Oxygenation of Ferrous Iron

亚铁的氧化

IIRON

WERNER STUMM and G. FRED LEE’ Department of Sanitary Engineering, Harvard University, Cambridge, Mass.

Oxygenation of Ferrous IronThis study of the reaction between ferrous iron and oxygen emphasizes the role oxygenation catalysts play in the deferrization process. Application of these catalysts may markedly increase the rate of over-all iron removal from natural watersBEARING WATERS are often deferrized by oxidizing the ferrous iron with dissolved oxygen and then removing the resulting ferric oxide floc by sedimentation or filtration. Ferrous iron in river water--e.g.. pollution by mine drainage or iron pickling wastes-not only disturbs the oxygen balance but reduces the p H and forms suspended matter as ferric oxide-Fe20, XH20, where X may vary. The capability of iron to undergo reversible oxidation or reduction reactions is rdated to the metabolic cycles of nearly all other elements of importance and especially to the distribution of oxygen in impounded surface waters. The corrosion of, and the subsequent occurrence of“red water” in a water distribution system with iron piping is influenced by the oxygenation rate of ferrous iron (22). A quantitative investigation of the reaction between ferrous iron and oxygen is needed in order to understand the many roles iron may play in natural waters. Special attention should be given to the kinetics of this reaction; the rate at which it proceeds depends on several parameters including pH, temperature, concentration of dissolved oxygen, and catalysts. The investigations reported here have shown that the rate of oxygenation of ferrous iron in bicarbonate solutions follows the law:

sions reached in these investigations have yielded valuable clues on the mechanism of the reaction between ferrous iron and oxygen, but they do not necessarily apply to bicarbonate-containing waters. Weiss (25) suggested that the oxygenation of ferrous iron proceeds in the following sequence:

+ 7 Fe(II1)+ HOi ( a j Fe(I1)+ HO.2+ Fe(II1)+Fe(I1)0 2

HzO? ( b )

+ HyOy+Fe(II1)+ H O’+ H2O (c) Fe(1I)+ H O .+Fe(II1)Fe(T1)

(2)

+ H20 ( d )

These reactions as written are not balanced with respect to hydrogen ions and are not intended to specify the dependence of the reaction rate on this variable. The rate determining step is believed to be Reaction 2.a thus indicating that the oxygenation rate is of the first order with respect to both the concentration of Fe(I1) and dissolved oxygen.-dFe(rl)= k[Fe(II)] (Os)

the hydrogen ion concentration is observed (70. 72), a very pronounced p H dependence is observed in less acidic solutions (1: 7 7, 13, 24). However, except for the very early study of Just (73), no quantitative investigations have been undertaken on the rate of ferrous iron oxygenation in the buffer system of natural waters. Solubility of Iron. Studies on iron oxidation in bicarbonate solutions are complicated by the slight solubility of ferrous and ferric iron in such solutions. Ferrous

constituents generally are more soluble than ferric constituents. Table I gives a survey of the equilibria that must be considered for the evaluation of their respective solubilities. In bicarbonate waters, the solubility of ferrous iron is controlled by the solubility of FeC03. and can be estimated by[Fe+*

+ Fe(OH)+]=

dt

(3)

--d[Fe(ll)l dt

= k[Fe(I1)]Po,(HO-)* (1)

Catalysts (especially Cu -2) in trace quantities, as well as anions which form complexes Lvith ferric iron (HzP(3-). increase the reaction rate significantly, while small amounts of Fe(III), C1-, S04-? have no effect on the reaction rate. Previous Studies Many reliable and extensive studies have been made in homogenous acid solutions (6, 72: 74, 20). The concluPresent address, Graduate School of Public Health, University of Pittsburgh, Pittsburgh, Pa.

The rate of oxidation is dependent upon the nature of the anion present, and increases as the complexing affinity of the anion to ferric iron increases (12). Thus under acid conditions, the rate is found to decrease in the following series: hydroxide, pyrophosphate (72), phosphate (76), chloride (20),sulfate (12, 14), and perchlorate (70). The rate law is in accordance with Equation 3 for the first four media in this list, but a second order dependence of the rate on the ferrous iron concentration has been found in sufficiently acid solutions of H& 0 4 and HCIOI. In addition to complexing agents, many substances, especially those which hasten the decomposition of peroxides in the presence of ferrous ironplatinum, charcoal, and cupric saltsaccelerate the oxidation of ferrous salt by oxygen (73). While in acid media only a small dependence of the oxygenation rate upon

where K, is the ion product of water and alk is the alkalinity[(HCO,-)+ 2(C0,-z)+(OH-)] in equivalents per liter and the numbered K values refers to the respective equation in Table I . Equation 4 is valid up to a p H of about 9 (23). In order to estimate the solubility of ferric iron: equilibria 6 to 9 in Table I must be considered. Experimental investigations (7, 9) in basic and neutral solutions (pH values 6 to 11) have shown that solubility of ferric hydroxide is markedly independent of the hydroxyl ion concentration. This observation has led to the proposal that Fe(OH)3 (s) is in equilibrium with undissociated molecular F e ( 0 H ) j (equilibrium 9: Table I). Experimental Method It is not feasible to study, within the pH range of natural waters, the kinetics of oxygenation of ferrous iron in homogenous solutions (Table I). Kinetic investigations in heterogenous solution are less amenable to rigid interpretation owing to the super-imposition of interfacial factors which operate at the phase boundaries on to the homogenous reacVOL. 53, NO. 20

FEBRUARY 1961

143

亚铁的氧化

Table I.

Solubility Equilibria

No.1 2

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