II. Thermokinetics of Sulphide Oxidation

     

Matthew Otwinowski

Non-linear waste rock modelling

II. THERMOKINETICS OF SULPHIDE OXIDATION

STOICHIOMETRIC EQUATIONS FOR PYRITE OXIDATION

(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 at pH less than 4.5.

(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 waste rock is such that self-buffering processes take place or when neutralizing minerals are added. Reaction (R4) eliminates reaction (R3) at high pH values.

At low pH values, reactions (R2) and (R3) provide a feedback loop discussed by Singer and Stumm (1970)¹ and Stumm and Morgan (1986)². (Ferrous iron produced in (R3) is utilized again in reaction (R2)).

KINETIC EQUATIONS FOR PYRITE OXIDATION

Oxidation at pH between 4 and 7

Kinetic equations for pyrite oxidation have a form³:

Rate constants k₁ and k₂ depend strongly on temperature. This is the dominant nonlinearity at the micro-scale, which is responsible for the strongly nonlinear thermodynamic behaviour of waste rock piles at the macro-scale.

The order p of the pyrite oxidation reaction with respect to the concentration of oxygen dissolved in water, [O₂], is equal to 1/2. For p=1/2, the reduction of the oxygen concentration has a smaller effect on the reduction rates then for p=1. (For p=1 the oxidation rates are reduced by 50% when the oxygen concentration is reduced by 50%; for p=1/2 the oxidation rates are reduced by 50% when the oxygen concentration is reduced by 75% - this is important for the design of covers).

TEMPERATURE AND pH DEPENDENCE OF REACTION RATES

The rate constants have been measured^4,5,6 at temperature T_o=20°C. Rate constants k₁ and k₂ are determined by measuring the molar oxidation rate per unit surface area of pyrite crystals and the molar concentration of oxygen dissolved in water (see also Table I). Pyrrhotite oxidizes much faster than pyrite.

For the activation energy of 70-90 kJ/mol reaction rates accelerate faster than by a factor of two per 10°C increase in temperature.

SATURATION CONCENTRATIONS OF OXYGEN 

DISSOLVED IN WATER

Table I. Saturation concentrations of dissolved oxygen in water, [O₂], at atmospheric pressure (1013.25 hPa) for different oxygen concentrations in the gaseous phase [O₂]_gas. The solubilities of oxygen are defined in terms of the mass of oxygen dissolved in one litre of the water in equilibrium with an atmosphere saturated with water vapour at various temperatures. The data are calculated following Benson and Krause (1980)⁷ and using hydrological tables published by Beer (1991)⁸.

Oxidation rates are controlled by temperature and oxygen concentration. Saturation values of oxygen dissolved in water decrease with temperature. This effect partly counteracts the temperature increase of oxidation rates. The dependence of [O₂] on [O₂]_gas and T was disregarded in previous models. We were able to derive an analytical formula which describes the values of [O₂](T,[O₂]_gas). This formula is used in the model. (We will use notation X=[O₂], Y=[O₂]_gas).

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¹ P.C. Singer and W. Stumm, Acid Mine Drainage: The Rate Determining Step, Science 167, 1121 (1970).

² W. Stumm and J.J. Morgan, Aquatic Chemistry. An Introduction Emphasizing Chemical Equilibria in Natural Waters, Wiley 1986.

³ M. Otwinowski, Quantitative Analysis of Chemical and Biological Kinetics for the Acid Mine Drainage Problem, MEND 1994.

⁴ M. A. McKibben and H.L. Barnes, "Geochimica et Cosmochimica Acta", 51, 793(1987).

⁵ B. Wehrli, In Aquatic Chemical Kinetics, W. Stumm ed., Wiley 1990.

⁶ R.V. Nicholson et al., "Geochimica et Cosmochimica Acta", 52, 1077(1988) & 54, 395(1990).

⁷ B.B. Benson and D. Krause Jr., "Limnology and Oceanography", 25, 662(1980).
⁸ T. Beer, Applied Environmetrics, Victoria, Australia (1991).