Oxidation of Pyrite Introduction

Pyrite which is found in a wide variety of geological formations is the most widespread and abundant of the sulphide minerals (Craig and Vokes, 1993). It is widely distributed and forms under extremely varied conditions. For example, it can be produced by magmatic (molten rock) segregation, or by hydrothermal solutions, and as stalactite growth. It occurs as an accessory mineral in igneous rocks, in vein deposits with quartz and sulphide minerals, and in sedimentary rocks, such as shale, coal, and limestone (Rafferty,

  • 2020). Pyrite is sometimes called fool's gold because of its similarity to gold in colour (brassy-yellow) and shape (King, 2020a; Rafferty, 2020). However, pyrite is quite easy to distinguish from gold, i.e., pyrite is much lighter, but harder than gold and cannot be scratched with a fingernail or knife (TMM, 2014). In addition, pyrite will tarnish when exposed to acid, whereas gold is nonreactive (IGS, 2011).
  • Oxidation Process

Pyrite is stable under anaerobic conditions, but is oxidised and dissolved to release soluble iron species and sulphuric acid when it comes in contact with oxygen and water (Satur et al., 2007). In fact, the oxidative dissolution of pyrite is one of the most extensively studied geochemical processes by many researchers (Lowson, 1982; Nordstrom and Alpers, 1999; Edwards et al., 2000; Esparia, 2008; Simate and Ndlovu, 2014). The pyrite oxidation and the factors affecting the kinetics of oxidation (02, Fe3+, temperature, pH, Eh, and the presence or absence of microorganisms) have been the focus of extensive study because of their importance in both environmental remediation and mineral separation (Blowes et al., 2003). As shown in Figure 3.1, the oxidation of pyrite follows a cycle of complex reactions (Stumm and Morgan, 1996; Banks et al., 1997; Ali, 2011; Buzzi et al., 2013) involving surface interactions with dissolved 02, Fe?+, and other mineral-based catalysts such as Mn02 (Blowes et al., 2003; Simate and Ndlovu, 2014). Ideally, several products are formed during the oxidation of pyrite including metastable secondary products such as ferrihydrite (5Fe203-9H20), schwertmannite (between Fe80s(0H)6S04 and Fe16Oi6(OH)]0(SO4)3), and goethite (FeO(OH)), as well as the more stable secondary jarosite (KFe3(S04)2(0H)6), and hematite (Fe203) depending on the geochemical conditions (Dold, 2010; Dold, 2014).

For pyrite oxidation, oxygen and ferric iron are the two possible available oxidants. However, the initial step and most important reaction is the oxidation of the pyrite (or sulphide) in the presence of atmospheric oxygen and water forming dissolved ferrous iron, sulphate, and hydrogen as shown in Equation 3.1 (Akcil and Koldas, 2006; Dold, 2010). The initial step of pyrite


Model for the oxidation of pyrite. (From Stumm and Morgan, 1996; Ali, 2011; Buzzi et al., 2013.) oxidation is also illustrated in Figure 3.1 (Stumm and Morgan, 1996; Banks et al., 1997; Ali, 2011; Buzzi et al., 2013). It must be noted that though two oxidisable species are present in pyrite (ferrous iron and sulphidic sulphur), it has been experimentally determined that irrespective of the mechanism (oxygen or ferric mediated), during the initial solubilisation of pyrite only the sulphidic sulphur is oxidised and the iron passes into solution in the ferrous state (Lowson, 1982).

In an environment which is sufficiently oxidizing (dependent on 02 concentration, pH greater than 3.5 and bacterial activity), the ferrous iron generated as shown in Equation 3.1 may be oxidised to ferric iron according to reaction 3.2 (Blowes et al., 2003; Akcil and Koldas, 2006; Udayabhanu and Prasad, 2010). However, according to Fripp et al. (2000), if the concentration of oxygen is low, reaction 3.2 will not occur until the pH reaches 8.5.

Once ferric iron is produced by the oxidation of ferrous iron (reaction 3.2), which is the case at low pH conditions and strongly accelerated by microbiological activities, then ferric iron also becomes an oxidant of pyrite (reaction 3.3) (Dold, 2010). In fact, an important factor in the oxidation of pyrite and the generation of acid mine waters is that Fe3+ is able to oxidise pyrite under anoxic subaqueous conditions at a much faster rate than does molecular oxygen (Espana, 2008). In other words, though oxygen is a primary oxidant, the ferric iron (Fe3+) resulting from the oxidation of ferrous iron is now recognised as a more powerful oxidant than oxygen even at nearneutral pH (Zdun, 2001).

It is noted that at pH < 3.5, reaction 3.2 is several orders of magnitude slower than reaction 3.1 (Espana, 2008). Therefore, the oxidation of Fe2+ by oxygen is usually considered as the rate-limiting step in pyrite oxidation (Singer and Stumm, 1970; Skousen et al., 1998). However, the presence of acidophilic bacteria such as Acidithiobacillus ferrooxidans and Leptospirillum ferrooxidans greatly accelerates (by a factor of around 106) the abiotic oxidation rate (Singer and Stumm, 1970; Nordstrom and Alpers, 1999; Espana, 2008), thus maintaining a high concentration of ferric iron in the system (Espana, 2008).

According to Singer and Stumm (1970), the overall process resulting from the combination of reactions 3.2 and 3.3 is traditionally known as the 'propagation cycle' and along with reaction 3.1 depicts a model by which pyrite oxidation initially starts by reaction 3.1 with oxygen as the oxidant at circum- neutral pH conditions, and as pH decreases to about 4 the oxidation of pyrite proceeds through reaction 3.3. It must be noted, however, that oxygen will always be required to replenish the supply of ferric iron according to reaction 3.2, so that the overall rate of pyrite oxidation is largely dependent on the overall rate of oxygen transport by advection and diffusion (Nordstrom and Alpers, 1999; Ritchie, 1994; Ritchie, 2003; Esparia, 2008).

It is noted further that at pH values between 2.3 and 3.5, ferric iron formed in reaction 3.2 may precipitate as Fe(OH)3 (and to a lesser degree as jarosite, H30Fe3(S04)2(0H)6) while simultaneously producing acid as shown in reaction 3.4 (Blowes et al., 2003; Akcil and Koldas, 2006; Dold, 2010; Dold, 2014). The hydrolysis and subsequent precipitation of Fe(OH)3 produces most of the acid in the whole process (Dold, 2010; Dold, 2014). If pH is less than 2, ferric hydrolysis products like Fe(OH)3 are not stable and Fe3+ remains in solution (Dold, 2010; Dold, 2014). Therefore, any remaining Fe3~ from reaction 3.2 that does not precipitate into Fe(OH)3 (or jarosite) from solution through reaction 3.4 may be used to oxidise additional pyrite, according to reaction 3.3 (Akcil and Koldas, 2006).

The process of pyrite oxidation relates to all sulphide minerals once exposed to oxidizing conditions (e.g., chalcopyrite, bornite, molybdenite, arsenopyrite, enargite, galena and sphalerite among others). In other words, while the principal sulphide mineral in mine wastes is pyrite, other sulphide minerals are also susceptible to oxidation releasing elements such as aluminium, arsenic, cadmium, cobalt, copper, mercury, nickel, lead, and zinc into the water flowing through the mine waste (Blowes et al., 2003). The oxidation of other sulphide minerals is discussed in the next sections.

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