## 3 Results and Discussion
### Uptake Experiment
The \(\text{Hg}^{0}_{\text{aq}}\) concentrations in the pure water, acidified water with nitric acid and seawater were, on average (\(\pm\)SD), \(150\pm 112\) pg L\({}^{-1}\), \(166\pm 51\) pg L\({}^{-1}\) and \(225\pm 41\) pg L\({}^{-1}\), respectively. These average \(\text{Hg}^{0}_{\text{aq}}\) concentrations were at a comparable level, but nearly one-half to one-third that of the equilibrated \(\text{Hg}^{0}_{\text{aq}}\) concentrations. This is most likely explained by the uptake solutions reaching the surface of the uptake solution, but not reaching equilibrium throughout the whole solution due to the slow diffusion of \(\text{Hg}^{0}_{\text{aq}}\). Our experiments exposed the standard gas in constant concentrations of GEM to the surface of the uptake solutions. Thus, it would take some time to reach equilibrium. On the other hand, many methods in the literature used for the determination of equilibrated concentrations bubbled standard GEM gas in solutions, and the physical mixing method would bring equilibrium rapidly. Meanwhile, the average \(\text{Hg}^{0}_{\text{aq}}\) concentrations in the acidified solution with sulfuric acid (48 \(\pm\) 23 pg L\({}^{-1}\)) were lower than those in the other uptake solutions. This difference can be explained by the chemical reaction of \(\text{Hg}^{0}_{\text{aq}}\) with sulfuric acid during the two-hour extraction of \(\text{Hg}^{0}_{\text{aq}}\). Because \(\text{Hg}^{0}_{\text{aq}}\), hydronium and sulfate ions coexisted in the solution even after GEM exposure, the \(\text{Hg}^{0}_{\text{aq}}\) was likely converted to \(\text{Hg}^{2+}_{\text{aq}}\) during \(\text{Hg}^{0}_{\text{aq}}\) extraction. The \(\text{Hg}^{2+}_{\text{aq}}\) in the uptake solutions showed interesting trends (Figure 4). The \(\text{Hg}^{2+}_{\text{aq}}\) concentrations in the seawater were lower than the \(\text{Hg}^{0}_{\text{aq}}\), while the \(\text{Hg}^{2+}_{\text{aq}}\) concentrations in the other solutions were higher than the \(\text{Hg}^{0}_{\text{aq}}\). The \(\text{Hg}^{2+}_{\text{aq}}\) concentrations in the pure water (pH 6.1-6.5), acidified water with nitric acid (pH 3.5-3.6) and seawater (pH 7.9-8.1) showed almost no time-dependent change as a two uncertainty range was considered for the slope, and the \(\text{Hg}^{2+}_{\text{aq}}\) concentrations in the acidified water with sulfuric acid (pH 3.3) increased significantly as a function of time up to 1068 pg L\({}^{-1}\), exceeding the upper limit of the equilibrated \(\text{Hg}^{0}_{\text{aq}}\) concentration. The series of results suggests the following: (1) the constrained oxidation of \(\text{Hg}^{0}_{\text{aq}}\) in seawater was likely associated with a high pH, rather than the high ionic strength, while some extent of the oxidation reaction of the \(\text{Hg}^{0}_{\text{aq}}\) took place in the pure water and acidified water with nitric acid, possibly with dissolved oxygen; (2) the oxidation reaction of \(\text{Hg}^{0}_{\text{aq}}\) in the acidified solution with sulfuric acid was associated not only with hydronium ions, but with sulfate ions. The first finding was based on the fact that the \(\text{Hg}^{0}_{\text{aq}}\) concentrations between all solutions were comparable, implying a low diffusivity of \(\text{Hg}^{0}_{\text{aq}}\) in seawater, due to the other dissolved ions not significantly lowering the GEM uptake under our experimental conditions, and the high-pH seawater containing dissolved oxygen did not oxidize the \(\text{Hg}^{0}_{\text{aq}}\) as much as the other solutions did. This explanation is also consistent with our general knowledge that \(\text{Hg}^{2+}_{\text{aq}}\) is stable in an acidic solution; that is, hydronium ions are likely a key player in the redox reactions of \(\text{Hg}^{0}_{\text{aq}}\), but it still seems to require other molecules to oxidize \(\text{Hg}^{0}_{\text{aq}}\). This is also consistent with our general knowledge that mercury is a less effective reducing element than hydrogen in the activity series. The second finding is also explained by this, and the oxidation is specific to sulfuric acid.