Characterization and oxidation of chromium(III) by sodium hypochlorite in alkaline solutions

Article abstract of DOI:10.1016/j.ica.2006.03.035

Chromium exists in nuclear waste sludges and is a problematic element in the vitrification process of high-level nuclear wastes. It is therefore necessary to treat the waste sludges to remove chromium prior to vitrification, by caustic leaching or oxidation of Cr(III) to Cr(VI). The objective of this study is to investigate the effect of oligomerization of Cr(III) on its oxidation by hypochlorite in alkaline solutions. Monomeric, dimeric and trimeric Cr(III) species in solution were separated by ion exchange. The kinetics of the oxidation of the separated species by hypochlorite in alkaline solutions was studied by UV/Vis absorption spectroscopy, and compared with the oxidation by hydrogen peroxide previously studied. Results indicate that hypochlorite can oxidize Cr(III) to Cr(VI) in alkaline solutions, but the rate of oxidation by hypochlorite is slower than that by hydrogen peroxide at the same alkalinity and concentrations of oxidants. The rate of oxidation of Cr(III) by both oxidants decreases as the concentration of sodium hydroxide is increased, but the oxidation by hypochlorite seems less affected by the degree of oligomerization of Cr(III) than that by peroxide. Compared with the oxidation by hydrogen peroxide where the major reaction pathway has an inverse order with respect to C_NaOH, the oxidation by hypochlorite has a significant reaction pathway independent of [OH^-].

Full text of DOI:10.1016/j.ica.2006.03.035

Inorganica Chimica Acta 359 (2006) 3237–3242
www.elsevier.com/locate/ica
Characterization and oxidation of chromium(III) by
sodium hypochlorite in alkaline solutions
a,
b
Huijian Jiang a,1, Linfeng Rao *, Zhicheng Zhang a,2, Dhanpat Rai
a
Lawrence Berkeley National Laboratory, Berkeley, CA 94720, USA
Pacific Northwest National Laboratory, Richland, WA 99352, USA
b
Received 4 January 2006; received in revised form 9 March 2006; accepted 24 March 2006
Available online 3 April 2006
Abstract
Chromium exists in nuclear waste sludges and is a problematic element in the vitrification process of high-level nuclear wastes. It is there-
fore necessary to treat the waste sludges to remove chromium prior to vitrification, by caustic leaching or oxidation of Cr(III) to Cr(VI). The
objective of this study is to investigate the effect of oligomerization of Cr(III) on its oxidation by hypochlorite in alkaline solutions.
Monomeric, dimeric and trimeric Cr(III) species in solution were separated by ion exchange. The kinetics of the oxidation of the sep-
arated species by hypochlorite in alkaline solutions was studied by UV/Vis absorption spectroscopy, and compared with the oxidation by
hydrogen peroxide previously studied. Results indicate that hypochlorite can oxidize Cr(III) to Cr(VI) in alkaline solutions, but the rate
of oxidation by hypochlorite is slower than that by hydrogen peroxide at the same alkalinity and concentrations of oxidants. The rate of
oxidation of Cr(III) by both oxidants decreases as the concentration of sodium hydroxide is increased, but the oxidation by hypochlorite
seems less affected by the degree of oligomerization of Cr(III) than that by peroxide. Compared with the oxidation by hydrogen peroxide
where the major reaction pathway has an inverse order with respect to CNaOH, the oxidation by hypochlorite has a significant reaction
pathway independent of [OHꢀ].
ꢀ 2006 Elsevier B.V. All rights reserved.
Keywords: Oxidation; Chromium(III); Oligomerization; Reaction rate; Hypochlorite
1. Introduction
cost of nuclear waste management, pretreatment tech-
niques must be developed to remove chromium from
HLW before vitrification.
Chromium in nuclear waste sludges is a problematic
constituent in the vitrification process of high-level nuclear
wastes (HLW). At the operating temperature of the glass
melter (1050–1150 ꢁC), even relatively small amounts of
chromium can result in the formation of a refractory spi-
nel-type crystalline solid phase. The formation and settling
of crystallites during the vitrification can adversely affect
the process operation. Consequently, the allowable limits
for chromium in the melter feed must be low. To decrease
the volume of vitrified glass and significantly reduce the
The nuclear wastes stored in underground tanks at Han-
ford, Washington State of the United States are alkaline
and consist of supernatant, salt cake and sludge. The
sludge, containing most of the actinides as well as chro-
mium (mostly Cr(III)), is the major source of the HLW.
Different strategies are proposed to remove chromium
from the sludge, including oxidative leaching under alka-
line conditions. However, leaching results of actual tank
sludge have indicated that the efficiency of removal is not
as high as expected, due to a poor understanding of the
chemical behavior of chromium in alkaline solutions. To
develop efficient pretreatment strategies for the removal
of chromium, systematic studies are needed to reveal the
chemical speciation and its effect on the redox reactions
of chromium in alkaline solutions.
*
Corresponding author.
E-mail address: LRao@lbl.gov (L. Rao).
Present address: A&Z Pharmaceutical Inc., 180 Oser Ave, Hauppauge,
1
NY 11788, USA.
2
Present address: Washington State University, Pullman, WA 99164,
USA.
0020-1693/$ - see front matter ꢀ 2006 Elsevier B.V. All rights reserved.
doi:10.1016/j.ica.2006.03.035

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H. Jiang et al. / Inorganica Chimica Acta 359 (2006) 3237–3242
It is known that hydrolysis of Cr(III) leads to formation
increasingly higher ionic strength (from 1 to 4 mol dmꢀ3
NaClO4) and HClO4 concentrations (from 0.01 to
0.04 mol dmꢀ3). The solutions of the separated Cr(III) spe-
cies were either immediately used for characterization and
redox studies or stored in a refrigerator prior to use. UV–
Vis absorption spectra of the solutions were obtained on a
Cary-5G spectrometer.
of various oligomers [1–8]. Recent studies have shown that
the degree of oligomerization, while depending on alkalin-
ity, temperature and aging time [9,10], affects the kinetics of
oxidation of Cr(III) by hydrogen peroxide [9] and persul-
fate [11]. Hydrogen peroxide and persulfate could oxidize
Cr(III) to Cr(VI) in alkaline solutions, but each has own
shortcomings. Hydrogen peroxide has a strong tendency
to decompose and generate gaseous products under alka-
line conditions. Persulfate, on the other hand, may gener-
ate sulfate that causes problems in glass vitrification and
increases the volume of nuclear wastes. Therefore, more
oxidants need to be investigated. In this paper, we report
the results of oxidation of Cr(III) by hypochlorite, another
oxidant that has wide applications in water treatment.
The dissolved species of Cr(III) hydroxide of different
degrees of oligomerization were separated by an ion
exchange chromatography technique developed previously
[3–10] and characterized by UV/Vis absorption spectros-
copy. Kinetic experiments were conducted to determine
the rate of oxidation of Cr(III) by hypochlorite in alkaline
solutions. The rate of oxidation by hypochlorite is com-
pared with that by hydrogen peroxide. The kinetic data,
in conjunction with the information from characterization,
help to reveal the effect of oligomerization on the rate of
oxidation of Cr(III).
2.3. Kinetic study of oxidation of Cr(III) by hypochlorite in
alkaline solutions
The oxidation kinetics were studied by UV/Vis absorp-
tion spectroscopy on an Ocean Optics ST 2000 or a Cary-
5G spectrometer, with most experiments carried out on the
former instrument. Appropriate amounts of NaClO and
NaOH were first mixed in a 1.0-cm cuvette to take the
background reading. Immediately after an aliquot of
Cr(III) solution was injected into the cuvette and rapidly
mixed, the absorbance of Cr(VI) at 372 nm (At) was mon-
itored as a function of time. When the reaction was com-
pleted, the absorbance at 372 nm was recorded as A .
1
Redox experiments were conducted with large excess of
sodium hypochlorite and hydroxide with respect to Cr
concentrations. At time t, [Crn(III)] (n = 1, 2 and 3 for the
monomer, dimer and trimer) was equal to ([Cr(VI)]1 ꢀ
[Cr(VI)]t)/n and proportional to (A1 ꢀ At)/n, where
[Cr(VI)]t and [Cr(VI)] are the concentrations of Cr(VI)
1
2. Experimental
at time t and at the completion of the reaction, respectively.
A statistical significance test [13] indicated that the kinetic
traces of ln(A1 ꢀ At) were essentially linear functions of
time, suggesting a first-order reaction with respect to
[Crn(III)]. The reaction rate is expressed as Eq. (1) and the
pseudo-first-order rate constant, kobs, were calculated from
the plots of ln(A1 ꢀ At) versus t.
2.1. Chemicals
All chemicals were reagent grade or higher and used
without further purification. Solutions were prepared with
deionized water from a Milli-Q system. Solutions of sodium
hydroxide were prepared from an oversaturated sodium
hydroxide solution and standardized by pH titration. Stock
solutions of Cr(III) were prepared by dissolving Cr(ClO4)3 Æ
6H2O (Aldrich) in water. UV–Vis absorption spectra of the
stock solutions were collected to confirm the absence of
Cr(VI) in the solutions. The Cr(III) solutions of different
acidity (to achieve different degrees of oligomerization)
were prepared by adding appropriate amounts of NaOH
solutions and stored in an inert-atmosphere (Ar) glove
box. The Cr(III) solutions used in the characterization
and the redox experiments were filtered prior to use, with
the 0.2 lm nylon Acrodisc filter (Gelman). Solutions of
sodium hypochlorite were prepared from a solution of
10% NaClO (Aldrich) and standardized by redox titrations
with sodium thiosulfate standard solutions [12] prior to use.
ꢀd½CrnðIIIÞꢁ=dt ¼ ð1=nÞd½CrðVIÞꢁ=dt ¼ kobs½CrnðIIIÞꢁ ð1Þ
The overall stoichiometry of the oxidation of Cr(III) by
hypochlorite was determined by using excess amounts of
sodium hypochlorite in the reaction mixture. After the oxi-
dation of Cr(III) to Cr(VI) was completed, the concentra-
tion of Cr(VI) was determined by UV–Vis absorption at
372 nm while the remaining hypochlorite was determined
by the titration method with sodium thiosulfate standard
solutions [12].
3. Results and discussion
3.1. Characterization of the oligomers of Cr(III) in solution
The optical absorption spectra of the separated Cr(III)
oligomer species in solution are shown in Fig. 1. As the
degree of oligomerization of Cr(III) increases, the absorp-
tion peaks are slightly red-shifted and intensified. The posi-
tions and intensities of the spectra are identical to those
observed in previous studies [9]. The structures of the
Cr(III) oligomers in these solutions have already been
determined by Extended X-Ray Absorption Spectroscopy
2.2. Separation and characterization of chromium(III)
oligomers
Chromium(III) oligomers were separated by an ion
exchange method using Sephadex SP C-25 cation exchange
resin (Aldrich) described in the literature [3–10]. The sepa-
ration was improved in this work by using eluents of

H. Jiang et al. / Inorganica Chimica Acta 359 (2006) 3237–3242
3239
reaction order with respect to [OH] was calculated to be
ꢀ1 or ꢀ0.5 with the two models, respectively. At present,
we do not have additional information to help select one of
the two models. However, the first model that results in an
integer reaction order (ꢀ1) appears more reasonable and it
is similar to the model used to interpret the oxidation of
Cr(III) by hydrogen peroxide [9]. Therefore, we elect to
include both pathways that are dependent on and indepen-
dent of [NaOH]. A general rate equation (Eq. (2)) can be sug-
gested for the oxidation of Cr(III) to Cr(VI) by hypochlorite.
The values of kOH, k and k0 are summarized in Table 1.
ꢀd½CrnðIIIÞꢁ=dt ¼ kobs½CrnðIIIÞꢁ
¼ kOH½CrnðIIIÞꢁ½NaClOꢁ
¼ k½CrnðIIIÞꢁ½NaClOꢁð1=½OHꢀꢁ þ k0Þ ð2Þ
3.2.2. Comparison of hypochlorite with hydrogen peroxide
Table 2 shows the comparison of the rate of oxidation of
Cr(III) by hypochlorite in this study with that by hydrogen
peroxide previously reported [9]. The results show the fol-
lowing similarities between the two oxidants: (1) both
hypochlorite and peroxide can oxidize Cr(III) to Cr(VI)
in alkaline solutions; (2) for both oxidants, the rate of oxi-
dation decreases as the concentration of sodium hydroxide
is increased. The results also show the following differences
between the two oxidants: (1) in general, the oxidation of
Cr(III) by hypochlorite is slower than that by peroxide at
the same alkalinity and concentrations of oxidants (see
the values of s in Table 2); (2) the reaction pathway inde-
pendent of [OHꢀ] is more significant in the oxidation by
hypochlorite than that by peroxide, suggested by the larger
k0 for hypochlorite; (3) the rate of oxidation by hypochlo-
rite seems less affected by the degree of oligomerization of
Cr(III) than that by peroxide.
Fig. 1. Optical absorption spectra of Cr(III) species separated by ion
chromatography: (top) monomer; (middle) dimer; (bottom) trimer.
(EXAFS) in the previous study [9]. Therefore, no EXAFS
experiments were repeated in this work. The structures of
the Cr(III) oligomers determined by EXAFS are also
shown in Fig. 1. Detailed discussions on the EXAFS data
and the structures are provided elsewhere [9].
The spectra of the Cr(III) stock solutions were checked
before and after the redox experiments that lasted for a few
months. No detectable changes were observed, confirming
that the separated Cr(III) species were stable during the
entire course of the redox experiments.
3.2.3. Overall reactions and possible mechanisms
It was determined in this study that three moles of
NaClO were consumed when two moles of Cr(III) were oxi-
dized. The stoichiometric ratio suggests that chloride,
instead of chlorine, is the end product of the reduction of
hypochlorite by Cr(III). However, such information is not
sufficient to reveal the actual reaction mechanism. At least
two possible mechanisms can be suggested: (1) hypochlorite
is reduced directly to chloride via a two-electron transfer
reaction; (2) hypochlorite is reduced to chlorine via a one-
electron transfer. The chlorine rapidly disproportionates
in alkaline solutions to chloride and hypochlorite and the
latter reacts with another equivalent of Cr(III). Both mech-
anisms would result in a stoichiometric ratio of 2:3 (Cr(III)/
hypochlorite). At present, we do not have other experimen-
tal evidence to distinguish between the two possible mecha-
nisms. Instead, we can only suggest the following overall
reactions between the monomer, dimer or trimer of Cr(III)
and hypochlorite in alkaline solutions. In the reactions, the
dominant species of hypochlorite is assumed to be ClOꢀ
because the pKa of HClO is about 7.4 [14].
3.2. Oxidation of Cr(III) by sodium hypochlorite
3.2.1. Reaction orders
Kinetic experiments were conducted at different concen-
trations of NaOH and NaClO under the pseudo-first-order
conditionswithrespecttoCr(III). Thepseudo-first-orderrate
constant (Eq. (1)), kobs, was calculated from the linear plot of
ln(A1 ꢀ At) versus t. The results are summarized in Table 1.
At constant CNaOH, kobs increased proportionately with
the increase of CNaClO as shown in Fig. 2, suggesting the reac-
tion is first-order with respect to CNaClO. The slopes of the
lines represent the values of kOH (Table 1). At constant
CNaClO, kOH decreased as the concentration of NaOH was
increased. Two models seem to interpret the results equally
well: one model assumes there are two reaction pathways
(dependent on or independent of [NaOH]); the other model
excludes the pathway that is independent of [NaOH]. The

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H. Jiang et al. / Inorganica Chimica Acta 359 (2006) 3237–3242
Table 1
Kinetic data for the oxidation of separated chromium(III) species by hypochlorite
CNaOH (mol dmꢀ3
)
CNaClO (mol dmꢀ3
)
kobs (minꢀ1
)
kOH ðminꢀ1 molꢀ1 dm3Þ
k (minꢀ1
)
kꢀ1 (molꢀ1 dm3)
ꢀ
Sample
Cr(III) monomer
0.104
5.60E ꢀ 03
1.12E ꢀ 02
2.80E ꢀ 02
5.60E ꢀ 02
1.12E ꢀ 02
2.80E ꢀ 02
5.60E ꢀ 02
1.12E ꢀ 01
2.98E ꢀ 02
5.60E ꢀ 02
1.12E ꢀ 01
1.68E ꢀ 01
2.98E ꢀ 02
5.60E ꢀ 02
1.12E ꢀ 01
1.68E ꢀ 01
0.0046
0.0089
0.0213
0.0477
0.835 0.021
0.368 0.003
0.282 0.003
0.260 0.006
0.212 0.004
0.068 0.002
0.070 0.004
0.058 0.006
2.0 0.2
0.250
0.500
0.750
1.00
0.0041
0.0107
0.0209
0.0407
0.0080
0.0149
0.0310
0.0476
0.0064
0.0141
0.0284
0.0454
5.03E ꢀ 02
1.12E ꢀ 01
1.68E ꢀ 01
2.23E ꢀ 01
0.0105
0.0260
0.0359
0.0462
Cr(III) dimer
0.104
0.250
0.500
0.750
1.00
1.12E ꢀ 02
2.79E ꢀ 02
5.58E ꢀ 02
1.12E ꢀ 01
1.12E ꢀ 02
2.79E ꢀ 02
5.58E ꢀ 02
1.12E ꢀ 01
2.79E ꢀ 02
5.58E ꢀ 02
1.12E ꢀ 01
1.68E ꢀ 01
5.58E ꢀ 02
1.12E ꢀ 01
1.68E ꢀ 01
2.23E ꢀ 01
0.010
0.022
0.048
0.104
0.909 0.019
0.542 0.026
0.361 0.005
0.282 0.005
0.246 0.002
2.7 0.2
0.0049
0.0091
0.029
0.063
0.010
0.019
0.042
0.060
0.016
0.031
0.050
0.061
5.58E ꢀ 02
1.12E ꢀ 01
1.68E ꢀ 01
2.23E ꢀ 01
0.013
0.028
0.041
0.055
Cr(III) trimer
0.104
0.250
0.500
0.750
1.12E ꢀ 02
2.79E ꢀ 02
5.58E ꢀ 02
1.12E ꢀ 01
1.12E ꢀ 02
2.79E ꢀ 02
5.58E ꢀ 02
1.12E ꢀ 01
2.79E ꢀ 02
5.58E ꢀ 02
1.12E ꢀ 01
1.68E ꢀ 01
5.58E ꢀ 02
1.12E ꢀ 01
1.68E ꢀ 01
2.23E ꢀ 01
0.010
0.020
0.040
0.055
0.704 0.050
0.440 0.005
0.301 0.007
0.209 0.002
2.9 0.5
0.006
0.012
0.025
0.049
0.010
0.015
0.035
0.050
0.012
0.023
0.036
0.046

H. Jiang et al. / Inorganica Chimica Acta 359 (2006) 3237–3242
3241
Table 1 (continued)
Sample
CNaOH (mol dmꢀ3
1.00
CNaClO (mol dmꢀ3
)
kobs (minꢀ1
)
kOH ðminꢀ1 molꢀ1 dm3Þ
k (minꢀ1
)
kꢀ1 (molꢀ1 dm3)
ꢀ
)
5.58E ꢀ 02
1.12E ꢀ 01
1.68E ꢀ 01
2.23E ꢀ 01
ꢀ
0.0113
0.0218
0.0337
0.0472
0
0.206 0.003
Ccr = 1.42 · 10ꢀ4 mol dmꢀ3. The error limits for kOH , k, and k represent the standard deviations at 95% confidence level. ꢀd[Crn(III)]/dt =
kobs[Crn(III)] = kOH[Crn(III)][NaClO] = k[Crn(III)][NaClO](1/[OHꢀ] + k0).
Fig. 2. Pseudo first-order rate constant, kobs, as a function of CNaClO at constant CCr (1.42 · 10ꢀ4 mol dmꢀ3) and different CNaOH: (a) monomer; (b) dimer; (c)
trimer. Symbols: CNaOH (mol dmꢀ3) = 0.104 (+), 0.250 (m), 0.500 (d), 0.750 (.), 1.00 (r). Detailed conditions of the experiments are provided in Table 1.
Table 2
Kinetic data for the oxidation of Cr(III) by hypochlorite and hydrogen peroxide
Cr(III) species
Hypochlorite
k (minꢀ1
Hydrogen peroxide [9]
)
k0 (molꢀ1 dm3)
sa (min)
k (minꢀ1
)
k0 (molꢀ1 dm3)
sa (min)
Monomer
Dimer
Trimer
0.068 0.002
0.070 0.004
0.058 0.006
0.0048 0.0026
2.0 0.2
2.7 0.2
2.9 0.5
2.6 4.3
490
390
440
327
6
0.07 0.09
0.0 0.1
1.34 0.04
0.53 0.01
0.3
9
32
11.3 0.3
1.33 0.04
0.032 0.002
Unseparated and aged in 8 M NaOH
5800
2040
The error limits in brackets for k, and k0 represent 3r.
a
Reaction lifetime s = 1/kobs. Values of s were calculated for [oxidant] = 0.01 M and [OHꢀ] = 1 M.
2CrðOHÞ4ꢀ þ 3ClOꢀ þ 2OHꢀ ! 2CrO42ꢀ þ 3Clꢀ þ 5H2O
to oxo-bridges through deprotonation and/or dehydration.
By conductometric titrations, they concluded that addition
of sodium hydroxide could cause a large increase in the
number of oxo-bridges formed [15].
ð3Þ
Cr2O2ðOHÞ42ꢀ þ 3ClOꢀ þ 2OHꢀ ! 2CrO42ꢀ þ 3Clꢀ þ 3H2O
ð4Þ
Though the overall reactions (Eqs. (3)–(5)) all contain
hydroxide on the reactant side, it does not mean that
hydroxide would necessarily facilitate the oxidation of
Cr(III) by hypochlorite in alkaline solutions. In fact, the
rate of oxidation actually decreases as CNaOH is increased
(Table 1). This observation could suggest that, either the
rate-determining step involves the weakening and release
of one hydroxyl group from the Cr(III) moiety upon the
attack by hypochlorite, or presumably the activated com-
plex between Cr(III) and hypochlorite involves chromium
2Cr3O4ðOHÞ43ꢀ þ 9ClOꢀ þ 6OHꢀ ! 6CrO42ꢀ þ 9Clꢀ þ 7H2O
ð5Þ
The formula of Cr(III) species in the above equations are
postulated based on the structures in acidic solutions
(Fig. 1), taking into consideration the likely conversion of
hydroxo-bridges to oxo-bridges in alkaline solutions. As
pointed out by Hall and Eyring [15], higher concentrations
of hydroxide facilitate the conversion of hydroxo-bridges

3242
H. Jiang et al. / Inorganica Chimica Acta 359 (2006) 3237–3242
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species that are protonated at one of the hydroxyl groups.
Alternatively, the observed decrease of reaction rate at
higher CNaOH could imply that the redox reaction occurs
through intermediate states involving the protonated spe-
cies of hypochlorite. Further mechanistic studies are
needed to reveal the reaction mechanism.
[7] L. Spiccia, W. Marty, Inorg. Chem. 25 (1987) 266.
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Acknowledgements
[9] L. Rao, Z. Zhang, J.I. Friese, B. Ritherdon, S.B. Clark, N.J. Hess,
D. Rai, J. Chem. Soc., Dalton Trans. (2002) 267.
[10] J.I. Friese, B. Ritherdon, S.B. Clark, Z. Zhang, L. Rao, D. Rai,
Anal. Chem. 74 (2002) 2977.
This work was supported by the US Department of
Energy, Office of Science, Office of Biological and Environ-
mental Remediation, under Contract No. DE-AC02-
05CH11231 at Lawrence Berkeley National Laboratory.
The authors express gratitude to the two anonymous review-
ers whose comments have helped to improve the manuscript.
[11] J.I. Friese, et al., unpublished data.
[12] A.I. Volgel, Vogel’s Textbook of Quantitative Inorganic Analysis,
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