1. Kinetic and Spectroscopic Studies of Co(II)- and Mn(II)-Substituted Catechol
Dioxygenases
Andrew J. Fielding
In my doctoral research, I studied the mechanism of O2 activation and catalysis by a class of
enzymes called catechol dioxygenases. These enzymes play an important role in the carbon-cycle,
in the biodegradation of aromatic compounds, including lignin from woody plants, and in the
bioremediation of the environment, by microorganisms, after an event such as an oil spill. I also
examined the degradation of halogenated aromatic compounds, by these dioxygenases.
More specifically, my doctoral research focused on understanding the mechanism of O2
activation and catalysis by two similar catechol dioxygenases: Fe(II)-homoprotocatechuate 2,3-
dioxygenase (Fe-HPCD) from B. fuscum and Mn(II)-MndD from A. globiformis. Both Fe-HPCD and
Mn-MndD cleave the same catechol substrate homoprotocatechuate at comparable rates.
Structurally, Fe-HPCD and Mn-MndD are nearly identical, yet employ metal ions with M(III/II)
standard redox potentials that differ by 0.79.1,2 One might expect that this large difference in
redox potentials would result in the
Mn(II) enzyme being significantly
slower, as Mn(II) is a poorer reducing
agent then Fe(II), for the initial
reductive O2 activation step (Scheme
1). However, metal-substitution
experiments showed that when Fe(II)
was swapped with Mn(II), and vice
versa, the resulting Mn-HPCD and Fe-
MndD exhibited no significant change
in activity. This suggests that the
mechanism of O2 activation by
extradiol catechol dioxygenases is
somewhat different then that of other
oxygenases.3
In my research, I prepared and characterized the cobalt-substituted variant of HPCD. Cobalt’s
(III/II) standard redox potential is 1.15 V higher than Fe(II) and thus would be expected to be a
very poor reducing agent for reductive O2 activation. While Co-HPCD shows a very low affinity for
O2, under O2 saturating conditions, it is three times more active then either Mn- or Fe-HPCD.4 To
rationalize these observations, a "self compensating rates" mechanism was proposed, where the
slower electron transfer from Co(II) to O2 to form the initial Co(III)-superoxide complex (C), is
compensated by the subsequent faster electron transfer from the bound catechol to Co(III) to form
a [Co(II)(semiquinone)-superoxide] (D), since Co(III) is expected to be a stronger oxidant then
Fe(III).4, 5
Using the Co-substituted enzyme I was also to trap and
characterize three O2 intermediates by employing electron poor
substrate analogs 4-nitrocatechol (4NC) and halogenated
catechols (4XC, X = Br, Cl, or F), to slow down the rate of catechol
oxidation. With 4NC a low-spin Co(III)-superoxide species,
intermediate C, was observed by EPR (Figure 1). This represents
the first example of an O2-adduct for a Co-substituted
oxygenase.6-8

2. Employing the 4XCs two intermediate species were
observed, the first species appears to be a
semiquinone radical species based on rapid freeze-
quench EPR experiments and was assigned to a low-
spin [Co(III)(semiquinone)peroxo]. The second
species, observed in stopped-flow experiments,
exhibits an intense chromophore (λmax = 460 nm,
Figure 2) upon mixing the anaerobic enzyme-substrate
complex with O2-saturated buffer, which then decayed
to the ring cleaved product (λmax = 380 nm). This
intermediate has been further characterized using
resonance Raman, EPR and Mössbauer, which suggest
that it is likely a [M(II)(quinone)peroxo] species.9
This intermediate is observable with all three M-HPCDs, it offered a unique opportunity to
compare the effect of the different properties of the metal on the rates of individual reaction steps.
The observed activity of Co-HPCD has augmented our understanding of O2 activation by
dioxygenases. This fundamental research has shown previously unobserved chemistry for cobalt,
which can provide a basis for practical application in catalyst designing. The metal-substitution
experiments have also allowed us to take advantage of the different spectroscopic properties of
each metal to fully characterize the electronic structure of the observed intermediates.10
References:
1 Miller, M. A.; Lipscomb, J. D. J. Biol. Chem. 1996, 271, 5524–5535.
2. Whiting, A. K.; Boldt, Y. R.; Hendrich, M. P.; Wackett, L. P.; Que, L. Jr. Biochemistry 1996, 35,
160–170.
3. Emerson, J. P.; Kovaleva, E. G.; Farquhar, E. R.; Lipscomb, J. D.; Que, L. Jr. Proc. Natl. Acad. Sci.
U.S.A. 2008, 105, 7347–7352.
4. Fielding, A. J.; Kovaleva, E. G.; Farquhar, E. R.; Lipscomb, J. D.; Que, L. Jr. J. Biol. Inorg. Chem.
2011, 16, 341–355.
5. Gunderson, W. A.; Zatsman, A. I.; Emerson, J. P.; Farquhar, E. R.; Que, L. Jr.; Lipscomb, J. D.;
Hendrich, M. P. J. Am. Chem. Soc. 2008, 130, 14465–14467.
6. Groce, S. L.; Miller-Rodeberg, M. A.; Lipscomb, J. D. Biochemistry 2004, 43, 15141–15153.
7. Smith, T. D.; Pilbrow, J. R. Coord. Chem. Rev. 1981, 39, 295–383.
8. Fielding, A. J.; Lipscomb, J. D.; Que, L., Jr. J. Am. Chem. Soc. 2012, 134, 796–799
9. Mbughuni, M. M.; K. Meier, K. K.; Eckard Münck, E.; Lipscomb J. D. Biochemistry 2012, 51,
8743–8754.
10. Fielding, A. J.; Lipscomb, J. D.; Que, L. J. Biol. Inorg. Chem. 2014, 19, 491–504.