In horticulture, lime sulfur (lime sulphur in British English; see American and British English spelling differences) is mainly a mixture of calcium polysulfides and thiosulfate (plus other reaction by-products such as sulfite and sulfate), formed by reacting calcium hydroxide with elemental sulfur, and is used in pest control. It can be prepared by boiling a suspension of poorly soluble calcium hydroxide (lime) and solid sulfur in water, together with a small amount of surfactant to facilitate the dispersion of these solids. After elimination of residual solids (flocculation, decantation, and filtration), it is normally used as an aqueous solution, which is reddish-yellow in color and has a distinctive offensive odor of hydrogen sulfide (H2S, rotten eggs).
Synthesis reaction The exact chemical reaction leading to the synthesis of lime sulfur is generally written as:
Ca(OH)2 + x/8 S8 → CaSx + by-products (S2O2−3, SO2−3, SO2−4) as reported in a document of the US Department of Agriculture (USDA). This reaction is poorly understood because it is vague and involves the reduction of elemental sulfur, and no reductant appears in the equation, while sulfur oxidation products are also mentioned. The initial pH of the solution imposed by poorly soluble hydrated lime is alkaline (pH = 12.5) while the final pH is in the range 11–12, typical for sulfides, which are also strong bases. When the hydrolysis of calcium sulfide is taken into account, the individual reactions for each of the by-products are:
1/2 S8 + H2O + 2 Ca(OH)2 → 2 H2S + CaS2O3 3/8 S8 + H2O + 2 Ca(OH)2 → 2 H2S + CaSO3 1/2 S8 + 2 H2O + 2 Ca(OH)2 → 3 H2S + CaSO4 However, elemental sulfur can undergo a disproportionation reaction, also called dismutation. The first reaction resembles a disproportionation reaction. The inverse comproportionation reaction occurs in the Claus process, which is used for desulfurization of oil and gas products in the refining industry:
H2S + 3/2 O2 → SO2 + H2O By rewriting the last reaction in the inverse direction, one obtains a reaction consistent with what is observed in the overall lime sulfur reaction:
3/8 S8 + 2 H2O → 2 H2S + SO2 In alkaline conditions, it yields:
3/8 S8 + 2 H2O + 6 OH− → 2 S2− + SO2−3 + 5 H2O and after simplification, or more precisely recycling of water molecules in the above reaction:
3/8 S8 + 6 OH− → 2 S2− + SO2−3 + 3 H2O Adding back 6 Ca2+ cations from hydrated lime for the sake of electroneutrality, one obtains the overall reaction. This last reaction is consistent with the overall lime sulfur reaction mentioned in the USDA document. However, it does not account for all the details, such as the production of thiosulfate and sulfate among the end products of the reaction. Nevertheless, it is a good first-order approximation, and it usefully highlights the overall lime sulfur reaction scheme because the chemistry of reduced or partially oxidized forms of sulfur is particularly complex, and all the intermediate steps or involved mechanisms are hard to unravel. Moreover, once exposed to atmospheric oxygen and microbial activity, the lime sulfur system will undergo a rapid oxidation, and its different products will continue to evolve and eventually enter the natural sulfur cycle. The presence of thiosulfate in the lime sulfur reaction can be accounted for by the reaction between sulfite and elemental sulfur (or with sulfide and polysulfides), and that of sulfate by the complete oxidation of sulfite or thiosulfate, following a more complex reaction scheme. More information on calcium thiosulfate production is described in a patent by Hajjatie et al. (2006). Hajjatie et al. (2006) expressed the lime sulfur reaction in various ways depending on the degree of polymerization of calcium polysulfides, but the following reaction is probably the simplest of their series:
3 Ca(OH)2 + 6 S → 2 CaS2 + CaS2O3 + 3 H2O where the S2−2 species corresponds to the disulfide anion −S−S− (with a covalent bond between the two sulfur atoms), also present in pyrite (FeS2), a Fe(II) disulfide mineral. They also successfully controlled this reaction to achieve the conversion of elemental sulfur into a quasi-pure solution of calcium thiosulfate.
Preparation
The New York State Agricultural Experiment Station recipe for the concentrate suggests starting with 80 lb of sulfur, 36 lb of quicklime, and 50 gal of water, equivalent to 19.172 kg of sulfur and 8.627 kg of calcium oxide per 100 liters of water. About 2.2:1 is the ratio (by weight) for compounding sulfur and quicklime; this ratio yields the highest proportion of calcium pentasulfide. If calcium hydroxide (builders' or hydrated lime) is used, an increase of one-third or more (to 115 g/L or more) may be used with the 192 g/L of sulfur. If the quicklime is 85%, 90%, or 95% pure, 101 g/L, 96 g/L, or 91 g/L is used, respectively; if impure hydrated lime is used, its quantity is increased to compensate, though in practice lime with a purity lower than 90% is rarely used. The mixture is then boiled for one hour while being stirred, and small amounts of water are added for evaporation.
… excerpt ends here. Continue reading the full article.
