Chemical Modeling of Aqueous Systems II - American Chemical Society

Chapter 34

Effect of Ionic Interactions on the Oxidation Rates of Metals in Natural Waters Frank J. Millero

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Rosenstiel School of Marine and Atmospheric Science, University of Miami, 4600 Rickenbacker Causeway, Miami, F L 33149

The reduced valence of metals in natural waters can be formed by photochemical, biochemical and geochemical processes. The longevity of these reduced metals will be influenced by the rates of oxidation with O and H O . This paper will examine how ionic interactions affect the rates of Cu(I) and Fe(II) oxidation in natural waters using measurements made as a function of pH, ionic strength and composition. The oxidation of Cu(I) is affected by the strong interactions of Cu with Cl-. The formation of CuCl2and CuCl 2- ion pairs causes the rates to decrease. At a constant Cl- con centration, the increase in the concentration of Mg causes the rates to decrease and the increase in the concentration of HCO3- causes the rates to increase. The oxidation of Fe(II) is affected by the strong interactions of Fe with OH-. The formation of F e O H and Fe(OH) ion pairs causes the rates to increase. Most other anions, with the exception of HCO3-, that form ion pairs cause the rates to decrease. 2

2

2

+

3

2+

2+

+

2

Trace metals in natural waters can participate in a wide variety of chemical reactions. Some of thes metals are needed, while others can be toxic to organisms. The bioavailability and toxicity of metals is strongly affected by the redox state and speciation of a given metal. Copper, for example, is toxic to many marine organisms when it is in the free or uncomplexed form. Iron and manganese are needed by phytoplankton. The oxidized forms of these metals are insoluble in seawater and are quickly scavenged from surface waters. Although thermodynamic speciation cal culations can yield the most probable redox form of a given metal, the form is normally controlled by the rates of oxidation and reduction of the metals. Recently a number of studies have been made on the rates of oxidation and reduction of copper and iron in natural waters (1-9). These measurements provide reliable rate equations for the oxidation of Cu(I) and Fe(II) as a function of pH, temperature, ionic strength and ionic composition. In the present paper, these results will be used to examine how ionic interactions affect the rates of Cu(I) and Fe(II) oxidation in natural waters. The effect of ionic interactions on the reaction rates of metals can be affected by the anions in solution (CI", S0 ", OH", etc.); while the anion reaction rates (e.g., HS") can be affected by the cations in solution ( H , N a , M g , etc.). Examples of the effect anions have on the rates of the reactions of cations are the forma tion of ion pairs 2

4

+

+

2 +

0097-6156/90/0416-0447$06.00/0 c 1990 American Chemical Society

Melchior and Bassett; Chemical Modeling of Aqueous Systems II ACS Symposium Series; American Chemical Society: Washington, DC, 1990.

CHEMICAL MODELING OF AQUEOUS SYSTEMS II

448 Cu Fe

+ 2 +

+ CI"—> CuCl° + OH" —> F e O H

(1) (2)

+

The formation of CuCl causes the rates of oxidation of Cu(I) to decrease (1-2.5-9). while the formation of F e O H causes the rates of oxidation of Fe(II) to increase (2-5.10). Examples of the effect of cations on the rates of reactions of anions are +

+

HS" + H — > H S 0 - + M g —> M g 0

(3) (4)

2

2 +

2

+ 2

The formation of H S causes the rates of sulfide oxidation to decrease (11), while the formation of the ion pair M g 0 causes the rates of 0 " to H 0 disproportionation to decrease (12). Our interest in the effects of ionic interactions on the rates of reactions has largely been for redox processes occurring in the oceans. These redox processes occur at oxic-anoxic inter faces and in the surface waters of the oceans due to photochemical processes (1). The formation of H 0 in surface waters (1.13.14) is thought to occur by the following reaction scheme 2

+


2

2

2

2

2

2

C + hv > C* C* —> C + e'(aq) 0 + e'(aq)—> 0 ' 20 - + 2H —> H 0 + 0

(5) (6) (7) (8)

+

2

2

+

2

2

2

2

The 0 " can react with the oxidized form of metals to produce reduced species 2

Fe Cu

3 +

2+

+ 0 "—> Fe + 0 + 0 — > Cu + 0 2

2 +

(9) (10)

2

+

2

2

The longevity of these reduced metals will depend on the rates of oxidation of Fe(II) and Cu(I) with 0 and H 0 2

2

2

Fe(II) + 0 — > Products Fe(II) + H 0 — > Products Cu(I) + 0 —> Products Cu(I) + H 0 — > Products

(11) (12) (13) (14)

2

2

2

2

2

2

This present paper will examine how the oxidation of Fe(II) and Cu(I) in natural waters can be affected by changes in the composition. These results will be explained in terms of ionic interac tions. The results are taken from measurements made in this laboratory (2.3.5-9.14.15). These papers should be examined for details of the methods used. CU(I) RATES OF OXIDATION The rates of oxidation of Cu(I) with 0 and H 0 have been made in seawater, NaCl and sea salts (6-9). The measurements were made as a function of pH and temperature. A comparison of the results for seawater and NaCl solutions at 25°C and pH = 8.0 are shown in Figures 1 and 2(^9). The large ionic strength dependence in NaCl and seawater solutions for the oxidation of Cu(I) with 0 and H 0 is related to changes in the CI" concentration. This chloride dependence can be analyzed by assuming that the various Cu(I) chloro complexes have different rates of oxidation (2). This gives 2

2

2

Cu

+

2

2

2

+H 0 2

2

CuCl + H 0 2

2

ko — > Products ki — > Products


(15) (16)

449

Oxidation of Metals in Natural Waters


MILLERO

-1

_1

Figure 2. The rate constant, (k, mol kg H 0 s ), for the oxidation of Cu(I) with H 0 in NaCl and seawater at 25°C and pH = 8.0. 2


2

2

450


CuCl - + H 0 2

2

2

2

CuCl " + H 0 3

2

2

—> Products

(17)

— > Products

(18)

1

-1

The observed rate constant is given by (k, mol" k g H 0 sec ) 2

k

k

= Vcu

a

k

k

a

1 9

+ l C u C l + 2"CuCl + 3 CuCl 2

( )

3


where the various rate constants, kj, are for the molar fractions a of species i. The substitution of the stepwise stability constants, β*, for the formation of the various ion pairs gives k/«Cu

=

k

C1

0 + M l * f "]

+

k

2/*2*[Clf + ...

(20)

Values of k / a for measurements of Cu(I) and 0 made in NaCl-NaC10 mixtures at a constant ionic strength are shown in Figure 3. The linear behavior indicates that C u and CuCl are the reactive species. Similar results for the measurements of Cu(I) + H 0 are shown in Figure 4 (9). Measurements made over a wider range of ionic strength (to 6m) indicate that the oxidation of CuCl " becomes important (7). Measurements in NaBr-NaC10 and NaI-NaC10 indicate that the ion pairs CuBr and Cul have similar rates of oxidation with CuCl (8). The differences of the overall rates in CI', Br" and I" solutions ( k c i B r l ) related to the differences in the stability constants ( £ > £ >£ ). The results for the oxidation of Cu(I) in seawater at the same CI" concentration are lower than the values in NaCl (see Figures 1 and 2). To elucidate these effects, measurements have been made on the rates of oxidation in NaCl, NaMgCl, NaCaCl, N a C l H C 0 , N a C l S 0 and NaMgClS0 solutions at their concentrations in seawater. The results are shown relative to the measurements in NaCl in Figure 5. The addition of S 0 " shows no effect, while the addition of M g and C a causes the rates to decrease. The addition of H C 0 " causes the rates to increase. The solution containing Na , M g , CI" and H C 0 " gives rates in agreement with the seawater results. Measurements of the M g and H C 0 " effects have been made over a wide range of ionic strengths (7). These results indicate that the effects are diminished at higher ionic strengths. The decrease in the rate of oxidation caused by the addition of M g or C a can be attributed to the slow exchange of MgL complexes with C u C u

2

4

+

2

2

2

4

>k

C u I

C u B r

>k

4

1&

C u C 1

3

4

4

2

4

2 +

2 +

3

+

2 +

3

2 +

3

2 +

2 +

2 +

MgL + C u

2 +

— > CuL + M g

2 +

(21)

This slow exchange may cause the overall oxidation rates of Cu(I) to be slower due to the back reactions of Cu(II) with H 0 (1). NTA, C 0 " , E D T A and B(OH) " give similar rates at con centration levels sufficient to complex C u in Na-Mg-Cl solutions. Thus, this exchange reaction may be common for natural ligands able to complex M g and C u , such as humic material. The increase in the rate of oxidation due to the addition of H C 0 " and C 0 " may be related to the formation of C u H C 0 ° which slows down the back reaction. A preferable pos sibility is that the increase is due to the formation of Cu(I) carbonate ion pairs 2

2

2

3

4

2 +

2 +

2 +

2

3

3

3

Cu Cu

+

+ HC0 " — > CuHC0 + C 0 " — > CuC0 " 3

+

(22) (23)

3

2

3

3

These ion pairs may be more reactive than CuCl. Differences in the back reactions of CuEDTA and C u C 0 3

+

CuEDTA + O ^ — > C u + 0 CuC0 + 0 3

_ < 2

+

~ > Cu + 0

2

+ EDTA 2

2

+ C0 " 3


(24) (25)

34.

MILLERO

451


1.2-

1.0-4 0.8ο

r

•

0.6-4 Downloaded by COLUMBIA UNIV on October 31, 2017 | http://pubs.acs.org Publication Date: December 7, 1990 | doi: 10.1021/bk-1990-0416.ch034

Ο

*Γ

0.40.24 0.00.0

—ι— 0.2

0.4

0.6

0.8

—ι— 1.0

1.2

[cr] 1

Figure 3. Values of k / a , (k, moH kg H 0 s" ), for the oxidation of Cu(I) with 0 in NaCl-NaC10 mixtures at I = 1.0 at 25°C. C u

2

2

4

[ci-] -1

1

Figure 4. Values of k / a , (k, mol kg H 0 s" ), for the oxidation of Cu(I) with H 0 in NaCl-NaC10 mixtures at I = 1.0 at 25°C. C u

2

4


2

2

Melchior and Bassett; Chemical Modeling of Aqueous Systems II ACS Symposium Series; American Chemical Society: Washington, DC, 1990. 2

Figure 5. The effect of various ions on the oxidation of Cu(I) with 0 at I = 0.7 at 25°C.


Η W

S3

§

> Ο

Ο Ο

w

ο

η

4t

34.

MILLERO

453


with 0 " or H 0 could also have different rates. If reaction (24) is faster than reaction (25), the overall rates of Cu(I) will be faster in C 0 solutions. Further kinetic and thermodynamic data are needed to elucidate the H C 0 " or C 0 ' effects. Even though the causes of how Mg > C a and H C 0 " affect the rates of oxidation of Cu(I) is uncertain, the correction of the seawater results for these effects yields rate constants that agree with the measured values in NaCl (see Figure 6). 2

2

2

2 -

3

2

3

3

2+

2 +

3

FE(II) RATES OF OXIDATION The rates of Fe(II) oxidation with 0 in natural waters have been studied by a number of workers (see references in 3). Measurements have also been made in artificial media (16.17.18). Stumm and Lee (1Q) showed that from pH = 5 to 7, the rate of oxidation of Fe(II) in water was given by (k, mol" kg H 0 min' ) 2

3

3

1


2

2

d[Fe(II)]/dt = -k[Fe(II)][OH"] [0 ]

(26)

2

The overall rate constant is given by 1 2

log k = 21.56 - 1545/T - 3.29 I / + 1.521

(27)

Millero et al. (3) have shown that this equation is valid in water and seawater from 0 to 50°C and salinity, S = 0 to 40. This second order dependence of the oxidation with OH" or H is shown in Figure 7 for the pseudo first order rate constant (kj) +

d[Fe(II)]/dt = -k^Feai)]

(28)

More recently, Millero and Izaguirre (16) have examined the pH dependence of the oxidation of Fe(II) in NaCl and NaC10 solutions as a function of ionic strength and temperature. The results at I = 6.0 are shown in Figure 8. These results clearly demonstrate that over a wide range of temperature (0 to 50°C) and ionic strength (0 to 6m), the rate of oxidation of Fe(II) is second order with respect to H or O H . As discussed elsewhere (2), this is related to the rate determin ing step for the oxidation being 4

+

-

Fe(OH) + 0 —> Fe(OH) 2

The 0

2

2

+

+0

2

-

(29)

2

produced reacts very quickly with the more predominant F e Fe Fe Fe

2 +

-

+ 0 + 2 H 0 — > Fe(OH) + + H 0 + H 0 + H 0 — > Fe(OH) + O H + H + O H + H 0 —> F e ( O H ) + H 2

2 +

2

2

2

+

2

2 +

2 +

2

2

2

+

2

+

2

+

2

(30) (31) (32)

The concentration of Fe(OH) is given by 2

Fe(OH) = [Fe(II)]/î [ O H f / a 2

2

(33)

F e

where a

2

F e

= (1 + ^ [ O H ] + β [ΟΗ]

+ β [ΟΗ]ψ

2

3

(34)

and β is the stepwise association constant. Thus, the overall rate equations can be simplified to ί

d[Fe(II)]/dt = -k'[Fe(OH) ][0 ] 2

(35)

2

where k' is related to k of Equation 26, β and a . 2

F e




454

Figure 7. Values for the pseudo first-order rate constants for the oxidation of Fe(II) with 0 in water and seawater as a function of pH at 25°C. 2


34.

MILLERO

455


The effect of pH on the oxidation of Fe(II) with H 0 is shown in Figure 9. The slope is near zero below a pH = 4.0 and has a value of 1.0 from 6 to 9. These results indicate that both F e and F e O H are reactive to oxidation by H 0 . From the comparisons shown in Figure 7, the rate of oxidation of Fe(II) in seawater is lower in seawater than in water at a given pH (on the free scale). As discussed elsewhere (2), this decrease has been attributed to the formation of ion pairs 2

2 +

+

2

Fe Fe


2

2 +

2

+

+ CI"— > FeCl + S 0 " —> FeS0 °

2 +

(36) (37)

2

4

4

that are less reactive to oxidation. Measurements of the overall rate constant in seawater and NaCl as a function of ionic strength (Figure 10) indicated that at higher ionic strengths, the seawater results are lower than the NaCl results. To investigate this decrease, we have made some measurements as a function of composition at a constant pH and ionic strength (0.7m). The results are given in Figure 11. The replacement of N a with M g (0.0547m) causes a decrease in the rate. This effect was not expected. This decrease in the rate due to M g may be related to the decrease in the concentration of C u H C 0 ° or CuC0 " species which, as discussed later, cause the rates to increase. The replacement of CI" by S 0 " (0.0293m) causes the rate of oxidation of Fe(II) to decrease. This can be attributed to the formation of the FeS0 ion pair. The decrease can be used to estimate the stability constant for the formation of FeS0 +

2 +

2 +

3

3

2

4

4

4

k

k

NaCl/ NaClS0

=

1

4

38

+ ^FeSO^S^Vl

C )

=

Our results yield ^FeS0 500 ± 125 which is in reasonable agreement with the infinite dilution value (β = 132), but larger than expected at I = 0.7 β = 10 (19). More reliable estimates will be made later. A solution made up of N a , M g , CI" and S 0 " gives the same results as an artificial seawater solution ( N a , M g , C a , K , CI", S0 ", Br", HC0 ") and real seawater (3). To determine the effect of ionic strength and temperature on the rates of oxidation of Fe(II), measurements were made in NaCl and NaC10 solutions (16). The 25°C results are shown in Figure 12. In dilute solutions the results are the same (16.4) within experimental error. At higher ionic strengths the results in NaCl are lower than in NaC10 . The decrease in the rates at low ionic strengths is similar to the seawater results (3) and the results in NaC10 of Sung and Morgan (18). This decrease is related to the effect of ionic strength on β as well as the rate con stant in pure water. The slope varies from -3.3 in seawater to -1.6 in NaCl or NaC10 . Since one would not expect the rate constant for Equation 29 to show an ionic strength dependence, one is forced to attribute these negative slopes to the reaction of F e and 0 ". The lower rate constants in NaCl at higher ionic strengths can be attributed to the forma tion of F e C l 4

+

2 +

2

4

+

2 +

2 +

+

2

4

3

4

4

4

2

4

2 +

2

+

k

k

cio / ci = l + / W C 1 1

(39)

4

=

The results from 4 to 6 m give jSpeCi 1·2 ± 0-5, which is in reasonable agreement with the litera ture value (19) of 1.0. The effect of various anions on the oxidation of Fe(II) (19) has been determined by making measurements in NaCl-NaX at a constant ionic strength. The results (16) are shown for log k as a function of the anion (X) in Figure 13. The addition of H C 0 " causes the rate to increase, while the addition of S 0 " and B(OH) " causes the rate to decrease. The anions, B r , N 0 " and C10 " cause the rate to increase slightly. The overall order of the rate constants are in the order HC0 "> >Br">N0 ">C10 ">Cl"> >S0 "> >B(OH) \ The large increase due to the addition of H C 0 " can be attributed to the formation of FeC0 3

2

4

3

4

4

2

3

3

4

4

4

3

3

Fe

2 +

2

+ C 0 " — > FeC0 3

3


(40)



456

Figure 8. Values for the pseudo first-order rate constants for the oxidation of Fe(II) with 0 in NaCl and NaC10 (I = 6.0 m) as a function of pH at 25°C. 2

4

7-1

1

1

1

1

•

j

1

1

•

1

1

1

1

•

1

1

j

1

r

·

1

el

1-1

OH 2

3

4

5

6

1 7

8

' 9

PH Figure 9. The effect of pH on the rate constant for the oxidation of Fe(II) with H 0 in seawater at 25°C. 2


2


34.

MILLERO


ο

Figure 10. The rate constant for the oxidation of Fe(II) with 0 in NaCl and seawater as a function of ionic strength at 25°C. 2


457

Melchior and Bassett; Chemical Modeling of Aqueous Systems II ACS Symposium Series; American Chemical Society: Washington, DC, 1990. 2

Figure 11. The effect of various ions on the oxidation of Fe(II) with 0 at I = 0.7 at 2 5 ° C .


2

S3

ci

ο

ο

r 2: ο

ο

r

w

a

00

en

34.

MILLERO

459


17.5

17.0

16.5


ο 16.0

15.5

15.0 3.0

Figure 12. The rate constant for the oxidation of Fe(II) with 0 in NaCl and NaC10 solutions as a function of ionic strength at 25°C. 2

4

16.5

16.0H

Ο

15.5

o 15.0-k o • + • Δ

14.5-

14.0 0.0

0.2

0.4

0.6

0.8

NaBr NaN0 Na2S04 NaHC0 NaB(0H)4 3

3

1.0

[x] Figure 13. The effect of various anions (X) on the oxidation of Fe(II) with 0 in NaClNaX solutions at I = 1.0 and 25°C. 2



460

l sn o t

which has a faster rate of oxidation than Fe(OH) . Since the value of FeCOo available at the present time, it is not possible to make a reliable estimate of the rate of oxidation of F e C 0 . The slight increase upon the addition of B r , N0 ~ and C10 " is related to the formation of F e C l ion pairs. The other anions have weaker interactions of these anions with F e compared to CI". The strong decreases in the rates due to the addition of S 0 " and B(OH) " can be attributed to the formation of FeS0 and F e B ( O H ) ion pairs. If one assumes that the ion pairs cannot be oxidized, the decrease in the rates can be given by 2

3

+

3

4

2 +

2

4

4

+

4

4

41

KCLAX = ι + / W * - ]

( ) =

anc

=

The experimental values of K a / k give log /?FeS0 ^ -0-1 * 1°8 /%eB(OH) 3.2 ± 0.1 at I = LO and 25°C. The value for FeS0 is in good agreement with the literature infinite dilution value (log / 3 ^ = 2.1 (19). To the best of our knowledge, values are not available for /?FeB(OH) - Our estimates, however, are in reasonable agreement with the /?'s for the formation of C u and Pb " complexes with B(OH) " (log β = 3.7 and 2.3, respectively (20). In conclusion, our recent studies of the rates of oxidation of Cu(I) and Fe(II) clearly demonstrate how the formation of ion complexes can affect the rates. Future rate studies for other metals are needed to examine how changes in the composition of natural waters can affect their rates of oxidation and reduction. x

4

4

4


FeS0

4

2 +

2

4

ACKNOWLEDGMENT The author wishes to acknowledge the support of the Office of Naval Research (N00014-87-G0116) and the Océanographie section of the National Science Foundation (OCE86-00284) for this study. LITERATURE CITED 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20.

Moffett, J.W.; Zika, R.G. Mar. Chem. 1983, 13 239-51. Millero, F.J. Geochim. Cosmochim. Acta 1985, 49, 547-53. Millero, F.J.; Sotolongo, S.; Izaguirre, M. Geochim. Cosmochim. Acta 1987, 51, 793-801. Moffett, J.W.; Zika, R.G. Environ. Sci. Technol. 1987, 21, 804-10. Millero, F.J.; Izaguirre, M.; Sharma, V.K. Mar. Chem. 1988, 22, 179-91. Sharma, V.K.; F.J. Millero. Environ. Sci. Technol. 1988, 22, 768-71. Sharma, V.K.; Millero, F.J. J. Solution Chem. 1988, 17, 581-99. Sharma, V.K.; Millero, F.J. J. Inorg. Chem. 1988, 27, 3256-59. Sharma, V.K.; Millero, F.J. Geochim. Cosmochim. Acta, in press. Stumm, W.; Lee, F.F. Environ. Sci. Technol. 1961, 53, 143-6. Millero, F.J.; Hubinger, S.; Fernandez, M.; Garnett, S. Environ. Sci. Technol. 1987, 21, 439-43. Millero, F.J. Geochim. Cosmochim. Acta 1987, 51, 351-53. Zika, R.G.; Saltzman, E.S.; Cooper, W.J. Mar. Chem. 1985, 17, 265-75. Zika, R.G.; Moffett, J.W.; Cooper, W.J.; Petasne, R.G.; Saltzman, E.S. Geochim. Cos mochim. Acta 1985, 49, 1173-84. Millero, F.J.; Sotolongo, S. Geochim. Cosmochim. Acta, in press. Millero, F.J.; Izaguirre, M. J. Solution Chem. 1988, 17, 581-99. Tamura, H.; Goto, K.; Nagayama, M. J. Inorg. Nucl. Chem. 1976, 38, 113-117. Sung, W.; Morgan, J J. Environ. Sci. Technol. 1980, 14, 561-68. Kester, D.R.; Byrne, R.H.; Liang, Y. In Marine Chemistry of the Coastal Environment; Church, T.M., Ed.; American Chemical Society: Washington, DC, 1975; ρ 56. van den Berg, M.G. Geochim. Cosmochim. Acta 1984, 48, 2613-17.

RECEIVED July 20, 1989


Chemical Modeling of Aqueous Systems II - American Chemical Society

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