This study examined the effect of initial manganous ion (Mn2+) concentration and oxidant
dose on the oxidation of Mn2+ to Mn oxide during water treatment. The oxidants studied were
chlorine dioxide (ClO 2 ), potassium permanganate (KMnO 4 ), and ozone. Initial Mn 2+
concentrations were 60, 200, and 1,000 µg/L, and the goal of the treatment was a final Mn2+
<10 µg/L. Bench-scale experiments were performed by applying different doses of each
oxidant to a raw surface water and measuring Mn2+ residuals over time. For all experiments,
the ambient raw water conditions were pH 7.0, 9oC, and total organic carbon = 3.4 mg/L.
Oxidation kinetics for low initial Mn2+ concentrations (60 and 200 µg/L) were significantly
slower than for high concentrations (1,000 µg/L). For the lower initial Mn2+ levels, ClO2 was the
only oxidant that consistently produced final Mn2+ <10 µg/L, generally within 60–120 s. All
oxidants produced final Mn2+ <10 µg/L when the initial Mn2+ concentration was 1,000 µg/L.
Oxidation with ozone consistently resulted in soluble Mn2+ >20 µg/L after 5 min, with the
formation of permanganate ion in many cases. For low Mn2+ concentrations, 90% oxidation
by KMnO4 required 15–30 min.
EFFECT OF
soluble Mn concentration
on
oxidation kinetics
BY DEAN GREGORY
AND KENNETH CARLSON
uring water treatment processes and in the distribution system, the
manganous ion (Mn2+) oxidizes to Mn oxide (MnO2(s)). The MnO2(s),
in turn, causes brown staining of plumbing fixtures and laundry and
“brown water” incidents. The Fort Collins (Colo.) Water Treatment
Facility (FCWTF) has experienced Mn2+ levels up to 450 µg/L in Horsetooth Reservoir, one of its two surface water supplies. The source of Mn2+ in the
reservoir is the chemical reduction of MnO2(s) in the lower depths of the reservoir, which becomes a reducing environment after stratification and deoxygenation of the reservoir’s hypolimnion.
The US Environmental Protection Agency has established a secondary maximum contaminant level of 50 µg/L total Mn for finished water because of its nuisance characteristics. The FCWTF, however, has experienced Mn problems from
effluent concentrations as low as 20 µg/L. Sly et al (1990) suggested that the
ultimate goal for Mn in finished water should be 10 µg/L, and the FCWTF has
adopted this goal.
Although Mn2+ oxidation has been practiced for years in water treatment, many
studies reported in the literature have not fully explored oxidation of low initial
Mn2+ concentrations and dose requirements—issues that may be important to
water treatment plants. This study sought to address these gaps in knowledge of
Mn2+ oxidation for potable water treatment.
D
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Dissolved Mn—µg/L
Dissolved Mn—µg/L
Dissolved Mn—µg/L
The objective of Mn removal processes
Bench-scale oxidation of Mn2+ by ClO2 for three initial Mn2+ concentrations
FIGURE 1
is to oxidize Mn2+ to MnO2(s), a solid
precipitate that can be removed by
solid/liquid separation processes such as
300% CIO2
400% CIO2
200% CIO2
sedimentation and filtration. Candidate
60
2+
Initial Mn = 60 µg/L
oxidants for the reaction include ozone
50
(O3), chlorine dioxide (ClO2), and potassium permanganate (KMnO4). Free chlo40
rine, an obvious alternative for water
treatment, is often impractical because
30
the Mn2+ oxidation reaction kinetics are
20
slow, requiring 2–3 h of contact time at
pH 8.0 (Knocke et al, 1990). For this rea10
son, chlorine was excluded as a candidate oxidant for this study.
0
0
50
100
150
200
250
300
The overall objective of the study was
Reaction Time—s
to evaluate ClO2, KMnO4, and O3 for
100% CIO2
200% CIO2
300% CIO2
the oxidation of Mn2+ to <10 µg/L. Spe220
cific objectives were to understand the
2+
Initial Mn = 200 µg/L
200
effect of initial Mn2+ concentration, par180
ticularly low initial concentrations, on
160
oxidant dose. Because this research was
140
conducted for the FCWTF, it also served
120
100
as an engineering analysis of Mn oxida80
tion for the water treatment plant, which
60
is a fairly typical conventional treatment
40
facility. The basic plant processes include
20
coagulation with alum added in a rapid0
mix basin, flocculation/sedimentation
0
50
100
150
200
250
300
(approximately 60-min detention time
Reaction Time—s
using inclined plates), rapid-rate filtra100% CIO2
tion with dual-media filters (sand/
1,100
2+
1,000
anthracite), and disinfection with chloInitial Mn = 1,000 µg/L
900
rine before distribution. The experimen800
tal conditions (pH, total organic carbon
700
[TOC], temperature, alkalinity) were
600
based on the typical value of each con500
stituent observed at the plant during the
400
5 µg/L
300
late summer/fall season when influent
200
Mn concentrations increase.
100
The other treatment plant parameter
0
important to this analysis was available
0
50
100
150
200
250
300
detention time between the raw water
Reaction Time—s
intake and the addition of alum in the
Relative stoichiometric doses are shown; temperature—9 C, pH—7.0, total organic
2+
rapid-mix basin. This reaction time was an
carbon—3.4 mg/L; CIO2—chlorine dioxide, Mn—manganese, Mn —manganous ion
important consideration because the
FCWTF had observed, while using a temporary KMnO4 system, that depression
µg/L. Before this study was conducted, the FCWTF used
of the pH from 7.0 to 6.3–6.4 in the rapid mix effectively
KMnO4 for Mn oxidation and removal, but results were
halted the KMnO4–Mn2+ reaction. During high-flow periods, the detention time is only 120–180 s. Detention time
often unsatisfactory in terms of finished water Mn conis approximately 180–300 s during periods of high Mn concentrations, particularly for low influent dissolved Mn
centration. Two criteria for selecting the appropriate chemlevels. The FCWTF decided to reassess its use of KMnO4
as well as to consider the alternatives of O3 and ClO2.
ical oxidant for the FCWTF were whether the Mn2+–oxidant reaction was completed within approximately 180 s
Another Mn removal option considered by FCWTF
and whether it produced a final Mn2+ residual of <10
was using MnO2(s)-coated filter media, or the “greeno
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99
TABLE 1
Water quality characteristics of Horsetooth Reservoir
Constituent
Range
pH
7.0–7.1
Temperature—oC
9–10
Alkalinity—mg/L as CaCO3
25–30
Total organic carbon—mg/L
3.4–3.6
SUVA*—L/mg-min
1.7
Dissolved oxygen—mg/L
1.8–2.1
Ambient Mn2+†—µg/L
10–13
Turbidity—ntu
3.5–4.5
*SUVA—specific ultraviolet absorbance
†Mn2+—manganese
sand” process, a method with demonstrated effectiveness
in removing relatively high concentrations of dissolved Mn
(Coffey et al, 1993; Knocke et al, 1991a). In this process,
Mn2+ adsorbs to the MnO2(s)-coated surface of the media
and is subsequently oxidized by free chlorine. This method,
however, appears to be most effective for groundwater
applications in which the influent dissolved Mn concentration is relatively high (>1 mg/L) and constant throughout the year. Because the dissolved Mn in Horsetooth
Reservoir is only an issue for approximately three months
of the year and the influent dissolved Mn levels vary during that period, preoxidation appeared to be a more effective alternative. Preoxidation could also provide treatment options for other challenging water quality
conditions such as taste and odor compounds.
Literature review. It has been shown that the oxidation
of Mn2+ by a reactive oxidant involves three mechanisms
(Morgan, 1967): solution-phase oxidation, adsorption
of the Mn2+ ion to MnO2(s), and surface-catalyzed oxidation of the sorbed Mn2+ ion. These mechanisms are
represented by Eqs 1–3.
k1
[Mn2+] + [Ox] → [MnO2]
[Mn2+]
(1)
k2
+ [MnO2] → [MnO2 ⬅ Mn]
k3
[MnO2 ⬅ Mn] + [Ox] → [MnO2]
(2)
(3)
in which MnO2 ⬅ Mn represents MnO2(s) with sorbed
Mn2+ ion (Van Benschoten et al, 1991) and Ox is oxidant.
Van Benschoten et al (1991) developed rate equations and a kinetic model based on Eqs 1–3 that predict
rapid oxidation of Mn2+ by ClO2 and KMnO4—nearly
complete oxidation of initial Mn2+ within 30 s. In a
study that focused on KMnO4, Carlson and Knocke
(1999) showed that the model could be expanded to
better characterize Mn2+ oxidation in a natural water
that contained significant concentrations of natural
organic matter (NOM). KMnO4 oxidation rates were
found to be significantly lower in that study. The rate
equations were expanded to include a term that
accounted for dissolved organic carbon (DOC) concentration, because the DOC concentration in the experiments Van Benschoten et al (1991) used to develop
their model was <1 mg/L. Carlson and Knocke (1999)
suggested that DOC interfered in the KMnO4–Mn2+
reaction by exerting a KMnO4 demand and possibly
complexing a small fraction of Mn2+, rendering it less
likely to be oxidized. This approach resulted in a model
that more closely characterized Mn 2+ oxidation by
KMnO4 in the raw surface water used in the study.
The work of Knocke et al (1990) is the most comprehensive investigation of Mn2+ oxidation for water treatment. In addition to their kinetic modeling work, the
authors examined the effect of pH, DOC, temperature,
and initial Mn2+ concentrations on the oxidation of Mn2+
by ClO2, KMnO4, O3, and other oxidants. They concluded that Mn2+ was oxidized rapidly by ClO2, KMnO4,
and O 3 in low-DOC waters over a pH range of pH
5.5–9.0.
Two factors, however, complicate the application of
these results to water treatment. One is that the lowest
concentrations used in the ClO2, KMnO4, and O3 experiments were 500, 200, and 850 µg/L, respectively. Experiments that examined the effect of independent variables
such as pH, DOC, and temperature used initial Mn2+
concentrations in the range of 900–1,000 µg/L. Many
water treatment facilities must practice Mn2+ oxidation
and removal at much lower concentrations, e.g., 50 µg/L,
to avoid aesthetic problems. Because the authors did show
that Mn2+ oxidation rates clearly decreased with decreasing initial Mn2+ concentrations, additional information on
the effect of lower concentrations could benefit the water
supply industry. A second factor that could complicate the
application of these results to water treatment plants is that
although oxidation rates appear to be rapid for all three
oxidants, the final Mn2+ residuals are often unacceptable, in terms of full-scale treatment. In the case of
KMnO4, Mn2+ residuals were generally in the range of
50–150 µg/L. As discussed earlier, FCWTF desires finished water Mn2+ <10 µg/L. In ClO2 experiments, Mn2+
residuals close to 0 µg/L were only observed when DOC
concentrations were <1 mg/L and initial Mn2+ concentrations were approximately 900 µg/L. For other conditions, i.e., lower initial Mn2+ concentrations and the presence of significant DOC, final Mn2+ residuals ranged
from 50 to 400 µg/L.
Paillard et al (1989) examined, among other things, the
effect of TOC on the oxidation of Mn2+ by O3. They
performed bench-scale experiments that simulated a conventional full-scale process, including sand filtration, with
2003 © American Water Works Association
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FIGURE 2
300% KMnO4
2+
Initial Mn = 60 µg/L
80
70
60
50
40
30
20
10
0
0
200
400
600
800
1,000
1,200
1,400
1,600
1,800
2,000
Reaction Time—s
MATERIALS AND METHODS
100% KMnO4
220
200
Dissolved Mn—µg/L
200% KMnO4
300% KMnO4
2+
Initial Mn = 200 µg/L
180
160
140
120
100
80
60
40
20
0
0
200
400
600
800
1,000
1,200
1,400
1,600
1,800
2,000
1,400
1,600
1,800 2,000
Reaction Time—s
100% KMnO4
1,000
900
Dissolved Mn—µg/L
Experimental water. The raw water
used for all experiments was from Horsetooth Reservoir in Fort Collins. This
water originates from high alpine sources
near the Continental Divide in the Rocky
Mountains. Table 1 provides the water
quality characteristics that were used for
all bench-scale experiments. Temperature (9 o C), pH (7.0), and TOC (3.4
mg/L) were nearly constant for all experiments. The specific ultraviolet
absorbance (SUVA) value for this water
is also provided in Table 1. SUVA, the
UV absorbance at 254 nm (m–1) divided
by the DOC concentration (mg/L), is an
indicator of the aromatic character of
NOM (Edzwald & Van Benschoten,
1990). The SUVA value increases with an
increasing fraction of aromatic or double bond–rich substances that tend to
absorb more strongly at 254 nm. The
SUVA value of 1.7 L/mg-m for Horsetooth Reservoir is relatively low and indicates that a significant fraction of the
NOM is hydrophilic in nature. The low
dissolved oxygen values shown in Table
1 are due to stratification of the reservoir
in late summer and fall.
Generation of oxidant stock solutions/analytical methods. Ozone was produced using a 1-lb/d (0.45 kg/d) pilotscale generator.1 A stainless-steel gas
conduit system was constructed to
deliver an O3/oxygen gas blend to the
Bench-scale oxidation of Mn2+ by KMnO4 for three initial Mn2+
concentrations
100% KMnO4
90
Dissolved Mn—µg/L
O3 applied to settled water. They showed
that for the same mg O3:TOC concentration ratio, increasing the TOC concentration increases the residual Mn2+
concentration. The initial Mn2+ concentration used was 250 µg/L.
Another study that investigated the
effect of organics on the oxidation of
Mn2+ by O3 offered a contradictory finding. Seby et al (1995) studied the influence of fulvic acid concentrations
(0.5–5.0 mg/L) on Mn2+ oxidation efficiency in a synthetic water. They reported
that 94% of initial Mn2+ was oxidized at
pH 8, regardless of fulvic acid concentration. The authors concluded the oxidation of Mn2+ by O3 has priority over
those of fulvic acids. These experiments
used a relatively high initial Mn2+ concentration of 660 µg/L.
2+
Initial Mn = 1,000 µg/L
800
700
600
500
400
300
200
100
0
0
200
400
600
800
1,000
1,200
Reaction Time—s
o
Relative stoichiometric doses are shown; temperature—9 C, pH—7.0, total organic
carbon—3.4 mg/L; KMnO4—potassium permanganate, Mn—manganese,
2+
Mn —manganous ion
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101
FIGURE 3
Bench-scale oxidation of Mn2+ by O3 for three initial Mn2+ concentrations
O3 = 0.2 mg/L
70
O3 = 0.5 mg/L
O3 = 1.5 mg/L
O3 = 2.0 mg/L
2+
Initial Mn = 60 µg/L
Dissolved Mn—µg/L
60
50
40
30
20
10
0
0
220
200
100
200
O3 = 0.5 mg/L
300
400
Reaction Time—s
O3 = 0.75 mg/L
500
O3 = 1.0 mg/L
600
700
O3 = 2.1 mg/L
2+
Initial Mn = 200 µg/L
Dissolved Mn—µg/L
180
about 60 mg/L to about 40 mg/L during this time. The O3 stock concentration was measured before each
experiment. The aqueous O3 concentration in the O3 stock solutions and
the residual O3 in experiments were
measured by the indigo trisulfonate
spectrophotometric method (Bader &
Hoigné, 1981).
ClO2 stock solutions were generated by a solid-phase ClO2 generator2
in which a 4% Cl2:96% N2 gas mixture is passed through a packed bed
of sodium chlorite. The gas was diffused into cold (2–3oC) DI water for
approximately 30 min, producing a
ClO2 stock solution concentration of
approximately 4,000 mg/L. The purity
of each stock solution was measured
according to the following formula:
160
% Purity = CClO2/(CClO2 + CChlorite
+ CFree chlorine)
140
120
100
in which C is the concentration of each
species in milligrams per litre. The purity
40
of all stock solutions was >99%. Stock
20
concentrations were measured daily
0
using amperometric titration (for stock
0
100
200
300
400
500
600
700
solutions 0.1 N titrant is used instead of
Reaction Time—s
the 0.00564 N titrant used for residual
analyses) and were found to be stable
O3 = 1.0 mg/L
O3 = 1.5 mg/L
O3 = 2.0 mg/L
O3 = 3.0 mg/L
when the solution was stored in the dark
1,000
and refrigerated. ClO2 residuals were
2+
900
Initial Mn = 1,000 µg/L
measured using the amperometric titra800
tion method (method 4500-ClO2 E),
700
(Standard Methods, 1998).
600
Stock solutions of KMnO4 (5,000
500
mg/L) were produced by dissolving
400
granular KMnO4(s) in DI water. The
300
concentration of the MnO 4– stock
200
solutions, which were refrigerated at all
100
times, did not decay significantly over
0
a period of several days. Fresh MnO4–
100
200
300
400
500
600
700
0
stock solutions were made weekly.
Reaction Time—s
Residual Mn analyses were perTemperature—9 C, pH—7.0, total organic carbon—3.4 mg/L; Mn—manganese,
2+
formed by atomic absorption specMn —manganous ion, O3—ozone
trometry (method 3500-Mn B) (Standard Methods, 1998). Samples were
acidified after filtration, stored in 125bench-scale apparatus located under a fume hood. The gas
mL plastic bottles, and refrigerated before analysis.
was bubbled into cold (2–3oC) deionized (DI) water in a
TOC concentrations were measured by an analyzer
1-L heat-resistant glass flask for 15 min. The O3 stock
that uses the persulfate–ultraviolet oxidation method3
(method 5310 C) (Standard Methods, 1998).
solutions decayed slowly when sealed and kept in an ice
Bench-scale protocol. The general procedure used for
bath. Each stock was used for several bench-scale experall experiments was to (1) determine the concentration of
iments over 3–4 h. The stock concentration changed from
80
Dissolved Mn—µg/L
60
o
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Dissolved Manganese Residual—µg/L
the oxidant stock solution, (2) inject the
O3 dosage versus dissolved Mn residual after 7-min reaction time
FIGURE 4
appropriate volume of the stock into the
sample water at t = 0 to achieve a specific applied dose, and (3) collect and
Mn = 0.06 mg/L
Mn = 0.2 mg/L
Mn = 1.0 mg/L
100
filter Mn samples to characterize the
kinetics of Mn2+ oxidation for O3, ClO2,
90
and KMnO4. For ClO2 and KMnO4,
80
the same stock solutions were often used
70
for several days, and the concentrations
60
generally remained constant. Because
O3 decays relatively rapidly, new stock
50
solutions were required after three or
40
four experiments.
30
The oxidant was injected into a
20
closed 1-L Erlenmeyer-type reaction flask
at t = 0 and was mixed for several sec10
onds to ensure an even distribution
0
0.5
1.0
1.5
2.0
2.5
3.0
3.5
throughout the sample volume. The
0
Ozone
Dose—
mg/L
water temperature was monitored and
maintained at approximately 9oC, the
Temperature—9 C, pH—7.0, total organic carbon—3.4 mg/L; Mn—manganese,
O3—ozone
average temperature of Horsetooth
Reservoir water during stratification,
throughout each experiment using a
kinetics of the O3–Mn2+ reaction. Rather, the intent was
water bath. All experiments were conducted at pH 7.0.
to characterize the latter stages of the interaction and,
The 1-L reaction flask was equipped with a bottle-top,
more important, to determine the equilibrium concenadjustable-volume dispenser that was flushed before each
trations of the dissolved Mn residuals. A majority of the
sample was dispensed. The filtration step used an ultraO3–Mn2+ interaction occurs within 60 s, as shown by
filtration apparatus with a 30K (30,000 AMU molecular
4
data from other researchers (Reckhow et al, 1991; Knocke
weight cut-off) ultrafiltration membrane. The ultrafiltration step was performed under pressure (approximately
et al, 1990), and by the Mn2+ oxidation and O3 decay data
provided in this article. The time interval between sam30 psi [207 kPa]), which was supplied by connecting a canples was dictated by the minimum time required for the
ister of compressed nitrogen gas to the apparatus. Filterfiltration process to be completed and the filter apparaing time was about 150 s per 25-mL sample. This filtratus to be cleaned and prepared for the next sample. In the
tion procedure ensured that only dissolved Mn species
o
The oxidant was injected into a closed 1-L Erlenmeyer-type reaction flask at t = 0 and was mixed for several
seconds to ensure an even distribution throughout the sample volume.
remained in the final filtrate sample (Carlson et al, 1997;
Reckhow et al, 1991). The pH of the filtered samples
was checked periodically. It remained virtually constant
throughout the experimental procedure. Before use, new
30K membranes were soaked in DI water to remove the
glycerin coating from their surfaces. Additionally, 500
mL of DI water was filtered through each new membrane
before it was used in experiments.
Sample collection. The sample times used were based
on oxidant-specific considerations such as the expected
kinetics of the oxidant–Mn reaction (whether or not a
quenching agent, which eliminates the oxidant residual at
the time of sample collection, was used) and the filtration
step required to separate dissolved from particulate and
colloidal Mn. In the case of O3, for example, sample
times were not intended to characterize the initial, rapid
O3 experiments, a quenching agent was not used (to avoid
reducing residual MnO4– to MnO2 (s)); therefore, samples
were filtered immediately. Further Mn oxidation may
have occurred during the filtration process for the sample collected at t = 1 min. For the samples collected at t
= 3, 7, and 11 min, further oxidation of Mn by O3 was
negligible.
Quenching agents were used for the ClO2 and KMnO4
experiments, and samples, therefore, did not need to be
filtered immediately as in the O 3 experiments. This
allowed samples to be collected more frequently. Preliminary experiments indicated that the ClO2–Mn reaction was far more rapid, particularly at lower initial
Mn2+ concentrations, than the MnO4––Mn reaction. As
a result, the duration of the ClO2 and KMnO4 experiments was 10 and 30 min, respectively. For ClO2, Mn
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103
samples collected at t = 30, 60, 90, 180, 300, and 600 s
adequately characterized the reaction kinetics within the
first few minutes as well as the equilibrium concentrations. For the KMnO4 experiments, samples were collected at approximately the same frequency for the first
10 min, but two additional Mn samples, at t = 1,200
and 1,800 s, were also obtained.
Quenching agents. An important aspect of the benchscale experiments was the selection of the quenching
(reducing) agent used to eliminate the oxidant residual and
halt the Mn 2+–oxidant reactions. Two features of a
quenching agent were desired: immediate and complete
reaction with the oxidant residual and the inability to
resolubilize previously formed MnO2(s).
In the O3 experiments, the primary reason for not
using a quenching agent was that one focus of this work
the ClO2 experiments. The results in Figure 2 illustrate
that oxidation of Mn2+ by KMnO4 is considerably slower
than by ClO2 in this water (note that different time scales
are used in Figures 1 and 2). For initial Mn2+ = 60 µg/L,
about 1,800 s were required for complete oxidation with
a 300% stoichiometric dose (the stoichiometry of the
reaction is 1.9 mg KMnO4/mg Mn2+). Even after this
reaction time, the final Mn2+ residual was >10 µg/L. For
initial Mn2+ = 200 µg/L, 1,200 s were required to produce
Mn2+ residuals near 10 µg/L. At initial Mn2+ = 1,000
µg/L, the oxidation rate and final Mn2+ residual were
comparable to the ClO2 results.
Figure 2 shows that the impact of initial Mn2+ concentration is more significant for KMnO4 than ClO2,
when considering the finished water Mn goal of 10 µg/L.
Final Mn2+ residual <10 µg/L was only observed when the
An important aspect of the bench-scale experiments was the selection of the quenching (reducing) agent
used to eliminate the oxidant residual and halt the Mn 2+–oxidant reactions.
was the formation of MnO4–, which appeared to occur in
a majority of the experiments. The use of a quenching
agent would have reduced some of the residual MnO4– to
MnO2(s), which would have been removed during the filtration process and made it impossible to determine the
true concentration of total dissolved Mn, not just Mn2+,
remaining in solution following ozonation.
In ClO2 experiments, a 0.1 N potassium iodide solution was a satisfactory quenching agent. For KMnO4
experiments, 0.00564 N phenylarsine oxide was used.
Experiments were conducted to determine the minimum
effective dose of each quenching agent to avoid overdosing the quenching agents and the potential to reduce
MnO2(s) to soluble forms of Mn.
RESULTS AND DISCUSSION
ClO2. Figure 1 shows the results of bench-scale oxidation experiments using ClO2 with three initial Mn2+
concentrations: 60, 200, and 1,000 µg/L. The relative
stoichiometric doses of ClO2 used ranged from 100 to
400% (the stoichiometry of the reaction is 2.5 mg
ClO2/mg Mn2+). Mn2+ was oxidized to <10 µg/L within
120 s for initial Mn2+ = 200 and 1,000 µg/L. At the lowest Mn2+ concentration of 60 µg/L, the dissolved Mn
concentration was <10 µg/L after 300 s. The only experimental conditions using ClO2 that resulted in a final
Mn2+ residual >10 µg/L were at the lowest initial Mn2+
level and a dose of 200% stoichiometry (0.3 mg/L).
Although oxidation was rapid for all experimental conditions, it was clear that increasing initial Mn2+ concentrations increased oxidation rates.
KMnO4. Figure 2 shows the results of bench-scale oxidation experiments using KMnO4. The initial Mn2+ concentrations were approximately the same as those used for
initial Mn2+ = 1,000 µg/L. The oxidation rates at this
concentration are far more rapid than for initial Mn2+ =
60 and 200 µg/L. The rapid oxidation rate observed in
experiments using an initial Mn2+ concentration of 1,000
µg/L indicates that solution-phase oxidation proceeds as
rapidly as surface-catalyzed oxidation in these conditions. These results suggest that for KMnO4, sorption of
Mn2+ onto MnO2(s) and subsequent surface oxidation
are the primary mechanism at low initial Mn2+ concentrations (when the solution-phase reaction proceeds
slowly). This was not the case for ClO2.
O3. Figure 3 shows the results of Mn2+ oxidation versus time by O3. The trendlines in Figure 3 are generally
flat after the 60-s point, indicating that the O3–Mn2+
reaction reached completion in less than 60 s. Figure 3 and
experimental observations indicate that the reaction is
rapid, probably reaching completion within 30 s. Final
Mn2+ residuals <10 µg/L were only observed in experiments using initial Mn2+ = 1,000 µg/L.
Figure 4 shows the effect of O3 dose on Mn2+ residuals after 7 min of reaction time for the three initial Mn2+
concentrations used in Figure 3. The 7-min samples were
used to represent equilibrium conditions. Figure 4 shows
that increasing O3 dosages above the optimum increases
Mn2+ residuals, and Mn2+ residuals from initial Mn2+ =
60 and 200 µg/L are always >20 µg/L. The increasing
Mn2+ residuals observed at higher O3 dosages appear to
be due to the formation of MnO4–, which is a dissolved
species and results from the oxidation of a small fraction of Mn(II) to Mn(VII). Mn2+ is essentially oxidized
beyond the desired Mn(IV) state to the undesirable dissolved state, Mn (VII).
Optimizing O3 dosages for minimum Mn2+ residuals
required a balance between overcoming O3 demand
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experimental pH, TOC, and temperature conditions showed that dosages
below this range increase Mn2+ residuals because of competing O3 demand
300% KMnO4
O3 = 1.5 mg/L
400% CIO2
from NOM (Gregory & Carlson,
90
2001). Increasing dissolved Mn resid80
uals observed for higher O3 dosages
70
were due primarily to the formation of
60
MnO4–.
50
Oxidation of Mn2+ by O3 proceeds
40
rapidly. The issue is that the potential
for MnO 4– formation exists for all
30
doses in which the stoichiometry for
20
2+
Initial Mn = 60 µg/L
the MnO4–-forming reaction (2.2 mg
10
O3/mg Mn2+) is satisfied. The distinc0
tive pink color of MnO4– was observed
0
100
200
300
400
500
600
in many experiments. In this study, to
Reaction Time—s
achieve effective Mn2+ removal for all
300% KMnO4
O3 = 1.0 mg/L
200% CIO2
but the highest initial Mn2+ concentra200
tion, the background demand exerted
180
by NOM required the use of O 3
160
dosages that were sufficient for the pro140
duction of MnO4–. As a result, dis120
solved Mn residuals <10 µg/L were not
100
attainable for initial Mn2+ concentra80
tions of 60 and 200 µg/L.
2+
Initial Mn = 200 µg/L
60
The formation of MnO4– in the O3
40
experiments implies that the dissolved
20
Mn residuals measured were due to
0
MnO4– or a combination of MnO4–
0
100
200
300
400
500
600
and Mn2+. Because the samples from
Reaction Time—s
these experiments were not quenched,
no distinction was made between the
O3 = 1.5 mg/L
100% CIO2
100% KMnO4
two species. Therefore, the results that
1,200
show “dissolved Mn” residuals for these
1,000
experiments include Mn2+ and MnO4–.
Comparison of oxidants. Figure 5
800
compares Mn2+ removal for the most
600
effective doses of each oxidant at each
initial Mn2+ concentration. When com2+
400
Initial Mn = 1,000 µg/L
pared with the results for KMnO4 and
O3, the data in Figure 5 indicate that
200
ClO2 was the most effective oxidant
0
for reducing dissolved Mn residuals to
0
100
200
300
400
500
600
<10 µg/L. Previous work (Knocke et
Reaction Time—s
al, 1990) suggested that Mn2+ oxidation
2+
rates by KMnO4 and O3 were similar to
Oxidant doses shown are the optimum at each initial Mn ; temperature—9 C, pH—7.0,
total organic carbon—3.4 mg/L; CIO2—chlorine dioxide, KMnO4—potassium
ClO2 and that all three were rapid,
2+
permanganate, Mn—manganese, Mn —manganous ion, O3—ozone
effective oxidants for Mn2+. These studies, however, generally used high iniexerted by NOM and overdosing, which causes the fortial Mn2+ concentrations (close to 1 mg/L), and the focus
mation of MnO4–. As a result, there is a relatively narrow
was not on attaining final Mn residuals of <10 µg/L. The
optimum dose—particularly for higher initial Mn2+ concurrent study emphasized attaining low dissolved Mn
centrations. Dosages above and below this range result in
residuals within a relatively short period of time (approxsignificantly higher dissolved Mn residuals. A previous
imately 180 s). Figures 1–3 and Figure 5 indicate that
study using Horsetooth Reservoir water and the same
there can be significant differences between the reaction
Bench-scale comparison of oxidation of Mn2+ by KMnO4, ClO2, and O3
at three initial Mn2+ concentrations
Dissolved Mn—µg/L
Dissolved Mn—µg/L
Dissolved Mn—µg/L
FIGURE 5
o
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105
MnO2 ≡ Mn—mg/L
MnO2 (s)—mg/L
+2
Mn /Mn
+2
0
rates and final Mn residuals for each
Results of modeling analysis showing effect of initial Mn2+ concentration on
FIGURE 6
oxidant.
concentrations of Mn2+, MnO2(s), and MnO2 ⬅ Mn2+ concentrations versus
From these results, ClO2 appeared
time for oxidation by KMnO4
to be the optimum alternative for two
reasons. First, when compared with the
KMnO4 results, the ClO2–Mn2+ reacInitial Mn = 0.06 mg/L
Initial Mn = 0.2 mg/L
tion kinetics were significantly more
Initial Mn = 1.0 mg/L
rapid (Figures 1 and 2), particularly for
1.2
initial Mn2+ = 60 and 200 µg/L. Sec1.0
ond, ClO 2 did not appear to create
–
MnO 4 , an undesirable dissolved
0.8
species, as did O3. Although MnO4–
was not directly measured in the ClO2
0.6
experiments, its distinctive pink color
0.4
was not observed in any of the ClO2
experiments, and the formation of
0.2
MnO4– from the reaction between ClO2
and Mn2+ has not previously been re0.0
0
300
600
900
1,200
1,500
1,800
ported in the literature. In the O3 expers
Time—
iments, the water turned pink when the
0.25
reaction stoichiometry of 2.2 mg O3/mg
2+
Mn was met. The formation of
MnO4– upon ozonation of water con0.20
taining Mn 2+ has previously been
reported (Gregory & Carlson, 2001;
0.15
Paillard et al, 1989).
The ClO2 results also indicated that
0.10
Mn2+ ions may be oxidized before the
NOM in Horsetooth Reservoir. This
0.05
hypothesis is supported by Figure 1 for
the cases of initial Mn 2+ = 200 and
0.00
0
300
600
900
1,200
1,500
1,800
1,000 µg/L. In both cases, ClO2 doses
Time—
s
of 100% stoichiometry (based on 2.45
mg ClO2/mg Mn2+) oxidized Mn2+ to
0.8
<5 µg/L. Significant ClO2 demand from
0.7
the NOM would have prevented the
2+
nearly 100% oxidation of Mn that
0.6
was observed.
0.5
Another hypothesis, which may
0.4
account for the low Mn 2+ residuals
observed in ClO2 experiments and the
0.3
higher residuals observed in the
0.2
KMnO4 experiments, is that ClO2 is
0.1
able to oxidize NOM-complexed Mn2+
ions and that KMnO 4, which has a
0
0
300
600
900
1,200
1,500
1,800
lower oxidation–reduction potential,
Time—s
may not. Although Mn2+ is not readAssumed oxidant dose—125% theoretical stoichiometry; KMnO4—potassium
ily complexed by NOM (unlike iron,
2+
permanganate, Mn—manganese, Mn —manganous ion, MnO2—manganese oxide
for instance), it has been shown that a
small fraction, probably <10%, will
complex with NOM (Salbu &
complexed fraction of the dissolved Mn and reduce
Steinnes, 1995; Knocke et al, 1990). If, for example,
Mn2+ concentrations to <5 µg/L.
5% of the Mn2+ in a 200-µg/L solution is complexed and
2+
unavailable for oxidation by KMnO4, the final Mn
Modeling results. To characterize the effect of initial
residual will be at least 10 µg/L. This concentration can
Mn2+ concentration on oxidation kinetics and examine the
be significant. ClO2, however, appears to oxidize the
relative concentrations of Mn2+, MnO2(s), and MnO2 ⬅
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Mn (MnO2 with sorbed Mn2+ ion) during the reaction
period, the kinetic model and equations developed by
Carlson and Knocke (1999) for oxidation by KMnO4
were used. This model was based on the autocatalytic
mathematical model originally developed by Van Benschoten et al (1991). Carlson and Knocke’s (1999) work
expanded the original rate expressions to include a term
to account for DOC concentration and were able to more
accurately characterize the KMnO4–Mn2+ reaction for
water treatment applications. The model in this study
used Carlson and Knocke’s second-order rate constants:
log k1 = 2.32, log k2 = 4.01, log k3 = 2.61, and log k4 =
–0.37 s–1 M–1. The rate constant k4 accounts for a firstorder reaction between organic matter and MnO4–. The
aim of this analysis was not to compare this study’s
observed Mn2+ removal results with model-predicted
results. Rather, the aim was to assist in the understanding
of the observed data, which indicated that initial Mn2+
centrations, however, the reaction proceeds relatively
slowly, and an excess of free MnO2(s) is not available—
contributing to the higher final Mn2+ residuals that were
observed for these conditions.
Effect of pH and temperature on Mn 2+ oxidation.
Although the effects of pH and temperature on Mn2+
oxidation were not directly investigated in this study, a
discussion of the effects of these variables will assist in
applying these results to other water quality conditions.
Mn2+ oxidation rates increase and final Mn2+ residuals decrease with increases in pH and temperature (Reckhow et al, 1991; Knocke, 1991b; Knocke et al, 1990).
Knocke (1991b) showed that for ClO2 and an initial
Mn2+ concentration of approximately 0.95 mg/L, oxidation was completed within 5 s at pH 7.0, whereas at
pH 5.5, complete oxidation required about 20 s. For
drinking water treatment, these differences may not be
significant if sufficient contact time is available. Tem-
The objectives of this study were to evaluate ClO 2, KMnO 4, and O 3 for the oxidation
of Mn 2+ to <10 µg/L and to understand the effect of initial Mn 2+ concentration and oxidant dose.
concentrations have a large effect on oxidation kinetics and
final dissolved Mn residuals and examine the relative
concentrations of Mn2+, MnO2(s), and MnO2 ⬅ Mn2+
throughout the reaction.
The model outputs (Figure 6) show the effect of initial
Mn2+ concentration on the relative concentrations of
Mn2+, MnO2(s), and MnO2 ⬅ Mn. The assumed KMnO4
dose was 125% of the theoretical stoichiometric requirement. Although the oxidation rates are somewhat slower
than those shown by the experimental data, the impact of
initial Mn2+ on oxidation kinetics is clearly illustrated
and is similar to the observed data presented in Figure 2.
Figure 6 shows that the MnO2 ⬅ Mn2+ complex is
the dominant species as the reaction proceeds. In the early
stages of the reaction, the concentration of free MnO2(s)
is low because Mn2+ quickly sorbs to newly formed
MnO2(s). This result was expected, given that the sorption
rate constant k2 (4.01 s–1 M–1) is approximately 1.7 orders
of magnitude greater than the solution-phase oxidation
constant k1 (2.32 s–1 M–1). As illustrated in Figure 6 at initial Mn2+ = 1.0 mg/L, only when the Mn2+ concentration becomes low does the MnO2(s) concentration increase
significantly. These results suggest that a mechanism that
accounts for the low final Mn2+ residuals observed in all
experiments using an initial Mn2+ = 1.0 mg/L would be
that an excess of free MnO2(s) serves to scavenge Mn2+
from solution. At high initial Mn2+ concentrations, the
reaction kinetics will be rapid because the concentration
of both reactants is high. As the reaction progresses and
the concentration of the reactants decreases, Mn2+ continues to be efficiently removed from solution because of
the sorption mechanism k2. For low initial Mn2+ con-
perature effects were similar. Complete oxidation with
ClO 2 was achieved within 60 s at 10 oC and within
approximately 10 s at 25oC.
For KMnO4, similar effects of pH and temperature
were observed. For pH 5.5 and 7.0, for example, there
were significant differences in the rate of oxidation, but
the final residuals were almost equal after 15 s. Knocke
et al (1990) observed that for ClO2 and KMnO4, second-order rate constants increased by approximately one
order of magnitude when the pH was increased from 6.0
to 7.0.
For O3, Reckhow et al (1991) also reported an increase
in oxidation rate with increasing pH. At pH 6, a secondorder rate constant of 5 × 103 M–1s–1 was observed,
whereas at pH 7 the observed rate was 2 × 104 M–1s–1.
CONCLUSIONS
The objectives of this study were to evaluate ClO2,
KMnO4, and O3 for the oxidation of Mn2+ to <10 µg/L
and to understand the effect of initial Mn2+ concentration
and oxidant dose. The following conclusions are based on
the specific raw water quality conditions of this study: pH
7.0, TOC = 3.4 mg/L, and temperature = 9oC.
ClO2 was the most effective oxidant for removing
lower initial Mn2+ concentrations (60 and 200 µg/L)
and for consistently meeting the treatment goal of <10
µg/L final Mn2+. Reaction kinetics between ClO2 and
Mn2+ were significantly faster than those observed for
the KMnO4–Mn2+ reaction. For ClO2, Mn2+ residuals
were reduced to <10 µg/L within 90 s at initial Mn2+ =
200 and 1,000 µg/L and within 180 s when the initial
Mn2+ = 60 µg/L. For KMnO4, the reaction time required
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GREGORY ET AL | PEER-REVIEWED | 95:1 • JOURNAL AWWA | JANUARY 2003
107
to produce similar results was approximately 30 min
for initial Mn2+ = 60 µg/L and 10–20 min for initial
Mn2+ = 200 µg/L.
For Mn2+ oxidation by O3, there is a relatively narrow dosage range in which minimum dissolved Mn
residuals can be achieved. Dosages below this range significantly increase Mn2+ residuals because of competing
O 3 demand from NOM. Dosages above this range
increase the potential for MnO4– formation and increase
dissolved Mn residuals. Optimum O3 dosages for the
lower initial Mn2+ levels resulted in final dissolved Mn
residuals >20 µg/L.
For the conditions of this research, Mn2+ oxidation
by KMnO4 did not provide acceptable results for the
two lowest initial Mn2+ concentrations. Oxidation of
Mn2+ to <10 µg/L by KMnO4 was possible only at initial Mn2+ = 1,000 µg/L. For an initial Mn2+ = 60 µg/L,
a reaction time of 30 min was required to achieve Mn2+
<20 µg/L. Thus, KMnO4 requires considerably longer
reaction time than O3 and ClO2 to achieve acceptable
Mn2+ oxidation.
For the highest initial Mn2+ concentration, 1,000 µg/L,
all oxidants met the treatment goal of final Mn2+ residuals <10 µg/L within approximately 180 s. The higher
concentrations of the reacting species, Mn2+, and oxidant cause an increase in the reaction rate for all three
Mn2+ removal mechanisms: solution-phase oxidation,
sorption of Mn2+ to MnO2(s), and surface oxidation.
ACKNOWLEDGMENT
The authors express gratitude to Kevin Gertig and Ben
Alexander of Fort Collins Utilities. Their enthusiastic
REFERENCES
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support allowed this research to be successful. This work
was funded by the city of Fort Collins.
ABOUT THE AUTHORS:
Dean Gregory5 is a PhD candidate in the Department
of Civil Engineering at Colorado State University, Fort
Collins, CO 80523; e-mail
[email protected]. He has a BA degree from
Colgate University in Hamilton, N.Y., and an MS
degree in environmental engineering
from Colorado State University. Gregory worked as a process engineer
for the Fort Collins Utilities investigating Mn issues and conducting
pilot-scale research on other treatment problems before beginning this
doctoral program. His work has
been published previously in Ozone Science & Engineering and Journal of Environmental Engineering.
Kenneth Carlson is an assistant professor in the
Department of Civil Engineering at Colorado State
University.
FOOTNOTES
1CDG
Technology, Bethlehem, Pa.
Technology, Bethlehem, Pa.
Instruments, model 800 TOC analyzer, Boulder, Colo.
4Amicon Inc., Beverly, Mass.
5To whom correspondence should be addressed
2CDG
3Sievers
If you have a comment about this article, please contact
us at [email protected].
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