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Hu Y, et al. Sci China Chem February (2012) Vol.55 No.2
Table 3 The mechanism and rate constant of oxidation of formamidine disulfide by hydrogen peroxide at pH 2.00
No.
R1
Reaction
Rate constants
3.30 × 104 s1
H2N(=NH)CS-SC(=NH)NH2+ H2O ⇌ (H2N)2C=S + H2N(=NH)CSOH
4.80 s1
1.94 × 104 s1
R2
R5
R6
R7
R8
R9
R10
H2N(=NH)CSOH+ H2O → (NH2)2CO + HSOH
H2O2 + (H2N)2C=S → H2N(=NH)CSOH + H2O
H2N(=NH)CS-SC(=NH)NH2+H2O2 → 2H2N(=NH)CSOH
H2N(=NH)CSOH + H2O2 → H2N(=NH)CSO2H + H2O
H2O2 + H2N(=NH)CSO2H → H2N(=NH)CSO3H + H2O
H2N(=NH)CSO3H + H2O → HSO3 + (H2N)2CO + H+
HSO3 + H2O2 → SO42 + H2O + H+
1.50 × 102 M1 s1
2.14 × 102 M1 s1
4.44 × 102 M1 s1
4.27 × 104 M1 s1
7.42 × 106 s1
1.48 × 105 M1 s1
which means that oxidation of FDS is the main reaction in
the presence of excess H2O2 at pH 1.502.50. However,
when the pH is greater than 3.00, the hydrolysis reaction
rate increases greatly with increasing pH and becomes
comparable with that of oxidation reaction.
4 Conclusion
In addition to TU, FSIA and FSOA, the formation of FSEA
and thiocynogen was newly detected by HPLC-MS during
hydrolysis and oxidation of FDS, and the HPLC technique
also provides real time tracking of the above sulfur species
for kinetic analysis. With a 10-step reaction mechanism
including hydrolysis equilibrium of formamidine disulfide
and irreversible hydrolysis of formamidine sulfenic acid,
experimental curves of different species can be effectively
simulated for FDS oxidation. The new results help to un-
cover the complicated mechanism of spatiotemporal
self-organization in the oxidation of TU. It is also helpful to
control the FDS stability in aqueous solution for practical
applications such as precious metal leaching, chemical syn-
thesis, and corrosion protection.
3.3 Mechanistic analysis and simulation
Several intermediates and the final products during the hy-
drolysis and oxidation of FDS were identified using HPLC
and MS. The oxidation of FDS by hydrogen peroxide was
accompanied by hydrolysis reaction, which produced TU,
FSIA, FSOA, two previously unreported intermediates,
FSEA and (SCN)2, and finally (H2N)2CO, SO42 and sulfur.
Since the FSEA peak during oxidation was higher than that
during hydrolysis as shown in Figure 4 in contrary to Figure
1, FSEA was produced not only from hydrolysis but also
from oxidation of FDS and TU. In our separate experiments,
we found that the consumption rates of FSOA in decompo-
sition and oxidation by hydrogen peroxide are of the same
when pH < 3, indicating that direct oxidation of FSOA can
be ignored. Table 3 gives the mechanism scheme for oxida-
tion of FDS, where R1 and R2 are hydrolysis equilibrium of
FDS and irreversible hydrolysis of FSEA, respectively. The
scheme does not include R3 and R4 because concentrations
of thiocynogen and FSEA were very low in FDS hydrolysis.
R5, R6, R7 and R8 are oxidation reactions that produce
FSEA, FSIA and FSOA, respectively. R9 is hydrolysis re-
action of FSOA followed by sulfite oxidation reaction R10.
The rate constants k1, k-1 and k2 have been fitted. Especially,
our fitted k1 value agrees with the values reported by Rábai
et al. [6] and Gao et al. [13] respectively. k5 and k7 were
reasonably adjusted according to pH difference between this
work (at pH 2.0) and the reference (at pH 1.5). k8 and k9
were determined with HPLC [25]. k10 were obtained from
Rabai et al. [26]. By pre-equilibrium approximation, the
rate constants of hydrolysis reactions (R1, R-1 and R2) were
calculated which are consistent with Table 1.
This work was partly supported by the National Natural Science Founda-
tion of China (21073232 & 50921002) and the Fundamental Research
Fund from the Chinese Central University (2010LKHX02).
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kh ≥ k1 k2/(k1 [FDS]0) = 2.66 × 105 s1 (pH 2)
Figure 5(a) also shows that the corresponding curves simu-
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