2.0 Oxidation Of Pyrite

Matthew Otwinowski

Scaling Analysis Of Acid Rock Drainage 

2.0 OXIDATION OF PYRITE

In this Chapter, the complex kinetics of pyrite oxidation is briefly summarized. We refer the reader to [SyT] and references therein for more details and further references.

The overall stoichiometry of pyrite oxidation may be described by the following reactions: 

Reactions (S1) and (S2) summarize several elementary chemical reactions which take place at low pH values (reaction (S1)) or high pH values (reaction (S2)).

In order to describe the oxidation kinetics, the intermediate steps in this reaction must be considered. The following reactions characterizing the oxidation of pyrite have been proposed [StM], [SiS], [TaW]:

The chemical network responsible for pyrite oxidation consists of four reactions:

(R1) The oxidation of pyrite by molecular oxygen to Fe²⁺ and sulphate. The oxidation of iron sulphide (pyrite) to sulphate (eq. (R1)) releases dissolved ferrous ions and acidity into the water.

(R2) Subsequently, the dissolved ferrous ions undergo oxidation to ferric ions (eq. (R2)). This is a slow reaction and viewed as the rate-limiting step determining the overall rate of pyrite oxidation.

(R3) Sulphide is oxidized again by ferric ion and acidity is released along with additional ferrous ions which may re-enter the reaction cycle via reaction (R2). This is regarded as a fast step.

(R4) Ferric ions hydrolyse to form insoluble ferric hydroxide (eq. (R3)), releasing more acidity to the stream. This reaction takes place only at high pH values which can be attained when the mineral composition of the waste rock is such that self-buffering processes take place or when neutralizing minerals are added.

Reactions (R2) and (R3) provide a feedback loop discussed by Singer and Stumm (p. 471 of ref. [StM]). Ferrous iron produced in (R3) is utilized again in reaction (R2). The number of reacting species during acid formation is greater than two. This leads to strong nonlinearities in the kinetic equations.

2.1 Oxidation at pH Between 4 and 7

The set of coupled kinetic differential equations has been derived in our previous study and has the form:

In our previous study we have presented numerical solutions to the set of equations (D1) for the experimentally determined values of the temperature and pH dependent rate constants k(T) and reactive surface area, S per unit volume, V.

2.2 Oxidation at pH Less Than 4 

At low pH values ferric hydroxide does not precipitate if concentrations of ferric iron are sufficiently small. This means that the ferric iron, Fe³⁺ (which is formed by oxidation of ferrous iron) oxidizes pyrite according to the reaction (R3). The set of coupled kinetic differential equations has the form:

The use of the full set of the kinetic equations in a physicochemical model is not practical, however, because of the required long computer time when transport processes are included. The main factors affecting the rates of acid production are concentration of oxygen, temperature and pH value. Because in our previous study no oscillatory or chaotic behaviour of the coupled chemical equations (D1) and (D2) has been found, the detailed kinetic equations can be replaced by much simpler effective equations suitable for the physico-chemical model. The most significant information can be obtained by analyzing the temperature dependence of the overall rates of oxygen consumption and the dependence of acid production rates on temperature. We have analyzed our previous numerical results and on that basis it is possible to derive effective kinetic equations for oxygen. The effective kinetic equations are subsequently used as a starting point for the scaling analysis. Numerical solutions to the kinetic equations (D1) and (D2) are presented in Appendix A. The previous report [SyT] discusses the agreement of the results obtained for the kinetic equations with laboratory data.