1. Effects of Ionic Strength and
pH on Arsenic Mobilization
from Arsenopyrite
Implications for Managed Aquifer Recharge
Chloe An
July 20, 2012
A Research Paper Presented to the Students and Teachers as Research Scientists Program at the
University of Missouri-St. Louis | Sponsored by LMI Aerospace, Inc./D3 Technologies, and the Solae
Company

2. ABSTRACT
Elevated arsenic concentrations are common in water systems that utilize managed aquifer
recharge (MAR) due to the oxidation of arsenopyrite, creating serious health implications. In this study,
the effects of ionic strength and pH on arsenic mobilization from arsenopyrite were investigated. The goal
was to determine the conditions under which arsenic mobilization is minimized. Triplicate batch reactors
were run for each condition, and arsenic levels and samples were measured using inductively coupled
plasma mass spectrometry (ICP-MS). In terms of pH, the lowest arsenic mobility was achieved at pH 5,
rather than pH 3 or 7, in the sodium nitrate system. This trend indicates that arsenic levels are affected by
the mobilization of arsenic from arsenopyrite as well as the sorption of arsenic by ferrihydrite. Therefore,
a balance is necessary between the two, resulting in lower levels of arsenic at the middle pH. On the other
hand, in the sodium chloride system, because the presence of chloride inhibits ferrihydrite formation, the
only trend seen was that of greater arsenic mobilization at lower pHs due to proton promoted dissolution
of arsenopyrite, evidenced by the lowest levels of arsenic at pH 7. However, the sodium nitrate system
yielded the highest arsenic levels at 10 mM, versus 1 or 100 mM. This suggests that further
experimentation may be necessary to understand how ionic strengths effects iron oxide nucleation or
growth, which may control arsenic mobilization.
INTRODUCTION
Managed aquifer recharge (MAR) is a water management technique that injects reclaimed water,
or wastewater with solids removed, into existing groundwater aquifers for later use (Dillon et al., 2009).
This technique can help restore the groundwater balance of an aquifer and provides a cost-effective means
of transforming reclaimed water into drinking or sanitation water without establishing a complex water
sanitation system (Dillon, 2005). The implications are huge—currently, the amount of fresh water
available is rapidly dwindling due to the increasing global population, and many around the world suffer
from either the lack of water or waterborne diseases from using contaminated water. MAR could help
address these issues without posing a huge cost on both the people and the environment.
However, one pressing issue does arise when using MAR—studies have found that injected water
may trigger the oxidation of arsenopyrite present in aquifer formation minerals (Jones & Pichler, 2007;
Wallis et al., 2010; Prommer & Stuyfzand, 2005; Dillon et al., 2009). This reaction mobilizes arsenic, a
toxic metalloid, known to cause skin and bladder cancer (Walker et al., 2005). This poses dangerous
health implications for those who drink the water. These dangers are evident—studies of West Bengal
and the Middle Ganga Plain, India, where arsenic levels in the drinking water are high, have pointed to
signs of skin lesions, liver disease, and even miscarriages and stillbirths for those who drank the
contaminated water (Rahman et al., 2005; Chakraborti et al., 2003).

3. Previous studies have looked into how to limit the oxidation of arsenopyrite, studying the effect
of temperature on the rate of dissolution and concentration of arsenic present (Neil, 2010), the effect of
concentration of dissolved oxygen, pH, and temperature on the rate of dissolution (Yu et al., 2007), and
looking at the rate of oxidation under circumstantial pH (Walker et al., 2005). The purpose of this study is
to determine the impact of water chemistry on the oxidation of arsenopyrite, focusing specifically on
ionic strength and pH, and to draw connections with arsenic mobilization during MAR.
MATERIALS AND METHODS
Sodium nitrate and sodium chloride solution
To prepare the sodium nitrate solution, the necessary mass of sodium nitrate was measured using
a Mettler Toledo XS105 scale for solutions with the ionic strengths of 1 mM, 10 mM, and 100 mM. The
pH of the solution was adjusted by adding a nitric acid or sodium hydroxide solution to achieve pHs of 3,
5, and 7.
The sodium chloride solution was prepared using the same method. However, high ionic
strengths in a sodium chloride solution often cause interference with the argon plasma in the inductively
coupled plasma mass spectrometer, and thus ionic strength levels tested for the sodium chloride solution
were 1mM and 10 mM.
Batch reactors
Triplicate batch reactors were used for each experimental condition, running simultaneously.
0.0500 grams of fine-grained arsenopyrite and 250 mL of the prepared solution were added to each
reactor. A 2 mL sample was taken from each reactor and filtered into a 5mL test tube using a 0.2 μm
polypropylene membrane syringe filter. A cap was added to prevent evaporation. Once all three 0 hour
samples were taken, magnetic stirring bars were added into the reactors and the reactors were set on a
stirring plate to be stirred continuously over the course of the reaction. Each reactor was sampled using
the same method at 0, 0.5, 1, 2, 3, 4, 5, 6, and 24 hours.
Inductively coupled plasma mass spectrometry
Arsenic mobilization was quantified from analyzing the batch reactor samples. The levels of
arsenic present were determined using inductively coupled plasma mass spectrometer (ICP-MS). Because
the samples are still exposed to air after they are taken, and thus further oxidation and subsequent
precipitation reactions could have occurred, 40 μL of nitric acid was added to each sample to dissolve any
precipitates as well as insure the samples won’t clog the machine. The samples were then run through the
machine, where they were ionized using an argon gas plasma flame and passed into a quadruple mass
spectrometer, which separates and identifies ions according to their mass/charge ratio (Thomas, 2001).

4. The intensity of counts for each ion was correlated to known concentrations of arsenic to create a linear
calibration curve. This relationship was then used to calculate the unknown concentrations in the samples.
RESULTS AND DISCUSSION
Observed variations in arsenic concentration between identical reactors are hypothesized to occur
because arsenopyrite samples used were a mixture of arsenopyrite and quartz. Thus, some reactors may
have received more arsenopyrite, and consequently released more arsenic, and vice versa. Therefore, from
the triplicate trials, one replicate was disregarded when the concentration evolution appeared to be
dramatically different from the other two. All of the graphs used to analyze the data represent duplicate
trials.
pH trends for sodium nitrate reactors
Figure 1a, which can be found in the appendix, represents the arsenic concentration evolution in a
1mM sodium nitrate solution with pHs of 3, 5, and 7. Figure 1b shows the same data, focusing on the 0 to
6 hour time period. Over the course of 24 hours in the 1 mM reactors, the arsenic concentration steadily
increased at all pH levels, with the pH 3 solution reaching the highest arsenic concentration of 1.37 μM,
and the pH 5 solution reaching the lowest arsenic concentration of 1.16 μM. The graphs show signs of a
pattern more evident in a 100 mM sodium nitrate solution—pH 5 yields lowest arsenic concentration of
the three pHs, only reaching a concentration of 1.16 μM. However, many of these error bars are
overlapping, particularly between pH 3 and pH 7, making it difficult to draw conclusions about the pH
trends for the 1 mM ionic strength solution.
Figure 2 shows the arsenic concentration evolution in a 10 mM sodium nitrate solution. Figure 2a
represents the data over a 24 hour period, while Figure 2b focuses in on the 0 to 6 hour period. Over the
course of the 10 mM experiment, arsenic concentration increased steadily at all pH levels, with the levels
of arsenic in the pH 3 solution dramatically greater than the levels of arsenic in the pH 5 and 7 solutions.
The arsenic levels in the pH 3 solution reached a high of 3.03 uM, while pH 5 and 7 leveled out at around
1.6 uM. Although the error bars for the arsenic levels at pH 5 and 7 overlap and prevents conclusions
about the pH trend being drawn, the 10 mM sodium nitrate graphs do show a drastic increase in arsenic
concentration at pH 3, which correlates with some similar trends seen in previous graphs.
Figure 3a shows the arsenic concentration evolution in a 100 mM sodium nitrate solution over 24
hours, and Figure 3b focuses on the data over the 0 to 6 hour period. Over a 24 hour period, the arsenic
concentration in the 100 mM reactors increased steadily at all three pH levels, reaching the highest levels
of arsenic at pH 3 with a concentration of 2.02 μM and the lowest levels of arsenic at pH 5 with a
concentration of 0.88 μM. These graphs highlight a pattern that appeared in the 1 mM solution as well—
the arsenic concentrations at pH 3 and 7 are drastically greater than the arsenic concentrations at pH 5. In

5. addition, for the 100 mM sodium nitrate solution, the error is small enough and the difference in
concentrations is large enough that there is little overlap between pH 5 and pH’s 3 and 7.
This trend can be explained by the relationship between pH and the formation of iron oxides,
specifically ferrihydrite, in the solution. As pH decreases, the levels of arsenic mobilized increases, as
shown by the highest levels of arsenic at pH 3, especially in 10 mM sodium nitrate solution. This is due to
the abundance of protons (H+
) at lower pHs and the proton-promoted dissolution of arsenopyrite (Neil et
al., 2011):
FeAsS +3H2O + 2H+
 Fe2+
+ H3AsO3 + H2S + H2 (1)
As pH increases to 5, arsenic mobility decreases. This can be due to less proton-promoted
dissolution or due to the formation of ferrihydrite, which can sorb large fractions of arsenic (Roberts et
al., 2004). The formation of ferrihydrite is more favorable in higher pH solutions due to the abundance of
hydroxide (OH-
) ions:
Fe3+
+ 3OH-
 Fe(OH)3 (2)
However, as pH increases further to pH 7, the ability of ferrihydrite to sorb the arsenic decreases
as well. Arsenic released by the arsenopyrite is released as an arsenate or arsenite anion, meaning that it
has a negative charge, and the ability of ferrihydrite to sorb arsenic depends on the attraction between
arsenic and the ferrihydrite surface. At higher pHs, more OH- will be present on the surface of the
ferrihydrite, and the arsenic anions will be less attracted to the less positively charged ferrihydrite. On the
other hand, at lower pHs, the surface of ferrihydrite is covered with more H+
ions, allowing for a strong
attraction between the ferrihydrite and the arsenic and thus more adsorption and less arsenic present in the
solution. Therefore, we observe the trend displayed by the graphs in the appendix—the levels of arsenic
are higher at the low and high pH levels. At pH 3, we found the most arsenic mobilization due to proton-
promoted dissolution (Equation 1) and a lack of ferrihydrite formation (Equation 2). At pH 7, while there
is likely to be more ferrihydrite formation, there would also be less arsenic sorption due to the prevalence
of negatively charge hydroxide ions on the surface, and thus a high level of arsenic can occur. However,
at pH 5, there seems to be a balance between these two interactions—there is some arsenic being released,
but sorption is also occurring, resulting in the lowest concentrations of arsenic.
Ionic strength effects for sodium nitrate reactors

6. Figure 4a represents arsenic concentration evolution at pH 3 with solutions at 1mM, 10 mM, and
100 mM over a 24 hour time period; figure 4b shows the same data over a 0 to 6 hour time period. Over
the 24 hours for the pH 3 reactors, arsenic levels increased steadily at all ionic strengths, with the arsenic
levels hitting a high of 3.03 uM in the 10 mM sodium nitrate solution and a low of 1.37 uM in the 1 mM
solution. Similar to the pH trend, these graphs show one level of ionic strength yielding dramatically
different levels of arsenic—the 10 mM solution yielded much higher levels of arsenic versus the 1 mM
and 100 mM solutions. However, the error bars between the 3 levels overlap slightly, and as a result
drawing concrete conclusions is difficult.
Figure 5a shows the arsenic concentration evolution at pH 5 over a 24 hour period, while figure
5b focuses on the same data over a 6 hour period. The arsenic levels increased steadily over all ionic
strengths within the pH 5 reactors, with the highest arsenic concentration found in 10 uM sodium nitrate
solution at 1.61 uM and the lowest arsenic concentration found in the 100 uM sodium nitrate solution at
0.88 uM, following the trend of highest arsenic concentrations present in the 10 uM solution seen in the
pH 3 system. In addition, while the error bars overlap slightly, there is not as much overlap between each
level.
Figure 6 shows the arsenic concentration evolution at pH 7, with 6a showing the data over 0 to 24
hour time period and 6b focusing on a 6 hour time period. At pH 7, the arsenic levels increased steadily
over all ionic strengths. The 100 mM sodium nitrate solution yielded the highest level of arsenic at 2.30
uM, while the 1 mM sodium nitrate solution yielded the lowest level of arsenic at 1.35 uM. This pattern
contradicts the previous two trends where the 10 mM solution yielded the highest level of arsenic.
The observed patterns for arsenic mobility can be explained in part by differences in the
mechanism of iron oxide formation at different ionic strengths. At low ionic strengths, iron oxide
formation is nucleation dominated, meaning that there is a large amount of small particles forming, but
little particle growth (Jun et al., 2010). Therefore, the particles present have a higher surface area, and
thus have a higher capacity for arsenic sorption, resulting in lower arsenic concentrations in a 10 mM
solution. This may help explain why arsenic levels are higher in a 1 mM solution than in a higher ionic
strength solution. However, these iron oxide formation studies were conducted at acidic conditions; there
is a caveat that different nucleation and growth behavior may occur at circumneutral pH.
pH effects for sodium chloride reactors
Figure 7a represents arsenic levels in a pH 3, 5, and 7 one mM solution over a 24 hour period,
while figure 7b shows the same data over a 6 hour period. The arsenic levels increase steadily over the
time of the reaction, with levels reaching a high of 2.26 uM at pH 5, and a low of 0.83 uM at pH 7. The
graph shows a pattern contradictory to that found in the sodium nitrate system—rather than the lowest

7. levels of arsenic coming from pH 5, which was often tied for highest with pH 3 in the sodium chloride
system, the arsenic levels at pH 7 are dramatically lower than the other two.
Figure 8 shows arsenic concentration evolution in 10 mM solution over a 6 hour time period. The
arsenic concentration increased over the duration of the experiment, reaching a maximum 0.33 uM at pH
5 and a minimum of 0.18 uM at pH 7. Again, this graph shows the lowest arsenic concentration at pH 7.
This trend can be explained by the interaction between the chloride and ferrihydrite—the presence
of chloride causes the formation of iron chloride complexes, and thus inhibits the formation of
ferrihydrite. Therefore, arsenic attenuation through adsorption by ferrihydrite doesn’t occur as easily, and
the pH trend, wherein the proton-promoted dissolution of arsenopyrite leads to higher levels of arsenic at
lower pHs, is more dominant.
Ionic strength effects for sodium chloride reactors
Figure 9 depicts the evolution of arsenic levels at pH 3 for 1 mM and 10 mM sodium chloride
over a 6 hour time period, while figure 10 shows levels at pH 5 and figure 11 shows levels at pH 7. For all
three systems, the 1 mM sodium chloride system yielded higher levels of arsenic than the 10 mM sodium
chloride solution.
There is a need for further investigations to understand the exact mechanisms of this trend.
Interestingly, compared to NO3
-
, the Cl-
containing system can be nucleation-dominated at acidic
conditions, resulting in smaller particles. These conditions may be optimal for adsorption due to their
greater surface area.
CONCLUSIONS AND RECOMMENDATIONS
This study illustrated many interesting trends regarding the effect of pH and ionic strengths on the
mechanism of arsenic release and aqueous mobility. In regards to pH, in the sodium nitrate system, the
lowest levels of arsenic existed in the pH 5 solution for all three levels of ionic strength, indicating a
necessary balance between arsenic mobilization and sorption. At lower pH levels, the release of arsenic
from arsenopyrite is promoted by the abundance of protons present in the solution and the sorption of
arsenic by the ferrihydrite is limited by reduced amounts of ferrihydrite produced. However, at higher pH
levels, while less arsenic is being released and more ferrihydrite is present, the ability for ferrihydrite to
sorb arsenic decreases with more OH-
present. pH 5 therefore represents a balance between these two
mechanisms. On the other hand, in the sodium chloride system, pH 7 yielded the lowest arsenic
concentration. This trend can most likely be explained by the fact that the presence of chloride often
causes the formation of iron chloride complexes and thus inhibits the formation of ferrihydrite. As a
result, less sorption occurs, and the main trend apparent in the system is the pH trend, where there is
higher dissolution at lower pHs due to proton-promoted dissolution.

8. In regards to ionic strength, while there were some signs of a trend present, further study is
needed to produce concrete conclusions. The pH 3 and pH 5 sodium nitrate systems indicated higher
levels of arsenic present in a 10 mM solution. Further investigation is necessary to look into how ionic
strengths affect arsenic mobilization and if a growth dominated solution limits arsenic mobilization in
some way.
From these results, recommendations for reclaimed water pretreatment can be made to adjust pH
levels and ionic strengths. However, there are some feasibility issues with this—we cannot simply change
an aquifer’s pH level on cue, and there are further implications for how this pretreated water would affect
other minerals present in the aquifer as well as native microbial activity.
In addition to the further investigation needed to determine the effect of ionic strength on arsenic
mobilization, future trials could be run to determine how pH affects arsenic mobilization in more complex
systems—a sodium nitrate system is much simpler than a wastewater system. Specifically, we are
interested in how the addition of organic compounds and bacteria affect the mechanics of arsenic release
and mobilization. Future experiments could also be run with a mixture of soil and arsenopyrite to see how
soil presence affects arsenic mobilization.
ACKNOWLEDGEMENTS
I would like to express my sincerest gratitude to my mentors, Ms. Chelsea Neil and Dr. Young-
Shin Jun, for allowing me to work in their lab and for their continued guidance during the research.
Special thanks also go to Mr. Michael Hope for his valuable comments during the writing process. I
would further like to express my appreciation to LMI Aerospace, Inc./D3 Technologies and The Solae
Company for their continued support of the STARS program.

9. APPENDIX
Figure 1a: Arsenic concentration evolution in 1mM sodium nitrate solution, 0 to 24 hours
Figure 1b: Arsenic concentration evolution in 1 mM sodium nitrate solution, 0 to 6 hours

10. Figure 2a: Arsenic concentration evolution in 10 mM sodium nirate solution, 0 to 24 hours
Figure 2b: Arsenic concentration evolution in 10 mM sodium nirate solution, 0 to 6 hours

11. Figure 3a: Arsenic concentration evolution in 100 mM sodium nitrate solution, 0 to 24 hours
Figure 3b: Arsenic concentration evolution in 100 mM sodium nirate solution, 0 to 6 hours

12. Figure 4a: Arsenic concentration evolution at pH 3 (sodium nitrate solution), over 0 to 24 hours
Figure 4b: Arsenic concentration evolution at pH 3 (sodium nitrate solution), over 0 to 6 hours

13. Figure 5a: Arsenic concentration evolution at pH 5 (sodium nitrate solution), 0 to 24 hours
Figure 5b: Arsenic concentration evolution at pH 5 (sodium nitrate solution), 0 to 6 hours

14. Figure 6a: Arsenic concentration evolution at pH 7 (sodium nitrate solution), 0 to 24 hours
Figure 6b: Arsenic concentration evolution at pH 7 (sodium nitrate solution), 0 to 6 hours

15. Figure 7a: Arsenic concentration evolution in 1 mM sodium chloride solution, 0 to 24 hours
Figure 7b: Arsenic concentration evolution in 1 mM sodium chloride solution, 0 to 6 hours

16. Figure 8: Arsenic concentration evolution in 10 mM sodium chloride solution, 0 to 6 hours
Figure 9: Arsenic concentration evolution at pH 3 (sodium chloride solution), 0 to 6 hours

17. Figure 10: Arsenic concentration evolution at pH 5 (sodium chloride solution), 0 to 6 hours
Figure 11: Arsenic concentration evolution at pH 7 (sodium chloride solution), 0 to 6 hours

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