KINETICS OF FES PRECIPITATION .1. COMPETING REACTION-MECHANISMS

The kinetics of the fast precipitation reaction between aqueous iron(II) and dissolved sulfide at 25 degrees C can be interpreted in terms of two competing reactions. The first may be represented by Fe2+ + H2S --> FeS(s) + 2H(+). This can be described by an observed rate law -d[aH(2)S]/dt = k(1)'[aH(2)S], where k(1)' is the observed first order rate constant with a value of 90 +/- 10 s (-1), [cH(2)S] is the concentration of dissolved H2S in moles per liter, and t is time in seconds. The rate law is consistent with an Eigen-Wilkins model of the process in which the rate is described by d[FeS]ldt = -d[aH(2)S]/dt = k(1)[aFe(2+)][aH(2)S], where aH(2)S and aFe(2+) are the formally dimensionless hydrogen sulfide and Fe(II) activities which are represented on a moles . liter(-1) scale for experimental and practical convenience. The logarithm of k(1), the theoretical Eigen-Wilkins reaction rate constant, has a value of 7 +/- 1 liters . mole(-1). s(-1). The second reaction may be represented by Fe2+ + 2HS(-) --> Fe(HS)(2)(s). The rate of this reaction may be described by an observed rate law of the form -d[aHS(-)]/dt = k(2)'[aHS(-)](2), where [aHS(-)] is the formally dimensionless bisulfide activity which is represented on a moles liter(-1) scale for experimental convenience. The observed second order rate constant, k(2)' has a value of 1.3 x 10(7) liters . mole(-1). s(-1) at 25 degrees C. The result is consistent with an Eigen-Wilkins model of the process in which k(2)' = k(2)[aFe(2+)][aHS(-)](2), where aFe(2+) is the dissolved Fe(II) activity and the logarithm of k(2), the Eigen-Wilkins reaction rate constant, has a value of 12.5 +/- 1 liters(2) . mole(-2). s(-1). The theoretical interpretation of both reactions suggests that the rates are direct functions of the ion activity products of the iron sulfide precipitates. The second stage of the reaction involves the condensation of Fe(SH)(2) to FeS with the release of dissolved sulfide back to solution: Fe(SH)(2)(s) --> FeS(s) + H2S. This reaction is relatively slow and results in a sinusoidal form superimposed on the Stage 1 concentration-time curve. Overall, the rate of removal of total dissolved sulfide from solution by these processes can be empirically modeled by -d[Sigma S]/dt = k(0)[Sigma S] and [Sigma S = [Sigma S](0) = e(-k0t) where [Sigma S] is the concentration of total dissolved sulfide at any time, [Sigma S](0) is the concentration of total dissolved sulfide at t = 0, and k(0) is a pseudo first order rate constant of 15 s(-1) where [Sigma Fe2+], the total dissolved iron(II) concentration, is between 10(-3) and 10(-4) M. Theoretically, the rates of both reactions are directly proportional to [aFe(2+)]. A good approximation for the rate of removal of total dissolved sulfide by the iron(II) (bi)sulfide precipitation processes in most environments can, therefore, be obtained using a value for k(0) of 15 x 10(-4)/[Sigma Fe2+]. Application of the rate laws to natural systems suggests that the relative dominance competing and independent of Sigma Fe2+. In environments with ppm or greater Sigma S concentrations (greater than or equal to 10(-3) M), the rate of sulfide removal is two magnitudes greater in neutral to alkaline systems than in systems with pH < 7. The bisulfide pathway resulting in the formation of Fe(SH)(2) dominates and the H2S pathway only dominates in acidic environments. The results suggest that, in these relatively sulfide-rich environments, a standing concentration of Fe(SH)(2) will be present and may constitute an important component in further reactions, such as pyrite formation. In contrast, in sulfide-poor systems with Sigma S concentrations at the sub-ppm (<10(-3) M) level, the rate is greater in neutral to acidic conditions and the H2S pathway, involving the direct formation of FeS, dominates in all environments with pH < 8.