5.4.3: Oxidation of sulfide and intermediate sulfur compounds
- Page ID
- 132821
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)Reduced sulfur compounds are those in which sulfur has an oxidation state less than \(+6\). That includes the sulfide produced by dissimilatory sulfate reduction and desulfurylation reactions as well as sulfur compounds with intermediate oxidation states (between \(-2\) and \(+6\)) such as elemental sulfur \(\left(\text{S}^{0}\right)\), thiosulfate \(\left(\text{S}_{2} \text{O}_{3}^{2-}\right)\), sulfite \(\left(\text{SO}_{3}^{2-}\right)\), dithionite \(\left(\text{S}_{2} \text{O}_{4}^{2-}\right)\), tetrathionate \(\left(\text{S}_{4} \text{O}_{6}^{2-}\right)\), and polysulfide compounds \(\left(\text{S}_{n}^{2-}\right)\). Intermediate sulfur species are collectively represented by \(\text{S}^{0}\) in Figure \(5.7\) and can be produced by incomplete oxidation of sulfide.
Chemotrophic and phototrophic microorganisms can use reduced sulfur species as electron donors in reactions that supply energy and fix carbon, respectively. Potential electron acceptors in the reactions include oxygen \(\left(\text{O}_{2}\right)\), nitrate, and ferric iron. Some example reactions with oxygen as the electron acceptor follow:
\[\begin{align} & \text{H}_{2} \text{S} + 2 \ \text{O}_{2} \longleftrightarrow \text{SO}_{4}^{2-} + 2 \ \text{H}^{+} \\ & \text{S}_{2} \text{O}_{3}^{2-} + \text{H}_{2}\text{O} + 2 \ \text{O}_{2} \longleftrightarrow 2 \ \text{SO}_{4}^{2-} + 2 \ \text{H}^{+} \\ & \text{SO}_{3}^{2-} + 0.5 \ \text{O}_{2} \longleftrightarrow \text{SO}_{4}^{2-} \\ & \text{S}^{0} + \text{H}_{2}\text{O} + 1.5 \ \text{O}_{2} \longleftrightarrow \text{SO}_{4}^{2-} + 2 \ \text{H}^{+} \end{align}\]
Microorganisms capable of sulfur oxidation are found within all three domains of life. Among Bacteria, well-recognized groups that contain species capable of chemotrophic sulfur oxidation include Thiobacillaceae and Beggiatoaceae (Ehrlich and Newman, 2009). Filamentous members of the Desulfobulbaceae, known as cable bacteria, can oxidize sulfide by transferring electrons from sulfide to oxygen or nitrate over centimeter distances in sediments (Müller et al., 2020; Sandfeld et al., 2020). Thus, they can help cycle sulfur by forming an electrical connection across redox zones. Anoxygenic phototrophic Bacteria capable of sulfur oxidation have been identified have been identified in four groups: green sulfur bacteria (Chlorobiaceae), purple sulfur and non-sulfur bacteria, Gram-positive Heliobacteria, and filamentous and gliding green bacteria (Chloroflexaceae) (Barton et al., 2014). Among Archaea, Sulfolobus and Acidianus are widely studied groups capable of sulfur oxidation (Ehrlich and Newman, 2009). Lastly, some fungi are also able to oxidize sulfur (Li et al., 2010).