Ligand
-
enhanced abiotic iron oxidation and the effects of chemical vs.
biological iron cycling in anoxic environments
Sebastian H. Kopf
1
, Cynthia Henny
2
, Dianne K. Newman
1,3,4
*
AVAILABLE SUPPORTING INFORMATION
•
Derivation of reaction equations
•
Table S1: pH of reactant solutions at the beginning and end of kinetic Fe(II)
oxidation experiments.
•
Table S2: Theoretical Fe(II) inorganic and organic speciation in bicarbonate
-
buffered f
reshwater medium at pH 7.
•
Table S3: Overview of rate constants reported for chemical oxidation of Fe(II)
by NO
2
-
.
•
Figure S1. Anaerobic growth of
Pseudogulbenkiania
sp. strain MAI
-
1.
•
Figure S2. MAI
-
1 growth on various Fe(II)
-
chelating ligands.
•
Figure S3. Ox
idation of Fe(II)
-
NTA in spent MAI
-
1 growth medium.
•
Figure S4. Oxidation test of Fe(II) in the presence of nitrite during sample
dilution for the ferrozine assay.
•
Figure S5. Reduction test of nitrite in the presence of Fe(II) (+ ligands) during
diazotizati
on for the nitrite assay.
•
Figure S6. Anaerobic growth and nitrite production of
Paracoccus
denitrificans
.
•
Figure S7: Model fits for abiotic Fe(II) oxidation by nitrite.
•
Figure S8. Evolution of N
2
O during reaction of Fe(II)
-
citrate with nitrite.
•
Figure S9. Absorption spectrum of Fe(II)
-
NTA solution reacted with nitrite.
Derivation of Fe(II) and NO
2
-
reaction equations:
d[Fe(II)]/dt =
-‐
2
k
app
[Fe(II)][NO
2
-‐
]
[Fe(II)] = [Fe(II)]
0
+ Δ[Fe(II)] ; [NO
2
-‐
] = [Fe(II)]
0
+ ½ Δ[Fe(II)]
d([Fe(II)]
0
+ Δ[Fe(II)]) / dt =
-‐
2
k
app
([Fe(II)]
0
+ Δ[Fe(II)]) ([Fe(II)]
0
+ ½ Δ[Fe(II)])
d
Δ[Fe(II)] / dt =
-‐
2
k
app
([Fe(II)]
0
+ Δ[Fe(II)]) ([Fe(II)]
0
+ ½ Δ[Fe(II)])
[eq. S1]
d[NO
2
-‐
]/dt =
-‐
k
app
[Fe(II)][NO
2
-‐
]
[NO
2
-‐
] = [NO
2
-‐
]
0
+ 2 Δ[NO
2
-‐
] ; [NO
2
-‐
] = [NO
2
-‐
]
0
+ Δ[NO
2
-‐
]
d([NO
2
-‐
]
0
+ Δ[NO
2
-‐
]) / dt =
-‐
k
app
([NO
2
-‐
]
0
+ 2 Δ[NO
2
-‐
]) ([NO
2
-‐
]
0
+ Δ[NO
2
-‐
])
d
Δ[NO
2
-‐
]
/ dt =
-‐
k
app
([NO
2
-‐
]
0
+ 2 Δ[NO
2
-‐
]) ([NO
2
-‐
]
0
+ Δ[NO
2
-‐
])
[eq. S2]
For 2:1 stoichiometry, Δ[Fe(II)] =
–
[Fe(II)]
ox
=
–
([Fe(II)]
0
–
[Fe(II)]
obs
) and Δ[NO
2
-‐
] =
–
[NO
2
-‐
]
red
=
–
([NO
2
-‐
]
0
–
[NO
2
-‐
]
obs
), and [S1] and [S2] integrate to yield:
!"
!!
!"#
!
=
!"
!!
!
!
!
!
!
!
!"
(
!!
)
!
!
!""
!
!"
!
!
!
=
!"
!
!
!
!
!"
!
!
!
!
!""
!
!
!
!
!
!
!"
!
!
!
!
!""
!
For 1:1 stoichiometr
y observed in the presence of NTA: The Fe
-‐
NTA
-‐
NO complex does not
appear to be reactive towards NO
2
-‐
such that [S1] describes Fe(II) oxidation even in the
presence of NTA, with the caveat that measured concentrations of Fe(II) (which include the
Fe(II)
-‐
NTA
-‐
NO
-‐
complex) require a correction for Fe
-‐
NTA
-‐
NO. Assuming all NO that is
generated complexes with Fe(II)
-‐
NTA such that it no longer participates in a redox reaction
with nitrite, but is still measured as Fe(II) by the
f
errozine assay and assuming the reac
tions
are coupled such that [Fe(II)]
ox
= [Fe(II)
-‐
NTA
-‐
NO], then Δ[Fe(II)] =
–
([Fe(II)]
ox
+ [Fe(II)
-‐
NTA
-‐
NO]) =
–
2 [Fe(II)]
ox
=
–
2 ([Fe(II)]
0
-‐
[Fe(II)]
obs
)
and Δ[NO
2
-‐
] =
–
[NO
2
-‐
]
red
=
–
([NO
2
-‐
]
0
–
[NO
2
-‐
]
obs
). This leads [S1] to integrate to:
!"
!!
!"#
!
=
!"
!!
!
!
!"
(
!!
)
!
!
!""
!
−
1
+
2
!
!"
(
!!
)
!
!
!""
!
SUPPORTING TABLES
Table S1
pH of reactant solutions at the beginning and end of kinetic Fe(II) oxidation experiments.
Condition
Start
End
Change
2mM
Fe(II) +
2mM
NO
2
-‐
7.03
6.88
-‐
0.15
+ 2mM NTA
7.00
7.12
0.12
+ 300mg/L PPHA
6.99
7.03
0.04
+ 100μM Citrate
6.95
7.02
0.07
+ 500μM Citrate
6.97
7.07
0.10
+ 2mM Citrate
6.96
7.06
0.10
+ 2mM Citrate + 300mg/L PPHA
6.94
7.13
0.19
Table S2
Theoretical Fe(
II) inorganic and organic speciation in bicarbonate
-‐
buffered freshwater medium at pH 7.
Species with relative abundance <
0.01% for all experimental conditions
are not shown
.
Species suggested to be relevant for Fe(II) oxidation by nitrite are
highlighted
in
gray.
2mM Fe(II)
5mM
Fe(II)
Ligand
none
PPHA
(300mg/L)
Citrate
(0.1mM)
Citrate
(0.5mM)
Citrate
(2mM)
Citrate + PPHA
(2mM+300mg/L)
NTA
(2mM)
Citrate
(10mM)
[Fe(II)
species
] / [Fe(II)
total
]
Fe
2+
26.66%
23.80%
25.75%
22.26%
11.49%
9.64%
1.89%
0.91%
Fe
-‐
OH
+
0.06%
0.05%
0.06%
0.05%
0.03%
0.02%
< 0.01%
< 0.01%
Fe
-‐
HCO
3
+
4.37%
3.91%
4.22%
3.64%
1.86%
1.57%
0.31%
0.13%
Fe
-‐
CO
3 (aq)
65.68%
58.82%
63.39%
54.60%
27.86%
23.44%
4.57%
1.93%
Fe
-‐
CO
3
-‐
OH
-‐
0.15%
0.14%
0.15%
0.13%
0.07%
0.06%
0.01%
< 0.01%
Fe
-‐
(CO
3
)
2
2
-‐
0.09%
0.08%
0.08%
0.07%
0.04%
0.03%
0.01%
< 0.01%
Fe
-‐
Cl
+
0.09%
0.08%
0.08%
0.07%
0.04%
0.03%
0.01%
< 0.01%
Fe
-‐
NH
3
2+
0.02%
0.02%
0.02%
0.01%
0.01%
0.01%
< 0.01%
< 0.01%
Fe
-‐
HPO
4
(aq)
0.32%
0.30%
0.32%
0.28%
0.16%
0.14%
0.03%
0.01%
Fe
-‐
H
2
PO
4
+
0.08%
0.07%
0.08%
0.07%
0.04%
0.03%
0.01%
< 0.01%
Fe
-‐
SO
4
(aq)
2.48%
2.24%
2.39%
2.06%
1.06%
0.90%
0.17%
0.07%
#
Fe
-‐
L
-‐
3.46%
16.73%
57.26%
55.67%
93.00%
96.79%
#
Fe
-‐
HL
0.01%
0.03%
0.09%
0.09%
< 0.01%
0.15%
Fe
-‐
HA
(complexed)
8.26%
7.04%
Fe::HA
(weakly bound)
2.23%
1.33%
#: Fe
-‐
L
-‐
= Fe
-‐
NTA
-‐
or Fe
-‐
C
itrate
-‐
, Fe
-‐
HL = Fe
-‐
HNTA or Fe
-‐
HCitrate
Table S3
Overview of rate constants reported for chemical oxidation of
Fe(II) by NO
2
-‐
.
Experimental conditions
Kinetic parameters
Source
pH
Temp
buffer
Order
Rate constant (k)
d[Fe(II)]/dt =
Reference
Oxidation by nitrite
Fe(II) as siderite (10g/L ~ 80mM)
6
25C
MES/PIPES/HEPES
2nd
1.00E
-‐
04
M
-‐
1
s
-‐
1
-‐
2 k [FeCO
3(s)
] [NO
2
-‐
]
Rakshit et al. (2008)
(
1
)
,
Fig. 5
Fe(II) as siderite (10g/L ~ 80mM)
6.5
25C
MES/PIPES/HEPES
2nd
6.39E
-‐
05
M
-‐
1
s
-‐
1
-‐
2 k [FeCO
3(s)
] [NO
2
-‐
]
Rakshit et al. (2008)
(
1
)
, Fig. 5
Fe(II) as siderite (10g/L ~ 80mM)
7.9
25C
MES/PIPES/HEPES
2nd
5.28E
-‐
05
M
-‐
1
s
-‐
1
-‐
2 k [FeCO
3(s)
] [NO
2
-‐
]
Rakshit et al. (2008)
(
1
)
, Fig. 5
Fe(II) as goethite
6.8
30C
carbonate
1st
3.18E
-‐
06
s
-‐
1
-‐
k [Fe(II)]
Weber et
al. (2001)
(
2
)
, T
able
3
Fe(II) as biogenic magnetite
6.8
30C
carbonate
1st
3.38E
-‐
05
s
-‐
1
-‐
k [Fe(II)]
Weber et al. (2001)
(
2
)
, T
able
3
Fe(II) as HFO
6.8
26
-‐
28
PIPES
3rd
3.83E+03
M
-‐
2
s
-‐
1
-‐
k [Fe(II)
diss
] [Fe(II)
bound
] [NO
2
-‐
]
Tai et al. (2009)
(
3
)
+2mM NTA
7
25C
carbonate
2nd
6.67E
-‐
03
M
-‐
1
s
-‐
1
-‐
2 k [Fe(II)] [NO
2
-‐
]
This study, Table 1
+2mM CIT
7
25C
carbonate
2nd
4.67E
-‐
03
M
-‐
1
s
-‐
1
-‐
2 k [Fe(II)] [NO
2
-‐
]
This study, Table 1
+10mM CIT,
P. denitrificans
spent medium
7
25C
carbonate
2nd
9.42E
-‐
03
M
-‐
1
s
-‐
1
-‐
2 k [Fe(II)] [NO
2
-‐
]
This
study, Table 2
+10mM CIT,
P. denitrificans
culture
7
25C
carbonate
2nd
1.06E
-‐
02
M
-‐
1
s
-‐
1
-‐
2 k [Fe(II)] [NO
2
-‐
]
This study, Table 2
Oxidation by nitrate
Fe(II) as green rust
8.25
25C
auto
-‐
titration
2nd
4.93E
-‐
05
M
-‐
1
s
-‐
1
-‐
8 k [Fe(II)
GR
] [NO
3
-‐
]
Hansen et al. (1996)
(
4
)
, Table 1
SUPPORTING FIGURES
Figure S1
Anaerobic growth
and concomitant Fe(II) oxidation
of
Pseudogulbenkiania
sp. strain MAI
-‐
1
in
freshwater medium amended with
10mM
nitrate
and different concentrations of Fe(II), NTA and acetate
, and a headspace containing
~3
%
hydrogen
.
In the presence of NTA
, up to
10mM Fe(I
I) is
oxidized within 24hours (in yellow), however, in the absence of NTA, neither growth nor Fe(II) oxidation
is
observed (in
green). Replicate
culture (duplicates or triplicates)
indicated with
solid, dashed and dotted line
s,
respectively.
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
OD700
Fe(II) [mM]
0.0
0.1
0.2
0.3
0.0
2.5
5.0
7.5
10.0
0
10
20
30
40
50
0
10
20
30
40
50
Time [hours]
Condition:
●
●
●
●
●
Abiotic ctrl: 10mM Fe(II) / 10mM NTA
H2 + 10mM Fe(II) / 10mM NTA
H2 + 4mM Fe(II)
H2 + 5mM Act + 5mM Fe(II) / 10mM NTA
H2 + 5mM Fe(II) / 10mM NTA
Figure S2
Growth of MAI
-‐
1 on various Fe(II) chelating ligands. The organism is grown
aerobically in freshwater medium in a 96 well plate (OD
600
is
measured every 5 minutes) with different ligands as the sole carbon source.
Citrate (Cit), humic acids (HA)
,
acetate (A
ct)
and diethylene
triamine pentaacetic acid (DTPA) can all serve as growth substrates for MAI
-‐
1. The strain’s ability to use siderophore desferioxamine
(DFO) as a carbon source is ambiguous. No growth could be observed in the presence of nitrilotriacetate
(NTA) as sole carbo source. This
makes NTA a suitable choice for
anaerobic growth experiments with MAI
-‐
1 as a chelator for Fe(II)
that does not
supply
extra
carbon.
Replicate cultures
indicated with
dashed and solid lines
, respectively
.
0.0
0.2
0.4
0.6
0
5
10
15
Time [hours]
O
D
600
Substrate:
150mg/L HA
1mM DTPA
5mM Act
5mM Cit
5mM DFO
5mM NTA
Figure S3
Oxid
ation of Fe(II)
-‐
NTA in spent MAI
-‐
1 growth medium. Triplicate cultures of
Pseudogulbenkiania
sp. strain MAI
-‐
1 (solid, dashed and
dotted line) were grown in freshwater medium amended with
10mM nitrate and
1.25mM acetate
,
with ~3% H
2
present in the headspace.
During growth of MAI
-‐
1 (upper left panel),
significant amounts of nitrite
accumulated in the medium (lower left panel)
.
Accumulated
nitrite was stable at the end of growth but upon addition of ~3mM Fe(II)
-‐
NTA to filer sterilized
spent medium, Fe(II) oxidation and
concomitant nitrite reduction could be observed (right panels).
●
●
●
●
●
●
●
●
●
●
●
●
●
●
OD700
NO2
−
[mM]
0.000
0.025
0.050
0.075
0.100
0.125
0
1
2
3
4
0
10
20
30
Time [hours]
●
●
●
●
●
●
NO2
−
[mM]
Fe(II) [mM]
2
2.5
3
3.5
2
2.5
3
3.5
0
5
10
15
Time [hours]
Figure S4
Oxidation test of Fe(II) in the presence of nitrite during sample dilution for the
f
erro
z
ine
(
5
)
assay.
The
f
erro
z
ine assay
often
includes an
acid dilution step prior to spectrophotometric determination of Fe(II) with the
f
erro
z
ine reagent
.
Acidificat
ion aids in the desorption of
strongly
coordinated
Fe(II) from mineral surfaces and other
strong
sorption sites and is an important preparative
step for environmental
samples.
However, at acidic pH, nitrite is protonated (pKa=3.4) to nitrous acid, which ca
n self
-‐
decompose to form reactive N
-‐
oxides
(
6
)
as
well
as
oxidize Fe(II) directly
(
7
,
8
)
. To
assess the effect of acidification in the presence of nitrite for our experimental setup, an
anoxic
freshwater solution containing ~650μM Fe(II) and ~1mM NO
2
-‐
was diluted
1:10
with 1M HCl, and Fe(II) concentrations were measured
after varying time interv
als using the
f
erro
z
ine assay (depicted in grey). Within 10 seconds of acidification, >20% of Fe(II) was oxidized
and could no longer be detected by the
f
erro
z
ine assay. After 1 minute, >60% of Fe(II) was lost. Without the acidification step (e.g. by
direc
t dilution of the sample with the
f
erro
z
ine reagent), Fe(II) concentrations did not significantly decrease within several minutes (black
line). Since our experimental conditions included relatively high concentrations of nitrite, but little to no risk of s
orptive loss of Fe(II), all
f
erro
z
ine measurements were conducted without acidification.
●
●
●
●
●
●
●
●
0
200
400
600
0.0
2.5
5.0
7.5
10.0
Time [min]
Fe (II) [
μ
M]
Diluted in:
●
●
ferrozine
HCL
Figure S5
Reduction test of nitrite in the presence of Fe(II) during incubation with sulfanilamide in phosphoric acid for the
nitrite
assay
used in this
study
.
To as
sess the effect of
free and chelated Fe(II) on the assay,
an
anoxic
freshwater solution containing ~
1.7mM
nitrite
was amended
with 2mM Fe(II) and no ligand / 2mM citrate / 2mM EDTA / 2mM NTA / 300mg/L PPHA, and immediately
diluted
1:10
with
1%
sulfanilamid
e in 5% phosphoric acid
for diazodization
. Nitrite concentrations were determined colorimetrically after varying time
intervals by addition of 0.1% N
-‐
1
-‐
napthylethylenediamine.
The true concentration of
nitrite
measured in the absenc
e of Fe(
II) is indicated
as a grey band with
95% confidence interval
s
.
As previously observed
(
9
)
, the presence of Fe(II)
-‐
EDTA leads to rapid disappearance of
nitrite
and significant underestimation of
nitrite
concentrations by this assay.
The addition of Fe(II) without a ligand, as well as with the
ligands used in this study
did not s
ignificantly affect the determination of nitrite by this assay (all measurements were conducted within 3
minutes of sulfanilamide addition to prevent nitrite loss).
●
●
●
●
●
●
0
500
1000
1500
2000
0
1
2
3
4
5
Time [min]
N
O
2
−
[
μ
M]
Ligand
●
●
CIT
EDTA
none
NTA
PPHA
Figure S6
NO
2
-‐
production by
P. denitrificans
(B) during anaerobic growth on succinate (A
). Samples for Fe(II) oxidation assays (
Figure 3
) were
taken after accumulation of ~5mM NO
2
-‐
for each biological replicate, respectively (grey shaded area indicated by arrow in panel B).
Experiment conducted in
biological
triplicate
s
. A
ll data are shown.
Time [hours]
0.0
0.5
1.0
1.5
2.0
0
2
4
6
8
10
OD
600
A
●
●
●
●
●
●
●
●
●
●
●
●
NO
2
-
[mM]
B
●
●
●
●
●
●
●
●
●
●
●
10
20
30
40
50
60
70
Figure S7
Model fits for abiotic Fe(II) oxidation by nitrite. Low citrate, no ligand, PPHA are best described by a zero
-‐
order (i.e. linear) reaction model
(linear least squares fit illustrated for these conditions instead of 2
nd
order decay).
Time [hou
r
s]
Concentration [
μ
M]
0
250
500
750
1000
1250
1500
1750
2000
F
e (II)
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
0
20
40
60
80
100
NO
2
ï
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
●
0
20
40
60
80
100
Ligand
●
none
●
2mM N
T
A
●
300mg/L PPHA
●
0.1mM CIT
●
0.5mM CIT
●
2mM CIT
●
2mM CIT + 300mg/L PPHA
Figure
S8
Evolution of N
2
O
in the headspace of sealed septum bottles
during
the
reaction of
5mM
nitrite with
~3
mM
Fe(II) complexed by citrate
vs.
NTA
(peaks normalized to Ar)
.
Retention times of the gases in the headspace were 2.2min (Ar), 3.0min (N2), 10.8min (N
2
O) and
12
-‐
13min
(CO
2
, poorly resolved).
The accumulation of N
2
O (gray band) as a reaction product could only be observed in the presence of citrate, but
not in the presence of NTA. Varying trace amounts of N
2
were present in the Ar/CO
2
headspace of the reaction vessels at the start of the
experiment but did not change significantly with reaction progress.
tp0
tp1
tp2
0 hrs
0 hrs
5 hrs
8 hrs
18 hrs
21 hrs
0.0%
2.5%
5.0%
7.5%
10.0%
0.0%
2.5%
5.0%
7.5%
10.0%
citrate
NTA
0
2
4
6
8
10
12
14
0
2
4
6
8
10
12
14
0
2
4
6
8
10
12
14
Retention time [min]
Relative TCD signal [x/Ar]
Figure S9
A
bsorption spectrum of
a
~3
mM Fe(II)
-‐
NTA solution
(dashed line)
after 950μM NO
2
-‐
was lost by abiotic oxidation of 1086μ
M Fe(II)
(21hrs data point in S3)
.
Fe(II)
-‐
NTA by itself does not absorb in this wavelength range. The oxidized Fe forms a complex with NTA that
absorbs light weakly with a characteristic peak at 470nm(dotted line).
Residual light absorption (solid line) af
ter accounting for the eff
ect
of Fe(III)
-‐
NTA in solution
is indicative of Fe(II)
-‐
NTA
-‐
NO
-‐
complex formation.
Characteristic absorption peaks of the Fe(II)
-‐
NTA
-‐
NO
-‐
complex (440nm and 600nm)
(
10
)
are
indicated in gray.
0.0
0.2
0.4
0.6
420
460
500
540
580
620
660
Wavelength [nm]
Absorbance [1/cm]