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I was reading a Jules Verne book recently, in which the following chemical reaction was done:

He placed a layer of branches and chopped wood, on which were piled some pieces of shistose pyrites ... This done, they set fire to the wood, the heat was communicated to the shist, which soon kindled ... Then new layers of buised pyrites were arranged so as to form an immense heap, the exterior of which was covered with earth and grass, several air-holes left.

(The shistose pyrates consisted of a mixture of coal, iron sulfide, and aluminum oxide)

After 10 days, the $\ce{FeS2}$ in the pyrite mixture was converted into iron sulfate. However, my experience with roasting iron pyrite involves sulfur dioxide vapors coming off and leaving behind iron. How would the iron sulfate be created?

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2 Answers 2

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Typically, roasting sulfides leads to the generation of sulfur dioxide and metal oxide. In the case of pyrite, this may be described by

$$\ce{4 FeS2 + 10 O2 -> 2 Fe2O3 + 8 SO2 }$$

The fire is needed to initiate the reaction. If well maintained, the heat generated by roasting may sustain the reaction.

However, and especially for marcasite as an other form of $\ce{FeS2}$ (polymorphism), it is known that exposure to humidity and air suffices to oxidize the sample in some parts, leading to the generation of small amounts of sulfuric acid. Because sulfuric acid is a stronger acid than $\ce{H2S}$, the former may displace the later. Eventually, the sample literally gets eaten away to yield $\ce{FeSO4}$ which is described as «pyrite decay».

The same source refers to a Conserv O Gram «Storage Concerns for Geological Collections» (April 1998, Number 11/2) by the National Parks' Museum Management Plan. This publication mentions $\pu{60 \%RH}$ as critical threshold (though, likely at ambient temperature, without statement how fast this process is). Possibly, the process observed in the museum doesn't happen in the kiln Verne described.

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    $\begingroup$ Yep, when moisture and oxygen are plentiful, "acid mine drainage" forms easily from the aerobic oxidation of iron pyrite to iron(II) sulfate. Iron(II) sulfate is highly water soluble, at least in acid. It is further oxidized by air to insoluble iron(III) oxides and hydroxides, but the abiotic reaction is slow at low pH. However, even at low pH, the reaction can be effected by specialized microbes. Under the right conditions, this can lead to thick scums of biofilm that grow at the acid/air interface in molar concentrations of iron and pHs of as low as zero or below. $\endgroup$
    – Curt F.
    Jan 6, 2022 at 2:38
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Wet-oxidation of pyrite lead to iron(II) sulfate.

$$\ce{2FeS2 + 2H2O + 7O2 ->[100 °C] 2FeSO4 + 2H2SO4}$$

iron(II) sulfate formed is eventually further oxidized to iron(III) sulfate, iron(III) oxide and iron(III) hydroxide. Note that the reaction is not that simple as it proceeds through many Fe(II) and Fe(III) intermediates:

enter image description here

There are different proposed reaction mechanism for this reaction. You can look up here, here and here.

References

  1. Pyrite Oxidation Mechanism by Oxygen in Aqueous Medium Egon Campos Dos Santos, Juliana Cecília de Mendonça Silva, and Hélio Anderson Duarte The Journal of Physical Chemistry C 2016 120 (5), 2760-2768 DOI: 10.1021/acs.jpcc.5b10949
  2. https://pubs.usgs.gov/of/1995/0389/report.pdf
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