Cuprous copper [Cu(I)] reacts with sodium persulfate (PDS) to generate sulfate radical SO4•, but it has been seldom investigated owing to its instability and difficulty in dissolving it. This study proposes a new method to regenerate Cu(I) from cupric copper [Cu(II)] by addition of hydroxylamine (HA) to induce the continuous production of radicals through active PDS, and investigates the resulting enhanced methyl orange (MO) degradation efficiency and mechanism in the new system. HA accelerated the degradation of MO markedly in the pH range from 6.0 to 8.0 in the HA/Cu(II)/PDS process. Both SO4• and hydroxyl radicals (•OH) were considered as the primary reactive radicals in the process. The MO degradation in the HA/Cu(II)/PDS process can be divided into three stages: the fast stage, the transitory stage, and the low stage. MO degradation was enhanced with increased dosage of PDS. Although high dosage of HA could accelerate the transformation of the Cu(II)/Cu(I) cycle to produce more reactive radicals, excess HA can quench the reactive radicals. This study indicates that through a copper-redox cycling mechanism by HA, the production of SO4• and •OH can be strongly enhanced, and the effective pH range can be expanded to neutral conditions.

INTRODUCTION

In situ chemical oxidation (ISCO) is a class of processes that have been developed over the past 20 years to treat soil and groundwater contaminated with organic pollutants (Al-Shamsi & Thomson 2013; Yan & Lo 2013). Of these ISCO processes, activated persulfate (PDS) is attractive due to its ability to generate powerful oxidizing sulfate radicals (SO4•), which can induce the production of another powerful and nonselective oxidizing species, hydroxyl radicals (•OH) (Liang & Su 2009). PDS anion (S2O82−, E0 = 2.01 V) is a strongly oxidizing agent, which is usually stable, and can be activated by heat, UV light, ultrasound, and transition ions (Mn+) to generate sulfate radicals (SO4•), as respectively described by Equations (1) and (2) (Anipsitakis & Dionysiou 2004; Adewuyi 2005; Waldemer et al. 2007). However, the process is not yet mature and its PDS activation properties deserve detailed study. 
formula
1
 
formula
2
Currently, many transition ions have been investigated for their PDS activation properties to produce SO4•, but not all the transition ions are efficient (Anipsitakis & Dionysiou 2004). Copper, one of the major redox-active transition metals (Balaz et al. 2002), is also considered to be one of the most efficient ions (Woods & Kolthoff 1965; Kolthoff & Woods 1966), but copper-catalyzed production of SO4• is seldom reported.

Copper typically occurs in either the cuprous (Cu(I)) or the cupric (Cu(II)) oxidation state in natural waters. However, naturally occurring copper is mostly in the Cu(II) oxidation state, because Cu(I) is unstable and easily oxidized to Cu(II) by oxygen or other oxidants, in aqueous solutions(Gonzalez-Davila et al. 2009). Nevertheless, Cu(II) has a low potential to activate PDS, while Cu(I) is capable of activating PDS to produce SO4• and Cu(II) following Equation (2) (Woods & Kolthoff 1965; Kolthoff & Woods 1966). Although Cu(I) could activate PDS markedly, the system has some intrinsic drawbacks, such as the instability of Cu(I), and the accumulation of Cu(II), which would cause the decline of oxidation rates (Yuan et al. 2012). In previous works, the redox cycling of copper could be strongly accelerated by hydroquinone and 2,3-dihydroxybenzoic acid (Liu et al. 2005; Yuan et al. 2013), which proved that it is a possible method to enhance copper activated PDS through accelerating the redox cycling of copper.

For the purpose of improving reaction rates, we could reasonably take into account that some reducing agents with low reaction rates to reactive species which accelerate the redox cycles of Cu(II) to Cu(I). NH2OH (HA) have been introduced into the Fenton process and the Fe(II)/PMS process to enhance the generation of reactive radicals and extend the pH range by accelerating the redox cycle of Fe(III)/Fe(II) (Chen et al. 2011; Zou et al. 2013). HA is also capable of reducing Cu(II) into Cu(I), so it was assumed that HA would also enhance the PDS/copper process (Adams & Overman 1909).

To our knowledge, there has been no literature to investigate the reduction of Cu(II) to Cu(I) in a PDS/copper system. Therefore, the aim of this study is to investigate the activation of PDS/copper by HA, focusing on the extension of the applied pH range, the role of HA, the identification of primary reactive oxidants, and proper dosage of main reagents in the process. Methyl orange was selected as the simple target compound in the Cu(II)/PDS process with the addition of HA.

MATERIALS AND METHODS

Materials

Hydroxylamine hydrochloride (HA, ≧99%), Sodium persulfate (PDS, ≧98.0%), methyl orange (MO), copper sulfate, absolute ethyl alcohol (ETOH), tert-butyl alcohol (TBA), phosphoric acid, monosodium phosphate, and sodium dihydrogen phosphate were of analytical grade and purchased from Sinopharm Chemical Reagent Co., Ltd.

Procedures

All experiments were performed in a 500 mL round-bottomed flask at 20 ± 1 °C. The reactor was operated in batch mode with rapid mixing provided at the bottom of the reactor. MO and Cu(II) with desired concentrations were spiked into the solution of the reactor buffered with 15 mmol/L phosphate buffer. Each run was initiated by adding the desired dosage of sodium persulfate and hydroxylamine. pH changed by less than 0.2 units during the process. Samples were withdrawn at predetermined intervals and then subjected to analysis with MO using a UV-vis spectrometer. Every experiment was carried out at least two times and the standard deviation obtained was less than 2.0%.

Analysis

Methyl orange was measured by a UV-vis spectrometer (MAPADA, UV-1800) at 463 nm using a 1 cm quartz cuvette. The pH changed less than 0.2 units during the process, and we produced the standard concentration–absorbency curve at different pH values for each experiment to work out the accurate MO degradation. The products of MO degradation were examined with the gas chromatography/mass spectroscopy (GC/MS) technique, operating on a QP2010Plus GC/MS analyzer. pH was measured by a pH meter (PHB-4).

RESULTS AND DISCUSSION

Effect of HA addition on MO degradation in the HA/Cu(II)/PDS process

Figure 1(a) shows the degradation of MO in PDS, Cu(II)/PDS, HA/PDS, and HA/Cu(II)/PDS processes. As shown, 29.6% and 30.5% of the MO was degraded in 60 min in the PDS process and the Cu(II)/PDS process, respectively, suggesting that Cu(II) can hardly activate PDS. Surprisingly, more than 60% of MO was degraded in 60 min in the Cu(II)/PDS process with the addition of HA. Note that only 24% of MO was degraded by the HA/PDS process. It was obvious that the addition of HA strongly enhanced oxidation in the Cu(II)/PDS process. As a typical reducing agent, HA could accelerate the transformation from Cu(II) to Cu(I). This allows Cu(I) to function as an activator of PDS to generate SO4• and •OH following Equations (3)–(6) (Kolthoff & Woods 1966; Pennington & Haim 1968; Hayon et al. 1972). The mechanism of the HA/Cu(II)/PDS process is shown in Figure 1(b).
Figure 1

(a) Time-dependent degradation of MO by PDS, Cu(II)/PDS, HA/PDS, and HA/Cu(II)/PDS processes. [Cu(II)]0 = 10 μM mol L−1, [PDS]0 = 1 m mol L−1, [HA]0 = 0.5 m mol L−1, [MO]0 = 18.3 μ mol L−1, pH = 7 ± 0.1. (b) The mechanism of HA/Cu(II)/PDS process.

Figure 1

(a) Time-dependent degradation of MO by PDS, Cu(II)/PDS, HA/PDS, and HA/Cu(II)/PDS processes. [Cu(II)]0 = 10 μM mol L−1, [PDS]0 = 1 m mol L−1, [HA]0 = 0.5 m mol L−1, [MO]0 = 18.3 μ mol L−1, pH = 7 ± 0.1. (b) The mechanism of HA/Cu(II)/PDS process.

According to the literature (James & Rod 1987; Yuan et al. 2012), different Cu(I) species may exist in the HA/Cu(II)/PDS process, such as Cu+, Cu3PO4, CuClOH, CuCO3, CuHCO3, CuCl, CuCl2, and CuCl32−. Thus, the Cu(I) mainly exist in the form of compounds, so Cu(I) can activate PDS in nearly neutral conditions. Nevertheless, it is possible that only part of the Cu(I) species can effectively activate PDS. This is most possibly due to the high activity of Cu+, CuCl, and CuClOH in inducing the generation of active radicals (James & Rod 1987; Yuan et al. 2012) and with varying pH value, the proportion of different Cu(I) species would be changed. 
formula
3
 
formula
4
 
formula
5
 
formula
6

The intermediates of MO degradation in the HA/Cu(II)/PDS process were examined by the GC/MS technique. The GC retention time (tR), molecular weight (MW), and main fragments are summarized in Table 1. The identified intermediate 5 indicates dissociation of the C–N bond close to the side of benzenesulfonic group, due to the attack of reactive radicals. The intermediates 1 and 4 are due to recombination of the radical fragments derived from opening the benzene ring by oxidation, which proved the strong oxidizability of the HA/Cu(II)/PDS process. Thus, the degradation of MO is not only because of the oxidation of the chromophoric group. Moreover, the GC chromatogram of the intermediates of MO degradation in the HA/Cu(II)/PDS process is shown in Figure S1 (see Supplementary Information, available in the online version of this paper). The absolute intensity of intermediates increased first and then decreased, which proved that the HA/Cu(II)/PDS process can further degrade the intermediates of MO degradation.

Table 1

GC/MS results

Intermediate tR (min) MW Main fragment ions (m/z
Methyl methacrylate (1) 3.61 100 100, 82, 69, 55, 41/40 
aniline (2) 5.11 93 92/91, 77, 65, 51, 31 
N,N-Dimethylbenzene-1,4-diamine (3) 8.25 136 121, 106/105, 91, 77, 65, 51, 33 
Oct-1-en-3-ol (4) 18.26 128 127, 113, 99, 85, 71, 57, 43/42 
4-Diazenyl-N,N-dimethylbenzenamine (5) 20.83 149 149, 121, 105, 93, 76, 66/65, 50 
Intermediate tR (min) MW Main fragment ions (m/z
Methyl methacrylate (1) 3.61 100 100, 82, 69, 55, 41/40 
aniline (2) 5.11 93 92/91, 77, 65, 51, 31 
N,N-Dimethylbenzene-1,4-diamine (3) 8.25 136 121, 106/105, 91, 77, 65, 51, 33 
Oct-1-en-3-ol (4) 18.26 128 127, 113, 99, 85, 71, 57, 43/42 
4-Diazenyl-N,N-dimethylbenzenamine (5) 20.83 149 149, 121, 105, 93, 76, 66/65, 50 

Effect of pH

Figure 2 shows the effect of pH on MO degradation in the HA/Cu(II)/PDS process. The effective pH value of degradation of MO ranges from 4 to 8. The MO degradation increased with the increase of pH from pH 3 to 7; a further increase of pH resulted in a decrease of MO degradation. These results show that the HA/Cu(II)/PDS process is efficient near neutral pH, which overcomes the drawbacks of Fenton or Fenton-like processes with an effective pH range under acidic conditions (Chen et al. 2011; Zou et al. 2013). When both pKa1 of HA and pH are equal to 5.96 (Robinson & Bower 1961), HA is in a neutral condition.
Figure 2

Effect of pH on MO degradation in the HA/Cu(II)/PDS process. [Cu(II)]0 = 10 μmol L−1, [PDS]0 = 1 mmol L−1, [HA]0 = 0.5 mmol L−1, [MO]0 = 18.3 μmol L−1, reaction time = 30 min, pH = 7 ± 0.1.

Figure 2

Effect of pH on MO degradation in the HA/Cu(II)/PDS process. [Cu(II)]0 = 10 μmol L−1, [PDS]0 = 1 mmol L−1, [HA]0 = 0.5 mmol L−1, [MO]0 = 18.3 μmol L−1, reaction time = 30 min, pH = 7 ± 0.1.

Beyond this point, the protonation ratios of HA decrease with the increase of pH. The fact that the degradation of MO increased with the decrease of the ratio of NH3OH+/HA as pH increased from 3 to 7 may prove that the unprotonated HA is more effective at reducing Cu(II) into Cu(I) to induce the generation of reactive radicals. However, the degradation of MO decreased with the increase of pH range from 7 to 8, which is different from the increase in degradation of MO observed at pH 3 to 7. Although low ratios of NH3OH+/HA can accelerate the transformation from Cu(II) to Cu(I) to produce more reactive radicals, high pH can increase the rate constant between HA with both SO4• and •OH via Equations (7) and (8) (Buxton et al. 1988; Neta et al. 1988). Moreover, the solubility product (Ksp) of Cu(OH)2 is 1.6 × 10−19, so a large percentage of Cu(II) translates to Cu(OH)2 which may not be effective in the HA/Cu(II)/PDS process at pH greater than 7.0 (Kang et al. 2014). In addition, the higher proportion of reactive oxidant •OH may decrease the degradation of MO due to the higher reaction rate between •OH and HA with pH increase. 
formula
7
 
formula
8

Identification of primary reactive oxidants

It has been reported that two different reactive radicals, SO4• and •OH, can be generated for catalyst-mediated decomposition of PDS systems (Liang & Su 2009). As shown in Equations (3)–(6), both radicals could be generated in the HA/Cu(II)/PDS process. However, the HA/Cu(II)/PDS process is complex, and •OH and SO4• are reactive intermediates with very short lifetimes. Thus, the existence and effect of SO4• and •OH are indirectly proved at different pH values through inhibition by absolute ethyl alcohol (ETOH) and tert-butyl alcohol (TBA). Owing to high rate constants with SO4• (1.6 × 107–7.7 × 107 mol−1Ls−1) (Neta et al. 1988) and •OH (1.2 × 109–2.8 × 109 mol−1Ls−1) (Buxton et al. 1988), EtOH is an effective quencher for both SO4• and •OH. In contrast, due to the high rate constant with •OH (6.0 × 108 mol−1Ls−1) (Buxton et al. 1988) and the much slower rate constant with SO4•(8.0 × 105mol−1Ls−1) (Neta et al. 1988), TBA is an effective quencher for •OH but not for SO4•. Based on these properties, the quenching experiment with EtOH and TBA could allow us to differentiate between the contributions of S2O82−, SO4• and •OH.

Figure 3 shows the inhibition effect of EtOH and TBA on the degradation of MO in the HA/Cu(II)/PDS process at pH 6–8 which were highly effective pH levels, as shown in Figure 2. Due to the addition of 30 mmol L−1 TBA, the inhibition of MO degradation in the HA/Cu(II)/PDS process increased with pH increase ranging from 6 to 8. It can be inferred that the proportion of •OH increased with pH increase ranging from 6 to 8. Moreover, with the pH increase, SO4• can react with OH to induce the production of •OH following Equation (4) which is significantly faster than Equation (5) (Liang & Su 2009). Due to the fact that the reaction rate constant of HA with •OH is higher than the rate constant between HA and SO4•, there are more reactive radicals reacting with HA to inhibit the degradation of MO with pH increase, which conforms to Figure 2. On the other hand, with the addition of 30 mmol L−1 EtOH, the degradation efficiency of MO was decreased from 51.74%, 61.34%, and 55.19% to 23.65%, 14.04%, and 10.09% at pH 6, 7, 8, respectively, which can mean that SO4• and •OH are two of the reactive oxidants in the HA/Cu(II)/PDS process and the addition of EtOH weakly inhibited the MO degradation which indicated the absence of radicals in the PDS process and the HA/PDS process (shown in Figure S2, see Supplementary Information, available in the online version of this paper). Hence, the primary reactive oxidants were PDS, SO4• and OH• in the HA/Cu(II)/PDS process.
Figure 3

Inhibition of radical scavengers on MO degradation in the HA/Cu(II)/PDS process at pH 6, 7 and 8. [Cu(II)]0 = 10 μmol L−1, [PDS]0 = 1 mmol L−1, [HA]0 = 0.6 mmol L−1, [MO]0 = 18.3 μmol L−1, [TBA]0 = 30 mmol L−1, [ETOH]0 = 30 mmol L−1, reaction time = 60 min.

Figure 3

Inhibition of radical scavengers on MO degradation in the HA/Cu(II)/PDS process at pH 6, 7 and 8. [Cu(II)]0 = 10 μmol L−1, [PDS]0 = 1 mmol L−1, [HA]0 = 0.6 mmol L−1, [MO]0 = 18.3 μmol L−1, [TBA]0 = 30 mmol L−1, [ETOH]0 = 30 mmol L−1, reaction time = 60 min.

Kinetics of MO degradation

In principle, the contaminant disappearance should be describable with a pseudo-first-order rate constant that is the sum of second-order terms for each oxidant as in Equation (9) where k″ represents the second-order rate constants for the reaction of the contaminant with each reactive intermediate (Waldemer et al. 2007). Under most conditions, the dominant term in Equation (9) is presumed to be the one involving SO4• (Kolthoff et al. 1951; Peyton 1993) which is in agreement with the HA/Cu(II)/PDS process in Figure 3, but it is not known how the relative significance of these terms varies with system parameters. 
formula
9

To investigate the kinetics of MO degradation, different dosages of HA in the HA/Cu(II)/PDS process were studied. The degradation of MO in the HA/Cu(II)/PDS process consists of three stages, which are the fast stage (k1), the transitory stage (k2) and the low stage (k3). Furthermore, the fast stage and the low stage agree with first-order kinetic models, and the transitory stage fits with second-order kinetic model. Rate constants (k, k1, k2, k3) and R2 (R2 > 0.964) of the three stages are shown in Table S1 and Table S2 (see Supplementary Information, available in the online version of this paper). Increased k1 and k2 were observed with the increase of HA concentration in the range of 0.1 mmol L−1 to 0.75 mmol L−1. The reaction rate of the fast stage depends on the initial HA concentration. After that, the increase of HA concentration resulted in decrease of k1 and k2. k3 was significantly less than k1 with no apparent relationship between the two values. From the comparison between Tables S1 and S2 (see Supplementary Information) the predominant increase of MO degradation in the HA/Cu(II)/PDS process occurs in the fast stage. k1 is significantly higher than k of the PDS process, the HA/PDS process and the Cu(II)/PDS process. On the contrary, k3 is lower than k in the same processes resulting in the activation of copper/PDS, which was facilitated in the fast stage by HA via acceleration of the Cu(II)/Cu(I) redox couple.

Effect of PDS, HA, and Cu(II) on MO degradation

To investigate the application of the HA/Cu(II)/PDS system, different dosages of PDS, HA, and Cu(II) were tested in the HA/Cu(II)/PDS process, as shown in Figure 4. With the increase of initial PDS concentration, MO degradation was strongly enhanced and optimal Cu(II) concentration is about 10μmol L−1 but the optimal HA concentration was increased with increase of PDS concentration. Thus, although proper HA concentration could increase the MO degradation in the Cu(II)/PDS process, the main contribution of MO degradation is PDS. Increased degradation of MO was observed with the increase of initial HA concentration. Exceeding the optimal concentration resulted in the decrease of MO degradation. It can be inferred that the more HA that was used to accelerate the transformation from Cu(II) to Cu(I) to activated PDS, the more reactive radicals were generated. However, a large amount of reactive radicals could be quenched by HA with high rate constants, especially when the HA concentration was high enough via Equations (7) and (8) (Buxton et al. 1988; Neta et al. 1988), resulting in decrease in MO removal. Moreover, increase of initial Cu(II) concentration could result in the increase of MO degradation but exceeding the optimal concentration resulted in the decrease of MO degradation. This is because Cu(I) may react with SO4• and •OH with a high rate constant following Equations (10) (Buxton et al. 1988) and (11) (Liang & Su 2009) to inhibit MO degradation in the HA/Cu(II)/PDS process. Therefore, a proper dosage of PDS, HA, and Cu(II) should be selected in the application in order to improve the degradation of probe compounds at the highest extent and reduced cost.
Figure 4

Effect of PDS, HA, and Cu(II) concentration on MO degradation in the HA/Cu(II)/PDS process. [MO]0 = 18.3 μmol L−1, reaction time = 90 min, pH = 7 ± 0.1. (a) [PDS]0 = 1 mmol L−1, [Cu(II)]0 = 1–20 μmol L−1, [HA]0 = 0.1–1.5 mmol L−1. (b) [PDS]0 = 0.1–1.5 mmol L−1, [Cu(II)]0 = 10 μmol L−1, [HA]0 = 0.1–1.5 mmol L−1. (c) [PDS]0 = 0.1–1.5 mmol L−1, [Cu(II)]0 = 1–20 μmol L−1, [HA]0 = 0.75 mmol L−1.

Figure 4

Effect of PDS, HA, and Cu(II) concentration on MO degradation in the HA/Cu(II)/PDS process. [MO]0 = 18.3 μmol L−1, reaction time = 90 min, pH = 7 ± 0.1. (a) [PDS]0 = 1 mmol L−1, [Cu(II)]0 = 1–20 μmol L−1, [HA]0 = 0.1–1.5 mmol L−1. (b) [PDS]0 = 0.1–1.5 mmol L−1, [Cu(II)]0 = 10 μmol L−1, [HA]0 = 0.1–1.5 mmol L−1. (c) [PDS]0 = 0.1–1.5 mmol L−1, [Cu(II)]0 = 1–20 μmol L−1, [HA]0 = 0.75 mmol L−1.

Furthermore, it should be noted that HA is a kind of toxic compound and the addition of HA in Cu(II)/PDS is far from practical. Our work just introduced an interesting phenomenon and proposed a preliminary interpretation that HA could accelerate the reaction rates and expand effective pH range. The ultimate aim is to find an environmentally friendly chemical with a feasible reaction rate with •OH and SO4• to reduce the cost and accelerate the generation of active radicals. 
formula
10
 
formula
11

CONCLUSION

This study sets up the HA/Cu(II)/PDS process to utilize HA to reduce Cu(II) to Cu(I) to induce the production of SO4• and •OH through the activation of PDS. The MO degradation in the HA/Cu(II)/PDS process can be divided into three stages: the fast stage, the transitory stage, and the low stage. The fast and low stages fit well with pseudo-first-order kinetic models, and the transitory stage agrees with a second-order kinetic model. HA can strongly enhance MO degradation in the HA/Cu(II)/PDS process, which countered the drawbacks of accumulation and sedimentation of Cu(II) and instability of Cu(I) and the HA/Cu(II)/PDS process excellently showed high degradation of MO in nearly neutral conditions, which enhanced the application of the HA/Cu(II)/PDS process. However, owing to the reactions between HA and reactive oxidants, the addition of HA may reduce the use of PDS in the HA/Cu(II)/PDS process. High dosage of PDS can enhance the MO degradation in the HA/Cu(II)/PDS process, and the optimal dosage of Cu(II) is about 10 μmol L−1. Although this study provided promising results, further research is required to better understand the problem. Future studies may not only be confined to HA; other reducing agents which can react with Cu(II) to generate Cu(I) may act similar to HA in the HA/Cu(II)/PDS process.

ACKNOWLEDGEMENTS

Appreciation and acknowledgment are given to the National Natural Science Foundation of China (No. 51508353), the Program for New Century Excellent Talents in University (NCET-11-0082), and the National Natural Science Foundation of China (No. 51008052).

REFERENCES

REFERENCES
Adams
M.
Overman
E.
1909
The reduction of copper sulphate with hydroxylamine
.
Journal of the American Chemical Society
31
,
637
640
.
Al-Shamsi
M. A.
Thomson
N. R.
2013
Treatment of organic compounds by activated persulfate using nanoscale zerovalent iron
.
Industrial & Engineering Chemistry Research
52
(
38
),
13564
13571
.
Anipsitakis
G. P.
Dionysiou
D. D.
2004
Radical generation by the interaction of transition metals with common oxidants
.
Environmental Science & Technology
38
,
3705
3712
.
Balaz
P.
Takacs
L.
Jiang
J. Z.
Soika
V.
Luxova
M.
2002
Mechanochemical reduction of copper sulfide
.
Journal of Metastable and Nanocrystalline Materials
13
,
257
262
.
Chen
L.
Ma
J.
Li
X.
Zhang
J.
Fang
J.
Guan
Y.
Xie
P.
2011
Strong enhancement on Fenton oxidation by addition of hydroxylamine to accelerate the ferric and ferrous iron cycles
.
Environmental Science & Technology
45
(
9
),
3925
3930
.
Gonzalez-Davila
M.
Santana-Casiano
J. M.
Gonzalez
A. G.
Perez
N.
Millero
F. J.
2009
Oxidation of copper(I) in seawater at nanomolar levels
.
Marine Chemistry
115
(
1–2
),
118
124
.
James
W. M.
Rod
G. Z.
1987
Reaction kinetics of hydrogen peroxide with copper and iron in seawater
.
Environmental Science & Technology
21
(
8
),
804
810
.
Kang
D. E.
Lim
C. S.
Kim
J. Y.
Kim
E. S.
Chun
H. J.
Cho
B. R.
2014
Two-Photon probe for Cu2+ with an internal reference: quantitative estimation of Cu2+ in human tissues by Two-Photon microscopy
.
Analytical Chemistry
86
(
11
),
5353
5359
.
Kolthoff
I. M.
Woods
R.
1966
Polarographic kinetic currents in mixtures of persulfate and copper (II) in chloride medium
.
Journal of American Chemistry Society
88
(
7
),
1371
1375
.
Kolthoff
I. M.
Medalia
A. I.
Raaen
H. P.
1951
The reaction between ferrous iron and peroxides. IV. Reaction with potassium persulfate
.
Journal of the American Chemical Society
73
(
4
),
1733
1739
.
Liang
C.
Su
H.-W.
2009
Identification of sulfate and hydroxyl radicals in thermally activated persulfate
.
Industrial & Engineering Chemistry Research
48
,
5558
5562
.
Neta
P.
Huie
R. E.
Ross
A. B.
1988
Rate constants for reactions of inorganic radicals in aqueous solution
.
Journal of Physical and Chemical Reference Data
17
(
3
),
1027
.
Pennington
D. E.
Haim
A.
1968
Stoichiometry and mechanism of the chromium-peroxydisulfate reaction
.
Journal of the American Chemical Society
90
,
3700
3704
.
Robinson
R.
Bower
V.
1961
The ionization constant of hydroxylamine
.
Journal of Physical Chemistry
65
,
1279
1280
.
Waldemer
R. H.
Tratnyek
P. G.
Johnson
R. L.
Nurmi
J. T.
2007
Oxidation of chlorinated ethenes by heat-activated persulfate: kinetics and products
.
Environmental Science & Technology
41
,
1010
1015
.
Yuan
X.
Pham
A. N.
Xing
G.
Rose
A. L.
Waite
T. D.
2012
Effects of pH, chloride, and bicarbonate on Cu(I) oxidation kinetics at circumneutral pH
.
Environmental Science & Technology
46
(
3
),
1527
1535
.

Supplementary data