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Metal Sulfide Reactions

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    At the very outset it is important for you to know that relatively little is understood about the fundamental reactions of metal sulfide minerals. As you will see below, the reactions can be rather complex because there are many possible intermediate products. The dissolution of pyrite (FeS2) and arsenopyrite (FeAsS) involves a series of oxidation reactions. Pyrite has been studied more intensely, and some aspects of its chemistry are well understood.8, 9, 10, 11 Arsenopyrite is much more reactive than pyrite and while there is much evidence that the mechanism of its oxidation is different from that of pyrite, the exact mechanism is not known. However, since pyrite is the most common and most widely studied metal sulfide, it is useful to review the chemistry of pyrite as a starting point for understanding arsenopyrite and other metal sulfides.

    In both pyrite and arsenopyrite, the oxidation states of the elements in the minerals are Fe2+, S1− and As1−. The iron is released in the +2 oxidation state (as ferrous ion), which is the same oxidation state that it has in the solid mineral. Because the −1 oxidation states for S and As are not stable in solution, the mineral reactions always involve oxidation of the S and As to some higher oxidation state in order to be released into solution. Consequently, most of the mechanistic chemistry of these metal sulfides is associated with oxidation of the sulfur species, and we refer to the reactions as “oxidative dissolution” to emphasize that oxidation must occur in order to dissolve at an appreciable rate. Because sulfur can be present in a number of different forms (sulfur, sulfite, sulfate, thiosulfate, and several other forms), measuring the sulfur species is not a good way of determining the overall rate of metal release. Arsenic is also very difficult to measure directly at the low concentrations needed. Consequently, in the first part of the design project you will be focusing on measuring the iron that is released when metal sulfides are exposed to aqueous solutions.

    Oxidation of sulfur

    The process of arsenopyrite oxidation can occur via several mechanisms, depending on what one assumes is the oxidizing agent and what oxidation state one assumes the products are in. In the environment, Fe3+ and O2 are the two main oxidizing agents. Ferric ion is the one of greatest interest. Because sulfur and arsenic both have several different stable oxidation states, one can write a number of different reactions depending on whether one is interested in the forms that are released from the mineral, or if one also includes subsequent reactions occurring in the water. For pyrite, the overall reaction is usually written:

    \[\ce{FeS2(s) + 14 Fe^3+ + 8 H2O \rightarrow 15 Fe^2+ + 16 H+ + 2 SO4^2-} \tag{1}\]

    Looking at this equation, you should notice that in order to balance the equation you need 23 moles of reactants! Clearly, this reaction does not occur by having 14 ferric ions and 8 water molecules simultaneously converge onto one spot on the FeS2 crystal lattice. Instead, it is generally believed that oxidation of pyrite occurs by sequential oxidation of the sulfur atoms at the surface to produce the thiosulfate anion, which is then released from the mineral surface into the aqueous phase, where it is later converted to sulfate.

    &\ce{FeS2(s) &&+ 2 Fe^3+ + H2O &&\rightarrow FeS2O(s) &&+ 2 Fe^2+ + 2 H+} \tag{2}\\
    &\ce{FeS2O(s) &&+ 2 Fe^3+ + H2O &&\rightarrow FeS2O2(s) &&+ 2 Fe^2+ + 2 H+} \tag{3} \\
    &\ce{FeS2O2(s) &&+ 2 Fe^3+ + H2O &&\rightarrow FeS2O3(s) &&+ 2 Fe^2+ + 2 H+} \tag{4}\\
    &\ce{FeS2O3(s) && \rightarrow Fe^2+ + S2O3^2- &&} \tag{5}

    Once released into water, the thiosulfate can be further oxidized to sulfate:

    \[\ce{S2O3^2- + 8 Fe^3+ + 5 H2O \rightarrow 8 Fe^2+ + 10 H+ + 2 SO4^2-} \tag{6}\]

    Adding Equations 2 through 6 gives the overall reaction in Equation 1.

    In the case of arsenopyrite, you can write a similar balanced equation, although in this case the arsenic only needs to be oxidized to the +3 oxidation state (As3+).

    \[\ce{2 FeAsS(s) + 14 Fe^3+ + 3 H2O \rightarrow 16 Fe^2+ + S2O3^2- + 6 H+ + 2 As^3+} \tag{7}\]

    The thiosulfate anion in solution might then be oxidized to sulfate, completing the formation of sulfuric acid. Remember, that while some aspects of the pyrite reaction are understood, much less is known about the overall mechanism of reaction for arsenopyrite. In particular, it is worthwhile noting that in pyrite the crystal structure has two sulfur atoms already closely bonded together in the S2 subunit, while in arsenopyrite the structure is more complicated. Thus, in arsenopyrite and most other sulfide-containing minerals it is not clear whether thiosulfate is important, or whether other pathways involving sulfite or other anions are most important.

    Understanding the fate of the sulfur-containing species is important because they control the overall reaction rate. Although the overall reaction should eventually produce sulfate (SO42-), there are many possible intermediates and reaction pathways. One particularly important intermediate is elemental sulfur. It is possible for the oxidation to stop at S0, leaving elemental sulfur (S8) as a product. Since sulfur is not very soluble in water, it remains on the surface of the mineral. If the sulfur forms a thin even coating on the mineral it could inhibit further reaction. However, there is good evidence that the sulfur on the surface is not even, but rather it covers only parts of the surface with isolated patches, leaving considerable areas free of sulfur and exposed to the oxidative environment.12, 13 Quantitative measurement of the amount of elemental sulfur present have yielded some clues as to possible reaction pathways.14, 15

    The roles of Fe3+ and O2 as oxidizers

    A key step in the reactions of all the sulfide minerals is that they release iron as ferrous ion (Fe2+). Oxygen present in the water can then oxidize Fe2+ back to Fe3+:

    \[\ce{2 Fe^2+ + \frac{1}{2} O2 + 2 H+ \rightarrow 2 Fe^3+ + H2O} \tag{8}\]

    This Fe3+ can again react with the mineral surface. Hence, the presence of oxygen greatly accelerates the rates of reactions such as 2 to 4. However, it does so in an indirect way, rather than reacting directly with the mineral surface. Under most conditions the regeneration of Fe3+ from Fe2+ and O2 (reaction 8) is believed to be the rate-limiting step in arsenopyrite oxidation and dissolution.

    It is also theoretically possible for O2 to react directly with the mineral surface, and one can write a number of different possible reactions, such as:

    \[\ce{FeAsS(s) + 5 \frac{1}{2} O2 + 3 H2O \rightarrow 2 Fe^2+ + 2 H3AsO3 + 2 SO4^2-} \tag{9}\]

    However, under most conditions this reaction is not as important, due in part to the fact that the solubility of O2 in water is quite low, especially at low pH.

    An additional factor that complicates the chemistry and affects the concentration of Fe3+ available in solution is the pH. Using Le Chatlier’s principle, you might predict that the overall dissolution reaction (Reaction 7) would proceed more easily at high pH than at low pH because Reaction 7 produces H+, which can then react with the available OH. In reality, the dissolution does not occur more readily at high pH because OH reacts Fe3+ in solution to produce Fe(OH)3 via the precipitation reaction:

    \[\ce{Fe^3+ (aq) + 3 OH- (aq) \rightarrow Fe(OH)3(s)} \tag{10}\]

    The equilibrium constant for this reaction is the inverse of the solubility product (1/Ksp) of Fe(OH)3 and is rather large, on the order of 1038. Hence, at high pH the concentration of Fe3+ is very, very low. Conversely, at low pH, the solution can contain more Fe3+(aq) and reactions 2 through 6 will occur more readily than they do at higher pH. Hence dissolution occurs more rapidly at low pH.

    Oxidation and the role of Fe2+

    Finally, the last unresolved question has to do with the reduced form of the oxidizer. Because the overall reactions are all energetically favorable, one would normally expect the back-reactions to not be important. However, if any of the steps involve an unusually slow step, then it is possible for the ferrous ion concentration to be significant. The most obvious way for this to happen would be if one of the steps involving an electron transfer step is close to equilibrium. The electron-donating ability of a solution is usually described in terms of the electrochemical reduction potential. The electrochemical potential for a system containing Fe2+ and Fe3+ will be given by:

    \[E=E^\circ_{\large{Fe^{2+}/Fe^{3+}}}-\dfrac{RT}{nF}\ln\dfrac{[Fe^{2+}]}{[Fe^{3+}]},\hspace{30px} \textrm{where n = 1.} \tag{11}\]

    If the slowest step in the overall reaction pathway involves a slow, nearly reversible exchange of electrons between the solid and aqueous Fe3+, then we might expect the overall reaction equilibrium to depend on the ratio of Fe3+ to Fe2+, rather than just the concentration of Fe3+ alone.

    Unanswered questions

    As you can see from the discussion above, the chemistry of arsenopyrite is complex and not all aspects are well understood. A scientist interested in learning how to understand, predict, and control how metal sulfide minerals affect the environment must first learn about the specific chemical interactions. Both a quantitative and qualitative study of the soil and rock samples is necessary to learn about the chemicals present, and about how these chemicals react to cause the problems discussed above. Some of the unanswered questions include:

    1. What oxidized forms of sulfur and arsenic are present on the surface?
    2. In what chemical forms are sulfur and arsenic released from the mineral surface?
    3. Is elemental sulfur an intermediate in the formation of thiosulfate or sulfate, or does it form via some type of parallel reaction pathway?
    4. Which chemical step or steps control the overall rate?
    5. How does the reaction rate depend on variables such as pH, oxidizing agents, temperature, etc.?
    6. Is the rate controlled only by the reactant Fe3+, or does the product Fe2+ have any effect on the overall rate of reaction?

    While answering all of these questions is certainly beyond the scope of what you can do in a Project Lab within a chemistry course, you should think about these questions as you are planning your project. Perhaps your experiments can yield some new information (not yet known!) that will further our understanding of the oxidative dissolution of arsenopyrite.

    This page titled Metal Sulfide Reactions is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by Contributor.

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