Chemical Reactions Of The Ferrous

The ferrous salt used as coagulant in water and wastewater treatment is copperas, FeS04 • 7H20. For brevity, this will simply be written without the water of hydration as FeS04. When copperas dissolves in water, it dissociates according to the following equation:

As in the case of alum, the ions must be rapidly dispersed throughout the tank in order to effect the complete coagulation process. The solid precipitate Fe(OH)2(s) and complexes are formed and expressing in terms of equilibrium with the solid Fe(0H)2(s), the following reactions transpire (Snoeyink and Jenkins, 1980):

Fe(0H)2(s) ^ Fe2++ 20H_ KSp,Pe( oh)2 — 10-145 (12.22)

Fe(0H)2(s) + 0H_ ^ Fe(OH)_ KFe(0H)3c — 10-5" (12.24)

The complexes are Fe0H+ and Fe(0H)_. Also note that the 0H_ ion is a participant in these reactions. This means that the concentrations of each of these complex ions are determined by the pH of the solution. In the application of the above equations in an actual coagulation treatment of water, conditions must be adjusted to allow maximum precipitation of the solid represented by Fe(OH)2(s). To allow for this maximum precipitation, the concentrations of the complex ions must be held to the minimum. The values of the equilibrium constants given above are at 25°C.

12.6.1 Determination of the Optimum pH

For effective removal of the colloids, as much of the copperas should be converted to the solid Fe(OH)2(s). Also, as much of the concentrations of the complex ions should neutralize the primary charges of the colloids to effect their destabilization. Overall, this means that once the solids have been formed and the complex ions have neutralized the colloid charges, the concentrations of the complex ions standing in solution should be at the minimum. The pH corresponding to this condition is the optimum pH for the coagulation using copperas.

Let spFeII represent all the species that contain the Fe(II) ion standing in solution. Thus, the concentration of all the species containing the ion is

[ spFe„ ] — [ Fe2+ ] + [ Fe0H+] + [ Fe(0H)_] (12.25)

All the concentrations in the right-hand side of the previous equation will now be expressed in terms of the hydrogen ion concentration. As in the case of alum, this will result in the expressing of [spFeII] in terms of the hydrogen ion. Differentiating the resulting equation of [spFeII] with respect to [H+] and equating the result to zero will produce the minimum concentration of spFeII and, thus, the optimum pH determined. Using the equilibrium reactions, Eqs. (12.22) through (12.24), along with the ion product of water, we now proceed as follows:

rT. 2+n {Fe2+} Ksp,Fe(OH)2 Ksp,Fe(OH)2{H } Ksp,Fe(OH)27nfH ]

7FeI1 7FeIl( OH") 7FeIlKw 7FeIlKw rp /-\tt+T {FeOH } ^-FeOHc ^-FeOHc

FeOHc H

7FeOHc 7FeOHc {OH } 7FeOHcKw 7FeOHcKw

{Fe(OH)") KFe(OH)3c{OH } KFe(OH)3cKw KFe(OH)3cKw

7Fen, 7eOHc, 7Fe(OH^c are, respectively, the activity coefficients of the ferrous ion and the complexes FeOH+ and [Fe(OH)- ]. Ksp,Fe(OH)2 is the solubility product constant of the solid Fe(OH)2(s). KFeOHc and K Fe(OH)3 c are, respectively, the equilibrium constants of the complexes FeOH+ and Fe(OH)-.

Equations (12.26) through (12.28) may now be substituted into Equation (12.25) to produce r Ksp,Fe(OH)2?H[H ] + KFeOHc7H[H+] + KFe(OH)3cKw

7FeIlKw 7FeOHcKw 7Fe(OH)3c7H[H ]

Differentiating with respect to [H+], equating to zero, rearranging, and changing H+ to Ho+pt , the concentration of the hydrogen ion at optimum conditions,

2Ksp,Fe(°H)27H I + 3 +J KFeOH c7h ^„+-,2 KFe(OH)3cKw rH0piJ +1 --^ HHopiJ - —-— (12.3°)

7FeIIK2w J p [TFeOHcKwp 7Fe(OH)3c7H

The value of [Hopt] may be solved by trial error.

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