the CH3Hg-thiol rate constants in Table 1, we calculated pseudo-first order photodegradaton rates ranging from 0.04 to 0.6 d−1, values that are close to rates quantified in lakes24. Overall, our results indicate that ecosystem mass balances for MeHg must consider water composition and MeHg speciation when estimating losses from photodegradation. With this new understanding of how photodemethylation occurs, we can improve efforts to predict mercury cycling in the aquatic environment and prevent bioaccumulation of MeHg in food webs.

Supplementary Material


We thank K. Linden for assistance with UV photolysis experiments and K. McNeill and B.M. Voelker for helpful discussions regarding this study. This work was supported by Duke’s Pratt School of Engineering and Duke’s Center for Comparative Biology of Vulnerable Populations funded by the National Institute of Environmental Health Science.


Author contributions

T.Z. performed all experiments and data analysis. H.H. conceived the study, supervised the research, and performed speciation calculations. Both authors drafted the manuscript.

Competing Financial Interests

The authors have no competing financial interests.


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