Aqueous Ferrate(V) and Ferrate(IV) in Alkaline Medium: Generation

Chapter 10

Aqueous Ferrate(V) and Ferrate(IV) in Alkaline Medium: Generation and Reactivity 1

2

Diane E. Cabelli and Virender K. Sharma Downloaded by UNIV OF OKLAHOMA on April 27, 2013 | http://pubs.acs.org Publication Date: July 25, 2008 | doi: 10.1021/bk-2008-0985.ch010

1

Chemistry Department, Brookhaven National Laboratory, Upton, NY 11973-5000 (email: [email protected]) Chemistry Department, Florida Institute of Technology, 150 West University Boulevard, Melbourne, FL 32901 *Corresponding author: [email protected] 2

This chapter reviews the generation of ferrate(V) and ferrate(IV) complexes in basic solutions. Ferrate(V) (Fe 0 ) is easily produced by the one-electron reduction of the relatively stable Fe O ion. Comparatively, generation of a ferrate(IV) complex via one-electron oxidation of Fe(III) is rather difficult, due to the relative insolubility of Fe(III) hydroxides and the slow oxidation rate. This has resulted in limited studies of the reactivity of ferrate(IV). The most studied aquated ferrate(IV) complex is ferrate(IV)-pyrophosphate. The reactivity of ferrate(IV) and ferrate(V) complexes with inorganic and organic substrates in alkaline solution is presented. The reactions of ferrate(IV)-pyrophosphate complex with pyrophosphate complexes of divalent metal ions are likely occurring through inner-sphere electron transfer. Ferrate(V) reacts with substrates predominantly via a twoelectron transfer process to Fe(III). The only known example of one-electron reduction of ferrate(V) is its reactivity with cyanide in which sequential reduction of Fe(V) to Fe(IV) to Fe(III) was demonstrated. The reaction of Fe(V) with cyanide thus provides an opportunity for selective and unambiguous production of quantitative amounts of Fe(IV) in aqueous media. v

3-

4

VI

2-

4

158

© 2008 American Chemical Society In Ferrates; Sharma, V.; ACS Symposium Series; American Chemical Society: Washington, DC, 2008.

159

Downloaded by UNIV OF OKLAHOMA on April 27, 2013 | http://pubs.acs.org Publication Date: July 25, 2008 | doi: 10.1021/bk-2008-0985.ch010

Introduction The higher oxidation states of iron (Fe(VI), Fe(V), and Fe(IV)) have been shown to be strongly oxidizing in enzymatic systems, where they can carry out aliphatic hydrogen abstraction (7-5). In addition, they have been postulated as intermediates in Fenton-type systems (6). Fe(VI) itself is relatively stable and has been shown to have potential as an oxidant in the so-called "green" treatment of polluted waters (7-9). By contrast, Fe(V) and Fe(IV) are relatively short-lived transients when produced in aqueous solution in the absence of strongly bonding ligands other than hydroxide, a feature that has limited studies of its reactivity. There has been an additional study suggesting that Fe(VI) might be useful in battery design (10-12). Finally, a very interesting study suggested that ferrate may be possible to oxidize insoluble chromium to chromate and thus serve to remove chromium contamination in the Hanford radioactive waste tanks (73). This paper summarizes the properties and reactivities of ferrate(V) and ferrate(IV) in alkaline medium.

Ferrate(V) We have been interested in the study of the reactivity of high valent aquo iron states, ferrate(V) with inorganic and organic systems as a tool for the degradation of pollutants in water. Ferrate(V) can be generated very easily in the presence of excess ferrate(VI) through the use of reducing radicals produced in pulse radiolysis according to the following scheme (14,15), where ROH is H 0 -/WW H(0.55), e '(2.65), OH(2.75), H O (0.72), H (0.45) N 0 + e - + H 0 -»OH + OH" + N H + OH" -> e ' + H 0 OH/O' + ROH -> H 0/OH- + ROH Fe(VI) + ROH -> Fe(V) + Product k = 9x 10 M"V (14) 2

aq

2

aq

2

2

2

I (1) (2) (3) (4)

2

2

aq

2

2

9

4

an alcohol (e.g. ethanol, isoproponol ortert-butanol)that will react with the OH radical to form a simple carbon-centered radical. The reaction of ferrate(VI) with the parent alcohol used to generate the alcohol radical that reduces ferrate(VI) (reaction 3) or with some other substrate of interest requires the use of a premix device. The premix pulse radiolysis apparatus consists of two glass syringes mounted in a double syringe drive. The ferrate(VI) solution in one syringe is mixed with the alcohol solution in the other syringe. The mixed solution is promptly injected into the optical cell and exposed to an ionizing pulse. In Figure 1, we show the visible spectral features of Fe(VI) (Fe 0 ") and Fe(V) (Fe 0 "). The ferrate(V) spectrum has a maximum at 380 nm (6 80nm 1460 M^cm"). This spectrum undergoes a blue shift with decreasing pH. Ferrate(V) absorbs much strongly in the UV region ((£ onm ~ 5000 M cm ). The overall decay of ferrate(V) in the acidic to basic pH range suggest three protonated forms (eqs 5-7) (16,17). vl

2

4

3

3

4

1

_1

27

In Ferrates; Sharma, V.; ACS Symposium Series; American Chemical Society: Washington, DC, 2008.

=

160 v

+

5.5 < pK > 6.5 pK *7.2 pK = 10.1

v

H Fe 0 o H + H Fe 0 ' H F e 0 ' » H + HFe 0 " 3HFe 0 - o H + Fe O„ 3

4

2

v

2

v

4

v

2

6

4

2

+

(5) (6) (7)

s

4

+

v

7

4

The decay of the ferrate(V) species is strongly dependent on pH (16,17), where the totally deprotonated Fe 0 * decays to a longer lived transient (ti ~ seconds) via afirst-orderprocess (reaction 8). However, as the pH is lowered, ferrate(V) disappears by second-order kinetics to form ferric ions and hydrogen peroxide (16). Scheme I describes the mechanism associated with the disappearance of ferrate(V) in alkaline medium, the mechanism takes deprotonated Fe(V) (Fe 0 ) all the way to Fe(III) (reactions 8-10). The second-order rate constant observed in the disappearance of Fe(V) (reactions 11 and 12) increases as the pH is lowered and is of the order of 10 M ' V (16). The pH effect can be explained by reaction between F e 0 ' and its conjugate acid, HFe 0 " or bimolecular dimerization of the monoprotonated species reaction (16). 3

4

/2


4

7

v

1

3

v

4

500

400

2

4

700

600

wavelength Figure 1. Visible spectra of Fe(V) and Fe(VI) in aqueous basic solution (14)

Scheme I First-OrderDecay: Fe 0 " + 2H 0 -> [Fe (;OH) (0) ] [Fe^(OH) (0) ] - -> [Fe (OH) (0 -)] " [Fe (OH) (0 -)f - Fe (OH) + H 0 v

3

v

4

3

2

4

3

4

2

2

HI

4

2

4

3

,rt

2

m

v

IV

2

2

HI

2

2Fe (0 ) 2

2Fe (0 -) + H 0 -> Feiii/(OH) + H 0 2

2

fast

lv

2

2

(9) (10)

2

Second-Order Decay: slow I IV 2 - F e =0 [-Fe (0 ")Fe -]

UI

(8)

2

ft,

,u

2

m

2

2


(id (12)

161 In the pH range from 3.6 to 7.0, the observed disappearance is predominantly first-order with rates that vary from 7.0 x 10V to 100 s" (Scheme II) (77). The first order process is described by the hydration of the tetrahedral ferrate(V) to a six-coordinate octahedral species prior to loss of peroxide. This aquation reaction dominates at the low and high pH range and competes with the dimerization when FeV) is either unprotonated or monoprotonated. 1

1

Scheme II First-Order Decay: H F e 0 ' + 2H +4H 0 -> Fe(OH) (H 0) + H 0 H Fe 0 " + H + 4H 0 -+ Fe(OH) (H 0) + H 0 v

2


l

k*5s £ - 1 5 0 s'

+

4

2

v

3

2

3

2

2

1

+

2

4

2

3

2

3

2

2

v

H Fe 0 (r ) + H 0 = (Fe(OH) ) (O ) 3

4

d

2

5

+

aq

h

+

(Fe(OH) ) + H = (Fe(OH) ) (Fe(OH) ) + H - Fe (aq) + H 0 5

aq +

4

4

+

aq

aq

m

2

2

(13) (14) (15) (16) (17)

By using a fast premixing apparatus synchronized with the accelerator pulse, experiments can be carried out to measure the reactivity of ferrate(V) with a substrate. If the concentration of the substrate is increased such that the substrate now reacts with the OH radical, then the reaction between ferrate(VI) and an oxidized substrate radical can be measured (reaction 18). Finally, conditions can be adjusted by saturating the solution with nitrogen instead of N 0 and eliminating the addition of alcohol. Here, the electron reduces the ferrate(VI) to ferrtae(V) and the substrate is oxidized by the OH radical (reactions 18 and 19). 2

Fe(VI) + e - -> Fc(V) S + O H -> S + H 0/OH aq

ox

_

2

(18) (19)

Table 1 highlight the reactivity of ferrate(V), which shows the kinetic parameters in the use of high oxidation state iron as an oxidant for inorganic and organic pollutants such as cyanides, sulfur-containing, carboxylic acids, and aromatic compounds (18). Gold is leached from the cyanide complex by applying soluble copper, which forms stronger complexes with cyanide than gold. The resulted copper cyanide complexes are highly toxic to aquatic life and are problematic because they are much more stable thanfreecyanide. The presence of copper cyanide complexes in gold industry thus presents the biggest concern in cyanide management (19). The destruction of cyanide by oxidation with ferrates is thus of great interest. The experiments on the reduction of ferrate(V) by cyanide have demonstrated sequential one-electron reductions of ferrate(V) to ferrate(IV) to ferrate(III) in aqueous media, where the order of reactivity is &(ferrate(V)) > &(ferrate(IV)) > &(ferrate(VI)) (20). Interestingly, complexation of cyanide with copper(I) enhanced the rate of the oxidation reaction (Table 1). The results demonstrate potential of ferrates in destroying cyanides in waste water.


162

Table 1. Ferrate(V) oxidation of compounds at pH = 12.4 in N 0 saturated 0.1 M phosphate buffer at 23-24 °C. 2

Pollutant (R) Cyanides Cu(I) cyanide Cyanide Thiocyanate

1

2


k, M-'s

Formula

7

1.35±0.02x 10 1.96±0.20x 10 6 . 3 7 ± 0 . 1 3 x 10

3

Cu(CN) " HCN, CNSCN" 4

4

1

Sulfur-containing compounds Cysteine HSCH CH(NH )COOH Cystine HOOCCH(NH )CH SSCH (NH )CHCOOH Thiourea NH CSNH Methionine CH SCH CH CH(NH )COO' 2

2

2

2

1

2

9

4.00 ± 0.80 x 10 1.95 ± 0.02 x 10

4

2

2

2

2

3

2

2

2

3

8.10 ± 0.40 x 10 1.58 ± 0.09 x 10

3

Carboxylic acids Glycine Alanine Aspartic Ketomalonic Tartaric Glycolic Malic Lactic Malonic Succinic Acetic

CH (NH )COOCH CH(NH )COOHOOCCH CH (NH )COO" C(OH) (COOH) HOOC(CHOH) COOH HOCH COOH HOOCCH(OH)CH COOH CH CH(OH)COOH CH (COOH) HOOCCH CH COOH CH COOH

(8.4 (3.1 (2.6 (1.4 (3.1 (7.2 (1.7 (1.6 (9.2 (2.0 (1.6

Aromatic compounds Histidine Phenylalanine Tyrosine Tryptophan Phenol Proline

C H N CH CH(NH )COO" C H CH CH(NH )COOHOC H CH CH(NH )COOC H NCH CH(NH )COOC H OH C H NCOO-

22.2 ± 0 . 1 x 10 9.5 ± 0.4 x 10 8.1 ± 0 . 2 x 10 9.3 ± 0.4 x 10 3.8 ± 0.4 x 10 0.1 ±0.01 x 10

1

2

3

+

2

3

+

3

3

+

2

2

3

2

2

2

2

2

3

2

2

2

2

3

3

6

3

2

5

6

6

5

4

7

2

2

6

8

2

2

4

2

2

2

2

± 0.6) x ±0.2)x ± 0.1) x ± 0.2) x ±0.2) x ± 1.0) x ± 0.2) x ± 0.2) x ± 1.0) x ± 0.2) x ± 0.2) x

AtpH= 12.2 At pH= 11.2 At pH = 9.0


6

10 10 10 10 10 10 10 10 10 10' 10

6

6

6

3

2

2

2

1

1

6

6

6

6

6

6


163 Thiocyanate (SCN") is used in processes such as the manufacture of thiourea, metal separation and electroplating as well as being formed in mining wastewater (27). Ferrate(V) reacts with thiocyanate via a two-electron pathway (22) with no observable ferrate(IV) formation. As has been seen generally, the reaction of Fe(V) with SCN' is significantly more rapid than that of Fe(VI) with SCN"; with a rate constant that is over two orders of magnitude faster (22). This suggests that ferrate(VI) oxidations may be accelerated in the presence of one-electron or two-electron reducing substrates, where the reactive high valent iron species is either ferrate(V) or ferrate(IV). Cysteine is the most reactive sulfur-containing compound (23) and its rate constants is three orders of magnitude higher than other compounds (Table 1). The reactivity of ferrate(VI) with sulfur compounds is determined by the nucleophilicity of the sulfiir atom in compounds. The sulfur center in cysteine is most readily oxidized by ferrate(V). The rate of oxidation of carboxylic acids by ferrate(V) varies with the nature of the substituent group at the -carbon atom of the acids, with rate constants ranging from 10 - 10 M' s" (Table 1). The rate constants decrease in the order ofa-NH > a-C-OH > a-C-H (24). 1

6

l

1

2

ferrate(V) + carboxylic acids -> Fe(III) + N H + a-keto acids

II

3

5

7

Fe(V) oxidation of aromatic compounds proceeds at rates 10 - 10 M'V (Table 1). A diamino dicarboxylic acid, histidine, reacts faster than other aromatic compounds (Table 1). 1

Ferrate(IV) 4

Iron(IV) complex with simple inorganic ligand, P 0 " in basic medium can be generated from the corresponding parent complexe by oxidation with OH/0" radical in aqueous solutions (reaction 20) (25). 2

3+

6

IV

6

[(P 0 ) Fe OH] " + OH -> [(P 0 ) Fe 0] " + H 0 2

7

2

2

7

2

2

7

7

A = 7.8 x 10 M'V

(20)

1

The spectrum of Fe(IV)-complex is shown in Figure 2. The spectrum has a peak at ^ = 430 nm (e = 1200 M ' W ) . The spectra of FeO(OH) , t(P20 ) Fe 0] ", and L Fe(IV) were found similar. On lowering the pH, 430 nm peak undergoes blue shift. The pyrophosphate complex of iron(IV), formed at pH > 10 is short lived (t = 100-600 msec) (25). This complex of iron(IV) disappears by a second-order process to form a Fe(III) pyrophosphate complex and molecular oxygen (reaction 21) (25). 1

2 n

n

IV

7

6

2

m

1/2

IV

m

6

2 [(P 0 ) Fe Of + 2 H 0 - » 2 [(P 0 ) Fe OH] " + Vi 0 2

7

2

2

2

7

2

2

(21)

Reactions of Fe(IV) with pyrophosphate complexes of divalent transition metal ions (reaction 22) have been studied (25). M(II) represents a divalent



164

pyrophosphate complex of a metal without assignment of the number of ligand L in a given complex. Fe(IV) + M(II) — Fe(III) + M(II)

(22)

The kinetics of these reactions was investigated by pulse irradiating solutions of Fe(III) pyrophosphate, which contained pyrophosphate complexes of M n , Fe , or Co . The rate constants are listed in Table 2. The Ni(II) and Cu(II) pyrophosphate complexes showed no reactivity and their upper limit rate constants are given in Table 2. 2+

2+

2+

Table 2. Reactivity of Fe(IV) pyrophosphate with some divalent pyrophosphate complexes ( L M ; L=P 0 ') at pH 10.0 (25). 2+

4

m

Metal Mn Fe Co

2+

2+

3+

2

Metal

1

k, M-'s"

6

2+

1.2 x 10 1.6 x 10 5.5 x 10

Ni Cu

6

2+

7

k, M-'s

1

2

Aqueous Ferrate(V) and Ferrate(IV) in Alkaline Medium: Generation

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