Low-pH Fe(II) oxidation influences the geochemistry of anoxic, Fe(II)-rich acid mine drainage (AMD). At many sites this is evidenced by the development of terraced iron formations (TIFs) formed via the oxidative precipitation of iron(III)-oxyhydroxysulfate minerals. In laboratory experiments using flow-through, continuously-stirred, chemostatic bioreactors, we measured GDM values ranging from 40-160 (g FeT d-1 m-2).

1. Bioreactors for low-pH Fe(II) oxidation also remove
remarkable amounts of total Fe
Bill Burgos, Yizhi Sheng, Brad Kaley –
Environmental Engineering
Appalachian Research Initiative
for Environmental Science (ARIES)

2. Rates of Fe(II) oxidation have been measured at many field
sites in the Appalachian Bituminous Coal Basin

3. “Natural” terraced iron
formations (TIFs)
“Engineered” (TIFs)

4. Rates of Fe(II) oxidation have also been measured at 3 field
sites in the Iberian Pyrite Belt (IPB)

5. Cerro Colorado, Rio Tinto mining district, Spain

6. Rio Tintillo, Huelva Province, Spain

7. AMD from the IPB has much higher conductance and dissolved metals,
and lower pH values compared to Appalachian sites
Cravotta (2008), Applied Geochemistry, 23:166-202 (n=99)
Sánchez-España et al. (2005), Applied Geochemistry, 20:1320-1356 (n=40)

8. We measured a suite of
geochemical parameters as
a function of distance
downstream from each
AMD source

9. We measured water velocity
at each sample location to
transform [conc]-vs-distance
to [conc]-vs-time
stopwatch tracer tape measure

10. 10
Assuming the stream reach behaves as a plug flow reactor,
and Fe(II) oxidation is a first-order reaction
[Fe(II)]top
[Fe(II)]bottom
kFe(II)
Accum = In – Out ± Rxn
d[Fe(II)]/dt *V = 0 – 0 – kFe(II)*[Fe(II)]*V
kFe(II) = LN(Fe(II)]bottom/Fe(II)]top) / travel time
travel time =
distance/velocity

11. -3
-2
-1
0
2.0 2.5 3.0 3.5 4.0 4.5
logkFe(II)(1/min)
pH
US sites
IPB sites
-1*pH
Fe(II) oxidation was fastest at lower pH values

12. -3
-2
-1
0
2.0 2.5 3.0 3.5 4.0 4.5
logkFe(II)(1/min)
pH
US sites
IPB sites
-1*pH
Fe(II) oxidation was fastest at lower pH values

13. -3
-2
-1
0
2.0 2.5 3.0 3.5 4.0 4.5
logkFe(II)(1/min)
pH
US sites
IPB sites
We enriched Fe(II)-oxidizers from two sites for lab tests
Scalp Level
fastest site
Brubaker Run
“average” site

14. Microbes were extracted from site sediments and enriched in
“fed-batch” mode where Fe(II) was re-spiked into reactor
52 54 56 58 60 62 64 66 68
0
5
10
15
200
5
10
15
0
210
420
630
52 54 56 58 60 62 64 66 68
TotalFe(T)Removal
Rate(mgFe(T)/L/h)
Time(d)
Oxidationrate
(mgFe(II)/L/h)
Fe(II)(mg/L)

15. A chemostatic (constant pH, T) flow-through bioreactor was
then used to measure Fe(II) oxidation rates under varied
conditions

16. Reactor pH and the influent [Fe(II)] were varied systematically
through a series of set points to measure Fe(II) oxidation rates
16
Brubaker Run Set points varied
Reactor pH
[Fe(II)]in = 300 mg/L
2.9, 2.6, 2.3, 2.6, 2.9, 3.2, 3.5, 3.8, 4.1, 3.8, 3.5
Influent [Fe(II)]
pH = 2.9
300, 80, 300, 600, 1200, 2400, 1200, 600
Scalp Level Set points varied
Reactor pH
[Fe(II)]in = 300 mg/L
2.7, 2.4, 2.1, 2.4, 2.7, 3.0, 3.3, 3.6, 3.9, 4.2
Influent [Fe(II)]
pH = 2.7
300, 66, 300, 600, 1200, 2400, 1200

17. 17
0 50 100 150 200 250 300
0.0
0.2
0.4
0.6
0.8
1.0
3.5
Pore volume
Total Fe (T)
Dissolved Fe(II)
Cout
/Cin
3.84.13.83.53.22.92.62.3pH= 2.9 2.6
Steady-state conditions for effluent [Fe(II)] were achieved for
each reactor pH set point (all experiments run at ΘH = 6 h)
Brubaker Run

18. 18
0 50 100 150 200 250 300 350 400 450 500 550
0.0
0.2
0.4
0.6
0.8
1.0
4.23.93.63.33.02.72.42.1
Cout
/Cin
Pore volume
Total Fe (T)
Dissolved Fe(II)
pH=2.7 2.4
Steady-state conditions for effluent [Fe(II)] were also achieved
with the Scalp Level chemostat

19. A general rate law can be written for biological Fe(II) oxidation
𝑅 𝑜𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛 = 𝑘 𝑏𝑖𝑜 𝐶 𝑏𝑎𝑐𝑡𝑒𝑟𝑖𝑎
𝛼
𝐷𝑂 𝛽
𝐻+ 𝛾
𝐹𝑒(𝐼𝐼) 𝛿
𝑈𝑛𝑖𝑡𝑠 𝑜𝑓 𝑘 𝑏𝑖𝑜 𝑎𝑟𝑒
𝑚𝑔 𝐹𝑒(𝐼𝐼)
𝐿 ∗ ℎ
𝑚𝐿
𝑐𝑒𝑙𝑙𝑠
−𝛼
𝐿
𝑚𝑔 𝑂2
−𝛽
𝐿
𝑚𝑜𝑙 𝐻+
−𝛾
𝐿
𝑚𝑔 𝐹𝑒(𝐼𝐼)
−𝛿
19
Simplified as follows:
• Cbacteria remained ~constant (107 cell/mL)
• DO averaged ~5 mg/L >> limiting concentration
• pH purposefully varied in experiments
𝑅 𝑜𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛 = −
𝑑 𝐹𝑒 𝐼𝐼
𝑑𝑡
= kFeII × [𝐹𝑒 𝐼𝐼 ]

20. 20
Assuming the bioreactor is completely-mixed, at steady state,
and Fe(II) oxidation is a first-order reaction
Q
[Fe(II)]in
Q
[Fe(II)]out
V
kFe(II)
Accum = In – Out ± Rxn
0 = Q*[Fe(II)]in – Q*[Fe(II)]out – kFe(II)*[Fe(II)]out*V
kFe(II) = ([Fe(II)]in – [Fe(II)]out)*Q
[Fe(II)]out*V

21. 21
-3
-2
-1
0
2.0 2.5 3.0 3.5 4.0 4.5
logkFe(II)(1/min)
pH
Brubaker Run chemostat
Scalp Level chemostat
-1*pH
Fe(II) oxidation was fastest at lower pH values

22. 22
-3
-2
-1
0
2.0 2.5 3.0 3.5 4.0 4.5
logkFe(II)(1/min)
pH
US sites
Brubaker Run chemostat
Scalp Level chemostat
Fe(II) oxidation was slower in the lab versus the field

23. 23
-3
-2
-1
0
2.0 2.5 3.0 3.5 4.0 4.5
logkFe(II)(1/min)
pH
US sites
Brubaker Run chemostat
Scalp Level chemostat
Lab rates were similar from sites with different field rates
Scalp Level
Brubaker Run

24. Reactor pH and the influent [Fe(II)] were varied systematically
through a series of set points to measure Fe(II) oxidation rates
24
Brubaker Run Set points varied
Reactor pH
[Fe(II)]in = 300 mg/L
2.9, 2.6, 2.3, 2.6, 2.9, 3.2, 3.5, 3.8, 4.1, 3.8, 3.5
Influent [Fe(II)]
pH = 2.9
300, 80, 300, 600, 1200, 2400, 1200, 600
Scalp Level Set points varied
Reactor pH
[Fe(II)]in = 300 mg/L
2.7, 2.4, 2.1, 2.4, 2.7, 3.0, 3.3, 3.6, 3.9, 4.2
Influent [Fe(II)]
pH = 2.7
300, 66, 300, 600, 1200, 2400, 1200

25. 0 50 100 150 200
0.0
0.2
0.4
0.6
0.8
1.0
600
Total Fe (T)
Dissolved Fe(II)
Cout
/Cin
Pore volume
Fe(II)in=
300mg/L
80 300 600 1200 2400 1200
Experiments were also conducted with variable
influent [Fe(II)] (all at pH 2.90 for Brubaker Run)

26. 26
0 500 1000 1500 2000 2500
0
1E-7
2E-7
3E-7
4E-7
5E-7
6E-7
7E-7
8E-7
Influent Fe(II) (mg/L)
chemostat_SL
chemostat_BR
OxidationRates(mol/L/s)
Fe(II) oxidation rates were highest at higher influent Fe(II)
concentrations
Scalp Level chemostat pH 2.7
Brubaker Run chemostat pH 2.9

27. 𝑅 𝑜𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛 = kbio
∗ 𝐻+
𝐾 𝐻+ + 𝐻+
𝐹𝑒(𝐼𝐼)
𝐾𝐹𝑒(𝐼𝐼) + 𝐹𝑒(𝐼𝐼)
𝑈𝑛𝑖𝑡𝑠 𝑜𝑓 kbio
∗
𝑎𝑟𝑒
𝑚𝑜𝑙 𝐹𝑒(𝐼𝐼)
𝐿 ∗ 𝑠
From all these chemostat experiments, we attempted to
parameterize a multi-Monod rate law for Fe(II) oxidation
27

28. 0.000 0.002 0.004 0.006 0.008
1.0E-7
1.2E-7
1.4E-7
1.6E-7
1.8E-7
2.0E-7
2.2E-7
2.4E-7
2.6E-7
OxidationRate(mol/L/s)
H+
(mol/L)
chemostat_SL
chemostat_BR
SL data fit
BR data fit
All data fit
Measured rates display “saturation effect” of [H+]
Scalp Level chemostat
Brubaker Run chemostat
KH+

29. 0 500 1000 1500 2000 2500
0.0
2.0E-7
4.0E-7
6.0E-7
8.0E-7
1.0E-6 chemostat_SL
chemostat_BR
SL data fit
BR data fit
All data fit
OxidationRate(mol/L/s)
Influent Fe(II) (mg/L)
Measured rates also display saturation effect of [Fe2+]
Scalp Level chemostat
Brubaker Run chemostat
KFe(II)

30. 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7
0
1E-7
2E-7
3E-7
4E-7
5E-7
6E-7
7E-7
8E-7
9E-7
OxidationRate(mol/L/s)
chemostat_SL_pH series
chemostat_BR_pH series
chemostat_SL_Fe(II) series
chemostat_BR_Fe(II) series
Linear Fit
y=8.93289*10
-7
x+4.01526*10
-8
r=0.91623
+
in
+
in
+
Fe(II) inH
[Fe(II) ][H ]
*
(K H ) (K Fe(II) ) 
We can now predict rate of low-pH Fe(II) oxidation based
solely on pH and influent [Fe2+]
k*
bio

31. aerationAMD Fe(II) oxidation
lime
neutralization
sedimentation discharge
low-pH Fe(II)
oxidizing
bioreactor
AMD
schwertmannite
removal
lime
neutralization
sedimentation discharge
A
B
Low-pH Fe(II) oxidation can be incorporated into an active
treatment plant

32. Low-pH Fe(II) oxidation can be
engineered like a conventional
wastewater treatment unit
Janneck, E., I. Arnold, T. Koch, J. Meyer, D. Burghardt and S. Ehinger (2010) “Microbial
synthesis of schwertmannite from lignite mine water and its utilization for removal of
arsenic from mine waters and for production of iron pigments.” Mine Water and
Innovative Thinking IMWA 2010, p. 131-134.
Previous failures caused by
accumulation of heavy mineral
precipitates can be overcome and
exploited by design modifications
Photograph of vertically-oriented growth
media totally encrusted with iron
minerals that are harvested and sold as
pigments or used for arsenic removal
Reactor schematic from group in Germany

33. Bioreactor systems proposed for low-pH Fe(II) oxidation and low-pH
Fe removal, followed by sulfate-reduction for metal recovery from
acid pit lakes
pH 2.1  2.3 NaOH  pH 3.5 pH 3.5  3.2
Hedrich and Johnson (2012) Bioresource Technology 106:44–49

34. 34
Fe(T) removal was calculated as a zero-order rate, and
normalized based on plan area of the bioreactor
Q
[Fe(T)]in
Q
[Fe(T)]out
V
GDM = g Fe(T) removed per day per m2 land area
GDM = ([Fe(T)]in – [Fe(T)]out)*Q
Atop
Atop

35. 2.0 2.5 3.0 3.5 4.0 4.5
0
20
40
60
80
100
120
140
160
180
GDM(gFe(T)/d/m2
)
chemostat_SL
chemostat_BR
pH
Fe(T) removal was exceptionally high ~pH 3
GDM = 4 – 8
sheet-flow
“chunk”
reactors
GDM = 20
settling pond
Scalp Level chemostat 300 mg/L
Brubaker Run chemostat [Fe(II)]in

36. Fe(T) removal increased with influent Fe(II)
GDM = 20
settling pond
0 500 1000 1500 2000 2500
0
100
200
300
400
500
GDM(gFe(T)/d/m2
)
Influent Fe(II) (mg/L)
chemostat_SL
chemostat_BR
Scalp Level chemostat pH 2.7
Brubaker Run chemostat pH 2.9

37. Terraced iron formations (TIFs) formed via low-pH Fe(II)
oxidation can be used in passive treatment systems

38. natural TIF Alkalinity
influent effluent
natural TIF Alkalinity
influent effluent
Scalp Level – very low emergent pH Upper Red Eyes – pH 4.1
2
3
4
5
pH[Fe2+]or[Fe3+][TotalFe]
distance
Fe2+
Fe3+
[Fe2+]or[Fe3+][TotalFe]
distance
Fe2+
Fe3+
2
3
4
5
pH

39. Conclusions
• Low-pH Fe(II) oxidation is fastest at lower pH and
higher [Fe(II)]
• Rates of Fe(II) oxidation can be predicted based on pH
and [Fe(II)]in
• At pH 3 and 300 mg/L Fe(II), suspended-growth
bioreactors can remove Fe(T) at 100 – 400 g Fe/d*m2
• Attached-growth bioreactors will likely remove even
more Fe(T)
• “Engineered” terraces are also effective in passive
treatment systems