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From a document on soil chemistry that I am reading:

The acid-base potential of a soil takes soil pH, total sulfur (sulfides and sulfates), and neutralization potential into account to determine the potential ability of a soil to generate acid.

Also, resources on well water quality usually mention only sulfides and sulfates, see, e.g., Indiana Department of Health — Sulfates & Sulfides in Well Water.

Why do these studies include sulfides $(\ce{S^2-})$ and sulfates $(\ce{SO4^2-}),$ but omit sulfites $(\ce{SO3^2-})? $ Don't sulfites occur in soil? Or is this just a question of terminology?

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    $\begingroup$ I think it's generally assumed that sulfite has a short lifetime in soil, decomposing to SO2 gas or getting oxidized to sulfate. See for example jstor.org/stable/42935238 (preview can be viewed for free). $\endgroup$
    – Andrew
    Mar 14, 2023 at 16:38

2 Answers 2

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I think Andrew is correct in his statement:

I think it's generally assumed that sulfite has a short lifetime in soil, decomposing to $\ce{SO2}$ gas or getting oxidized to sulfate.

Sulfur is found in natural soils as sulfide minerals like pyrites, marcasite, and greigite, and in sulfate forms like gypsum, anhydrite, barite, and jarosite. The S available for plants in agricultural ecosystems is in a dynamic storage, which can be schematically represented as the flow of sulfur in soil (where APS is adenosine 5′-phosphosulfate; Ref.1):

Sulfur circle

As correctly stateed in Ref.1:

In agricultural practice, sulfur has broad use in the form of sulfate fertilizers and, to a lesser extent, as sulfite biostimulants. In agricultural practice, sulfur has broad use in the form of sulfate fertilizers and, to a lesser extent, as sulfite biostimulants. When used in the form of bulk elemental sulfur, or micro- or nano-sulfur, applied both to the soil and to the canopy, the element undergoes a series of changes in its oxidation state, produced by various intermediaries that apparently act as biostimulants and promoters of stress tolerance. The final result is sulfate $\ce{S^6+}$, which is the source of sulfur that all soil organisms assimilate and that plants absorb by their root cells. The changes in the oxidation states of sulfur $\ce{S^0}$ to $\ce{S^6+}$ depend on the action of specific groups of edaphic bacteria. In plant cells, $\ce{S^6+}$ sulfate is reduced to $\ce{S^2-}$ and incorporated into biological molecules. $\ce{S^2-}$ is also absorbed by stomata from $\ce{H2S}$, COS, and other atmospheric sources. $\ce{S^2-}$ is the precursor of inorganic polysulfides, organic polysulfanes, and $\ce{H2S}$, the action of which has been described in cell signaling and biostimulation in plants. $\ce{S^2-}$ is also the basis of essential biological molecules in signaling, metabolism, and stress tolerance, such as reactive sulfur species (RSS), SAM, glutathione, and phytochelatins.

It is common that when sulfur containing minerals such as pyrites (a sulfide) exposed to atmosperic oxygen, they undergoes oxidation process and exists in oxidation states ranging from +6 to −2 (Table 1), with the most oxidized state in the form of sulfate $(\ce{SO4^2-})$, which is the chemical form that plants absorb from the soil to feed themselves with $\ce{S}$. That oxidation process may undergoes through sulfite formation, but it may only act as an intermediate in the presence of abundantly excess oxygen.

On the other hand, when sulfite is introduced to soil as $\ce{CaSO3}$, laboratory experiments has confirmed that two transformations of $\ce{CaSO3}$ occurred in soil systems in two ways: (1) decomposition to produce $\ce{SO2}$ gas, and (2) oxidation to calcium sulfate. Conversion to $\ce{SO2}$ occurred in solution and soil at low $\mathrm{pH}$, and acid soils treated with $\ce{CaSO3}$ were initially toxic to seedling root growth. The degree of toxicity was time-dependent, with reduction in toxicity occurring as $\ce{CaSO3}$ oxidized to calcium sulfate. Soil reaction also influenced toxicity, and at soil $\mathrm{pH}$ levels above 6, little seedling toxicity was evident (Ref.2). That may conclude this answer.

References:

  1. Laura Olivia Fuentes-Lara, Julia Medrano-Macías, Fabián Pérez-Labrada, Erika Nohemí Rivas-Martínez, Ema Laura García-Enciso, Susana González-Morales, Antonio Juárez-Maldonado, Froylán Rincón-Sánchez, and Adalberto Benavides-Mendoza, "From Elemental Sulfur to Hydrogen Sulfide in Agricultural Soils and Plants," Molecules 2019, 24(12), 2282 (17 pages) (DOI: 10.3390/molecules24122282).
  2. K. Dale Ritchey, Thomas B. Kinraide and Russell R. Wendell, “Interactions of calcium sulfite with soils and plants,” Plant and Soil 1995, 173(2), 329-335 (DOI: https://www.jstor.org/stable/42947538).
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This article(ref.) gives you your answer:

While techniques are available for determining the concentration of most S-oxides in soil, determination of soil sulphite has been largely neglected. This is probably because the ion is considered too labile to make a significant contribution to the S-content of most soils. An exception may be soils exposed to heavy atmospheric pollution where sulphite enters in rainfall or directly by absorption of $\ce{SO2}$

Ref.: WAINWRIGHT, M., and J. JOHNSON. “DETERMINATION OF SULPHITE IN MINERAL SOILS.” Plant and Soil 54, no. 2 (1980): 299–305. http://www.jstor.org/stable/42935238.

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    $\begingroup$ I wrote the answer then I noticed that the article was already mentioned in comments. So, I made the answer communtiy-wiki. $\endgroup$ Mar 14, 2023 at 17:07

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