Oxygenation of Ferrous Iron(3)

145

亚铁的氧化

401

1

\

&

5-

I3b 1’0

\io3Io

40

50 1

610

70 1

*b

TlME.min

Figwe 3. C U+~catalyzes the oxygenation reaction20.5’

C. Po,= const.

pH= 6.0

with results obtained in acid systems by Cher and Davidson (6). The results of these authors have been accounted for by the following m

echanism: Fe(I1) CU(I)

+ Cu(I1) 7+ Fe(II1)CuV) ( a )

tion rate. The addition of various amounts of complex formers such as glycine, tetramine, and ethylene diamine tetraacetate did not affect the reaction rates. The latter two complexing agents markedly retard the oxygenation reaction in the absence of CuT2. Copper complexes appear to be as effective catalysts as the aqueous copper ions. Some results of a preliminary survey on the effects of other catalysts are represented in Figure 4 illustrating that other metal ions (MnT2, Co-2) as well as anions (besides OH-) which form complexes with F e ( I I I ) ( H? P 0 4 - . metaphosphate) are capable of hastening the oxygenation reaction. This figure shows that the reaction in Cambridge tap water which contains no analytically detectable C U+ is considerably faster than~ the reaction in synthetic water of similar mineral composition. Addition of chloride and sulfate (10 mg. per liter) had no effect on the oxidation rate. Practical Aspects Oxidation of ferrous iron is only one step in the deferrization process. Flocculation and sedimentation (or filtration) certainly may codetermine the over-all iron removal rate in surface waters and in iron removal plants. Many constituents present in natural waters may accelerate or decelerate the oxidation as well as the flocculation reaction. It can be inferred from our experiments, however. that the oxidation reaction may be the controlling factor up to pH values around 7. In the more alkaline p H range, flocculation may be the slowest step. At pH values above 8, diffusion or the oxygen rather than the chemical reaction determines the oxidation rate. In many natural waters organic matter may stabilize ferric oxide colloids or may increase the Fe(II1) solubility by complex formation. Generally, iron removal is very slow under such conditions (5, 79). Most organic impurities are known to hasten oxygenation reactions; under these conditions, difficulties in iron removal are apparently due to slow flocculation and cannot

necessarily be overcome by an increase in pH. Coagulation of iron(II1) is fast in many natural waters within the pH range 5 to 7 (4>5, 27) and ferric oxide removal by sand filtration is best within the same p H range (27). Under such conditions, however, oxidation rates are slow. Here the application of oxidation catalysts-e.g. Cu+2 may prove very efficient. For a given p H and oxygen concentration, the addition of as little as 0.02 mg. per liter reduces the oxygenation time by a factor of 5. Acknowledgment The authors wish to thank Carol Sinclair and N. S. Graham for excellent technical assistance. Literature Cited(1) Abel, E.. Z . Elektrochem. 59, 903 (1955). (2) American Public Health Association,

+

0 2

f

Cu(l1) HO,

+( b ) (8)

The HO’? is assumed to react further with Fe(I1) as in the Weiss mechanism (Equation 2). The data obtained do not permit the formulation of a meaningful rate expression which would account satisfactorily for

the catalytic effect of copper. However, severalqualitative statements can be made. The same proportionate acceleration of the oxygenation rate was observed at any pH value within the pH range investigated. The runs made in the presence of Cu-2 show the same temperature dependence as those made in the absence of C u f 2 . The catalytic effect of C U+~ dependent upon both is the initial ferrous and ferric iron concentrations. Higher initial ferrous iron concentrations and/or an incipient addition of ferric iron causes a n increased rate of reaction. There is a saturation concen~ tration of C U+ for any given initial ferrous iron above which further addition of Cu+2 has little or no effect on the reac-

New York.“Standard Methods for the Examination of Water Sewage and Industrial Wastes,“ 9th and 10th Cdition, 1946. 1955. (3) American Society of Testing Materials, Philadelphia, Pa..“Manual on Industrial Water.’‘ D. 264. 1954. (4) Bartow, E.,’Black, A. P., Sandsbury, W. E., Trans. Am. Soc. Czd Engrs.

100, 263 (1935). (5) Camp, T. R., Root, D. A, Bhoota, B. V., J . Am. Water W o r k s Assoc. 32,1913 (1940). (6) Cher, M., Davidson, N.. J . Am. Chem. SOC. 793 (1955). 77, (7) Feitknecht, W., Z. Elektrochem. 63, 34 (1959). (8) Feitknecht, W., Keller, G.; Z. anorg. u. allgem. Chem. 262, 61 (1950). (9) Gayer, K. H.? Woontner, L., J . Phys. Chem. 6C, 1569 (1956). (IO) George, P., J . Chem. SOC. 1954,4349. (11) Holluta, J., Eberhardt, M., Jb. Vom Wasser 24, 79 (1957). (12) Huffman, R. E., Davidson, N.. J . Am. Chem. Soc. 78, 4836 (1956). (13) Just: G., Z., Physik. Chemie 63, 385 (1908). (14) Lamb, A . P., Elder, L. W., . Am. J Chem. SOG. 137 (1931). 53, (15) Latimer, W. M.,”Oxidation Potentials,” p. 222, Prentice Hall, New York, 1953. (16) Lee, G. F., Stumm. W.,. A m . W a t e r J M’orks Assoc. (in press). (17) Leussing, 0. L.. Kolthoff, J. M., J . Am. Chem. SOC. 2476 (1953). 75, (18) Milburn. R. M.. Ibid.. 79, 537 (1957). (19) Moore, E. W., Snow. E.. A, j .:Yew Engl. Water Works Assoc. 56, 320 (1942). (20) Posner, A. M., Trans. Faraday SOC. 49, 382 (1953). (21) Stanley, D. R.: P h . D . thesis. Harvard University. Cambridge. Mass., 1952. (22) Stumm, W.: Proc. Am. Soc. C i d Eners. San. Engr. D i t . 86, 21 (November1 o .. . ., m

Figure 4. Insidious trace quantities of impurities may strongly accelerate the oxidation reaction

(23) S;Umm,\V.. Lee. G. F., Rerm. suisse d’Hydrologie 22, 295 (1960). (24) Todt, F., Stoklassa, K.. Z . Elektrochem. 5 8, 354 (1954). (25) Weis …… 此处隐藏：4134字，全部文档内容请下载后查看。喜欢就下载吧 ……